College-Level
Organic Chemistry
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TABLE OF CONTENTS Preface........................................................................................................ 1 Chapter 1: Chemical Bonding in Organic Chemistry .................................... 5 Carbon-Based Chemistry ................................................................................................. 5 Orbital Theories ............................................................................................................... 7 Carbon Hybridization .................................................................................................... 10 d Orbital Hybridization.................................................................................................. 12 Multiple Bonding ........................................................................................................... 12 Writing Organic Molecular Structures .......................................................................... 13 Bonding Trends in Organic Chemistry .......................................................................... 14 Constitutional Isomers .................................................................................................. 16 Organic Molecular Charges ............................................................................................17 Resonance Chemistry .................................................................................................... 19 Key Takeaways ............................................................................................................... 22 Quiz ................................................................................................................................ 23 Chapter 2: Basic Organic Molecular Structures ......................................... 27 IUPAC Nomenclature .................................................................................................... 27 Alkyl Halides .................................................................................................................. 30 Alkenes and Alkynes ...................................................................................................... 30 Alcohols .......................................................................................................................... 31 Ethers ............................................................................................................................. 32 Aldehydes ....................................................................................................................... 32 Ketones ........................................................................................................................... 34 Carboxylic Acids ............................................................................................................. 35
Esters .............................................................................................................................. 36 Amines ........................................................................................................................... 36 Functional Groups ......................................................................................................... 37 Stereochemistry and Isomers ........................................................................................ 38 Diastereomerism ............................................................................................................ 39 Enantiomer .................................................................................................................... 40 Key Takeaways ............................................................................................................... 42 Quiz ................................................................................................................................ 43 Chapter 3: Organic Solvent Chemistry ...................................................... 46 Types of Solvents ........................................................................................................... 46 Nonpolar Solvents.......................................................................................................... 47 Solvation......................................................................................................................... 52 Key Takeaways ............................................................................................................... 54 Quiz ................................................................................................................................ 55 Chapter 4: Alkanes, Alkenes, and Alkynes ................................................. 58 Alkanes ........................................................................................................................... 58 Alkyl Groups................................................................................................................... 59 Alkoxides or Alkoxy Groups .......................................................................................... 60 Chemical Properties of Alkanes ..................................................................................... 61 Cycloalkanes................................................................................................................... 62 Alkenes ........................................................................................................................... 63 Physical Properties of Alkenes ....................................................................................... 66 Alkynes ........................................................................................................................... 67 Alkyne Reactivity ........................................................................................................... 68
Key Takeaways ............................................................................................................... 70 Quiz .................................................................................................................................71 Chapter 5: Aldehydes, Ketones, and Carboxylic Acids ............................... 75 Naming Aldehydes ......................................................................................................... 77 Naming Ketones ............................................................................................................. 78 The Carbonyl Group.......................................................................................................80 Reactivity of Aldehydes and Ketones ............................................................................ 82 Natural Occurrence of Ketones and Aldehydes............................................................. 84 Carboxylic Acids ............................................................................................................. 85 Fatty Acids...................................................................................................................... 87 Properties of Carboxylic Acids .......................................................................................88 Key Takeaways ............................................................................................................... 89 Quiz ................................................................................................................................ 90 Chapter 6: Aromatic Compounds .............................................................. 93 Introduction ................................................................................................................... 93 Nomenclature of Aromatics ........................................................................................... 94 Benzene Chemistry ........................................................................................................ 98 Aromatic Reactions .......................................................................................................101 Halogenation of Benzene ............................................................................................. 102 Nitration of Benzene .................................................................................................... 103 Sulfonation of Benzene ................................................................................................ 104 Friedel-Crafts Reaction ................................................................................................ 105 Key Takeaways ............................................................................................................. 108 Quiz .............................................................................................................................. 109
