College Level Chemistry by AudioLearn

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TABLE OF CONTENTS Preface........................................................................................................ 1 Chapter One: Atoms and Molar Mass .......................................................... 5 Atoms and Subatomic Particles ....................................................................................... 5 Atomic Number................................................................................................................ 7 Atomic Mass Number ...................................................................................................... 9 Isotopes .......................................................................................................................... 10 Isobars ............................................................................................................................. 11 Valency ............................................................................................................................ 11 Atomic Structure Theories ............................................................................................. 13 Periodic Table and Periodic Law ................................................................................... 13 Atomic Radius .................................................................................................................17 Ionization Energy........................................................................................................... 18 Concept of Mole ............................................................................................................. 19 Stoichiometry ................................................................................................................. 19 Equivalent Weight and Mole Ratio ............................................................................... 22 Key Takeaways ............................................................................................................... 23 Quiz ................................................................................................................................ 24 Chapter Two: Atomic Structure ................................................................. 28 Electron Configuration .................................................................................................. 28 Pauli Principle ................................................................................................................ 34 Aufbau Principle ............................................................................................................ 34 Electron Affinity ............................................................................................................. 36 Electronegativity ............................................................................................................ 38

Dipole Moments ............................................................................................................. 39 Magnetic Properties in Atoms ....................................................................................... 41 Key Takeaways ............................................................................................................... 42 Quiz ................................................................................................................................ 43 Chapter Three: Properties of Gases ........................................................... 47 Gases and their Properties ............................................................................................. 47 A Word on Temperature ................................................................................................ 51 Gas Volume .................................................................................................................... 51 Pressures and Gases....................................................................................................... 52 Gas laws.......................................................................................................................... 53 Boyle’s Gas Law.............................................................................................................. 54 Charles’ Law ................................................................................................................... 55 Avogadro’s Law .............................................................................................................. 55 Partial Pressures in Gases .............................................................................................. 57 Kinetic Theory ................................................................................................................ 58 Graham’s Law of Effusion.............................................................................................. 59 Key Takeaways ............................................................................................................... 60 Quiz ................................................................................................................................ 61 Chapter Four: Thermochemistry............................................................... 65 Energy in Chemistry and Physics .................................................................................. 65 Properties of Heat in Chemistry .................................................................................... 66 Laws of Thermodynamics .............................................................................................. 67 Heat Capacity ................................................................................................................. 69 Calorimetry .....................................................................................................................71

Enthalpy and Energy ..................................................................................................... 73 Rules of Thermochemistry............................................................................................. 75 Key Takeaways ............................................................................................................... 76 Quiz ................................................................................................................................ 77 Chapter Five: Chemical Bonding and Molecular Structure ........................ 81 Compounds and Bonding .............................................................................................. 81 Ions ................................................................................................................................. 84 Covalence .......................................................................................................................88 Molecular Orbital Theory .............................................................................................. 90 Shapes of Molecules ....................................................................................................... 91 Hydrogen Bonding ......................................................................................................... 94 Bonding in Metals .......................................................................................................... 95 Key Takeaways ............................................................................................................... 96 Quiz ................................................................................................................................ 97 Chapter Six: Chemical Formulas and Reactions ...................................... 100 How Reactions Occur................................................................................................... 100 Types of Chemical Reactions ....................................................................................... 102 Writing Reactions ........................................................................................................ 103 Limiting Reagents ........................................................................................................ 106 Rates of Reactions ........................................................................................................ 107 Energy of Activation..................................................................................................... 108 Chemical Equilibrium .................................................................................................. 109 Key Takeaways .............................................................................................................. 112 Quiz ............................................................................................................................... 113

Chapter Seven: Solids and Liquids ........................................................... 117 Phases of Matter ........................................................................................................... 117 Water Condensation, Boiling, and Evaporation .......................................................... 121 Crystals ......................................................................................................................... 125 Liquids.......................................................................................................................... 129 Liquid Forces ................................................................................................................ 131 Key Takeaways ............................................................................................................. 135 Quiz .............................................................................................................................. 136 Chapter Eight: Solutions and Colloids ..................................................... 140 The Ideal Solution ........................................................................................................ 140 Nonideal Solutions ....................................................................................................... 145 Electrolyte Solutions .................................................................................................... 145 Mixing Solutions .......................................................................................................... 148 Expression of Concentrations ...................................................................................... 148 Colligative Properties ................................................................................................... 150 Anomalous Colligative Properties ................................................................................ 151 Solution Stoichiometry ................................................................................................ 152 Reactions in Solutions ................................................................................................. 152 Colloids......................................................................................................................... 154 Key Takeaways ............................................................................................................. 156 Quiz ...............................................................................................................................157 Chapter Nine: Oxidation-Reduction Reactions ......................................... 161 Oxidation and Reduction Reactions ............................................................................. 161 Redox Reactions in Common Situations ..................................................................... 166

Key Takeaways ............................................................................................................. 167 Quiz .............................................................................................................................. 168 Chapter Ten: Acids and Bases .................................................................. 172 Theories of Acids and Bases .........................................................................................172 pH Scale ........................................................................................................................ 177 Neutralization Reactions ............................................................................................. 179 Titration ....................................................................................................................... 179 Buffers ........................................................................................................................... 181 Hydrolysis .................................................................................................................... 182 Key Takeaways ............................................................................................................. 184 Quiz .............................................................................................................................. 185 Chapter Eleven: Electrochemistry ........................................................... 189 Electrochemistry Basics ............................................................................................... 189 Properties of Electric Cells ........................................................................................... 189 Faraday’s Law .............................................................................................................. 196 Nernst Equation ........................................................................................................... 197 Electromotive Force ..................................................................................................... 199 Key Takeaways ............................................................................................................ 200 Quiz .............................................................................................................................. 201 Chapter Twelve: Organic Chemistry and Biochemistry............................ 205 Nomenclature .............................................................................................................. 205 Alkanes .........................................................................................................................206 Cycloalkanes................................................................................................................ 208 Alkenes and Alkynes ....................................................................................................209

Benzene Derivatives ..................................................................................................... 210 Isomerism ..................................................................................................................... 211 Structural Isomerism ................................................................................................... 212 Functional Groups ....................................................................................................... 215 Basics of Biochemistry ..................................................................................................217 Proteins ........................................................................................................................ 218 Carbohydrates ..............................................................................................................220 Nucleic Acids ................................................................................................................ 224 Lipids ............................................................................................................................ 227 Key Takeaways ............................................................................................................. 229 Quiz .............................................................................................................................. 230 Summary ................................................................................................ 234 Course Questions and Answers ............................................................... 238

PREFACE This course on college chemistry provides a way for the college student pursuing an interest in chemistry or possibly already struggling with chemistry to understand this discipline in a much clearer way. Chemistry is a broad subject, ranging from “physical chemistry” or the study of atoms, molecules, chemical reactions, and the properties of solids, liquids, and gases, to things like “organic chemistry” and “biochemistry” that deal with the topic of the chemistry specific to carbon-based molecules, which are the molecules of living things. In some ways, the majority of chemistry is about the physical properties of atoms and molecules in singular form as well as the physical properties of molecules in different states, including the molecular properties of gases, solids, and liquids. Most atoms and molecules are not inert and interact with one another in specific ways. This will be a major focus in this course as we learn about how different substances exist in natural and experimental situations. Chapter one in the course introduces chemistry in its most basic form by covering atoms and their nature. Everything in the universe is made from different types of atoms. As it turns out, atoms have specific sub-particles that make up the basic structure of the atom. There are also numerous different atoms in nature, as we’ll discuss when talking about the periodic table, which outlines the relationships between different types of atoms. Finally, there is a discussion on the concept of the “mole” as it exists in understanding quantities of atoms and molecules. The focus of chapter two is the structure of atoms. As it turns out, there are specific rules that define how subatomic structures are organized. This chapter will cover the arrangement of electrons, and both the Pauli principle and Aufbau Principle, which define electron relationships inside the atom. Electrons have a natural affinity for the proton they surround but also allow for affinity of different types of atoms in the formation of larger molecules. This will be discussed in this chapter.

The main topics of chapter three in the course are the properties of gases. There are specific laws that describe the behavior of gases in chemistry, which are covered in this chapter. The study of gases involves the relationship between volume and pressure in a gas and, in the situation of mixed gases, a relationship between the different gases in a given volume. The behavior of gases in relation to the ideal gas law is covered in this chapter as well as the study of the kinetics of gas molecules in a gaseous solution. Chapter four discusses chemistry as it relates to heat, also called “thermochemistry”. There is an intricate connection between molecules and the temperature of the atoms in the molecules. Concepts of heat capacity and calorimetry are discussed as they relate to a specific molecule. The enthalpy or the total heat in a given system is discussed as it relates to the energy of the system. As in all of physical chemistry, there are specific rules related to the properties of substances at different temperatures, which are covered in this chapter. Chapter five in the course gets into the subject of chemical bonding and molecular structure. Unless an atom is completely inert, it will have the natural tendency to interact with other atoms. This is certainly the case with ionic substances in chemistry. The chapter will study the important topic of covalence and covalent bonding between molecules and how this should look from an atomic perspective. Because of the different properties of atoms, they exist within molecules in specific three-dimensional arrangements, which is covered in this chapter. The topic of bonding in metallic atoms is also covered as part of this chapter. Chapter six focuses on the inevitability of chemical reactions and on what chemical formulas look like. Molecules and atomic substances tend not to stay separate from one another when put together in the same physical space and reactions will follow a specific pattern. The rates of different reactions and why they occur or do not occur are discussed as well as the issue of the energy of activation necessary to take a group of molecules in one form and turn them into molecules of a different form. Exactly what happens in chemical reactions is covered as part of this chapter. The focus of chapter seven in the course is the behavior of solids and liquids. Molecules can exist in crystalline form, depending on the physical circumstances. There are certain

molecular forces in play when substances are in crystalline form, which are discussed in this chapter. In the same way, liquids have certain properties unique to that state that depend on the atomic nature of the molecule and on the temperature of the liquid. The different forces in play when substances are in liquid form are covered in the chapter. The study of liquids continues in chapter eight with a discussion of solutions and colloids. Solutions involve the dissolution of a substance in a liquid medium. Exactly how this is successfully done is covered as part of this chapter as well as what it means to indicate the “concentration” of a solute in solution. A solution is a unique medium for reactions to take place, as it exhibits certain properties that keep things in a given solution and that help drive reactions within the solution. Colloids have certain properties that will also be discussed. Two significant types of reactions are covered in chapter nine in the course—that being oxidation and reduction reactions, which are opposing but interrelated chemical reactions. Because these reactions happen in a balanced way, they are often discussed together. Taken together, an oxidation-reduction reaction is any chemical reaction in which the oxidation number of a molecule, atom, or ion is altered by gaining or losing an electron. The specific properties of acids and bases are discussed in chapter ten in the course. Methods involved in the titration of acids and bases are discussed as part of this chapter as are properties of strong and weak acids and bases. Buffers are weak acids or weak bases that prevent significant changes in the pH of a solution; how these substances work in acid-base chemistry is covered. Finally, the action of hydrolysis is an important part of the discussion in this section. The focus of chapter eleven in the course is electrochemistry, or the study of electricity as it applies to chemical reactions. Faraday’s law, which states that the amount of substance produced at an electrode is directly proportional to the quantity of charge flowing through an electrochemical cell, is discussed. The basics of electricity in chemical reactions are important to understand because electricity plays a role in many chemical interactions. The different properties of electrical cells and what makes them up is a part of this chapter’s discussion.

Organic chemistry and biochemistry are courses of their own; however, they are covered together in chapter twelve in the course because they are inter-related and important aspects of the study of chemistry. Organic chemistry is carbon-based chemistry and involves a variety of different types of molecules and reactions typically seen in living things. The nature of organic molecules is discussed as well as the functional groups that shape organic molecules and their behavior. Finally, biochemistry and its principles are covered as these types of molecules, too, follow the basic chemical principles already laid out in this course.

CHAPTER ONE: ATOMS AND MOLAR MASS This chapter introduces chemistry in its most basic form by covering atoms and their nature. Everything in the universe is made from different types of atoms. As it turns out, atoms have specific sub-particles that make up the basic structure of the atom. There are also numerous different atoms in nature, as we’ll discuss when talking about the periodic table, which outlines the relationships between different types of atoms. Finally, there is a discussion on the concept of the “mole” as it exists in understanding quantities of atoms and molecules.

ATOMS AND SUBATOMIC PARTICLES Most people think of the atom as the basic form of matter. The truth is that atoms have a specific structure and are themselves made up of subatomic particles. Knowing the structure of the atom is highly important as it aids in the understanding of the world as we know it and in the understanding of how substances behave in natural and experimental circumstances. The atom is made from one or more protons, neutrons, and electrons. Protons are considered to be positively charged subatomic structures, electrons are negatively charged subatomic particles, and neutrons (as you might surmise) are neutral. These will vary greatly in their size as well. Electrons are the smallest subatomic particle in basic chemistry (but not in quantum chemistry, which is beyond the scope of this course). The weight of an electron is about 9.11 x 10-31 kilograms and, as mentioned, it is negatively charged. It is the negative charge that keeps the electron attracted to the positively charged proton(s) as the natural course of electricity is that positive attracts negative and vice versa. What you should know is that, in the case of atoms that have an extra electron, the net charge on the atom is negative and it’s called a negatively charged ion. In the same way, if an atom has lost an electron, the net charge on the atom is positive, leading to a positively charged ion. We will discuss ions in detail in this course so you need to 5

understand this concept. Any ion, whether it be positively charged or negatively charged, will be automatically attracted to an ion of the opposite charge in order to have a neutrally charged substance. Protons will have a positive charge, with an approximate weight of 1.67 x 10-27 kg, which is 1836 times the weight of an electron. What’s important to know is that the number of protons in an atom is referred to as the atomic number. This will be important to remember as you study the periodic table as the atoms are organized by atomic number. Neutrons have a neutral charge and a weight of 1.69 x 10-27 kilograms, which makes them the heaviest of the subatomic particles in the atom. They have a mass that is 1839 times the mass of an electron. These particles are the least well understood of the subatomic particles. All atoms will have a neutron except for hydrogen, which has just one proton and one electron. Electrons are not known to have any specific internal structure, making them elementary particles. Both protons and neutrons have substance, being made up of other types of elementary particles known as quarks. There are two known types of quarks in atoms, which are charged particles. There are two “up quarks” each with a charge of +2/3 and one “down quark” with a fractional charge of -1/3. Adding this up, you get a total charge of +3/3 or +1. Neutrons have one up quark and two down quarks, with a neutral charge. Because quarks are charged, they have strong forces that bind them. The force that is involved in their connection involves gluons, which are elementary particles that create what is called gluon forces. There are gauge bosons, which are elementary particles that interact with physical forces. Gluons are a type of gauge boson. It should be understood that, in atoms, forces are mediated by elementary particles. Aside from gluon forces, there is a nuclear force, which binds the protons and neutrons together within the nucleus. All of the protons and neutrons together are referred to as the nucleus of the atom. These are also referred to as the nucleons (collectively). Nucleons are bound together by a short-ranged attractive force called the “residual strong force”. Atoms themselves are

much larger than the nucleus because electrons expand the total size of the atom. Figure one depicts what an atom might look like structurally:

Figure 1.

So, basically, every atom has a nucleus in the center that has at least one proton in it and at least one electron, which is “orbiting” the nucleus at a fixed energy level with a certain quantum spin value. We will talk more about this quantum spin value and about “orbitals” of electrons in the next chapter. Electrons contribute much to the charge of the atom but are elementary particles of their own. As you may have surmised, an elementary particle is one that cannot be further broken down. Quarks are elementary particles, gluons are elementary particles, and the different gauge bosons (beyond the scope of this chapter) are all elementary particles, while neutrons and protons are not elementary particles.

ATOMIC NUMBER As mentioned, the atomic number is defined as the number of protons in an atom. In chemistry, it is represented by the letter Z. An element is a substance that has a specific atomic number. For example, hydrogen, the smallest element, has just one proton, having a Z value of 1. The number of protons or the atomic number defines the element. Atoms, by definition, have a neutral charge so that the number of protons is equal to the number of electrons. Carbon has six protons and six electrons, so its atomic number is

six. Similarly, oxygen has an atomic number of 8 because it has 8 protons and 8 electrons. An atomic number is always going to be a whole number as it reflects an actual number of protons and ranges from 1 to 118 (although only the first 98 actually exist as relatively stable elements in nature). The atomic number helps identify the element and will form the basis of arranging the elements in the periodic table, which is arranged according to increasing order of atomic number. While the atomic number determines some aspect of the properties of the element, the electrons help determine the actual bonding behavior of the element as you will soon learn. You can learn the atomic number of any element by looking at its symbol on the periodic table of the elements. For example, next to the letter describing the element, such as Magnesium or Mg, there will be a number that defines the atomic number. Figure 2 shows the elemental symbol of Magnesium along with its atomic number 12:

Figure 2.

Another way of determining the atomic number is by looking at the isotope symbol. An isotope is defined as any of two or more forms of the same element that contain equal numbers of protons but different numbers of neutrons in their nuclei, and therefore differ in relative atomic mass but not in chemical properties. Isotope notation consists of both the atomic mass (to be discussed next) and atomic number written to the left of the chemical symbol. The atomic mass is written as a superscript, and the atomic number as a subscript. For example, carbon -14 is an isotope of carbon and is written as the carbon symbol “C,” with the atomic mass, 14, written as a superscript, and its atomic

number, 6, written as a subscript. ( 146C )

. The atomic number will always

be the lower of the two numbers.

ATOMIC MASS NUMBER The atomic mass number or just “mass number” of an atom or element is the sum of the protons and neutrons. This is nearly equal to the actual atomic mass of the atom. You can determine the number of neutrons in the substance by subtracting the atomic number from the mass number. The mass number is important in the understanding of isotopes. The isotope mass is referred to by the letter “A”. An isotope of an element is one that has the same atomic number but has a different mass number than another isotope. In general, there is a typical mass number for an atom or element that is relatively stable in nature. For example, with an atomic number of 6, carbon normally has a mass number of 12, giving it six neutrons and six protons. There also exists Carbon-14, which is “heavier” carbon, having two extra neutrons in it than what’s generally found. Different isotopes of the same element have the same atomic number but a different mass number. So, the mass number, represented by the letter A, is the sum of the protons and neutrons—together referred to as the nucleon. The neutron number will change, while the proton number of the element will stay the same. The actual atomic mass of the atom will be close to the mass number but will not be exactly the same because of the nearly negligible weight of the electrons. The mass number A will equal Z (the atomic number and number of protons) plus N (the number of neutrons). Radioactivity is a possible property of atoms. A radioactive isotope is one that has an unstable nucleus. It tends instead to release subatomic particles in order to become more stable, releasing energy called radioactivity or radiation in the process. Atoms or elements can exist in nature in both radioactive and non-radioactive form, although the non-radioactive isotope is found to a greater degree in nature because it is inherently more stable.

ISOTOPES Isotopes do not have to be completely unstable. An example is the hydrogen atom. There are three different forms of this element in nature: protium, deuterium, and tritium. The proton number is, of course, the same, while the neutron number changes. There are no neutrons in protium, one neutron in deuterium, and two neutrons in tritium. Figure 3 represents the structure of these three isotopes:

Figure 3.

Carbon also has three isotopes: Carbon-12, Carbon-13, and Carbon-14, representing different atomic masses. The stable form is carbon-12, while carbon-14 is considered radioactive. Other common isotopes seen in chemistry include the following: •

Oxygen—there are three isotopic forms, being O-16, O-17, and O-18

•

Uranium—there are two isotopic forms, which are U-235 and U-238

•

Chlorine—there are two isotopic forms, which are Cl-35 and Cl-37

•

Fluorine—there are three isotopic forms, which are F-17, F-18, and F-19

There are 275 isotopes made from 81 stable elements with more than 800 natural and artificially-created radioactive isotopes. This means that, while there are just 118 elements, many will have different isotopes found. Most isotopes will behave the same chemically, except for hydrogen isotopes because they vary greatly in total weight

relative to one another. There are, however, different physical properties of the different elements; they can be separated using diffusion techniques and fractional distillation. There are different uses for isotopes in science. Carbon dating, for example, makes use of the decay properties of carbon-14 to tell the age of specific fossils and uranium-235 is used as a type of fuel in nuclear reactors. There are radioactive isotopes used in medicine, such as Arsenic-74, used in tumor detection, Cobalt-60, used in certain cancer treatments, and iodine-131, used in the management of thyroid diseases.

ISOBARS Atoms belonging to different element types but that have the same atomic mass or mass number are referred to as isobars. They have a different atomic number because they belong to different elements but the atoms weigh the same. This is because they have neutrons of different quantities to add up to the same atomic mass. The size of the nucleon will be roughly the same between isobars. Because they are different elements, they will not behave the same chemically. They will, however, have similar physical properties.

VALENCY Valency in an atom is a measure of the ability of an atom or molecule to combine with other atoms or molecules. It is basically a measurement of the element’s ability to react with others. As you’ll see later, electrons are arranged in an atom in different shells or orbitals. Those electrons that are present in the outermost orbit of an atom are called the valence electrons. These are the electrons that participate in chemical reactions as they are the more energetic electrons. We will also look at different schemas of atomic structures; however, the Bohr-bury scheme best explains the concept of valency. According to the scheme, the outermost orbit contains a maximum of 8 electrons. If this orbit is completely filled, the electron has no available energy or affinity for other atoms. As you’ll see, certain elements are considered inert (such as the noble gases) because their outer orbitals are full and there is no room for reactivity with other elements. 11

When the outermost shell has eight electrons in it, an atom is said to have a “complete octet”. If it is incomplete, there is room for other electrons to be “shared” with other atoms, resulting in chemical reactivity. Hydrogen is tricky because it has just one electron in its outermost shell. It needs to lose one to be stable, giving it a valency of one. This is because the first orbital in hydrogen is considered “full” at two electrons. Magnesium has two electrons in its outermost shell. This gives it a valency of 2. Similarly, fluorine has a valency of 1 but not because it has one electron in its outer shell. It has seven, which are too many to lose. Instead, it has to gain one, which would be easier. It has the same valency as an element that has an electron to donate. It goes both ways so the goal is to get the essential “octet” in the outer shell. There is positive valency and negative valency, as you’ll soon understand. Nitrogen is an example of a substance that has a valency of 3. This is because it has seven electrons (an atomic number of seven), in which two electrons fill the first orbital (which only needs two) and five electrons fill the second orbital (which needs eight). This leads to a need for 3 electrons for a stable octet. In ammonia or NH3, this stability is achieved because each hydrogen atom shares its only electron, filling the nitrogen outer orbital. The periodic table will show which atoms or elements are stable and which have certain valencies. You can also determine the valency by using the atomic number, knowing that the number of electrons is the same as the number of protons. You also need to know that the first orbital is different and takes only 2 electrons, while the rest will take eight. In looking at the inert gas neon with an atomic number of ten, there are two electrons for the first orbital and eight for the outermost orbital. This completes an octet because two plus eight equals ten. If there are one to three electrons in the outer orbital, the tendency is to lose or donate those electrons, giving it a positive valency. If there are four to seven electrons in the orbital, the tendency is to want to accept or gain electrons. It is easier to accept an electron than it is to donate one so the valency is usually determined by subtracting the number of electrons by eight in such cases.

All the metal elements on the periodic table will have a valency of plus one, while all the elements in column 17 of the periodic table (such as fluorine and chlorine) will have a valency of negative one. Noble gases in column 18 will have a valency of zero and are considered inert. Some elements have more than one active orbital so their valencies are difficult to determine by their atomic number. These include gold, copper, and iron. There is a difference between the valency or the combining capacity of an atom and the oxidation number. The valency is directly related to the number of valence electrons in the outer orbital and how many electrons necessary to gain or lose. The oxidation number is the actual charge the atom can carry. For nitrogen, the valency is three but the oxidation number can vary from -3 to +5.

ATOMIC STRUCTURE THEORIES There have been many theories about atomic structure with most disproven over time. This has left scientists with the Bohr theory. In this theory, there is a nucleus orbited by electrons but they do not simply orbit around. They exist only in discrete distances from the nucleus in orbitals that have different energies. There are as many orbitals as are necessary to fill them. The lowest orbital contains two electrons, while the outer orbitals contain eight electrons. The Bohr theory has many flaws but it does help to understand the behavior of atoms when combining with other atoms.

PERIODIC TABLE AND PERIODIC LAW The periodic law basically states that the chemical and physical properties of the elements recur in a predictable way as long as the elements are arranged in order of increasing atomic number. As long as the elements are ordered in a specific way, their behavior can be predicted. This Periodic Law is what has led to the development of the periodic table used in modern times. This law was independently elaborated by Dmitri Mendeleev and Lothar Meyer in 1869. It allowed for the arrangement of elements in the order they are currently in even though the understanding of why the arrangement existed was not then understood. Figure 4 explains what the periodic table looks like: 13

Figure 4.

Periodic Law allows for certain properties to be explained. The first is the atomic radius and ion radius. These are different from one another but follow the same trend. It is the measure of the size of a different atom or ion. The atoms are arranged in increasing order of atomic radius, increasing from left to right and from top to bottom on the periodic table. The ionization energy of the element is a measure of how easy it is to remove an electron from an atom. This increases from left to right across the table and decreases when moving down a column on the table. The electron affinity is how easily an electron accepts an electron. The alkaline earth elements have a low electron affinity. The halogens, on the other hand, have a high electron affinity. Noble gases have a nearly zero electron affinity because they have a zero valency and full outer orbitals. Electronegativity is most closely connected to the electron affinity. Electronegativity is how easily an element attracts electrons in order to form a chemical bond. These will both decrease moving down a group and will increase moving across a period. Electropositivity is essentially the opposite trend compared to electronegativity. Figure 5

shows an illustration of what the electronegativity trend is for the elements on the periodic table.

Figure 5.

In looking at the periodic table, you can see that the elements are arranged in order of atomic number. The rows are called periods and the columns are called groups. Elements of the same group will have similar characteristics. There are eighteen groups in a specific arrangement. Hydrogen and helium are in the top row but are on opposite sides of the table. This is because they are the smallest atoms and fill the closest orbital, which only contains two electrons. There would be only eight groups if it weren’t for the 10 groups of transition metals in the middle of the table. This matches well with the eight electrons filling up the eight spots in the outer orbital. The first group of metals (the alkali metals) has a valency of one because one electron fills the orbital that should have eight in it. On the far side of the table are the halogens, which lie just to the left of the noble gases on the far right-hand side of the table. The halogens have seven electrons in the orbital and are the most likely to want to accept an electron. While it sounds hard, you should memorize as many elements on the periodic table as possible, as well as their positions. You should memorize the one to two-letter initials or the chemical symbol that define each element on the table. 15

Most elements are considered “metals”. These are found on the left-hand side of the periodic table. The far right-hand side contains the nonmetals (except for hydrogen, which is not a metal). There are some elements that have the properties of both metals and nonmetals; these are called metalloids or semi-metals. These run in a zigzag line from the upper part of group 13 to the bottom part of group 16. All metals except mercury are generally solid, while nonmetals can be solid, liquid, or gaseous at room temperature. The different metals are the alkali metals, the alkaline earth metals, the transition metals, the basic metals, the lanthanides, and the actinides. On the far right-hand side are the nonmetals, the halogens, and the noble gases. Let’s look at what these elements look like: •

Alkali metals—this is the first group (column) on the periodic table and includes Lithium, Sodium, Potassium, Rubidium, Cesium, and Francium. These tend to take a +1 charge when ionized.

•

Alkaline earth metals—this is the second group on the periodic table and includes Beryllium, Magnesium, Calcium, Strontium, Barium, and Radium. These tend to take a +2 charge when ionized.

•

Transition metals—these are thirty-eight different metals in groups three through twelve of the periodic table that often have different ionization charges. For example, there can be Fe (iron) with a 2+ charge and Fe (iron) with a 3+ charge.

•

Basic metals—these are eleven metals that take up portions of groups 13 through 16 of the periodic table and include elements like aluminum, gallium, and indium.

•

Lanthanides—these include fifteen uncommonly seen elements that are seen along the bottom of the chart with atomic masses of between 57 and 71.

•

Actinides—these include fifteen uncommonly seen elements that are listed along the bottom of the chart with atomic masses of between 89 and 103. Not all of these are seen in nature nor are they stable.

•

Metalloids—these elements zigzag along a line that includes parts of groups 13 through 17. They have properties of metals and nonmetals and include things like Boron, Silicon, and Arsenic.

•

Nonmetals—these include hydrogen and several elements that are partially on groups 14 through 16. They include carbon, nitrogen, oxygen, phosphorus, sulfur, and selenium.

•

Halogens—these include all of the elements in group 17 of the periodic table that includes fluorine, chlorine, bromine, and iodine, among others. They usually take a -1 charge when ionized.

•

Noble gases—these include group 18 and are all gaseous elements, including helium, neon, argon, krypton, and xenon.

ATOMIC RADIUS As mentioned, the atomic radius is increased as one goes down the periodic table from left to right and from top to bottom. The actual atomic radius cannot be determined because there is no way to find an electron in space so there is an indistinct boundary to an atom. The best way to do this instead is to measure the internuclear distance (between two nuclei). A covalent radius is one-half the distance between the nuclei of two identical atoms. On the other hand, the ionic radius is one-half the distance between the nuclei of two ions in an ionic bond. The cation (positively charged ion) will be smaller, while the anion (negatively charged ion) will be larger. The metallic radius is one-half the distance between nuclei of two adjacent atoms in a crystalline metallic substance. You should know that the atomic radii are determined in Standard International units or SI units of the nanometer and picometer or nm and pm, respectively. One nanometer is 1 x 10-9 meters, while the picometer is 1 x 10-12 meters. One needs to understand the concept of penetration and screening. Penetration is the distance the electron is from its nucleus. Screening is the phenomenon in which the inner electrons block the outer electrons from the positive charge of the nucleus. This

allows the outer electrons to be more easily borrowed to other atoms. The effective nuclear charge is the atomic number with the number of shielding electrons separated from it. In other words: Z(effective)= Z (atomic number) minus S (number of shielding electrons) An example of this would be chlorine: It has an atomic number of 17 and 10 shielding electrons (those protected in stable orbitals). The effective Z value is +7 because there are seven electrons left over in an outer orbital that is not complete with a filled octet. The value of +7 is the effective nuclear charge. This means that the nucleus is pulling the outer electrons closer to the nucleus, leading to a smaller atomic radius. Sodium has a much lesser nuclear charge at +1 because its atomic number is 11 (having 10 shielding electrons). Its pull on the electrons is less. Moving from left to light across the periodic table, the nucleus has a greater pull on the outer electrons and the atomic radius goes down. Moving down the periodic table means larger atomic numbers and a decreased pull on the outer valence electrons and a larger atomic radius.

IONIZATION ENERGY This is an important feature of the periodic table. In order to remove an electron from an atom, you need enough energy to remove an electron from the positive charge of the nucleus. The IE or I or Ionization energy is the energy necessary to get rid of an electron from the atom or ion. This is always a positive number. The “first IE” is the energy needed to remove the first valence electron, while the “second IE” is the energy required to remove the second valence electron. There is a third and fourth IE as well. An example is sodium. The first IE is that to make a sodium ion Na+ (Sodium Plus) and an electron, while the second IE is the energy to take Na+ to make Na2+ and the second electron. Going up a group increases the ionization energy because the atom is smaller and has fewer shielding electrons and a stronger ionization energy.

CONCEPT OF MOLE We will introduce in this section the concept of matter and how it is measured. Matter is everything—including solids, liquids, and gases. Matter consists of all the atoms and molecules that exist in nature. The three major states of matter include gases, liquids, and solids. Solids have atoms very close to one another with very strong intermolecular forces keeping the atoms held firmly in one position with only vibratory motion within the particle. There is a definite shape to metals, such as iron, aluminum, and other metallic substances. Liquids have weaker intermolecular forces with particles that are capable of a minimal amount of fluid. There is a definite volume but not a definite shape. They take the shape of the container in which they are held, including water and mercury. Gases have very weak intermolecular forces between the different molecules that are free to move, having no specific shape or volume. There is a large space between molecules when compared to solids and liquids. Examples include hydrogen, methane gas, and oxygen gas. Of course, substances can be mixtures of atoms, being classified as pure elements and compounds (which are more heterogeneous than pure elements).

STOICHIOMETRY Stoichiometry involves the calculation of various quantities or products of a chemical reaction. This long word comes from the Greek words “stoichion”, which means element, and “metry”, which means to measure. Stoichiometry involves the relative numbers of atoms and molecules which are involved in a chemical reaction.

For a reaction to be balanced, both sides of a chemical equation should have an equal number of elements. Three concepts you need to know about in balancing equations are these: •

Formula mass. This is the total sum of the atomic masses of the different atoms in a molecule. For the substance Na2S, there are two sodium atoms and one sulfur atom. Adding the two atomic masses, you get 2 times 23 and one times 32, which together is seventy-eight. The formula mass of this is then 78.

•

Mole. A mole is the unit of measurement for amount of substance in the International System of Units (SI). This is a relatively new definition. A mole is defined as the amount of a chemical substance that contains exactly 6.02214076 × 1023 (Avogadro's constant) constitutive particles, such as atoms, molecules, ions or electrons.

•

Avogadro number. This is defined as the number of atoms seen in exactly 12 grams of carbon-12, which is approximately 6.022 x 1023.

•

Molar Mass. This is the sum of the total mass of all of the atoms that make up a molecule per mole.

Take a look at the reaction between sodium and hydrochloric acid. You cannot make the reaction be one sodium atom plus one molecule of hydrochloric acid or HCl. This would lead to sodium chloride but would have a hydrogen ion left over that cannot form anything. You need to have hydrogen gas at the end which is H2. This means that the reaction must go like this: 2Na plus 2HCl leads to 2 NaCl (sodium chloride) plus one H2 molecule. This completely balances the equation. When we talk about this equation, we say that two moles of sodium react with two moles of hydrochloric acid, leading to two moles of sodium chloride and one mole of H2 gas. Figure 6 shows what this looks like: 2Na + 2HCl → 2NaCl + H2 Figure 6.

So, a mole is basically a unit of measurement used because things like grams or kilograms just don’t fit into doing a chemical reaction. Using the actual number of molecules or ions would be clumsy and wouldn’t be an easy way to describe a chemical equation. Grams and nanograms also wouldn’t work because it isn’t convenient. You should be able to convert grams to moles and vice versa. For example, how many grams is in a certain amount of a molecule. You need to learn the atomic masses of the different elements. For example, how many grams of carbon dioxide is in 0.2 moles of CO2? To do this, you need to know that the mass of carbon is 12.01 grams per mole and the mass of oxygen is 16.0 grams per mole. This leads to CO2 being 44.01 grams per mole. Taking this number, you get 0.2 multiplied by 44.02 grams/mole to get 8.80 grams. Figure 7 shows sodium along with its atomic number and atomic mass listed above the chemical symbol:

Figure 7.

So, in calculating the grams of a sample, you need to use this equation: the grams of a sample equal the molar mass multiplied by the number of moles. This means that number of moles of a substance is the grams divided by the molar mass or, in the case of an atom, its atomic mass.

EQUIVALENT WEIGHT AND MOLE RATIO The equivalent weight or gram equivalent is the mass of a given substance that will combine with a fixed quantity of another substance. The equivalent weight of a specific element is the mass of the element that will combine with 1.008 grams of hydrogen or 8.0 grams of oxygen or 35.5 grams of chlorine. It can also be determined by an elements atomic weight divided by its valence number. Molar ratio is the ratio of two molecules in a chemical reaction. For example, the equation 2 moles of H2 gas plus one mole of O2 gas goes to 2 H2O molecules is a balanced equation, which leads to a molar ratio of 1:2 between oxygen and H2O because it takes one O2 molecules to make 2 H2O molecules in the balanced equation. The ratio for hydrogen and H2O is 1:1 because they both have equal mole numbers on opposite sides of the equation. The equation ozone or O3 going to oxygen O2 is completely unbalanced as it is because that would leave an extra Oxygen atom left over in the equation. You cannot calculate the mole ratio in an unbalanced equation. But if you balance it you get 2 moles of O3 going to 3 moles of O2 with a mole ratio of 2:3.

KEY TAKEAWAYS •

An atom consists of at least one proton, usually some neutrons, and electrons.

•

Protons and neutrons have sub-particles called quarks that are held together by gluons.

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Protons are positively charged, neutrons are neutral, and electrons are negatively charged.

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There are orbitals around the atomic nucleus containing specific numbers of electrons in them.

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The valency of an atom is determined by how many electrons are in the outer orbital.

•

The periodic table is set up so that elements are arranged according to their behavior chemically.

•

You as a student need to memorize as much of the periodic table as you can in order to be able to do calculations in chemistry.

•

A mole is defined as the amount of a substance that contains exactly 6.02214076 × 1023 particles, such as atoms, molecules, ions or electrons.

QUIZ 1. Which ion has lost an electron? a. PO4b. NH4+ c. Cld. OHAnswer: b. Without knowing these ions well yet, you can recognize that losing an electron means losing an electric charge that is negative. This leads to a positively charged ion of NH4+. 2. What defines the atomic number of a specific atom? a. The number of electrons in the atom b. The number of protons and neutrons in the atom c. The number of protons in the atom d. The number of protons and electrons in the atom Answer: c. The atomic number is defined as the number of protons that exist in the atom. 3. What is the mass number of an atom or element? a. The number of electrons in the atom b. The number of protons and neutrons in the atom c. The number of protons in the atom d. The number of protons and electrons in the atom Answer: b. The number of protons and the number of neutrons in an atom is considered the atomic mass or mass number of the atom.

4. If the atomic number of aluminum is 13 and the mass number is 27. How many neutrons are in the aluminum atom? a. 12 b. 13 c. 14 d. 27 Answer: c. The mass number is going to be the atomic number plus the neutron number. This means that the number of neutrons is 27 minus 13 or 14. 5. What aspect is the same in different isobars? a. Atomic number b. Neutron number c. Atomic mass d. Proton number Answer: c. These substances have atoms that have the same atomic mass because the total number of protons and neutrons together is the same. Their atomic number will be different as they are of different elements. 6. Isobars have the same number of what particles per atom? a. Protons b. Neutrons c. Electrons d. Nucleons Answer: d. Remember that nucleons are the sum total of the neutrons and protons; the number of nucleons is the same in atoms that are isobars of each other.

7. What trait of an element is most closely related to the electronegativity of an atom? a. Ionization energy b. Valency c. Atomic radius d. Electron affinity Answer: d. The electronegativity is most closely related to the electron affinity because both are related to the willingness of an atom to accept an electron. 8. Which symbol is the symbol for iron on the periodic table? a. Au b. Fe c. I d. Ag Answer: b. The chemical symbol for iron is Fe, which explains the name often used to describe iron in a molecule as being “ferrous” or “ferric”. 9. Which is not a nonmetal? a. Carbon b. Oxygen c. Potassium d. Nitrogen Answer: c. With the exception of potassium, each of these is a nonmetal and have different properties; potassium is an alkali metal and is not a nonmetal.

10. Which of these is not a noble gas? a. Argon b. Hydrogen c. Helium d. Xenon Answer: b. Argon, Helium, and Xenon are noble gases; however, hydrogen is listed as a nonmetal and, instead of being relatively inert, has highly reactive properties.

CHAPTER TWO: ATOMIC STRUCTURE The focus of this chapter is the structure of atoms. As it turns out, there are specific rules that define how subatomic structures are organized. This chapter will cover the arrangement of electrons, and both the Pauli principle and Aufbau Principle, which define electron relationships in the atom. Electrons have a natural affinity for the proton they surround but also allow for affinity of different types of atoms in the formation of larger molecules. This will be discussed in this chapter.

ELECTRON CONFIGURATION We’ve briefly touched on the idea of atoms and electrons as they fit into orbitals. This chapter will go on to look at quantum chemistry and how it explains the activity of atoms and elements in chemistry. So, how do you come to understand what quantum chemistry is all about? In chemistry and in physics, a quantum is described as a discrete packet of energy or matter. The plural of quantum is quanta. A quantum means the minimum value of any physical property of a substance. An example of this is the electron. The quantum of charge of an electron is the charge of the electron. This charge can only change in certain increments or energy levels. There is no such thing as half of a charge. In light, a quantum of light is called a photon. All energy, such light, electromagnetic, and others, is absorbed or emitted in certain quanta or “packets” of energy. In the history of quantum mechanics, there is classical mechanics, “old quantum mechanics”, and “new quantum mechanics”. Classical mechanics is what you see in physics: there are forces, acceleration equations, etc. Classical mechanics works for large things that, by chemical standards, aren’t moving very fast. In classical mechanics, things like energy and velocity can have any value from zero to infinity. An item being dropped over a distance will not fall in a jerky fashion but will fall smoothly. This isn’t what is found in quantum mechanics. According to “old quantum mechanics”, there are distinct stepwise patterns of movement in things like electron 28

movement. Electrons don’t pass through a smooth trajectory around an atomic nucleus. They move in quanta, in one place then in the next, without continuous movement. Figure 8 describes the old theory of electron travel, which is wholly incorrect:

Figure 8.

According to “old quantum mechanic theory”, particles are also considered to be waves. This means that, if you try to look at the location of an electron, it cannot be done because it is a wave and is always moving. As soon as you look for its location, it will have moved. This is what’s described in the Heisenberg Uncertainty Principle (although the exact principle is more mathematical than just indicated). Simplified, you can’t identify the position of a wave. So, instead of looking at the orbit of an electron, we look at orbitals, which are defined as wave functions rather than pathways. Orbitals have specific shapes as defined by mathematical probability or the probability of finding the electron in a given point in space. Figure 9 shows the shapes of some orbitals around the nucleus:

Figure 9.

