Young Ji International School / College
Page 1
Chapter 1
Chemistry and Technology
Chemistry is so central in our lives that we read about it in the news every day. You will see this in the following recent breakthroughs. Technology is the application of science to improve the quality of human life. Cell phones, compact discs, a variety of processed food, and the Internet are some of the products of technology. Technology can bring about even a change in one‘s physical appearance. Our for cosmetic procedures launched some years ago. Cloning has emerged gradually over the past decade. Each advance has been startling enough, prompting ethical debates.
A Flashback 1. Alchemy Is an influential philosophical tradition whose practitioners have, from antiquity, claimed it to be the precursor to profound powers. The defining objectives of alchemy are varied but historically have typically included one or more of the following goals: the creation of the fabled philosopher's stone; the ability to transmute base metals into the noble metals (gold or silver); and development of an elixir of life, which would confer youth and longevity. 2. Iatrochemistry (or chemical medicine) Is a branch of both chemistry and medicine. The word "Iatro" was the Greek word "doctor" or "medicine." Having its roots in alchemy, iatrochemistry seeks to provide chemical solutions to diseases and medical ailments. This area of science has fallen out of use in Europe since the rise of modern establishment medicine. However, iatrochemistry was popular between 1525 and 1660, especially in Flanders. Its most notable leader was Paracelsus, an important Swiss alchemist of the 16th century. Iatrochemists believed that physical health was dependent on a specific balance of bodily fluids. Agrochemical therapies and concepts are still in wide use in South Asia, East Asia and amongst their diasporas communities worldwide. 3. The phlogiston theory is an obsolete scientific theory that postulated a firelike element called phlogiston, contained within combustible bodies and released during combustion. The name comes from the Ancient Greek υλογιστόν phlogistón (burning up), from υλόξphlóx (flame). It was first stated in 1667 by Johann Joachim Becher. The theory attempted to explain burning processes such as combustion and rusting, which are now collectively known as oxidation. 4. The history of chemistry represents a time span from ancient history to the present. By 1000 BC, civilizations used technologies that would eventually form the basis to the various branches of chemistry. Examples include Young Ji International School / College
Page 2
extracting metals from ores, making pottery and glazes, fermenting beer and wine, extracting chemicals from plants, for medicine and perfume, rendering fat into soap, making glass, and making alloys like bronze. The protoscience of chemistry, alchemy, was unsuccessful in explaining the nature of matter and its transformations. However, by performing experiments and recording the results, alchemists set the stage for modern chemistry. The distinction began to emerge when a clear differentiation was made between chemistry and alchemy by Robert Boyle in his work The Sceptical Chymist (1661). While both alchemy and chemistry are concerned with matter and its transformations, chemists are seen as applying scientific method to their work. Chemistry is considered to have become an established science with the work of Antoine Lavoisier, who developed a law of conservation of mass that demanded careful measurement and quantitative observations of chemical phenomena. The history of chemistry is intertwined with the history of thermodynamics, especially through the work of Willard Gibbs. The Atomic Bomb 1. Enrico Fermi (Italian: [enˈri.ko ˈfeɾ.mi]; 29 September 1901 – 28 November 1954) was an Italian physicist, best known for his work on Chicago Pile-1 (the first nuclear reactor), and for his contributions to the development of quantum theory, nuclear and particle physics, and statistical mechanics. He is one of the men referred to as the "father of the atomic bomb".[4] Fermi held several patents related to the use of nuclear power, and was awarded the 1938 Nobel Prize in Physics for his work on induced radioactivity by neutron bombardment and the discovery of transuranic elements. He was widely regarded as one of the very few physicists to excel both theoretically and experimentally. 2. Lise Meitner- coined the term nuclear fission to mean the splitting of the atomic nucleus. 3. Julius Robert Oppenheimer (April 22, 1904 – February 18, 1967) was an American theoretical physicist and professor of physics at the University of California, Berkeley. He is among the persons who are often called the "father of the atomic bomb" for their role in the Manhattan Project, the World War II project that developed the first nuclear weapons. The first atomic bomb was detonated on July 16, 1945, in the Trinity test in New Mexico; Oppenheimer remarked later that it brought to mind words from the Bhagavad Gita: "Now I am become Death, the destroyer of worlds." Polymer Industry 1. Wallace H. Carothers- an American DuPont chemist led a research team that invented nylon. 2. Karl Ziegler- a German chemist shared a Nobel Prize with Guilio Natta on their work with polymers, high density polyethylene or high impact plastics. 3. Charles Goodyear- made serenities discovery of vulcanization of rubber. Young Ji International School / College
Page 3
Discussion Box 1. In your own opinion, which technology helps society most and why? a. Fiber optics b. Neon lights c. Vulcanization of rubber d. Polymerization to form plastic wares 2. Tell how is chemistry related to other sciences like: a. Agriculture b. Food technology c. Medicine d. Pharmacology e. Biology f. Geology
Chapter 2
Measurement
You must have measured quantitative such as temperature, length and volume before. When you take a measurement, you always compare it with a reference standard. The International System of Units (SI) For a long time, the metric system was the standard of measurement. But in 1960, a system of units called the International System of Units (SI) was established by the 11th general Conference on Weights and Measures. Derived from the French words System Internationale, the SI is built upon a set of seven metric units called the base units of SI.
Some SI and Non-SI Units of Measurements Quantity SI Unit Non-SI Unit Length meter foot Volume cubic meter (m3) liter (L) Mass kilogram pound (lb) Density gram per cubic centimeter pound per Cubin inch (g/cm3) (lb/in3) gram per milliliter (g/mL) Temperature Kelvin degree Celsius (°C) Time second (s) hour (h) Pressure Pascal (Pa) atmosphere Energy joule (J) calorie (cal) Mass and Weight The amount of matter in an object is called its mass. In SI, the unit of mass is kilogram (kg). in chemistry, a kilogram is relatively large unit of measurement. The more commonly used unit of mass is gram (g) which is 1/1000 of kilogram.
Young Ji International School / College
Page 4
The weight of an object is equal to the force of gravity on the object. This varies with location. For example, a 1.00-kg object in Baguio City would weight about 0.99 kg in Metro manila although the mass is the same in both locations. Temperature Temperature measures the hotness or coldness of an object. It also determines the directions of the flow of heat. A glass of hot milk fells hot because heat is transferred from the hot glass to your hand. Heat always flows from an object at higher temperature to another object at lower temperature. There are three temperature scales being used. These are the Celsius (formerly Centigrade), the Fahrenheit, and the Kelvin scales. Their units are °C (degrees Celsius), °F (degrees Fahrenheit), and K (kelvins). Two commonly-fixed points are the temperature at which water freezes and the temperature at which water boils at standard atmospheric pressure. On the Celsius scale, the freezing point of water is 0°C and its boiling point is 100°C. The interval between these points is divided into 100 equal parts. On the Fahrenheit scale, the freezing point of water is 32°F and the boiling point is 212°F. The interval between these points is divided into 180 equal parts. On the Kelvin scale, the freezing point of water is 273 K and the boiling point is 373 K. Notice that on the Kelvin scale is the same as in the Celsius scale. The zero point on the Kelvin scale, 0-K, is called the absolute zero temperature and its equal to -273 °C. To express °C in °F, use T °F = 9/5 T °C + 32 To express °F in °C, use
T °F = 9/5 (T °C - 32) To express °C in K, use TK = T °C + 273 Example 1 The normal body temperature is 37°C. What is this in a Fahrenheit thermometer? T °F = 9/5 T °C + 32 = (37) + 32 = 66.6 + 32 Example 2
T °F = 98.6 °F
The temperature for roasting chicken is at 350°F. What is this temperature on the Celsius scale? Young Ji International School / College
Page 5
T °C = (T°F -32) = (350 -32) = (318) T °C = 177°C Exercise Mark has a high fever with a body temperature of 104°F. What is his temperature in? a. °C? __________________________ b. K? _________________________ Scientific Notations Scientific notation (also referred to as "standard form" or "standard index form") is a way of writing numbers that are too big or too small to be conveniently written in decimal form. Scientific notation has a number of useful properties and is commonly used in calculators and by scientists, mathematicians and engineers. Standard decimal notation Scientific notation 2
2×100
300
3×102
4,321.768
4.321768×103
−53,000
−5.3×104
6,720,000,000
6.72×109
0.2
2×10−1
0.000 000 007 51
7.51×10−9
In scientific notation all numbers are written in the form
(a times ten raised to the power of b), where the exponent b is an integer, and the coefficient a is any real number (however, see normalized notation below), called the significand or mantissa. The term "mantissa" may cause confusion, Young Ji International School / College
Page 6
however, because it can also refer to the fractional part of the common logarithm. If the number is negative then a minus sign precedes a (as in ordinary decimal notation). Decimal floating point is a computer arithmetic system closely related to scientific notation.
Exercise A. Change the following to scientific notation 1. 1000 2. 0.000 706 3. 0.000 000 2 4. 36 000 000 B. Convert the following to whole number or decimal form: 1. 3.0 x 104 2. 2,162 x 10-5 3. 1.03 x 105 4. 1.17 x 10 -13 5. 1.6 x 104
Chapter 3 Derived Units The combinations of these units are called derived units. For example, speed is defined s the distance divided by the time required to travel that distance. Thus the SI unit for speed is meters per second (m/s) An important measurement usually used in a chemistry laboratory is the volume. Volume has the unit (length)3; thus, the basic unit of volume is cubic meters (m3). However, this is a very large unit, so a more commonly used volume unit is the cubic centimeter (cm3). Another unit of volume is the liter (L). There are 1000 milliliters (mL) in a liter and 1 mL is equal to cm3. Another commonly derived unit is density. Density is defined as the amount of mass of an object divided by its volume.
Density = mass Volume
or
d=m v
The SI derived unit for density is kilogram per cubic meter (kg/m 3). But this unit is too large for most applications, so density is usually expressed in units of grams per cubic centimeter (g/cm3. Densities of Some Common Substances Air Young Ji International School / College
Density (g/cm3) 0.0013 Page 7
Feather Ice Water Brick Aluminum Steel Silver Gold
0.0025 0.92 1.00 1.84 2.70 7.80 10.50 19.30
Specific gravity, another derived unit, is the ratio of the density of a substance to the density of a reference substance at the same temperature. Water is commonly used as the reference substance. Note that the units in the ratio cancel, thus specific gravity has no units. Specific gravity is a common test used in hospitals, such as in urine samples for the diagnosis of certain aliments like diabetes. It is also used to determine the amount of acid in car batteries and the condition of antifreeze car radiations.
Precision and Accuracy All measurements from any measuring device are never exact. The degree of error depends on the measuring instruments and the person making the measurement. There are two ways of checking these errors. Precision is one way of checking measurement. It tells how close several measurements are to the same value. When this happens, the measurement is said to be precise, but not necessarily. Example There students measure the mass of a piece of coppers that weights 3.000 g.
Trial 1 Trial 2 Trial 3 Trial 4
Student A 2.970 g 2.971 g 2.976 g 2. 969 g
Student B 3.000 g 3.001 g 3.002 g 3.001 g
Student C 2.972 g 2.965 g 2. 985 g 2.974 g
If there is a significant difference among the measurements, the precision is low (or poor). If the measurements differ in a small amounts, the precision is high) or good). Thus students B have most precise measurements. Accuracy is another way of checking measurements. It tells how close a measurement is to the true or accepted value. A measurement that is accurate, or has a high accuracy is one that is close to the true value. So, in the table, the student B also has the most accurate measurements being closet to the true value of 3.000 g.
Young Ji International School / College
Page 8
Significant Figures In any measurements, there is always some degree of uncertainty due to experimental error. For example, the mass determined using a simple balance is 4.82 g. in both instances, there is a degree of uncertainty in the final (last) digit-the 8 in 4.8 and the 2 in 4.82. The preceding digits are certain or definite and the also digits are uncertain or estimations. To determine the number of significant figures obtained from measurements, follow these rules. 1. Any digit that is not zero is significant. Thus, 125 has 3 significant figures; 1234 has 4 significant figures. 2. Zeros between nonzero digits are significant. Thus 105 has 3 significant figures; 1 001 has 4 significant figures. 3. Zeros at the beginning of a number are not significant. They merely indicate the position of the decimal point. Thus, 0.06 has 1 significant figure; 0.0045 has 2 significant figures. 4. Zero at the end of a number and the after the decimal point are significant. Thus, 0.10 has 2 significant figures; 42.00 has 5 significant figures; 0.00900 has 3 significant figures. 5. Zeros at the end of a number may or may not be significant. Thus, 300 may have 1 significant figure (3), 2 significant figures (30), or 3 significant figures (300) depending on the sensitivity of the measuring instruments. We use scientific notations to clear this ambiguity.
Example Determine the number of significant figures in each of the following measurements: a. b. c. d. e.
25 mL 5 002 mm 0.050 kg 95.00 m 0.8007 g
2 significant figures 4 significant figures 2 significant figures 4 significant figures 4 significant figures
Exercise
a. b. c. d. e.
Round off 24.58165 to indicated number of significant figures. 6 significant figures _________________ 5 significant figures _________________ 4 significant figures _________________ 3 significant figures _________________ 2 significant figures _________________
Scientific Notation Many numbers used in chemistry are either very large or very small that the number of zeros becomes difficult to handle. For accuracy and the convenience, these numbers are written in scientific notation. A number in scientific notation has Young Ji International School / College
Page 9
two parts. The first is the number between 1 and 10 (N). the second part is a power 10 (10y) as shown below N x 10y The exponent y can be a positive or a negative integer. To write numbers in scientific notation, move the decimal point such that you have a number between 1 and 10. Determine the correct exponent by counting the number of times the decimal point is moved. When the decimal point is moved to the right, the exponent is written as a negative number. When the decimal point is moved to the left, the exponent is written as a positive number. Example 0.00567
5.67 x 10-3
456000
4.56 x 105
Exercise Write the following numbers in scientific notation: a. 52 400 _______________________________ b. 1 005 000 ____________________________ c. 0.000 543 ____________________________ d. 0.000 000 870 ________________________ This is a compendium of temperature conversion formulas and comparisons among eight different temperature scales, several of which have long been obsolete. Celsius (centigrade) from Celsius
Fahrenheit [°F] = [°C] × 9⁄5 + 32
Kelvin
Rankine
[K] = [°C] + 273.15
to Celsius
[°C] = ([°F] − 32) × 5⁄9
[°C] = [K] − 273.15
[°R] = ([°C] + 273.15) × 9⁄5 [°C] = ([°R] − 491.67) × 5⁄9
Delisle
[°De] = (100 − [°C]) × 3⁄2
[°C] = 100 − [°De] × 2⁄3
Newton
[°N] = [°C] × 33⁄100
[°C] = [°N] × 100⁄33
Réaumur [°Ré] = [°C] × 4⁄5 Young Ji International School / College
[°C] = [°Ré] × 5⁄4 Page 10
[°Rø] = [°C] × 21⁄40 + 7.5
Rømer
[°C] = ([°Rø] − 7.5) × 40⁄21
Fahrenheit from Fahrenheit
Celsius [°C] = ([°F] − 32) × 5⁄9
Kelvin
[K] = ([°F] + 459.67) × 5⁄9
Rankine [°R] = [°F] + 459.67
to Fahrenheit
[°F] = [°C] × 9⁄5 + 32
[°F] = [K] × 9⁄5 − 459.67
[°F] = [°R] − 459.67
Delisle [°De] = (212 − [°F]) × 5⁄6
[°F] = 212 − [°De] × 6⁄5
Newton [°N] = ([°F] − 32) × 11⁄60
[°F] = [°N] × 60⁄11 + 32
Réaumur [°Ré] = ([°F] − 32) × 4⁄9
[°F] = [°Ré] × 9⁄4 + 32
Rømer [°Rø] = ([°F] − 32) × 7⁄24 + 7.5 [°F] = ([°Rø] − 7.5) × 24⁄7 + 32 Kelvin from Kelvin
Celsius
[°C] = [K] − 273.15
Fahrenheit [°F] = [K] × 9⁄5 − 459.67
Rankine
[°R] = [K] × 9⁄5
Young Ji International School / College
to Kelvin
[K] = [°C] + 273.15
[K] = ([°F] + 459.67) × 5⁄9
[K] = [°R] × 5⁄9
Page 11
Delisle
[°De] = (373.15 − [K]) × 3⁄2
[K] = 373.15 − [°De] × 2⁄3
Newton
[°N] = ([K] − 273.15) × 33⁄100
[K] = [°N] × 100⁄33 + 273.15
Réaumur [°Ré] = ([K] − 273.15) × 4⁄5
Rømer
[K] = [°Ré] × 5⁄4 + 273.15
[°Rø] = ([K] − 273.15) × 21⁄40 + 7.5 [K] = ([°Rø] − 7.5) × 40⁄21 + 273.15
Rankine from Rankine
Celsius
[°C] = ([°R] − 491.67) × 5⁄9
Fahrenheit [°F] = [°R] − 459.67
to Rankine
[°R] = ([°C] + 273.15) × 9⁄5
[°R] = [°F] + 459.67
Kelvin
[K] = [°R] × 5⁄9
[°R] = [K] × 9⁄5
Delisle
[°De] = (671.67 − [°R]) × 5⁄6
[°R] = 671.67 − [°De] × 6⁄5
Newton
[°N] = ([°R] − 491.67) × 11⁄60
[°R] = [°N] × 60⁄11 + 491.67
Réaumur [°Ré] = ([°R] − 491.67) × 4⁄9
[°R] = [°Ré] × 9⁄4 + 491.67
Rømer
[°Rø] = ([°R] − 491.67) × 7⁄24 + 7.5 [°R] = ([°Rø] − 7.5) × 24⁄7 + 491.67
Newton from Newton
Young Ji International School / College
to Newton
Page 12
Celsius
[°C] = [°N] × 100⁄33
Fahrenheit [°F] = [°N] × 60⁄11 + 32
[°N] = [°C] × 33⁄100
[°N] = ([°F] − 32) × 11⁄60
Kelvin
[K] = [°N] × 100⁄33 + 273.15 [°N] = ([K] − 273.15) × 33⁄100
Rankine
[°R] = [°N] × 60⁄11 + 491.67 [°N] = ([°R] − 491.67) × 11⁄60
Delisle
[°De] = (33 − [°N]) × 50⁄11
Réaumur [°Ré] = [°N] × 80⁄33
Rømer
[°Rø] = [°N] × 35⁄22 + 7.5
[°N] = 33 − [°De] × 11⁄50
[°N] = [°Ré] × 33⁄80
[°N] = ([°Rø] − 7.5) × 22⁄35
Réaumur from Réaumur
Celsius
[°C] = [°Ré] × 5⁄4
Fahrenheit [°F] = [°Ré] × 9⁄4 + 32
to Réaumur
[°Ré] = [°C] × 4⁄5
[°Ré] = ([°F] − 32) × 4⁄9
Kelvin
[K] = [°Ré] × 5⁄4 + 273.15 [°Ré] = ([K] − 273.15) × 4⁄5
Rankine
[°R] = [°Ré] × 9⁄4 + 491.67 [°Ré] = ([°R] − 491.67) × 4⁄9
Delisle
[°De] = (80 − [°Ré]) × 15⁄8 [°Ré] = 80 − [°De] × 8⁄15
Newton
[°N] = [°Ré] × 33⁄80
Young Ji International School / College
[°Ré] = [°N] × 80⁄33
Page 13
Rømer
[°Rø] = [°Ré] × 21⁄32 + 7.5 [°Ré] = ([°Rø] − 7.5) × 32⁄21
Rømer from Rømer
Celsius
[°C] = ([°Rø] − 7.5) × 40⁄21
Fahrenheit [°F] = ([°Rø] − 7.5) × 24⁄7 + 32
to Rømer
[°Rø] = [°C] × 21⁄40 + 7.5
[°Rø] = ([°F] − 32) × 7⁄24 + 7.5
Kelvin
[K] = ([°Rø] − 7.5) × 40⁄21 + 273.15 [°Rø] = ([K] − 273.15) × 21⁄40 + 7.5
Rankine
[°R] = ([°Rø] − 7.5) × 24⁄7 + 491.67 [°Rø] = ([°R] − 491.67) × 7⁄24 + 7.5
Delisle
[°De] = (60 − [°Rø]) × 20⁄7
[°Rø] = 60 − [°De] × 7⁄20
Newton
[°N] = ([°Rø] − 7.5) × 22⁄35
[°Rø] = [°N] × 35⁄22 + 7.5
Réaumur [°Ré] = ([°Rø] − 7.5) × 32⁄21
[°Rø] = [°Ré] × 21⁄32 + 7.5
Using Scientific Method The scientific method is a body of techniques for investigating phenomena, acquiring new knowledge, or correcting and integrating previous knowledge. To be termed scientific, a method of inquiry must be based on empirical and measurable evidence subject to specific principles of reasoning. The Oxford English Dictionary defines the scientific method as "a method or procedure that has characterized natural science since the 17th century, consisting in systematic observation, measurement, and experiment, and the formulation, testing, and modification of hypotheses." The chief characteristic which distinguishes the scientific method from other methods of acquiring knowledge is that scientists seek to let reality speak for itself, supporting a theory when a theory's predictions are confirmed and challenging a Young Ji International School / College
Page 14
theory when its predictions prove false. Although procedures vary from one field of inquiry to another, identifiable features distinguish scientific inquiry from other methods of obtaining knowledge. Scientific researchers propose hypotheses as explanations of phenomena and design experimental studies to test these hypotheses via predictions which can be derived from them. These steps must be repeatable to guard against mistake or confusion in any particular experimenter. Theories that encompass wider domains of inquiry may bind many independently derived hypotheses together in a coherent, supportive structure. Theories, in turn, may help form new hypotheses or place groups of hypotheses into context. Scientific inquiry is intended to be as objective as possible in order to minimize bias. Another basic expectation is the documentation, archiving and sharing of all data collected or produced and of the methodologies used so they may be available for careful scrutiny and attempts by other scientists to reproduce and verify them. This practice, known as full disclosure, also means that statistical measures of their reliability may be made. The scientific method is an orderly plan of thinking and doing things most especially in solving a problem.