Chapter 7: Alcohols and Alkyl Halides ...................................................... 113 Nomenclature of Alcohols ............................................................................................ 113 Reactivity of Alcohols.................................................................................................... 117 Alcohol Dehydration ..................................................................................................... 118 Oxidation of Alcohols ................................................................................................... 120 Reactivity of Alkyl Halides ........................................................................................... 120 Glycols .......................................................................................................................... 122 Key Takeaways ............................................................................................................. 123 Quiz .............................................................................................................................. 124 Chapter 8: Ethers, Epoxides, and Esters.................................................. 128 Ethers ........................................................................................................................... 128 Physical Properties of Ethers ....................................................................................... 129 Reactions with Ethers .................................................................................................. 130 Epoxides ....................................................................................................................... 132 Esters ............................................................................................................................ 134 Ester Reactions .............................................................................................................137 Polyester Formation .................................................................................................... 140 Saponification .............................................................................................................. 140 Key Takeaways .............................................................................................................. 141 Quiz .............................................................................................................................. 142 Chapter 9: Enols and Enolates ................................................................ 146 Introduction to Enols and Enolates ............................................................................ 146 Enolate reactions ......................................................................................................... 150 Acidic Alpha Halogenation of Ketones and Aldehydes ............................................... 150
Basic Alpha-Halogenation of Ketones and Aldehydes ................................................. 151 The Haloform Reaction ............................................................................................... 152 Alkylation of Enolates .................................................................................................. 153 The Aldol Reaction of Aldehydes ................................................................................. 154 The Aldol Reaction of Ketones .....................................................................................155 Conjugate Reactions .................................................................................................... 156 Michael Addition Reaction ...........................................................................................157 Key Takeaways ............................................................................................................. 158 Quiz .............................................................................................................................. 159 Chapter 10: Sulfur-Containing Organic Compounds................................ 163 Nomenclature of Sulfur Compounds ........................................................................... 163 Oxidation of Alcohols using DMSO ............................................................................. 166 Thiols ............................................................................................................................ 167 Sulfides ......................................................................................................................... 168 Synthesis of Sulfides .................................................................................................... 169 Key Takeaways .............................................................................................................. 171 Quiz ...............................................................................................................................172 Chapter 11: Nitrogen-containing Organic Molecules ................................. 176 Nomenclature of Amines ............................................................................................. 176 Physical Properties of Nitrogenous Compounds......................................................... 178 Alkylation of Ammonia ................................................................................................ 180 Reduction of Nitrogenous Compounds ........................................................................ 181 Reductive Amination via Imines ................................................................................. 183 Amine Reactions .......................................................................................................... 183
Preparing Amides ........................................................................................................ 184 Nitrosation of Amines .................................................................................................. 185 Key Takeaways ............................................................................................................. 187 Quiz .............................................................................................................................. 188 Chapter 12: Carbohydrates in Organic Chemistry .................................... 192 Nomenclature and Naming of Carbohydrates ............................................................ 192 Glycosides .................................................................................................................... 195 Reducing Sugars .......................................................................................................... 196 Substituted Sugars ....................................................................................................... 197 Alpha and Beta Isomers ............................................................................................... 197 Reactions of Carbohydrates ......................................................................................... 198 Alkylation of Carbohydrates ........................................................................................ 198 Acylation of Carbohydrates ......................................................................................... 199 Reduction of Carbohydrates ........................................................................................ 199 Oxidation of Carbohydrates ......................................................................................... 199 Hydrolysis of Sugars ................................................................................................... 200 Glycosidic Bond Formation ........................................................................................ 200 Key Takeaways ............................................................................................................. 201 Quiz ..............................................................................................................................202 Chapter 13: Amino Acids, Proteins, and Peptides in Organic Chemistry .. 206 Amino Acids .................................................................................................................206 Amino Acid Stereochemistry .......................................................................................209 Amino Acid Synthesis ..................................................................................................209 Reactions of Amino Acids ............................................................................................ 210