Figure 10 shows another picture of an orbital on an x and y-axis:

Figure 10.

The pictures usually show an orbital in two different colors. This is because a wavelength can be positive or negative; it is important to understand this because it helps determine how an electron helps to form bonds with other atoms. Electrons are a lot like the waves of an ocean, where there are peaks of the waves and troughs of the waves. When two waves interact, they can “add positively” to one another when each wave is at its peak, “add negatively” from one another when each wave is at its trough, and cancel each other out, when one is at its peak and one is at its trough. As you have learned, there are different orbital shapes and different numbers, called s, p, d, and f. The smallest orbital is orbital s, which is seen in helium and hydrogen, as well as Li+2 (lithium two-plus). When there is more than one electron, they interact with one another and will split the orbitals into slightly different energies. In any given shell, the energy of the orbitals depends on the angular momentum of the orbitals s, p, d, and f in order of lowest to highest energy. Figure 11 shows the energy of the different orbitals:

Figure 11.

The simplest and lowest-energy orbital is the 1s orbital. It has no nodes. A node is the point at which the probability of finding an electron is zero. There is only one 2s orbital per shell and the s sub-orbital can only hold 2 electrons, as long as they have different quantum spin number . S orbitals have the ability to be involved in bonding. The p orbitals are the second set of orbitals. Each orbital has one angular node with three choices for the magnetic quantum number, indicating three p orbitals differently oriented. Since each orbital can accommodate two electrons having opposite spins, the p orbital capacity is six (because two times three is six). P orbitals are also often involved in bonding. Each d orbital has two nodes with five choices of magnetic quantum number, which leads to five different d orbitals. Remember that each of these orbitals can have two electrons of opposite spin, giving rise to a total of 10 electrons total capacity. Atoms with d orbitals tend to be larger so they are usually involved in bonding only in inorganic chemistry situations. The first set of f orbitals is the 4f subshell. There are seven possible magnetic quantum numbers, leading to seven f orbitals. As each can hold two electrons with opposite spins, there are fourteen f electrons in total capacity.

When looking at orbitals, the principal quantum number n says how big the orbital is and how many nodes the orbital has. The n values go from n = 1 and on, but there are only so many orbitals because there are only so many elements. The number of spherical nodes is equal to n minus the angular momentum number minus one. The s orbital is spherical and therefore has no nodes (where there isn’t the possibility of an electron). What is known fairly clearly is that in each shell of orbitals, the orbitals with the lowest energy are the most stable so that electrons line up in the s orbital first, followed by the p orbital, the d orbital, and finally, the f orbital. Orbitals with the same energy are referred to as being “degenerate”. The more things happening around an atom (such as a magnetic field, neighboring atoms, added electrons) the fewer orbitals in each shell become degenerate. This means that, when you make a chemical bond, the orbitals are no longer degenerate. So, to determine the electron configuration for an element, you need to know how many electrons there are in the atom. In general, the atomic number equals the number of electrons unless the atom is an ion that has added or subtracted electrons. There are stable orbitals that are completely filled and orbitals that contribute to the valence electrons we’ve already covered. Valence electrons are the outer shell electrons that will participate in electronic bonding. Figure 12 shows an image of the electrons around an atom:

Figure 12.

If there are many electrons, they must first go into the orbitals closest to the nucleus. When these get filled up, they go outward to fill higher energy orbitals. As we have already discussed, each suborbital can only contain 2 electrons of opposite spin. There are two spin numbers, -1/2 or +1/2. According to the Bohr model, the number of electrons per shell follows a specific pattern, which is set up clearly in the periodic table: 2, 8, 8, 18, 18, 32, and 32. Take a look at the periodic table again. In the first period, there are two elements in the period (hydrogen and helium), in the next two periods, there are eight elements each, in the next two periods, there are 18 elements each, and in the last two periods, there are 32 elements each.

PAULI PRINCIPLE We’ve already talked about what the Pauli Principle or the Pauli Exclusion Principle is but we didn’t call it that. According to the principle, in any given atom or molecule, no two electrons can have the same four electronic quantum numbers. Because each orbital can contain only two electrons, they must have opposite spins. One must have an up spin of +1/2, while the other one is said to have a down spin of -1/2. For hydrogen, there is just one electron, so it is said to have an electron configuration that is 1s1, while for the helium atom, there are two electrons, so the configuration is 1s2. The configuration for Beryllium is 1s22s2, with two up-spin electrons and two down-spin electrons.

AUFBAU PRINCIPLE The Aufbau Principle isn’t named after anyone in particular but is named after the German word “aufbauen”, which means to build up. What this means is that you can use the knowledge of orbitals to build up to higher energy levels as the atom gets bigger. You should be able to use this principle to build the number of electrons in the atom. Because of the Pauli Exclusion Principle, no sub-orbital can have an electron with the same spin or the same set of quantum numbers, so they have to move to higher energy levels. Without this principle and the Pauli exclusion principle, the electrons would just cluster as close to the proton as possible, making the atomic radius not change or even get smaller as the proton charge increases. Because there are orbitals, the electrons have places to go, even though they are at higher energy levels. We can use the Aufbau Principle to determine electron configurations. There is a pattern that can be identified when determining electron configurations, which is described as follows:

Remember that Z is the atomic number and that this is also the number of electrons in the atom: •

Hydrogen—Z = 1, with the configuration of 1s1

•

Helium—Z = 2, with the configuration of 1s2

•

Lithium—Z = 3, with the configuration of 1s22s1

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Beryllium—Z = 4, with the configuration of 1s22s2

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Boron—Z = 5, with the configuration of 1s22s22p1

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Carbon—Z = 6, with the configuration of 1s22s22p2

This goes on up to neon, with a Z of 10 and an electron configuration of 1s22s22p6, which fills up the p orbitals at the second level. What you should know is that, with three p orbitals and with higher energy orbitals, they do not fill up an orbital and then another orbital. They “spread themselves out” so that, when there are three electrons in the p orbitals, one will be in each of the three p sub-orbitals, until they are filled out before adding an electron to one of the three orbitals that so far have just one each. By the third period, things get a little more complicated. The first level is filled and the 2p orbitals are filled. Now comes the 3-level or third level. It starts with the 3s1 electron and then to the 3s2 electron before filling up the 3p orbitals for a total of 10 more spots to fill in the third period. This leads up to Argon, which fills up the entire third level of the first three levels. By the fourth period, there are the 4s orbitals, the 4 p orbitals, and then the 4d orbitals. The lower energy 4s orbitals get filled up first and then 3d orbitals. Notice how there are ten elements across the transitional metals in the fourth period. This is because these atoms are filling up the 5 d orbitals, which are of a lower energy level than the 4p orbitals. The transitional elements are referred to as the d-block elements because the d orbitals of the previous number get filled up first, before the p orbitals corresponding to the period or “row”. While the terms d-block element and transitional element are not technically the same, they can be considered the same for the purposes of this exercise. With the d orbitals, they are filled up just the same way as in p orbitals in that the electrons will spread out over the orbitals until the 5 orbitals are filled and then will go 35

up to fill up the orbitals that already have an electron in them. This leads to a total of 10 electrons in the five orbitals per period. As always, the spins of the electrons will be opposite when they do inhabit the same orbital. You need to know that it isn’t exactly written in the way you’d think. Rather than write Scandium, for example as 1s22s22p63s23p64s23d1, which would make the most sense, it is instead written with the 3 level orbitals clumped together and the four level orders clumped together, like this: 1s22s22p63s23p63d14s2. This is made even more confusing by the fact that the sequence for chromium is written differently, according to the levels that get filled up first, which makes the chromium sequence go like this: 1s22s22p63s23p63d54s1. So, the rules for filling up electrons in orbitals are that first the Periodic Table is used to find the atomic number or the total electron number. Next, fill up the 1s, 2s, 2p, 3s, 3p, 4s, 4d, and 4p orbitals except that in copper and chromium, the pattern is broken up and has the d orbital and s orbital pattern switched as was described.

ELECTRON AFFINITY The definition of electron affinity is the change in energy (as listed in kilojoules or kJ per mole) in the gaseous phase when an electron is added to the atom in order to form a negative ion. It is basically the likelihood that a neutral atom will gain an electron. It is through chemical equations that atoms gain or lose electrons. A reaction that releases energy is called an exothermic reaction, while a reaction that absorbs energy is called an endothermic reaction. Because energy is released in an exothermic reaction, the energy is given a negative sign. The opposite is true of an endothermic reaction, which is given a positive sign. When an electron is added to a neutral atom, this results in the release of energy, giving the reaction a negative sign. This holds true only for the first electro affinity so first electron affinities are negative. How it looks is this: Atom + an electron goes to a negatively charged ion plus energy (which makes it exothermic). When it comes to adding a second electron; however, there is more energy required to add an electron to what is already a negatively charged

ion. This overwhelms the energy that is given off when the electron is added so the net energy is positive or endothermic. Ionization energy, which we’ve already talked about, is essentially the opposite of electron affinity. This concept of electron affinity is usually seen for those electrons in groups 16 and 17, as these are the most likely to gain an electron to fill their orbitals. The first electron affinity is described as the amount of energy released when one mole of a gaseous atom becomes a mole of a negatively charged gaseous atom, having taken on one electron. The number is different for every atom. An example is chlorine, which has an electron affinity of -349 kJ per mole. According to convention, this is a negative sign because energy is released. While the halogens have a highly negative electron affinity, metals do not. They require a great deal of energy to take on an electron and are much more likely to form cations, which are positively charged ions. Metals, unlike halogens, have a weaker pull on their electrons and will more easily lose them. When nonmetal elements gain electrons, the change in energy is usually negative because they give off energy in an exothermic reaction to form an anion (a negativelycharged ion). Nonmetals, as you know, have more valence electrons than metals do so it is easier to fill a stable octet in the outer shell. Nonmetals will have a greater electron affinity than is true of metals. The electron affinity is highest for chlorine than with any other atom—even higher than with fluorine, which technically should have the highest affinity when you look at the periodic table. We will talk about why this is true in a minute. The general pattern of electron affinity is for the affinity to increase as one goes across the periodic table from left to right as well as when the atoms get smaller. The fewer valence electrons an atom has, the less likely it is to want to add another one and the more likely it is to want to give the ones it has away. Larger atoms do not have as strong of a pull from the protons because of the shielding effect of existing electrons. The stronger the attraction of the proton (the smaller the atom), the more energy is released when the electron is held onto.

Fluorine has an electronic signature of 1s22s22px22py22pz1. It has just 9 protons in the nucleus, with an incoming electron entering the 2-level of orbitals. It has only 2 shielding electrons. Chlorine has an electronic signature of 1s22s22p63s23px23py23pz1 with 17 protons and 10 shielding electrons. The reason why the electron affinity is higher for chlorine than for fluorine is because fluorine is so small that it is crowded with electrons already so there is slightly more repulsion of these electrons in fluorine than it is with chlorine. The same thing happens between oxygen and sulfur, which are of the same group but reverse the expected trend. The issue of second electron affinity is only seen in Group 16 elements oxygen and sulfur. Both of these will form a -2 (negative two) anion. The second electron affinity is described as the energy required to add a second electron to a mole of gaseous ions to produce a mole of negative two anions. This takes energy because, in a sense, one is adding a negative electron to something that is already negative. For oxygen, the first electron affinity is -142 kj per mole. The second electron affinity is +844 kJ per mole, which makes it an endothermic, energy-requiring process.

ELECTRONEGATIVITY Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons. Fluorine is the most electronegative element, having an assigned value of 4.0. This is the value given to it by the Pauling Scale. When two atoms of the same element come together, they have the same electronegativity and therefore have electrons that are on average halfway between the two atoms. This is seen in H2 gas or Cl2 gas. In most cases, however, there is not an equal sharing of electrons between two substances bonded together. If one element or atom is slightly more electronegative than another element, they are described as not sharing electrons completely equally. The way this looks is that there will be one atom that is called delta-positive (meaning slightly positive) and one that is delta-negative (meaning slightly negative). In such cases, there will be a “covalent” bond, which is a bond that shares electrons from between the two atoms. On the other

hand, when one electron is highly negative (taking on more of the electrons) and one is highly positive, the bond is considered “ionic”. In fact, there is a spectrum of bonding, with no clear-cut division of bonding. A pure covalent bond is one between identical atoms. Even with highly ionic bonding that is seen in things like sodium chloride, one atom does not have complete control over the electron. There is always some sharing happening. There is some bonding that is somewhat ionic and some that is covalent. When it comes to electronegativity, there is a general trend from the lower left to the upper right-hand corner of the periodic table (not counting the noble gases). It isn’t perfect so you need to memorize the trend of electronegativity, going from highest to lowest, including the following: •

Fluorine

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Oxygen

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Chlorine

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Nitrogen

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Bromine

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Iodine

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Sulfur

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Carbon

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Hydrogen

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Metals

DIPOLE MOMENTS Dipole moments happen when there is a separation of charge. This can occur between two ions in an ionic bond or in the situation of a non-pure covalent bond. It happens whenever there is a difference in electronegativity between two atoms. The larger the

difference in electronegativity, the larger the dipole moment. The dipole moment also increases as the proton and electron are separated. In addition, the dipole moment also is affected by the size of the atom and the pull of the protons on the electrons. It is measured in Debye units, which is equal to the distance between the charges multiplied by the actual charge. H20 or water is a molecule that has a dipole moment, which has a plus-delta on the hydrogen ions and a two-minus delta on the oxygen. The dipole moment exists because the molecule is not straight so the delta charges do not cancel each other out. By convention, the arrow of the dipole moment goes from a positive side to the negative side (which is the opposite of that seen in physics). Because it isn’t a straight molecule, there is a dipole moment over the whole H2O molecule and not just between the atoms, making this a polar molecule. Polar molecules tend to align themselves in one of three ways: •

In an electric field

•

With respect to one another

•

With respect to ions

This will become important to understand when thinking of why water is such a good solvent for different salts. Being polar means that it can dissolve other polar molecules, including ions like sodium and chloride in water. Another comparison of dipole moment is in the molecule CO2. This, unlike with water, is a straight molecule. There is a dipole moment that separates the charge between the carbon and each oxygen molecule; however, because it is a straight molecule, the net dipole moment is zero because they cancel each other out. With water, the angle of the molecule is 104.5 degrees, making it angle enough to have a net dipole moment and making water a polar molecule. Figure 13 is a depiction of what the water molecule looks like:

Figure 13.

MAGNETIC PROPERTIES IN ATOMS The magnetism of any system refers to the magnetic dipole moment. This applies to things you can see, such as a magnet, as well as things you cannot see, like atoms and electrons. As you will find out, there are different types of magnetism, including diamagnetism, paramagnetism, anti-ferromagnetism, and ferromagnetism. Paramagnetism refers to the magnetism of an atom that has at least one unpaired electron. Because of the unpaired electrons, there will be a dipole moment and an attraction to positively-charged particles. According to Hund’s Rule, electrons must occupy each orbital of equal energy before pairing up with another electron in the same orbital. This can lead to an atom with many unpaired electrons. These are paramagnetic substances. Diamagnetism involves having paired electrons. When electrons fill up the orbitals, paired electrons will tend to cancel each other out because they have opposite spins. These substances will be weakly repelled by a magnetic field. The electron configuration or electron signature will tell if a substance is paramagnetic or diamagnetic. This is why it pays to understand how to figure out the electronic signature of an atom or element. You can see if the orbitals are unpaired or paired.

KEY TAKEAWAYS •

There are specific orbitals that determine the likely place that electrons will be in an atom.

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The electronic signature will be determined by the orbitals that are filled from the lowest energy to the highest energy possible.

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The Pauli principle states that electrons in the same orbital must have opposite magnetic spins.

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The Aufbau principle allows for the building of electrons from the lowest to highest energy.

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Electron affinity is related to electronegativity and tends to increase in the upper right-hand corner of the periodic table.

•

A polar molecule is one with a net positive dipole moment.

QUIZ 1. What is the first orbital filled at every level? a. s b. p c. d d. f Answer: a. The s orbital, and the only orbital in a hydrogen or helium atom, is the first orbital filled. This is a spherical orbital and is at the lowest energy level. 2. Which is not an orbital name in understanding the orbitals around atomic nuclei? a. p b. d c. e d. f Answer: c. The orbitals in the atoms that have been discovered thus far only go up to the following, in order: s, p, d, and f. There is no e orbital. 3. What charge is seen on an electron that has an up-spin? a. +1 b. -1 c. +1/2 d. -1/2 Answer: c. An up-spin for an electron is a spin of +1/2, while a downspin electron has a spin of -1/2.

4. According to the Pauli Exclusion Principle, what must be the case? a. Electrons must be in the lowest energy state possible. b. Electrons must fill the p orbitals first c. Electrons will stay at closest to the nucleus as possible. d. Two electrons in the same sub-orbital cannot have the same spin. Answer: d. According to the Pauli Exclusion Principle, electrons cannot have the same spin if they are to be in the same sub-orbital. 5. What type of element is seen in the d-block on the periodic table? a. Alkali metal b. Transitional metal c. Halogen d. Noble gas Answer: b. The transitional metals are in the d-block because they fill up the d orbitals of a lower energy class than they do of the p orbitals. For example, the 3d orbitals get filled after the 4s orbitals but before the 4 p orbitals. It’s all about the energy level of the orbital. 6. Which period has the higher energy d-orbital type in it? a. First period b. Second period c. Third period d. Fourth period Answer: d. The fourth period starts the d block elements, which are technically thought of as being a higher energy orbital.

7. Which atom has the highest electron affinity among other atoms? a. Fluorine b. Chlorine c. Bromine d. Potassium Answer: b. Chlorine has the highest electron affinity among the other atoms, even though it isn’t in the top right-hand corner of the periodic table, which is the natural trend of electron affinity. 8. Which group has the greatest electron affinity among the elements? a. Group 15 b. Group 1 c. Group 18 d. Group 17 Answer: d. The group that has the greatest affinity for an electron is group 17 because they need that extra electron to fill their orbitals. 9. What kind of bonding is seen when there are two atoms of the same type bonded together? a. Hydrogen bonding b. Covalent bonding c. Ionic bonding d. Pure covalent bonding Answer: d. When there is a bond between the atoms and the two atoms are identical, this is called pure covalent bonding.

10. What kind of bonding is seen when there is one atom that is highly electronegative and one that is not very electronegative? a. Hydrogen bonding b. Covalent bonding c. Ionic bonding d. Pure covalent bonding Answer: c. Ionic bonding is seen when an atom is highly electronegative and bonds with another atom that is not very electronegative.

CHAPTER THREE: PROPERTIES OF GASES The main topics of this chapter are the properties of gases. There are specific laws that describe the behavior of gases in chemistry, which are covered in this chapter. The study of gases involves a relationship between volume and pressure and, in the situation of mixed gases, a relationship between the different gases in a given volume. The behavior of gases in relation to the ideal gas law is covered in this chapter as well as the study of the kinetics of gas molecules in a gaseous solution.

GASES AND THEIR PROPERTIES Many centuries ago, in the 1600s, the invention of a sensitive scale determined that gases were matter that had weight. Air, oxygen, and other gaseous substances have weight and are made up of particles—completely separate things from vacuums, which have an absence of particles. This led to the definition of gases as being particles that have few interactions and interferences with each other (but, as you’ll see, the operative term is “few” as gases do have interactions with each other). Gases, by definition, have no definite volume or shape. They fill whatever volume they are contained in and, unlike liquids, they do not have a distinct upper surface. Gases have low densities when compared to liquids and solids. Liquid water (the molar volume) is 1000 times denser than the molar volume of gaseous water (at one atmosphere and at 298 Kelvin temperature). Gases, interestingly enough, all behave the same way when it comes to pressure and temperature, expanding or contracting by an amount that can be predictable. This is not true of liquids and solids. In this chapter, we’ll study the macroscopic and microscopic characteristics of gases. Gases, as you probably know, have molecules that are in continuous motion, striking the inner walls of the container they are in. The reversal of direction when they do this doesn’t lose the kinetic energy of a gas but will exert a force on the container walls. The

force divided by the total surface area of the container walls is called the gas pressure. Figure 14 describes the differences between gases, liquids, and solids:

Figure 14.

The pressure of a gas can also be described as the pressure that needs to be applied internally by a piston in order to keep the gas within the container. Imagine a container with a piston on top. The force can be defined as the weight of the piston necessary to keep the piston from pushing out of the container. Pressure is equal to this force divided by the cross-sectional area of the piston. There is a specific unit of pressure that has been defined in the SI (standard international) system. This is called the Pascal or Pa. This is the force of one newton per square meter. At the surface of the earth, the air pressure is 9.81 Pascal units. In chemistry, gas is also measured millimeter of mercury or mm Hg. This is defined as on millimeter of level distance in a mercury barometer at 0 degrees Celsius. In meteorology, the pressure unit most commonly used is the bar (think of isobars in studying meteorology). One bar equals 750.06 torr equals 0.987 atmospheres.

When studying the air pressure on earth, you can consider this as being a column of air of a certain square area. As one moves higher in the atmosphere, this column will have decreased pressure. This means that there is a direct relationship between the height of the column and the pressure of the gas. The classical barometer measures pressure of gases. It is the height of a column of liquid that can be supported by the atmosphere. The liquid is usually mercury contained within a column that is topped by a vacuum. Metallic mercury is a liquid, making it the ideal barometric fluid. The height of mercury at sea level is shown in figure 15:

Figure 15.

Certainly, water can be used as a barometer but you would have to look at the density of the two liquids. The density of mercury is 13.6 grams per cubic centimeter with a height at atmospheric pressure of 76 centimeters in a 1 centimeter squared cross section. Because water is just a fraction of the density of mercury, the same sized barometer of 1 cm cross-sectional area would be 10.3 meters, making it impractical to be used as a barometer. Figure 16 is an example of a U-tube barometer, in which one end is open to the atmosphere and the other is connected to a measuring device.

Figure 16.

There are two types of manometers. In the closed-tube manometer, gas is let onto one side of the U, leading downward pressure on the mercury and raises it on the opposite side of the U. The other side of the U is closed with a vacuum above the mercury column. The pressure difference is the difference in height of mercury on both sides of the U-shape, measured as the torr pressure of the gas in the chamber. A pressure gauge is used to measure the pressure in the system. In an open tube manometer, the U shape has gas on one side at a certain pressure, mercury in the U and no top to the other side of the U (which is instead open to the atmosphere). The gas pressure is then measured against atmospheric pressure.

A WORD ON TEMPERATURE In general, in chemistry and in physics, if two bodies are adjacent and one body is at a higher temperature, the tendency is for heat to flow from the warmer side to the colder side until the temperatures are the same. This is what a thermometer is based on. The change in temperature is calibrated on a thermometer, such as a mercury thermometer. A thermometer makes use of some temperature-dependent quantity, such as a liquid’s density. Again, mercury is used for convenience. The reading on the temperature scale is completely arbitrary and has to be calibrated. Celsius or “centigrade” uses the freezing point of water as the zero point and the boiling point of water as being 100 degrees at one atmosphere of pressure. Fahrenheit is an older measurement, in which water freezes at 32 degrees Fahrenheit and boils at 212 degrees Fahrenheit. One-degree Fahrenheit is 5/9 of a degree Celsius. There is also absolute temperature. This is the Kelvin scale, which uses the Celsius scale in the sense that a degree means the same thing, except that zero is -273 degrees Celsius, which is the absence of any thermal energy. In engineering, the Rankine scale is used, which is an absolute scale based on the Fahrenheit degree.

GAS VOLUME The volume of a gas is basically the space in which gas molecules are free to move. Mixed gases will occupy the same volume. A gas burette can measure the volume of a gas by trapping it above mercury in a calibrated tube. The SI unit of volume is the cubic meter but in chemistry, this is impractical so that liter and milliliter are used. Another unit used in chemistry that you should know about is the cubic centimeter. A cubic centimeter or cc is the same as one milliliter or ml of a substance. You should know that just reporting the volume of a gas is not very helpful. The volume will vary according to the temperature and the pressure of the gas. What is commonly done in chemistry is to report the volume of a gas at ambient temperature and standard atmospheric pressure, standardizing it to 25 degrees Celsius or 0 degrees Celsius.

PRESSURES AND GASES Pressures can be determined and are described in liquids and gases alike. With gases, there is pressure exerted on everyone by the atmosphere but it is something we have become accustomed so it is generally ignored. Water pressure, as is seen with swimming, is greater than air pressure and is something that can be felt when moving through water. As mentioned, pressure is the force divided by the area on a firm substance against it. When a balloon is blown up, the balloon expands because its pressure is greater than atmospheric pressure against the walls of the balloon. When the balloon pops, the air flows from an area of high pressure to an area of low pressure. In water pressure, the pressure increases as one goes deeper, and in air pressure, the pressure decreases at high altitudes. At 20,000 feet above sea level, air pressure is cut in half because half of the whole atmosphere of the earth is above and below this level. At sea levels, the atmospheric pressure is about 14.7 pounds per square inch. Car tires need a pressure of about 30 pounds per square inch. Bicycle tires have a pressure of about 60 pounds per square inch or psi. Standard atmospheric pressure at sea level is exactly 760 millimeters of mercury of mm Hg; 760 torr is one atmosphere, which is the air pressure arbitrarily set at sea level. One atmosphere is about 1 x 105 Pascal (which is the SI unit for pressure---the same as about 1 kilogram per meter per second squared). In reality, one pascal is equivalent to one newton (1 N) of force applied over an area of one meter squared (1 m2). The barometric pressure is measured in mm Hg (or millimeters of mercury). The barometer was invented by Evangelista Torricelli in the early 1600s, and uses mercury as the standard of weight. As you have seen by the U-shaped tube in figure q6, the mercury is put into a U-shaped tube and is left open on one end to measure the pressure of a certain gas against the atmospheric pressure. When measuring the atmospheric pressure force equals mass times acceleration. The acceleration is based on the Earth’s gravitational pull, which is about 9.8 meters per second squared.

GAS LAWS There are three fundamental gas laws that have been developed to describe the relationship between temperature, pressure, and volume of gas. Before getting into the three gas laws (plus the ideal gas law), you need to know the difference between the ideal gas or “perfect” gas and real gases. The ideal gas has three properties: •

It has particles that are so small that they do not occupy any space of their own.

•

It has a constant, straight line, and random movement of particles.

•

Particles only collide elastically with the walls of the container and with each other with no forces between the particles of the gas.

Real gases differ from ideal gases in two specific ways. First, as the pressure increases, the volume does not approach zero because, in fact, the molecules do occupy space and have a certain volume. Second, there are intermolecular forces between gas molecules, particularly at low temperatures when molecular motion slows down. At high temperatures and even at moderate temperatures, the intermolecular forces are considered negligible, however. Gases do behave more ideally at high temperatures. This is because there is expansion of the gas so the molecular volume (the volume of each molecule) does not play a major role. In addition, the molecules are moving so fast that they do have an elastic behavior when bumping against each other and against the walls of the container they are in. Gases are most non-ideal when they are at cold temperatures, which allow for attractive or repulsive forces to take place between molecules. They also do not behave ideally at high pressures because the volume of the molecules becomes a factor. The molar volume can be worked out of a gas if you know its density at a particular pressure and temperature. For gases, the density is highly dependent on the temperature and pressure. This means that when referring to the density of gases, we are talking about the density at STP, where STP means standard temperature (0 degrees Celsius) and pressure (o ne atmosphere). Standard gas density is the density of gas at STP. The official IUPAC unit for gas density Is kilograms per meter squared, which is also one gram per liter. 53

This standard gas density is used to measure the molar mass of a gas. The density of helium, for example, is about 0.18 grams per liter at STP. The mass of a mole of helium is 4 grams, making the molar volume or the volume a mole of Helium about 22.4 liters. As it turns out, most of the gases have a molar volume at STP of about this same amount. For this reason, when using an equation that uses molar volume, 22.4 liters is used, even though this isn’t 100 percent accurate. The most ideal gas is one with the smallest molecules and the lowest chance of any type of intermolecular forces. This gas is helium because it has an atomic number of 2 and doesn’t have any forces between molecules. It is also a single atom, unlike hydrogen, which is an H2 gas consisting of two molecules together. Larger gas molecules will take up more space, deviating from the ideal gas, and other gaseous molecules will often interact with other gases, leading to a less ideal situation.

BOYLE’S GAS LAW Boyle’s gas law was defined by Robert Boyle in 1662 and involved the relationship between pressure and volume, assuming the temperature and amount of gas (n) is constant. What the law states is that the Pressure is inversely proportional to Volume. This graphs out according to Figure 17:

Figure 17.

The law states that P1 times V1 = P2 times V2 so that, if you know any volume and pressure at a specific point you can calculate the pressure and volume at a different point, if the temperature and number of molecules is the same.

CHARLES’ LAW This was devised in 1787 by French physicist Jacques Charles. This is a correlation between the temperature and the volume, assuming that the pressure (P) and Amount of Gas (n) are constants. According to the Law, Volume equals a constant multiplied by the Temperature. The constant depends on the number of gas particles and the pressure. This is a linear graph that indicates that volume increases with temperature in a linear fashion.

AVOGADRO’S LAW This law was developed in 1811 by Amedeo Avogadro, which correlates the amount of gas particles (n) as it relates to the volume when the temperature and pressure remain constants. According to this law, volume is directly proportional to the number of gas particles. In other words, volume is equal to a constant multiplied by the number of molecules. This is a linear relationship. There is another law that isn’t as well known, called the Amontons’ Law, which indicates that the pressure is directly proportional to temperature, given a constant number of moles of a gas and an unchanging volume. Ideal Gas Law The problem with the laws governing gases is that there is no such thing as the ideal gas. Nevertheless, chemists in history have managed to devise an equation that uses aspects of all the previous gas laws combined, even though, strictly speaking, there is no gas that will follow this law during all circumstances. The law must assume that particles do not interact with one another and that atoms do not take up space. The ideal gas law or the “ideal gas equation” combines four variables: pressure (P), volume (V), number of moles of gas (n), and temperature (T). There is a constant R

called the gas constant. According to the law PV = nRT or Pressure multiplied by volume is equal to the number of moles of gas multiplied by the temperature (which is also multiplied by the constant R). Mathematically, this is also described as this: PVnRT = 1. This term is also described as the compression factor, which determines the ideality of a given gas molecule. Pressure can be listed in atmospheres (atm), Torr, Pascal units (Pa), or mm Hg. Volume is listed in liters or meters cubed. Number of molecules (n) is the moles of a gaseous substance. Temperature is listed in Kelvin units. The gas constant R is unfortunately different according to changing pressure and volume. The value of R is different according to what units are used in determining the answer. For example, R = 0.082057 Liters atmospheres per mole per kelvin degrees. This number will be different, depending on the units for volume and pressure. What this means is that you have to use the constant that applies to the pressure and volume units you are using but not on the number of molecules or temperature because mole and degrees Kelvin are used in all the constants. At high pressures, the ideal gas law falls apart because molecules take up space and the situation gets crowded. The volume that the molecules exist within is effectively less. In such cases, the total space available is the total volume minus the actual volume of the gas molecules. For a given pressure, a real gas will end up taking up a greater volume than can be predicted by the ideal gas law because of this additional molecular space. As mentioned earlier as well, the ideal gas law falls apart at lower temperatures because the molecules move less and have a greater chance of interacting with one another in such “slow” situations. This will decrease the volume the gas will take up at the same pressure. The van der Waals equation will best predict the behavior or real gases. This incorporates the effect of the gas molecule volume as well as the temperature effect. The equation is complex but suffice it to say that the equation takes into account corrections for low temperatures and high pressures when dealing with gases.

PARTIAL PRESSURES IN GASES When gases are together in a mixture, they each have a partial pressure. According to Dalton’s law (which is also called the Law of Partial Pressures), the total pressure exerted by a mixture of gases is equal to the sum of the partial pressures of the gases in the mixture. According to the kinetic theory of gases (which will be further elaborated on soon), gases will diffuse in the container it is in in order to fill up the space it exists in and there is no force of attraction between the molecules. In other words, the different molecules in a mixture of gases are so far apart that they act independently from one another and do not interact. In an ideal gas, the collisions are between the gas and the container and not between each other so that the partial pressure isn’t affected by the other molecules in the gaseous chamber. According to Dalton’s law, the total number of moles in a gaseous substance equals the sum of the moles of each gaseous molecule in a mixture. The mole ratio or Xi can be calculated by knowing the partial pressure of a certain gas and the total pressure of a certain mixture of gases. It describes the fraction of the mixture that is made up by a given gas. If oxygen, for example, takes up 4 atmospheres in a system consisting of 10 atmospheres, the mole ratio of oxygen is 4 divided by ten or 0.4. As you can imagine, the sum of the mole ratios of all the gases in a mixture must equal to one. According to the ideal gas law, all gases behave exactly the same and their behavior is independent of repulsive and attractive forces. The only thing that is important in partial pressures is the number of molecules and not any other aspect of the nature of the gases themselves. In order to determine the partial pressure of a gas in a mixture, you need to know the molar mass of the molecule (the atomic mass) so that you can calculate the number of moles of the substance from the grams that might be given in an equation. This is why it is a good idea to have a periodic table that gives the atomic weight or atomic mass of the element.

KINETIC THEORY The kinetic theory or kinetic-molecular theory of gases helps to explain the behavior of a hypothetical ideal gas. According to the theory, gases are made from tiny particles that move in a random and straight-line fashion. They do not move slowly and they do not stop moving at normal temperature, making collisions with each other and with the walls of the container they are in. This was the first theory that is based primarily on molecular collisions rather than static forces between the molecules. According to the kinetic theory, different sizes of the gaseous particles can give them different individual speeds. According to this theory, the following things are true: •

Gases consist of constantly and randomly moving particles that travel in a straight line.

•

Gas particles are point masses that have no volume of their own.

•

There are no attractive or repulsive forces in place.

•

Gas pressure comes directly from collisions between the molecules and the walls of the container.

•

No energy is lost in the movement or collisions of the molecules.

•

There is no time for the collision itself compared to time between collisions.

•

The kinetic energy of a gas is a measure of the Kelvin temperature.

•

Individual gas molecules behave differently by size but the total kinetic energy is an average of their speeds.

•

The kinetic energy is directly proportional to the temperature.

•

All gases at the same temperature will have the same average kinetic energy.

•

Lighter gas molecules move faster than heavier gas molecules.

The temperature of a gas is the measure of its kinetic energy. The kinetic energy of a particle is related to its velocity according to this equation: Kinetic energy equals 12

multiplied by the mass of the molecule multiplied by the square of its velocity. The thermal energy of a system is the total kinetic energy of all particles of a system. As you can see by the equation, the kinetic energy is closely related to its velocity. Kinetic energy theory helps to lay down the three major laws of gases as already described.

GRAHAM’S LAW OF EFFUSION The process of effusion involves the escape of gas molecules through a small hole into a vacuum. The assumption is that, during the exit, there are no collisions on the molecule that is exiting. According to this law, the lighter molecules will exit first because they are moving faster. The relationship has been described in the law stating that the rate of effusion of a gas is inversely proportional to the square root of the mass of its particles. Diffusion is more complicated than effusion because there are intermolecular forces involved in diffusion that are not presumed to be a factor in effusion. The molecules that diffuse across an area will collide with existing molecules in the space and will not exit into a vacuum that is the case with effusion. The truth is that, along with effusion, molecules will differ in their rate of diffusion in that larger molecules will diffuse at a decreased rate compared to lighter and less dense molecules.

KEY TAKEAWAYS •

There are gas laws that exist which correlate with certain values of gases holding the other values constant.

•

The ideal gas law encompasses all gas laws and intertwines pressure, volume, temperature, and number of molecules.

•

The pressure of a gas is related to the collision of molecules against each other and against the container the gas is in.

•

The partial pressures of separate gases add together to make the total pressure of the gaseous mixture.

•

Kinetic energy is directly related to the temperature of a gas.

•

Diffusion and effusion are related and both are dependent on the size of the molecule.

QUIZ 1. What is the best way to describe the pressure of a gas? a. The weight of the gas when it becomes liquid b. The force of the gas on a container c. The mass of one mole of a gas d. The force of the gas per a certain square area of a container Answer: d. Pressure of a gas is defined as its force per square area against a container, which is different depending on things like the temperature of the gas. 2. Pascal units are SI units involving what measurement? a. Pressure of a gas b. Force of a gas c. Weight of a gas d. Volume of a gas Answer: a. The Pascal unit or Pa of a gas is the pressure of a gas. In the measurement of air pressure at sea level, the pressure of air is 9.81 Pa. 3. What is the measurement known as a Pascal? a. One kilogram per meter squared of force b. One newton of force per meter squared c. One atmosphere at sea level d. One millimeter of mercury at sea level Answer: b. The definition of a Pascal is one newton of force applied over one meter squared. This is the SI unit of measurement for chemistry purposes and is abbreviated as Pa.

4. There are three things described in the three gas laws for real gases. What is not described in these three laws? a. Pressure b. Volume c. Temperature d. Kinetic energy Answer: d. The three gas laws that have been described include measurements of pressure, volume, and temperature of gases but not kinetic energy. 5. Boyle’s law of gases can predict the relationship between what two factors of gases? a. Pressure and Temperature b. Temperature and number of molecules c. Pressure and Volume d. Volume and Temperature Answer: c. Boyle’s law indicates that pressure and volume are inversely proportional to each other when the number of molecules and temperature stays the same. 6. Charles’ Law predicts the relationship with what two factors of gases? a. Pressure and Temperature b. Temperature and number of molecules c. Pressure and Volume d. Volume and Temperature Answer: d. The volume and temperature are linearly related to one another in Charles’ law so that, as the temperature increases, so does the volume.

7. The mole ratio of a given gas can best be described as what? a. A comparison of the weight of a gas compared to the weight of other gases in a mixture. b. A comparison of the weight of a gas compared to the total weight of the mixture. c. The ratio of gas A in a system versus gas B in a system by number of molecules. d. The ratio of the molecule number of a gas compared to the total number of molecules in a mixture. Answer: d. The mole ratio is based on the number of molecules; it is the number of molecules of a gas in a mixture of gases compared to the total number of molecules in the mixture. 8. In measuring gases, how many torr units equal to one atmosphere of the gas? a. 273 b. 760 c. 1080 d. 100 Answer: b. The equation that explains this is that 760 torr (or 760 millimeters of mercury) is the same as one atmosphere, which is the atmospheric pressure at sea level. 9. At a given temperature, the kinetic energy of a gaseous mixture is based on what factor? a. The average of the kinetic energies of the different molecules. b. The sum of the kinetic energies of the different molecules. c. The multiplication of the different kinetic energies of the molecules. d. It is a given number that is irrespective of the different molecules. Answer: a. The kinetic energy in a gaseous mixture is the average kinetic energies of the different molecules.

10. According to Graham’s law, the rate of effusion of a gas molecule is related to what? a. The square of the mass of the molecule b. The square root of the mass of the molecule c. Inversely to the square root of the mass of the molecule d. Inversely to the square of the mass of the molecule Answer: c. The rate of effusion is inversely proportional to the square root of the mass of the molecule. If a molecule is 16 times heavier than another molecule it has one-fourth the rate of effusion of the other molecule.

CHAPTER FOUR: THERMOCHEMISTRY This chapter discusses chemistry as it relates to heat. There is an intricate connection between molecules and the temperature of the atoms in the molecules. Concepts of heat capacity and calorimetry are discussed as they relate to a specific molecule. The enthalpy or the total heat in a given system is discussed as it relates to the energy of the system. As in all of physical chemistry, there are specific rules related to the properties of substances at different temperatures, which are covered in this chapter.

ENERGY IN CHEMISTRY AND PHYSICS Before getting into thermochemistry, you need to understand energy in its basic form. Energy can take on many forms and has a great deal of effect on the matter it acts upon. Energy can be seen, as in the different sources of heat energy, or invisible, as in that of electrical energy and microwave energy. In chemistry, we can speak of chemical energy, but there are other forms of energy, such as radiant energy, thermal energy, electrical energy, and nuclear energy. These types of energy forms can be passed from one form to another. Thermal energy is important in chemistry because it results directly from molecular or atomic motion. Radiant energy comes from light sources, radio waves, and microwaves—some of which will become thermal as substances absorb energy from the sun or from microwave exposure. Electrical energy comes from the flow of electrons through particles. Nuclear energy takes the energy stored in the nucleus of an atom, while chemical energy can be “stored” in certain molecular bonds. Each form of energy, except for radiant energy, involves what’s called potential energy or stored energy. Kinetic energy is not stored but comes from the movement of an object. While these forms interconnect, there is the “law of conservation of energy” meaning that the total amount of energy cannot be changed as energy cannot be created or destroyed. Energy is the capacity for work. Mechanical work is the energy necessary to move an object over a specific distance. The formula in physics is this: work is equal to force

multiplied by distance. It can also be described as mass multiplied by acceleration multiplied by distance (as force is mass multiplied by acceleration). Heat, on the other hand, is thermal energy or energy that can go from one object to another. This net transfer of heat only stops when the two objects arrive at the same temperature. Any form of energy can be described along the same lines as kinetic energy because forms of energy can be interconverted. The SI unit of energy is called the joule or J. One joule is a kilogram meter squared per second squared. Most scientists use kilojoules as a joule is such a small form of energy. Your takeaway from this is that all forms of energy can be expressed this way. A common non-SI unit of energy is called the calorie which is the amount of energy needed to raise the temperature of a gram of water from 14.5 degrees Celsius to 15.5 degrees Celsius. The different temperatures are used because the amount of energy required to do this actually varies according to the elevation. Since a calorie is a small quantity, the term kilocalorie is used, which is what is described when looking at the socalled “calories” in food items. Another way to describe a calorie is in terms of the number of joules it involves: 1 calorie is equal to exactly 4.184 Joules.