1. 2. 3. 4. 5.
The following are steps in the scientific method. Knowing and understanding the problem. Finding ways to solve your problem; Selecting the best solution by reasoning; Conducting an experiment to see if the solution is correct; and Verifying your result by further experiments.
1. Process The overall process involves making conjectures (hypotheses), deriving predictions from them as logical consequences, and then carrying out experiments based on those predictions to determine whether the original conjecture was correct. There are difficulties in a formulaic statement of method, however. Though the scientific method is often presented as a fixed sequence of steps, they are better considered as general principles. Not all steps take place in every scientific inquiry (or to the same degree), and are not always in the same order. As noted by William Whewell (1794–1866), "invention, sagacity, [and] genius" are required at every step. 2. Formulation of a question The question can refer to the explanation of a specific observation, as in "Why is the sky blue?", but can also be open-ended, as in "How can I design a drug to cure this particular disease?" This stage frequently involves looking up and Young Ji International School / College
Page 15
evaluating evidence from previous experiments, personal scientific observations or assertions, and/or the work of other scientists. If the answer is already known, a different question that builds on the previous evidence can be posed. When applying the scientific method to scientific research, determining a good question can be very difficult and affects the final outcome of the investigation. 3. Hypothesis A hypothesis is a conjecture, based on knowledge obtained while formulating the question that may explain the observed behavior of a part of our universe. The hypothesis might be very specific, e.g., Einstein's equivalence principle or Francis Crick's "DNA makes RNA makes protein‖, or it might be broad, e.g., unknown species of life dwell in the unexplored depths of the oceans. A statistical hypothesis is a conjecture about some population. For example, the population might be people with a particular disease. The conjecture might be that a new drug will cure the disease in some of those people. Terms commonly associated with statistical hypotheses are null hypothesis and alternative hypothesis. A null hypothesis is the conjecture that the statistical hypothesis is false, e.g., that the new drug does nothing and that any cures are due to chance effects. Researchers normally want to show that the null hypothesis is false. The alternative hypothesis is the desired outcome, e.g., that the drug does better than chance. A final point: a scientific hypothesis must be falsifiable, meaning that one can identify a possible outcome of an experiment that conflicts with predictions deduced from the hypothesis; otherwise, it cannot be meaningfully tested. 4. Prediction This step involves determining the logical consequences of the hypothesis. One or more predictions are then selected for further testing. The more unlikely that a prediction would be correct simply by coincidence, then the more convincing it would be if the prediction were fulfilled; evidence is also stronger if the answer to the prediction is not already known, due to the effects of hindsight bias . Ideally, the prediction must also distinguish the hypothesis from likely alternatives; if two hypotheses make the same prediction, observing the prediction to be correct is not evidence for either one over the other. (These statements about the relative strength of evidence can be mathematically derived using Bayes' Theorem.) 5. Testing This is an investigation of whether the real world behaves as predicted by the hypothesis. Scientists (and other people) test hypotheses by conducting experiments. The purpose of an experiment is to determine whether observations of the real world agree with or conflict with the predictions derived from an hypothesis. If they agree, confidence in the hypothesis increases; otherwise, it decreases. Agreement does not assure that the hypothesis is true; Young Ji International School / College
Page 16
future experiments may reveal problems. Karl Popper advised scientists to try to falsify hypotheses, i.e., to search for and test those experiments that seem most doubtful. Large numbers of successful confirmations are not convincing if they arise from experiments that avoid risk. Experiments should be designed to minimize possible errors, especially through the use of appropriate scientific controls. For example, tests of medical treatments are commonly run as double-blind tests. Test personnel, who might unwittingly reveal to test subjects which samples are the desired test drugs and which are placebos, are kept ignorant of which are which. Such hints can bias the responses of the test subjects. Furthermore, failure of an experiment does not necessarily mean the hypothesis is false. Experiments always depend on several hypotheses, e.g., that the test equipment is working properly, and a failure may be a failure of one of the auxiliary hypotheses. Experiments can be conducted in a college lab, on a kitchen table, at CERN's Large Hadron Collider, at the bottom of an ocean, on Mars (using one of the working rovers), and so on. Astronomers do experiments, searching for planets around distant stars. Finally, most individual experiments address highly specific topics for reasons of practicality. As a result, evidence about broader topics is usually accumulated gradually. 6. Analysis This involves determining what the results of the experiment show and deciding on the next actions to take. The predictions of the hypothesis are compared to those of the null hypothesis, to determine which is better able to explain the data. In cases where an experiment is repeated many times, a statistical analysis such as a chi-squared test may be required. If the evidence has falsified the hypothesis, a new hypothesis is required; if the experiment supports the hypothesis but the evidence is not strong enough for high confidence, other predictions from the hypothesis must be tested. Once a hypothesis is strongly supported by evidence, a new question can be asked to provide further insight on the same topic. Evidence from other scientists and experience are frequently incorporated at any stage in the process. Depending on the complexity of the experiment, much iteration may be required to gather sufficient evidence to answer a question with confidence, or to build up many answers to highly specific questions in order to answer a single broader question. Areas of Investigation 1. Devise a shortcut method for converting temperature from Celsius scale to Fahrenheit scale mentally. 2. Make a comparative diagram of the Celsius and Fahrenheit scales of the thermometer with the calibrations going down as low -40 C in both scales. Inform the class what you have discovered. 3. Design an activity on measuring the expansion of water when frozen. Try your design by actual experimentation at home. 4. Make a concept map on the basic and integrated science processes. Young Ji International School / College
Page 17
5. There are so many problems around us. We have problems in the environment like problems on pollution and misuse of natural resources. Identify a problem and propose possible solutions using the scientific method. The basic science processes are: 1. Observing 2. Measuring 3. Inferring 4. Predicting 5. Classifying 6. Collecting data 7. Recording data The Integrated Science processes are: 1. Interpreting data 2. Controlling data 3. Defining operationally 4. Formulating hypothesis 5. Establishing Space-time relationship 6. Experimenting Observation is the core of experimentation. It involves the use of the senses namely: sense of touch, smell, taste, hearing, and sight. A more accurate observation results from maximizing the use of the senses. We usually make two kinds of observation- qualitative and quantitative. Quantitative observation consists of numbers taken by various measurements of the system. Quantitative observation consists of general description of the attributes or information about the system. Hypothesis is a proposal explanation for set of observations. It is a tentative solution to the problem at hand after enough information has been gathered; hence, this entire wild guess has to be tested by conducting experiments. Theory is a unifying principle that explains a body of facts. The principle has been verified by repeated experimentation. When a theory surpassed the test of time, it is elevated into a scientific law. Experimental law or scientific law is a concise verbal or mathematical statement of relationship between phenomena that are always under the same conditions. Discussion Box 1. What is the difference between accuracy and precision? 2. Which of the three types of error can be avoided? How? 3. Is there a need for an international agreement for units of measurement to be used by each nation? If so, why?
Young Ji International School / College
Page 18
4. Compare and contrast scientific notation and decimal notations. Why was there a preference for scientific notation/ 5. Classify the following as hypothesis, theory or law. a. A stitch in time saves nine. b. Diamond has pores so small that fire cannot enter, hence it does not burn. c. Masses of substances before and after reaction remain the same. d. For every action there is an equal and opposite reaction. e. Materials rich in phlogiston burns. Areas of Investigation 1. Devise a shortcut method for converting temperature from Celsius scale to Fahrenheit scale mentally. 2. Design an activity on measuring the expansion of water when frozen. Try your design in the by actual experimentation at home. 3. Make a concept map on the basic and integrated science processes. 4. There are so many problems around us. We have problems on pollution and misuse of natural resources. Identify a problem and propose possible solutions using the scientific method.
Chapter 4 
Properties of Elements
Mass versus weight
If one were to stand behind this girl at the bottom of the arc and try to stop her, one would be acting against her inertia, which arises from mass, not weight.
In everyday usage, the mass of an object is often referred to as its weight though these are in fact different concepts and quantities. In scientific contexts, mass refers loosely to the amount of "matter" in an object (though "matter" may be difficult to define), whereas weight refers to the force experienced by an object due to gravity. In other words, an object with a mass of 1.0 kilogram will weigh approximately 9.81 Newton‘s (Newton is the unit of force, while kilogram is the unit of mass) on the surface of the Earth (its mass multiplied by the gravitational field strength). Its weight will be less on Mars (where gravity is weaker), more on Saturn, and negligible in space when far from any significant source of gravity, but it will always have the same mass.
Young Ji International School / College
Page 19
Objects on the surface of the Earth have weight, although sometimes this weight is difficult to measure. An example is a small object floating in a pool of water, or even a dish of water, which does not appear to have weight since it is buoyed by the water; but it is found to have its usual weight when it is added to water in a container which is entirely supported by and weighed on a scale. Thus, the "weightless object" floating in water actually transfers its weight to the bottom of the container (where the pressure increases). Similarly, a balloon has mass but may appear to have no weight or even negative weight, due to buoyancy in air. However the weight of the balloon and the gas inside it has merely been transferred to a large area of the Earth's surface (in fact the entire surface, eventually), making the weight difficult to measure. The weight of a flying airplane is similarly distributed to the ground, but does not disappear. If the airplane is in level flight, the same weightforce is distributed to the surface of the Earth as when the plane was on the runway, but spread over a larger area. A better scientific definition of mass describes it as having inertia, the resistance of an object to being accelerated when acted on by an external force. Gravitational "weight" is the force created when a mass is acted upon by a gravitational field and the object is not allowed to free-fall, but is supported or retarded by a mechanical force, such as the surface of a planet. Such a force constitutes weight. This force can be added to by any other kind of force. For example, in the photograph, the girl's weight, subtracted from the tension in the chain (respectively the support force of the seat), yields the necessary centripetal force to keep her swinging in an arc. If one stands behind her at the bottom of her arc and abruptly stops her, the impetus ("bump" or stopping-force) one experiences is due to acting against her inertia, and would be the same even if gravity were suddenly switched off. While the weight of an object varies in proportion to the strength of the gravitational field, its mass is constant (ignoring relativistic) as long as no energy or matter is added to the object. Accordingly, for an astronaut on a spacewalk in orbit (a free-fall), no effort is required to hold a communications satellite in front of him; it is "weightless". However, since objects in orbit retain their mass and inertia, an astronaut must exert ten times as much force to accelerate a 10-ton satellite at the same rate as one with a mass of only 1 ton. On Earth, a swing set can demonstrate this relationship between force, mass, and acceleration. If one were to stand behind a large adult sitting stationary on a swing and give him a strong push, the adult would temporarily accelerate to a quite low speed and then swing only a short distance before beginning to swing in the opposite direction. Applying the same impetus to a small child would produce a much greater speed. 
Mixture
Young Ji International School / College
Page 20
In chemistry, a mixture is a material system made up of two or more different substances which are mixed but are not combined chemically. A mixture refers to the physical combination of two or more substances on which the identities are retained and are mixed in the form of solutions, suspensions, and colloids. Mixtures are the one product of a mechanical blending or mixing of chemical substances like elements and compounds, without chemical bonding or other chemical change, so that each ingredient substance retains its own chemical properties and makeup. Despite that there are no chemical changes to its constituents, the physical properties of a mixture, such as its melting point, may differ from those of the components. Some mixtures can be separated into their components biophysical. Azeotropes are one kind of mixture that usually poses considerable difficulties regarding the separation processes required to obtain their constituents (physical or chemical processes or, even a blend of them). 
Chemical substance
Steam and liquid water are two different forms of the same chemical substance, water. In chemistry, a chemical substance is a form of matter that has constant chemical composition and characteristic properties.[1] It cannot be separated into components by physical separation methods, i.e. without breaking chemical bonds. It can be solid, liquid, gas, or plasma. Chemical substances are often called 'pure' to set them apart from mixtures. A common example of a chemical substance is pure water; it has the same properties and the same ratio of hydrogen to oxygen whether it is isolated from a river or made in a laboratory. Other chemical substances commonly encountered in pure form are diamond (carbon), gold, table salt (sodium chloride) and refined sugar (sucrose). However, in practice, no substance is in practice entirely pure, and chemical purity is specified according to the intended use of the chemical. Chemical substances exist as solids, liquids, gases or plasma, and may change between these phases of matter with changes in temperature or pressure. Chemical reactions convert one chemical substance into another. Young Ji International School / College
Page 21


An inorganic compound is a compound that is not considered "organic". Inorganic compounds are traditionally viewed as being synthesized by the agency of geological systems. In contrast, organic compounds are found in biological systems. Organic chemists traditionally refer to any molecule containing carbon as an organic compound and by default this means that inorganic chemistry deals with molecules lacking carbon.[1] An organic compound is any member of a large class of gaseous, liquid, or solid chemical compounds whose molecules contain carbon. For historical reasons discussed below, a few types of carbon-containing compounds such as carbides, carbonates, simple oxides of carbon (such as CO and CO2), and cyanides are considered inorganic. The distinction between organic and inorganic carbon compounds, while "useful in organizing the vast subject of chemistry... is somewhat arbitrary."