Terminology for Proteins and Peptides ........................................................................ 211 Amide Bond................................................................................................................... 211 Disulfide Bonds ............................................................................................................. 211 Sequencing Amino Acids ............................................................................................. 212 Protein Structure ......................................................................................................... 212 Alpha Helices and Beta Pleats ..................................................................................... 213 Peptide Synthesis ......................................................................................................... 213 Key Takeaways ............................................................................................................. 215 Quiz .............................................................................................................................. 216 Chapter 14: Lipids in Organic Chemistry ................................................. 220 Fatty Acids and Triglycerides ......................................................................................220 Saturated Fatty Acids ................................................................................................... 221 Unsaturated Fatty Acids (all have 18 carbons)............................................................ 221 Micelles ........................................................................................................................ 222 Phospholipids............................................................................................................... 223 Prostaglandins ............................................................................................................. 224 Terpenes ....................................................................................................................... 225 Steroids ........................................................................................................................ 226 Key Takeaways ............................................................................................................. 228 Quiz .............................................................................................................................. 229 Chapter 15: Nucleic Acids and Nucleosides in Organic Chemistry............ 233 Nucleic Acids ................................................................................................................ 233 Pyrimidines and Purines.............................................................................................. 233 Nucleosides .................................................................................................................. 234
Nucleotides................................................................................................................... 236 Nucleic Acids ................................................................................................................ 236 Nucleic Acids as Bioenergetic Molecules..................................................................... 237 DNA (Deoxyribonucleic Acid) ..................................................................................... 239 Key Takeaways ............................................................................................................. 242 Quiz .............................................................................................................................. 243 Summary ................................................................................................ 247 Course Questions and Answers ................................................................251
PREFACE In this course, we attempt to bring the basics of organic chemistry to the student needing to understand the nomenclature, chemical properties, and reactivity of carbonbased molecules. While there are far too many organic molecules in nature to discuss in any organic chemistry course, there are specific ways to clearly identify these molecules as well as certain trends in how these various molecules behave in chemical reactions. By the end of the course, you will understand how to identify and name organic molecules, their physical properties, and how they chemically interact with one another in a variety of types of chemical reactions. Chapter one in the course begins the study of organic chemistry by introducing how organic molecules are put together. There really isn’t any difference between the way the atoms in organic molecules are put together and the way other chemical molecules are put together but it is worth reviewing, even if you have studied chemistry in the past. This chapter will focus on orbital theory and the particulars of organic molecular bonding as well as the shorthand involved in writing out organic molecular structures. Finally, the chapter talks about resonance chemistry as it applies to organic molecules. Chapter two in the course covers the basics of nomenclature in identifying organic molecules. Because there are innumerable organic molecules and because they are based on just a few different types of atoms, there needs to be a way to identify what each molecule looks like by name alone. This leads to a discussion of the IUPAC nomenclature and coverage of the different functional groups in organic chemistry. You will need to understand how to name the different molecules you see in organic molecules, which is covered in this chapter. In addition, there will be a discussion of stereochemistry as it applies to organic molecules. The topic of Chapter three in the course is organic solvent chemistry. For students who have participated in regular chemistry experiments, the solvent has typically been water. In organic chemistry, the solvent may or may not be water because many aspects of organic chemistry involve nonpolar substances that do not dissolve in water. Solvents
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may or may not participate in chemical reactions themselves but are important to the chemistry of different molecules. Issues regarding solvation and solutions in organic chemistry are also covered in this chapter. Chapter four begins a series of chapters on the different organic compounds and their properties. You will have learned about alkane, alkene, and alkyne nomenclature in previous chapters so the focus of this chapter is to learn more about these compounds and how they interact with one another and with other organic compounds. These substances can have a variety of different configurations, which need to be discussed as part of this chapter. The focus of Chapter five in the course is the chemistry of aldehydes, ketones, and carboxylic acids. This is the first time the chemistry of oxygen comes into play in this course. Aldehydes and ketones are discussed together because they have very similar chemistry and reaction types. Carboxylic acids are also oxygen-related because they have a COOH side chain as their defining characteristic. They also have great reactivity and are seen in nature as fatty acids and other biochemically-important molecules. In all cases, you will come to understand their nomenclature, their physical properties, and some of the most important chemical reactions associated with these molecules. Chapter six introduces the structure and chemistry of aromatic compounds. All aromatic compounds consist of a cyclic compound that carries resonance. The most common aromatic compound is benzene, which is very stable and has chemistry unique to the molecule. In this chapter, the nomenclature and chemistry of aromatic compounds will be covered as well as the different reactions that are seen in organic chemistry with these types of molecules. The topic of Chapter seven in the course is the chemistry of alcohols and alkyl halides. Alcohols are organic compounds that have a hydroxyl group as its major functional group, often represented with the general formula of ROH, where R can be any number of organic chemistry alkyl groups. The hydroxyl group is highly reactive so that there are any number of reactions that can occur at this functional group. The chapter also covers the chemistry of alkyl halides, which are alkyl groups that have one or more
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halogen side chains attached to it. The halogen is also highly reactive, with many possible chemical reactions associated with it. Chapter eight in the course focuses on the organic chemistry associated with ethers, epoxides, and esters. Ethers and epoxides are related to one another in that certain types of cyclic ethers are referred to as epoxides. In both types of molecules, the general formula is ROR’, involving a variety of R side chains. These are molecules commonly seen in perfumes, industrial compounds, waxes, oils, and dyes. Esters are also commonly used in industry, being a part of the making of many products—the most common of which are the polyesters. The topic of Chapter nine is the structure and chemistry of enols and enolates. Enols are also referred to as alkene alcohols, which are alkenes that have an alcohol group added to one of the carbon atoms. These are first alkenes but, chemically-speaking, they should be considered important for their electron-donating capacity. Enols can be mixed with alkali substances to make enolates, which are the conjugate bases of enols. Both of these types of molecules are best known for the many different types of reactions they participate in, which are covered in this chapter. Chapter ten in the course changes nomenclature and reactions in organic chemistry to include molecules that contain sulfur. Sulfur compounds are somewhat similar to oxygen-containing molecules in that they belong to the same group but sulfur is a great deal larger than oxygen, leading to slightly different chemical reactivity unique to these molecules. The nature and chemistry of thiols and sulfides is discussed as part of this chapter. Chapter eleven places a focus on the different nitrogen-containing molecules in organic chemistry. These types of molecules are not only important in basic organic chemistry; they are also important in numerous biochemical processes. The main type of molecule discussed will be the amine compounds, which are considered organic derivatives of ammonia. Like ammonia itself, amine compounds will have a certain degree of basicity, which leads to nucleophilicity of the nitrogen compounds in organic compounds.