PROPERTIES OF HEAT IN CHEMISTRY Heat or thermal energy will flow from one place to another in a system. The reaction of chemical substances is referred to as the system, while there can also be heat energy involved with the surroundings, the container the reaction is taking place in, and even the rest of the universe. Heat energy doesn’t always stay within the system but can transfer to the surroundings. You need to understand that there are three different systems seen in chemistry. An open system can exchange matter and energy with the surroundings, such as is seen when water boils in a pan. The reaction of water liquid going to water vapor does not stay in the system but goes into the environment. In a closed system, energy is exchanged between types of matter but not with the surroundings. Isolated systems exchange no matter and no energy with the surroundings. There is practically no system that is truly isolated. The amount of energy lost in a reaction is the same as the amount

of energy (potential or heat energy) gained by another aspect of the reaction (because of the law of conservation of energy). As mentioned earlier in this course, chemical reactions can be exothermic and will give off heat, or endothermic and will store or take on energy. Some chemical reactions can be written in such a way as to account for the heat given off by stating that there is heat as part of the equation. A thermochemical equation is one that has heat put into it or given off as a byproduct of the reaction. An example of an endothermic reaction is that of turning solid H2O to liquid H2O. This reaction requires the input of heat and is said to be endothermic.

LAWS OF THERMODYNAMICS There are three laws of thermodynamics that apply to all physical and chemical systems. The first law has been described, which is that energy cannot be created or destroyed. In chemistry, it also involves the concept that in all systems, the tendency of all chemical (and nonchemical) systems is to move toward the state of the lowest possible energy (which in chemistry, is the most stable system). This means that potential energy is more likely to become kinetic energy. In biochemical systems, catabolic reactions are those that take a bigger molecule (which stores energy in its chemical bonding) to be metabolized into smaller molecules plus CO2 and water. Because these are biological systems, however, the energy is given off as heat to a mild degree but goes into creating high-energy molecules like adenosine triphosphate or ATP, so that the energy is not destroyed or created but is transferred into another form of potential energy, although these molecules, too, will go on to become lower energy molecules. Even the building of molecules like glucose from CO2 and water, which are endothermic, requires the input of energy from the universe in the form of photons of light energy from the sun to make photosynthetic organisms build glucose and other molecules. In such cases, this is an open system because it requires the input of energy (photons) from the universe. In such cases, energy is not created because energy from the universe is necessary.

In an example from the previous chapter, there is pressure-volume work done by an expanding gas. In the equation for this work equals negative P multiplied by the change in volume. What this means is that, an expanding gas at the same pressure uses up energy as the volume increases. A contracting gas holds potential energy, as it will spontaneously expand when the container is breached. Think of the potential energy in a balloon under pressure. According to the first law of thermodynamics, there are two kinds of energy processes, heat and work, that can lead to a change in the internal energy of a system. If energy flows into a system from the outside, both the change in heat and the change in work are positive. If there is work done by the system or heat given off, the change in heat and work are negative. The second law of thermodynamics says that the entropy of a system always increases. Entropy is the measure of disorder of a system. The best way to think of the second law of thermodynamics is to think of small children in a neat and tidy playroom. The natural course of things is that the room will become increasingly messy until energy is put back into the system from an external (or internal source) who can return the room to its former state. Even in this system, the entropy of the universe is always increasing. Figure 18 shows the idea of entropy in a system:

Figure 18.

The third law of thermodynamics states that the entropy of a system will approach a constant value as the temperature approaches absolute zero, such as the entropy is zero at absolute zero (which is -273 degrees Celsius or 0 degrees Kelvin). At absolute zero, the entropy of a pure crystalline substance is zero.

HEAT CAPACITY The heat capacity or thermal capacity is a measurable number that is equal to the ratio of heat added to or removed from a given object to the resulting temperature change. The unit for this is joules per Kelvin (which is the SI unit). The specific heat of an object is the amount of heat necessary to raise an object of one kilogram of mass up 1-degree Kelvin. Heat capacity by itself is considered an “extensive property” of matter because it is directly proportional to the size of the system. If expressing the phenomenon as an “intensive property”, the heat capacity is also divided by the amount of the substance (the mass, number of molecules, or volume). The term “specific” is an intensive property of a substance because it refers to a certain mass of a substance. The specific heat capacity in SI units is Joules per kilogram per degree Kelvin. Water, for example, has a heat capacity of 4.186 joules per gram per degree Celsius. In chemistry, an intensive property of heat is specified relative to a mole of a substance, such as the molar heat capacity, which has the units of joules per mole per degree Celsius. In engineering and other scientific circles, there is the volumetric heat capacity, which is the heat capacity per unit of volume, giving it the units of Joules per cubic meter Kelvin. Calories are sometimes used in place of joules in certain industries. The specific heat capacity in this system is one calorie per kilogram per degree Celsius. There are different ways to measure the heat capacity. It involves adding a known amount of heat to a substance and measuring the change in temperature. This only really works well in the measurement of solids. Gases and liquids are held to a constant pressure or constant volume—most likely constant volume because there is energy necessary in gases to hold them to a constant pressure that interferes with the specific

heat capacity. The specific heat capacity of certain substances will vary according to the temperature. Measuring the specific heat at constant volume can be difficult for liquids and solids. Small changes in temperature will require a very rigid and strong containing vessel because large pressures are required to keep the volume of a solid or liquid

constant

as heat is added to the system. For this reason, it is better to keep the pressure constant, allowing material to expand or contract, in order to measure the heat capacity of the substance. Because the heat capacity is in Joules/kelvin, the heat capacity of a substance at absolute zero would be infinite, which violates the third law of thermodynamics. For this reason, the heat capacity at absolute zero is said to be zero. This involves a simple calculus term in which the heat capacity is calculated as the limit of heat capacity as the change in temperature approaches zero. Molecules have many internal vibrations with potential energy stored in the different degrees of freedom of the molecules in a given sample. The more internal degrees of freedom of a substance, the greater is its specific heat capacity at high enough temperatures to overcome the quantum effects of the molecule, which are greater effects near absolute zero. Each independent degree of freedom allows for greater storage of thermal energy. So, what are degrees of freedom? There is the free 3-dimensional ability of substances to move in space when heat (or any type of energy) is added to the system. This is called translational kinetic energy. There is also rotational kinetic energy of a substance (mainly an atom or molecule). There are only three degrees of translational kinetic energy, corresponding to three dimensions (x, y, and z dimensions). Rotational energy depends on the size and degree of inertia of an atom (which is small for a small single atom but bigger for molecules and larger atoms). Molecules are polyatomic and will have rotational energy. This will allow these types of substances to have a fourth degree of freedom. Finally, there are internal vibrational degrees of freedom of an atom. While there are three degrees of translational freedom, there are also three rotational degrees of freedom, which also involves the three degrees in three-dimensional space.

This isn’t seen in singular molecules and is only seen in two degrees in linear molecules because they rotate only in two degrees. For this reason, a diatomic molecule has only two degrees of freedom in the rotational sense. This gives the diatomic molecule like N2 or diatomic nitrogen a total of five degrees of freedom. The degrees of freedom corresponding to translations and rotations are called the rigid degrees of freedom, since they do not involve any deformation of the molecule. Vibrational degrees of freedom are not rigid with these degrees of freedom depending on the shape and characteristics of the molecule. Vibrational degrees involve deformations and changes between atoms of a molecule. An example is nitrous oxide or N2O, which is linear so it has two rotational degrees of freedom but 8 vibrational degrees of freedom. This leads to a total of 13 degrees of freedom (2 rotational, 3 translational, and 8 vibrational) and a greater heat capacity.

CALORIMETRY Calorimetry means “heat-measurement” in Latin and Greek, which basically means that this is the science of measuring the amount of heat in a substance. It can measure the amount of heat consumed (in an endothermic process), the amount of heat produced (in an exothermic process), or the amount of heat dissipated by a sample. Calorimetry can be used in a vast number of chemical and biological systems as it can be used whenever heat is exchanged, generated, or lost. A calorimeter is a device that measures the heat of a reaction. Reactants can be mixed in a calorimeter, allowed to react, and then there will be a measurement of the temperature difference before and after the reaction. A calorimeter does not have to be fancy or complex. It simply has to contain a reaction and measure temperatures. A calorimeter can be operated under constant pressure (such as atmospheric pressure) or constant volume. It can measure the heat capacity of the calorimeter itself by adding heat to it and measuring the amount of heat that is required to raise the temperature of the entire calorimeter by 1-degree Kelvin. This will be an extremely small number, requiring very sensitive thermometers.

A bomb calorimeter is an explosive-proof steel container called the bomb that does not change in volume during the reaction. It is submerged in liquid (usually water) that absorbs the heat of the reaction taking place. The volume does not change so that the heat evolved under constant volume is the change in temperature multiplied by a constant. When using water in a calorimeter, the change in temperature of the water after a chemical reaction will be detected and will be used to calculate the energy added to or given off by a system. In bomb calorimetry, an object is placed in a chamber along with air or oxygen. An electrical signal is used to ignite the object in order to burn it completely. The energy given off when the item is completely combusted is transferred to the surrounding water, changing its temperature and becoming a measurable amount. Figure 19 shows what a calorimeter looks like:

Figure 19.

What you need to know is that, while the reaction takes place inside the chamber, the measurement of the energy or heat produced or taken up by the reaction is of the surrounding water and not the change in temperature of the chamber. The energy of the chamber is equal to the mass multiplied by the change in temperature multiplied by the specific heat of water: 4.18 joules per gram per degree Celsius.

ENTHALPY AND ENERGY Both heat and work are different representations of energy transfer mechanisms. Work transfers energy such as when moving an object from place to place, while heat can transfer from one place or substance to another. The unit of work is Joules, as you remember, and it doesn’t matter what type of energy is involved. Reactions can occur at a constant pressure and reactions can occur at a constant volume. First consider the situation of a constant pressure reaction. The change in energy of a system involving the reaction, called the delta U, is equal to the flow of heat minus the pressure multiplied with the change in volume. If the reaction happens in a closed vessel, t here is no change in volume so the change in energy must equal the heat flow in the reaction. The problem with this is that most reactions are not carried out in completely sealed “constant volume” situations. They happen in open containers at a constant pressure. In such cases, the initial reaction does consider that a change in volume might occur. This introduces a new concept called enthalpy, defined by the initial H. Enthalpy is similar to energy but is not exactly the same because it takes into account the shrinkage and growth of a reaction. Enthalpy is the internal energy plus the pressure multiplied by the volume. The change in enthalpy of a system accounts for the change in internal energy added to the change in pressure-volume of the system. The enthalpy is the heat flow in constant pressure situations. State functions are those functions of a system that depend only on the state of the situation, such as its pressure, temperature, composition, and amount of substance but not on its past history. Internal energy and enthalpy are also state functions of a substance. This means that the internal energy and enthalpy changes depend only on

the initial state and the final state and not how the substance achieved that change. It could take five steps to get from enthalpy A to an enthalpy B or just one step; the change in enthalpy will be the same. The enthalpy change will be negative for an exothermic reaction, while the enthalpy change will be positive for an endothermic reaction. If the change in enthalpy is negative, the reaction will be considered to go downhill energetically speaking; if the change in enthalpy is positive, the energy will go uphill energetically speaking. With this in mind, you need to know that bond breaking in chemistry ALWAYS involves an input of chemical energy (positive enthalpy), while bond-making ALWAYS involves the release of energy (negative enthalpy). Negative enthalpy equals exothermic, while positive enthalpy equals endothermic. Reversing any chemical reaction will change the sign of the enthalpy. Ice, for example, will absorb heat when it melts and electrostatic interactions are broken, while liquid water must release heat when it freezes. The magnitude of the enthalpy does not change in these types of reactions—only the sign as being plus or minus will change. The magnitude of the enthalpy changes according to the mass of the substances in the reaction. Larger reactions will produce more heat or will take on more heat, depending on the type of reaction. It can be described as a change in heat energy per mole of substance in a specific reaction. In the reaction: 2 moles of Hydrogen gas (H2) plus one mole of Oxygen gas (O2) goes to 2 moles of H2O (water) This is an endothermic reaction, which requires 570 kilojoules of energy. Rather than just use the term kilojoule to determine the enthalpy it is written as 570 kilojoules per mole of Oxygen gas or 285 kilojoules of energy per mole of hydrogen gas, using the stoichiometry of the equation in order to determine a particular enthalpy for an equation (thus standardizing the enthalpy).

RULES OF THERMOCHEMISTRY As you have learned, thermochemistry helps to describe the energy and heat necessary for a chemical reaction or physical change to occur. The absorption or release of energy may result in a phase change (from solid to liquid, from liquid to gas, etc.). Thermochemistry helps to explain whether or not a reaction will go spontaneously. The rules that govern thermochemistry include the following: •

The amount of enthalpy change will vary proportionally with the amount of substrate or product in the reaction.

•

The amount of enthalpy change is the same for opposite reactions so only the sign of the enthalpy amount will change.

•

The magnitude of enthalpy change does not depend on the number of steps required to achieve a given end product as long as the beginning substrate and end product are the same.

There is also the second rule of Lavoisier and Laplace, which states that processes that are endothermic have a positive enthalpy and processes that are exothermic have a negative enthalpy. The third rule or Hess’ law states that enthalpy is a state function and does not depend on the path of equation or equations.

KEY TAKEAWAYS •

Heat is a form of energy, of which the Joule is the standard SI unit.

•

The calorie can be used to describe heat but is a non-SI unit for heat.

•

The calorimeter will measure the heat released or absorbed in a given reaction or in the combustion of a given substance.

•

Enthalpy is a state function of a certain substance, like pressure and number of molecules.

•

The change in enthalpy is positive in an endothermic reaction and negative in an exothermic reaction.

QUIZ 1. Energy that comes from the splitting of certain atoms is referred to as what kind of energy? a. Chemical energy b. Nuclear energy c. Thermal energy d. Electrical energy Answer: b. Nuclear energy comes from the splitting of certain atoms, which is why it is sometimes referred to as atomic energy. 2. Which type of energy is not considered potential energy? a. Radiant energy b. Electrical energy c. Nuclear energy d. Chemical energy Answer: a. Each of these forms of energy is considered potential energy, except for radiant energy, which involves the movement of particles rather than potential energy. 3. Which law of thermodynamics refers to entropy at absolute zero? a. Law of conservation of energy b. First law of thermodynamics c. Second law of thermodynamics d. Third law of thermodynamics Answer: d. The third law of thermodynamics indicates that entropy approaches zero at absolute zero, which is 0 degrees Kelvin or -273 degrees Celsius.

4. The tendency of the universe to become increasingly disordered is called what? a. Entropy b. Enthalpy c. Work d. Heat capacity Answer: a. Entropy is a measure of a system to become increasingly disordered or to reach the lowest energy state, which is the same for all practical purposes, to be disordered. 5. Which measure of heat changes with the amount of the substance? a. Heat capacity b. Volumetric heat capacity c. Specific heat capacity d. Molar heat capacity Answer: a. The heat capacity will change as it depends on the volume or amount of the substance. In other words, the heat capacity of a substance will double if twice the amount of the substance is being looked at. This is not the case with the other measures of heat, which are intrinsic measures of heat. 6. If using calories instead of joules, what is the specific heat capacity of water in calories per kilogram per degree Kelvin? a. 1 b. 4.32 c. 10 d. 532 Answer: a. The specific heat capacity of water in units of calories per kilogram per degree Kelvin,

is 1.

7. What is measured in order to determine the energy given up by a reaction in a bomb calorimeter? a. The change in temperature inside the bomb chamber b. The change in temperature in the surrounding air around the calorimeter c. The change in temperature in the surrounding water around the chamber d. Both the change in temperature of the surrounding water and in the chamber Answer: c. What’s measured is the change in temperature of the surrounding water around the chamber and not the temperature change in the chamber itself. 8. When is enthalpy the same as the heat flow in a system? a. Under all systems or reactions b. Under constant pressure situations c. When there is pressure volume work done in a system d. When heat is not lost by the system to the environment Answer: b. The enthalpy is the same as the energy flow when the reaction is done under constant pressure. 9. In standardizing enthalpy of an equation for one mole of substrate A going to 2 moles of end product B, what are the units to describe this enthalpy change? a. Joules b. Joules per mole of the substrate A c. Joules per mole of the end product B d. Joules per mole of either the substrate or the end product Answer: d. Enthalpy standardized for the stoichiometry of the equation can be represented as joules per mole of either substrate A or end product B.

10. The equation: one mole of substrate A goes to two moles of end product B plus 100 joules of heat can be described as -100 joules per mole of substrate A. What is the description of the reverse reaction? a. It is still -100 joules per mole of substrate A. b. It is -50 joules per mole of substrate B c. It is +50 joules per mole of substrate B d. It is +100 joules per mole of substrate B. Answer: c. Because the reaction is A goes to 2B and is exothermic, the enthalpy is negative. The reverse reaction is endothermic so the enthalpy will be positive and will be half of 100 joules because of the stoichiometry of the reaction.

CHAPTER FIVE: CHEMICAL BONDING AND MOLECULAR STRUCTURE This chapter gets into the subject of chemical bonding and molecular structure. Unless an atom is completely inert, it will have the natural tendency to interact with other atoms. This is certainly the case with ionic substances in chemistry. The chapter will study the important topic of covalence and covalent bonding between molecules and how this should look from an atomic perspective. Because of the different properties of atoms, they exist within molecules in specific three-dimensional arrangements, which is covered in this chapter. The topic of bonding in metallic atoms is covered as part of this chapter as well.

COMPOUNDS AND BONDING Compounds are the structure that an atom is most likely to remain stable in. An example is sodium. Sodium is an atom found everywhere in nature, but unlike noble gases like argon and helium, it does not exist in nature in elemental form. Why is this the case? It is because sodium in nature reacts explosively with water and reacts in oxygen, making it impossible to store the element except in mineral oil or kerosene. Instead, sodium is found in ionic form, referred to as Na+. This is the stable chemical form of sodium in nature. Noble gases are stable, as you recall, because they have what’s referred to as a stable octet, with all of their s and p sub-orbitals filled. These eight valence electrons are all that are necessary to be a stable compound. Elements that gain electrons because they have a so-called need to fill their valence electron orbitals with eight electrons are called anions and are negatively charged in nature. Elements that donate electrons in an ionic bond are positively charged and are called cations. The bond between cations and anions is called an ionic bond. Covalent bonding will be discussed shortly. This happens when there is a relative sharing of valence electrons. Only nonmetals are capable of this type of bonding.

With ionic bonding, there is no such thing as one cation bonding with one anion. Ions will arrange themselves into organized patterns so that negatively-charged ions intermingle with positively-charged ions in a matrix. This organized structure is referred to as a lattice structure. This basically happens only in crystalline or solid form and there is no movement between the atoms of the structure. The ionic bond occurs between metals and nonmetals with the exchange of electrons being a permanent situation. As the ionic compound is dissolved or melted, the particles will separate into ions rather than maintain their “atomic” configuration. The tendency to become an ion is, for many ions, so overwhelming that the element does not exist in any other way. As solids, these are brittle compounds that break apart easily. Figure 20 shows what an ionic bond in lattice form looks like:

Figure 20.

This electrostatic attraction between cations and anions is extremely strong—so strong that a compound like sodium chloride will take 800 degrees Celsius to melt and 1500 degrees Celsius to boil. In solution, ionic compounds will conduct electricity as they will form a negative pole and a positive pole with anions on one side and cations on the other side. Because water is polar, they will be mobile atoms in an aqueous (water) solution. Covalent bonding is another way to achieve the desired octet. A nonmetal or metalloid can form a covalent bond. This cannot happen with metals because they have so few electrons to share in terms of their valency. The term covalent refers to the sharing of “valent” electrons. Covalent compounds have low melting points, being liquids or gaseous at room temperature. They also often do not dissolve in water or, if they do, they cannot conduct electricity. Figure 21 shows several covalently bonded molecules with polar and nonpolar covalent bonding:

Figure 21.

Metallic bonding is a third type of bonding. In this type of bonding, there is a delocalization of the electron that can be tied to more than one nucleus. There is what’s referred to as an “electron sea” with multiple electrons flowing around different nuclei. This is why metals are malleable, conduct electricity, and are lustrous. The process of losing or gaining electrons will completely change the chemical properties of a given atom and element. Sodium becoming positively charged by giving up an electron is no longer volatile and explosive. It will not go back to becoming 83

sodium again once it has gained an electron. Chlorine in gaseous form is poisonous but it is harmless as a chloride anion. Organic molecules are made from carbon, oxygen, and hydrogen (for the most part) and act very differently depending on their arrangement, with molecules taking on their own physical properties and chemical properties. When we talk of compounds, we are talking about substances that cannot normally be separated and that have their own unique melting point, boiling point, and other physical properties. Elements and compounds are substances considered to be pure. These can be differentiated from mixtures, which consist of more than one type of molecule. A mixture will physically combine but will not chemically combine. The law of constant composition is what differentiates a mixture and a compound. According to the law, in a compound, the molar ratio of substances will always stay the same. Examples are CO2 and H2O. These will always stay at the same ratio in any situation. There will always be one carbon atom to two oxygen atoms in CO2 and there will always be two hydrogen atoms to one oxygen atom in H2O. Brass is another example; however, this is a mixture of copper and zinc. The ratio of copper and zinc will change and it will still be called brass. This is not a compound. We’ve been referring to compounds by their molecular formula already and by now, hopefully you have memorized or plan to memorize the different chemical symbols from the periodic table of the elements. When looking at a compound like aluminum oxide, it is not written as AlO, which would not be accurate. This is because the molecule requires two atoms of aluminum and three atoms of oxygen in order to make the molecule. It is then written as Al2O3. Similarly, calcium hydroxide is written as Ca for calcium and OH for hydroxide but there are parentheses around hydroxide and the number 2 after that because it takes two OH subunits to make the molecule that becomes Ca(OH)2.

IONS As mentioned, the ionic bond is the somewhat magnetic bonding between the cation that is positively charged and an anion that is negatively charged. The anion usually is a nonmetal, and the cation is usually a metal, such as sodium and potassium. The changeover of electrons makes an electrostatic attraction between the charged

molecules. Electrostatic attraction is that between a positively charged ion and a negatively charged ion. The attraction does not have to be one-to-one , such as in the case of an atom that needs two electrons, such as sulfur and an atom that has one electron to give. This can lead to two sodium ions connecting to one sulfur ion, to make N2S. There would still be an electrostatic attraction but there would be three atoms attracted instead of two. You should look at the periodic table and memorize which ions are likely to give up electrons to become cations and which ions are more likely to receive electrons to become anions. All of the metals in the first group of the periodic table (lithium, sodium, potassium, and even hydrogen sometimes) will become a +1 cation because they have an electron to spare in their outer s orbital. All metals in the second group of metals (beryllium, magnesium, calcium, etc.) will lose two valence electrons and will become +2 cations. The transition metals will vary and this will be discussed later. The 3A atoms (boron, aluminum, and gallium) will form a +3 cation. The 5A nonmetals like nitrogen and phosphorus will form a -3 anion. The 6A family of nonmetals will form a -2 anion (such as oxygen and sulfur) and the 7A family of nonmetals will form -1 anions (such as fluorine, chlorine, bromine, and iodine). The 8A family involves the noble gases; these are stable and will not form ions. The anions often end in “-ide” when they are in anionic form. In other words, fluorine becomes fluoride and oxygen becomes oxide (and so forth). There are also what’s called “polyatomic ions”. These are different from the monoatomic ions we have already seen and talked about. They are a group of covalently bonded atoms that behave like a single ion. Nearly all of these are anions versus cations. We’ll list these and they should be memorized: •

Ammonium cation—NH4+

•

Hypochlorite anion—ClO-

•

Chlorite anion—ClO2-

•

Chlorate anion—ClO3-

•

Perchlorate anion—ClO485

•

Nitrite anion NO2-

•

Nitrate anion—NO3-

•

Bicarbonate anion—HCO3-

•

Hydroxide anion—OH-

•

Acetate anion—C2H3O2-

•

Permanganate anion—MnO4-

•

Cyanide anion—CN-

•

Thiocyanate anion—SCN-

•

Sulfite anion—SO32-

•

Sulfate anion—SO42-

•

Carbonate anion—CO32-

•

Peroxide anion—O22-

•

Oxalate anion—C2O42-

•

Silicate anion—SiO32-

•

Thiosulfate anion—S2032-

•

Chromate anion—CrO42-

•

Dichromate anion—Cr2O72-

•

Phosphate anion—PO43-

Figure 22 shows these cations and anions in a table:

Figure 22.

The transition metals can form ions with different charges, making them unique in making differently-charged cations. Iron can lose two electrons to make a Fe2+ cation or can lose three electrons to make a Fe3+ cation. The convention is to insert a Roman numeral in order to define the charge on the transition metal cation. Fe2+ would be Fe with a roman numeral two after it, and so forth. These will be the case also with copper, tin, nickel, gold, and lead. In the case of iron oxide for example, there are two forms that are chemically very different from one another. FeO is different from Fe2O3. For this reason, FeO is called iron (II) oxide and Fe2O3 is named iron (III) oxide. When writing an ionic formula, you need to fully balance the molecule so that the charge on the cation or cations must balance the charge on the anion or anions. When you mix calcium and fluoride, for example, you need to know that calcium has a 2+ charge and fluoride has a -1 charge. This means that the combined molecule is CaF2. It is always written with the cation first and the anion second. The Roman numeral is used when writing for transition elements that can have a different charge (which is not the case for all metallic elements). As mentioned, things like oxygen ions become oxide, 87

and chlorine ions become chloride. Use parentheses if necessary to clear up the number of atoms in the molecule.

COVALENCE Covalent bonding, as you know, happens between nonmetals and involves the overlapping of orbitals rather than the complete transfer of electrons and the electrostatic attraction between two ions. The shared electrons will occupy the valence orbitals of both atoms at the same time, creating a strong bond, called the covalent bond. This gives each atom an effective octet. Lattice structures do not occur with covalent bonding. There are many more covalently bonded substances when compared to ionic substances. In H2 for example, there is one electron in each 1s orbital and, when this is shared, they both have an essentially full 1s orbital. There can be polyatomic molecules that have both covalent and ionic bonding. For situations like sodium nitrate or NaNO3, there is a covalent bonding between the nitrogen and oxygen molecules; however, there is a charge left over, which is a -1 charge. This then makes an ionic bond with sodium, which has a +1 charge. The Lewis structure is one way to identify covalent bonding, which is based on the eight valence electrons around most of the atoms that are involved in this type of bonding (except for hydrogen, which follows the duet rule). In drawing a Lewis structure, you need to first look at the number of valence electrons around the atom on the periodic table. You then place these electrons as dots around the larger molecule. When they are bonded with another molecule and sharing part of their electrons in an orbital, you draw a line connecting the two atoms. Figure 23 shows what the Lewis Structure looks like:

Figure 23.

Sometimes, this will lead to one line between the atomic structures with single bonds. Two lines can be drawn with a “double bond” and three lines can be drawn with triple bonds. Figure 24 describes three Lewis structures for water, diatomic oxygen, and diatomic nitrogen:

Figure 24.

This Lewis structure, also called the Lewis dot structure, can be drawn for any type of covalent bond structure. Again, this involves bonding between nonmetals and nonmetals or metalloids only.

MOLECULAR ORBITAL THEORY According to molecular orbital theory, chemical bonding happens when the attractive forces between an electron and two nuclei adjacent to one another overcomes the repulsion of the nuclei with each other. Imagine that there is a region of space between two nuclei that represents the “binding region”. Areas not involved in bonding and where the electron is not located is called the “anti-binding region”. The orbital or orbitals that participate in the binding process become the molecular orbitals of the atom and are the molecular orbitals of the molecule as a whole. The bonding molecular orbital is one that has become involved in the bonding process. As the 1s orbital of the hydrogen atom, for example, becomes close to another 1s orbital of another hydrogen atom, these two molecular orbitals are available for bonding. When these two atoms ultimately share an electron so that there are two electrons in the molecular orbital, it becomes a bonding molecular orbital that is shared between the two atoms. It is also referred to as a sigma orbital. So far, we have referred to an orbital as having a specific shape. So, what is the shape of the orbital when it has combined? It goes back to what we know of orbitals and electrons. When orbitals combine, they can change their shape but, as in any orbital, the electrons must have opposite spin numbers, with one having a +1/2 charge and one having a -1/2 charge. If they are out of phase with each other (both positive or both negative), this is called an antibonding orbital. If the two are in phase with each other, it is called a bonding orbital. Antibonding orbitals will be similar to the canceling out of two wave forms in the ocean water with a nodal surface where the bond cannot occur because the wave does not occur. This is an unfavorable situation. The bonding orbital does not have a nodal surface so the electrons can interact within the space. Remember that a node is a place

where an electron has zero probability of being within the orbital so a bonding orbital must have no nodes. In addition, in dihydrogen, there is a reduction in the potential energy between the two nuclei by 270 kJ/mole or a decrease in the repulsion of the two nuclei by 270 kJ/mole when one electron is between the two nuclei. When two electrons are added into the bonding orbital, one would expect a reduction in the repulsion to be twice that or -540 kJ/mole. This isn’t the case, though, and it’s actually less than this—only about 452 KJ/mole. Why this is the case is because the two electrons repel each other and this repulsion is added to the equation. This is why it is easier than one would think to split dihydrogen or H2 into two atoms of hydrogen ion or H+. Things get more complicated in larger molecules but, suffice it to say that with any molecular orbital, there is a favorable bonding orbital and an unfavorable anti-bonding orbital created. The bond will happen primarily with a bonding molecular orbital but will not happen with regard to the anti-bonding orbital. This predicts whether or not a bond between two atoms happens.

SHAPES OF MOLECULES Lewis structures are helpful in determining the arrangement of electrons and bond pairs in that they predict whether a bond will occur. In reality, however, Lewis structures are by nature two dimensional while molecules themselves are three dimensional. Often, it is the three-dimensional shape of a molecule that best predicts the characteristics of the molecule. We’ve hinted at 3D shapes of molecules when we talked about polarity and the fact that water is a polar molecule because it is not a straight-lined molecule. It is the VSEPR theory or valence shell electron pair revulsion theory that best predicts the shape and geometry of the molecule. The basis of this theory is that, in order to minimize electron-pair repulsion activity, these electron pairs around the central large atom need to get as far apart from one another as possible. Figure 25 describes this when looking at the methane molecule:

Figure 25.

Methane is considered to be tetrahedral with a bond angle of 109.5 degrees between any two hydrogen atoms with the carbon molecule in the middle. Figure 26 depicts the ammonia molecule, which is different and is considered to be trigonal pyramidal:

Figure 26.

In this case, there are four pairs of electrons around the central atom; however, one pair is not bonded with anything. The electrons will be 109.5 degrees apart from one another, giving a different shape to the molecule. With water, this is not a straight molecule either. It has a central atom and two lone pairs of electrons, making a bent molecule.

This makes the five possible shapes of small molecules to be linear, trigonal planar, bent, trigonal pyramidal, and tetrahedral. We’ve talked a little bit about covalent bonds being polar or nonpolar. This can be described by looking at the ability of an atom to attract the shared electrons between the atoms, called “electronegativity”. When two atoms combine, their electronegativities with respect to one another can determine how much each molecule wants to attract that electron pair. If there is equal sharing because the two molecules are the same (as in dihydrogen or dioxygen), the respective electronegativities between the two molecules is the same so this is a nonpolar covalent bond. The extreme of that is the ionic bond. In between the nonpolar covalent bond and the ionic bond is what’s called the polar covalent bond. This is a bond that is covalent but that involves two atoms with different electronegativities. There is complete sharing of the electrons but the attraction is not equal. There will not be a completely +1 charge and -1 charge; however, there will be what’s called a “fraction of a charge” or sigma positive and sigma negative charges. There is a dipole moment between the two molecules (with the arrow pointing to the more electronegative atom or element). The difference in electronegativity must be between 1 and 1.7 in order to have a polar covalent bond. An electronegativity difference greater than this is an ionic bond. Polarity can determine the molecule’s boiling point, melting point, surface tension, and solubility. The Lewis structure will help determine this polarity versus non-polarity. Nonpolar molecules have no unshared pairs of electrons, while polar molecules are asymmetrical, having different electronegativities or lone pairs of electrons. Polar covalent molecules are more likely to have the following: •

Higher melting points than nonpolar molecules

•

Higher boiling points than nonpolar molecules

•

Greater water solubility than nonpolar molecules

•

Lower vapor pressures than nonpolar molecules

HYDROGEN BONDING Whenever a hydrogen atom is bonded to a highly electronegative atom (like nitrogen, oxygen, or fluorine), this will form a highly polar bond. The electronegative atom takes on a negative partial charge and the hydrogen atom takes on a partial positive charge. This is what happens between polar covalent molecules but happens to a greater degree with hydrogen bonding. Hydrogen bonding is found extensively with water . The hydrogen atoms bind to the oxygen atom in a polar covalent bond; however, the hydrogen atoms between two different molecules in a solution of water are also involved in hydrogen bonding. This is partially why water has the polar properties it has. This bonding affects the crystalline structure of ice, which creates a hexagonal lattice shape. Figure 27 shows the basic structure of water in its various states:

Figure 27.

The hexagonal structure of ice makes ice less dense than liquid water so that water as a solid actually floats on the surface of water instead of sinking, which is the case with nonpolar molecules that form a solid shape at lower temperatures. It is also the

hydrogen bonding that contributes to the high boiling point of water compared to similar molecules that are not involved with this type of bonding. Hydrogen bonding is involved with protein and DNA molecules in living things. It is what gives DNA its double helical shape. This type of bonding is what makes the characteristics of polymers like plastic and nylon molecules. Hydrogen bonding increases water solubility, increases the boiling point, and increases the melting point. It accounts for the high surface tension of water.

BONDING IN METALS Metallic bonding has briefly been discussed. It involves the sharing of multiple valence electrons between many metallic (so called positively-charged) ions. The electrons form the glue that connects the different metal atoms together and gives metals their structure. Metallic bonding is different from all other types of bonding already discussed. It is based on the fact that metals have a low ionization energy and is why these have high melting and boiling points. The freedom of the electrons in this “soup” means that electrons can have energy pass through them in the form of electric currents. This makes just metals, graphite, and ionic compounds the only things that can conduct electricity. Metals can also conduct heat faster than other molecules. Metals will have a minimum of one valence electron that isn’t shared per se. Instead, the outer atomic orbitals will overlap, leading to something similar to covalent bonding. One exception to the typical metallic bonding seen in metals is that of mercury or Hg2+ molecules that form covalent bonds with other mercury ions. This accounts for the different properties of mercury when compared to other metals. Alloys represent solutions of different types of metals.

KEY TAKEAWAYS •

The focus of this chapter was on molecular structure and molecular bonding types.

•

There can be ionic bonding, covalent bonding, metallic bonding, and hydrogen bonding.

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In ions, electrons are “given away” to other atoms and the transfer is essentially irreversible because of the stability of the ion when compared to the atom that is not in ionic form.

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Covalent bonding can be polar or nonpolar, depending on the relative electronegativity of the atoms involved in the bond.

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Hydrogen bonding involves a partial connection between hydrogen and electronegative molecules.

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Metallic bonding involves bonding between metals that have shared and overlapping orbitals.

QUIZ 1. Cations and anions can bind together in what are called what kinds of bonds? a. Covalent bonds b. Metallic bonds c. Hydrogen bonds d. Ionic bonds Answer: d. The type of bonding that goes on with cations and anions is called ionic bonding. 2. What type of bonding only occurs between nonmetals that share electrons? a. Covalent bonds b. Metallic bonds c. Hydrogen bonds d. Ionic bonds Answer: a. Covalent bonding only occurs between nonmetals and other nonmetals. There is a relative sharing of electrons between the atoms in this type of bonding. 3. Which molecule is considered covalently bonded? a. CH4 b. NaCl c. KPO4 d. KOH Answer: a. Each of these is considered an ionically-bonded molecule except for CH4, which is methane gas. Methane gas is covalently bonded.

4. Which type of bonding is considered an electrostatic attraction? a. Covalent b. Nonionic c. Ionic d. Metallic Answer: c. Ionic bonding is the electrostatic attraction between a positively-charged cation and a negatively-charged anion. 5. What is the charge on the hypochlorite ion? a. +1 b. -1 c. -2 d. -3 Answer: b. Hypochlorite is a -1 anion, as is chlorite, chlorate, and perchlorate. 6. Which anion has a -3 charge? a. Hydroxide b. Sulfate c. Peroxide d. Phosphate Answer: d. Phosphate is the only polyatomic anion that has a -3 charge. 7. What is the molecular formula for aluminum sulfide? a. AlS b. Al2S3 c. AlS2 d. AlS3 Answer: b. Al2S3 is the right answer because Aluminum has a +3 charge and sulfide has a -2 charge. In order to balance this, there needs to be two aluminum atoms for every three sulfide atoms.

8. Lead IV has is an ion that can combine with oxygen to make lead oxide. What is the molecular formula for this? a. Pb(IV)O b. Pb2O2 c. PbO4 d. PbO2 Answer: d. Lead IV has a 4+ charge that will ultimately fit well with oxygen, which has a -2 charge. This leads to PbO2 or Pb(IV)O2 as the molecular formula. 9. Sodium nitrate has what type of bonding? a. Ionic bonding b. Metallic bonding c. Ionic and covalent bonding d. Covalent bonding Answer: c. There is combined ionic and covalent bonding, which involves ionic bonding between sodium and nitrate as well as covalent bonding between the nitrogen and oxygen molecules in the nitrate ion. 10. Which atom follows the duet rule rather than the octet rule? a. Hydrogen b. Oxygen c. Carbon d. Nitrogen Answer: a. Hydrogen follows the duet rule because it is the only atom that does not involve a p orbital so it only has a 1s orbital with two electrons desired by the orbital.

CHAPTER SIX: CHEMICAL FORMULAS AND REACTIONS This chapter focuses on the inevitability of chemical reactions and on what chemical formulas look like. Molecules and atomic substances tend not to stay separate from one another when put together in the same physical space and reactions will follow a specific pattern. The rates of different reactions and why they occur or do not occur are discussed as well as the issue of the energy of activation necessary to take a group of molecules in one form and turn them into molecules of a different form. Exactly what happens in chemical reactions is covered as part of this chapter.

HOW REACTIONS OCCUR Chemical reactions happen whenever substrates or reactants are given the energy and opportunity to react within the same physical space. According to the collision theory, there needs to be collisions between atoms and molecules to such a degree that the kinetic energy is given to break chemical bonds and reform new ones. Besides needing to be in the same physical vicinity, the molecules must be moving fast enough to collide with sufficient energy. This is why heat is one of the more common ways that reactions can be coaxed into happening. Heat equates with kinetic energy when it comes to molecular interactions and chemical reactions. In addition, molecules need to collide in the right orientation for a reaction to occur. Molecules are not simply amorphous blobs that mix together. The collision needs to happen with the molecules lined up with one another for the breakage of bonds and the remaking of other bonds. There is a “reactive site” on a molecule that needs to be acted on in order for the molecular bond to be broken or changed in some way. So, in a reaction AB + C goes to AC + B, there must be a reaction site or reactive site on AB that interacts with C in a specific way in order for the combination of CAB to break the bond and become AC instead (because AC is inherently more stable than AB under the reaction circumstances. This requires energy, which can be kinetic energy or heat 100

but can be the difference in energy between AB and AC that drives the reaction in that direction. The previously mentioned reaction does not usually happen as simply as that. While it can proceed with one step, many reactions are multi-step processes in which there are intermediates. What you have already learned, however, when discussing enthalpy of reactions, is that it doesn’t matter how many intermediate steps there are in a reaction. There is a beginning enthalpy and an ending enthalpy, with the change in enthalpy that is independent of the intermediates. If a given reaction AB + C goes to AC + B is considered exothermic, heat is given off and released as part of the reaction. The reactants exist at higher energy states than is true of the products, with no choice but to have heat released as a result. There may be an activation energy, which will be discussed further in a few minutes, which is the energy necessary to get even an exothermic reaction going. Figure 28 describes an exothermic reaction that includes an activation energy:

Figure 28.

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In the figure, you’ll see that EA is the activation energy—necessary to start the reaction process, while “delta H” is the change in enthalpy in the exothermic reaction. There is a period of time called the transition state, at the top of the activation energy hill, in which there may be intermediates in the reaction process. If the reaction is endothermic, heat will be absorbed and the delta H will be positive. There will still be an activation energy to get the reaction going but the energy of the products is higher than the energy of the reactants. The absorbed energy is stored in the bonds made by the products as potential energy (that could be reversed as the reaction goes in the reverse). In both cases, the activation energy is gotten back as the reactant energy goes to the product energy.

TYPES OF CHEMICAL REACTIONS There are several general kinds of chemical reactions that can occur, which are different in substantive form: •

Combination reactions—this is when two or more reactions form a single product. An example of this is the reaction where 2 sodium (Na) molecules plus chlorine gas (Cl2) react to make sodium chloride (NaCl). Even the burning of coal or solid carbon (C) plus oxygen (O2) goes to make CO2 or carbon dioxide. Technically, there can be more than one product as long as things are combined in the reaction.