Dietary elements (commonly known as dietary minerals or mineral nutrients) are the chemical elements required by living organisms, other than the four elements carbon, hydrogen, nitrogen, and oxygen present in common organic molecules. The term "dietary mineral" is archaic, as the substances it refers to are chemical elements rather than actual minerals. Chemical elements in order of abundance in the human body include the seven major dietary elements calcium,phosphorus, potassium, sulfur, sodium, chlorine, and magnesium. Important "trace" or minor dietary elements, necessary for mammalian life, include iron, cobalt, copper, zinc, manganese, molybdenum, iodine, bromine, and selenium (see below for detailed discussion). Over twenty dietary elements are necessary for mammals, and several more for various other types of life. The total number of chemical elements that are absolutely needed is not known for any organism. Ultra trace amounts of some elements (e.g., boron, chromium) are known to clearly have a role but the exact biochemical nature is unknown, and others (e.g. arsenic, silicon) are suspected to have a role in health, but without proof. Most chemical element that enter into the dietary physiology of organisms are in the form of simple compounds. Larger chemical compound of elements need to be broken down for absorption. Plants absorb dissolved elements in soils, which are subsequently picked up by the herbivores that eat them and so on, the elements move up the food chain. Larger organisms may also consume soil (geophagia) and visit salt licks to obtain limiting dietary elements they are unable to acquire through other components of their diet. Bacteria play an essential role in the weathering of primary elements that result in the release of nutrients for their own nutrition and for the nutrition of others in the ecological food chain. One element, cobalt, is available for use by animals only Young Ji International School / College
Page 22
after having been processed into complicated molecules (e.g., vitamin B12), by bacteria. Scientists are only recently starting to appreciate the magnitude and role that microorganisms have in the global cycling and formation of biominerals. Methods of separating the Components of a Mixture 
Mechanical separation techniques
Types of mechanical separator Materials separated Separators Liquid from liquid
Settling tanks, liquid cyclones, centrifugal decanters, coalescers
Gas from liquid
Still tanks, deaerators, foam breakers
Liquid from gas
Settling chambers, cyclones, electrostatic precipitators, impingement separators
Solid from liquid
 1. 2. 3. 4. 5.
Filters, centrifugal filters, clarifiers, thickeners,
Flotation (historically spelled floatation) involves phenomena related to the relative buoyancy of objects. The term may refer to: Flotation, any material added to the hull of a watercraft to keep the hull afloat Flotation process, in process engineering, a method for the separation of mixtures Froth flotation, a process for separating hydrophobic from hydrophilic materials Dissolved air flotation (DAF), a water treatment process Induced gas flotation, a water treatment process that clarifies wastewaters (or other waters) by the removal of suspended matter such as oil or solids

Decantation is a process for the separation of mixtures, by removing a layer of liquid, generally one from which a precipitate has settled. 1. The purpose may be either to produce a clean decant, or to remove undesired liquid from the precipitate (or other layers). 2. If the aim is to produce a clean solution, a small amount of solution must generally be left in the container, and care must be taken to prevent any precipitate from flowing with the solution out of the container.
Young Ji International School / College
Page 23
Filtration is commonly the mechanical or physical operation which is used for the separation of solids from fluids (liquids or gases) by interposing a medium through which only the fluid can pass. The fluid that passes through is called the filtrate. Oversize solids in the fluid are retained, but the separation is not complete; solids will be contaminated with some fluid and filtrate will contain fine particles (depending on the pore size and filter thickness). Filtration is also used to describe some biological processes, especially in water treatment and sewage treatment in which undesirable constituents are removed by absorption into a biological film grown on or in the filter medium as in slow sand filtration. Centrifugation is a process that involves the use of the centrifugal force for the sedimentation of heterogeneous mixtures with a centrifuge, used in industry and in laboratory settings. This process is used to separate two immiscible liquids. More-dense components of the mixture migrate away from the axis of the centrifuge, while less-dense components of the mixture migrate towards the axis. Chemists and biologists may increase the effective gravitational force on a test tube so as to more rapidly and completely cause the precipitate ("pellet") to gather on the bottom of the tube. The remaining solution is properly called the "supernate" or "supernatant liquid". The supernatant liquid is then either quickly decanted from the tube without disturbing the precipitate, or withdrawn with a Pasteur pipette. Evaporation is a type of vaporization of a liquid that occurs from the surface of a liquid into a gaseous phase that is not saturated with the evaporating substance. The other type of vaporization is boiling, which is characterized by bubbles of saturated vapor forming in the liquid phase. Steam produced in a boiler is another example of evaporation occurring in a saturated vapor phase. Evaporation that occurs directly from the solid phase below the melting point, as commonly observed with ice at or below freezing or moth crystals (napthalene or paradichlorobenzine), is called sublimation. Sublimation is the transition of a substance directly from the solid to the gas phase without passing through an intermediate liquid phase. Sublimation is an endothermic phase transition that occurs at temperatures and pressures below a substance's triple point in its phase diagram. The reverse process of sublimation is desublimation, or deposition. Distillation is a process of separating the component substances from a liquid mixture by selective vaporization and condensation. Distillation may result in essentially complete separation (nearly pure components), or it may be a partial separation that increases the concentration of selected components of the mixture. In either case the process exploits differences in the volatility of mixture's components. In industrial chemistry, distillation is
Young Ji International School / College
Page 24
a unit operation of practically universal importance, but it is a physical separation process and not a chemical reaction. Commercially, distillation has many applications. For example:
In the fossil fuel industry distillation is a major class of operation in obtaining materials from crude oil for fuels and for chemical feedstock. Distillation permits separation of air into its components — notably oxygen, nitrogen, and argon — for industrial use. In the field of industrial chemistry, large ranges of crude liquid products of chemical synthesis are distilled to separate them, either from other products, or from impurities, or from unreached starting materials. Distillation of fermented products produces distilled beverages with high alcohol content, or separates out other fermentation products of commercial value.
An installation for distillation, especially of alcohol, is a distillery. The distillation equipment is a still.
Laboratory display of distillation: 1:A source of heat 2: Still pot 3: Still head 4: Thermometer/Boiling point temperature 5: Condenser 6: Cooling water in 7: Cooling water out 8:Distillate/receiving flask 9: Vacuum/gas inlet 10: Still receiver 11: Heat control12: Stirrer speed control 13: Stirrer/heat plate 14: Heating Young Ji International School / College
Page 25
(Oil/sand) bath 15:Stirring means e.g. (shown), boiling chips or mechanical stirrer 16: Cooling bath.
Fractional distillation is the separation of a mixture into its component parts, or fractions, such as in separating chemical compounds by their boiling point by heating them to a temperature at which one or more fractions of the compound will vaporize. It is a special type of distillation. Generally the component parts boil at less than 25 °C from each other under a pressure of one atmosphere. If the difference in boiling points is greater than 25 °C, a simple distillation is used.
Polymerization and Plastic
Polymerization is a process of reacting monomer molecules together in a chemical reaction to form polymer chains or three-dimensional networks. There are many forms of polymerization and different systems exist to categorize them.
An example of alkene polymerization, in which each styrene monomer's double bond reforms as a single bond plus a bond to another styrene monomer. The product is polystyrene.
A plastic material is any of a wide range of synthetic or semi-synthetic organic solids that are moldable. Plastics are typically organic polymers of high molecular mass, but they often contain other substances. They are usually synthetic, most commonly derived from petrochemicals, but many are partially natural.
Household items made of various types of plastic Young Ji International School / College
Page 26
Vulcanization of Rubber Vulcanization (or vulcanization) is a chemical process for converting natural rubber or related polymers into more durable materials via the addition of sulfur or other equivalent curatives or accelerators. These additives modify the polymer by forming cross-links (bridges) between individual polymer chains. Vulcanized materials are less sticky and have superior mechanical properties. It was Charles Goodyear from new haven, Connecticut who discovered the vulcanization process by serendipity. He accidentally dropped a mixture of natural rubber and sulfur on a hot stove. Fermentation Fermentation is a metabolic process that converts sugar to acids, gases and/or alcohol. It occurs in yeast and bacteria, but also in oxygen-starved muscle cells, as in the case of lactic acid fermentation. Fermentation is also used more broadly to refer to the bulk growth of microorganisms on a growth medium. French microbiologist Louis Pasteur is often remembered for his insights into fermentation and its microbial causes. The science of fermentation is known as zymology.
Fermentation in progress: Impurities formed by CO2 gas bubbles and fermenting material. Technology Update Nanotechnology is the manipulation of matter on an atomic, molecular, and supramolecular scale. The earliest, widespread description of nanotechnology referred to the particular technological goal of precisely manipulating atoms and molecules for fabrication of macro scale products, also now referred to as molecular nanotechnology. A more generalized description of nanotechnology was subsequently established by the National Nanotechnology Initiative, which defines nanotechnology as the manipulation of matter with at least one dimension sized from 1 to 100 nanometers. This definition reflects the fact that quantum mechanical effects are important at this quantum-realm scale, and so the definition shifted from a particular technological goal to a research category inclusive of all types of research and technologies that deal with the special properties of matter that occur below the given size threshold. It is therefore common to see the plural Young Ji International School / College
Page 27
form "nanotechnologies" as well as "nanoscale technologies" to refer to the broad range of research and applications whose common trait is size. Because of the variety of potential applications (including industrial and military), governments have invested billions of dollars in nanotechnology research. Areas of Investigation 1. People in sugar cane farms crudely prepare sugar for their own consumption. If you were among them, how would you prepare sugar in small scale using a crude process at home? 2. Design an activity where you can use column chromatography or paper chromatography to separate the pigments of: a. Turmeric or luyang dilaw b. Duhat c. Sibukaw 3. How would you extract oil from ilang-ilang flowers or rose petals/ 4. Interview a medical technologist from the blood bank and inquire how blood platelets are separated from the whole blood. 5. Design a method of separating iodine crystals from granulated glass. 6. Investigate how sugar is milled from sugar cane and the various treatments that finally make it granulated and refined. Find also the other products you can obtain from it. 7. Make a home investigation to separate the components of the following: a. b. c. d.
Salt from sea water Different color components from milk Potassium alum from murky water Iron filings mixed with sulfur
Discussion Box 1. What are some naturally-occurring mixtures in the environment? 2. How would you separate the components of a discharged dry cell? 3. What is desalination? What global problem will arise if sea water will be desalinated in large scale? 4. Enumerate the steps in water treatment and purification.
Young Ji International School / College
Page 28
Lesson 5
Atoms, Elements, and Compounds
Atoms of Antiquity Free will in antiquity was not discussed in the same terms as used in the modern free will debates, but historians of the problem have speculated who exactly was first to take positions as determinist, libertarian, and compatibility in antiquity. There is wide agreement that these views were essentially fully formed over 2000 years ago. Candidates for the first thinkers to form these views, as well as the idea of a non-physical "agent-causal" libertarianism, include a. Democritus (460-370) The materialist philosophers Democritus and his mentor Leucippus were the first determinists. They claimed that all things, including humans, were made of atoms in a void, with individual atomic motions strictly controlled by causal laws. Democritus said: "By convention (nomos) color, by convention sweet, by convention bitter, but in reality atoms and a void." Democritus wanted to wrest control of man's fate from arbitrary gods and make us more responsible for our actions. But ironically, he and Leucippus originated two of the great dogmas of determinism, physical determinism and logical necessity, which lead directly to the traditional and modern problem of free will and determinism. Leucippus dogmatically declared an absolute necessity which left no room in the cosmos for chance. The consequence is a world with but one possible future, completely determined by its past. The Pythagoreans, Socrates, and Plato attempted to reconcile an element of human freedom with material determinism and causal law, in order to hold man responsible for his actions. b. Aristotle (384-322)
Young Ji International School / College
Page 29
The first major philosopher to argue convincingly for some indeterminism was probably Aristotle. He elaborated the four possible causes (material, efficient, formal, and final). Aristotle's word for these causes was ἀιτία, which translates as causes in the sense of the multiple factors responsible for an event. Aristotle did not subscribe to the simplistic "every event has a (single) cause" idea that was to come later. Then, in his Physics and Metaphysics, Aristotle also said there were "accidents" caused by "chance (τυχή)." In his Physics, he noted that the early physicists had found no place for chance among their causes. c. Epicurus (341-270) It is with Epicurus and the Stoics that clearly in deterministic and deterministic positions are first formulated. Writing one generation after Aristotle, Epicurus argued that as atoms moved through the void, there were occasions when they would "swerve" (clinamen) from their otherwise determined paths, thus initiating new causal chains. Epicurus argued that these swerves would allow us to be more responsible for our actions (libertarianism), something impossible if every action was deterministically caused.
d. Chrysippus (280-207) e. Carneades (214-129)
Dalton’s Atomic Theory John Dalton was an English chemist, meteorologist and physicist. He is best known for his pioneering work in the development of modern atomic theory, and his research into color blindness (sometimes referred to as Daltonism, in his honor).
Young Ji International School / College
Page 30
In the early 1800‘s, John Dalton determined that each chemical element is composed of unique type of atom, and that the atoms different by their masses. The following were his assumptions about the atom: 1. Matter consists of indivisible and indestructible particles called atoms. 2. All the atoms of a particular element are identical in mass, size, shape, and other properties. 3. Atoms cannot be changed into atoms of other elements or be destroyed in chemical reactions. 4. Atoms combine in a definite ratio to from a compound. 5. The relative numbers and kinds of atoms in a given compound do not change. Fundamental Laws of Chemical Combinations 
The Law of Conservation of Mass The law of conservation of mass, or principle of mass conservation, states that for any system closed to all transfers of matter and energy (both of which have mass), the mass of the system must remain constant over time, as system mass cannot change quantity if it is not added or removed. Hence, the quantity of mass is "conserved" over time. The law implies that mass can neither be created nor destroyed, although it may be rearranged in space, or the entities associated with it may be changed in form, as for example when light or physical work is transformed into particles that contribute the same mass to the system as the light or work had contributed. The law implies (requires) that during any chemical reaction, nuclear reaction, or radioactive decay in an isolated system, the total mass of the reactants or starting materials must be equal to the mass of the products. The concept of mass conservation is widely used in many fields such as chemistry, mechanics, and fluid dynamics. Historically, mass conservation was discovered in chemical reactions by Antoine Lavoisier in the late 18th century, and was of crucial importance in the progress from alchemy to the modern natural science of chemistry.
Young Ji International School / College
Page 31
Antoine Lavoisier's discovery of the Law of Conservation of Mass led to many new findings in the 19th century. Joseph's law of definite proportions and John Dalton's atomic theory branched from the discoveries of Antoine Lavoisier. Lavoisier's quantitative experiments revealed that combustion involved oxygen rather than what was previously thought to be phlogiston. Discovering Electrons and Electrical Charges Electric charge is the physical property of matter that causes it to experience a force when placed in an electromagnetic field. There are two types of electric charges: positive and negative. Positively charged substances are repelled from other positively charged substances, but attracted to negatively charged substances; negatively charged substances are repelled from negative and attracted to positive. An object is negatively charged if it has an excess of electrons, and is otherwise positively charged or uncharged. The SI derived unit of electric charge is the coulomb (C), although in electrical engineering it is also common to use the ampere-hour (Ah), and in chemistry it is common to use the elementary charge (e) as a unit. The symbol Q is often used to denote charge. The early knowledge of how charged substances interact is now called classical electrodynamics, and is still very accurate if quantum effects do not need to be considered.
Electric field of a positive and a negative point charge.
The Cathode Ray Tube (CRT) The cathode ray tube or (CRT) is a vacuum tube containing one or more electron guns, and a fluorescent screen used to view images.[1] It has a means to accelerate and deflect the electron beam(s) onto the screen to create the images. The images may represent electrical waveforms(oscilloscope), pictures (television, computer monitor), radar targets or others. CRTs have also been used as memory devices, in which case the visible light emitted from the fluorescent material (if any) is not intended to have significant meaning to a visual observer (though the visible pattern on the tube face may cryptically represent the stored data). The CRT uses an evacuated glass envelope which is large, deep (i.e. long from front screen face to rear end), fairly heavy, and relatively fragile. As a matter of Young Ji International School / College
Page 32
safety, the face is typically made of thick lead glass so as to be highly shatterresistant and to block most X-ray emissions, particularly if the CRT is used in a consumer product.
Cutaway rendering of a color CRT: 1. Three Electron guns (for red, green, and blue phosphor dots) 2. Electron beams 3. Focusing coils 4. Deflection coils 5. Anode connection 6. Mask for separating beams for red, green, and blue part of displayed image 7. Phosphor layer with red, green, and blue zones 8. Close-up of the phosphor-coated inner side of the screen
X-rays Detected
Wilhelm Conrad Röntgen was a German physicist, who, on 8 November 1895, produced and detected electromagnetic radiation in a wavelength range today that was known as X-rays or Röntgen rays, an achievement that earned him the first Nobel Prize in Physics in 1901. In honor of his accomplishments, in 2004 the International Union of Pure and Applied Chemistry (IUPAC) named element 111, roentgenium, a radioactive element with multiple unstable isotopes, after him.
Wilhelm Roentgen (1711-1793) Young Ji International School / College
Page 33
Science Ideas
Radioisotopes are used tracers or diagnostic agents for tumors and cancer. These are also used to stop abnormal cell growth or to destroy cancer cells. Half-life is the time required for the reactant to decrease to half of its initial concentration. Radioactivity is the spontaneous emission of particles or rays from a radioactive substance. Atomic mass is the average mass of the difference isotopic forms of an element. Isotopes are atoms of the same element with the same atomic number but have different mass numbers. The total atomic masses of different isotopes of an element give its average atomic mass, which are not whole numbers. Mass number is the total number of protons and neutrons in the nucleus of an atom.
Areas of Investigation 1. Make a report on the life history of John Dalton. 2. Search for the elements having the greatest number of isotopes. 3. Look for materials that show transfer of charges as you rub them. What is electrostatics? 4. Give some benefits of radioactivity to humans. 5. Suppose you operate a hospital and would like to quit the business, how would you properly dispose the radioactive materials that you used as tracers for therapy? Discussion Box 1. 2. 3. 4. 5. 6.