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Chapter twelve in the course begins to make sense of what you will learn in prior chapters on simpler molecules and applies it to more complex biochemical molecular structures. Sugars and carbohydrates are basically organic molecules that come from the phrase “carbon hydrates”. They contain only carbon, oxygen, and hydrogen atoms and have a specific generic formula, based on whether they are simple sugars, disaccharides, or more complex polysaccharides. The main focus of this chapter is to use organic molecular principles that will make sense after you know the basic reaction types involved. The focus of Chapter thirteen in the course is to bring on more of the biochemistry involving organic chemistry principles by talking about the organic chemistry of amino acids, oligopeptides, and proteins. These are molecules that have nitrogenous compounds as the basis of their chemistry and that, like carbohydrates, exist as monomer units and polymers or polypeptides. These will also have reactions at their functional units, which involve a variety of different side chains and parts of the parent chain. Chapter fourteen in the course studies the organic chemistry associated with lipids. The term “lipid” is a broadly reaching term that applies to a wide variety of molecules that are called lipids because of their biochemical nature and their lack of solubility in water. They can range from fatty acids to triglycerides to more complex molecules that are complicated to synthesize and even to understand how they are synthesized in body systems and in organic chemistry models. Lipids have poor solubility in water but are much more soluble in chloroform, benzene, ether, and acetone, which are either nonpolar or weakly polar. The focus of Chapter fifteen and the final chapter of the course is the organic chemistry and biochemistry of nucleic acids, which are deoxyribonucleic acids and ribonucleic acids, commonly called DNA and RNA. These are molecules that contain ribose or deoxyribose sugars, phosphate groups, and nucleic acid bases, which are seen in paired form with DNA and sometimes with RNA. The transcription of nucleic acids and the translation to proteins is covered as part of the chapter as these are not generally made synthetically in an organic chemistry laboratory.
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CHAPTER 1: CHEMICAL BONDING IN ORGANIC CHEMISTRY This chapter begins the study of organic chemistry by introducing how organic molecules are put together. There really isn’t any difference between the way the atoms in organic molecules are put together and the way other chemical molecules are put together but it is worth reviewing, even if you have studied chemistry in the past. This chapter will focus on orbital theory and the particulars of organic molecular bonding as well as the shorthand involved in writing out organic molecular structures. Finally, we will introduce resonance chemistry as it applies to organic molecules.
CARBON-BASED CHEMISTRY As we have discussed, organic chemistry is carbon-based chemistry. There are many features of carbon that make organic chemistry relatively unique when it comes to molecular structures. Carbon-carbon bonds are completely strong. Consider that when it comes to carbon in diamond form, there is a tightly-bound diamond lattice that provides the substance its strength. Graphite is also carbon but in a different shape. Figure 1 shows the difference between diamond and graphite.
Figure 1
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What makes carbon unique among the many other atoms that can bond with each other is that it has the capacity to bond strongly with itself as well as with other atoms, making a multitude of possible molecules—with all of the bonds being potentially strong. The backbone of organic chemistry is hydrocarbons, which are strongly-bonded carbon atoms—both with each other and with hydrogen atoms. Why is carbon so versatile in its ability to bond with such a wide variety of elements (including metal salts, halogens, oxygen, and hydrogen)? Part of it is because it is such a small element and part of it is because it has four valence electrons around it. Its choices are that it could lose four electrons or gain four electrons (which would cause too much repulsion of so many electrons) if it were to make a salt. Instead, it shares its valence electrons with other atoms. This is what it “prefers” to do. Figure 2 shows what it does with methane and ethane:
Figure 2.