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Decomposition reactions—these types of reactions are the opposite of combination reactions. In this type of reaction, a single compound breaks down into two or more simpler substances. An example of this is the breakdown of water into hydrogen gas and oxygen. The breakdown of hydrogen peroxide into oxygen gas and water is a decomposition reaction.

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Single displacement reactions—in this type of reaction, a more active element kicks out a less active element from a compound. An example reaction is Zinc solid plus Copper sulfate goes to copper solid plus zinc sulfate. In this reaction, zinc displaces copper in its reaction with the sulfate ion. Zinc is more reactive than copper and the reverse reaction is not likely to occur. 102

•

Double displacement reactions—these are also referred to as metathesis reactions; in them, two types of ions are displaced. These usually happen in solution with an insoluble precipitate formed (in precipitation reactions) or water formed (in neutralization reactions). The precipitate will leave the solution, which will drive the solution forward. In a neutralization reaction, there is double displacement between and acid and a base to yield salt plus water.

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Combustion reactions—this is basically the burning of a substance or the reaction of a carbonaceous substance with oxygen.

Heat is the main product of this,

along with carbon dioxide and water. The combustion of hydrocarbons, such as propane fuel yields CO2 and heat. •

Redox reactions—this is an overarching term that involves the exchange of electrons from one substance to another. We will talk more about redox reactions later.

WRITING REACTIONS In writing reactions, there are specific rules to follow in order to write an accurate assessment of what goes on in a reaction. The left side of any chemical equation, is where the reactants are listed. These are written in what’s called the stoichiometric coefficients. Basically, this means that the reaction is written so that it becomes clear how many moles of one substance mix with how many moles of another substance. In other words, the reaction proceeds as if molecules react with molecules rather than x grams of a substance mixes with y grams of another substance. In addition, after the molecular formula for a reactant or end product, you must put the form of the reactant and end product: You write (l) for liquid, (s) for solid, and (g) for gas. Many reactions will say (aq), which means “aqueous” in those reactions that involve a chemical in solution with water. The number of a given substance or element must be the same on the left-hand side of the equation as is the case on the right-hand side of the equation.

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An example of how to balance an equation is shown in Figure 29:

Figure 29.

In the figure, you can see that the equations are completely balanced with respect to the numbers of elements or atoms on both sides of the arrow. You need to know how to balance these types of equations as they lie at the root of chemistry and how you understand chemical reactions. So, in the first case in the figure, you have hydrogen gas and nitrogen gas going to methane. Because of the fact that methane is NH3, you need to have at least three hydrogen atoms on the left side. Because you can’t have half of a hydrogen gas, you must put 3H2 gas molecules on the left, leading to the mixture with N2 gas that must yield 2 molecules of NH3 (or methane). In real cases of chemical reactions, there will be the possibility of a limiting reactant that will affect the reaction. This will not affect the equation—only the actual number of moles that can be gotten from the reaction. This is also sometimes referred to as the limiting reagent. In order to have knowledge of the limiting reagent, you need to transfer the number of grams of a substance in an equation and turn it into the number of moles. In the reaction of H2 gas and O2 gas, you can balance the equation that it takes 2 moles of H2 gas and one mole of O2 gas to make 2 moles of H2O. If you have 4 grams of H2 (weighing 2.016 grams per mole, you get 1.98 moles of hydrogen gas). If you do the reaction stoichiometrically-speaking, you need 0.99 moles of oxygen. If you have 20 104

grams of oxygen, this amounts to 0.625 moles (less than the 0.99 moles necessary). This means that oxygen is the limiting agent. You also need to understand the idea of theoretical yield. Theoretically, there will be a total reaction in which all of the reactants go to all of the products. This is almost never the case. This leads to the calculation of the percent yield, which is the actual yield divided by the theoretical yield multiplied by 100 percent. It will always be a number less than 100 percent. When reactions happen in a solution, the reaction is often thought of as the number of moles of a substance divided by the volume of the solution in liters. This refers to as the Molarity, which is the number of moles of the solute divided by the volume of the solution in liters. So, if there is one mole of a solute in a liter of water, the molarity is one. To do a calculation, imagine putting 100 grams of table salt (NaCl) in 50 milliliters of water. The number of moles (with 22.99 grams per mole plus 35.45 grams per mole of sodium and chloride, respectively), you get a total of 1.7 moles of sodium chloride. Divide this by 50 milliliters of water or 0.05 liters, you get a molarity of 34.2 moles per liter. There can be a physical change in the properties of a reactant during the chemical reaction. While there will be a physical change, it does not change the actual substance itself—only its form. There may be the transformation of a liquid to a gas or an aqueous solution of something into a precipitated solid. Again, this is a physical change but not a chemical change. When a precipitation reaction occurs, the solid is called a precipitate. The remaining liquid is referred to as the “supernate”. In such a reaction, the equation is written similar to Figure 30 as shown in your manual:

Figure 30.

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In the example, you must make it clear by adding the nature of the different substances in the reaction. In other words, if it is a solid, you need to put an “s” after the solid and an “aq” after substances that are in solution, in this case, aqueous refers to a solution in water specifically. In a neutralization reaction or acid-base reaction, an acid and base are mixed together. The acid, by definition, produces H+ or hydrogen ions. The base or alkaline substance produces OH- or hydroxide ions. These naturally produce water as an end product. The acid and base substances themselves will produce a type of salt. For example, sodium hydroxide (NaOH) and hydrochloric acid (HCl) make water and sodium chloride (NaCl). Water is also a natural end product. In a combustion reaction, a fuel reacts with an oxidizing agent (usually oxygen), with heat energy released. These are, by nature, exothermic. The energetic substance used in these types of reactions is usually a hydrocarbon, which involves carbon, oxygen, and hydrogen mainly. The end product is going to be CO2 and water.

LIMITING REAGENTS As mentioned, the limiting reagent or limiting reactant is the reactant that gets completely used up in a reaction, therefore determining when the reaction stops. It can be determined directly from knowing the number of grams of a reactant and knowing the molecular weight of the reactant. In any given reaction, there will often be one reactant that will have fewer “moles” than the other, making it the limiting reactant. The other will be the in-excess reactant. There are two approaches to determining the limiting reagent and amount of end product made, including the following: •

Approach 1: Find the balanced chemical equation and convert all the grams to moles in the reaction. Find out which reactant has fewer moles. Use this to determine the number of moles of product that can be made by the limiting reactant moles. Reconvert the end product back into the number of grams of end product that can be made. You can also calculate the grams of the in-excess reactant this way. 106

•

Approach 2 is to balance the equation and convert the information into moles. Use stoichiometry to determine the number of moles and mass of the product made. The limiting reactant is the one that produces the fewer number of moles of the end product, while the in-excess product is the one that produces more moles of the end product.

An example is Na2O2 reacting with water. If 78 grams of Na2O2 is mixed with 29.4 grams of H20, which is the limiting reagent? By adding the atomic weights of two sodium atom and two oxygen atoms, you get a molar weight of 77.96 grams per mole, which leads to 1.001 moles of Na2O2. Doing the same thing with oxygen, you have a molecular weight of 18 grams per mole, leading to 1.633 moles of water. The reaction is a 1:1 ratio of the two reagents, which leads to the limiting reagent being the Na2O2 by about 0.6 moles. This would lead to about 0.6 moles of water left over, which could be calculated in grams by multiplying 0.6 moles by 18 grams per mole or about 10.8 grams of water left.

RATES OF REACTIONS Chemical reactions need different lengths of treatment. The study of reaction rates is called “chemical kinetics”. There are five different things that determine the rate at which a reaction occurs: •

The concentration of the reactants in the reaction. Higher concentrations lead to a faster rate of reaction, particularly in an aqueous environment. If you think of the collision theory, it makes sense that higher concentrations mean more collisions between molecules.

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Temperature. This makes sense, too, as it increases the kinetic energy of the molecules, driving the reaction forward.

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Physical characteristics of the reactants. If a solid is in powder form, this will increase the rate of reaction versus a solid in block form, which will have a lesser surface area to volume ratio.

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•

The presence of a catalyst or inhibitor. Catalysts, like enzymes in biochemical reactions, will increase the rate of a reaction, while an inhibitor will slow down the rate of a reaction.

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Light. This is true of certain reactions, such as those in photosynthesis, that will require light energy to drive the reaction.

With gases, it is also the pressure which drives the reaction. In higher pressure environments in which two gaseous reactants combine or interact, the higher pressure means that there is a greater “concentration” of reactants and a greater chance that the gas molecules will collide with one another. You can define the rate of a reaction as the change in concentration over time. It is expressed as a change in reaction or product divided by the change in time. The expressions for reactants are given as a negative sign, while the rate for the products is a positive sign. The stoichiometry of the reaction needs to be taken into account when describing a rate reaction. So, if there is a reaction in which X + 2Y = Z, the rate of the disappearance of X is half of the rate of the disappearance of Y (and both of these rates will be negative).

ENERGY OF ACTIVATION The activation energy has already been discussed. It applies to chemical reactions and to nuclear reactions in physics. It is measured in joules or kilojoules per mole or kilocalories per mole (with the standard for chemical reactions being kilojoules per mole. It represents the difference between the translational energy and the starting energy of the equation, with the translational energy being the energy of the intermediates that then go on to achieve the end product energy. The rate of reaction will generally increase with increasing temperature but this is not universally true. In certain reactions such as gaseous reactions without a border, increasing temperature will pull the molecules away from each other so that the rate of reaction will not increase or may even decrease.

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A catalyst is something that will change the rate of the reaction by decreasing the activation energy. It will not change the energy of the reactants or the energy of the end products but will only change the activation energy. Figure 31 describes what this looks like:

Figure 31.

What a catalyst does is modify the transition state energy so that the overall activation energy is lowered. This is what’s seen in enzymes that make it more likely to have the transition state occur. Catalysts will not be consumed as part of the reaction so that they will continue to drive the reaction forward. Because a catalyst does not get consumed, small amounts of any given catalyst are necessary in order to drive the chemical reaction.

CHEMICAL EQUILIBRIUM So far, we’ve talked about reactions as though they go from one point to another point but do not go back and do not stop until the limiting reactant has been used up. The concept of chemical equilibrium involves the idea that, at the time of the end of the reaction, there will be equality in the rate of the forward reaction and in the rate of the

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reverse direction. So, when describing a reaction, the convention of using a forward arrow to represent the reaction is instead a forward and reverse arrow, which takes into account that the equilibrium state has been achieved. It can only be achieved in a closed system. Catalysts will decrease the activation energy but will have no effect on the equilibrium point. According to Le Chatelier’s principle or the Equilibrium law, a system will have a disturbance in the equilibrium (such as a change in the concentration of reactants or products, temperature, pressure, or volume) and will have a new equilibrium state. There will often be an observable property or measured physical property in the reaction when it has achieved equilibrium. This can be a change in density, color, pressure, or concentration of substances in the system. As mentioned, the system must be closed in order to have equilibrium. This is because, if one or more of the products (or reactants) escape, there will be a shift of the reaction so that one part of the equation does not participate in the reaction, driving the reaction in one direction. In equilibrium, the process of “reacting” does not cease; rather, there is equivalency of the driving forces of both reactions (forward and reverse). The chemical reaction in which xA + yB goes to zC, where x, y, and z are the stoichiometric constants in the reaction and A, B, and C, refer to the concentration of the different reactants and products, the equilibrium constant can be referred to using Figure 32:

Figure 32.

Note that the coefficients of products and reactants are powers of the concentration values. While the concentrations of the products and reactants may change, the ratio of these concentrations will be the same so that the value of K will remain constant. This is why it is called the “equilibrium constant”. If the value of K is greater than 1, the predominance of products will occur; when K is equal to one, neither side is favored; 110

when K is less than 1, there will be a predominance of reactants. If the K is zero, this implies that no reaction whatsoever is occurring. There are two types of equilibrium reactions. The first is a homogeneous reaction. This is one in which the states of matter are all the same. It can be liquid, gaseous, or solid but all reactants and products are in the same state. An example of this would be nitrogen gas plus hydrogen gas going to ammonia or NH3. This will ultimately reach a state of equilibrium. The second is a heterogeneous reaction, in which the states of matter are different from one another. This can happen in an aqueous solution when a precipitate happens or even when a solid is dissolved into aqueous ions. One example, is the decomposition of baking soda or sodium hydrogen bicarbonate (NaHCO3) into sodium carbonate solid, water vapor, and CO2 gas. So, how do you get to the equilibrium constant? It is not a constant that is a given number that is the same for all reactions. It must be calculated by measuring the concentrations of products and reactants in a system and using the equation in figure 32. The equilibrium constant is called a constant because it determines the rate of a reaction in an ideal reaction that involves a set of reactants and a set of products. When the reaction is gaseous, the concentration is not used; instead, the partial pressures of the different gaseous entities is used in the equilibrium constant equation. The equilibrium constant works well for gaseous homogeneous reactions and for solutions in water. It is different in heterogeneous reactions. In these, solids and pure liquids are excluded from the equilibrium constant. The activities of solids and liquids are considered to be one for mathematical purposes so they do not affect the overall K value. Solvents also have a mathematical equivalent of one. So, when doing the calculation of K, there is nothing added to the equation by solvents, solids, and liquids. If there is more than one step in the reaction, each will have its own K value. For mathematical and practical purposes, when there are multiple steps, the K values are not added together but are multiplied together.

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KEY TAKEAWAYS •

Things like heat, concentration, surface area, and catalysts will increase the likelihood of a reaction occurring.

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The activation energy does not change the overall enthalpy of a reaction.

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The limiting reagent can be determined by knowing the stoichiometry of the equation and the amount of starting and ending products in the reaction.

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There are several types of reactions that differ in the types of beginning reactants and end products.

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Reactions are in equilibrium when the rate of reaction in the forward direction is the same as the rate of reaction in the reverse direction.

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QUIZ 1. Why does heat help aid in the occurrence of a chemical reaction? a. It cause molecules to line up in the right orientation. b. It weakens chemical bonds. c. It allows for more mixing of reactants. d. It increases the kinetic energy of the reactants Answer: d. Heat will increase the kinetic energy of the reactants, increasing the chance of collisions between the molecules in the reaction. 2. What is true of the activation energy in an exothermic reaction? a. It will be a negative number in this type of reaction b. It is independent of the fact that the reaction is exothermic c. It will be zero in an exothermic reaction d. It will contribute to the change in enthalpy of the reaction if positive Answer: b. The activation energy in an exothermic reaction is independent of the fact that the reaction is exothermic. 3. In what type of reaction is heat the primary end product? a. Combustion reaction b. Precipitation reaction c. Decomposition reaction d. Neutralization reaction Answer: a. A combustion reaction is basically a burning of a combustible product with heat as the major end-product.

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4. In what type of reaction is water the primary end product? a. Combustion reaction b. Precipitation reaction c. Decomposition reaction d. Neutralization reaction Answer: d. In a neutralization reaction, water and salt are the end products of the neutralization of acids and bases. 5. In this reaction: 2 HNO3 (aq) + Sr(OH)2 (aq) → Sr(NO3)2 (aq) +2 H2O (l), what is the acid in this equation? a. Sr(OH)2 b. HNO3 c. Sr(NO3)2 d. H2O Answer: b. HNO3 is nitric acid, which is the acid that goes into this neutralization reaction, leading to a salt and water. In this case, the base is strontium hydroxide. 6. In the reaction, two moles of nitric acid plus one mole of Strontium hydroxide goes to one mole of strontium nitrate, a total of 0.85 moles of the salt are created. What is the percent yield of product with respect to strontium hydroxide? a. 0.85 percent b. -0.85 c. 85 percent d. 1.85 Answer: c. The percent yield will be 0.85 moles divided by an expected yield or theoretical yield of 1 mole to get 85 percent yield.

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7. What is most true of a catalyst in a chemical reaction? a. It will increase the likelihood of an exothermic reaction. b. It will be consumed as part of the reaction process. c. It will change the overall enthalpy of the reaction. d. It will change the activation energy of the reaction. Answer: d. A catalyst will not change the overall enthalpy of the reaction and does not get consumed as part of the reaction. It will change the characteristics of the transition phase in the reaction so that it will decrease the activation energy. 8. What is the role of an enzyme in biochemical reactions? a. They change the stoichiometry of the reaction. b. They will catalyze certain biochemical reactions. c. They change the energy state of the beginning substrates, favoring the reaction. d. The can inhibit or favor a biochemical reaction. Answer: b. The role of the enzyme in biochemical reactions is to act as a catalyst in biochemical reactions. They will not inhibit a biochemical reaction nor will they change the stoichiometry or the energy states of the reactants and products. 9. Which law of thermodynamics refers directly to an increase in the energy of a given system? a. Law of conservation of energy b. First law of thermodynamics c. Second law of thermodynamics d. Third law of thermodynamics Answer: c. The second law of thermodynamics states that entropy in a system or in the universe, unless acted on by a force, is always increasing.

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10. What is the amount of heat in a system that is necessary to raise an object of one kilogram of mass up 1-degree Kelvin? a. Heat capacity b. Specific heat c. Calorie d. Thermal capacity Answer: b. The specific heat of a substance such as water (or any other substance) is the amount of heat necessary to raise a kilogram of the substance by one-degree Kelvin.

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CHAPTER SEVEN: SOLIDS AND LIQUIDS The focus of this chapter is the behavior of solids and liquids. Molecules can exist in crystalline form, depending on the physical circumstances. There are certain molecular forces in play when substances are in crystalline form, which are discussed in this chapter. In the same way, liquids have certain properties unique to that state that depend on the atomic nature of the molecule and on the temperature of the liquid. The different forces in play when substances are in liquid form are also covered in the chapter.

PHASES OF MATTER So far, we’ve studied gases and their properties. Gases, like liquids and solids (or crystals) are considered “phases” that an element or molecule can be in. Water is a common molecule that can occur in crystal or solid form, liquid form, and gaseous form. Solid water is, of course, ice, while gaseous water is referred to as water vapor. At particular combinations or pressure and temperature, one will see one of the three main phases of water (or any molecule) being thermodynamically stable. Water and other substances will have a clearly defined melting point and boiling point. At temperatures above absolute zero, all atoms and molecules have kinetic energy that keeps them in constant states of motion. As we have studied, the different states of motion or degrees of freedom are translational, vibrational, and rotational states of motion. The greater the temperature of the system, the more numbers of states of motion that will be active with the energy dispersed among these states. When matter is solid, the molecules are in fixed locations so that the kinds of motion allowed are fewer than in liquids and gases. As temperature rises, technically anything can be gaseous, allowing for greater numbers of quantum states. Changes in state will always change the enthalpy and the internal energy of a given molecule. Regardless of the state that a substance is in, the idea is that the molecules will try to spread themselves out as much as they can. What this means is that molecules

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have free energy, which is the escaping tendency of the molecules from the phase they’re in. Let’s go back to gases for a minute. Gases will have an escaping tendency, which is basically the pressure of the gas and its tendency to want to leave the container it’s in. The greater the pressure, the more frequently the molecules will collide with the container so that they could get outside of the container. With liquids and even with solids, there is not complete confinement to the state the molecules are in. While the thermal energy (kinetic energy) is less, there is still the chance that a molecule near the surface of the liquid or solid will escape into the gaseous phase. This tendency can be measured by putting the solid or liquid into an evacuated container and measuring the gradual increase in gas pressure in the container. This does not occur immediately but will eventually reach an equilibrium inside the vacuum. The pressure at equilibrium in such a condition is referred to as the “vapor pressure”. If the container were to be left open in the situation with water, for example, the vapor pressure will not be reached and the liquid would just evaporate as gaseous water, would leave the liquid and would be lost to the atmosphere. In addition, the vapor pressure will change in a nonlinear fashion with temperature. The vapor pressure relationship of several different liquids is shown

Figure 33.

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in Figure 33:

The normal boiling point of a liquid is the temperature at which the liquid is in equilibrium with its vapor at a partial pressure at one atmosphere or 760 torr. The curve of water in the figure also explains why water boils at a lower temperature in higher elevations because the temperature-vapor pressure curve will be along the same lines as the curve indicates. As you can see, the vapor pressure varies greatly over the range of temperatures, particularly when it comes to water. So, the definition of normal boiling point is the temperature at which the vapor pressure is the same as that of the atmosphere. This vapor pressure curve ends at what’s called the “critical point”. This is shown in figure 34, which is a phase diagram of water:

Figure 34.

At the critical point, the particles of a substance in a closed container are vaporizing at such a rapid rate that the density of the vapor and liquid are equal. This forms a “supercritical fluid” with an absence of the surface tension. This is a critical pressure and critical temperature, so that the condensation of a gas can never occur above the critical point. The gas will become very dense but will not form a liquid because of the high temperature and pressure.

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Relative humidity involves the partial pressure of water in air at a certain temperature. In fact, the ratio of the partial pressure in air to its vapor pressure (equilibrium) at any given temperature is referred to as its “relative humidity”. Air is not denser as it is felt by humans but is, in fact, less dense. It is uncomfortable because it inhibits the vaporization of water through sweating. The relative humidity in low temperatures is higher but not when this same air is brought into warmer temperatures of buildings, leading to a decreased relative humidity inside buildings. The dew point is the temperature at which the relative humidity is 100 percent. It is the temperature at which the partial pressure of water is equal to its partial pressure at a given higher temperature and relative humidity. The temperature of the air must be at or below the dew point for water to leave the atmosphere as rain. Many solid salts have water molecules within their crystal lattices, with compounds known as hydrates resulting. These solid hydrates will have a vapor pressure that correlates with an equilibrium equation in which the hydrate salt goes to the anhydrous salt and water vapor. Figure 35 shows this equation:

Figure 35.

In the figure, the reaction takes the hydrated salt of strontium chloride hexahydrate and allows it to go to the anhydrous salt. The vapor pressure of this hydrate is 8.5 torr at 25 degrees Celsius. This is the pressure at which the two solids can coexist. If the partial pressure of water is greater than 8.4 torr at this temperature, the sample will absorb moisture from the air and will turn it into a hydrate. When fully hydrated, water will make up 40 percent of the mass of the hydrated salt. In a hydrated salt situation, if the partial pressure of water in the air is less than the vapor pressure of the hydrate, the hydrated salt will lose moisture, forming its anhydrous form. This will cause the crystal to break up its crystalline form into a powder on the surface of the crystal—a process that is known as efflorescence.

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WATER CONDENSATION, BOILING, AND EVAPORATION Condensation is the transformation of the vapor form to the liquid form at a temperature that, in water in nature, occurs below the boiling temperature. This is why dew forms at lower temperatures. The boiling point can be described as the “normal boiling point”, which is the transfer of water as liquid to H2O gas at a temperature of 100 degrees and at 1 atmosphere or 760 torr. This “normal boiling point” is not the boiling point at lower pressures and at higher pressures. As mentioned, the boiling point of water occurs at a lower pressure, such as at higher altitudes. Steam droplets will form above a pot of boiling water because the temperature will be lower above the pot of water so that some of the gas turns back into droplets that are seen as steam. What you need to know, though, is that, at the boiling point, vapor will not usually condense into a liquid. In addition, the liquid will not boil at its boiling point but needs a temperature above that in order to boil. We’ll talk more about that later. Finally, you need to know that the dew point is equivalent to the condensation point and both are equivalent to the boiling point at 100 degrees Celsius at 760 torr. If water is boiling and becomes steam, if the droplets are very small and the partial pressure of water is less than the vapor pressure, the tendency will be for the droplets to evaporate. When in droplet form, the surface area to volume ratio is reduced and the curvature means that each molecule faces fewer attraction forces from its nearest neighbor. This also supports evaporation. In the situation of a bubble inside a liquid such as water, the tendency is for the bubble to collapse. This is a situation where a bubble, a hole in a liquid, has molecules just on the outer surface (in the water) that curves inward, making an increase in nearestneighbor attractions. The temperature must rise within the liquid above the boiling point in order for the bubble to expand and not to collapse. There is also the hydrostatic pressure of the weight of the water above the bubble that has formed. This too needs to be overcome in order to have the bubble expand and not collapse. Water must be superheated in order for it to actually “boil” and have expansion

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of the bubbles forming at the bottom of a pot of water. Superheating involves the heating of water above its boiling point (under higher pressure circumstances). So, if tiny drops have forces that tend to evaporate them, how do they condense and form rain, for example? In the example of a vapor cooled in a container, the condensation happens on the inner surface of the container because it is in that location that there are greater intermolecular attractions between water molecules that will stabilize droplets. This also explains why condensate appears on the outside of an icy glass or on grass. Droplets of steam above a pot of boiling water

happen because of

dust particles in the air. Precipitation in the environment happens because of particles in the air that “attract” water in liquid form. This can be dust, volcanic ash, meteoric dust, and smoke. This will stabilize the small droplets until they grow to a size that causes rain, snow, or fog to develop. These droplets are microscopic but have a macroscopic impact on the development of rain and other precipitation. Fog itself can be irritating—not because of the moisture—but because of the dust particles that underlie the moisture. What is the difference between the boiling point of a liquid and the evaporation of a liquid? When a liquid evaporates, it does so from the surface of the liquid. When liquid boils, it boils from the interior of the liquid and the bubbles are propelled to the surface by the fact that they have a lower density than liquid water. These bubbles expand as they rise because they aren’t affected by the hydrostatic pressure of being deeply under the water, which makes them even more buoyant. Because boiling involves the superheating of liquids, there can be an explosion of the liquid if it is disturbed as the liquid is superheated in a container. When water is microwaved, superheating can occur throughout the liquid so that when a powder is added to it, such as hot cocoa powder, the liquid can suddenly boil and will potentially burn the individual. This doesn’t happen when water is boiled on the stovetop, where it boils from the bottom up. Sublimation involves the direct vaporization of a solid to its gaseous form. This happens with any solid, including ice that will evaporate in cold, dry weather. It will involve a measurable number less than its melting point. This will be a different number

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depending on the vapor pressure. It is seen when dry ice or CO2 solid sublimes at 1 atmosphere pressure without going through the melting phase. The temperatures and pressures at which a given phase of a substance is considered stable is considered a property of the substance (with temperature and pressure being important factors). A phase diagram can be written of any element or molecule. Figure 36 showed a phase diagram for water and figure 36 shows a phase diagram for a hypothetical substance:

Figure 36.

A phase diagram is somewhat distorted because of the wide variations in temperature and pressure and the need to use one diagram to describe some crucial features of phase changes. The phase diagram goes from low temperature to high temperature and from low pressure to high pressure. There are several lines noted, including the melting line that depends on the temperature primarily. The sublimation curve involves the curve at which solids go to gaseous form. The vapor pressure curve involves liquids going to gases. There are two points to note on the phase diagram. The first is the triple point. This is the equilibrium point of a certain temperature and pressure, when the vapor pressures

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of solids, gases, and liquids will be identical. The second point is the critical point as mentioned already. This is the endpoint of the vapor pressure curve with no separate liquid phase present. Above this point, the substance is referred to as a fluid or a “supercritical fluid” involving high temperatures and pressures. There are several crucial things to know on the phase diagram. These are the normal melting temperatures and the normal boiling temperatures of a substance. These exist as the pressure is 1 atmosphere. Water is unique in that the slope of the melting curve is dependent on pressure, meaning that the melting point of ice decreases with increasing pressure. If subjected to high pressure, ice at zero degrees will melt and ice floats on liquid water, which is an unusual property compared to other substances. In addition, water can be supercooled, which is the phenomenon of “freezing rain”. This is a nonstable state of water in which it can be cooled below its freezing point for a period of time. The triple point of water is just 0.0075 degrees above the freezing point, allowing for all three phases to coexist at this temperature. Above the critical point of water, which is 374 degrees Celsius, no separate liquid phase can exist in water. Interestingly, dry ice is dry because its triple point is at 5.11 atmospheres so under normal pressures, there can be no liquid CO2; it all goes to the gaseous phase under the pressures seen in normal use circumstances. The critical point of CO2 is only 31 degrees Celsius. This is why CO2 extinguishers must be very strong to withstand the pressure of 73 atmospheres in which the CO2 in the fire extinguisher would be entirely vaporized. Supercritical CO2 fluid is used as a solvent to remove the caffeine from coffee beans. Every substance has a phase diagram that will follow similar rules but that are unique to the substance. Helium, for example, has a very unique phase diagram in which sublimation is not possible. It can only be frozen at very high temperatures and there are two liquid phases (helium I and helium II). Helium II is essentially a superfluid, which behaves differently from a quantum perspective when compared to ordinary atoms.

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CRYSTALS Crystals or solid forms of substances will have “faces”. These will have certain geometrical relationships, resulting in symmetry in crystals. Different solid molecules will have different crystalline structures. Sodium chloride is an easy one because if forms a cubic structure that can be seen with a simple magnifying glass. Figure 37 shows what a cubic crystal looks like molecularly:

Figure 37.

Irregularities or differences in the pattern of crystal shapes are referred to as “crystal habits”. With sodium chloride, the general shape of things is that it grows into a completely cubic shape. Partial melting and impurities can create different habits of the cubic shape. While a crystal can become broken or distorted, it will have the same angles between its corresponding faces. This is referred to as the law of constant angles. When a crystal is cleaved, it will cleave along particular cleavage planes according to the law of constant angles. The cleavage angles are determined by the molecular structure of the molecules in solid form. The different possible ordered shapes include the cube, the rectangle, the parallelogram, the rhombus, and the hexagonal shape. The rhomboid shape is similar to the hexagonal shape, with the rhomboid shape being simpler. These are the two-dimensional shapes that crystals can take. As you know, snowflakes take on a hexagonal shape. Why is this the case? It’s because water crystals or ice forms a hexagonal pattern when ice crystalizes. They form unique

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shapes but, in reality, snowflakes are entirely based on the hexagon. Figure 38 shows the hexagonal basis of the snowflake or ice crystal:

Figure 38.

In reality, there are fourteen different lattice shapes, which are referred to as Bravais lattices. These shapes or “crystal systems” can be defined by the length of the sides of the 3-dimensional shapes and the angles that exist between the sides. The different crystal systems in solid elements and molecules include the following: •

Cubic—there are three possible cubic shapes that can be made molecularly.

•

Tetragonal—these involve two equal sides and one unequal side. There are two possible molecular ways to make this shape.

•

Orthorhombic—there are four possible molecular ways to make this. The angles are all 90 degrees but the side length are different.

•

Hexagonal—there is one way to create this shape. It is created from three rhombic prisms.

•

Trigonal (rhombohedral)—there is one way to make this shape molecularly. The sides are all the same length but one of the angles is not 90 degrees.

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•

Monoclinic—this is a three-dimensional shape in which none of the sides are equal.

•

Triclinic—this is a shape in which no side is equal and no angle is 90 degrees.

Figure 39 shows what the different shapes of crystals look like:

Figure 39.

There are four types of crystals as described by the particles that make them up. These include: 1) ionic crystals, 2) metallic crystals, 3) covalent network crystals, and 4) molecular crystals. Ionic crystals include things like sodium chloride (NaCl) and calcium fluoride (CaF2). Metallic crystals include mercury (Hg) and sodium (Na). Covalent network crystals can be things like diamonds (carbon) and silicon oxide (SiO2). Molecular crystals can include iodine (I2) and ammonia (NH3).

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Ionic crystals will have alternating positively-charged cations and negatively-charged anions. These can be polyatomic or monoatomic ions. These are typical of group 1 or group 2 metals along with group 16 and group 17 nonmetals or polyatomic ions. They have high melting points and are extremely brittle. They do not conduct electricity as solids but do conduct electricity when in aqueous solutions. Figure 40 shows an ionic crystal from a molecular standpoint:

Figure 40.

Metallic crystals consist of a grouping of metal cations in a soup of valence electrons. These are good conductors of electricity because of the delocalized electrons in the group that do not belong to any particular nucleus. These have a wide range of melting points. Covalent network crystals consist of atoms that are covalently bonded to their nearest neighbor. These include things like diamonds, quartz, many metalloids, and certain metal oxides. These have very high melting and boiling points and are extremely hard solids. They do not conduct electricity.

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Molecular crystals are held together by weak intermolecular forces. They can be held together by hydrogen bonds or other intermolecular forces that are much weaker than ionic and covalent network crystalline forms. They lack free electrons and so do not conduct electricity well. They have much lower boiling and melting points when compared to other crystal types.

LIQUIDS So far, we’ve discussed phase changes as properties of different substances. Liquids in particular have unique properties that depend on the nature of the intermolecular relationships. In this section, we will talk about these things, including the surface tension, capillary action, viscosity, and boiling point. Surface tension is described as the energy required to increase the surface area of a liquid by a specific amount with the SI units for this being in joules per meter squared. The bonding of water involves hydrogen bonding, yielding a high surface tension of water compared to other liquids. Mercury is even higher with a surface tension six-times that of water. With water, a glass of water can actually be overfilled with the level of water higher than the rim of the glass. Paper clips can float on the surface of water, despite the density, because of the surface tension of water, which prevents the “breakage” of the surface of the water. There are other ways to measure the surface tension besides as joules per square meter. Another is dyne per centimeter where one dyne is 1 x 10-5 Newtons. The stronger the intermolecular forces, the higher the surface tension. High surface tension leads to high boiling points. Mercury is the main anomaly to this in that it has a high surface tension but not a very boiling point . In the case of mercury, the surface tension is high because of the metallic forces between the molecules. Soaps and detergents can disrupt the normal intermolecular attractions seen in water molecules so these can decrease the surface tension. These are also referred to as surfactants. Surfactants are used to fight fires in large-scale firefighting because they can spread across burning surfaces.

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Intermolecular forces also play a role in determining the capillary action of a liquid. Capillaries are small tubes, like straws, that have forces causing water to rise to higher than the level expected when a tube is put into a larger bath of water. The smaller the diameter of the tube, the higher the liquid rises. It involves both the cohesive forces of the liquid molecules with each other and the adhesive forces that bind the liquid to the surface of the tube. The molecular makeup of the capillary makes a difference as well. Glass contains silicon hydroxide, which is polar and attracts water molecules (that are also polar). If the adhesive forces are greater than the cohesive forces, the liquid will rise in the tube against the force of gravity. Mercury has a negative capillary action, in part because of its lack of adhesive force with glass. The meniscus (the shape of the upper surface of a column of liquid) is concave with water but is convex with mercury because of the different phenomena involved. Figure 41 shows the capillary action in water and mercury:

Figure 41.

Viscosity is the resistance of liquid to flow. Certain liquids have a low viscosity (like water and gasoline), while others (like molasses) have a high viscosity. It can be

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measured in several ways. For example, it can be measured as the time it takes for a certain quantity of liquid to flow through a narrow vertical tube. It can also be measured as the time it takes for a solid to fall through a given volume of the liquid. This is measured in poise units. The higher the number, the higher the viscosity. Strong intermolecular forces will increase the viscosity. Long flexible molecules will have higher viscosities. This is why long-chain hydrocarbons like certain motor oils have high viscosity. Hot liquids will have lower viscosity because of increased kinetic energy. The vapor pressure of a liquid is the pressure of a vapor when it is in equilibrium with its more condensed phases at a given temperature and in a closed system. The vapor pressure of a liquid will increase with temperature until the liquid boils. As mentioned, the normal boiling point of a liquid is that at 1 atmosphere of pressure. This is exactly 100 degrees Celsius with water. Of course, the actual boiling point is pressuredependent. Molecules of a liquid need a greater kinetic energy (higher temperature) to escape under higher pressure. This is how cooking in pressure cookers can be so successful.

LIQUID FORCES Liquid forces are more similar to that of solids than they are to gases. It’s the intermolecular force between the molecules that affects the properties of the liquid. These forces will be less than that seen in covalent bonding. Using water as an example, it takes 927 kJ of energy to break the covalent bond between a hydrogen ion and the hydroxyl ion but only about 41 kJ of energy to overcome intermolecular forces in order to convert liquid water to gaseous water. The melting/freezing points of solids and the boiling points of liquids are determined by intermolecular forces. These forces are primarily electrostatic in nature—between positively-charged and negatively-charged molecules. These forces fall off rapidly as intermolecular distance falls so they become more important between molecules in liquid and solid form versus those in gaseous form (except, of course, at high pressures). Ionic forces are referred to as Coulombic forces because they involve the phenomena of “like repels like” and “opposites attract”. These forces are very strong—even stronger

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than covalent bonding so that it takes greater energy to break apart sodium and chloride ions than it does to break two hydrogen atoms together. As we’ve talked about, water is a polar molecule, having a dipole moment across the entire molecule with a partial negative charge assigned to the oxygen atom and a partial positive charge assigned to each hydrogen atom. This polarity affects the ability of ions to dissolve in a liquid, leading to hydrated ions. The strength of the ion-dipole attraction depends on the magnitude of the dipole moment and on the charge density of the ion. The charge density of the ion is its charge divided by its volume. This means that smaller positive ions have larger charge densities than negative ions. The largest charge density belongs to Hydrogen ions, which is why it exclusively found as H3O+ (hydrated hydrogen ions) in water. Ion-dipole interactions always are negative (attractive) because an ion will line up to the aspect of the molecule that is attractive to it. The potential of the attraction in ion-dipole interactions drops off faster than is the case with ion-ion interactions. There is the inverse square law with ion-dipole reactions and a linear relationship in ion-ion interactions. There are also dipole-dipole interactions, in which polar molecules that have a partial positive charge on one end and a partial negative charge on the other end will have an interaction with themselves. While there are attractive and repulsive forces in polar liquids, the overall forces are positive in liquids that have molecules which can move freely to align themselves in whatever direction they want to. These are weaker forces than are seen in any other type of bonding (such as ionic bonding). The interaction decreases as the distance r between the two molecules to the sixth power or interaction equals 1/r6. Being liquid doesn’t just apply to polar molecules. There are interactions between nonpolar molecules like benzene and hexane, which are liquid at room temperature. Depending on the pressure and temperature, even noble gases can be liquid. But what types of interactions are involved in this? These involve what are called London dispersion forces, which involve transient distributions of electrons within the atoms themselves, producing attractive forces between molecules.

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These London dispersion forces involve non-equal distribution of electrons at any given point in time, creating an instantaneous dipole moment within the atom itself. These are weak forces that fall off rapidly with increasing distance and are much less than any ionion force, ion-dipole force, or dipole-dipole force. London dispersion forces get stronger with larger molecules because these molecules can shift their electrons to a greater degree than smaller molecules. The London dispersion forces explain why there is a general trend toward higher boiling points with liquids that have a greater molecular mass and a greater surface area. The shape of the molecule will determine how much of a molecule can interact with its neighboring molecules at any given period of time as well. Long and thin molecules have a greater surface area to interact with other molecules than short and fat molecules. Molecules that are polar can have London dispersion forces as well. Dipole-dipole interactions in small polar molecules are much greater than London dispersion forces. These two forces are collectively referred to as van der Waals forces, although the term can apply to any weak intermolecular force. We’ve talked about hydrogen bonding briefly but will talk about it here as it represents a type of intermolecular bonding force—a strong dipole-dipole force between hydrogen bond donors and hydrogen bond acceptors in two different molecules. Hydrogen bonding accounts for why HF, H2O, and NH3, have such high boiling points compared to less electronegative molecules. The large difference in electronegativity causes a partial positive charge on the hydrogen atom and a partial negative charge on the electronegative atom (oxygen, fluorine, or nitrogen). This leads to a strong dipole-dipole interaction between these molecules and intermolecular forces between the two molecules. Figure 42 shows hydrogen bonding in an H2O molecular situation:

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Figure 42.

Hydrogen bonding accounts for the high boiling temperature of water as well as its high surface tension and high heat of vaporization. Water has a high viscosity compared to similar liquids because of its intermolecular forces. Its cohesion is greater because of its polarity and it can dissolve many things because of its polarity.

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KEY TAKEAWAYS •

Phase changes depend on the pressure and temperature of the molecule.

•

A phase change diagram involves the changes of phases including solids, gases, liquids, and supercritical fluids.

•

The normal boiling point is the boiling point at 1 atmosphere.

•

Solids will form crystals of different shapes and types.

•

Liquids have certain properties that are based on their intermolecular forces.

•

There are several types of intermolecular forces that affect the properties of the molecule.

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QUIZ 1. What most affects the vapor pressure of a liquid? a. Its polarity b. Its temperature c. Its volume d. Its melting point Answer: b. The temperature of a liquid will increase the vapor pressure of a liquid. 2. What is the point at which a liquid is in equilibrium with its vapor at a partial pressure of 760 torr called? a. Critical point b. Vapor pressure c. Normal boiling point d. Freezing point Answer: c. The normal boiling point is the point at which a liquid is in equilibrium with its vapor at 760 torr or one atmosphere of partial pressure. 3. What is not an equivalent number at 100 degrees and 760 torr? a. Condensation point b. Boiling point c. Dew point d. Critical point Answer: d. Each of these is the same value under the conditions of 100 degrees and 760 torr with the exception of the critical point where there is no possibility to have the gas become liquid.

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4. What does the superheating of water involve? a. Raising the temperature so all of the liquid boils off. b. The raising of the temperature above the critical point of water. c. The raising of the temperature of water above its boiling point. d. The raising of solid water above its melting point. Answer: c. Superheating of water involves raising its temperature above its boiling point, which usually requires increasing the pressure circumstances of the water being superheated. 5. What is the critical point of water? a. 0 degrees Celsius b. 100 degrees Celsius c. 374 degrees Celsius d. 742 degrees Celsius Answer: c. The critical point of water is 374 degrees, above which it is a superfluid rather than a liquid or gaseous substance. 6. A phase diagram for an element is a function of what two factors? a. Pressure and Temperature b. Pressure and boiling point c. Temperature and sublimation point d. Boiling point and volume Answer: a. There is a relationship between temperature and pressure, which creates the phase diagram for all substances in nature.