State the assumptions of the Modern Atomic Theory Give the three laws, which lend applications of Dalton‘s atomic theory. What information did the cathode ray tube (CRT) experience reveal? Define half-life. Differentiate a nuclear reaction from chemical reaction. Differentiate cathode rays, ultraviolet rays, and x-rays from one another.
The Atomic structure
Young Ji International School / College
Page 34
The red dots are protons, the black dots are neutrons, and the blue dots are electrons. Any atom is composed of a little nucleus surrounded by a "cloud" of electrons. In the nucleus there are protons and neutrons. However, the term "atom" just refers to a building block of matter; it doesn't specify the identity of the atom. It could be an atom of carbon, or an atom of hydrogen, or any other kind of atom. This is where the term "element" comes into play. When an atom is defined by the number of protons contained in its nucleus, chemists refer to it as an element. All elements have a very specific identity that makes them unique from other elements. For example, an atom with 6 protons in its nucleus is known as the element carbon. When speaking of the element fluorine, chemists mean an atom that contains 9 protons in its nucleus. 
Atom: A fundamental building block of matter composed of protons, neutrons, and electrons.

Element: A uniquely identifiable atom recognized by the number of protons in the nucleus.
Despite the fact that we define an element as a unique identifiable atom, when we speak, for example, 5 elements, we don't usually mean those 5 atoms are of the same type (having the same number of protons in their nucleus). We mean 5 'types' of atoms. It is not necessary there are only 5 atoms. There may be 10, or 100, etc. atoms, but those atoms belong to one of 5 types of atoms. I'd rather define 'element' as 'type of atom'. I think it is more precise. If we'd like to refer to 5 atoms having the same 6 protons in their nucleus, I'd say '5 carbon atoms' or '5 atoms of carbon'. It is important to note that if the number of protons in the nucleus of an atom changes, so does the identity of that element. If we could remove a proton from nitrogen (7 protons), it is no longer nitrogen. We would, in fact, have to identify the atom as carbon (6 protons). Remember, elements are unique and are always defined by the number of protons in the nucleus. The Periodic Table of the Elements shows all known elements organized by the number of protons they have. An element is composed of the same type of atom; elemental carbon contains any number of atoms, all having 6 protons in their nuclei. In contrast, compounds are composed of different type of atoms. More precisely, a compound is a chemical substance that consists of two or more elements. A carbon compound contains some carbon atoms (with 6 protons each) and some other atoms with different numbers of protons.
Young Ji International School / College
Page 35
Compounds have properties different from the elements that created them. Water, for example, is composed of hydrogen and oxygen. Hydrogen is an explosive gas and oxygen is a gas that fuels fire. Water has completely different properties, being a liquid that is used to extinguish fires. The smallest representative for a compound (which means it retains characteristics of the compound) is called a molecule. Molecules are composed of atoms that have "bonded" together. As an example, the formula of a water molecule is "H2O": two hydrogen atoms and one oxygen atom.
Properties of Matter Properties of matter can be divided in two ways: extensive/intensive, or physical/chemical.
without changing the chemical's identity. The freezing point of a substance is physical. When water freezes, it's still H2O.
Extensive properties depend on the amount of matter that is being measured. These include mass and volume. Intensive properties do not depend on the amount of matter. These include density and color.
Physical properties can be measured
Chemical properties deal with how one chemical reacts with another. We know that wood is flammable because it becomes heat, ash, and carbon dioxide when heated in the presence of oxygen.
STATE OF MATTER One important physical property is the state of matter. Three are common in everyday life: solid, liquid, and gas. The fourth, plasma, is observed in special conditions such as the ones found in the sun and fluorescent lamps. Substances can exist in any of the states. Water is a compound that can be liquid, solid (ice), or gas (steam).
Young Ji International School / College
Page 36
The ice in this picture is a solid. The water in the picture is a liquid. In the air there is water vapor, which is a gas.
The states of matter depend on the bonding between molecules. 
Solids
Solids have a definite shape and a definite volume. Most everyday objects are solids: rocks, chairs, ice, and anything with a specific shape and size. The molecules in a solid are close together and connected by intermolecular bonds. Solids can be amorphous, meaning that they have no particular structure, or they can be arranged into crystalline structures or networks. For instance, soot, graphite, and diamond are all made of elemental carbon, and they are all solids. What makes them so different? Soot is amorphous, so the atoms are randomly stuck together. Graphite forms parallel layers that can slip past each other. Diamond, however, forms a crystal structure that makes it very strong. 
Liquids
Liquids have a definite volume, but they do not have a definite shape. Instead, they take the shape of their container to the extent they are indeed "contained" by something such as beaker or a cupped hand or even a puddle. If not "contained" by a formal or informal vessel, the shape is determined by other internal (e.g. intermolecular) and external (e.g. gravity, wind, inertial) forces. The molecules are close, but not as close as a solid. The intermolecular bonds are weak, so the molecules are free to slip past each other, flowing smoothly. A property of liquids is viscosity, the measure of "thickness" when flowing. Water is not nearly as viscous as molasses, for example. 
Gases
Gases have no definite volume and no definite shape. They expand to fill the size and shape of their container. The oxygen that we breathe and steam from a pot are both examples of gases. The molecules are very far apart in a gas, and there are Young Ji International School / College
Page 37
minimal intermolecular forces. Each atom is free to move in any direction. Gases undergo effusion and diffusion. Effusion occurs when a gas seeps through a small hole, and diffusion occurs when a gas spreads out across a room. If someone leaves a bottle of ammonia on a desk, and there is a hole in it, eventually the entire room will reek of ammonia gas. That is due to the diffusion and effusion. These properties of gas occur because the molecules are not bonded to each other. 
In gases, intermolecular forces are very weak, hence molecules move randomly colliding with themselves, and with the wall of their container, thus exerting pressure on their container. When heat is given out by gases, the internal molecular energy decreases; eventually, the point is reached when the gas liquefies.

Valence electron
Four covalent bonds. Carbon has four valence electrons and here a valence of four. Each hydrogen atom has one valence electron and is univalent.
In chemistry, a valence electron is an electron that is associated with an atom, and that can participate in the formation of a chemical bond; in a single covalent bond, both atoms in the bond contribute one valence electron in order to form a shared pair. The presence of valence electrons can determine the element's chemical properties and whether it may bond with other elements: For a main group element, a valence electron can only be in the outermost electron shell. In a transition metal, a valence electron can also be in an inner shell. An atom with a closed shell of valence electrons (corresponding to an electron configuration s2p6) tends to be chemically inert. An atom with one or two valence electrons more than a closed shell is highly reactive, because the extra valence electrons are easily removed to form a positive ion. An atom with one or two valence electrons fewer than a closed shell is also highly reactive, because of a tendency either to gain the missing valence electrons (thereby forming a negative ion), or to share valence electrons (thereby forming a covalent bond). Like an electron in an inner shell, a valence electron has the ability to absorb or release energy in the form of a photon. An energy gain can trigger an electron to move (jump) to an outer shell; this is known as atomic excitation. Or the electron can even break free from its associated atom's valence shell; this is ionization to form a Young Ji International School / College
Page 38
positive ion. When an electron loses energy (thereby causing a photon to be emitted), then it can move to an inner shell which is not fully occupied.
The number of valence electrons The number of valence electrons of an element can be determined by the periodic table group (vertical column) in which the element is categorized. With the exception of groups 3–12 (the transition metals), the units digit of the group number identifies how many valence electrons are associated with a neutral atom of an element listed under that particular column.
The periodic table of the chemical elements Electron configuration For a main group element, the number of valence electrons that it may have depends on the electron configuration in a simple way. However, for a transition metal, the relationship is more complex. For a main group element, a valence electron is defined as an electron that resides in the electronic shell of highest principal quantum number n.[1] For example the electronic configuration of phosphorus (P) is 1s2 2s2 2p6 3s2 3p3 so that there are 5 valence electrons (3s2 3p3), corresponding to a maximum valence for P of 5 as in the molecule PF5; this configuration is normally abbreviated to [Ne] 3s2 3p3, where [Ne] signifies the core electrons whose configuration is identical to that of the noble gas neon. However, this simple method does not work for transition metals, which have incomplete d (i.e., 3d, 4d or 5d) subshells, whose energy is normally comparable with that of an
s electron. Instead, a valence electron for a transition metal is
Young Ji International School / College
Page 39
defined as an electron that resides outside a noble-gas core. For example, manganese (Mn) has configuration 1s2 2s2 2p63s2 3p6 4s2 3d5; this is abbreviated to [Ar] 4s2 3d5, where [Ar] denotes a core configuration identical to that of the noble gas argon. In this atom, a 3d electron has energy similar to that of a 4s electron, and much higher than that of a 3s or 3p electron. In effect, there are possibly seven valence electrons (4s2 3d5) outside the argon-like core; this is consistent with the chemical fact that manganese can have an oxidation state as high as +7 (in the permanganate ion: MnO− 4). The farther right in each transition metal series, the lower the energy of an electron in a d subshell and the less such an electron has the properties of a valence electron. Thus, although a nickel atom has, in principle, ten valence electrons (4s23d8), its oxidation state never exceeds four. Because the number of valence electrons which actually participate in chemical reactions is difficult to predict, the concept of the valence electron is less useful for a transition metal than for a main group element; as mentioned already, the d electron count provides a more useful tool for understanding the chemistry of a transition metal.

Chemical reactions
The number of electrons in an atom's outermost valence shell governs its bonding behavior. Therefore, elements whose atoms can have the same number of valence electrons are grouped together in the periodic table of the elements. As a general rule, a main group element (except hydrogen or helium) tends to react to form a closed shell, corresponding to the electron configuration s2p6. This tendency is called the octet rule, because each bonded atom has eight valence electrons including shared electrons. The most reactive kind of metallic element is an alkali metal of group 1 (e.g., sodium or potassium); this is because such an atom has only a single valence electron; during the formation of an ionic bond which provides the necessary ionization energy, this one valence electron is easily lost to form a positive ion (cation) with a closed shell (e.g., Na+ or K+). An alkaline earth metal of Group 2 (e.g., magnesium) is somewhat less reactive, because each atom must lose two valence electrons to form a positive ion with a closed shell (e.g., Mg2+). Within each group (each periodic table column) of metals, reactivity increases with each lower row of the table (from a light element to a heavier element), because a heavier element has more electron shells than a lighter element; a heavier element's valence electrons exist at higher principal quantum numbers (they are Young Ji International School / College
Page 40
farther away from the nucleus of the atom, and are thus at higher potential energies, which means they are less tightly bound). A nonmetal atom tends to attract additional valence electrons to attain a full valence shell; this can be achieved in one of two ways: An atom can either share electrons with a neighboring atom (a covalent bond), or it can remove electrons from another atom (an ionic bond). The most reactive kind of nonmetal element is a halogen (e.g., fluorine (F) or chlorine (Cl)). Such an atom has the following electron configuration: s2p5; this requires only one additional valence electron to form a closed shell. To form an ionic bond, a halogen atom can remove an electron from another atom in order to form an anion (e.g., F−, Cl−, etc.). To form a covalent bond, one electron from the halogen and one electron from another atom form a shared pair (e.g., in the molecule H–F, the line represents a shared pair of valence electrons, one from H and one from F). Within each group of nonmetals, reactivity decreases with each lower rows of the table (from a light element to a heavy element) in the periodic table, because the valence electrons are at progressively higher energies and thus progressively less tightly bound. In fact, oxygen (the lightest element in group 16) is the most reactive nonmetal after fluorine, even though it is not a halogen, because the valence shell of a halogen is at a higher principal quantum number. In these simple cases where the octet rule is obeyed, the valence of an atom equals the number of electrons gained, lost, or shared in order to form the stable octet. However there are also many molecules which are exceptions, and for which the valence is less clearly defined.
Electrical conductivity Valence electrons are also responsible for the electrical conductivity of an element; as a result, an element may be classified as a metal, a nonmetals, or semiconductors (or metalloid). A metal is an element with high electrical conductivity or malleability when in the solid state. In each row of the periodic table, the metals occur to the left of the nonmetals, and thus a metal has fewer possible valence electrons than a nonmetal. However, a valence electron of a metal atom has a small ionization energy, and in the solid state this valence electron is relatively free to leave one atom in order to associate with another nearby. Such a "free" electron can be moved under the influence of an electric field, and its motion constitutes an electric current; it is responsible for the electrical conductivity of the metal. Copper, aluminum, silver, and gold are examples of good conductors.
Young Ji International School / College
Page 41
A nonmetallic element has low electrical conductivity; it acts as an insulator. Such an element is found toward the right of the periodic table, and it has a valence shell that is at least half full (the exception is boron). Its ionization energy is large; an electron cannot leave an atom easily when an electric field is applied, and thus such an element can conduct only very small electric currents. Examples of solid elemental insulators are diamond (an allotrope of carbon) and sulfur. A solid compound containing metals can also be an insulator if the valence electrons of the metal atoms are used to form ionic bonds. For example, although elemental sodium is a metal, solid sodium chloride is an insulator, because the valence electron of sodium is transferred to chlorine to form an ionic bond, and thus that electron cannot be moved easily. A semiconductor has an electrical conductivity that is intermediate between that of a metal and that of a nonmetal; a semiconductor also differs from a metal in that a semiconductor's conductivity increases with temperature. The typical elemental semiconductors are silicon and germanium, each atom of which has four valence electrons. The properties of semiconductors are best explained using band theory, as a consequence of a small energy gap between a valence band(which contains the valence electrons at absolute zero) and a conduction band (to which valence electrons are excited by thermal energy).
Quiz 1. Give the symbol of the element that fits the following descriptions: a. The alkaline earth metal in the sixth period b. A noble gas in the third period c. An alkali metal with the highest atomic mass d. The semi-metal in Group 3A e. A halogen which is only liquid at normal condition f. The sixth-period representative element with properties similar to Be g. The radioactive noble gas h. The Group 5A semimetal in the fifth period 2. Predict which element in each pair has larger/higher atomic radius ionization energy, ionic radius. a. Li or Na b. As or Se c. Pb or Bi d. Mg or Al e. Cs or Ba 3. State the number of valence electrons for each representative elements a. N b. O Young Ji International School / College
Page 42
c. d. e. f. g. h. i. j. k. l.
P B F Ar Cs Ga Sb Ca Ne Mg
I. Multiple Choices: Encircle the correct answer. 1. Chemistry came from the Greek word a. chemeia
b. chemistra
c. kemistra
d. kimia
2. Chemistry is the science that deals with a. matter
b. gas
c. elements
d. structures
c. composition
d. all of these
3. Chemistry it deals with matter. a. structure
b. change
4. He is a German physician and chemist who coined the term phlogiston. a. Robert Boyle Priestley‘s
b. George Ernest Stahl
c. Johan Becher
d. Joseph
5. An English natural philosopher was among the first to use the scientific method in testing his theories. a. Robert Boyle Priestley‘s
b. George Ernest Stahl
c. Johan Becher
d. Joseph
6. He was the first to do quantitative observations to explain burning. a. Antoine Laurent Lavoisier b. Joseph Preistley‘s
c. George Ernest Stahl
7. It is the application of chemistry concepts and principles to produce basic commodities for man to make his life easy and comfortable. a. chemical advances
b. chemical technology
c. chemical analysis
8. An Italian who was awarded a Nobel Prize in physics built the first atomic pile and produced the first controlled chain reaction. a. Wallace Carothers Goodyear
b. Enrico Fermi
Young Ji International School / College
c. Karl Ziegler
d. Charles
Page 43
9. He is an American Dupont chemist, led a research team that invented nylon. a. Wallace Carothers Goodyear
b. Enrico Fermi
c. Karl Ziegler
d. Charles
10. He is a German chemist, shared a Nobel Prize with Giulio Natta on their work with polymers, high density polyethylene or high impact plastics. a. Wallace Carothers Goodyear
b. Enrico Fermi
c. Karl Ziegler
d. Charles
11. He made a serendipitous discovery of vulcanization of rubber. a. Lise Meitner
b. Enrico Fermi
c. Charles Goodyear d. Alan Heeger
12. It refers to the ―closeness‖ among measured data. a. precision
b. precise
c. accuracy
d. measured
13. It is a measured data is ______ when it is close to the accepted or true value. a. precision
b. precise
c. accurate
d. measured
14. It refers to an individual measurement, which gives the same result each time you make the measurement no matter how many times the measurement is repeated. a. precision
b. precise
c. accuracy
d. measured
15. It is a proposal or suggested explanations for an observation and must be testable. a. observations
b. theory
c. hypothesis
d. all of these
16. What error is committed when using a 100-mL graduated cylinder in measuring a gallon of water? a. Random error
b. Gross error
c. Systematic error d. Absolute error
17. What is a quantitative observation? a. The temperature is 30°C.
c. It is a hotter day
b. Yesterday was cooler than today
d. All of these
18. How did Edison discover the filament of his bulb? a. serendipity
b. trial and error
c. by intuition
d. by wild guess
19. Failure to focus your eyes directly at the level of the lower meniscus while the volume of the liquid will cause due to _______. Young Ji International School / College
Page 44
a. miscalculation
b. misconception
c. parallax
d. refraction
20. The reproducibility of a measurement refers to ________. a. precision
b. accuracy
c. reliability
d. mass
II. Identification. Answer the following questions. Write the correct answer on the space provided. __________ 1. It is a concise mathematical expression of observations; or a statement of the observations made by many investigators to which no clear exceptions are known. __________ 2. It is a kind of skills that developed through interpreting analyzing, and manipulating information. __________ 3. It is a conceptual model based on intensive experimentation. __________ 4. It is an internationally-accepted system of measurement. __________ 5. Are digits that are definitely known plus one estimated value? __________ 6. The factor-unit or factor-label method is a scheme in problem-solving wherein the units go with the numerical values during calculations, manipulated that other units are cancelled out, leaving behind the desired one. __________ 7. It is an orderly plan of thinking and doing things and most especially in solving problems. __________ 8. It is arises when behavior is blocked because a desired end is not at once attainable. __________ 9. It is the core of experimentation. __________ 10. Are separated and purified in many ways like filtration, decantation, evaporation, flotation, sublimation. III. Conversion problems 1. A piece of metal has a measured volume of 10 cm3. What is its volume in m3? 2. If a poultry farm harvests 5 eggs / minute, how many eggs could be collected per hour? Per week? Per month? 3. Change 325 cm3 to liter. 4. The normal atmosphere pressure at sea level is 760 mm Hg. Express this pressure in terms of: a. cm Hg b. atmosphere c. millibar Conversion factors: 1 atm = 760 mm of Hg = 76 cm of hg = 1013 mb = 760 torr 1L = 1000 cm3 1m = 100 cm IV. Match Column A with Column B. Write only the letter of your answer. Column A Column B Young Ji International School / College
Page 45
_____1. Mutagens _____ 2. Coal and ores _____ 3. Nuclear fission _____ 4. Acid rain _____ 5. Carcinogenic extracts _____ 6. Hair dyne _____ 7. LPG _____ 8. Pyroclastic materials _____ 9. Oxyacetylene tanks _____10. Food additives _____11. CBB-Cement-Bonded Boards _____12. Fertilizers and pesticides
a. metallurgist b. geneticists c. cosmetologists d. architects e. mining engineers f. food technologies g. house wives h. volcanologists i. pharmacologists j. welders k. agriculturists l. ecologists m. physicists
V. Classify the following chemical systems as substance or mixture, and as homogeneous. 1. paint 2. perfume 3. iodized salt 4. nail polish 5. distilled water 6. vinegar 7. air 8. table salt 9. acetone 10. fresh milk 11. soft drinks 12. Helium 13. sea water 14. mayonnaise 15. cup of coffee VI. Give examples of the following. 1. base 2. acid 3. nonmetal 4. metal 5. compound 6. element 7. colloid 8. solution 9. chemical change 10. physical change 11. heterogeneous mixture 12. homogeneous mixture 13. extensive property 14. intensive property 15. gold
Young Ji International School / College
Page 46
VII. Answer the following questions. (5 points each) 1. Explain why an astronaut having a mass of 75 kg achieves weightlessness on board the skylab orbital space laboratory. 2. How chemistry help our daily life of living? 3. What are the 5 scientific methods?