By sharing electrons with other atoms, there is a decrease in repulsions between electrons of the valence shell that are compensated for by the positive charge of the atom it is associated with (in the case of ethane and methane, this will be hydrogen or another carbon atom). It should be noted that the electrons are not necessarily shared equally, even in situations where the bond is technically “covalent”. More of the negative charge, for example will be on the fluorine atom when it bonds to carbon because the fluorine atom is more electronegative. What this means is that there is a gradual shift from purely ionic to purely covalent bonding with carbon bonding with other atoms being generally somewhere in between. Molecules like lithium hydride are purely ionic, while molecules like H2 are purely covalent (because there is equal sharing of electrons). Lithium hydride becomes a salt6
like ionic compound with a high melting point, while methane, being non-salt-like, has a boiling point of -161 degrees Celsius. Methane or CH4 is mainly a covalently-bonded molecule. There is little electrostatic attraction between methane molecules (they are nonpolar) so that they boil at a very low temperature to become a gas. Hydrogen fluoride boils, on the other hand, 200 degrees higher than methane gas. This is because it is more ionic and has connections between the Hydrogen and Fluorine atoms through not only ionic bonding but also with hydrogen bonding between the HF molecules in chains and rings. This makes the boiling point higher. Because methane is nonpolar, it is inert to almost all reagents that could remove the hydrogen ion under anything but the most extreme conditions. It is therefore difficult to generate a CH3+ or CH3- ion, as these would be very reactive and just wouldn’t last long. In other words, you couldn’t make CH4 plus HF to make CH3F plus H2. This would require a cation of CH3+, which is not very stable in nature.
ORBITAL THEORIES While there are many ways to describe and write covalent bonds (which will be described later in this chapter), there are some concrete theories about how covalent bonds exist in organic molecules. While there is ionic bonding in organic chemistry, the most important bonding in this type of chemistry involves covalent bonds. You may have learned that covalent bonding involves the sharing of electrons but you may not know exactly how this happens. This introduces the topic of the valence bond theory, which describes how bonding happens in covalent molecules. According to this theory, there are two atomic orbitals around a pair of atoms—each orbital of which contains one electron. In sharing orbitals, the orbital pair will contain a stable set of two electrons. The H2 molecule is the simplest case of a covalently-bonded molecule. There are two spherical orbitals (1s orbitals) in each hydrogen atom, each with one electron in it. In the bonding of the two atoms, the electrons no longer are anywhere within the sphere but spend more time in that part of each sphere between the two nuclei, which holds the
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atoms together. According to quantum chemistry, the two electrons must occupy a shared orbital space in order to form a bond. The hydrogen atoms are too far apart, the 1s orbitals cannot overlap in order to form a covalent bond. As they overlap, a bond will begin to form, lowering the potential energy of the system. There will be attraction of the opposite electron to the opposite nucleus. They cannot, however, get too close or there will be repulsion of the nuclei with each other, which increases the potential energy of the system. A balance is had when there is an optimal distance between the nuclei. In a sense, there is a defined “optimal” distance between two bonding nuclei, which is when the potential energy is the lowest, meaning that the combined repulsive and attractive forces add up to the greatest amount of attraction. This optimal internuclear distance is referred to as the “bond length”. The difference between the optimal “bonded” energy and the separated energy is the “bond energy”. Every covalent bond has a certain strength and bond length. As an example, the carboncarbon bond is 1.5 Angstroms long with an Angstrom being 10-10 meters in length. Bonds aren’t the straight sticks as they are depicted in drawing chemical bonds. Instead, they are more like springs, which have the ability to bend, extend, and compress. According to the valence bond theory, there are two assumptions about bonds: 1) the strength of a given bond is directly proportional to the degree of overlap of the bond (in other words, the greater the overlap, the more stable is the bond). 2) An atom is able to use different combinations of atomic orbitals in order to maximize the overlap of the orbitals used by bonded atoms. Maximum overlap occurs between orbitals of the same spatial orientation and similar energies. Bonding can take place between beryllium and hydrogen to make BeH2 which, according to the periodic table of the elements in figure 3, has an atomic number of 4:
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Figure 3
According to the atomic number and the atomic orbital approach, beryllium has a 1s22s2 electron configuration. It has apparently filled its 2s orbital subshell, leaving behind no apparent reason why it should want to overlap with the singly occupied 1s orbital associated with the hydrogen atom. How do these two atoms overlap to make BeH2? One way to do this is to excite a 2s electron on beryllium to allow for a partially empty 2p orbital that absolutely could bind with hydrogen. Doing this is called “promotion”. What this looks like is seen in figure 4:
Figure 4.