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7. Surface tension of liquid can be described as what it takes to break the surface of the liquid by another force. What is the SI unit for surface tension? a. Joules b. Joules per liter c. Joules per square meter d. grams per square meter Answer: c. This is the energy that it takes to break the surface of a liquid; it is described as the energy per square area or, in the case of SI units, it is joules per square meter. 8. What property of water is most affected by surfactants? a. Surface tension b. Viscosity c. Boiling point d. Capillary action Answer: a. Surfactants interfere with the intermolecular forces and can therefore reduce the surface tension of water, making these substances good for firefighting and in using water for cleaning. 9. With ion-dipole interactions between ions and polar molecules like water, what is the relationship between the strength of the interaction and the distance between the molecules? a. There is a positive relationship between the two that is linear. b. There is a negative relationship between the two that is linear. c. There is a positive relationship between the two that is nonlinear. d. There is a negative relationship between the two that is nonlinear. Answer: d. There is a negative relationship between the two that is nonlinear. The interaction falls off with increasing distance in a nonlinear fashion. This is because of the inverse square law.

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10. Which type of intermolecular force is the weakest? a. dipole-dipole forces b. dipole-ion forces c. ion-ion forces d. They are roughly the same Answer: a. Dipole-dipole forces are the weakest of those listed, falling off by a factor of r to the sixth power, where r is the distance between the molecules.

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CHAPTER EIGHT: SOLUTIONS AND COLLOIDS The study of liquids continues in this chapter in a discussion of solutions and colloids. Solutions involve the dissolution of a substance in a liquid medium. Exactly how this is successfully done is covered as part of this chapter as well as how to demonstrate the “concentration” of a solute in solution. A solution is a unique medium for reactions to take place, as it exhibits certain properties that keep things in a given solution and that help drive reactions within the solution. The properties of colloids are also discussed in this chapter.

THE IDEAL SOLUTION We’ve talked about ideal gases before but there are also “ideal solutions” which do not technically have the properties of real solutions. An ideal solution is a solution in which the change in enthalpy of the solution versus the solvent (or liquid) alone is zero or the “delta H is zero”. The solution in which the change in enthalpy is close to zero is considered a more ideal solution. What you will find out, however, is that most solutions have nonideal characteristics. There are laws that determine the behavior of solutions. There is Raoult’s law, which states that the vapor pressure of a solution is equal to the vapor pressure of the pure solvent at the same temperature scaled by the mole fraction of the solvent present. It states that the partial vapor pressure of each component of an ideal mixture of liquids is equal to the vapor pressure of the pure solvent multiplied by its mole fraction in the solution. In other words, if a non-volatile solute (zero vapor pressure, does not evaporate) is dissolved into a solvent to form an ideal solution, the vapor pressure of the final solution will be lower than that of the solvent. The decrease in vapor pressure is directly proportional to the mole fraction of solute in an ideal solution. So, what does this mean? It was discovered that, when a substance is dissolved in a solution (such as the dissolution of sodium chloride into water), the vapor pressure of the solution will decrease. The vapor pressure change depends on two factors: 1) the

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mole fraction of the amount of dissolved solute present, and 2) the original vapor pressure of the pure solvent. Remember that the vapor pressure is the vapor formed above a liquid (or solid) in which the vapor (or gaseous form) is in equilibrium with the solid form (at any given temperature). This is a constant that is independent of the amount of substance present. Raoult’s law only works for ideal solutions, similar to laws that apply to ideal gases versus real or non-ideal gases. The biggest difference in ideal gases versus ideal solutions is that ideal solutions have intermolecular forces (which we’ve already talked about in the last chapter). While Raoult’s law only applies to ideal solutions, it works fairly well for dilute solutions under many conditions. In reality, the decrease in vapor pressure will be greater than that which can be calculated by Raoult’s law for extremely dilute solutions. Adding a solute to a solution will lower the vapor pressure because the additional solute particles will fill the gaps between the solvent particles and will take up space so that less solvent will be on the surface of the liquid. This means that less of the solvent will break free to enter the gaseous phase. Figure 43 will show what the vapor pressure means:

Figure 43.

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It is the pressure you get at equilibrium when the number of particles sticking to the to surface of a liquid is the same as the number of particles escaping into the gaseous phase. So, in a solution that is 50 percent solute and 50 percent solvent, only half of the solution will be available on the surface of the liquid so this would theoretically cut the vapor pressure in half. The tendency of these particles to leave the liquid to make a gas would be decreased but the tendency of gaseous particles to want to re-stick to the liquid would stay the same. The overall effect is that the saturated (equilibrium) vapor pressure is lower in a solution. Figure 44 shows the behavior of solutions according to Raoult’s law:

Figure 44.

As you can see, the tendency of the vapor pressure will decrease as the solution becomes more concentrated but that it works for nonideal solutions when the solutions are very dilute. In an ideal solution, it would take exactly the same amount of energy for the solvent to break away from the solution as it would in a pure solvent, ignoring the fact that there are intermolecular interactions between the solvent and solution that would make this energy different. This is the main thing that Raoult’s law ignores. The nature of the solute also plays a role in whether or not Raoult’s law comes into play. What really matters is the number of moles of substance you put into the solution or the number of particles of solute in solution. In ionic substances, like sodium chloride, there

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are two moles of particles when one mole divides into two ions. This is twice the number of moles of particles as would be seen in the case of say, a sugar, dissolving in solution. You have to double the mole fraction in the calculation if a salt divides into two portions in solution. Raoult’s law affects the phase diagram of a solvent. Remember the phase diagram of a pure solvent shows the normal boiling point and normal freezing point of a solvent. So, what happens when things go into solution? If the solute itself is non-volatile, the curve will be shifted to the right so that the pressure will always be lower than the solvent itself at a particular temperature. Figure 45 shows the phase diagram of a solution versus a solvent.

Figure 45.

The triple point has moved in the figure as shown and, as you can imagine, it shifts the boiling point of the solution to the right or to a higher temperature. In addition, it will shift the freezing points of the solution to the left or to a lower temperature. We’ll talk more about that in a few minutes. One can calculate the vapor pressure of a solution in a couple of ways, depending on whether there is volatility of the solute. If the solute is volatile, it will exert its own vapor pressure, which will affect the overall vapor pressure of the solution. This is not the case

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in nonvolatile solutes. In nonvolatile solutions, you need to determine the change of vapor pressure for the solvent itself. Let’s use an example. You have 1.5 moles of sugar dissolved into two liters of water. The normal vapor pressure of water is 23.8 mm Hg at 25 degrees Celsius. What is the change in vapor pressure? First you need to know that 2 liters of water equals 2000 grams of water. The molar mass of water is 18.02 grams/mole so the number of moles is 110.9 moles of water. This leads to the calculate mole fraction of water to be 0.979 by dividing the moles of water by the total moles of water plus sugar together. Multiplying this by the vapor pressure of water at 25 degrees, you get 0.979 multiplied by 23.8 mm Hg to get 23.3 mm Hg. If the solute is an electrolyte, you need to double the numbers to reflect the fact that one mole of a salt equals two moles of ions in solution. This means that electrolytes have a bigger effect on the change in vapor pressure of a solution versus a non-electrolyte in the same number of moles added. In volatile solutes, you have to account for the partial vapor pressure of the solute as well. The sum of the vapor pressure is the vapor pressure of the solute plus the vapor pressure of the solvent. You take the vapor pressures of the two components, multiply by the mole fractions by each component and add the two partial pressures. In figure 44, we referred to Henry’s law but haven’t yet explained it. It was formulated in 1803 and states that, at a constant temperature, the amount of gas that can dissolve into a volume of liquid is directly proportional to the partial pressure of the gas in equilibrium with the liquid. Unlike that of Raoult’s law, it works mainly at high concentrations of gas in comparison to the liquid. This works, for example, when air is completely saturated over water with water as the solvent. It does not work at high pressures because gases will push on the liquid and will go into solution. It only works when the molecules are at equilibrium. It doesn’t work if there is a chemical reaction between the solute gas and the solvent.

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NONIDEAL SOLUTIONS Both Henry’s law and Raoult’s law apply to ideal solutions only. There are a variety of forces that act on real mixtures, making it difficult to predict the properties of these types of solutions. There are two types of “nonideal” solutions, including when there are intermolecular forces that are less strong between the solute and solvent than that of the solvent alone and when there are intermolecular forces that are stronger than those of the solvent itself. The solvent can always be described as the major component of a solution, while the solute is the minor component of the solution. In situations where there are weaker forces between dissimilar molecules than between similar molecules, the resulting solution will have a larger enthalpy so the process is considered to be endothermic. This will mean that the forces that attract the solutes exceed the force that attracts the solutes to the solvent AND that the forces that attract the solvent molecules to each other also exceed s the solute to solvent attraction forces. A nonideal situation also exists when the intermolecular forces between dissimilar molecules are greater than those between similar molecules. These interactions will release energy making the solution-making process exothermic. There is an activity coefficient, which is a correctional factor without a unit that takes into account the nonideal characteristics of a solution. This activity coefficient is a number that centers around the value of 1. If there is no interaction, the coefficient is 1. If there is repulsion between the molecules, the coefficient is greater than one. If there is attraction between the molecules, the coefficient is less than one.

ELECTROLYTE SOLUTIONS There are attractive forces that make electrolyte solutions possible. The force of attractions is described by Coulomb’s law, which states that: The magnitude of the electrostatic force of attraction or repulsion between two charged molecules is directly proportional to the product of the magnitudes of charges and inversely proportional to the square of the distance between them. If the charges are opposite, there will be attraction and if the charges are the same, there will be repulsion.

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Because ions are not stable on their own, no ion can be studied separately. In biological systems, all ions will have a counterion (a cation or anion) that balances its charge. There is no possibility of measuring the enthalpy and entropy of an ion like hydrogen ions, so there is a reference point that has been established. The change in enthalpy, entropy, and Gibbs energy for a hydrogen ion is said to be zero. Remember that the Gibbs energy is the energy associated with a chemical reaction that can be used to do work or the sum of its enthalpy (H) plus the product of the temperature (Kelvin) and the entropy (S) of the system. A solution can be a homogeneous mixture of at least two substances in the gas phase, the liquid phase, or the solid phase. The enthalpy change in a solution refers to the amount of heat that is released or absorbed during the dissolution or dissolving process of a substance at a constant pressure. This is called the enthalpy of solution (which can be positive and endothermic or negative and exothermic), expressed as “Delta H of the solution). In order go into solution, the intramolecular forces between the solute molecules must be disrupted. This can be called the delta H one (ΔH1). This will always be endothermic because it requires the breaking of bonds. The second process requires the breaking up of the intermolecular forces that hold the solvent together. This will also be an endothermic process resulting in a positive delta H2. The third process is the mixing of the two substances. When this occurs, there will be a delta H3 or a change in enthalpy stemming from the mixing of the two substances. This will be a negative number because it will be exothermic. The change in enthalpy of the solution is a sum of all three enthalpy changes. It can be endothermic, exothermic, or zero. If the enthalpy change is zero, the solution is ideal. If it is anything other than that, it is a non-ideal solution. Figure 46 shows the process of mixing sodium chloride to make a solution:

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Figure 46.

So, what it all means is that in an ideal solution the sum of the strengths of the intermolecular forces between the solute and the solvent will be zero. The sum of the endothermic forces of breaking up the solvent and the solute as well as the opposite force of attraction between the solute and solvent will be zero so the negative forces and positive forces will balance each other out. Using the rule of like dissolves like with the making of ionic solutions, there must be two things to look at. The first is the strength of the ion-dipole forces of attraction between water and the ionic compounds. The second is the strength of the interionic bond of the ionic compound itself. In order for the ionic compound to make a solution, the iondipole forces need to be greater than the interionic bonds or ion-ion forces.

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The ionic compound becomes surrounded by water when it goes into an aqueous solution, becoming hydrated with the water’s negative ends attracted to the positive ion and water’s positive end attracted to the negative ion. These forces of attraction will overwhelm the interionic bonding so that the molecules will be surrounded by water molecules. In the case of an ion dissolution, the breaking of the ionic compound is endothermic but hydrating both of the ions will be exothermic. The sum of these steps will lead to the enthalpy of solution.

MIXING SOLUTIONS The process of mixing two or more substances involves increasing the entropy of the solution. There will be an inevitable heat transfer, referred to as the delta H of the solution or the enthalpy of the solution. Remember that enthalpy is a state function so that the change in enthalpy of a solution will be equal to the enthalpy of the products minus the enthalpy of the reactants. This would indicate that the process of dissolution or the mixing of substances is similar to a chemical reaction, just like phase changes and regular chemical reactions. In an ideal situation, or an “ideal solution” there will be no change in enthalpy when substances mix together. The solvent and solute will mix at random. As mentioned, nonideal situations exist if the enthalpy is positive or the enthalpy is negative, meaning there are differences between the intermolecular forces of similar molecules versus dissimilar molecules. If the enthalpy change is considerably positive, the mixture will be nonhomogeneous or heterogeneous (as in a mixture of oil and water).

EXPRESSION OF CONCENTRATIONS Solutions are considered homogeneous mixtures in which there is one solvent and at least one solute. The solvent’s “concentration” is considered to be considerable, while the solute is considered to be dissolved into the solvent. There are different ways to describe the concentration of a solute, including these commonly used ways: •

Mass percent—this is a percent based on the mass of solute divided by the mass of the solution and multiplied by 100 percent. 148

•

Volume percent—this is a percent based on the volume of solute divided by the volume of solution and multiplied by 100 percent.

•

Mass/volume percent—this is the mass of the solute divided by the volume of solution, such as grams divided by milliliters, multiplied by 100 percent. One example of this is 0.9 percent sodium chloride solution, which is considered “normal saline” in medicine. This is a weight/volume percent because 0.9 grams is dissolved in 100 milliliters of solution. It approximates the mass/mass percent because the density of a dilute aqueous solution is nearly 1 gram per milliliter.

•

Proof—this is how alcohol content is measured in the US. It is a weight/weight calculation of alcohol to water solution, multiplied by 2 to get a maximum of 200 proof.

The mass/mass or volume/volume percent must match grams to grams and milliliters to milliliters in order to get the correct percent. Another way to describe the percent in very dilute solutions is to use the “part” system. This would be something like 1 part of something to 30 parts of another. In very dilute solutions, things like parts per million or parts per billion are used in order to describe a dilute solution. One part per million or ppm is the equivalent of 1 gram per liter of solution. One part per billion is the same as a microgram of solute per liter of solution. This is represented as ppb. Parts per trillion or ppt is the same as 1 nanogram of a solute per liter of solution. When dealing with certain equations, such as PV = nRT (one of the gas equations), the mole fraction of a substance can be used. This is the number of atoms out of the total number of atoms or molecules as a fraction. The sum of the mole fractions of the different substances in a mole fraction situation will equal one. The mole percent is the same as the mole fraction but it is multiplied by 100 percent. Molarity or capital M is the number of moles per liter of a given solution. It must be calculated by knowing the number of grams and the molar mass of a given molecule in order to calculate the molarity. The molality or small m is the number of moles of solute per kilogram of the solvent (not the solution itself). There is a big difference between molarity and molality. Volume varies with temperature so molarity varies with 149

temperature, while molality does not. This makes molality a better measurement when temperature is a factor in the equation.

COLLIGATIVE PROPERTIES The colligative properties of a solution will depend on the numbers of solvent and solute particles and not on the nature of the particles. They involve the vapor pressure decrease, boiling point elevation, freezing point depression, and osmotic pressure. They are based on the solute concentration and a given constant. The colligative property of boiling point elevation fits the relationship of this: boiling point elevation equals the solute concentration (molality) multiplied by a constant, in this case, the boiling point constant. The vapor pressure change is based on the mole fraction of the solute and the vapor pressure of the pure solvent. The freezing point depression is based on the molality of the solute and the freezing point constant. The osmotic pressure is based on the molarity of the solute. Note that the freezing point and boiling point changes are based on molality as this is not temperature-dependent. Boiling points of solutions will all be higher than that of the pure solvent. This difference is proportional to the concentration of the solute particles. The change in boiling point equals the boiling point constant multiplied by the molality of the solute. The boiling point constant is also called the ebullioscopic constant, which is dependent on the properties of the solvent. The freezing point depression is a greater phenomenon than the boiling point elevation; it is also based on the molality of the solute as well as a constant that varies according to the solvent. Like the boiling point elevation, the change in temperature is the constant multiplied by the molality of the solute. Both the freezing point depression and boiling point elevation can be predicted by the phase diagram changes of a solution versus a solvent. The boiling point of any substance is reached when the chemical potential of the pure solvent in liquid form reaches that of the chemical potential of the pure vapor. It will be a higher temperature in the solution versus the solvent form. The freezing point is reached when the chemical potential of the pure solvent as a liquid reaches that of the 150

pure solvent as a solid. This will be a lower temperature in solution and explains why salt is added to roads to melt the ice. The maximum freezing point is 0 degrees Fahrenheit, which is -18 degrees Celsius under high concentrations of sodium chloride. As you can imagine, Calcium chloride will be used for lower temperatures because it divides into three ions, rather than two because its chemical formula is CaCl2 (breaking into one ion of calcium and 2 ions of chloride). Osmosis is the diffusion of a fluid through a semipermeable membrane. Only the solvent is able to pass through the membrane and not the solute. The osmotic pressure of a solution is dependent upon the molar concentration of the particles in solution and is the pressure difference needed to stop the flow of a solvent across a semipermeable membrane. It is calculated based on the ideal gas constant (0.0821 Liters atmospheres per mole per degree kelvin), which is designated as R. This makes the osmotic pressure equal to MRT, where M is the molar mass of the solute, R is the gas constant, and T is the temperature in kelvin. Vapor pressure involves the ability of the liquid to switch to the gas phase at the top of the liquid surface. Regardless of the solvent, when it comes to a solution, the vapor pressure will be lowered. This is also going to be independent of the solute. As a colligative property, it depends on how many solute particles will be dissolved in the solution. It will be directly proportional to the concentration of the solute. Because it uses Raoult’s law, it is proportional to the mole fraction of the solute rather than on molality or molarity. Vapor pressure of the solution will be the mole fraction of the solvent multiplied by the vapor pressure of the solvent alone. The mole fractions of solvent and solute will add up to 1.

ANOMALOUS COLLIGATIVE PROPERTIES Anomalous colligative properties are colligative properties that deviate from the norm. Colligative properties are those properties of solutions that rely on the concentration of solute particles in an ideal solution, such as vapor pressure lowering, boiling point elevation, osmotic pressure, and freezing point depression. In these properties, there is a direct relationship between the concentration and the property seen. Some solutes

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have a greater effect on colligative properties than what is normally expected. In ionic solutions of sodium chloride, for example, the effect is twice what can be expected with urea, which is just one solute and not two ions.

SOLUTION STOICHIOMETRY Calculations involving reactions in solution are carried out in the same way as we’ve already discussed except that instead of masses, the volumes of the solutions of known concentrations are used to determine the number of moles of the reactants. The equations must be balanced according to the number of moles of the reactant needed and the number of moles of product made. Instead of the molar mass used in typical reactions involving grams of different substances, the molarity is used as this represents moles per liter (which is multiplied by the number of liters in the reaction). The same concept of limiting reactants applies to solution reactions.

REACTIONS IN SOLUTIONS Some reactions take place entirely in solution. Others involve precipitation reactions, which involve two or more soluble components that take place in solution, resulting in a substance that precipitates outside of solution. This can be used to determine how much of a given substance used in the reaction because the precipitate can be dried and weighed, using the molar weight and stoichiometry in order to determine how much of the reactants are in the reaction. One example is silver nitrate (AgNO3) and potassium chloride (KCl) to make KNO3 (which is aqueous) and silver chloride (AgCl), which is insoluble. When this reaction takes place, one would weigh the silver chloride, determine the number of moles in the dried weight of the silver chloride, and then determine the number of moles and molarity of the reactants. The solubility is the relative ability of any solute (liquid, solid, or gas) to dissolve into a solvent in order to form a solution. The thing that plays the largest role in solubility is temperature, although pressure can affect the solubility of gases. Highly polar ionic compounds will be more soluble in polar solutions, like water but won’t be soluble in 152

benzene and chloroform (which are nonpolar). A solution is considered saturated when no additional solute can be dissolved in the solution. The range is from “fully miscible” or infinitely soluble substances, like ethanol in water to completely insoluble, such as silver chloride in water. Miscibility can be defined as the property of a substance to mix in all proportions (that is, to fully dissolve in each other at any concentration), forming a homogeneous solution. Temperature is greatly related to the solubility of many substances. As water molecules (or any solvent) raises in temperature, they vibrate more quickly and are better able to interact with and break apart the solute, increasing the solubility of a substance in a solvent. This is seen with solids but is the opposite with gases. Pressure has little effect on the solubility of solids or liquids in water. It has a great effect on the solubility in gases. Think about soda, for example. The pressure in the soda can is high when it is sealed but is low when it is open, resulting in a hiss of gas because the gas is coming out of solution. As the can temperature increases, the solubility of CO2 in the water is going to decrease. Like tends to dissolve like so polar molecules will dissolve better in polar solutions and nonpolar molecules will dissolve better in nonpolar solutions. Figure 47 shows the different abilities of certain cation and anion salts in water:

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Figure 47.

Salts of alkali metals and ammonium, and salts containing nitrate and acetate are always soluble. Carbonates, sulfates, hydroxides, phosphates, and heavy metal salts are typically insoluble in water.

COLLOIDS A colloid is one of three main types of mixtures (with solutions and suspensions being the other two). It is a solution with particles that range from 1 to 1000 nanometers in diameter but that are able to remain evenly dispersed throughout the solution. These are also referred to as colloidal dispersions because they will not settle to the bottom of the container.

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As you can see, a colloid is going to be similar to a solution but involves particles that are not seen with a naked eye but that are larger than a molecule. Particles larger than 1000 nanometers will be a suspension and particles less than 1 nm will be a solution. There are four types of colloid classifications: •

Sol—this is a colloidal suspension with solid particles in a liquid.

•

Emulsion—this is a colloid between two liquids.

•

Foam—this is a colloid where gas particles are trapped in a liquid or solid

•

Aerosol—this is a liquid or solid dispersed in gas.

When the solvent is water, the resulting colloid is referred to as a hydrocolloid. Hydrocolloids are often used to make medical dressings. A simple way of determining whether or not a mixture is a colloid or not is to examine the Tyndall effect. When light is passed through a solution, it will pass through it directly. When light passes through a colloid, it will be scattered in all directions, making it readily seen. Fog is an example of a colloid so a headlight beam can be seen through it because of the Tyndall effect. This reflection of light is seen only in colloids. Another example is attempting to pass the colloid through a semipermeable membrane. Colloids will not be able to pass through the semipermeable membrane.

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KEY TAKEAWAYS •

Raoult’s law applies to ideal solutions and relates to the vapor pressure of a solution versus that of a pure solvent.

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Colligative properties involve properties that change in a solution based on the number of molecules in the solution rather than on the characteristics of the solute.

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There are a number of ways that the concentration of a solution can be described—based on weight, volume, and number of moles of the substances.

•

Solubility of a substance is based on its polarity, the polarity of the solvent, temperature, and pressure.

•

Colloids are mixtures of larger molecules between 1 nanometer and 1000 nanometers in a substance in another phase or type of phase.

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QUIZ 1. What is the application of Raoult’s law of solutions in non-ideal solutions? a. It can determine the kinetic energy of the solutes in solution. b. It can determine the actual concentration of solutes in solution. c. It can determine the enthalpy of a solution versus the pure solvent. d. The law cannot be applied to any non-ideal solution. Answer: d. While Raoult’s law can determine the vapor pressure of a solute as an expression of its mole fraction, this is not true of non-ideal solutions so it cannot be applied to real solutions in chemistry. 2. What is the behavior of the vapor pressure of a solution versus the pure solvent in chemistry? a. The vapor pressure of a solution will increase with increased solutes in solution. b. The total vapor pressure of the solution will stay the same in a solution versus the pure solvent. c. The total vapor pressure of a solution will decrease versus vapor pressure of the pure solvent. d. There is equal effect on the vapor pressure of different solutes, regardless of their relative concentrations. Answer: c. The total vapor pressure of a solution will decrease versus the vapor pressure of the pure solvent in solution chemistry.

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3. When does Henry’s law fall apart in determining the solubility of a gas in solution? a. At high pressures b. Before equilibrium c. When gases are reactive with the solvent d. At each of these situations Answer: d. Henry’s law falls apart and does not apply in high pressure situations, before equilibrium, and when gases are reactive with the solvent. 4. According to Coulomb’s law, what is the relationship between the force of attraction or repulsion between two molecules in relation to their distance between them? a. Directly proportional to the distance between them b. Inversely proportional to the square of the distance between them c. Directly proportional to the square of the distance between them d. Inversely proportional to the distance between them Answer: b. The forces between two molecules according to Coulomb’s law is inversely proportional to the square of the distance between them. 5. What will happen to a mixture if the change in enthalpy is large and positive? a. The mixture will become a solution. b. The mixture will be heterogeneous. c. The mixture will heat up after mixing. d. The mixture will require heating in order to go into solution. Answer: b. The mixture will be heterogeneous, such as the mixture of oil and water because the enthalpy change will be too large to overcome.

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6. You measure a milligram of a solute into a liter of solution. What is the best way to represent this concentration? a. Moles per liter b. Parts per trillion c. Parts per billion d. Parts per million Answer: d. In 1 milligram of solute into a liter of solution, there is 1 ppm or one part per million of solute. This makes this a good way to describe this sort of situation in terms of concentration. 7. What is not a colligative property of a solute in a solution? a. Osmotic pressure b. Vapor pressure elevation c. Boiling point elevation d. Freezing point depression Answer: b. Each of these is a colligative property of a solute in a solution except for vapor pressure elevation. In actuality, it is vapor pressure lowering that is seen with solutes in solution. 8. Which colligative property of a solute is based on the molarity of the solute versus some other measure of concentration? a. Boiling point elevation b. Osmotic pressure c. Vapor pressure depression d. Freezing point depression Answer: b. The osmotic pressure of the solution is based on the molarity of the solute. The molality is used for those colligative properties that are temperature-dependent and the vapor pressure depression is based on the mole fraction of the solute.

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9. Which colloidal type is a liquid suspended in another liquid? a. Sol b. Aerosol c. Foam d. Emulsion Answer: d. An emulsion is a liquid suspended in another liquid. By definition, this is a colloid that has particles between 1 and 1000 nanometers that does not fall due to gravity if it is stable. 10. Which colloidal type is a solid or liquid particle suspended in a gaseous solvent? a. Sol b. Aerosol c. Foam d. Emulsion Answer: b. An aerosol will have a solid or liquid that is suspended in a gas and is typical of aerosols you can think of, such as aerosols in aerosol cans.

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CHAPTER NINE: OXIDATION-REDUCTION REACTIONS Two main types of reactions are covered in this chapter, that being oxidation and reduction reactions—which are opposing but interrelated chemical reactions in chemistry. Because these reactions happen in a balanced way, they are often discussed together. Taken together, an oxidation-reduction reaction is any chemical reaction in which the oxidation number of a molecule, atom, or ion is changed by gaining or losing an electron.

OXIDATION AND REDUCTION REACTIONS Originally, the term oxidation was used to describe reactions in which an element reacts specifically with oxygen. An example of this would be when two magnesium atoms combine with oxygen to make 2 magnesium oxide atoms. The opposite reaction would be called a reduction reaction, in which magnesium solid is made from magnesium oxide. This isn’t exactly an opposite reaction because it can involve MgO (magnesium oxide) plus carbon solid at 2000 degrees Celsius to make magnesium solid plus CO2. Historically speaking, oxidation was the term used to describe any reaction in which oxygen gas plus a metal forms a metal oxide. This expanded to involve oxygen-like substances and finally to all reactions in which something loses electrons. The term “reduction” was used because metal oxides lose weight when metal is extracted from them. This involves the gain of electrons (which is the opposite of what one would think). Instead, it can be thought of as a reduction of an electric charge from positive to negative. Since the identification of electrons, it has become clear that oxidation and reduction reactions involve more than just oxygen and involve the transfer of electrons from one atom to another. In the oxidation reaction where magnesium goes to magnesium oxide, there is a loss of electrons from the magnesium atom. This has led to the idea that

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oxidation involves the loss of electrons by an atom, while reduction involves the gain of electrons by an atom (although this is hypothetical). Because electrons are neither created nor destroyed in any type of chemical reaction, it is virtually impossible to have true oxidation without reduction. This essentially renames the concept of these types of reactions so that a redox reaction is one in which the oxidation number of atoms are changed. Consider, the reaction of carbon monoxide or CO gas + H2O gas to make Hydrogen (H2) gas and CO2 gas. Figure 48 shows what this looks like on a Lewis dot structure:

Figure 48.

The total number of valence electrons stays the same in the figure; however, the oxidation state of carbon goes from +2 to +4, and the oxidation state of hydrogen goes from +1 to zero. So, what is the oxidation state? This is also called the “oxidation number”, which describes the degree of oxidation (loss of electrons) in an atom. This can be a positive number, a negative number, or zero. It is the imagined charge that an atom would have if it were a 100 percent ionic compound, with no covalent component. The highest known oxidation number is +9 in the tetroxoiridium (IX) cation, which is IrO4+, while the lowest oxidation number is -4 for carbon in methane gas or chromium in the complex anion referred to as [Cr(CO)4]4-. What this means is that the term “oxidation state” or “oxidation number” is a bit of a misnomer because the oxidation number can change without involving oxygen at all. The oxidation state of an atom does not also represent a true “charge change” in the atom at all. The oxidation state of a given compound depends somewhat on the electronegativity of the atom. Carbon is unique in that, with four valence electrons, it

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can act as a -4-oxidation number atom or a +4-oxidation number atom (or anything in between). The oxidation state is represented in inorganic chemistry by a Roman numeral placed after the name of the element or as a superscript after the symbol of the element. What’s important to think about is that ionic bonding and covalent bonding are extremes of a continuum of bonding. There will always be some covalent characteristics of ionic compounds and some ionic characteristics of covalent bonding. This makes it simpler to determine the oxidation number and the oxidation status of an atom or atoms. In truth, while the oxidation number of MgO is considered to be +2, the reality is that it is more like +1.5 because the interaction is not truly ionic. Figure 49 shows the periodic table of the elements with the groups clearly identified as these are important to understanding oxidation and redox reaction:

Figure 49.

For the groups IA and IIA active metals, the difference between the oxidation state of the metal atoms and the charge on each atom is extremely small and can be ignored. This means that in an inorganic reaction involving sodium as an element, for example, its oxidation number will go to +1 when it is oxidized and its charge is +1. When it comes to group IIIA and group IVA, however, these will form compounds that are largely

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covalent. Aluminum bromide is not Al3+ ions interacting with Br- ions. It exists covalently as Al2Br3, which is covalent, despite the electronegativity of bromine. Transition metals can have even more confusing pictures. Manganese oxide or MnO is basically ionic; however, Mn2O7 is another molecule of manganese and oxygen that is covalent. This doesn’t mean that manganese in the Mn2O7 molecule is a +7 ion. Instead, it is considered to have an oxidation number of +7. When magnesium reacts with oxygen to make MgO, the magnesium acts as a “reducing agent” to give electrons to oxygen, thereby reducing the oxygen gas. On the other hand, the O2 molecule gains electrons to oxidize the magnesium. It acts as the oxidizing agent to magnesium. This is very clear when the substance being acted on is oxygen and magnesium. It becomes less clear when the oxidizing agent is not oxygen. It’s better to think of it as gaining or losing electrons with oxidizing agents gaining electrons and reducing agents losing electrons. The oxidizing agent gets reduced and the reducing agent gets oxidized. What is usually the case then is that metals mostly act as reducing agents and are themselves oxidized. The reaction in which copper plus oxygen plus heat turn into copper (II) oxide is one in which copper is oxidized. That much is clear. If Copper (II) oxide gets mixed with H2 gas, however, the reaction becomes Copper (II) oxide + H2(g) goes to copper solid plus water. In such cases, copper is the oxidizing agent and hydrogen gas is the reducing agent. Figure 50 shows this:

Figure 50.

Every reducing agent is linked to an oxidizing agent and vice versa. This means that every time something is oxidized (gains electrons), it then can become a reducing agent

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so that it loses electrons). These are called conjugate reducing agents and oxidizing agents. The conjugate reducing agent to O2 gas (which is always an oxidizing agent) is the O2- ion. Strong reducing agents include all of the active metals in the group (I)A metals. They do have conjugates but they are weak oxidizing agents—so weak that they almost never actually oxidize anything under normal circumstances. This is why you don’t see sodium as a metal in nature and why it reacts explosively with water. Another example is oxygen or O2 gas. This is such a strong oxidizing agent that O2- is a very poor reducing agent. There are those atoms that are less strong of an oxidizer than oxygen and that are nearly equally likely to be oxidizing agents as they are reducing agents. Every strong oxidizing agent has a weak conjugate reducing agent and every strong reducing agent has a weak conjugate oxidizing agent. The strength of the metal as a reducing agent can be determined by whether anything happens when the metal is mixed with the salt of another metal. An example is the reaction where aluminum metal plus iron oxide goes to aluminum oxide to iron. Because this reaction happens at all, the aluminum gets oxidized and is the stronger reducing agent, while iron oxide is the stronger oxidizing agent. If the opposite reaction were the natural course of the reaction, the results would be different. Enough heat applied to the opposite reaction, however, can run the reverse reaction. Oxidation and reduction do not generally happen alone. These would be “half reactions”; it takes the oxidation of one thing and the reduction of another thing in order to have a single and complete reaction. The reaction must be balanced with respect to the electrical charge on both sides of the reaction. In addition, true oxidation and true reduction do not occur in most cases because the actual situation is an increase in oxidation state or a decrease in oxidation state. This is what happens when the reaction involves primarily covalent bonds.

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REDOX REACTIONS IN COMMON SITUATIONS Hydrogenation could also be used instead of reduction in biochemistry and organic chemistry because it involves increasing the hydrogen number to the carbon atoms. When unsaturated fats become saturated with hydrogen atoms, this is reduction. When glucose gets broken down in to CO2, it is called oxidation. In organic chemistry, there can be the stepwise oxidation of a hydrocarbon, which starts as a hydrocarbon (fully reduced) and goes to make an alcohol, then an aldehyde or ketone, then a carboxylic acid, and finally a peroxide, which is highly oxidized. In metal chemistry, the rusting and corrosion of metals involve the oxidation of the metal to form a metal oxide, such as with iron (III) oxide, which is rust. Redox reactions are highly important in biological processes such as the entire process of cellular respiration, in which glucose, which is C6H12O6, is oxidized by oxygen to make CO2 and water. Of course, this does not depend on oxidation alone because oxygen gets reduced to form water. While photosynthesis is not the reverse of cellular respiration, it does involve the creation of glucose and oxygen from water and CO2 but it takes light energy to do this. The whole process of storing biological energy happens because of redox reactions. Energy substances like nicotinamide adenine dinucleotide or NAD+ get reduced to make things like NADH, which contributes greatly to the creation of a biological proton gradient, which helps to make adenosine triphosphate or ATP, which is the energy currency of all cellular organisms. A reduced biological substance in biological systems is one that has energy, while an oxidized biological substance has spent energy. Free radical reactions in biology are redox reactions in which free oxygen species are created in order to destroy pathogens by attaching electrons to any molecule in the pathogen. The problem is that these can build up and affect the human host, becoming dangerous if they do not attach to another redox molecule, such as an antioxidant. Antioxidants are, in effect, reducers that get rid of oxygen free radicals.

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KEY TAKEAWAYS •

Redox reactions combine the oxidation of one atom and the reduction of another.

•

Reduction means the gaining of one or more electrons in a reaction.

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Oxidation means the losing of one or more electrons in a reaction.

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Reduction is always accompanied by oxidation.

•

Redox reactions are found in things like the rusting and corrosion of metals as well as in biological systems.

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QUIZ 1. Which reaction is considered a reduction reaction? a. 2Mg + O2 leads to 2MgO b. 2MgO + Carbon leads to 2Mg + CO2 c. 2Cu + O2 leads to 2CuO2 d. 2Al + Fe2O3 leads to Al2O3 + 2Fe Answer: b. Each of these is an oxidation reaction or a balanced redox reaction with the exception of the reaction in which MgO goes to magnesium solid, which is a true reduction reaction. 2. What is the best definition of an oxidation number? a. The number of electrons the atom is sharing with other atoms in a covalent bond b. The number of oxygen atoms attached to a compound c. The number of hydrogen atoms lost from a compound d. The imagined charge that an atom would have if it were a 100 percent ionic compound Answer: d. The oxidation number is the imagined charge on an atom if it were a 100 percent ionic compound. This makes it easier to determine whether something is oxidized or not. 3. What is the range of oxidation numbers for carbon as an atom in a reaction? a. 0 to +4 b. -4 to 0 c. 0 to +4 d. -4 to +4 Answer: d. The range of oxidation numbers for carbon is -4 for methane (a fully reduced carbon atom) to +4 for CO2, which is a fully oxidized carbon atom).

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4. Which number reflects the most reduced state or reduction of an atom? a. -3 b. 0 c. +2 d. +4 Answer: a. The most reduced state of an atom is when it has an oxidation number as low as possible. This makes -3 the most reduced state of an atom. 5. In the reaction where calcium solid mixes with hydrogen gas goes to CaH2, which is the reducing agent? a. Calcium b. Hydrogen gas c. CaH2 d. None of these Answer: a. In this reaction, calcium gets oxidized to a theoretical +2 charge, meaning that it becomes the reducing agent to hydrogen, donating two electrons. Calcium is a (II)A atom, making it more likely to get oxidized than reduced. 6. What does the oxidation number of sodium change to when it mixes with Chlorine gas to make sodium chloride? a. -1 b. 0 c. +1 d. +2 Answer: c. In this case, sodium is a (I)A metal, meaning that it becomes positively charged. In such cases, this positive charge becomes its oxidation number.

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7. The reaction of aluminum and sodium chloride going to aluminum chloride plus sodium does not occur. What does it say that the stronger oxidizing agent must be? a. Aluminum chloride b. Sodium c. Sodium chloride d. Aluminum Answer: a. Aluminum chloride must be the stronger oxidizing agent because it cannot be made from oxidizing aluminum with salt to make aluminum (III)chloride. Instead, it is reducible to make aluminum when mixed with sodium, which is a strong reducing agent. 8. Which atom is the strongest reducing agent? a. Calcium b. Iron c. Aluminum d. Sodium Answer: d. This can be predicted by looking at the periodic table. Sodium, being a (I)A metal is most reactive and is the strongest reducing agent of all of these elements. The weakest reducing agent in the list would be iron. 9. What is the chemical reaction that occurs when an iron pipe gets rusty? a. It involves the oxidation of a solid metal b. It involves the reduction of a solid metal c. It involves the gain of electrons by solid metal d. It involves the gain of oxidation number of a metal oxide Answer: a. The rusting of iron metal into metal oxide by its reaction with oxygen involves the oxidation of the solid metal.

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10. Which organic chemical is most highly oxidized? a. Alcohols b. Carboxylic acids c. Hydrocarbons d. Peroxides Answer: d. Hydrogen peroxide is H2O2 and is highly oxidized—making peroxides the most oxidized of all of the listed organic chemicals.

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CHAPTER TEN: ACIDS AND BASES The specific properties of acids and bases are discussed in this chapter. Methods involved in the titration of acids and bases are discussed as part of this chapter as are properties of strong and weak acids and bases. Buffers are weak acids or weak bases that prevent significant changes in the pH of a solution; how these substances work in acidbase chemistry is covered. Finally, the action of hydrolysis is an important part of the discussion in this section.

THEORIES OF ACIDS AND BASES While we haven’t talked much about acids and bases, they are very important to chemistry, biology, and biochemistry. There are three major ways to describe acids and bases. The first is the Arrhenius definition of acids and bases, first developed in 1884. In it, acids produce H+ ions in solution, while bases produce -OH ions in solution. Hydrogen ions are not stable so they develop hydronium ions, which are H3O+ ions. This means that, in any equation involving the dissociation of an acid involves the addition of water so that, when HCl goes to H+ and Cl-, water is included to make H3O+ instead. This theory works for certain acids and bases but, of course, there are acids and bases that do not contain hydroxide ions and things like F- and NO2- that can make basic solutions in water. Nevertheless, it can easily explain reactions like NaOH goes to Na+ and OH-, which is a strong base reaction. Bare hydrogen ions just cannot exist. This would technically be just a bare hydrogen nucleus and simply isn’t stable. This high charge density will attract anything it can, including the water molecule. In such cases, H2O acts as a base and collects the hydrogen proton to make the hydronium ion. As you can imagine, there are problems with this theory. There is the weak base ammonia that will release OH- ions in the presence of water but doesn’t contain OH- by itself. The reaction is shown in Figure 51: 172

Figure 51.

Another reaction of sodium hydroxide versus ammonia shows how the bases dissociate and interact with a strong acid; in this case, it is HCl or hydrochloric acid. This is depicted in Figure 52:

Figure 52.