Chapter 6
Mystery of the Atom
The atom is the smallest unit that defines the chemical elements and their isotopes. Every material object, or substance that can be touched and felt, is made up of atoms. Everything that is solid, liquid, or gas is made up of atoms. Atoms are tiny; their size is typically measured in picometers (trillionths of a meter). A single strand of human hair is about one million carbon atoms wide. Every atom is composed of a nucleus made of protons and neutrons (hydrogen-1 has no neutrons). The nucleus is in turn surrounded by a cloud of electrons. The electrons in an atom are bound to the atom by the electromagnetic force, and the protons and neutrons in the nucleus are bound to each other by the nuclear force. Over 99% of the atom's mass is in the nucleus. The protons have a positive electric charge, the electrons have a negative electric charge, and the neutrons have no electric charge. Normally, an atom's electrons balance out the positive charge of its protons to make it electrically neutral. If an atom has a surplus or deficit of electrons, then it will have an overall charge, and is called an ion. The number of protons in the nucleus determines what chemical element the atom belongs to (e.g. all copper atoms contain 29 protons). The number of neutrons determines what isotope of the element it is. The electron cloud of the atom determines the atom's chemical properties and strongly influences its magnetic properties. Atoms can attach themselves to each other by chemical bonds to form molecules, network solids, metal alloys, crystals, and other solid solutions. The tendency for atoms to bond and break apart is responsible for most of the physical changes we observe in nature, and this is studied by the science of chemistry. Atoms and sub-atomic particles behave in peculiar ways that cannot be explained through the classical laws of physics. The field of quantum mechanics was developed to explain the structure and behavior of atoms. Atoms are not the only basic structure in the Universe, but they do comprise all the types of matter than can be seen and touched. In addition to atomic matter that is made of atoms, and also certain free subatomic particles, it is believed that the Universe contains an even larger amount of dark matter, which is not made of atoms, and is composed of particles of yet unknown type. The idea that matter is made up of discrete units is a very old one, appearing in many ancient cultures such as Greece and India. The word "atom", in fact, was coined by ancient Greek philosophers. However, these ideas were founded in philosophical and theological reasoning rather than evidence and experimentation. Young Ji International School / College
Page 47
As a result, their views on what atoms look like and how they behave were very incorrect. They also couldn't convince everybody, so atomism was but one of a number of competing theories on the nature of matter. It wasn't until the 19th century that the idea was embraced and refined by scientists, when the blossoming science of chemistry produced discoveries that only the concept of atoms could explain.
First evidence-based theory
Various atoms and molecules as depicted in John Dalton's A New System of Chemical Philosophy (1808).
The structure of atoms
The gold foil experiment Dalton also believed atomic theory could explain why water absorbs different gases in different proportions. For example, he found that water absorbs carbon dioxide far better than it absorbs nitrogen. Dalton hypothesized this was due to the differences in mass and complexity of the gases' respective particles. Indeed, carbon dioxide molecules (CO2) are heavier and larger than nitrogen molecules (N2). In 1827, botanist Robert Brown used a microscope to look at dust grains floating in water and discovered that they moved about erratically, a phenomenon that became known as "Brownian motion". This was thought to be caused by water molecules knocking the grains about. In 1905 Albert Einstein produced the first mathematical analysis of the motion. French physicist Jean Perrin used Einstein's Young Ji International School / College
Page 48
work to experimentally determine the mass and dimensions of atoms, thereby conclusively verifying Dalton's atomic theory. The physicist J. J. Thomson, through his work on cathode rays in 1897, discovered the electron, and concluded that they were a component of every atom. Thus he overturned the belief that atoms are the indivisible, ultimate particles of matter. Thomson postulated that the low mass, negatively charged electrons were distributed throughout the atom in a uniform sea of positive charge. This became known as the plum pudding model. In 1909, Hans Geiger and Ernest Marsden, under the direction of Ernest Rutherford, bombarded a metal foil with alpha particles to observe how they scattered. They expected all the alpha particles to pass straight through with little deflection, because Thomson's model said that the charges in the atom are so diffuse that their electric fields could not affect the alpha particles much. However, Geiger and Marsden spotted alpha particles being deflected by angles greater than 90°, which was supposed to be impossible according to Thomson. To explain this, Rutherford proposed that the positive charge of the atom is concentrated in a tiny nucleus at the center of the atom. While experimenting with the products of radioactive decay, in 1913 radiochemistry Frederick Soddy discovered that there appeared to be more than one type of atom at each position on the periodic table. The term isotope was coined by Margaret Todd as a suitable name for different atoms that belong to the same element. J.J. Thomson created a technique for separating atom types through his work on ionized gases, which subsequently led to the discovery of stable isotopes.
The Bohr model of the atom, with an electron making instantaneous "quantum leaps" from one orbit to another. This model is obsolete. Helium atom
Young Ji International School / College
Page 49
A chemical symbol is a code for a chemical element. It is usually derived from the name of the element, often in Latin.
An illustration of the helium atom, depicting the nucleus(pink) and the electron cloud distribution (black). The nucleus (upper right) in helium-4 is in reality spherically symmetric and closely resembles the electron cloud, although for more complicated nuclei this is not always the case. The black bar is one angstrom (10−10 m or 100 pm).
This is an example of an atomic symbol. The text boxes explain where the numbers are derived from. Only the first letter is capitalized. For example, "He" is the symbol for helium(English name, not known in ancient Roman times), "Pb" for lead (plumbum in Latin), "W" for tungsten (wolfram in German, not known in Roman times).Temporary symbols assigned to newly or not-yet synthesized elements use 3-letter symbols based on their atomic numbers. For example, "Uno" was the temporary symbol for hassium which had the temporary name of unniloctium and "Uuo" is the symbol for ununoctium (temporary name) with the atomic mass 118. Chemical symbols may be modified by the use of pretended superscripts or subscripts to specify a particular isotope of an atom. Additionally, appended superscripts may be used to indicate the ionization or oxidation state of an element. They are widely used in chemistry and they have been officially chosen by the International Union of Pure and Applied Chemistry. There are also some historical symbols that are currently not official any more. Attached subscripts or superscripts specifying a nucleotide or molecule have the following meanings and positions:
The nucleon number (mass number) is shown in the left superscript position (e.g., 14N)
The proton number (atomic number) may be indicated in the left subscript position (e.g., 64Gd) If necessary, a state of ionization or an excited state may be indicated in the right superscript position (e.g., state of ionization Ca2+). In astronomy, non-ionised atomic hydrogen is often known as "H I", and ionised hydrogen as "H II". The number of atoms of an element in a molecule or chemical compound is shown in the right subscript position (e.g., N2or Fe2O3) A radical is indicated by a dot on the right side (e.g., Cl· for a chloride radical)
In Chinese each chemical element has an ideograph, usually created for the purpose, as its symbol (see Chemical elements in East Asian languages).
Young Ji International School / College
Page 50
The Periodic Table, elements being denoted by their symbols A list of current, dated, as well as proposed and historical signs and symbols is included here with its signification. Also given is each element's number, atomic or the atomic mass of the most stable isotope, group and period numbers on the periodic table, and etymology of the symbol. Antimatter atoms are denoted by a bar above the symbol for their matter counterpart, so e.g. H is the symbol for antihydrogen. Pictographic symbols The following is a list of pictographic symbols employed to symbolize elements known since ancient times (for example to the alchemists). Not included in this list are symbolic representations of substances previously called elements (such as certain rare earth mineral blends and the classical elements fire and water of ancient philosophy) which are known today to be multi-atomic. Also not included are symbolic representations currently used for elements in other languages such as the Traditional Chinese elements. Modern alphabetic notation was introduced in 1814 by JĂśns Jakob Berzelius.
Chemical symbol
ďż˝
Original name
Modern name
Atomic number
Origin of symbol
Hydrogen
Hydrogen
1
Daltonian symbol circa 1808.
Sulfur
Sulfur
16
Alchemical symbol.
Pallas
Sulfur
16
Alchemical symbol.
Sulfur
Sulfur
16
Alchemical symbol.
Young Ji International School / College
Page 51
♁
Sulfur
⊛
Magnesium Magnesium 12
Alchemical symbol.
♂
Mars
Iron
26
Alchemical symbol.
Stellae Fixae
Copper
29
Pre-16th century alchemical symbol.
Venus
Copper
29
Alchemical symbol.
Copper
Copper
29
Alchemical symbol.
Copper
Copper
29
Daltonian symbol circa 1808.
Arsenic
Arsenic
33
Alchemical symbol.
�
Arsenic
Arsenic
33
Alchemical symbol.
☽
Luna
Silver
47
Alchemical symbol.
�
Silver
Silver
47
Alchemical symbol.
♃
Jupiter
Tin
50
Alchemical symbol.
♁
Antimony
Antimony
51
Alchemical symbol.
Antimony
Antimony
51
Alchemical symbol.
♀
©
Sulfur
Young Ji International School / College
16
Daltonian symbol circa 1808.
Page 52
☾
Platinum
Platinum
78
Alchemical symbol.
☉
Platinum
Platinum
78
Alchemical symbol.
Uranus
Platinum
78
Alchemical symbol.
Sol
Gold
79
Alchemical symbol from the 16th century.
Sol
Gold
79
Alchemical symbol from 1700 through 1783.
Gold
Gold
79
Alchemical symbol.
Pisces
Mercury
80
Pre-16th century alchemical symbol.
Neptunus
Mercury
80
Alchemical symbol from the 17th century.
☿
Mercurius
Mercury
80
Alchemical symbol from 1700 through 1783.
♄
Saturnus
Lead
82
Alchemical symbol circa 1783.
Taurus
Bismuth
83
Alchemical symbol.
☼
�
Properties of Elements and their Uses: A closer Look on the main Group of Elements
Young Ji International School / College
Page 53
All elements are arranged in the periodic table. To recall, elements of similar characteristics are grouped together and are written vertically on the periodic table. Elements written in one vertical column belong to one group or family. There are eight families or groups and are collectively named as the group A elements, the representative or main group of elements. They are groups I-A, 2-A, 3-A, 4-A, 5-A, 6A, and 7-A, and 8-A. these elements are also aligned horizontally from left to right and are collectively called periods or series. The representative metals have valence electrons in their outermost s or p orbitals. All elements in Group 1-A (except H) and 2-A are metals. The heavier members of groups 3-A, 4-A and 5-A are called post-transition metals. The metallic property increases from top to bottom within group/family and decreases from right to left within a period or series. 
The Alkali Metals (Group 1-A)
The alkali metals found on Group 1-A of the periodic table are very reactive metals hence, do not occur freely in nature. These metals have only one valence electron. They are ever ready to lose that one electron in ionic bonding with other element to become stable. Alkali metals are malleable, ductile, and good conductors of heat and electricity. These metals are softer than most other metals. They produce hydrogen gas when reacted with water. Cesium and francium are the most reactive elements that they explode if exposed to water or even water vapor in the air. The members of this family are lithium, sodium, potassium, and rubidium. These metals should be stored under anhydrous nonpolar liquids like mineral oils or kerosene. General Properties 1. Alkali metals have long melting and boiling points resulting from their weak bonding forces. 2. They are excellent electrical and thermal conductors because of their loosely held outer electrons. 3. They ionize with low energy light. That is why cesium is used in photoelectric cell. 4. All alkali metals have positive (+) oxidation number. Their compounds are ionic in nature. 5. Several members of this group have similarities with other element. 
The Alkali Earth metals (Group 2-A)
These groups of metals, just like 1-A are not found free in nature. Each member has two electrons in the highest occupied energy level. They have an oxidation number of +2 and are very reactive. The members are beryllium, magnesium, calcium, strontium, barium and radium. Most 2-A metal compounds are Young Ji International School / College
Page 54
ionic but that beryllium have covalent character because it resembles Group 3-A, aluminum. Calcium and magnesium are abundant in the Earth‘s crust as carbonates and sulfates while all known radium isotopes are extremely rare and reactive.
Boron family (Group 3-A) Boron is at the top of the family and is a nonmetal. Its melting point is 2300°C, very high, because it crystallizes as a covalent solid. All Group 3-A elements are solids. Aluminum through thallium form metallic crystals with lower melting points.
Carbon family (Group 4-A)
Elements in Group 4-A are carbon, silicon, germanium, tin, lead, and Rutherfordium. The most commonly used elements are carbon and silicon. The main component of organic substances is carbon. Graphite, amorphous carbon, diamond, and fullerenes are pure forms of carbon and are called its allotropes. Allotropes are different structural modifications of an element. Silicon is shiny, blue gray, a brittle and with a high melting point metalloid. It is second to oxygen in abundance in the earth‘s crust, at about 87 percent of which occurs as silica.
Nitrogen family (Group 5-A)
Nitrogen and phosphorous are nonmetals, arsenic is predominantly nonmetallic; antimony is more metallic; and bismuth is definitely metallic. The oxidation states of the Group 5-A elements range from -3 to +5.
Oxygen family ( Group 6-A)
The oxygen family includes sulfur, selenium, tellurium, and oxygen. Oxygen, sulfur, and selenium are clearly nonmetallic; Tellurium is a metalloid and forms metal-like crystals; and polonium is a nonmetal, which has 29 radioactive isotopes. Except for oxygen, the elements under this group can be bonded covalently to as many as six other atoms.
The Halogens (Group 7-A) The halogens are five nonmetallic elements found in group 7A of the Periodic table. The term halogen means ‗salt former‖ and compounds containing halogens are called salts. The halogens exist at room temperature in three different phases. 1. Soild- Iodine and Astatine 2. Liquid- Bromine 3. Gas – Fluorine and Chlorine
Young Ji International School / College
Page 55
Properties of Halogens 1. They have high electronegativities indicating that they attract electrons strongly. 2. They from ionic compounds with metals Fluorine is actually, the most electronegative element. 3. They are the most reactive nonmetals. 4. They are obtained by oxidation of halide salts. Uses of Noble Gases or Inert gases 1. Helium- is used for filling of observations balloons and other lighter air craft. 2. Argon- provides an inert atmosphere for welding; also for filling incandescent light bulbs that inhibits vaporizations of tungsten filament and blackening of bulbs. 3. Neon- produces a bright red-orange glow even at low pressure and moderate electric current; used for neon signage in advertisements. 4. Krypton- used in airport runway and approach lights, gives longer life to incandescent light than argon but is more expensive. 5. Xenon- used with Krypton mixture in high intensity for short exposure photographic tubes; fast response to the electric current. 6. Radon- used for radioactive therapy of cancerous tissues.
The d-Transition Elements Transition elements refer to the elements in the middle of the periodic table. They provide transition between the ―base former‘ on the left and the ―acid former‘ on the right. The term also refers to both d and f transition elements (d and f orbitals are being filled).
General Properties of transition elements 1. All are metals, ductile, malleable and conduct electricity and heat. 2. Most are harder and more brittle, have higher melting points and heats of vaporization than non-transition elements. 3. Their ions and compounds are colored. 4. They from many complex ions. 5. They exhibit multiple oxidation states with few exceptions. 6. Many are paramagnetic, like iron, cobalt, and nickel. Uses of Transition Elements 1. Copper is used as a construction material, specifically for electrical wrings. 2. Gold, silver, and copper are used for making coins. Gold and silver are also useful in making jewelry because of their luster, malleability, and ductility. 3. Chromium is used for automobile parts and for electroplating. Young Ji International School / College
Page 56
4. 5. 6. 7. 8. 9.