In 1923, the Bronsted-Lowry definition of acids and bases was uncovered. In such cases, acids are identified as proton donors, while bases are proton acceptors. There are compounds that act as Bronsted-Lowry acids and bases together, which are called amphoteric molecules. The acid no longer has to give an H+ ion and a base no longer has to give a hydroxide (OH-) ion in order to be an acid or a base. This gives rise to conjugate acids and bases, which are similar to the conjugate reducers and oxidizers seen in the previous chapter. For example, in the reaction where HCl plus NH3 goes to NH4+ + Cl-, there is the acid HCl and the “conjugate base” which would be Cl-. In the same way NH3 is the base, while NH4+ is the conjugate acid. In this case, there is a basic salt, such as NaF or sodium fluoride. It interacts with water to generate OH- ions because it rapidly takes the hydrogen ion as an electron acceptor. What is introduced in the Bronsted-Lowry equation is the concept of the acid ionization constant of an acid as well as the base ionization constant: in deriving this constant you get these equations: In acid dissociation, acid HA goes to A- + H+ ion so the acid ionization constant becomes the product of the concentration of A-, the concentration of H+, and the concentration of HA. In the same way, the base dissociation becomes base B plus water 173

goes to HB+ plus OH-. The base constant is the product of the HB+ ion, OH- ion and B. Figure 53 shows what this looks like:

Figure 53.

The key feature of this is that, at equilibrium, every acid has a conjugate base and vice versa. This gives rise to the Kw or equilibrium constant for water, which is the product of the concentration of hydrogen and hydroxide ions for water. This is a known constant at 25 degrees Celsius, which is 1 x 10-14. What this means is that water is a known acid and a known base and that the dissociation constants for acids and the conjugate bases multiplied together equals the equilibrium constant for water. Strong acids are those compounds that essentially ionize completely in an aqueous solution, dissociating readily into H+ ions and its matching anion. In fact, there aren’t very many common strong acids that do this. Strong acids include the following: •

HClO4 (perchloric acid)

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HClO3 (chloric acid)

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HBr (hydrogen bromide or hydrobromic acid)

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CH3O-

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NH2-

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HCl (hydrochloric acid)

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HNO3 (nitric acid)

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H2SO4 (sulfuric acid) 174

•

HI (hydrogen iodide or hydroiodic acid)

Weak acids include CH3COOH (acetic acid), HCN (hydrogen cyanide), H2SO3 (sulfurous acid), HF (hydrofluoric acid), CHOOH (formic acid), HNO2 (nitrous acid), H2PO4 (phosphoric acid) and H2S. Strong acids produce strong bases as conjugates, while weak acids produce weak conjugate bases. Strong bases will dissociate completely or nearly completely in an aqueous solution. There are strong bases, which primarily come from Group IA and Group IIA metals. Strong bases include these: •

LiOH (lithium hydroxide)

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NaOH (sodium hydroxide)

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KOH (potassium hydroxide)

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RbOH (rubidium hydroxide)

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CsOH (cesium hydroxide)

•

Ca(OH)2 (calcium hydroxide)

Weak bases include all other bases, including NH3 (ammonia), CH3NH2 (methylamine), and C5H5N (pyridine). Remember that each of these has a corresponding conjugate acid, regardless of whether it is a strong acid, weak acid, strong base, or weak base). A third theory is the Lewis theory of acids and bases, which takes into account the electron pair donors and acceptors. This makes acids acceptors of an electron pair and bases donors of electron pairs. This is a theory and definition that does not mention anything about the hydrogen atom or proton. It does account for this type of reaction: Ammonia (NH3) and boron trifluoride (BF3) will be in equilibrium with H2NBF3. This does not involve an actual transfer of a hydrogen atom, making this a Lewis acid and Lewis base transaction. NH3 (ammonia) has a single pair of electrons, while BF3 (boron trifluoride) does not have a complete octet. The combination of the two allows for a complete octet for this combined molecule. Boron has six electrons around it and has two more that it can hold. This makes BF3 a Lewis acid that will accept the pair of

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electrons from the nitrogen ion in NH3. What this involves regarding the H2NBF3 molecule is a covalent bond between nitrogen and boron. The Lewis theory of acids and bases applies best to gaseous and solid acids and bases, which act somewhat differently from liquid acids and bases. It involves two substances (an electron donor and an electron acceptor) that make a coordinate covalent bond, making a substance that is neither an acid or nor a base but that is relatively stable covalently. This covalently bonded structure is known as an “adduct”. Lewis acids are called “electrophiles” and Lewis bases are called “nucleophiles” because they donate an electron pair to a nucleus. Now, this becomes more crucial to understand. Under the Lewis system, all cations are considered Lewis acids because they can accept electrons. This includes Cu2+, Fe2+, and Fe3+. Any atom, molecule or ion that has an incomplete octet of electrons are electrophiles and will be Lewis acids, including AlF3 and BF3. Any molecule that has more than 8 valence shell electrons can be electrophiles and will be Lewis acids, including SiBr4 and SiF4. Molecules that have multiple bonds (like CO2 and SO2) are Lewis acids because they are electron acceptors or electrophiles. On the other hand, Lewis bases will donate an electron pair. As nucleophiles, they “crave” the positive nucleus. Any atom, molecule, or ion with a lone pair of electrons can be a Lewis base. These will have a “highest occupied molecular orbital” or HOMO that interact with the LUMO (lowest unoccupied molecular orbital” of the Lewis acid. While both acids and bases have HOMOs and LUMOs, only the HOMO is under consideration for the bases and only the LUMO is considered for the acids . According to quantum theory, the acids are at a lower energy level because they have vacant orbitals and bases are at a higher energy level because they have

a pair of electrons to donate.

An amphoteric molecule is any molecule that can act as a base and as an acid. This is true of water which, as you have seen will have the ability to donate a hydrogen ion or donate a hydroxyl ion. It will react with NH3 (ammonia) to make a hydroxyl ion plus NH4+. It will interact with HCl to make a chloride ion and H3O+. In such cases, this is both an acid and base. Another amphoteric molecule is Al(OH)3 because it can act as an

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acid and a base, depending on the environment it is in. Figure 54 describes amphoteric molecules:

Figure 54.

PH SCALE The pH is very important when it comes to recognizing whether or not a substance is an acid or base in solution chemistry. This is based on the concentration of hydrogen ions in a solution. The basic formula is that the pH equals the negative log of the hydrogen ion concentration. It is generally thought of as a scale from 1 to 14 but it is somewhat incorrect, as you’ll soon see. A pH of 7 is directly neutral, while a pH of 1-7 is acidic and a pH of 7-14 is considered basic. Figure 55 shows the pH scale:

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Figure 55.

Because water is amphoteric, it can be considered acidic or basic. The fact of the matter is that two H2O molecules can make one H3O+ molecule and one OH- molecule. This is referred to as autoionization or self-ionization with equal concentrations of H3O+ and OH-. This means that the molarity of each of these ions is 1 x 10-7. Remember that the Kw of water is the multiple of each of these two ions so that it is 1 x 10-14. This number should stick out because the maximum of the pH scale is 14. If an acid is added to water, the equilibrium shifts to a decrease in OH- ions. If a base is added to water, the equilibrium shifts to the right and there are more OH- ions. Because of the Kw of water, the pKw is going to be 14. The pH becomes the negative logarithm of the molarity of hydrogen ions in water, while the pOH becomes the negative logarithm of the hydroxide ion concentration. The pH scale is considered logarithmic. This means that every increase or decrease of the pH by one is a change in concentration by tenfold of the hydrogen ion. A pH of 3 is 100 times more acidic than a pH of 5.

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NEUTRALIZATION REACTIONS A neutralization reaction happens when any acid and any base react with one another to make water and a salt. When this happens, the pH is going to be 7 if a strong acid and a strong base are neutralized; the end result will be a pH of less than 7 if the combination is a strong acid and a weak base; the end result will be a pH of more than 7 if the combination is a weak acid and a strong base. It will always be equal weights of acids and bases involved. The amount necessary will be the amount of an acid that will yield one mole of protons and the amount of a base needed to yield one mole of OH- ions. An example is when HCl (a strong acid) and NaOH (a strong base) get combined to make water and NaCl salt. The real reaction is H+ plus OH- equals water or H2O. This completely neutralizes the reaction and results in a pH of 7.

TITRATION Titration is one way to determine the concentration of an acid and a base. It is based on the idea of a titration curve that goes from an acidic pH to a basic pH, with a sharp curve. Figure 56 shows what different titration curves can look like:

Figure 56.

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In titration, you have an acid or base in a flask and slowly add the opposite to the flask, looking for when the pH makes a sharp difference. In the process, you need to know that the M1V1 = M2V2, where M stands for the molarity of the solutions and V is the volume in liters. If you know the molarity and volume of three of these things, you can determine the molarity of the unknown solution. If you don’t know the molarity of the starting solution but you know the volume and you know the volume and molarity of the second solution, you can determine the molarity of the unknown. As you can see, these titration curves are very unique in that, at a certain point in the volume/pH curve, you will get a sharp change in the pH from acidic to very basic when an acid is titrated with a base. The equivalence point will be 7.0 with strong acids and bases but will be on the basic side for weak acids and on the acidic side for weak bases. Regardless of the actual pH, this point will be marked by a sudden change in the pH with just a tiny change in volume. To do a titration, you’ll need a strong acid or a strong base to do the titration with and you’ll need to know the molarity of the strong acid or base. Once you know the volume of the solution being titrated and the titration amount that causes the sudden pH change, you will know the molarity of the unknown acid or base. You can determine other things as well, such as the pH when starting the titration, during the titration, and after the titration. So, how can you measure the pH change in a titration reaction? There are certain organic substances that can be added to the solution that will change color at a certain pH change. Phenolphthalein will change color at pH 8.3, being colorless at lower pH levels and pink at higher pH levels. There are a number of pH indicators that will change in colors over specific pH changes. An example is crystal violet that changes color at pH zero to two from yellow to blue. At the far end of the spectrum is alizarin yellow R that changes color between pH 10 and 12. Figure 57 shows litmus paper and the color changes seen at different pH levels:

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Figure 57.

BUFFERS Buffers are weak mixtures of weak acids and its conjugate base or a mixture of a weak base and its conjugate acid. These are helpful in biological systems because they resist changes in pH when small amounts of strong acids or bases are added. Acetic acid and sodium acetate together are a buffer solution that will resist wide pH shifts. Ammonia and ammonium chloride also represent a buffer system. Buffers are buffers because they contain weak acids or bases and their salts. They can “soak up” or combine with hydroxide ions or hydronium (H3O+) ions to keep the pH roughly the same within certain limits. In biological systems, there is a phosphate buffer system, protein buffer systems, and the carbonic acid system—each of which keeps the pH steady in biological systems. It doesn’t matter whether OH- is added or hydronium ion is added, there will be the prevention of a pH change. Buffer solutions are not unlimited in their capacity to maintain a constant and steady pH. If there is an acetic acid/acetate buffer system and a lot of base is added, the buffering capacity will be exhausted, eve if not all of the buffer is exhausted. As the buffering action is depleted, the buffering action will not be as effective. The more buffering acids/bases and salts that are added to the system, the better the buffering capacity will be.

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A good buffer mixture will have about equal concentrations of the acid or base plus the corresponding salt. When about 90 percent of the buffer pair has been used up, the buffer becomes much less effective. For example, with acetic acid buffer systems, the initial pH will be about 4.75. When there is a situation in which the acetic acid concentration is about 10 percent of the acetate concentration, the pH will change by 1 pH unit. Past that, the pH will increase dramatically. This means that it will work until 90 percent of the mixture is sodium acetate and not acetic acid. When the pH of the system is less than 7, weak acids and salts are better buffers. When the pH of the system is greater than 7, weak bases and their salts are better buffers. In blood systems, when the pH is around seven, the carbonic acid (H2CO3) and bicarbonate ion (HCO3-) system is prominent. In such cases, the pH is about 7.4 with pH changes greater than 0.4 being potentially fatal.

HYDROLYSIS Remember that a salt is made when an acid and base are mixed together and when hydroxide ions and hydrogen ions mix to form water. The bystander in all of this is the salt ions and salt solution. When weak acids and weak bases react, the relative strength of the conjugated acid-base pair determines the pH of the solution. A salt solution can be neutral, acidic, or basic by itself because of the nature of its ions. In other words, there are neutral salt ions, acidic ions, and basic ions. This is in keeping with the idea that, according to the Lewis theory, molecules not directly having H+ or OH- ions will have acidic or basic potential because they can donate or need electron pairs. Neutral cations include sodium, potassium, rubidium, cesium, magnesium, strontium, barium, and calcium. Neutral anions include chloride, bromide, iodide, and perchlorate. Acidic cations include ammonium, lead (II), and aluminum (III). Acidic anions include HSO4- and H2PO4-. Basic anions include (HPO2)4-, (PO3)4-, F-. NO2-, CN-, and HCO3-. There are no basic cations. Any salt made from a strong acid and a weak base will be an acidic salt. An example of this is ammonium chloride (made from HCl and NH4OH). The ammonium chloride 182

leaves behind an acid cation in the form of NH4+ that itself interacts with water in an activity called “hydrolysis”, which leads to ammonia plus H3O+. A basic salt is formed between a weak acid and a strong base. An example of this is sodium acetate, which is made from the weak acid acetic acid and the strong base sodium hydroxide. Acetate is itself a basic anion that undergoes hydrolysis to make acetic acid plus hydroxide ion (and leads to a basic situation). As mentioned before, the salt of a weak acid and a weak base might be neutral, acidic, or basic, depending on the K constant. If the K constant of the acidic cation is greater than that of the basic anion, this will be an acidic salt. If the K constant of the acidic cation is less than that of the basic anion, the salt will be a basic salt. If the K constant of the acid is the same as the K constant of the base, the resultant salt will be a neutral salt.

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KEY TAKEAWAYS •

There are three key definitions of acids and bases: Arrhenius, Bronsted-Lowry, and Lewis theories.

•

There are strong acids and strong bases that completely dissociate in water as well as weak acids and bases that do not completely dissociate.

•

The pH of a solution is the negative log of the hydrogen ion concentration. The pH scale goes from 1 to 14 and is a logarithmic scale.

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Titration involves slowly neutralizing an acid or base until the pH changes abruptly from acidic to basic or vice versa.

•

Neutralization involves combining an acid and a base to get a salt. The salt itself can be acidic basic or neutral in pH.

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QUIZ 1. In the Arrhenius definition of acids, what does the hydrogen ion become when an acid dissociates in solution? a. H2O b. H3O+ c. H+ d. OHAnswer: b. The H+ ion in an acid does not remain stable so that it combines with water to make a hydronium ion or H3O+ ion. 2. Which of the following is a base? a. HSO4 b. HCl c. NH3 d. KCl Answer: c. NH3 is a weak base because it is a proton acceptor, becoming NH4+. 3. What is the equilibrium constant for water? a. 1 b. 1 x 1014 c. 1 x 10-7 d. 1 x 10-14 Answer: d. The equilibrium constant for water is 1 x 10-14, which is an important number because it is the product of the base constant times the acid constant for an acid and its conjugate base.

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4. Which is considered a strong acid? a. Perchloric acid b. Acetic acid c. Ammonium d. Formic acid Answer: a. While it is difficult to remember all of these acids and bases, it is a good idea to remember as many of the strong bases and strong acids as possible. Each of these is a weak acid except for perchloric acid, which is a strong acid. 5. In understanding Lewis acids and bases, when these are combined, what do they form? a. Bronsted-Lowry compounds b. Nucleophiles c. Adducts d. Electrophiles Answer: c. In understanding the Lewis acid-base system, there are nucleophiles and electrophiles, which combine to make adducts, which are covalently bonded molecules that comprise these acids and bases combined. 6. The molecule Al(OH)3 or aluminum (III) hydroxide is an amphoteric molecule. What does this mean technically? a. It is a strong acid b. It is a strong base c. It is a weak base d. It can be an acid or a base Answer: d. Aluminum (III) hydroxide is considered amphoteric because it has the ability to act as an acid or a base depending on the environment it is placed in, based on the Lewis theory of acids and bases.

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7. Which of the following molecules is a Bronsted-Lowry base but not an Arrhenius base? a. NH3 b. NaOH c. Ca(OH)2 d. KOH Answer: a. While each of these is a base, only NH3 does not directly give an OH- ion to the solution, making it a Bronsted-Lowry base but not an Arrhenius base. 8. A weak acid and a weak base are involved in a neutralization reaction. What will the pH be at the time of neutralization? a. 7 b. Less than 7 c. Greater than 7 d. It depends on the Ka and Kb Answer: d. This pH will be different depending on these K constants. If the Ka is greater, the pH will be less than 7; if the Kb is greater, the pH will be greater than 7. 9. Which is not an example of a buffer system? a. Ammonia b. Acetate c. Sulfuric acid/sodium sulfate d. Phosphate Answer: c. With the exception of sulfuric acid/sodium sulfate, each of these is a buffer system because it is a system involving a weak acid or weak base plus the corresponding salt. Sulfuric acid and sodium sulfate together involve strong acids and salt solutions, which would not make a very good buffer system.

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10. If the H+ concentration is 0.00001 M/l, what is the pH? a. 3 b. 4 c. 5 d. 6 Answer: c. The pH is the negative log of the hydrogen concentration. A concentration of 0.00001 M/l is 10 to the -5 power, the negative log of which is 5.

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CHAPTER ELEVEN: ELECTROCHEMISTRY The focus of this chapter is electrochemistry, or the study of electricity as it applies to chemical reactions. Faraday’s law, which states that the amount of substance produced at an electrode is directly proportional to the quantity of charge flowing through an electrochemical cell, is discussed. The basics of electricity in chemical reactions are important to understand because electricity plays a role in many chemical interactions. The different properties of electrical cells and what makes them up is also a part of this chapter’s discussion.

ELECTROCHEMISTRY BASICS Thus far, you have learned about chemistry as it relates to micro-electrical situations because all chemical reactions ultimately involve electrons and, in some cases, the shifting of electrons from one element to another. Redox reactions are a part of electrochemistry in particular because they directly involve the movement of electrons. Remember that, when a substance loses an electron, its oxidation state is increased while, when a substance gains an electron, its oxidation state is decreased. Consider the redox reaction where H2 plus F2 go to 2HF. This actually involves two half-reactions, in which H2 goes to 2H+ plus 2 negative electrons and in which F2 plus 2 negative reactions go to 2 F-. The first is an oxidation reaction and the second is a reduction reaction. While this involves a simple chemical reaction, it is actually electrochemistry at work.

PROPERTIES OF ELECTRIC CELLS In any electrochemical process, electrons can flow from one chemical substance to another, simply through a balance oxidation-reduction reaction. In an electrical cell, a redox reaction will occur when electrons are transferred from one oxidized substance to one reduced substance.

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This leads to two electrochemical definitions: 1) the reductant is the substance losing electrons (the same as the “reducing agent”) and 2) the oxidant is the substance gaining electrons (the same as the “oxidizing agent”). The potential energy is determined by the potential difference between the valence electrons in atoms of different elements. Remember that orbitals in one atom will have a higher energy status than orbitals in another atom so it makes for a “potential difference”. Remember, too, that it is impossible to have reduction without oxidations , especially in an electrochemical situation. Redox reactions will involve two half-reactions—one with reduction and one with oxidation. When zinc metal interacts with bromide gas (both at zero oxidation numbers), they will make zinc (II) ions and bromine (2-) ions. Zinc gains an oxidation number by +2 and two bromine atoms lose an oxidation number by -2. The reduction half reaction is the bromine gas going to Br2- with the addition of two electrons, while the oxidation half reaction is the gaining of two electrons when zinc metal goes to Zn2+ ion. In this case, the reductant is zinc, and the oxidant is bromine gas. The half-reactions are just written this way because they show what happens to the reactants but it doesn’t, of course, go this way in reality. Figure 58 describes this redox reaction:

Figure 58

This is always going to be balanced with respect to the electrons gained and lost. So far, we have assumed that these reactants are next to one another in solution. In the acidbase reactions we discussed in the previous chapter, the acid and base are together in solution. With redox reactions, however, it is completely possible to physically separate the oxidation and reduction half reactions as long as there is a circuit of “electricity”

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connecting the two half reactions, such as exists with a wire between them. Electrons will flow from the reductant to the oxidant through the electrical reaction. This leads to the definition of an electrochemical cell. An electrochemical cell is an apparatus used to generate electricity from a spontaneous redox reaction or one that uses electricity to drive a nonspontaneous redox reaction by separating the oxidation part of the reaction from the reduction part of the reaction into two separate physical spaces. There are two types of electrochemical cells. The first is the galvanic cell, discovered in the late 1700s. It is also called a voltaic cell that uses the energy released by a spontaneous redox reaction. In a sense, it generates electricity. The other type of cell is called an electrolytic cell. This consumes electricity from an electrical source, using it to drive a nonspontaneous redox reaction. Both of these cells use the concept of the Gibbs free energy. The Gibbs free energy, as you’ll remember, is the thermodynamic potential that can be used to calculate the maximum of reversible work that may be performed by a system at a constant temperature and pressure. It is related to the enthalpy because it is the enthalpy minus the product of the temperature and entropy. In a galvanic cell, the change in Gibbs free energy is less than zero so electricity is generated. In an electrolytic cell the change in Gibbs free energy is positive so energy is necessary. Figure 59 is what a galvanic cell looks like:

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Figure 59.

Figure 60 shows what an electrolytic cell looks like as it involves a battery that adds energy to the system in order to drive the reaction:

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Figure 60.

The two types of cells will have two electrodes that consist of solid metals connected to an external circuit, which provides an electrical connection between two parts of the system. Oxidation (the half-reaction) occurs at the anode, while the reduction halfreaction occurs at the cathode. Electrons flow from the anode to the cathode. There is also an electrolyte solution that allows ions to transfer between the electrode apartments in order to maintain electrical neutrality in the system. Another slight difference between galvanic cells and electrolytic cells is that, in galvanic cells, there is a porous membrane between the two halves that contain the electrolyte solution, while in electrolytic cells, there is no porous membrane and the electrolyte solution mixes freely. In both cases, electrical neutrality is maintained by the electrolyte in the system. So, how do galvanic cells work? Take the example of zinc solid plus cupric ions (Copper II ions), which interact to make Zinc (II) ions and copper solid. A zinc rod is placed in copper II sulfate solution). Zinc gets dissolved to make zinc sulfate and copper solid is made. In a regular reaction, this gives off heat because it is exothermic and happens spontaneously. This reaction happens extremely quickly. Now, place the reaction in a galvanic cell. A copper strip Is placed into a beaker with a molar solution of copper II ions and a zinc strip is inserted into a different beaker with a

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molar solution of zinc (II) ions. The strips are electrodes, connected by a wire and a salt bridge is created between the two beakers (which is a U-shaped tube containing a gelled or concentrated liquid electrolyte). Figure 61 shows the salt bridge in a voltaic (galvanic) cell:

Figure 61.

The ions in the salt bridge are chosen specifically because they won’t precipitate out of solution and won’t be oxidized or reduced themselves. This is why Na+ or K+, Cl-, NO3-, or SO42- are often used in salt bridges. The ions in the salt bridge do not have to be the same as those in the redox reaction; they are just there to maintain electrical neutrality.

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When the circuit is closed, the reaction is spontaneous. Zinc goes to Zn2+ in the anode (because it gives away electrons) and Cu2+ goes to Cu metal in the cathode (because it receives electrons). The zinc strip dissolves and the copper strip gains mass. The salt bridge maintains electrical neutrality and the whole thing will allow electrons to pass through a wire between the two beakers. This “electricity” can be measured in a voltmeter. The delta G or change in Gibbs free energy is negative. The salt bridge maintains the neutrality by allowing equal charges of ions to pass to one beaker or the other so that electrical neutrality is maintained throughout the system. Again, the actual identity of the electrolytes in the salt bridge isn’t important as long as the salt doesn’t interfere with the reaction in any way. Without the salt bridge, the reaction would not proceed very long. The voltage or potential of the cell to do work is the energy necessary to move a charged particle in an electric field. The anode will be negatively charged and the cathode will be positively charged in this kind of cell (because it attracts electrons). If the reaction goes forward, the cathode will be the oxidized metal and the anode will be the reduced metal. The electrode does not have to be a true part of the redox reaction. It can involve something like platinum, which conducts electricity but doesn’t get oxidized or reduced. Cell phase diagrams are used to identify what is going on in any electrical cell. Figure 62 shows the notation used to describe a voltaic or galvanic cell:

Figure 62.

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The junction potential is the resistance to the flow of charge at a boundary, which can affect the way the reaction smoothly goes. This can be reduced by having just one electrolyte solution and two electrodes without a salt bridge. In such cases, there is no double line in the phase diagram because there is no actual barrier situation going on in the cell. These aren’t as helpful in producing electricity because the reaction happens too quickly.

FARADAY’S LAW It is possible to measure the rate at which charges are transferred in any electrical situation via an ammeter. What an ammeter measures the current in Amperes (also referred to as Amps or A). Unlike a voltmeter, ammeters have electrons pass through the devices in order to measure or “clock” them as they go by. The total amount of electrical charge that has passed through the circuit can be calculated by the simple relationship of: charge equals current over time. This is the equivalent of Coulombs, which is determined as amp-seconds. These types of calculations allow us to connect the reaction stoichiometry to actual electrical measurements. This leads us to Faraday’s law, which will be explained in a minute. A voltmeter can be placed across the electrical conduit or “wire” while the ammeter is part of the circuit. Electrons technically do not pass through the voltmeter, simply continuing along the wire. The voltmeter measures the difference in electric potential between two points of a conducting wire when an electric current of one ampere dissipates one watt of power between those points. One volt also equals to the potential difference between two parallel, infinite planes spaced 1 meter apart that create an electric field of 1 newton per coulomb. Its units are volts or amperes multiplied by the number of ohms of resistance (current times resistance). It accounts for the resistance in the wire. Both the ammeter and the voltmeter can be said to be polarized. This means that they have negative and positive terminals marked on them because the electrons can only go

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in one direction. This is especially true in direct current (DC) situations, in which electrons go from one place to another in an electrochemical cell. According to Faraday’s law of electrolysis, the amount of substance made at each electrode in an electrochemical cell is directly proportional to the quantity of charge that flows through the cell. Because substances have different oxidation/reduction number increases or decreases in terms of the electrons per atom or ion that are gained or lost, the amount of substances produced will not be produced in the same molar amounts. When these ratios are factored into the equation, the law will be correct in all situations.

NERNST EQUATION As you have determined, the energy of a chemical system is what drives the electrons to move such that the driving force gives rise to the cell potential of the galvanic cell. This energy is also highly related to the chemical equilibrium of the reaction. The Nernst equation helps tie this all together. Remember that energy can take on many forms. It can be mechanical, radiation (photons), chemical energy, nuclear energy, thermal energy, and electrical energy. These can change back and forth so that, as you have seen, chemical energy can be used to create electrical energy. It is not always a perfect transformation, which is why batteries will heat up when the chemical reaction is going to make electrical energy. According to the Nernst equation, we must assume that the energy of the cell is equal to the potential energy of reduction minus the potential energy of oxidation. In order to drive the reaction forward, the change in energy must be a negative number. The change in energy or delta G (which means Gibbs free energy) is equal to the negative product of the number of electrons transferred, Faraday’s constant (which is 96,500 calories per mole), and the potential difference across the cell. Figure 63 shows this equation:

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Figure 63.

The general Nernst equation relates to the concentrations of the different reactants and end products. It also depends on the temperature in Kelvin. According to the standard conditions of 25 degrees Celsius (298 degrees kelvin), the Nernst equation becomes what is depicted in figure 64:

Figure 64.

Interestingly, the Nernst equation indicates that a battery can be built by simply using the same reactants and products but having a higher concentration in one side of the cell and a lower concentration of reactants and products on the opposite side of the cell. This is called a “concentration cell”. In the equation, it predicts whether or not a reaction will or will not happen, based primarily on the change in potential energy over the course of the reaction.

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ELECTROMOTIVE FORCE The electromotive force or EMF is the maximal potential difference between the two electrodes in a voltaic cell or galvanic cell. This is directly related to the tendency of any element, molecule, or ion to gain or lose electrons. An example is the common cell involving zinc, zinc ions, copper, and copper ions, which is driven toward the making of copper from copper ions (reduction) and the loss of zinc into zinc ions (oxidation). This reaction is driven forward. The EMF of this reaction is 1.100 Volts under standard conditions, which means standard temperature (25 degrees Celsius) and 1 Molar solution (being an ideal solution). This has led to the standard electromotive force or the standard cell potential for this type of cell. So, how can this be determined mathematically? You need to review the reduction potential of Zinc and Copper. One of the two metals will have a greater propensity to be reduced than another, which will drive the reaction in a forward direction so that electrons will go across the wire in order to reduce the metal that “wants” the electrons more and is more likely to be reduced. This difference in reduction potential will equal the EMF, which as you have just learned is 1.10 volts under standard conditions between copper and zinc. There is a table that can be utilized to determine the standard reduction potential of a given metal atom. This will be different for different types of metals. Using the table, you can determine the EMF for the cell and can determine if putting the two metals into electrodes to make an electric cell will actually drive the reaction forward and create electrical potential, making electricity happen. If the cell potential is negative, the reaction is reversed.

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KEY TAKEAWAYS •

Electrochemistry involves the way in which certain chemical reactions (redox reactions) are used to generate electricity.

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Galvanic or voltaic cells will generate electricity, while electrolytic cells will require electricity in order to drive the reaction.

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Faraday’s law relates to the amount of substance produced at an electrode being directly related to the charge released by the other electrode.

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The electric potential of a cell is related to the difference in the reduction potentials of the different metals involved in the cell.

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QUIZ 1. What does the ammeter measure in units? a. Charge b. Amperes c. Joules d. Coulombs Answer: b. The ammeter measures amperes or the electrical charge that passes through the system per second. 2. What are the units for Coulombs? a. Joules per second b. Amp-seconds c. Volts d. Ohms Answer: b. The units for coulombs are amp-seconds, which can be calculated using an ammeter. 3. In a reaction where zinc metal plus bromine gas goes to zinc 2+ ions and Bromine 2- ions, what is the reductant in this situation? a. Zinc b. Bromine gas c. Zinc ions d. Bromine ions Answer: a. The zinc is the reductant because it reduces something else, getting oxidized at the same time.

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4. In a reaction where Zinc metal plus iodine goes to zinc (II) ions plus 2 iodide ions, what is the oxidant? a. Zinc b. Iodine c. Zinc ions d. Iodide ions Answer: b. The iodine is the oxidant in the reaction because it gets reduced but oxidizes something else. 5. What statement is the same for both a galvanic cell and an electrolytic cell? a. It involves the spontaneous passage of electrons through a wire. b. It involves separate reduction and oxidation reactions. c. It involves a spontaneous redox reaction. d. Electricity is generated by the cell. Answer: b. The only thing that is true for both the galvanic cell and an electrolytic cell is that they both involve separate oxidation and reduction reactions. The rest of the statements cannot be applied to both types of cells. 6. What is the main difference between a galvanic cell and an electrolytic cell? a. The electrolytic cell involves electrolytes and the galvanic cell does not. b. The galvanic cell involves metals and electrolytic cells do not. c. The galvanic cell involves redox reactions and electrolytic cells do not. d. The galvanic cell generates electricity, while the electrolytic cells use up electricity. Answer: d. The main difference is that the galvanic cell generates electricity while the electrolytic cell uses up electricity. They both involve redox reactions and many involve metals and metal ions.

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7. What term does not equate with the term “galvanic cell”? a. Electric cell b. Electrolytic cell c. Voltaic cell d. Redox cell Answer: b. The electrolytic cell is completely different from the galvanic cell. The other terms, “electric cell”, “redox cell”, and “voltaic cell” are the same thing as the galvanic cell. 8. What would be the least appropriate ion in a salt bridge in a voltaic cell? a. K+ b. Na+ c. NO3d. Zn2+ Answer: d. The salt ions in the salt bridge must not precipitate out of solution and must not be able to be oxidized or reduced by themselves. Zinc is reducible so it would not be used easily in a voltaic cell. 9. What is the junction potential in an electrical cell? a. It is the electrical resistance in the electrical wire. b. It is the potential difference between the two beakers in a cell. c. It is the resistance to the flow of charge at the barrier or the “salt bridge”. d. It is the difference in charge from the cathode to the anode. Answer: c. The junction potential is the resistance to the flow of charge at the barrier or the salt bridge between the beakers.

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10. According to Faraday’s law, under what circumstances would the amount of a metal gained at an electrode in an electric cell not be equal to the amount of metal lost at the other electrode? a. When the concentrations of the electrolytes are different in each half of the cell b. When the charge on one atom in the electrode does not equal the charge on the other atom’s electrode. c. When the molar mass of one is greater than the molar mass of the other. d. It will never be an unequal mass according to this law. Answer: b. When the charge on one atom is not equal to the charge on the other atom, the amount of metal would not be the same. The law states that the amount of substance made at each electrode in an electrochemical cell is directly proportional to the quantity of charge that flows through the cell. The law will be accurate if the stoichiometry of the equation is taken into account.

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CHAPTER TWELVE: ORGANIC CHEMISTRY AND BIOCHEMISTRY Organic chemistry and biochemistry are courses of their own; however, they are covered together in this chapter because they are inter-related and important aspects of the study of chemistry. Organic chemistry is carbon-based chemistry and involves a variety of different types of molecules and reactions typically seen in living things. The nature of organic molecules is discussed as well as the functional groups that shape organic molecules and their behavior. Finally, biochemistry and its principles are covered as these types of molecules, too, follow the basic chemical principles already laid out in this course.

NOMENCLATURE Before you know anything about organic chemistry, you need to understand how organic molecules are named. There are huge numbers of organic molecules and it would be impossible to name each of them separately as is done with the periodic table of the elements. While most of organic chemistry comes from carbon, oxygen, hydrogen, and (to a lesser extent) nitrogen, there are many ways these can combine together. This leads to the need for specific guidelines for naming organic molecules. Common organic chemicals were given names before any systematic naming system was developed. You need to know these basic structures: •

Methane—CH4

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Butane—C2H10

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Acetone—CH3COCH3

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Toluene—CH3C6H5

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Acetylene—C2H2

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Ethyl alcohol—C2H5OH

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A rational nomenclature is clearly necessary based on the number of carbon atoms and the different structures involved. There are specific functional groups that we will talk about later, which help to determine the reactivity of the molecule. This is where the IUPAC system has become so helpful in naming organic compounds. You should be able to write the formula based on its IUPAC nomenclature name. The IUPAC name will have three essential features: •

A root or base indicating a ring of carbon atoms or major chain of carbon atoms

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A suffix that tells the type of functional groups involved in the molecule

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Names of those groups that aren’t hydrogen that complete the molecule’s structure

ALKANES Alkanes are hydrocarbons that have no double or triple bond functional groups attached to the carbon atoms. Remember that carbon has the ability to bind to four other molecules, making it very versatile. With four spots, a carbon atom can bind with hydrogen or another carbon atom in short or lengthy chains. They can be arranged in rings or in chains. For this reason, these are also called “cycloalkanes”. These are part of a larger group of “aliphatic compounds”, which are called aliphatic because they do not have aromatic rings (which will be explained soon). Alkanes can be named as ending with “-ane”. You should memorize the basic chemical formula for all of the alkanes up to decane because they are the basis of many IUPAC nomenclature names. They are listed solely by the number of carbon atoms associated with them. The following is a list of the unbranched alkanes by carbon number: •

One—methane

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Two—ethane

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Three—propane

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Four—butane

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Five—pentane 206

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Six—hexane

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Seven—heptane

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Eight—octane

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Nine—nonane

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Ten—decane

All of these are based on the formula Cn + H (2n+2), where n is the number identified by the prefix. For example, with simple methane, the structure will be CH4. This is the highest possible hydrocarbon ratio, representing a fully hydrogenated or fully saturated carbon structure. The problem comes when there are alkane isomers, which involve a branched chain of carbon atoms that are also fully saturated. This starts with any molecule bigger than butane (C4H10). How do you name all the possible branched possibilities? This involves what are called “alkyl groups”, which are named groupings of carbon atoms and hydrogen atoms. They are named in several ways—a few of which are named here: •

Methyl—CH3-

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Ethyl—C2H5-

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Propyl—CH3CH2CH2-

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Butyl—C5H9-

So, what are the IUPAC rules for naming these alkanes? These are the basic rules: •

Find the longest continuous carbon chain and name it.

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Find and name the groups attached to the chain.

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Number the chain in order, starting at the end nearest a substituent group.

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Designate the location of each substituent group by number and name.

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Name the groups in alphabetical order, using the full name (the prefixes di, tri, tetra, etcetera, aren’t used in the alphabetizing scheme).

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An example would be 2-methylpentane, which is a pentane molecule with a methyl group at the second carbon molecule. Halogens can be added like fluoro-, chloro-, bromo-, and iodo- for the halogen side-groups. When these are added, the compound is called an alkyl halide. C2H5Cl can be ethyl chloride or chloroethane. It means the same thing. Of course, side groups can contain halogens, too, such as trichloromethyl (CCl3-) and bromomethyl (BrCH2).

CYCLOALKANES These involve at least one carbon ring associated with it. The simplest one is a single unsubstituted carbon ring with the general formula of the ring being CnH2n because carbon is bonded to itself with a single bond in a ring. The simplest one is cyclopropane, which contains 3 carbons in a “triangular ring” that has six hydrogen atoms associated with it. These are named like the alkanes, with cyclobutane, cyclopentane, and cyclohexane named for the number of carbon atoms added to the ring. The rules are similar: •

Name the ring first unless the ring is itself a substituent group on a larger alkane.

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Start the numbering with the first substituted carbon atom.

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Go around the ring so that the next substituted group has the lowest possible number.

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If there are a lot of substituted groups, list them in alphabetical order.

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More than one ring involves a polycyclic compound, which can be bicyclic, tricyclic, etcetera.

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Rings can be separated, have one common carbon atom, one common bond, or two common atoms.

Spiro rings have one common atom; fused rings have a common bond, and bridged rings have two common atoms. If no carbon atoms are in common, this is called an isolated ring system. Figure 65 shows some of the cycloalkanes:

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Figure 65.

ALKENES AND ALKYNES This is a situation where there are double or triple bonds. These are unsaturated hydrocarbons. An alkane is saturated. An alkene has at least one double bond; an alkyne has a triple carbon bond. What does this mean? It means that a molecule identified as C5H8 can have more than a dozen different molecule shapes associated with it. According to the rules, the -ene suffix can be an alkene or cycloalkene. The longest chain must include both carbon atoms of the double bonded chain. Numbering starts from the end nearest to a double bonded carbon atom. The double-bonded carbons are labelled number 1 and number 2. The vinyl group is the simplest double bond with C2H3-.

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Alkynes end in the -yne suffix. It can be an alkyne and cycloalkyne compound. The longest chain that contains the triple bond should be used as the root name. Numbering starts at the nearest triple bond carbon atom. The simplest one is the ethynyl group or the C2H group.

BENZENE DERIVATIVES These can be more difficult to name. There are some that use the benzene ring and just one substitution. Others have more substitutions. Figure 66 shows what the benzene ring looks like:

Figure 66.

Examples include chlorobenzene, bromobenzene, nitrobenzene, and more difficult ones, like toluene (which is a methyl group attached to a benzene ring). Two common ones are the phenyl group, which is abbreviated with a Ph and the benzyl group, which is abbreviated with a Bn. Basically, these need to be memorized as they are not IUPAC rules. Figure 67 shows a number of these benzene ring-affiliated molecules:

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Figure 67.

ISOMERISM As you have seen, different combinations of carbon and hydrogen atoms can have different basic structures. With isomerism, you can see that there are more structures than you can imagine. An isomer is a molecule that has the same number of atoms but different structural or spatial arrangements of the atoms within the molecule. This leads to more than ten million organic compounds. There are five different types of isomerism—many more than the simple structural isomerism we have already discussed (in which there is a structural difference between two molecules. The two broad groups of isomerism are structural isomerism and stereoisomers. Let’s look at the different types of isomerism.

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STRUCTURAL ISOMERISM There are three main structural isomers, which are chain isomers, position isomers, and functional group isomers. Structural isomerism can lead to a diverse number of these types of isomers so that, if there are 40 carbon atoms to a molecule, there would be about 62 billion possible structural isomers. In chain isomers, the main carbon skeleton is not the same between the isomers. This can involve branched chains or continuous chains of carbon atoms. Figure 68 shows what structural isomers look like:

Figure 68.

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Position isomers are based on the position of a functional group. The example in figure 68 is an example of position isomers. The name is going to be slightly different because of the position of a methyl group, an ethyl group, etcetera. Functional isomers differ according to the type of functional group added to the chain. The chain can go from an alkane, for example, to a cycloalkane with the same number of atoms. There can be different functional groups that make up the different structures. With functional isomers, the structure differs only in the type of functional group added to the molecule. Stereoisomerism There are two types of stereoisomerism: geometric isomerism and optical isomerism. These aren’t rearrangements of molecules in 2 dimensions; rather, it is threedimensional changes in molecules in space. Because they look the same in 2D, they are more difficult to describe. Figure 69 shows what a stereoisomer of an organic substance looks like:

Figure 69.

Geometric isomerism is actually a term that isn’t used in the IUPAC system. Instead, it is referred to as cis-trans isomerism. It most commonly involves carbon double bonding. These bonds are different from single carbon bonds because they cannot rotate as freely as single bonds. This allows two different molecules to be made, depending on the side chains and functional groups added to them. A Cis-isomer or Z-isomer has the functional groups on the same side of the double bond, while the trans-isomer or Eisomer has the functional group on the opposite side of the double bond. This is better pictured in figure 70: 213

Figure 70.

Optical isomers are what are described in figure 71. These come in pairs and look like mirror images in a sense. They aren’t completely superimposable but are more like the left hand and the right hand, which aren’t directly mirror images of one another. These are lettered as S-isomers (meaning “left” in Latin or “sinister”) or R-isomers (meaning “right” in Latin or “rectus”). The letter is put in front of the name in order to describe which one it is. They are also referred to as enantiomers. Figure 71 shows the Senantiomer and R-enantiomer of a simple carbon molecule:

Figure 71.

Isomers of all kinds are important to chemistry because they have different chemical and physical properties. This applies even to enantiomers, which have very similar 214

appearances. An example is the compound called “carvone” found as an R-enantiomer in mint leaves, contributing to its unique smell. The S-enantiomer is the compound found in caraway seeds—clearly having a different odor. Other molecules can be drugs that have different enantiomers that react very differently in the body, depending on their stereochemistry. Thalidomide can have an Senantiomer and an R-enantiomer. The R-enantiomer is safe but the S-enantiomer is teratogenic. It turns out that the embryo can switch the enantiomer in utero, causing serious birth defects. Figure 72 shows the molecule of teratogen. The enantiomer happens within the two rings as shown:

Figure 72.

FUNCTIONAL GROUPS We’ve talked about functional groups. These are add-on molecular structures that markedly change the character of the molecule. A functional group can be a single atom, like Chlorine and Bromine or a group of atoms, such as -COOH. The following list describes the different names of chemical compounds based on their functional groups (some of which we’ve already talked about): •

Alkanes—RH (where R is the functional group)

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Alkenes—RR’C-CR’’R’’’

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Alkynes—RCR’

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Arenes—those that have an aromatic side chain

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Alkyl halides—R plus a halogen

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Aryl halides—a benzene-like compound and a halogen

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•

Alcohol—anything with an OH as a functional group

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Phenols—an aromatic compound with an OH as a functional group

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Ethers—two side chains connected by an oxygen molecule

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Aldehydes—A side chain that has a CHO attached to it

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Ketone—A side chain that has a CO attached to it

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Carboxylic acid—a chain that has a COOH attached to it

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Ester—two side chains that have a CO2 between them

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Amide—two side chains that have a CONH between them

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Amine—a chain that has an NH2 attached to it

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Nitrile—a chain that has an CN attached to it

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Nitro—an aromatic compound with an NO2 attached to it

The last four of these have nitrogen side chains attached to the main chain. The rest will have halogens, hydrocarbon, or oxygen-containing compounds as the functional groups. While these seem like a lot to remember, you’ll need to know these compounds if what you are studying is going to be organic chemistry. This means not only knowing the IUPAC nomenclature but also knowing what these side chains actually look like chemically. When dealing with aromatic compounds (benzene-like compounds), you need to know what the different functional groups mean as they attach to the benzene chain. These include ortho-, meta, and para- positions of groups. Think of benzene having numbered carbon atoms from 1 through 6. In a meta- arrangement, the side chains are next to each other in a 1,2-arrangement. Figure 73 shows the ortho-arrangement of methyl groups in o-xylene or 1,2-dimethyl benzene:

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Figure 73.

The molecules of m-xylene or meta-xylene are similar with the methyl groups at the 1,3 position, while the para configuration of p-xylene involves the 1,4 positions or “opposite” configurations of the methyl groups—referred to as opposite because the para position has methyl groups on opposite sides of the hexagonal benzene molecule. These terms are used to describe configurations of side groups on aromatic compounds for the most part.

BASICS OF BIOCHEMISTRY Biochemistry is basically organic chemistry applied to living things. All of living chemistry involves complex organic molecules that interact with one another to have the structural (anatomic) and biochemical (physiological) aspects of life. All of living processes can be explained in complex biochemical ways. In this section, we won’t cover all of biochemistry but will outline the different types of organic molecules that make up life. These mainly include proteins, carbohydrates, nucleic acids, and lipids.

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PROTEINS Proteins can be small (called oligopeptides) or large (called polypeptides or proteins). These are all based on the same set of 20 different amino acids that are connected to form peptides. Amino acids are compounds that have an amino side chain and a carboxylic acid side chain. Figure 74 illustrates the shape of an amino acid:

Figure 74.

The amino acids in proteins are referred to as alpha-amino acids because the amino group is attached to the alpha carbon of the carboxylic acid compound. About half of these amino acids in humans can be synthesized while 9 of them are essential amino acids, which must be gotten from the diet. Figure 75 shows the 20 amino acids seen in nature:

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Figure 75.

It should be noted that there are actually 22 amino acids in nature; however, the last two have been recently found and are only found in minimal amounts in living things. These two are selenocysteine and pyrrolysine. Amino acids are usually written as neutral molecules, which they are but not in the way you’d think. Most exist as zwitterions, which are negatively charged on one end and positively charged at the other. This explains the extreme reactivity of amino acids and the fact that they tend to connect to one another. This is shown in figure 76:

Figure 76.

The R chain is what makes an amino acid unique from other amino acids. There are simple side chains and very complex side chains. The side chain confers the ability of an

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amino acid to be hydrophilic (water-loving) or hydrophobic (lipid-loving)—having a major effect on the ability of the amino acid to interact with other amino acids. They can add to the charge on the amino acid and can help to create the three-dimensional shapes of proteins. You need to know that, aside from the simple amino acid glycine, the molecule will be chiral. Chiral means that there can be enantiomers, which in biochemistry are referred to as L-enantiomers and D-enantiomers (for left-sided and dextro- or right-sided enantiomers). Chemists use the same structural characteristics as glyceraldehyde to identify D-amino acids and L-amino acids because it structurally resembles amino acids. Figure 77 shows the difference between D-amino acids and L-amino acids:

Figure 77.

Nearly all amino acids are found in L-form except for a few amino acids seen in bacteria and in some antibiotics.

CARBOHYDRATES All carbohydrates have carbon, hydrogen, and oxygen atoms in them and nothing else. Examples include fiber, glycogen, and starch (which are large molecules) as well as sugars, such as glucose (a simple monosaccharide) and sucrose (a disaccharide). According to organic chemistry, carbohydrates can be described as polyhydroxy aldehydes, shown in figure 78:

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Figure 78.

Green plants undergo photosynthesis to create glucose, known as C6H12O6 from CO2 (carbon dioxide) and water using solar energy. It can be determined that it takes 686 kilocalories of solar energy to make a mole of glucose in reactions. After glucose is made, it is stored as starch molecules or in cellulose, which makes up the carbohydrate cell wall of plants. Cellulose cannot be broken down by humans for energy but starch can be broken down to be used as food. In cellular respiration, oxygen is used to break the glucose back down to CO2 and water (plus energy that is used for cellular activities). Cellular respiration is a series of reactions that start with glucose and end with CO2 and water. Humans can, of course, break down fats and amino acids, but they must be broken down into components that are fed into the cellular respiration process. Figures 79 and 80 show some of the biochemical reactions that take place in cellular respiration:

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Figure 79.

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Figure 80.

In the same way, carbohydrates can be used to make nucleic acids, lipids, and many proteins, being the base molecules that go into making these complex organic molecules.

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Simple carbohydrates are called monosaccharides and include glucose, fructose, and galactose. Disaccharides include lactose and sucrose. Figure 81 shows the basic sugars seen as monosaccharides or disaccharides in nature:

Figure 81.

NUCLEIC ACIDS Nucleic acids are the main genetic material molecules in living things. There are two types: DNA and RNA. DNA is a long chain of unique nucleotides, which are the subunits of DNA and RNA. Nucleotides can be broken down into a pentose (5-carbon) sugar, phosphoric acid, and a nitrogenous base. Figure 82 shows what the nucleotides look like:

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Figure 82.

The DNA of a typical mammal cell has about 30 billion nucleotides (per cell). The main differences between DNA and RNA include two things: 1) the pentose sugar in RNA is ribose when the pentose sugar in DNA is deoxyribose; and 2) there are a couple of changes in the bases with the thymine base exchanged for uracil in RNA. The bases are collectively referred to as pyrimidines or purines. Figure 83 shows the purines and pyrimidines in DNA and RNA:

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Figure 83.

There is a bond between the C1’ of the pentose sugar and the N1 of the pyrimidine base or the N9 of the purine base that connects the molecules together (with the loss of water in the equation). The bases have an ability to connect through hydrogen bonding with each other in the double helical shape that makes up DNA. In this structure Adenine always connects with Thymine in DNA, while Guanine always connects with Cytosine in DNA. Uracil would bind with Adenine in double-stranded RNA; however, RNA is not often seen double stranded except in viruses. You should know that, aside from being the monomer units of RNA and DNA, nucleotides have other significant properties. Adenosine triphosphate and adenosine diphosphate are ATP and ADP, respectively, which are highly important energy molecules in physiology. Other important energy molecules and coenzymes include FAD (flavin adenine dinucleotide) and NAD+ (nicotinamide adenine dinucleotide). Figure 84 shows what ATP looks like:

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Figure 84.

LIPIDS Lipids can be fatty acids, which have a carboxylic acid end and a long hydrocarbon chain. Three fatty acid molecules are attached to glycerol to form what are known as “triglycerides”. These are the storage form of fatty acids. As you may know, there are saturated fatty acids, which have all the hydrocarbons saturated with hydrogen atoms, monounsaturated fatty acids, which have the loss of hydrogen atoms because of a double bond between carbon atoms, and polyunsaturated fatty acids, which have many double bonds in the fatty acid carbon chain. Figure 85 is a sampling of some fatty acids:

Figure 85.

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Besides fatty acids, there are cholesterol-based molecules that are also called lipids. The basic structure of cholesterol is shown in figure 86:

Figure 86.

While cholesterol looks nothing like a fatty acid, it is still considered a lipid and is the backbone molecule to many lipid-soluble molecules, such as estrogen, progesterone, and testosterone—all important molecules in biochemical reactions in higher-order animals.

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KEY TAKEAWAYS •

Organic chemistry involves primarily carbon, oxygen, and hydrogen atoms in molecules, although there are nitrogen, chlorine, bromine, and other molecules as well.

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There is specific nomenclature that defines all organic molecules based on their basic structures, even though many common molecules will have common names.

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There are specific functional groups that help define an organic molecule and that control whether or not they are reactive in an organic chemical reaction.

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The main biochemical molecules are proteins, carbohydrates, nucleic acids, and fatty acids or lipids.

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QUIZ 1. Organic chemistry involves several different elements interacting in different ways. Which is the least likely element to be seen in organic molecules? a. Carbon b. Hydrogen c. Oxygen d. Nitrogen Answer: d. Organic molecules can consist of any of these but nitrogen is the least likely to be seen in an organic molecule. 2. An organic molecule defined as one that has no aromatic rings is referred to as what? a. Aliphatic b. Hydrocarbon c. Alkane d. Cycloalkane Answer: a. An aliphatic molecule is one that does not contain an aromatic ring. It does have the capability of forming a ring itself that is not aromatic. 3. What is the structure called bromomethyl? a. Br2CHb. Br3Cc. BrCH2d. Br4Answer: c. You have to assume that, unless it is specified, there will be just one bromine atom added as part of the methyl group. This leads to BrCH2- as the only real option.

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4. What type of carbon bonding is seen with an alkyne organic molecule? a. Single carbon bonds b. At least one double bond c. Two double bonds d. At least one triple bond Answer: d. An alkyne is identified as having a least one triple bond connecting two carbon atoms. The simplest one would be C2H2. 5. A cis-isomer is what type of isomer? a. Functional isomer b. Geometric isomer c. Mirror image isomer d. Position isomer Answer: b. Cis-isomers are a type of geometric isomer that is different from a trans-isomer (which is also a type of geometric isomer). 6. What is the opposite of a trans-isomer? a. R-enantiomer b. S-enantiomer c. E isomer d. Z isomer Answer: d. The Z isomer is the same thing as a cis-isomer, which is the opposite of a trans isomer. Enantiomers are different from the geometric isomers that involve E or trans isomers or Z isomers.

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7. Which of the following functional groups in organic chemistry do not have a nitrogen as part of the functional group? a. Amide b. Amine c. Nitrile d. Carboxylic acid Answer: d. Carboxylic acid compounds have COOH molecules as part of them. The rest have some functional group that has nitrogen in the group. 8. What type of molecule is identified as being an RCO configuration? a. Ketone b. Aldehyde c. Carboxylic acid d. Amide Answer: a. The ketone is an RCO molecule with a double bond between the carbon atom and an oxygen molecule. This gives the carbon atom two “spots” for attaching to some other side chains. 9. What is not a part of an amino acid? a. Ketone side chain b. Carboxylic acid side chain c. Amino side chain d. R side chain Answer: a. Amino acids have an R side chain, an amino side chain and a carboxylic acid side chain. They do not generally have a ketone side chain as a general part of their structure (although technically, a ketone side chain could be part of an R chain, which can be a variety of things).

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10. Carbohydrates each contain a molecule of these atoms, except for what atom? a. Carbon b. Hydrogen c. Oxygen d. Nitrogen Answer: d. The thing that makes carbohydrates unique is that they only contain carbon, hydrogen, and oxygen, with most carbohydrates in nature being C6 carbohydrates, having six carbon atoms.

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SUMMARY Hopefully, the course on college chemistry you have just studied has provided a way for you who have been pursuing an interest in chemistry or who have possibly already been struggling with chemistry to understand this discipline in a much clearer way. As you have seen, chemistry is a wide-reaching subject, ranging from “physical chemistry” or the study of atoms, molecules, chemical reactions, and the properties of solids, liquids, and gases, to things like “organic chemistry” and “biochemistry” that deal with the topic of the chemistry specific to carbon-based molecules, which are the molecules of living things. As you have hopefully determined in studying this course, the majority of chemistry is about the physical properties of atoms and molecules in singular form as well as the physical properties of molecules in different states, including the molecular properties of gases, solids, and liquids. Most atoms and molecules are not inert and interact with one another in specific chemical and physical ways. This has been the major focus in this course as we learned about how different substances exist in natural and experimental situations. Chapter one in the course introduced chemistry in its most basic form by covering atoms and their nature. Everything in the universe is made from different types of atoms. As it turns out, atoms have specific sub-particles that make up the basic structure of the atom. There are also numerous different atoms in nature, as we talked about when introducing the periodic table, which outlined the relationships between different types of atoms. Finally, there was a discussion on the concept of the “mole” as it exists in understanding quantities of atoms and molecules. The focus of chapter two was the structure of atoms. As you have seen, there are specific rules that define how subatomic structures are organized. This chapter looked at the arrangement of electrons, and both the Pauli principle and Aufbau Principle, which define electron relationships inside the atom. Electrons have a natural affinity for the

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proton they surround but also allow for affinity of different types of atoms in the formation of larger molecules. This concept was also covered in the chapter. The main topics of chapter three in the course were the properties of gases. There are specific laws that describe the behavior of gases in chemistry, which were introduced and discussed in this chapter. The study of gases involves the relationship between volume and pressure in a gas and, in the situation of mixed gases, a relationship between the different gases in a given volume. The behavior of gases in relation to the ideal gas law was covered in this chapter as well as the study of the kinetics of gas molecules in a gaseous solution. Chapter four talked about chemistry as it relates to heat, also called “thermochemistry”. There is an intricate connection between molecules and the temperature of the atoms in the molecules. Concepts of heat capacity and calorimetry were discussed as they relate to a specific molecule. The enthalpy or the total heat in a given system was discussed as it relates to the energy of the system. As in all of physical chemistry, there are specific rules related to the properties of substances at different temperatures, which were also covered in this chapter. Chapter five in the course got into the subject of chemical bonding and molecular structure. Unless an atom is completely inert, it will have the natural tendency to interact with other atoms. This is certainly the case with ionic substances in chemistry. The chapter looked at the issue of covalence and covalent bonding between molecules and how this should look from an atomic perspective. Because of the different properties of atoms, they exist within molecules in specific three-dimensional arrangements, which was covered in this chapter. The topic of bonding in metallic atoms was also covered as part of this chapter. Chapter six focused on the inevitability of chemical reactions and on what chemical formulas look like. Molecules and atomic substances tend not to stay separate from one another when put together in the same physical space and reactions will follow a specific pattern. The rates of different reactions and why they occur or do not occur were discussed as well as the issue of the energy of activation necessary to take a group of molecules in one form and turn them into molecules of a different form. Exactly what

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happens in chemical reactions was covered as part of this chapter, including the fact that reactions ultimately have an equilibrium state. The focus of chapter seven in the course was the behavior of solids and liquids. Molecules can exist in crystalline form, depending on the physical circumstances. There are certain molecular forces in play when substances are in crystalline form, which was covered in this chapter. In the same way, liquids have certain properties unique to that state that depend on the atomic nature of the molecule and on the temperature of the liquid. The different forces in play when substances are in liquid form were also discussed in the chapter. The study of liquids continued in chapter eight with a discussion of solutions and colloids. Solutions involve the dissolution of a substance in a liquid medium. Exactly how this is successfully done was covered as part of this chapter as well as what it means to indicate the “concentration” of a solute in solution. A solution, as you have discovered, is a unique medium for reactions to take place, as it exhibits certain properties that keep things in a given solution and that help drive reactions within the solution. Colloids have certain properties that were also discussed. Two significant types of reactions were covered in chapter nine in the course—that being oxidation and reduction reactions, which are opposing but interrelated chemical reactions. Because these reactions happen in a balanced way, they are often discussed together. Taken together, an oxidation-reduction reaction is any chemical reaction in which the oxidation number of a molecule, atom, or ion is altered by gaining or losing an electron. You hopefully have come to understand why these reactions are related and how to balance them. The specific properties of acids and bases were discussed in chapter ten in the course. Methods involved in the titration of acids and bases were discussed as part of this chapter as were properties of strong and weak acids and bases. Buffers, as you have learned, are weak acids or weak bases that prevent significant changes in the pH of a solution; how these substances work in acid-base chemistry was covered. Finally, the topic of hydrolysis was an important part of the discussion in this section.

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The focus of chapter eleven in the course was electrochemistry, or the study of electricity as it applies to chemical reactions. Faraday’s law, which states that the amount of substance produced at an electrode is directly proportional to the quantity of charge flowing through an electrochemical cell, was discussed. As you have learned, the basics of electricity in chemical reactions are important to understand because electricity plays a role in many chemical interactions. The different properties of electrical cells and what makes them up was a part of this chapter’s discussion. While the subjects of organic chemistry and biochemistry are courses of their own, they were instead covered together in chapter twelve because they are important aspects of the study of chemistry. Organic chemistry is carbon-based chemistry and involves a variety of different types of molecules and reactions typically seen in living things. The nature of organic molecules was discussed as well as the functional groups that shape organic molecules and their behavior. Finally, biochemistry and its principles were covered as these types of molecules, too, follow the basic chemical principles already laid out in this course.

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COURSE QUESTIONS AND ANSWERS 1. Which subatomic particle weighs the most? a. Neutrons b. Protons c. Electrons d. Protons and neutrons weigh the same Answer: a. Neutrons are the heaviest subatomic particle; however, their weight is very close to that of the proton with a total weight of 1.69 x 10-27 kilograms. 2. How many quarks are in a proton? a. There are two up quarks and two down quarks b. There is one up quark and two down quarks c. There are two up quarks and one down quark d. There are two up quarks for a positive charge Answer: c. There are two up quarks with a charge of +2/3 each and one down quark with a charge of -1/3 so that there is a net charge of +1. 3. What mediates the force that holds quarks together in a proton? a. Gluons b. Nuclear force c. Residual strong force d. Gauge boson Answer: a. Gluons are subatomic elementary particles that mediate the force that holds quarks together in a proton and in a neutron.

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4. Currently, what is the highest known atomic number? a. 53 b. 74 c. 98 d. 118 Answer: d. The highest known atomic number is 118, from the element Organesson, which does not technically exist in nature but has been experimentally created. It is an unstable element, prone to radioactive decay because of the high number of protons and neutrons in the element. 5. Tritium is an isotope of the hydrogen atom. How many neutrons are found in tritium? a. 0 b. 1 c. 2 d. 3 Answer: c. Tritium has two neutrons in it and one proton, giving it a mass number of three and an atomic number of one. 6. Which form of carbon is considered to be radioactive carbon? a. Carbon-11 b. Carbon-12 c. Carbon-13 d. Carbon-14 Answer: d. Carbon-14 is the radioactive isotope of carbon because it isn’t stable and will decay over time.

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7. Which element has an isotope that is used in dating fossils? a. Carbon b. Fluoride c. Uranium d. Cobalt Answer: a. Carbon-14 is used for carbon dating in paleontology because it decays at a specific rate that can determine the timing of the origin of fossils and other carbon-based substances. 8. An isotope varies from its other forms by the number of what substance in each atom? a. Quarks b. Neutrons c. Electrons d. Protons Answer: b. Neutron number is what separates an isotope from the other isotopes of the same element. 9. An atom has seven electrons in its outer shell. What is its valency? a. 1 b. 4 c. 7 d. 8 Answer: a. The valency is one because it needs one electron in order to make a full octet. The same would be true of an element that has an electron to “give away” in order to make a full octet.

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10. How many electrons fill an outer orbital in an atom in order for the atom to be considered stable? a. 4 b. 8 c. 12 d. 16 Answer: b. The outer orbital, in order to be considered stable, usually needs a full octet or eight electrons per outer shell. 11. What is the valency of the noble gases? a. 0 b. 2 c. 1 d. 3 Answer: a. Noble gases do not react with other elements as their orbitals are full, leading to a zero valency. 12. The atomic number of sulfur is 16. What is its valency? a. 0 b. 1 c. 2 d. 3 Answer: c. The valency of sulfur is two because two electrons fill the first orbital and eight electrons fill the outer orbital. This leads to six left, which does not fill out an octet, leaving 2 missing spots for a valency of 2.

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13. The electronegativity of an atom is highest in what part of the periodic table? a. Upper left b. Upper right c. Lower left d. Lower right Answer: b. The electronegativity of an atom or element is highest in the upper right-hand portion of the periodic table. 14. The atomic number of an atom or element is the greatest in what part of the periodic table? a. Upper left b. Upper right c. Lower left d. Lower right Answer: d. The atomic number is highest in the lower right-hand portion of the periodic table. 15. Which is not an alkali metal? a. Lithium b. Iron c. Potassium d. Sodium Answer: b. Iron is a transitional metal and does not take the typical +1 charge when ionized; each of the others is an alkali metal.

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16. Which of the following is a halogen? a. Fluorine b. Oxygen c. Nitrogen d. Carbon Answer: a. Fluorine is a halogen that lies along the seventeenth group of the periodic table. It will typically take a -1 charge when ionized, which is a characteristic of halogens. 17. Which of these chemical symbols is incorrect? a. K is potassium b. I is iodine c. Au is gold d. Co is copper Answer: d. The one that is incorrect is “Co is copper” because copper’s chemical symbol is Cu and Co represents Cobalt instead. 18. What is the metallic radius in an element? a. Half the distance between the nuclei of an anion and cation. b. The diameter of a metallic atom divided by two. c. Half the distance between the nuclei of two metallic crystalline substances. d. The distance between the nucleus and outer electron in a metallic atom. Answer: c. The actual diameter and radius of an atom cannot be determined directly so it is considered half the distance between the nuclei of two metallic crystalline substances.

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19. The element Potassium has an atomic number of 19. What is its effective atomic radius? a. +1 b. +2 c. +4 d. +7 Answer: a. The atomic radius is +1 because the Atomic number is 19 and the number of shielding electrons is 2 + 8 + 8 or 18. This leads to an equation so 19 minus 18 or +1. 20. What is the best explanation of the ionization energy of an atom? a. It is a measure of the reactivity of the atom. b. It goes up as the atomic radius goes up. c. It is the energy required to lose an electron. d. It is the energy required to gain an electron. Answer: c. The ionization energy is the energy required to lose an electron and increases as one goes up a specific group or column in the periodic table because of fewer shielding electrons in smaller atoms. 21. If the atomic mass of carbon is 12 and the atomic mass of oxygen is 16, what is the formula mass of carbon dioxide? a. 28 b. 32 c. 40 d. 44 Answer: d. The formula mass of carbon dioxide is 44 because the atomic mass of carbon is 12 added two of the atomic masses of oxygen or two times 32, which leads to 44.

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22. What is the definition of Avogadro number or Avogadro’s constant? a. 2 x 1012 b. 6.022 x 1023 c. 6.022 x 10-23 d. 1.2 x 1024 Answer: b. Avogadro’s constant or Avogadro number is a specific number that is 6.022 x 1023. 23. What type of element has been used to define Avogadro’s number? a. Carbon b. Hydrogen c. Neon d. Potassium Answer: a. Avogadro’s constant is defined as the number of atoms in exactly 12 grams of carbon-12, which is approximately 6.022 x 1023. 24. Which orbital series collectively has the largest capacity for electrons? a. s b. p c. d d. f Answer: d. There are seven f orbitals each with the capacity for 2 electrons of opposite spin. This leads to a total of 14 electrons in this orbital series.

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25. How many elements are there in the first period of the periodic table? a. 2 b. 8 c. 18 d. 32 Answer: a. There are just two elements in the first period, which are hydrogen and helium. This fits with the idea of two electrons in the first shell of the atom. 26. How many elements are there in the third period of the periodic table? a. 2 b. 8 c. 18 d. 32 Answer: b. There are eight elements in the third period filling up the different s and p orbitals. 27. How many s sub-orbitals are there in the s shell of an atom? a. 2 b. 4 c. 5 d. 7 Answer: a. There are only 2 s sub-orbitals, making up the 1s and 2s orbitals.

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28. How many p sub-orbitals are there in the p shell of an atom? a. 2 b. 3 c. 5 d. 7 Answer: b. There are 3 sub-orbitals with a total of 6 electrons possible in the p shell. Each sub-orbital must contain two electrons of opposite spin. 29. How many sub-orbitals are there in the d shell of an atom? a. 2 b. 3 c. 5 d. 7 Answer: c. There are 5 sub-orbitals called d-orbitals, which collectively can contain a total of 10 electrons. These are seen in the larger atoms. 30. What statement best describes an exothermic reaction? a. It is any reaction in which energy is given off. b. It is any reaction in which energy is absorbed. c. It is any reaction in which electrons are taken on by an atom. d. It is any reaction in which electrons are removed from an atom. Answer: a. An exothermic reaction is any reaction in which energy is given off; it may or may not involve the gaining or losing of an electron.

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31. Which type of element has the greatest electron affinity? a. Alkali metal b. Noble gas c. Halogen d. Transitional metal Answer: c. Halogens have the highest electron affinity because they want to gain an electron to fill their p orbital. 32. What element is the most electronegative? a. Fluorine b. Carbon c. Hydrogen d. Oxygen Answer: a. Fluorine is the most electronegative and the most reactive of all the elements. 33. What does the Pauling scale measure? a. Ionization energy b. Electron affinity c. Electronegativity d. Atomic periodicity Answer: c. The Pauling Scale arbitrarily assigns a value to an element according to its degree of electronegativity.

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34. What type of bonding has the greatest dipole moment created? a. Hydrogen bonding b. Covalent bonding c. Ionic bonding d. Pure covalent bonding Answer: c. Ionic bonding has a highly electronegative atom with a low electronegative atom, giving it a greater dipole moment or separation of charge. 35. Which molecule is considered highly polar? a. H20 b. Carbon tetrachloride c. Acetone d. Ethanol Answer: a. Water is a highly polar molecule and therefore interacts strongly with other polar molecules, particularly in solutions with water as a solvent. 36. A substance that has paired electrons that cancel each other out is said to have what type of magnetism? a. Paramagnetism b. Diamagnetism c. Ferromagnetism d. Anti-ferromagnetism Answer: b. A substance that has paired electrons in the orbitals so that they cancel each other out is said to be diamagnetic or to have diamagnetism.

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37. The O atom has 2s22p4 as the electron configuration. How many unpaired electrons are there? a. 1 b. 2 c. 3 d. 4 Answer: b. Oxygen have four p electrons in the outer shell. With three sub-orbitals in the p series or p shell, there should be 6 to fill the p shell and there are only 4, leaving two unpaired electrons. This is why oxygen can form two bonds with other atoms. 38. There are many different measurements that define pressure. What is the unit of pressure used in meteorology to define gas or “air” pressure? a. Bar b. Atmosphere c. Torr d. Pascal Answer: a. Each of these is a measurement of pressure of gas, with the bar being used in meteorology to measure the pressure of air. 39. When two gases of different temperature are adjacent to each other, what tends to happen? a. The temperatures of the gases remain the same unless they mix. b. The cold will flow to the warmer area to make the warm area colder but the cold area will stay at the same temperature. c. The warm area will stay the same temperature and yet the cold area will become warmer. d. Heat will flow from the warm to the cold area until the temperatures are equal.

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Answer: d. In such cases, the heat will flow from the warm to the cold area until the temperatures are the same, even if the gases do not technically mix. 40. How many degrees Celsius is absolute zero on the Kelvin scale? a. 0 degrees Celsius b. -273 degrees Celsius c. -475 degrees Celsius d. -521 degrees Celsius Answer: b. Absolute zero on a Celsius scale is -273 degrees. The actual degree increments are the same. 41. One-degree difference on a scale is the same for all the different temperature scales except for which one? a. Fahrenheit scale b. Celsius scale c. Centigrade scale d. Kelvin scale Answer: a. The Fahrenheit scale is different than any of the other scales such that 1-degree Fahrenheit is 5/9ths of the other scales’ degrees. 42. What is not a property of an ideal gas? a. The molecules do not occupy space of their own because they are so small. b. The molecules do not interact with one another. c. The molecules only interact at very low temperatures. d. Molecules move in a constant straight line randomly. Answer: c. These are all properties of an ideal gas with no interaction with each other, even at lower temperatures, which is a feature of real gases that have molecules that interact particularly at lower temperatures when there is little molecular movement.

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43. Under which situation is a gas most likely to behave non-ideally? a. High pressure b. Low pressure c. High volume d. High temperature Answer: a. Under high pressure situations, the gases will not behave ideally because the volume of the gas molecules will come into play. 44. When determining the density of gases, the STP is referred to as the standard temperature and pressure. What is the STP referring to? a. 25 degrees Celsius and 1 mm Hg of mercury b. 0 degrees Celsius and one atmosphere of pressure c. 25 degrees Celsius and one atmosphere of pressure d. 0 degrees Kelvin and 1 mm Hg of mercury Answer: b. STP is considered to be 0 degrees Celsius and one atmosphere of pressure. 45. Which gaseous atom behaves the most ideally as gases behave? a. Neon b. Hydrogen c. Oxygen d. Helium Answer: d. Helium has the smallest molecular size and has molecules that do not interact with one another, making this gas behave the most ideally.

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46. Avogadro’s Law predicts the relationship between what to factors of gases? a. Pressure and Temperature b. Volume and number of molecules c. Pressure and Volume d. Volume and Temperature Answer: b. Avogadro’s number assumes that the pressure and the temperature stay the same and indicates a linear relationship between the volume and the number of particles. 47. The ideal gas law uses aspects of what other gas law in determining it? a. Avogadro’s law b. Boyle’s law c. Charles’ law d. All gas laws Answer: d. The ideal gas law uses aspects of all three gas laws in determining the values of pressure, temperature, volume, and number of particles. 48. Why will gases behave differently than ideally at lower temperatures? a. Because there are increased forces between the molecules b. Because the area is more crowded at low temperatures c. Because gases will turn to liquid at low temperatures d. Because gases will precipitate into a solid at low temperatures Answer: a. Gases will behave differently than ideally at lower temperatures because there are increased forces between the molecules when they operate at lower speeds and aren’t actively “bumping” against each other in an elastic way.

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49. Which gas law involves the Law of Partial Pressures? a. Avogadro’s law b. Charles’ law c. Dalton’s law d. Boyle’s law Answer: c. Dalton’s law applies specifically to the law of partial pressures of gases. The other laws do not directly relate to the partial pressures of the different gas molecules. 50. In determining the partial pressure of a gas in a mixture, what aspect of the gas is most important to consider? a. The number of molecules of the gas b. The molar weight of the gas c. The polarity of the gas molecule d. The temperature of the system Answer: a. The number of molecules of the gas is the only thing that is important and not the nature of the gaseous molecule or the temperature of the system. 51. The partial pressure of Gas A is P1 and the partial pressure of Gas B is P2. Gas A is twice as large molecularly when compared to Gas B. What is the total pressure of the mixture of gas A and gas B? a. 2P1 plus P2 b. 2P1 multiplied by P2 c. P1 multiplied by P2 d. P1 plus P2 Answer: d. The total pressure will always be the sum of the two pressures, regardless of the size or character of the two gases.

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52. What is not a postulate of kinetic molecular theory of gases? a. Gases are point molecules with no mass. b. Gas molecules travel in a straight line. c. Gas pressure is made from static forces between the molecules. d. No energy is lost in the collision of molecules. Answer: c. The gas pressure is made from collisions of the molecules within the container and is not based on static forces between the molecules. 53. The kinetic energy of a gas is most related to what factor? a. Gas pressure b. Temperature c. Molar mass d. Molecular size Answer: b. The temperature is the most important factor in determining the kinetic energy of a gas. 54. What is the difference between gaseous effusion and gaseous diffusion? a. They are considered the same in an ideal gas. b. Diffusion allows for intermolecular collisions and effusion does not. c. Diffusion involves real gases while effusion involves ideal gases. d. Diffusion applies only to molecules of the same size, while effusion does not. Answer: b. Diffusion allows for the fact that intermolecular collisions occur, while effusion does not allow for these types of collisions.

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55. What is the SI unit of energy? a. Kilocalorie b. Calorie c. Kilojoule d. Joule Answer: d. The joule is the SI unit of energy but because it is such a small value, the term kilojoule is often used instead for practical purposes. All forms of potential and kinetic energy can be expressed in joule form. 56. About how many joules are equal to one calorie when measuring energy? a. 1 b. 3 c. 4 d. 7 Answer: c. Calories and joules are both measures of energy, with one calorie being equal to 4.184 joules. Calories are usually thought of in terms of thermal energy but both forms will apply to any type of energy source. 57. What gets exchanged between systems in an isolated chemical reaction? a. Neither energy nor matter b. Energy but not matter c. Matter but not energy d. Both energy and matter Answer: a. Technically, neither energy nor matter gets exchanged outside of the system in an isolated chemical reaction; however, this type of thing does not occur in reality as energy and matter can be leached out of the system.

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58. What is not considered an endothermic reaction or reactions? a. Solid water going to liquid water b. The building of glucose from CO2 and water c. The breakdown of glucose into CO2 and water d. Liquid water going to water vapor Answer: c. The breakdown of glucose into CO2 and water gives off heat, which is important in the heat seen in warm-blooded individuals. 59. Which law of thermodynamics indicates that entropy is always increasing? a. Law of conservation of energy b. First law of thermodynamics c. Second law of thermodynamics d. Third law of thermodynamics Answer: c. The second law of thermodynamics states that, in any given system, entropy is always increasing. This is true of the entropy of the entire universe. 60. What is the definition given to the amount of heat necessary to raise an object of one kilogram of mass up 1-degree Kelvin? a. Heat capacity b. Specific heat c. Calorie d. Thermal capacity Answer: b. The specific heat is the amount of heat necessary to raise a kilogram of mass up one-degree Kelvin.

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61. What is the universal SI unit for heat capacity? a. Joules b. Joules per second c. Joules per Kelvin squared d. Joules per Kelvin Answer: d. This unit of heat capacity can also be said to be joules per degree in which the degree means the degrees Kelvin (which is the same increment as degrees Celsius). 62. What is the universal SI unit for volumetric heat capacity? a. Calories per gram per kelvin b. Joules per mole per Kelvin c. Joules per meter cubed per Kelvin d. BTUs per meter cubed per Kelvin Answer: c. The SI unit for volumetric heat capacity is the joule per meter cubed per kelvin. This is often used in scientific and engineering circles to deal with the heat capacity in a certain volume of a substance. 63. What is the specific heat capacity of water in joules per gram per degree kelvin? a. 1 b. 4.19 c. 6.73 d. 114.2 Answer: b. The most accurate measurement of the specific heat capacity of water is 4.186 or about 4.19 joules per gram per degree Kelvin. This measurement is often used rather than the caloric specific heat capacity measurement for water.

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64. How many translational degrees of freedom does a gaseous substance have, which play a role in its heat capacity? a. 1 b. 2 c. 3 d. 4 Answer: c. There are three translational degrees of freedom, which are the three axes in space, the x-axis, the y-axis, and the z-axis. 65. Heat capacity is related to the number of degrees of freedom of a molecule. Which type of degrees of freedom can contribute the most to the heat capacity? a. Rigid degrees of freedom b. Translational degrees of freedom c. Rotational degrees of freedom d. Vibrational degrees of freedom Answer: d. There can be many vibrational degrees of freedom of a molecule but only a maximum of 3 translational degrees and three rotational degrees of freedom, corresponding to the three dimensions. 66. What is one thing that a calorimeter does not measure? a. Force on a system b. Energy of a system c. Heat of a reaction d. Heat capacity of a substance Answer: a. Because heat and energy are essentially the same, heat, energy, and heat capacity can be measured in a calorimeter; however, force is not measured in a calorimeter.

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67. What is not a state function of a substance? a. Enthalpy b. Pressure c. Number of atoms d. Heat flow Answer: d. With the exception of heat flow, each of these is a state function that depends on the condition of the substance rather than on the change in a substance. 68. When it comes to bond-breaking in chemical reactions, what can be said about these reactions? a. They involve a negative enthalpy situation and require an input of energy b. They involve a negative enthalpy situation and release energy. c. They involve positive enthalpy and require an input of energy. d. They involve positive enthalpy and release energy. Answer: c. Positive enthalpy is the same thing as an input of energy and an endothermic reaction, which is true of bond breaking. 69. When it comes to making bonds in chemical reactions, what can be said about these types of reactions? a. They involve a negative enthalpy situation and require an input of energy b. They involve a negative enthalpy situation and release energy. c. They involve positive enthalpy and require an input of energy. d. They involve positive enthalpy and release energy. Answer: b. Negative enthalpy is the same as releasing energy, which is true of bond-making and of an exothermic reaction.

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70. A reaction takes place in which substrate A goes to end products B and C in one step and gives off a certain amount of energy +H per mole of substrate A. It takes three chemical steps to turn the two end products back into the same amount of substrate A. What is the H change in the reverse reactions? a. The reverse three reactions will have a -H energy change. b. The reverse three reactions will have a -3H energy change. c. The reverse three reactions will have a +H energy change. d. The reverse three reactions will have a +3H energy change. Answer: a. The enthalpy per mole of substrate A going to the two end products is +H. The reverse reactions will always be equal to -H, regardless of the number of steps it takes to get there. 71. What does Hess’ rule indicate in thermochemistry? a. Energy cannot be created or destroyed. b. Enthalpy is a state function. c. The state of enthalpy is always increasing. d. It takes more energy to cause an endothermic reaction than is given off in an exothermic reaction. Answer: b. Hess’ rule indicates that enthalpy is a state function. This means that the change in entropy depends solely on the enthalpy of the starting product and the enthalpy of the end product and does not depend on the pathway to get to the end product. 72. What type of bonding is least likely to become involved in the conduction of electricity? a. Covalent bonds b. Metallic bonds c. Hydrogen bonds d. Ionic bonds

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Answer: a. Covalently bonded molecules tend not to dissolve well in water and do not conduct electricity very well. Both ionic bonding and metallic bonding will lead to the conduction of electricity. 73. Which is not a characteristic of metals involved in metallic bonding? a. The sharing of electrons with multiple nuclei b. Malleability c. Low melting point d. Luster Answer: c. With the exception of a low melting point, each of these is a characteristic of metals bonded with metallic bonding. They are, for the most part, solids at room temperature. 74. Which type of bonding is considered a weak connection between molecules and atoms? a. Ionic bonding b. Hydrogen bonding c. Metallic bonding d. Covalent bonding Answer: b. Hydrogen bonds are weak connections between molecules and atoms. Of these, the strongest bond occurs in covalent bonds, which do not separate without a great deal of energy involved. 75. What is not a difference between a compound and a mixture? a. A compound has a specific composition of different atoms and a mixture does not. b. Compounds are bonded chemically and mixtures are not. c. Compounds have unique physical properties and mixtures do not. d. Compounds involve just one type of atom and mixtures do not. Answer: d. Both compounds and mixtures can have more than one type of atom but mixtures do not have unique physical properties, chemical bonding between substances, or the same composition of atoms.