Tungsten is used for lightbulb filaments. Titanium is used for bicycle frame and as aircraft construction material. Zinc is used for batteries. Silver is the best conduct for electricity and is widely used a jewelry. Cobalt is used for paints and pigments, for properties, and ink making. Iron is a component of the blood and is used extensively as a construction material.
Technology Update 1. Spectroscopy is an instrument technique used to determine the chemical composition, structure and mass, or concentration of an unknown compound. This technique takes advantage on the absorption and emission of energy of atoms or molecules. 2. The emission spectrum of a chemical element or chemical compound is the spectrum of frequencies of electromagnetic radiation emitted due to an atom or molecule making a transition from a high energy state to a lower energy state. The energy of the emitted photon is equal to the energy difference between the two states. There are many possible electron transitions for each atom, and each transition has a specific energy difference. These collections of different transitions, leading to different radiated wavelengths, make up an emission spectrum. Each element's emission spectrum is unique. Therefore, spectroscopy can be used to identify the elements in matter of unknown composition. Similarly, the emission spectra of molecules can be used in chemical analysis of substances. 3. Absorption spectroscopy refers to spectroscopic techniques that measure the absorption of radiation, as a function of frequency or wavelength, due to its interaction with a sample. The sample absorbs energy, i.e., photons, from the radiating field. The intensity of the absorption varies as a function of frequency, and this variation is the absorption spectrum. Absorption spectroscopy is performed across the electromagnetic spectrum. 4. Infrared spectroscopy (IR spectroscopy) is the spectroscopy that deals with the infrared region of the electromagnetic spectrum, that is light with a longer wavelength and lower frequency than visible light. It covers a range of techniques, mostly based on absorption spectroscopy. As with all spectroscopic techniques, it can be used to identify and study chemicals. For a given sample which may be solid, liquid, or gaseous, the method or technique of infrared spectroscopy uses an instrument called an infrared spectrometer (or spectrophotometer) to produce an infrared spectrum. A basic IR spectrum is essentially a graph of infrared light absorbance (or transmittance) on the vertical axis vs. frequency or wavelength on the horizontal axis. Typical units of frequency used in IR spectra are reciprocal centimeters (sometimes called wave numbers), abbreviated as cm−1. Units of IR wavelength are commonly given in microns, abbreviated as Îźm, which are related to wave numbers in a reciprocal way. A Young Ji International School / College
Page 57
common laboratory instrument that uses this technique is a Fourier transform infrared (FTIR) spectrometer. Two-dimensional IR is also possible as discussed below. 5. Ultraviolet–visible spectroscopy or ultraviolet-visible spectrophotometry (UV-Vis or UV/Vis) refers to absorption spectroscopy or reflectance spectroscopy in the ultraviolet-visible spectral region. This means it uses light in the visible and adjacent (near-UV and near-infrared [NIR]) ranges. The absorption or reflectance in the visible range directly affects the perceived color of the chemicals involved. In this region of the electromagnetic spectrum, molecules undergo electronic transitions. This technique is complementary to fluorescence spectroscopy, in that fluorescence deals with transitions from the excited state to the ground state, while absorption measures transitions from the ground state to the excited state. Areas of Investigation 1. Atmosphere is usually exhibited by oxides of some metalloids where they exhibit the properties of an acid in the presence of a base and a base in the presence of an acid. Ask the laboratory assistance of your chemistry laboratory and identify those substances that are amphoteric. 2. Test for compounds belonging to the transition metals and identify their colors. Why do these metals usually form colored salts? 3. Get some samples of familiar transition metals at home or from the laboratory and test their magnetic properties. List them down. Write the electron configuration of these metals and discover an explanation of their magnetic behavior. 4. Make a concept map of the transition elements considering their inherent properties and uses. 5. List and describe the known allotropes of carbon. 6. What is a doping agent? How is doping being done? Discussion Box 1. Why do metals especially in family 1-A and 2-A, lose electrons? 2. Why are transition metals paramagnetic? 3. Give the major differences between the representative elements and the transition metals? 4. What are allotropes? Give examples.
Young Ji International School / College
Page 58
Chapter 7
Chemical Bonds
The Lewis Electron Dot Structure It should be noted that when atoms bond, only their outermost or the valence shell electrons are involved. Lewis Electron Dot Structure (LEDS) can be used to show how the valence electrons of an atom interact with another atom. The LEDS consists of the chemical symbol, which represents the nucleus and dots to represent the valence or outermost electrons. Linus Pauling introduced the concept of electro negativity by assigning values from 0.8 to 4.0. In the periodic table, electronegativity values are increasing from left to right in a period and decreasing within a family. Electron affinity follows the same trend. TYPES OF CHEMICAL BONDS There are three types of chemical bonds that arise from atomic interactions. These are ionic, covalent, and metallic bonds. 7. Ionic bonds are formed by actual electron transfer. This occurs between atoms with large differences in tendency to lose or gain electrons. An ionic bond is formed between elements of group 1A or 2A with nonmetals. A metal atom has low IE, hence, can easily give up its one or two valence electrons to a nonmetallic atom. The nonmetal atom, on the other hand, has high electron affinity, hence, higher tendency to gain electrons. The electrostatic force of attraction between the positive and negative ions formed after electron transfer keeps the ion-pair together. During electron transfer, the electron donor (metal atom) becomes a positively-charged ion (cation) because it has more protons than electrons. Likewise, the electron acceptor (nonmetal atom) becomes a negatively charged ion (anion) because it has more electrons than protons. 8. Covalent bonds are formed between atoms by sharing their valence electrons in order for each atom to have eight valence electrons. Types of Covalent Bonds a. Nonpolar Covalent Bonds are formed by equal sharing of electrons. This made possible by bonding identical atoms (equal electronegativity values) to form a molecule. Two different atoms may also form a nonpolar covalent bond if the difference in their EN values is less than 0.4. For example, hydrogen Young Ji International School / College
Page 59
atom has one valence electron. Bonding between two hydrogen atoms would mean sharing each atom‘s only electron with each other. The two H-atoms each has 1s orbital coming together and bonding at an effective distance where the forces of negatively-charged electrons are overcome by attractive forces between two nuclei and the electron clouds overlap. The bond holding the two hydrogen atoms to from a hydrogen molecule is called nonpolar covalent bond. Similarly, two oxygen atoms form an oxygen molecule (O2). The oxygen atom has a configuration of 1s22s22p4 with six valence electrons. It needs two or more electrons to complete its 2p orbital and attain a closed configuration like neon (1s22s22p6). A covalent bond is a chemical bond that involves the sharing of electron pairs between atoms. The stable balance of attractive and repulsive forces between atoms when they share electrons is known as covalent bonding.[1] For many molecules, the sharing of electrons allows each atom to attain the equivalent of a full outer shell, corresponding to a stable electronic configuration. Covalent bonding includes many kinds of interactions, including Ďƒ-bonding, Ď€bonding, metal-to-metal bonding, agostic interactions, and three-center two-electron bonds. The term covalent bond dates from 1939. The prefix co- means jointly, associated in action, partnered to a lesser degree, etc.; thus a "co-valent bond", in essence, means that the atoms share "valence", such as is discussed in valence bond theory. In the molecule H2, the hydrogen atoms share the two electrons via covalent bonding. Covalence is greatest between atoms of similar electro negativities. Thus, covalent bonding does not necessarily require that the two atoms be of the same elements, only that they are of comparable electronegativity. Covalent bonding that entails sharing of electrons over more than two atoms is said to be delocalized.
A covalent bond forming H2 (right) where two hydrogen atoms share the two electrons Lone Pair In chemistry, a lone pair is a valence electron pair which is not shared with another atom and is sometimes called a non-bonding pair. Lone pairs are found in the outermost electron shell of atoms. They can be identified by using a Lewis Young Ji International School / College
Page 60
structure. Electron pairs are therefore considered lone pairs if two electrons are paired but are not used in chemical bonding. Thus, the number of lone pair electrons plus the number of bonding electrons equals the total number of valence electrons around an atom.
Lone pairs (shown as dots) in the Lewis structure of hydroxide
Lone pairs are a concept used in VSEPR theory which explains the shapes of molecules. They are also referred to in the chemistry of Lewis acids and bases. However not all non-bonding pairs of electrons are considered by chemists to be lone pairs. Examples are the transition metals where the non-bonding pairs do not influence molecular geometry and are said to be stereo chemically inactive.
Examples
Lone pairs in ammonia (A), water (B) and hydrogen chloride (C) A single lone pair can be found with atoms in the nitrogen group such as nitrogen in ammonia, two lone pairs can be found with atoms in the halogen group such as oxygen in water and the halogens can carry three lone pairs such as in hydrogen chloride. In VSEPR theory the electron pairs on the oxygen atom in water from the vertices of a tetrahedron with the lone pairs on two of the four vertices. The H–O– H bond angle is with 104.5°, less than the 109° predicted for a tetrahedral angle, and this can be explained by a repulsive interaction between the lone pairs.
Polarity of molecules
While the molecules can be described as "polar covalent", "nonpolar covalent", or "ionic", this is often a relative term, with one molecule simply being more Young Ji International School / College
Page 61
polar or more nonpolar than another. However, the following properties are typical of such molecules. A molecule is composed of one or more chemical bonds between molecular orbital of different atoms. A molecule may be polar either as a result of polar bonds due to differences in electro negativity as described above, or as a result of an asymmetric arrangement of nonpolar covalent bonds and non-bonding pairs of electrons known as a full molecular orbital.
Polar molecules
The dipole moment of the water molecule A polar molecule has a net dipole as a result of the opposing charges (i.e. having partial positive and partial negative charges) from polar bonds arranged asymmetrically. Water (H2O) is an example of a polar molecule since it has a slight positive charge on one side and a slight negative charge on the other. The dipoles do not cancel out resulting in a net dipole. Due to the polar nature of the water molecule itself, polar molecules are generally able to dissolve in water. Another example includes sugars (like sucrose), which have many polar oxygen–hydrogen (OH) groups and are overall highly polar.
The ammonia molecule, polar as a result of its molecular geometry. The red represents partially negatively charged regions. The hydrogen fluoride, HF, molecule is polar by virtue of polar covalent bonds — in the covalent bond electrons are displaced towards the more electronegative fluorine atom. Ammonia, NH3, molecule the three N–H bonds have only a slight polarity (toward the more electronegative nitrogen atom). However, the molecule has two lone electrons in an orbital, that points towards the fourth apex of the approximate tetrahedron, (VSEPR). This orbital is not participating in covalent bonding; it is electron-rich, which results in a powerful dipole across the whole ammonia molecule.
Young Ji International School / College
Page 62
In the ozone, O3, molecule the two O–O bonds are nonpolar (there is no electro negativity difference between atoms of the same element). However, the distribution of other electrons is uneven — since the central atom has to share electrons with two other atoms, but each of the outer atoms has to share electrons with only one other atom, the central atom is more deprived of electrons than the others (the central atom has a formal charge of +1, while the outer atoms each have a formal charge of −1/2). Since the molecule has a bent geometry, the result is a dipole across the whole ozone molecule.
Metallic bonding Metallic bonding occurs as a result of electromagnetism and describes the electrostatic attractive force that occurs between conduction electrons (in the form of an electron cloud of delocalized electrons) and positively charged metal ions. It may be described as the sharing of free electrons among a lattice of positively charged ions (cations). In a more quantum-mechanical view, the conduction electrons divide their density equally over all atoms that function as neutral (non-charged) entities. Metallic bonding accounts for many physical properties of metals, such as strength, ductility, thermal and electrical resistivity and conductivity, opacity, and luster. Metallic bonding is not the only type of chemical bonding a metal can exhibit, even as a pure substance. For example, elemental gallium consists of covalentlybound pairs of atoms in both liquid and solid state—these pairs form a crystal lattice with metallic bonding between them. Another example of a metal–metal covalent bond is mercurous ion (Hg2+2).
Van der Waals forces include attractions and repulsions between atoms, molecules, and surfaces, as well as other intermolecular forces. They differ from covalent and ionic bonding in that they are caused by correlations in the fluctuating polarizations of nearby particles (a consequence of quantum dynamics). Intermolecular forces have four major contributions: 1. 2. 3. 4.
Dipole –dipole forces Ion-dipole forces Induced dipole Dispersion or London forces
Young Ji International School / College
Page 63
A repulsive component resulting from the Pauli Exclusion Principle that prevents the collapse of molecules Attractive or repulsive electrostatic interactions between permanent charges (in the case of molecular ions), dipoles (in the case of molecules without inversion center), quadrupoles (all molecules with symmetry lower than cubic), and in general between permanent multipoles. The electrostatic interaction is sometimes called the Keesom interaction or Keesom force after Willem Hendrik Keesom. Induction (also known as polarization), which is the attractive interaction between a permanent multipole on one molecule with an induced multipole on another. This interaction is sometimes called Debye force after Peter J.W. Debye. Dispersion (usually named after Fritz London), which is the attractive interaction between any pair of molecules, including non-polar atoms, arising from the interactions of instantaneous multipoles. Returning to nomenclature, different texts refer to different things using the term "van der Waals force." Some texts describe the van der Waals force as the totality of forces (including repulsion); others mean all the attractive forces (and then sometimes distinguish van der Waals-Keesom, van der Waals-Debye, and van der Waals-London). All intermolecular/van der Waals forces are anisotropic (except those between two noble gas atoms), which means that they depend on the relative orientation of the molecules. The induction and dispersion interactions are always attractive, irrespective of orientation, but the electrostatic interaction changes sign upon rotation of the molecules. That is, the electrostatic force can be attractive or repulsive, depending on the mutual orientation of the molecules. When molecules are in thermal motion, as they are in the gas and liquid phase, the electrostatic force is averaged out to a large extent, because the molecules thermally rotate and thus probe both repulsive and attractive parts of the electrostatic force. Sometimes this effect is expressed by the statement that "random thermal motion around room temperature can usually overcome or disrupt them" (which refers to the electrostatic component of the Van der Waals force). Clearly, the thermal averaging effect is much less pronounced for the attractive induction and dispersion forces. The Lennard-Jones potential is often used as an approximate model for the isotropic part of a total (repulsion plus attraction) Van der Waals force as a function of distance. Van der Waals forces are responsible for certain cases of pressure broadening (van der Waals broadening) of spectral lines and the formation of van der Waals molecules. The London-van der Waals forces are related to the Casimir effect for dielectric media, the former being the microscopic description of the latter bulk property. The first detailed calculations of this were done in 1955 by E. M. Lifshitz. Young Ji International School / College
Page 64

Its main characteristics are:1. They are weaker than normal covalent ionic bonds. 2. Van der Waals forces are additive and cannot be saturated. 3. They have no directional characteristic. 4. They are all short - range forces and hence only interactions between nearest need to be considered instead of all the particles. The greater is the attraction if the molecules are closer due to Van der Waals forces.
Van der Waals forces are independent of temperature except dipole - dipole interactions.
Attractive interactions resulting from dipole-dipole interaction of two hydrogen chloride molecules Van der Waals forces include attractions and repulsions between atoms, molecules, VSEPR theory Valence shell electron pair repulsion (VSEPR) theory is a model used, in chemistry, to predict the geometry of individual molecules from the number of electron pairs surrounding their central atoms. It is also named Gillespie– Nyholm theory after its two main developers. The acronym "VSEPR" is occasionally pronounced "vesper" or "vuh-seh-per". The premise of VSEPR is that the valence electron pairs surrounding an atom tend to repel each other, and will therefore adopt an arrangement that minimizes this repulsion, thus determining the molecule's geometry. The sum of the number of atoms bonded to a central atom and the number of lone pairs formed by its nonbonding valence electrons is known as the central atom's steric number. VSEPR theory is usually compared with valence bond theory, which addresses molecular shape through orbitals that are energetically accessible for bonding. Valence bond theory concerns itself with the formation of sigma and pi bonds. Molecular is another model for understanding how atoms and electrons are assembled into molecules and polyatomic ions. VSEPR theory has long been criticized for not being quantitative, and therefore limited to the generation of "crude" (though structurally accurate) molecular geometries of covalently-bonded molecules. However, molecular mechanicsforce fields based on VSEPR have also been developed. 
Description
Young Ji International School / College
Page 65
VSEPR theory, occasionally pronounced "vesper" or "vuh-seh-per", is used to predict the arrangement of electron pairs around non-hydrogen atoms in molecules, especially simple and symmetric molecules, where these key, central atoms participate in bonding to 2 or more other atoms; the geometry of these key atoms and their non-bonding electron pairs in turn determine the geometry of the larger whole. The number of electron pairs in the valence shell of a central atom is determined after drawing the Lewis structure of the molecule, and expanding it to show all electron-pair bonds and lone pairs *of* electrons. For the purposes of VSEPR theory, the multiple electron pairs in a double bond or triple bond are treated as though they were a bond with single pair of electrons. In cases where a molecule can be depicted by two or more resonance structures, these structures generally differ only by the interchange of double and single bonds, so that they have the same steric number and therefore the same VSEPR model. The electron pairs are assumed to lie on the surface of a sphere centered on the central atom and tend to occupy positions that minimize their mutual repulsions by maximizing the distance between them. Gillespie has emphasized that the electron-electron repulsion due to the Pauli exclusion principle is more important in determining molecular geometry than the electrostatic repulsion. The number of electron pairs, therefore, determines the overall geometry that they will adopt. For example, when there are two electron pairs surrounding the central atom, their mutual repulsion is minimal when they lie at opposite poles of the sphere. Therefore, the central atom is predicted to adopt a linear geometry. If there are 3 electron pairs surrounding the central atom, their repulsion is minimized by placing them at the vertices of an equilateral triangle centered on the atom. Therefore, the predicted geometry is trigonal. Likewise, for 4 electron pairs, the optimal arrangement is tetrahedral. The difference between lone pairs and bonding pairs may also be used to rationalize deviations from idealized geometries. For example, the H2O molecule has four electron pairs in its valence shell: two lone pairs and two bond pairs. The four electron pairs are spread so as to point roughly towards the apices of a tetrahedron. However, the bond angle between the two O-H bonds is only 104.5°, rather than the 109.5° of a regular tetrahedron, because the two lone pairs (whose density or probability envelopes lie closer to the oxygen nucleus) exert a greater mutual repulsion than the two bond pairs. 