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76. Which element will not form a +1 charged cation? a. Sodium b. Calcium c. Lithium d. Potassium Answer: b. Each of these will form a +1 cation except for calcium, which forms a +2 cation. Those that form a +1 cation exist in the 1A family or first group of elements in the periodic table. 77. Which element is most likely to form a -1 anion? a. Aluminum b. Oxygen c. Nitrogen d. Chlorine Answer: d. Chlorine is in the 7A family and needs just one electron to form an octet. For this reason, it is most likely to become a -1 anion. 78. Which element will most likely form a +2 cation? a. Magnesium b. Aluminum c. Potassium d. Cesium Answer: a. Magnesium is willing to give up two electrons for a stable octet so it is most likely to become a +2 cation. 79. Which element is most likely to react similarly to nitrogen in a chemical reaction? a. Fluorine b. Sulfur c. Phosphorus d. Oxygen

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Answer: c. Phosphorus and nitrogen both have five electrons, requiring three to make a -3 anion so they behave similarly, at least when it comes to bonding. 80.Which polyatomic ion is a cation versus being an anion? a. Ammonium b. Sulfate c. Cyanide d. Acetate Answer: a. Ammonium is positively charged, making it the only polyatomic cation. The others are negatively charged. 81. What is the difference between hypochlorite, chlorite, chlorate, and perchlorate? a. They have different charges b. They differ in the number of oxygen atoms c. They differ in the number of chlorine atoms d. There is no difference between them Answer: b. These are polyatomic ions that have the same charge but differ in the number of oxygen atoms attached to the chlorine atom. The number of chlorine atoms is the same in all four cases. 82. What is the charge on the ammonium polyatomic ion? a. +1 b. -1 c. -2 d. -3 Answer: a. The charge on the ammonium polyatomic ion is +1. This is because the overall charge on the nitrogen ion is -3 plus the four +1 charges in the hydrogen ion, making this ion a 4+ cation.

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83. What is the formula for aluminum chloride? a. AlCl b. AlCl2 c. AlCl3 d. Cl2Al Answer: c. The formula is AlCl3 because aluminum has a +3 charge and chloride has a -1 charge so it takes three chlorides to balance the molecular formula. 84. What is the molecular formula for magnesium sulfate? a. MgSO4 b. Mg2SO4 c. MgS2O8 d. Mg(SO4)2 Answer: a. This involves magnesium, which naturally has a +2 charge and sulfate or SO4, which naturally has a -2 charge. These are balanced when the overall charge is equal, which makes MgSO4 the correct answer. 85. How do you name the compound Na2O? a. Disodium oxide b. Sodium dioxide c. Sodium oxide d. Oxygen sodium Answer: c. While it is two molecules of sodium and one molecule of oxygen in this molecular formula, it is still called sodium oxide with the inference of the formula being Na2O because of the charges on the respective atoms.

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86. What is the name of the molecular formula NaC2H3O2? a. Sodium hydroxide b. Sodium acetate c. Disodium hydroxide d. It cannot be named with the information given Answer: b. The atomic symbol for sodium is Na and the anion C2H3O2 is the chemical symbol for acetate. The sodium ion is a +1 charge, while acetate is negatively charged at -1. This leads to sodium acetate as the appropriate name for this. 87. What is the name of the molecular formula CuCl2? a. Cupric chloride b. Copper (III) chloride c. Cuprous chloride d. Copper (II) chloride Answer: d. Because copper can be charged differently, it is important to look at the Chloride ion in the equation, which has a -2 charge. The matching copper ion must be a +2 charge. This leads to the conventional way to write this, which is Copper (II) chloride. 88. What is the orbital called in an atom that is involved in a connection with another atom to form a molecule using covalent bonding? a. Anti-bonding region b. Molecular orbital c. p orbital d. Bonding molecular orbital Answer: d. The bonding molecular orbital is the molecular orbital that is involved in the covalent bond between two atoms. A molecular orbital is not necessarily involved in bonding but is said to be available for bonding.

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89. Which molecule is most likely to be considered a polar covalently-bonded molecule? a. O2 b. H20 c. NaCl d. KOH Answer: b. Water or H2O is the only polar covalent molecule in the list. O2 is a nonpolar covalent molecule and both KOH and NaCl are ionic molecules. 90. What is not a characteristic of a polar molecule versus a nonpolar molecule? a. Higher boiling point b. Higher melting point c. Higher vapor pressure d. Greater solubility in water Answer: c. Each of these except having a higher vapor pressure is true of polar molecules

when compared to nonpolar molecules.

91. Hydrogen bonding is described as bonding between hydrogen and what other type of molecule? a. Nonpolar molecules b. Other hydrogen molecules c. Carbon molecules d. Electronegative molecules Answer: d. Hydrogen bonding happens between hydrogen and other electronegative molecules; it represents a partial charge difference between hydrogen and any electronegative atom that is also bonded with another atom in a covalent bond.

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92. Which type of bonding in molecules creates what can be referred to as an “electron soup”? a. Metallic bonding b. Hydrogen bonding c. Covalent bonding d. Ionic bonding Answer: a. Metallic bonding involves metals with low ionization energies that share a group of electrons with other atoms in an “electron soup” that easily conducts electricity. 93. In what type of bonding is there an overlap of orbitals more so than a sharing of electrons? a. Covalent bonding b. Metallic bonding c. Hydrogen bonding d. Ionic bonding Answer: b. Metallic bonding involves an overlapping of orbitals rather than a true sharing of electrons. The electrons are relatively free to move between different orbitals of different atoms. 94. What is true of the activation energy in an endothermic reaction? e. It is much greater in an endothermic reaction than in an exothermic reaction. a. The energy put in during an endothermic reaction is not regained. b. The energy is regained before the final end product energy is achieved. c. It is not necessary to have an activation energy for this type of reaction. Answer: c. In an endothermic reaction, there is an activation energy that is regained before the final end product is achieved.

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95. A reaction in which a substance is synthesized from two or more precursor substances is called what kind of reaction? a. Combination reaction b. Decomposition reaction c. Single displacement reaction d. Double displacement reaction Answer: a. A combination reaction is one in which a single substance or lesser number of substances is synthesized from a greater number of reactants. 96. The breakdown of hydrogen peroxide into oxygen gas and water is an example of what type of reaction? a. Precipitation reaction b. Neutralization reaction c. Combustion reaction d. Decomposition reaction Answer: d. This is the breakdown of hydrogen peroxide or H2O2 into two simpler components, which is the definition of a decomposition reaction. 97. When a reactive metal removes a less reactive metal salt to make a salt and a solid form of the less reactive metal, this is what type of reaction? a. Combination reaction b. Neutralization reaction c. Single displacement reaction d. Double displacement reaction Answer: c. This is a single displacement reaction in which a reactive metal displaces a less reactive metal to make a salt of the more reactive metal and a solid form of the less reactive metal.

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98. In what type of reaction does a substance in solution leave the solution as a solid? a. Combustion reaction b. Precipitation reaction c. Decomposition reaction d. Neutralization reaction Answer: b. In a precipitation reaction, at least one substance in solution is said to “precipitate” out of solution, driving the reaction forward. 99. Which chemical reaction is written correctly? a. H2 + N2 → NH3 b. 2H2 + 2N2 → 2NH3 c. 2H2 + N2 → 2NH3 d. 2H2 + N2 → 2NH3 Answer: c. The goal is to balance the equation so that there is an equal number of moles of each element on both sides. This is the basis of stoichiometry of chemical reactions. 100.

If you balance the equation in which Magnesium and Oxygen gas goes to

Magnesium oxide, how many moles of magnesium oxide are there on the right-hand side of the equation? a. 1 b. 2 c. 3 d. 4 Answer: b. Because oxygen is a diatomic gas (meaning O2 gas) and Magnesium oxide is MgO, you need to balance the MgO to account for the fact that there is just one O2 gas on the left, making the necessity of having 2 MgO on the right side of the equation.

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101.

In a reaction involving two reactants leading to an end product, what can

be said of the theoretical yield? a. It is a number based on a complete reaction occurring. b. It is always 100 percent. c. It will be above 100 percent. d. It will always be less than the percent yield. Answer: a. The theoretical yield is not a percent but an actual number of grams or other number of a substance that is based on a complete reaction occurring. It is always equal to or more than the actual yield. 102.

When writing an equation, what defines a molecule that is in a solution of

water? a. l b. aq c. s d. g Answer: b. Aqueous, or aq, in particular defines a substance that is in solution specifically in water. Liquid or “l” defines a liquid and not a solution in water. 103.

In a neutralization reaction, what are the reagents that begin the reaction?

a. Acids b. Water c. Acids plus bases d. Salts plus water Answer: c. This reaction always starts with acids and bases that get “neutralized”, turning into salts plus water.

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104.

What are the main end products of a combustion reaction?

a. CO2 and water b. Carbon and oxygen c. Hydrocarbons d. CO2 and hydrogen gas Answer: a. The end product of this type of reaction is always going to be CO2 and water because the hydrocarbon bonds are broken and combined with oxygen to give water and CO2 as an endpoint. 105.

What type of reaction is this: C3H6O3 + O2 goes to CO2 (g) +H2O (g)?

a. Endothermic decomposition reaction b. Exothermic decomposition reaction c. Endothermic combustion reaction d. Exothermic combustion reaction Answer: d. This is a combustion reaction in which the end result is CO2 and water plus heat, making it exothermic. 106.

If you have 2 moles of H2 gas and when added to O2 gas, it leads to 0.8

moles of water, what is the limiting reagent? Hint: the reaction is 2H2 + O2 leads to 2H2O a. There are equal amounts of the reagents b. H2O c. H2 d. O2 Answer: d. If there were two moles of H2 gas and 1 mole of O2, it should go to 2 moles of water but, as only 0.8 moles of water are made, there must have been only 0.4 moles of O2 in the equation, so oxygen or O2 gas is the limiting reactant or reagent.

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107.

What factor is least likely to play a role in the rate of reaction in a gaseous

reaction? a. Volume b. Pressure c. Temperature d. Catalysts Answer: a. Pressure, temperature, and catalysts will play a role in a gaseous reaction; however, the volume, if it doesn’t affect the pressure, will not affect the rate of reaction. 108.

The reaction A + 2B = C involves a certain rate expression. If the rate

expression for C is 2 moles per hour, what are the rate reactions for A and B. a. 2 moles per hour for A and two moles per hour for B b. -2 moles per hour for A and -4 moles per hour for B c. ½ mole per hour for A and 1 mole per hour for B d. -1/2 moles per hour for A and -2 moles per hour for B Answer: b. The rate of reaction for reactants is always going to be a negative number and will reflect the stoichiometry of chemical reaction. This means that -2 moles per hour is the rate of reaction for A and -4 moles per hour is the rate of reaction for B. 109.

What will most likely lower the activation energy of a reaction?

a. Temperature b. Pressure c. Catalysts d. Concentration Answer: c. Catalysts are the only thing that will actually lower the activation energy by modifying the transition states of the reaction.

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110.

What happens at chemical equilibrium in a reaction?

a. The rate of the forward reaction equals the rate of the reverse reaction. b. The rate of forward reaction exceeds the rate of the reverse reaction. c. There is equal quantities of reactants and products. d. All of the limiting reactants have been used up. Answer: a. The equilibrium in a chemical reaction has been achieved when the rate of the forward reaction equals the rate of the reverse reaction. 111.

What factor will not change the state of equilibrium in a chemical

reaction? a. Concentration of reactants b. Gas pressure c. The presence of a catalyst d. Temperature change Answer: c. The presence of a catalyst does not change the state of equilibrium in a chemical reaction; however, each of the others represent disturbances that, according to Le Chatelier’s principle, will change the state of equilibrium. 112.

If there is a group of reactants and products and the equilibrium constant

is zero, what does it say about the products? a. There is a predominance of products b. There is a predominance of reactants c. The ratio of reactants and products is roughly equal d. There are no products and there is no reaction Answer: d. If the equilibrium constant is zero, there is no product or products. The equilibrium has an equal number of reactants and products when the K value (equilibrium constant) is one and not when it is zero, which implies no reaction is occurring.

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113.

Which value is most likely to affect the equilibrium constant or K constant

of an equation? a. Amount of solvent in the reaction b. Concentration of reactant in the reaction c. Amount of solid in the reaction d. Amount of liquid in the reaction Answer: b. The amount of solvent, solid, and liquid in a reaction do not affect the equilibrium constant as their “concentration” is said to equal 1. 114.

What happens to a liquid above the critical point on a pressure-

temperature curve? a. It cannot exist. b. It will condense from a gas to a liquid. c. It will boil. d. It will become solid. Answer: a. Above the critical point on a pressure-temperature curve, a liquid cannot exist and the resulting gas will just become denser. 115.

What is the relative humidity of water at a certain temperature?

a. The vapor pressure of water at a certain temperature b. The percent of water in the atmosphere c. The partial pressure of water in the air d. The ratio of the partial pressure of water to the vapor pressure Answer: d. The relative humidity is the ratio of the partial pressure of water in the air to the vapor pressure of water at a certain temperature. High relative humidity indicates less dense air and not denser air as it is commonly felt to be.

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116.

What is the relative humidity at the dew point of water?

a. 0 percent b. 50 percent c. 100 percent d. It cannot be determined Answer: c. At the dew point of water, the relative humidity is 100 percent so that water will fall as rain and so that dew will form on certain objects, such as grass. It usually occurs when the air temperature drops so that the relative humidity of humid air will reach 100 percent. 117.

What happens during the process of efflorescence in a crystal of a hydrated

salt? a. The salt will lose its water, resulting in a powder on the crystal surface. b. There will be a liquid surface on the crystal surface. c. The salt will begin to liquify and melt. d. The salt will become more gaseous. Answer: a. The salt will form a powdered surface due to a loss of water from its hydrated form. In such cases, the powder is the anhydrous form of the salt. 118.

What most contributes to the development of rain in the atmosphere?

a. Changes in air temperature b. Impurities in the air c. High air pressure d. Changes in air pressure Answer: b. Impurities in the air will cause small droplets to form and grow, causing rain. Without these impurities, it would be more difficult for droplets to grow to a critical size.

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119.

In chemistry and physics, what is sublimation?

a. The melting of a solid under high pressure situations b. The solidifying of a gaseous substance into liquid c. The solidifying of a liquid into a solid d. The transfer of a solid into a gas without a liquid intermediary Answer: d. Sublimation involves the transfer of a solid into a gas without a liquid intermediary. It can be seen when solid CO2 or dry ice becomes gaseous CO2 without a liquid intermediary. 120.

What vapor pressure is lower at the triple point on a phase diagram?

a. The vapor pressure of the gas phase b. The vapor pressure of the liquid phase c. The vapor pressure of the solid phase d. The vapor pressures are the same for each phase Answer: d. The triple point recognizes the point at which the vapor pressures of all three phases are the same. It happens at a specific combination of the temperature and pressure along the phase diagram of a given substance. 121.

At very high temperatures and very high pressures, what state is a

substance found in? a. Liquid b. Supercritical fluid c. Gas d. Solid Answer: b. At very high temperatures and very high pressures, the curve for vapor pressure breaks down and you end up with a supercritical fluid in which gas is dense but does not meet the criteria for being a true liquid.

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122.

What is the preferred crystal shape for sodium chloride?

a. Tetragonal b. Cubic c. Hexagonal d. Orthorhombic Answer: b. Salt or sodium chloride is considered to have a cubic shape. This means that all sides are equal and that the angles are 90 degrees between sides. 123.

Which type of crystalline solid is the best conductor of electricity?

a. Ionic b. Covalent molecular c. Metallic d. Molecular Answer: c. Metallic crystalline solids have free electrons and are therefore good conductors of electricity. 124.

Which type of crystalline solid has the strongest intermolecular

connections? a. Ionic b. Covalent molecular c. Metallic d. Molecular Answer: b. Covalent molecular crystals will have covalent bonds between molecules so that they are exceedingly strong and difficult to divide.

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125.

Which type of crystalline solid has the weakest intermolecular

connections? a. Ionic b. Covalent molecular c. Metallic d. Molecular Answer: d. The molecular forces between the different molecules in this type of structure are extremely weak, making them have lower melting points than some other types of crystal structures. 126.

Which plays the least role in the phenomenon of capillary action of

liquids? a. Intermolecular cohesive forces b. Adhesive forces c. Force of gravity d. Freezing point Answer: d. Each of these will play a strong role in the phenomenon of capillary action of liquids except for the freezing point of a liquid. Gravity plays a role but is, of course, constant for all liquids. 127.

What will cause a liquid to have a lower viscosity?

a. High intermolecular forces b. Low temperatures c. Short molecules d. High number of poise units Answer: c. Each of these indicate a high viscosity situation with the exception of short molecules, such as short hydrocarbons, which will have a decreased viscosity.

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128.

What most affects the boiling point of a liquid?

a. Surface tension b. Pressure c. Volume d. Density Answer: b. The boiling point of a liquid is highly pressure-dependent and will, for example, greatly affect the temperature at which water can boil. 129.

Which ion has the greatest charge density?

a. Sodium ion b. Hydroxyl ion c. Lithium ion d. Hydrogen ion Answer: d. The smallest, most positively charged molecule or ion will have the greatest charge density, which leads to hydrogen ions having the greatest charge density. This greatly adds to its willingness to be hydrated by H2O molecules in water. 130.

What happens with London dispersion forces among molecules?

a. Smaller molecules have greater London dispersion forces. b. Nonpolar molecules have greater London dispersion forces. c. Larger molecules have greater London dispersion forces. d. Metals have greater London dispersion forces. Answer: c. London dispersion forces are very weak forces that mainly apply to nonpolar molecules. The larger the molecule, the greater is the London dispersion force potential in the molecule.

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131.

Van der Waals forces involve which two forces?

a. Covalent bonding forces and ionic bond forces b. Dipole-dipole forces and London dispersion forces c. Hydrogen bonding forces and dipole-ion forces d. Dipole-ion forces and ion-ion forces Answer: b. Both dipole-dipole forces and London dispersion forces are types of van der Waal forces. Covalent bonding force is not an intermolecular force and ion-ion forces, hydrogen bonding, and dipoleion forces are not, strictly speaking, types of van der Waals forces. 132.

Which molecule has the ability to connect to itself with hydrogen bonding?

a. CH4 b. C2H6 c. H2 d. CH3OH Answer: d. Only methanol or CH3OH involves a connection between a hydrogen atom and an electronegative atom (oxygen). This allows for hydrogen bonding. The others do not involve an electronegative atom. 133.

When is Raoult’s law of partial pressures in a solution most accurate to the

behavior of real solutions? a. When the solvent is not water. b. When the solvent is highly polar. c. When the solute is ionic. d. When the solution is very dilute. Answer: d. When the solution is very dilute, there are fewer intermolecular forces between the solvent and solute, so Raoult’s law is most accurate.

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134.

What happens to the freezing point and boiling point of a solution versus

the solvent when a solvent becomes a solution? a. The boiling point is increased and the freezing point is decreased. b. The boiling point and freezing point are increased. c. The boiling point and freezing point are decreased. d. The freezing point is increased and the boiling point is decreased. Answer: a. When looking at the phase diagram, the boiling point is increased and the freezing point is decreased. 135.

You add one mole of solute (nonionic) to 99 moles of water. What is the

mole fraction of water after this is done? a. 100 b. 0.99 c. 1 d. It cannot be determined Answer: b. This is calculated by adding the total moles of solute to solution, which is 100 moles. The mole fraction of water becomes 99 divided by 100 or 0.99. 136.

What is the main difference between Henry’s law and Raoult’s law when it

comes to solutions? a. Raoult’s law applies to nonideal solutions and Henry’s law applies to ideal solutions. b. Raoult’s law applies to gases in solutions and Henry’s law applies to all solutes in solution. c. Raoult’s law applies to all solutes and Henry’s law applies to gases in solutions. d. Henry’s law applies to nonreactive gases and Raoult’s law applies to reactive gases. Answer: c. Raoult’s law applies to all solutes and Henry’s law applies to gases and their solubility in liquids.

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137.

In putting ions into solution, what is true of the change in enthalpy of

separating ions and separating solvent molecules? a. The change in enthalpy will be negative for both of these. b. The change in enthalpy will be positive for ion separation but negative for separating solvent molecules. c. The change in enthalpy will be positive for both of these. d. The change in enthalpy will be negative for ion separation but positive for separating solvent molecules. Answer: c. Both of these reactions will be endothermic because it requires energy to break intermolecular bonds so the change in enthalpy will be positive for both. 138.

What will the change in enthalpy be for an ideal solution when dissolving a

solute into the solvent? a. It will be zero. b. It will be 1. c. It will be any number greater than zero. d. It will be any number less than zero. Answer: a. For an ideal solution, the change in enthalpy will be zero when a solute is dissolved into a solvent to make a solution. 139.

What determines whether or not an electrolyte goes into an aqueous

solution? a. The process must be exothermic rather than endothermic. b. The ion-dipole forces must be less than the ion-ion forces. c. The ion-dipole forces must be equal to the ion-ion forces. d. The ion-dipole forces must exceed the ion-ion forces. Answer: d. In order for the electrolyte to go into solution, the iondipole forces must exceed the ion-ion forces. Otherwise the forces would lead in the direction of keeping the ionic substance together rather than in solution.

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140.

When the mixing of a solvent and solute is exothermic, what can be said

about it? a. The division of the solute molecules must be exothermic. b. The division of the solvent molecules must be exothermic. c. The enthalpy of the dissolution process is going to be negative. d. The enthalpy of the dissolution process is going to be positive. Answer: c. The enthalpy of the dissolution process is going to be negative but it doesn’t necessarily say anything about the division of solute and solvent molecules from one another—processes that tend to be endothermic. 141.

You know the mole fraction of gas A is 0.4 and the mole fraction of gas B is

0.2. What is the mole fraction of the remaining gas in the system? a. 2 b. 0.2 c. 0.4 d. 40 Answer: c. The mole fraction of all the substances in a given mixture of substances in a system always equals one. This leads to a mole fraction of 0.4 remaining after adding the mole fractions of the other gases in the system. 142.

When is it best to use molality to describe a solution versus the molarity?

a. When temperature is a factor in the equation. b. When pressure is a factor in the equation. c. When dealing with very small units of solute. d. When dealing with gases versus liquids. Answer: a. Molarity is based on solvent volume, which is temperaturedependent; however, molality is not temperature-dependent.

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143.

What are the units of molality in determining the amount of something?

a. Moles per liter b. Parts per million c. Moles per mole d. Moles per kilogram Answer: d. The units of molality are in moles per kilogram of solvent. 144.

What are the units used to describe molarity?

a. Moles per liter b. Parts per million c. Moles per mole d. Moles per kilogram Answer: a. The units of molarity are in moles of solute per liter of solution. 145.

The boiling point elevation is based on what aspect of the molality of the

solute? a. Inversely proportional to the concentration b. Inversely proportional to the square of the concentration c. Directly proportional to the square of the concentration d. Directly proportional to the concentration Answer: d. The boiling point elevation is directly proportional to the concentration of the solute, in particular, its molality. 146.

Which ion would result in a lower freezing point depression in water?

a. CaCl2 b. NaCl c. KCl d. NaNO3 Answer: a. The colligative property of freezing point depression depends solely on the number of particles and not on the nature of the

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particle. For this reason, the effect will be greater in an ionic compound that breaks into more ions, which is the case for CaCl2 versus the other compounds. 147.

What does the change in vapor pressure when a solvent becomes a

solution most depend on? a. The nature of the solvent b. The nature of the solute c. The number of particles of solute in solution d. The baseline vapor pressure of the solvent Answer: c. The vapor pressure change will be completely dependent on the number of particles of solute in solution but will be completely independent of the natures of the solvent and solution. 148.

What range of particles constitutes a colloid or colloidal dispersion?

a. 1-10 nanometers b. 1000-10,000 micrometers c. 0.1-10 nanometers d. 1-1000 nanometers Answer: d. There is a broad range of what constitutes a colloid or colloidal dispersion. It will be 1 to 1000 nanometers, with anything larger than this being a suspension. 149.

What is the solvent when a colloid is referred to as a hydrocolloid?

a. Water b. Toluene c. Acetic acid d. Oil Answer: a. A hydrocolloid is any type of colloid that has, as the solvent, water. It can be a gas, solid, or another liquid making up the rest of the colloid.

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150.

Whipped cream in a can is a colloid. What type of colloid is it?

a. Sol b. Aerosol c. Foam d. Emulsion Answer: c. It is gas particles trapped in a liquid, so this makes it a foam colloid. 151.

What factor least affects the solubility of a solid solute in water?

a. Polarity of the solute b. Polarity of water c. Temperature d. Pressure Answer: d. Pressure is more involved in the solubility of a gas and is least affecting the solubility of a solid in water. The polarity of the solute and water along with the temperature will play a larger role. 152.

What does miscibility relate to in chemistry?

a. Vapor pressure b. Solubility c. Volatility d. Precipitation Answer: b. Miscibility is the property of a substance to mix in all proportions (that is, to fully dissolve in each other at any concentration), forming a homogeneous solution. It involves the solubility of the substance.

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153.

What happens to the solubility of gases in liquids when temperature and

pressure change? a. The solubility decreases with temperature but increases with pressure b. The solubility decreases with temperature and pressure c. The solubility increases with temperature and decreases with pressure d. The solubility increases with temperature and pressure Answer: a. The solubility decreases with temperature but increases with pressure. Think about a can of soda, which has increased solubility of the CO2 gas under greater pressure and decreased temperature conditions. 154.

Which cation or anion is least soluble in water when forming a salt?

a. Sodium b. Potassium c. Phosphate d. Chloride Answer: c. Phosphates will be least soluble in water; however, the other ions will be highly soluble in water in most cases. 155.

What is the range of oxidation numbers in chemistry?

a. 0 to 4+ b. -4 to +9 c. -8 to +8 d. -4 to 0 Answer: b. The range of oxidation numbers is from -4 to +9, although much smaller numbers, including 0 are possible.

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156.

What is the oxidation number for carbon as an element?

a. 0 b. +4 c. -4 d. It depends on how many valence electrons it has at any given point in time. Answer: a. The oxidation number of any pure element will be zero as it is not sharing electrons with anything. 157.

In the reaction of strontium plus 2 waters goes to Strontium hydroxide

Sr(OH)2 to Hydrogen gas, what is the oxidation number change of strontium. a. -2 to +2 b. 0 to +1 c. 0 to +2 d. -2 to 0 Answer: c. The oxidation number of strontium alone is going to be zero. This goes to a molecule of strontium and two hydroxide atoms. Because this effectively makes strontium a +2 ion, it increases the oxidation number from 0 to +2. 158.

What is the oxidation number of manganese in the molecule known as

Mn2O7? a. -7 b. -2 c. +2 d. +7 Answer: d. Given that the oxidation number of oxygen is going to be -2 and the fact that there are seven of them and two atoms of manganese, the oxidation number of manganese is going to be +7.

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159.

What is the oxidizing agent when the reaction is this: 2K + NH3 becomes

KNH2 + H2 gas

a. Potassium b. Ammonia c. KNH2 d. H2 gas Answer: b. In this case, ammonia becomes somewhat of an NH2 negative ion, oxidizing potassium to become somewhat of a K+ ion. This makes ammonia or NH3 an oxidizing agent. 160.

What is the conjugate reducing agent to O2 gas?

a. H2O b. H2 gas c. O2- ion d. It depends on the reaction Answer: c. The conjugate reducing agent to O2 gas is the O2- ion because the O2 gas is an oxidizing agent and the O2- ion is the reducing agent. These are conjugate agents. 161.

What type of reaction is it when unsaturated fats get saturated?

a. It can be called oxidation and dehydrogenation. b. It can be called reduction and hydrogenation. c. It can be called oxidation and hydrogenation. d. It can be called reduction and dehydrogenation. Answer: b. This is a biochemical process in which hydrogen gets added to unsaturated fats in the process of hydrogenation, which involves reduction from a chemical perspective.

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162.

Which molecule is the most reduced in organic chemistry?

a. Methanol b. Ethanol c. Methane d. 1-Propanol Answer: c. Nothing can be more reduced when it comes to carbon in organic chemistry than methane, which is CH4. The others are alcohols that, by definition, contain oxygen and therefore could be reduced more completely. 163.

What is the reduction part of cellular respiration in biological systems?

a. Glucose going to CO2 b. CO2 going to glucose c. O2 going to H2O d. H2O going to O2 Answer: c. The redox reaction that happens in cellular respiration involves the oxidation of glucose and the reduction of oxygen to make H2O and CO2. The O2 going to H2O is the reduction part of this redox reaction series. 164.

What is the purpose of antioxidants in biological species as they apply to

redox reactions? a. They help reduce oxygen to make water. b. They help to reduce free radicals to produce reduced organic molecules. c. They oxidize glucose to make energy and CO2. d. They help to kill pathogens by reducing their biological molecules. Answer: b. Antioxidants are molecules that reduce oxygen free radical species by taking on their electrons in organic reactions.

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165.

What is the conjugate oxidizing agent to sodium metal?

a. Oxygen b. H2O c. Sodium hydroxide d. Sodium ion Answer: d. The conjugate oxidizing agent for any metal will be its metal ion. This will be a very weak oxidizer because metals like sodium are strong reducers, themselves preferring their oxidized form. 166.

In the reaction when calcium plus water go to make calcium hydroxide

plus hydrogen gas, which is the main oxidizing agent? a. Hydrogen gas b. Calcium c. Water d. Calcium hydroxide Answer: c. In this reaction, as it proceeds forward, calcium is the reducing agent and water is the main oxidizing agent. 167.

Which of the following is considered a strong acid?

a. HCl b. NH4+ c. NH3 d. KCl Answer: a. HCl or hydrochloric acid is a strong acid because it easily dissociates into H+ and Cl-. NH4+ is considered a weak acid because it does dissociate but not to the degree that hydrochloric acid does. According to the Bronsted-Lowry definition, it is a very strong proton donator because Cl- is very electronegative.

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168.

What would a conjugate base be to hydrofluoric acid?

a. Sodium chloride b. Potassium hydroxide c. Hydronium ion d. Fluoride Answer: d. The conjugate base is the ion that mixed with water in order to make HF and OH-, making it a very weak conjugate base. 169.

What can be said about the acidity and alkalinity of water?

a. It is a basic substance b. It is an acidic substance c. It is both an acid and a base d. It is neither an acid nor a base Answer: c. Water is considered both an acid and a base because it generates hydrogen ions and hydroxide ions. 170.

Which is a weak base?

a. Sodium hydroxide b. Cesium hydroxide c. Methylamine d. Lithium hydroxide Answer: d. Methylamine, which is CH3NH2, will mix with water to make hydroxide ions. For this reason, it is a base but it is not a strong base as are the hydroxides associated with the Group IA metals, such as sodium, lithium, and cesium.

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171.

Which statement is not true about the theories on acids and bases?

a. Acids are considered to be proton donors. b. Bases are considered to be electron pair acceptors. c. Acids give off hydrogen ions. d. Bases are considered to be electron pair donors. Answer: b. There is the Arrhenius, Bronsted-Lowry, and Lewis theories of acids and bases. In the Lewis theory, bases are electron pair donors and are not electron pair acceptors. 172.

According to the Lewis theory of acids and bases, what makes ammonia

(NH3) a weak base? a. It can give off an OH- ion when mixed with water. b. It does not have an OH- ion associated with it so it makes a weak base. c. It can accept a hydrogen ion. d. It has a free pair of electrons to donate. Answer: d. According to the Lewis theory of acids and bases, ammonia has a free pair of electrons to donate, making it a Lewis base, even though it does not have an OH- ion attached to it. 173.

You have two aqueous solutions and one has a pH of 7 and one has a pH of

4. How many times more or less is the concentration of hydrogen ions in the pH 7 solution versus the pH 4 solution? a. 100 times less concentrated b. 1000 times less concentrated c. 100 time more concentrated d. 3 times more concentrated Answer: b. The scale is negative and logarithmic with respect to the concentration of hydrogen ions so a pH of 7 is 1000 times less concentrated with respect to hydrogen when compared to a pH of 3.

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174.

Which molecule is a Bronsted-Lowry base?

a. NH4+ b. HP3O4 c. HPO42d. CH3NH3+ Answer: c. Bases are considered proton acceptors and acids are considered proton donators (according to the Bronsted-Lowry theory). Each of these will donate a proton except for HPO42- breaks into PO43- plus H+, making it acidic under this theory. 175.

What is considered a neutral substance under the pH scale?

a. 0 b. 1 c. 7 d. 14 Answer: c. A pH of 7 has equal concentrations of hydrogen and hydroxide ions, making it completely neutral. 176.

When doing a titration reaction, when is the equivalence point when

titrating an acid with a strong base? a. At the point when there is a pH of 7. b. When a slight change in volume results in a large pH increase. c. When a slight change in volume results in a large pH decrease. d. When all of the hydroxide ion is used up. Answer: b. In a titration reaction, small amounts of base are added at a time; when a slight change in volume results in a large pH increase, this will be the equivalence point in a titration curve. It will not necessarily be at pH 7 when this occurs.

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177.

In doing a titration reaction, what is it that you are usually trying to

determine? a. The molarity of the solution you are titrating with b. The volume of the solution you are titrating with c. The molarity of the solution you are titrating d. The volume of the solution you are titrating Answer: c. Usually, you know the volume of the solution you are trying to titrate and you can figure out the volume of solution you are titrating with as well as its molarity. This means that what you are trying to determine is the molarity you are titrating. From this, you can determine what the substance is because you can know the number of grams you’ve added to make the solution in the first place; you can also know the number of grams of solute is in the solution after knowing its molecular weight and molarity. 178.

What is the purpose of phenolphthalein in acid-base chemistry?

a. It can determine the sudden change in pH through a color change at the equivalence point. b. It can determine the exact pH of a strong acid. c. It can determine the exact pH of a strong base. d. It can be used to determine the molarity of a weak acid or base. Answer: a. This is a pH indicator that will change color at the equivalence point in a titration reaction, making it a good additive to these types of reactions.

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179.

If the pH of a solution in water is 4, what is the pOH?

a. 8 b. 10 c. 11 d. 13 Answer: b. The pH plus the pOH is always going to be 14. With this in mind, if the pH is equal to 4, the pOH is going to be 10 because 14 minus 4 equals 10. 180.

Weak acid buffers like acetic acid tend to work to maintain the pH until

how much of the acid is remaining? a. 10 percent b. 40 percent c. 65 percent d. 90 percent Answer: a. This type of buffer will work until just about 10 percent of the acid/salt pair is still acetic acid. Anything less than this will result in a marked pH change. 181.

Which cation is considered to be a neutral cation (having no effect on the

pH of a solution)? a. Pb(II) b. Al(III) c. NH4+ d. Ca2+ Answer: d. Each of these is an acidic cation, being acidic as a salt is made from it, except for calcium ion, which is a neutral cation.

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182.

What is the voltage of a system?

a. The same as amperes b. The same as amperes per second c. The same as amperes multiplied by the resistance d. The same as the current divided by the charge Answer: c. Voltage depends on the resistance involved and is the same as the amperes multiplied by the resistance. 183.

In an electrochemical reaction, the electrons go where?

a. They are shared with the oxidized and reduced substance. b. They arrive at the oxidized substance. c. They leave the reduced substance. d. They leave the oxidized substance and arrive at the reduced substance. Answer: d. In an electrochemical reaction, the electrons leave the oxidized substance and arrive at the reduced substance. 184.

What type of reaction is happening in an electrochemical cell?

a. Combustion reaction b. Acid-base reactions c. Combination reactions d. Redox reactions Answer: d. Electrochemical cells are all about redox reactions because they actually involve the transfer of electrons from one part of the cell to another.

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185.

What statement is not true of a galvanic cell?

a. It requires a battery as part of the cell. b. Electricity can be demonstrated through a voltmeter. c. It involves a spontaneous redox reaction. d. It involves an anode and a cathode. Answer: a. Galvanic cells do not require a battery but make electricity instead that is derived from the spontaneous redox reaction. 186.

What happens between the anode and the cathode in a galvanic cell?

a. Ions pass from the anode to the cathode b. Ions pass from the cathode to the anode c. Electrons pass from the anode to the cathode d. Electrons pass from the cathode to the anode Answer: c. In such cases, it is the electrons that pass from one point to another, in this case, from the anode to the cathode. 187.

What is the main purpose of the electrolyte solution in an electric cell?

a. The main purpose is to maintain electrical neutrality in the cell. b. The main purpose is to allow ions to flow from the anode to the cathode. c. The main purpose is to accept electrons in the anode side. d. The main purpose is to accept electrons in the cathode side. Answer: a. The electrolyte solution contains both the anode and the cathode. Its main purpose is to maintain electrical neutrality in the cell. 188.

What is not true about the salt bridge in a voltaic cell?

a. Ions can pass in either direction in the cell. b. The ions can be any nonreactive salt solution. c. The negative ions are attracted to the anode. d. The ions cannot precipitate out of solution. Answer: c. The negative ions are attracted to the cathode and not the anode because the cathode is positively charged and the anode is

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negatively charged. The ions can be any nonreactive salt and the ions can go in any direction. They don’t all flow to the anode side or the cathode side. 189.

A reaction goes from tin solid to nitric oxide (NO) gas in a galvanic cell

that uses a platinum metal strip in the nitric oxide side. What is the cathode in this situation? a. The tin b. The platinum c. The nitric oxide gas d. The salt bridge Answer: a. Anything that donates electrons will be the cathode because it will be positively charged. If the reaction goes in a positive direction (which it does in a galvanic cell), the cathode will be the tin. 190.

The standard EMF or standard electromotive force of a cell depends on all

of these things except for what? a. Being at standard temperature b. Being an ideal solution c. Being at a one molar concentration d. Being at 0 degrees Kelvin Answer: d. The standard EMF of a given cell involves being at a standard (25 degrees) temperature, being an ideal solution, and being at a one molar concentration. It will not be standard at 0 degrees Kelvin.

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191.

What most drives the EMF of a cell to make the cell galvanic rather than

electrolytic? a. The relative reduction ability of the two metals. b. The molar mass of the two metals. c. The atomic weight of the two metals. d. The type of salt in the salt bridge of the electric cell. Answer: a. The relative ability of the two metals to be reduced determines the EMF of the cell. If there is a big difference, the electromotive force or EMF will be greater in the cell. 192.

How many carbon molecules are there in the hydrocarbon called butane?

a. One b. Two c. Three d. Four Answer: d. Butane has four carbon atoms and 2 multiplied by 4 plus 2 or 10 hydrogen atoms around this carbon chain. 193.

If there are 6 carbon atoms in a fully saturated hydrocarbon, how many

hydrogen atoms will there be around these? a. 10 b. 14 c. 16 d. 20 Answer: b. A six-carbon chain that is fully saturated will have a total of 14 hydrogen atoms around it. This fits with the 2 multiplied by n + 2 arrangement of hydrogen to carbon atoms.

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194.

What shape is benzene chemically?

a. Triangular b. Square c. Hexagon d. Octagon Answer: c. Benzene is a carbon ring that is in the shape of a hexagon. 195.

What type of isomerism involves mirror images of molecules?

a. Stereoisomerism b. Functional isomerism c. Position isomerism d. Chain isomerism Answer: a. Each of these is a structural isomer, except for stereoisomers, which can involve mirror images of different, otherwise identical, molecules. 196.

An organic molecule that has a halogen added to it as a side chain or

functional group can be called what? a. Phenol b. Alkyl halide c. Aldehyde d. Alcohol Answer: b. An alkyl halide is an alkene or similar compound that has Bromine, Chlorine, or Iodine added to it making it having a halogen as an R group.

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197.

An aromatic compound with an OH side chain is called what kind of

compound? a. Aldehyde b. Alcohol c. Phenol d. Ketone Answer: c. Phenol in its most basic form is a benzene (aromatic) ring plus an alcohol or OH side chain. 198.

How many other side chains are available to carbon when a CN functional

group is attached to the molecule? a. Zero b. One c. Two d. Three Answer: b. Because there is a triple bond between nitrogen and carbon (by the virtue of nitrogen having three valence spots available), this leaves only one spot available for an R chain attaching to the carbon chain. 199.

What would the configuration be in a benzene molecule if the side groups

are on the 1,4-carbon atoms on the molecule? a. Ortho b. Meta c. Para d. Trans Answer: c. The para configuration involves an aromatic molecule in a 1,4-carbon configuration. This means that 1,4-dimethylbenzene is also called para-dimethylbenzene or referred to colloquially as p-xylene.

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200. Which molecule is a disaccharide ? a. Glucose b. Galactose c. Sucrose d. Fructose Answer: c. Sucrose is a disaccharide made from glucose and fructose. The rest are monosaccharides that cannot be broken down any further except through glycolysis and cellular respiration.

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College Level Chemistry

Articles inside

Summary

Quiz

Key Takeaways

Lipids

Carbohydrates

Nucleic Acids

Proteins

Benzene Derivatives

Basics of Biochemistry

Structural Isomerism

Isomerism

Alkenes and Alkynes

Nernst Equation

Quiz

Key Takeaways

Cycloalkanes

Electromotive Force

Quiz

Key Takeaways

Faraday’s Law

Hydrolysis

Buffers

pH Scale

Quiz

Quiz

Key Takeaways

Redox Reactions in Common Situations

Key Takeaways

Crystals

Colloids

Anomalous Colligative Properties

Colligative Properties

Quiz

Liquid Forces

Liquids

Water Condensation, Boiling, and Evaporation

Key Takeaways

Chemical Equilibrium

Quiz

Energy of Activation

Rates of Reactions

Limiting Reagents

Writing Reactions

Types of Chemical Reactions

Quiz

Key Takeaways

Hydrogen Bonding

Bonding in Metals

Shapes of Molecules

Covalence

Molecular Orbital Theory

Quiz

Key Takeaways

Rules of Thermochemistry

Enthalpy and Energy

Calorimetry

Heat Capacity

Laws of Thermodynamics

Properties of Heat in Chemistry

Quiz

Graham’s Law of Effusion

Key Takeaways

Kinetic Theory

Partial Pressures in Gases

Boyle’s Gas Law

Gas laws

Pressures and Gases

Quiz

Magnetic Properties in Atoms

Electronegativity

Key Takeaways

Electron Affinity

Quiz

Ionization Energy

Atomic Mass Number

Equivalent Weight and Mole Ratio

Isotopes

Key Takeaways