AXE method
The "AXE method" of electron counting is commonly used when applying the VSEPR theory. The A represents the central atom and always has an implied subscript one. The X represents the number of ligands (atoms bonded to A).
Young Ji International School / College
Page 66
The Erepresents the number of lone electron pairs surrounding the central atom. The sum of X and E is known as the steric number.
Molecule Type
Shape
AX2E0
Linear
BeCl2, HgCl2, CO2
AX2E1
Bent
NO− 2, SO2, O3, CCl2
AX2E2
Bent
H2O, OF2
AX2E3
Linear
XeF2, I−3, XeCl2
AX3E0
Trigonal planar
BF3, CO2−3, NO− 3, SO3
AX3E1
Trigonal pyramidal
NH3, PCl3
AX3E2
T-shaped
ClF3, BrF3
Young Ji International School / College
Electron arrangement
Geometry
Examples
Page 67
AX4E0
Tetrahedral
CH4, PO3− 4, SO2− 4, ClO− 4, TiCl4, XeO4
AX4E1
Seesaw
SF4
AX4E2
Square planar
XeF4
AX5E0
Trigonal bipyramidal
PCl5
AX5E1
Square pyramidal
ClF5, BrF5, XeOF4
AX5E2
Pentagonal planar
XeF−5
AX6E0
Octahedral
SF6, WCl6
AX6E1
Pentagonal pyramidal
XeOF− 5, IOF2−5
Young Ji International School / College
Page 68
Pentagonal bipyramidal
IF7
AX8E0
Square antiprismatic
IF− 8, ZrF4− 8, ReF−8
AX9E0
Tricapped trigonal prismatic(as drawn) OR capped square antiprismatic
ReH2−9
AX7E0
† Electron arrangement including lone pairs, shown in pale yellow ‡ Observed geometry (excluding lone pairs) When the substituent (X) atoms are not all the same, the geometry is still approximately valid, but the bond angles may be slightly different from the ones where all the outside atoms are the same. For example, the double-bond carbons in alkenes like C2H4 are AX3E0, but the bond angles are not all exactly 120°. Likewise, SOCl2 is AX3E1, but because the X substituents are not identical, the XAX angles are not all equal. As a tool in predicting the geometry adopted with a given number of electron pairs, an often used physical demonstration of the principle of minimal electron pair repulsion utilizes inflated balloons. Through handling, balloons acquire a slight surface electrostatic charge that results in the adoption of roughly the same geometries when they are tied together at their stems as the corresponding number of electron pairs. For example, five balloons tied together adopt the trigonal bipyramidal geometry, just as do the five bonding pairs of a PCl5 molecule (AX5) or the two bonding and three non-bonding pairs of a XeF2 molecule (AX2E3). The molecular geometry of the former is also trigonal bipyramidal, whereas that of the latter is linear. Molecular geometry
Young Ji International School / College
Page 69
Geometry of the water molecule Molecular geometry is the three-dimensional arrangement of the atoms that constitute a molecule. It determines several properties of a substance including its reactivity, polarity, phase of matter, color, magnetism, and biological activity.[1][2] The angles between bonds that an atom forms depend only weakly on the rest of molecule, i.e. they can be understood as approximately local and hence transferable properties.
Bonding
Molecules, by definition, are most often held together with covalent bonds involving single, double, and/or triple bonds, where a "bond" is a shared pair of electrons (the other method of bonding between atoms is called ionic bonding and involves a positive cation and a negative anion). Molecular geometries can be specified in terms of bond lengths, bond angles and torsional angles. The bond length is defined to be the average distance between the centers of two atoms bonded together in any given molecule. A bond angle is the angle formed between three atoms across at least two bonds. For four atoms bonded together in a chain, thetorsional angle is the angle between the plane formed by the first three atoms and the plane formed by the last three atoms. Types of molecular structure
Linear: In a linear model, atoms are connected in a straight line. The bond angles are set at 180°. A bond angle is very simply the geometric angle between two adjacent bonds. For example, carbon dioxide and nitric oxide have a linear molecular shape. Trigonal planar: Just from its name, it can easily be said that molecules with the trigonal planar shape are somewhat triangular and in one plane (flat). Consequently, the bond angles are set at 120°. An example of this is boron trifluoride. Bent: Bent or angular molecules have a non-linear shape. A good example is water, or H2O, which has an angle of about 105°. A water molecule has two pairs of bonded electrons and two unshared lone pairs. Tetrahedral: Tetra- signifies four, and -hedral relates to a face of a solid, so "tetrahedral" literally means "having four faces". This shape is found when
Young Ji International School / College
Page 70
there are four bonds all on one central atom, with no extra unshared electron pairs. In accordance with the VSEPR (valence-shell electron pair repulsion theory), the bond angles between the electron bonds are arccos(−1/3) = 109.47°. An example of a tetrahedral molecule is methane (CH4). Octahedral: Octa- signifies eight, and -hedral relates to a face of a solid, so "octahedral" literally means "having eight faces". The bond angle is 90 degrees. An example of an octahedral molecule is sulfur hexafluoride (SF6). Trigonal pyramidal: A trigonal pyramidal molecule has a pyramid-like shape with a triangular base. Unlike the linear and trigonal planar shapes but similar to the tetrahedral orientation; pyramidal shapes require three dimensions in order to fully separate the electrons. Here, there are only three pairs of bonded electrons, leaving one unshared lone pair. Lone pair – bond pair repulsions change the bond angle from the tetrahedral angle to a slightly lower value. An example is NH3 (ammonia).
Discussion Box 1. Write the Lewis Electron Dot Structure (LEDS)f the following: a. Na = 11 b. Mg = 12 c. Br = 35 d. Sc = 21 2. Tell the kind of bonds formed in the following: a. Cr b. MgCl2 c. O3 d. HCl e. CaO f. Cl2 3. Tell whether the molecules are polar or nonpolar. a. Br2 b. SF3 c. PCI5 d. H2O e. CHCI3 4. What is the Octet Rule? 5. Name the different kind of Van der Waals forces. 6. Give the difference between a polar and a nonpolar bond. 7. What makes metals good conductors of heat and electricity?
Young Ji International School / College
Page 71
Chapter 8 
Changes in Matter
Chemical Formula: A Chemist’s Shorthand
Chemical formula is a way of expressing information about the proportions of atoms that constitute a particular chemical compound, using a single line of chemical element symbols, numbers, and sometimes also other symbols, such as parentheses, dashes, brackets, and plus (+) and minus (−) signs. These are limited to a single typographic line of symbols, which may include subscripts and superscripts. A chemical formula is not a chemical name, and it contains no words. Although a chemical formula may imply certain simple chemical structures, it is not the same as a full chemical structural formula. Chemical formulas are more limiting than chemical names and structural formulas. Is a representation made up of symbols of elements and the numerical subscripts showing the type and number of each atom present in unit of the substance? It represents the compound including the fixed ratio of the atom of elements that combine to become a compound. A molecular formula represents the actual number of atoms of each element in a compound. Example: CO2
A structural formula shows the actual number of atoms, their relative placement and the bond between them. Example: O=C=O
Diatomic molecules are molecules composed of only two atoms, of either the same or different chemical elements. The prefix di- is of Greek origin, meaning "two". If a diatomic molecule consists of two atoms of the same element, such as hydrogen (H2) or oxygen (O2), then it is said to be homonuclear. Otherwise, if a diatomic molecule consists of two different atoms, such as carbon monoxide (CO) or nitric oxide (NO), the molecule is said to be heteronuclear. A compound is represented by a chemical formula. We could say that a symbol is to element as formula is to compound. To write the formula of a compound, oxidation numbers of the elements should be known. Before a chemical formula can be written, a mastery of the symbols of the elements is needed. More importantly, a mastery of the names of polyatomic ions and their corresponding charges or oxidation numbers is supposed to be memorized by heart. Young Ji International School / College
Page 72
The formula of a compound results from a combination of elements in proper ratio that the sum of electrical charges of the ions or atoms that comprise the chemical formula is zero. Oxidation Numbers The oxidation state, often called the oxidation number, is an indicator of the degree of oxidation (loss of electrons) of an atom in a chemical compound. Conceptually, the oxidation state, which may be positive, negative or zero, is the hypothetical charge that an atom would have if all bonds to atoms of different elements were 100% ionic, with no covalent component. This is never exactly true for real bonds. [Oxidation state] is defined as the charge an atom might be imagined to have when electrons are counted according to an agreed-upon set of rules: 1. the oxidation state of a free element (uncombined element) is zero 2. for a simple (monatomic) ion, the oxidation state is equal to the net charge on the ion 3. hydrogen has an oxidation state of +1 and oxygen has an oxidation state of −2 when they are present in most compounds. Exceptions to this are that hydrogen has an oxidation state of −1 in hydrides of active metals, e.g. LiH, and oxygen has an oxidation state of −1 in peroxides, e.g. H2O2. 4. the algebraic sum of oxidation states of all atoms in a neutral molecule must be zero, while in ions the algebraic sum of the oxidation states of the constituent atoms must be equal to the charge on the ion. Polyatomic Ions Polyatomic ions consist of two or more atoms bonded covalently and have a net positive or negative charge. It has its own identity and stays as a unit in its interaction with other ions. An example of a polyatomic ion is the hydroxide ion - consisting of one oxygen atom and one hydrogen atom, hydroxide has a charge of −1. Its chemical formula is OH−. An ammonium ion is made up of one nitrogen atom and four hydrogen atoms: it has a charge of +1, and its chemical formula is NH4+. Polyatomic ions are often useful in the context of acid-base chemistry or in the formation of salts. A polyatomic ion can often be considered as the conjugate acid/base of a neutral molecule. For example, the conjugate base of sulfuric acid (H2SO4) is the polyatomic hydrogen sulfate anion (HSO4-). The removal of another hydrogen ion yields the sulfate anion (SO42-). The following tables give examples of commonly encountered polyatomic ions.
Young Ji International School / College
Page 73
Anions
Acetate (ethanoate)
CH3COO− or C2H3O−2
Acetylide
C2−2
Benzoate
C6H5COO−or C7H5O−2
Carbonate
CO2−3
Chromate
CrO2−4
Citrate
C6H5O3−7
Cyanide
CN−
Hypochlorite
ClO−
Chlorite
ClO−2
Chlorate
ClO−3
Perchlorate
ClO−4
Dichromate
Cr2O2−7
Dihydrogen phosphate
H2PO−4
Hydrogen carbonate (bicarbonate) HCO−3
Young Ji International School / College
Page 74
Hydrogen sulfate (bisulfate)
HSO−4
Hydrogen phosphate
HPO2−4
Hydroxide
OH−
Nitrite
NO−2
Nitrate
NO−3
Peroxide
O2−2
Permanganate
MnO−4
Phosphate
PO3−4
Sulfite
SO2−3
Sulfate
SO2−4
Cations
Ammonium
NH+4
Phosphonium PH+4
Hydronium
H3O+
Fluoronium
H2F+
Young Ji International School / College
Page 75
Mercury(I)
Hg2+2
Tropylium
C7H+7
Guanidinium C(NH2)+3
Binary compound (two nonmetals) A binary compound is a chemical compound that contains exactly two different elements. Examples of binary ionic compounds include calcium chloride (CaCl2), sodium fluoride (NaF), and magnesium oxide (MgO), whilst examples of a binary covalent compounds include water (H2O), and sulfur hexafluoride (SF6). Binary acids In the group of binary compounds of hydrogen a binary acid will include a hydrogen atom attached to another atom that will typically be in the 17th group of the periodic Table. These include chlorine, fluorine, bromine, iodine, and astatine. Others such as sulfur, tellurium, polonium, selenium, and arsenic are also included. The naming convention is: ―Hydro-‖ + Nonmetal + ―-ic‖ + ―acid‖ An example is HCl: hydrochloric acid. If the acid is in a gaseous form or an anhydrous form, the "-ic" is replaced by "-ide" and the "acid" suffix is removed.
Binary covalent compounds Nonmetal X + Nonmetal Y + "-ide." Add the appropriate Latin prefix to each element name to denote the number of atoms of each element present in a molecule of the compound. This method is generally not used with ionic compounds (see below). For example, K2O is usually not called dipotassium monoxide; it is simply potassium oxide. The reason that it is called potassium oxide is that potassium oxide is a binary ionic compound, thus it follows the rules for binary ionic compounds. P4O6, however, would be tetraphosphorus hexoxide. Some elements beginning with vowels (Oxygen, for
Young Ji International School / College
Page 76
example) replace the vowel ending of its prefix; mono- + Oxide = Monoxide, O4 = Tetroxide, O5 = Pentoxide, and so on. Binary ionic compounds A binary ionic compound is a salt consisting of only two elements in which both elements are ions, a cation (which has a positive charge) and an anion (which has a negative charge). When naming these compounds, its composition must be considered. Type 1 binary ionic compounds are those in which the cation has only one form, or charge. Type 2 binary ionic compounds are those in which the cation can have multiple forms. 1. The cation (which is a metal in most cases) is listed first and the anion (which is a nonmetal in most cases) second. 2. The cation takes the name of its elemental form. For example, Li+ would be called "Lithium". 3. The anion name uses the first part of its elemental name, and subsequently adding the suffix "-ide". For example, Br- would be called "Bromide" .Examples: LiF (which is composed of Li+ cation and F− anion) = lithium fluoride BaO (which is composed of Ba2+ cation and O2- anion) = barium oxide. Metals used are transition metals except for Al3+, Zn2+, Ag+. 1. The steps follow those of Type 1 Binary Ionic compounds however, since the cation can take on multiple charges, the charge must be written within parentheses in Roman numerals after stating the cation name. Examples: CoO (which is composed of Co²+ cation and O²- anion) = cobalt(II) oxide FeN (which is composed of Fe³+ cation and N³- anion) = iron(III) nitride Note that there is another way to name Type 2 ionic compounds that is not as common. This involves using an alternate, Latin name for the cation. Common Type 2 cations include Iron, Copper, Cobalt, Tin, Lead, and Mercury. When naming binary compounds with polyatomic ions: 1. The cation is listed first and the anion second. 2. The polyatomic ion names must be memorized. 3. No extra suffixes are added. Examples: NaCN (which is composed of Na+ cation and CN- polyatomic anion) = sodium cyanide NH4CN (which is composed of NH4+ polyatomic cation and Cl- anion) =ammonium chloride Cation + Anion + "-ide" (for anions consisting of individual elements, such as nitride) When multiple compounds are possible, the oxidation state of the cation is added after it in Roman numerals (copper (II) sulfide), or the cation's stem is used with a -ous or -ic suffix (cupric sulfide). Young Ji International School / College
Page 77
Chemical Equations A chemical equation is the symbolic representation of a chemical reaction wherein the reactant entities are given on the left-hand side and the product entities on the right-hand side. The coefficients next to the symbols and formulae of entities are the absolute values of the stoichiometric numbers.
A chemical equation consists of the chemical formulas of the reactants (the starting substances) and the chemical formula of the products (substances formed in the chemical reaction). The two are separated by an arrow symbol ( , usually read as "yields") and each individual substance's chemical formula is separated from others by a plus sign. As an example, the equation for the reaction of hydrochloric acid with sodium can be denoted: 2 HCl + 2 Na → 2 NaCl + H2
Symbols are used to differentiate between different types of reactions. To denote the type of reaction:
"
" symbol is used to denote a stoichiometric relation.
" " "
" symbol is used to denote a net forward reaction. " symbol is used to denote a reaction in both directions. " symbol is used to denote an equilibrium.
Physical state of chemicals is also very commonly stated in parentheses after the chemical symbol, especially for ionic reactions. When stating physical state, (s) denotes a solid, (l) denotes a liquid, (g) denotes a gas and (aq) denotes an aqueous solution. If the reaction requires energy, it is indicated above the arrow. A capital Greek letter delta ( ) is put on the reaction arrow to show that energy in the form of heat is added to the reaction. is used if the energy is added in the form of light.
1. Before balancing an equation, check but each formula is written correctly. 2. Balancing is done by placing coefficients on the left side of the formulas to make sure that the same number of atoms of each element
Young Ji International School / College
Page 78
is found on both sides of the equation. NEVER alter the formula by adding or changing subscripts. 3. Always consult the Activity series of metals and nonmetals for replacement reactions. An equation should always be balanced in terms of number of atoms of each element at both side of the equation. To do this, count the existing number of an atom of each element and write the appropriate number before the formula to obtain the desired number of atoms that would balance. This number is referred to as coefficient. Coefficients are placed at the left side of the formula. For simple equations, balancing can be done by simple inspection but for oxidation-reduction (redox) equation more complicated steps are to be followed.
Types of Chemical Reaction
Color changes - Different combinations of molecules reflect light differently. A color change indicates a change in molecules.
Heat content changes - In all chemical reactions, the heat content of the reactants and the heat content of the products is never the same. Sometimes the difference is great and can be easily detected. At other times, the difference is slight and more difficult to detect.
Gas produced - Whenever a gaseous product forms in a liquid solution, bubbles can be seen. A colorless gas produced in a reaction of solids is much harder to detect.
Precipitate forms - Precipitates are insoluble products formed by a reaction taking place in a liquid solution. This insoluble product will eventually settle to the bottom, but might immediately appear by turning the clear solution cloudy.
Most chemical reactions can be placed into one of five basic types: 1. Decomposition Reactions
A compound breaks into parts.
compound → element + element
2H2O → 2H2 + O2
Some decomposition complications with heat:
Some acids, when heated, decompose into an acidic oxide and H 2O.
Young Ji International School / College
Page 79
H2SO3 → SO2 + H2O
Metallic hydroxides, when heated, decompose into a metallic oxide and H 2O. Ca(OH)2 → CaO + H2O
Metallic carbonates, when heated, decompose into a metallic oxide and CO2. Li2CO3 → Li2O + CO2
Metallic chlorates, when heated, decompose into metallic chlorides and O 2. 2KClO3 → 2KCl + 3O2
2. Synthesis Reactions
Elements are joined together.
element + element → compound
2H2 + O2 → 2H2O
Compounds are joined together
compound + compound → compound
6CO2 + 6H2O → C6H12O6 + 6O2
3. Single Displacement Reactions
A single element replaces an element in a compound.
element + compound → element + compound
Zn + 2HCl → H2 + ZnCl2
4. Double Displacement Reactions
An element from each of two compounds switch places.
compound + compound → compound + compound
Young Ji International School / College
Page 80
H2SO4 + 2NaOH → Na2SO4 + 2H2O
5. Combustion Reactions
A hydrocarbon (a compound containing only carbon and hydrogen) combines with oxygen.
The products of combustion are always carbon dioxide and water.
hydrocarbon + oxygen → carbon dioxide + water
CH4 + 2O2 → CO2 + 2H2O
When metallic substances combine with oxygen, the result is an oxidationreduction reaction. The rusting of iron - 4Fe + 3O2 → 2Fe2O3
Chemical reactions can be classified in other ways as well: Neutralization Reactions
Special types of double displacement reactions that involve the reaction between an acid and base to form a salt and water.
acid + base → salt + water
Heat is usually given off in neutralization reactions.
A suspension of solid magnesium hydroxide in water is widely used as an antacid to neutralize excess stomach acid: Mg(OH)2 (s) + 2HCl (aq) → MgCl2 (aq) + 2H2O (l)
Oxidation-Reduction Reactions
Any reaction in which elements experience a change in oxidation number.
one atom gains e&minus and another atom looses e&minus S + O2 → SO2
In the reaction above, sulfur and oxygen both have an oxidation number of zero before the reaction. After the reaction, sulfur is +4 and oxygen is −2.
Precipitation Reactions Young Ji International School / College
Page 81
Aqueous reactions that involve the formation of a precipitate (solid).
soluble compound + soluble compound → insoluble compound 2KI (aq) + Pb(NO3)2 (aq) → 2KNO3 (aq) + PbI2 (s)
The physical state symbol (aq) says the reaction is taking place in a water solution. The physical state symbol (s) says the lead (II) iodide is a solid therefore insoluble in the solution.
1. Combination or Synthesis or Direct Union. Reactants combined to form one or more products General Form: A + B ---------- AB 2. Decomposition A compound is broken down into simpler forms by use of light (photolysis), electricity (electrolysis) heat or flame (pyrolysis0 and water (hydrolysis) General Form: AB --------- A + B
3. Substitute or Single displacement or Replacement General Form: AB + C -------- A + CB If C is a more active metal than A Or AB + C -------- AC + B If C is a more active nonmetal than B
Let us try this! Predict the product and balance the equation: 1. Br 2 + CaCI2 ----------2. Zn + Hg(NO3)2 ----------3. K + H2O ----------4. Al + HCI ----------5. Fe + CuSO4 -----------
4. Double decomposition or Double Displacement or Metathesis
Young Ji International School / College
Page 82
General Form:
AB
+
CD
------------- AD + CB
Predicting Products of Chemical Reactions Although products formed from chemical reactions can only be determined in the laboratory, somehow the probable products can be predicted of the combining abilities, oxidation number, and other relationships are known. Products between a metal and nonmetal could be determined, knowing that a metal transfer its electrons, thus acquiring a positive charge and nonmetal accepting electrons become negatively charged. For example: Sodium metal reacts with chlorine (nonmetal) to form sodium chloride. Science Ideas
The six states of matter are solid, liquid, gas, plasma, Bose-Einstein condensate Chemical change involves a change in the chemical composition of the substance whereby new substances are formed with a different set of intensive properties. A physical change is one which involves only a change in phase without altering the chemical composition of the substance. Chemical equation is a set of symbols and formulas that represents the chemical change that occurs in the reaction. Reactants and product are the components of a chemical equation. The reactant at the left side of the arrow is the substances being changed while the products, at the right side of the arrow, are the new substances or substances formed. Oxidation number is the positive or negative number representing the electrons lost or gained by an atom or group of atoms during the chemical reaction. The Activity Series is an arrangement of elements based on their reactivity. This is used in replacement reactions. There are four types of chemical reactions, namely synthesis or combination. Decomposition, replacement and metathesis or double decomposition. Diatomic molecules are made up of two similar kinds of atoms. A molecule is an electrically neutral group of two or more atoms held together by chemical bonds. Molecules are distinguished from ions by their lack of electrical charge. However, in quantum physics, organic chemistry, and
Young Ji International School / College
Page 83
 
biochemistry, the term molecule is often used less strictly, also being applied to polyatomic ions. A molecular formula shows the actual number of atoms of each element in a compound or molecule. It describes s the composition of a substance. A structural formula shows the number of atoms and the bonds formed between them and the relative placement and connections of atoms in the molecule.
Areas of Investigation 1. Conduct a home activity where you will investigate what would happen to sulfur (nonmetal) when heated in air and the oxide is treated with water. You will need some red and blue litmus paper for this activity. 2. From the activity predict what is formed when you repeat the same procedure, this using a metallic element like magnesium ribbon. 3. How would you prove that water from our faucet is chlorinated from source? 4. Try to put vinegar to some pulverized eggshells. What kind of change results? Give some evidences of the changes to justify your answer.
Discussion Box 1. Name the following binary compounds: a. AuBr3 b. CaI2 c. Cu2O d. NO2 e. MgF2 f. SO3 2. Name the following acids. a. H2SO3 b. HCN(aq) c. H3PO4 d. HCIO2 e. HNO3 f. HI (aq) 3. Write the formula of the following: a. carbon tetrachloride b. hypochlorous acid c. disphosphorous pentoxide Young Ji International School / College
Page 84
d. e. f. g. h.
tin (IV) chloride hydrogen sulfide carbonic acid mercuric oxide chlorous acid
Chapter 9
Laws Governing Chemical Changes
The Law of Conservation of Mass The law of conservation of mass, or principle of mass conservation, states that for any system closed to all transfers of matter and energy (both of which have mass), the mass of the system must remain constant over time, as system mass cannot change quantity if it is not added or removed. Hence, the quantity of mass is "conserved" over time. The law implies that mass can neither be created nor destroyed, although it may be rearranged in space, or the entities associated with it may be changed in form, as for example when light or physical work is transformed into particles that contribute the same mass to the system as the light or work had contributed. The law implies (requires) that during any chemical reaction, nuclear reaction, or radioactive decay in an isolated system, the total mass of the reactants or starting materials must be equal to the mass of the products. A French chemist, Antoine Lavoisier found that mass just like energy is concerned in a chemical reaction. For any natural course of change in matter the amount of substances before and after the change remains equal. The law of Conservation of Mass states that the total mass of reactants and total mass of products formed is equal. Consequently, the numbers of atoms are equal to the number of atoms on the product side. For instance if we have hydrogen gas reacting with oxygen gas to form water, then: a. in terms of mass 2H2
+
O2 -------- 2H20
(4)(1 amu) + (2)(16) = (4)(1amu) + (2)(16amu) 4 amu +
32 qmu = 4 amu + 32 amu 36 amu = 36 amu
b. in terms of number of atoms
Young Ji International School / College
Page 85
2H2 + O2 -------- 2H20 4 atoms + 2 atoms = 4 atoms + 2 atoms 6 atoms
= 6 atoms
The Law of definite Composition In a chemical reaction, substances react in proportional amounts. The reactants of a particular reaction would not just combine in any amount available. Rather, a fixed amount of one substance reacts with a definite amount of the other to form a product. French chemist Joseph Louis Proust (1754-1826) stated that the amounts of constituents in a particular compound are always definite and constant. This is the essence of the Law of Definite Composition, which states that there should be a definite amount of one substance to react with certain amount of another to form a compound and the masses are related to each other in ratios of small whole numbers. This idea of Proust came out eight years earlier than Dalton‘s atomic theory. For instance, water is analyzed and the sample yields the same ratio of hydrogen to oxygen, which is 1:8. This is the fixed ratio of the components of water because the ratio of the atom in a given compound is fixed. The compound water will always have this ratio
Sample Problem 2 Eight (8) g of hydrogen is made to react with 66 g of oxygen. The mass ratio of hydrogen and oxygen is 1:8 and given the equation. 2H2
+
O2 ----------- 2H2O
a. How many grams of water is formed? b. How many grams of hydrogen and oxygen reacted completely with each other? c. Is there an excess amount unreached? How many grams? d. Is there a limiting reagent? If so, what and how many grams? e. Does it conform to the law of Conservation of Mass?
Sample Problem 3 a. Suppose 20 g of Substance A reacted completely with 60 g of Substance B. what is the mass of the product? Does it conform to the law of Conservation of Mass? Young Ji International School / College
Page 86
b. Having in mind the above ratio of A and B. given 90 g of b, how many grams of A will completely react with B? How much of the product (AB) is formed? Does it still conform to the Law of Conservation of Mass? c. If 15 g of A and 50 g of B are allowed to react, what is the total mass of product formed? Exercise 1 1. Given 6.4 g of sulfur was burned in air to form 12.8 g of sulfur dioxide. How much oxygen from the air was used in the reaction? What is the ratio between sulfur and oxygen? If 50 g of sulfur will be burned with 100 g of oxygen in air, how many grams of sulfur dioxide will be formed? Is there a limiting reagent? Excess reagent? What are these?
The law of multiple Proportions Another example of the law can be seen by comparing ethane (C2H6) with propane (C3H8). The weight of hydrogen which combines with 1 g carbon is 0.252 g in ethane and 0.224 g in propane. The ratio of those weights is 1.125, which can be expressed as the ratio of two small numbers 9:8.
The law of multiple proportions demonstrated with oxygen and 1.00 gram of nitrogen Different Compounds of Carbon and Oxygen Compounds
Masses
Mass ratio C/O
Ratio of O in CO and Co2
CO
12 g-16 g
ž
1:2
CO2
12 g-32 g
3/8
Young Ji International School / College
Page 87
Formula Mass/Molecular Mass A mass formula is an equation or set of equations in physics which attempts to predict the mass or mass ratios of the subatomic particles. An important step in high energy physics was the discovery of the Gell-Mann– Okubo mass formula predicting relationships between masses of the members of SU (3) multiples. The development of an accurate mass formula is one of several fundamental aspects to developing a working theory of everything, which is expected to overcome the incompatibilities between current classical and quantum physics theories.
A formula mass or molecular mass is the sum of the atomic mass of the atoms or ions in a molecule or formula unit of a compound. Since atomic masses are expressed in atomic mass unit (amu), formula masses are commonly expressed in amu.
Formula Mass/ Molecular Mass Calculation Sample Problem 1 Lets us calculate the molecular mass of water (H2O). Element H O
Number of Atoms x Atomic Mass 2 x 1 amu 1 x16 amu
Total Mass 2 amu 16 amu
Molecular Mass of H2O = 18 amu
Sample Problem 2
Determine the formula mass of table salt (NaCI)
Element Na CI
Number of Atoms x Atomic Mass 1 x 23 amu 1 x 35 amu
Young Ji International School / College
Total Mass 23 amu 35 amu
Page 88
Formula Mass of NaCI = 58 amu
Formula Mass/Molecular Mass
Mole is a unit of measurement used in chemistry to express amounts of a chemical substance, defined as the amount of any substance that contains as many elementary entities (e.g., atoms, molecules, ions, electrons) as there are atoms in 12 grams of pure carbon-12 (12C), the isotope of carbon with relative atomic mass of exactly 12 by definition. This corresponds to the Avogadro constant, which has a value of 6.02214129(27) ×1023 elementary entities of the substance. It is one of the base units in the International System of Units; it has the unit symbol mol and corresponds with the dimension symbol N.
The Mole as Counting Unit
We always quantify things in our everyday living. We used to count the number of teaspoon of sugar and of coffee powder for our cup of coffee, the number of close friends or relatives we have, the pieces of our jewelry and many others.
However, if the number of things to be counted involves a great number of pieces such as the number of rice grains or banana fruits harvested in a hectare of farm, counting by pieces then is impractical. Instead, we use ―cavan‖ or ―bunch‖ for such cases.
Since its adoption into the International System of Units in 1971, there have been a number of criticisms of the concept of the mole as a unit like the meter or the second:
the number of molecules, etc. in a given lump of material is a fixed dimensionless quantity that can be expressed simply as a number, so does not require its own base unit; the SI thermodynamic mole is irrelevant to analytical chemistry and could cause avoidable costs to advanced economies; the mole is not a true metric (i.e. measuring) unit, rather it is a parametric unit and amount of substance is a parametric base quantity;
Young Ji International School / College
Page 89

the SI defines numbers of entities as quantities of dimension one, and thus ignores the ontological distinction between entities and units of continuous quantities.

Other units called "mole"
Chemical engineers use the concept extensively, but the unit is rather small for industrial use. For convenience in avoiding conversions, some American engineers adopted the pound-mole (noted lb-mol or lbmol), which is defined as the number of entities in 12 lb of 12C. One lb-mol is equal to 453.59237 mol In the metric system, chemical engineers once used the kilogram-mole (noted kg-mol), which is defined as the number of entities in 12 kg of 12C, and often referred to the mole as the grammole (noted g-mol), when dealing with laboratory data. Late 20th century chemical engineering practice came to use the kilomole (kmol), which is identical to the kilogram-mole, but whose name and symbol adopt the SI convention for standard multiples of metric units. Concentrations expressed as kmol/m3 are the same numbers as those in mol/dm3 or molarity conventionally used by chemists for bench measurements, which is convenient for scale-up.
Scale basis Scale basis relative to 12C = 12Relative deviation from the 12C = 12 scaleAtomic mass of hydrogen = 11.00794(7)−0.788%Atomic mass of oxygen = 1615.9994(3)+0.00375%Relative atomic mass of
16
O=
1615.9949146221(15)+0.0318%
Molar Mass A mole of any substance contains 6.022 x 1023 numbers of particles. How can you relate this value then to atomic mass unit? The atomic mass of 1 mole of any substance contains the number of particles as there are exactly 12 grams of C-12, which is the basis of atomic mass unit (amu)
The mass of a mole of any substance expressed in grams is known as molar mass. Since, numerical values of atomic masses of elements expressed in amu are numerically equivalent to their masses in grams (g) for every mole, then the Young Ji International School / College
Page 90
numerical values of molecular formula masses of compounds are also equal to their molar masses.
Molar Volume at STP One mole of N, gas, one mole of O2, or one mole of H2 gas will each occupy a volume of 22. 4 L when measured at 0°C and I atmosphere pressure, this condition is known as Standard Temperature and Pressure (STP). One mole of gas occupied a volume of 22. 4 L at STP. This is known as the molar volume at STP. Sample Problem What volume will 88 grams of CO2 gas occupy at STP?
Given mass CO2 = 88g Molar Volume at STP =22.4 L Let‘s try this! If 18 g of H2O (g) and 28 g of N2 gas were measured at STP: a. b. c. d. e.
How many liters will 18 g of H2O (g) occupy? How many liters will 28 g of N 2 (g) How many liters will 1 mole of H2O (g) occupy? How many liters will 1 mole of N2 (g) occupy? Justify your answer.
Mole Relationship 1. Convert mass to mole (mass-mole relationship)
Sample Problems How many moles of CO2 are in 132 grams of the gas? Atomic masses:
C = 12, 0 =16
Given mass CO2
= 132 g
Find: Number of moles a. Find the molecular mass of CO2 which is equal to mass per mole: C = 1 x 12 amu Young Ji International School / College
=12 amu Page 91
2 O = 2 x 16 amu Molecular mass of CO2 Molar mass CO2
= 32 amu = 44 amu = 44 g
2. Convert mole to mass (mole-mass relationship) Sample problem A chemical reaction requires 2,5 mol of potassium chlorate, (KCIO3). What mass of KCIO3 is needed?
3. Convert number of particles to Mole (Number of particles (Avogadro‘s No. J –mole relationship) Sample Problem A sample of CaCI2 contains 4 x10 30 formula units. How many moles of CaCI2 are in the sample? Given:
Formula units CaCI2 = 4 x1030
Find:
Number of moles in 4 x 1030 formula units of CaCl2
Areas of Investigation 1. The mineral albite contains sodium, aluminum, silicon, and oxygen. Look from the formula of albite and find its percentage composition. 2. Construct a concept map on the mole concept. 3. If a given the option, would you use MSG? find out how safe it is as a food additive. 4. Research on the percentage of fluoride present in toothpaste. 5. Find the percentage composition of glass.
Discussion Box 1. At STP, how many moles of atoms and molecules are in 70 grams of chlorine gas? 2. Convert 180 grams of water vapor to: Young Ji International School / College
Page 92
a. number of moles b. number of molecules c. molar volume at STP
Young Ji International School / College
Page 93