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In this excited state (having an unmatched 2s1 and 2p1 orbital), Beryllium could certainly bind with hydrogen in two unmatched orbitals according to this theory. Research, however, has shown that these two bonds are identical. How can this be? Basically, the only way this makes sense is through the process of hybridization. What this involves is combining orbitals that are not equivalent but that are properly oriented to form bonds. These new combinations are called hybrid orbitals because they are made by hybridizing two or more atomic orbitals from the same atom. This leads to a unique hybrid orbital in beryllium that has this energy pattern, as shown in figure 5:
Figure 5.
This obviously leads to the ability of the beryllium to bind with the hydrogen atom in order to make BeH2. This will produce a linear BeH2 molecule. Both promotion and hybridization require an input of energy; however, this is made up for when beryllium bonds with 2 hydrogen atoms. It means that, in situations not associated with beryllium, some compounds are so unstable because they have such a high energy necessary to make the hybrid orbital. This energy is not made up for by the energy recouped in the bonding effort.
CARBON HYBRIDIZATION Carbon has six electrons, with a 2s22p2 configuration. Based on this, one would expect that it would be likely to bond with just two other atoms; we know, on the other hand, that it bonds with four other atoms to make covalent bonds. On the second level, there is one 2s orbital and 3 2p orbitals, which can be hybridized to make four degenerate sp3 10
hybrid orbitals that contain a single electron available for bonding. Figure 6 shows what this hybridization looks like:
Figure 6.
The energy of the hybridized orbitals is somewhere in the middle of the ground state and the excited state of carbon. This leads to orbitals that are in the shape of a tetrahedron with 109.5-degree angles between them. These are equivalent in energy and allow for a tetrahedral molecule, which is the exact shape of methane. The amount of energy released in any hybrid orbital situation increases with the number of bonds formed. In the case of carbon, more energy is released in the formation of four bonds versus just two, making CH4 so much more stable than CH2 or CF2. Intermediates like CH2 are so reactive that they only form temporarily under certain experimental circumstances. Valence bond theory also explains why molecules such as NH3 and H2O are formed. With nitrogen, the 2s22p3 electron configuration becomes hybridized to make four sp3 hybrid orbitals. One of the orbitals is completely occupied with a pair of electrons, while the remaining three are single and are available for bonding.
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D ORBITAL HYBRIDIZATION Hybridization is not restricted to the s and p atomic orbitals. In period three, the 2d plus 3s and 3p orbitals can be used to create a hybrid orbital, such as is seen in molecules such as SF6 and PF5. In such cases, the five hybrid orbitals created are not the same. Three will form a triangular array of orbitals that are 120 degrees from each other. The other two are at 90 degrees to the first three and are at 180 degrees to each other.
MULTIPLE BONDING So, what about multiple bonding situations, such as is seen in ethylene (C2H4)? What happens is that, out of a 2s22p2 situation, we can promote one of the 2s electrons to a 2p orbital and then create three hybridized sp2 orbitals that each have an electron in it and the promoted 2p orbital. This, as you can see by figure 7, requires energy.
Figure 7.
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Figure 8 shows what the ethylene molecule looks like from an orbital perspective:
Figure 8.
The two CH bonds come from the sp2 + s orbitals and the double bond comes from the sp2 + Sp2 orbital set. This leaves two 2p orbitals that have a single electron each. They overlap to form a pi-bond which, together with the C-C bond will make a double bond. This pi-bond is also seen in figure 8. As you know, carbon can form triple bonds with itself, using a similar pattern to make acetylene (C2H2). In such cases, the ground state gets promoted as in the case of ethylene but hybridizes to make 2 sp hybrid orbitals and has two remaining 2p orbitals. These 2p orbitals will connect with the 2p orbitals on the opposite carbon atom two make 2 pi-bonds along with the C-C bond it already makes. This leads to a triple bond and a bond left over for a hydrogen atom on each side. The tight bonding makes for a shorter carbon-carbon bond length and a higher bond energy than is seen for single and double-bonded carbon structures like ethane and ethylene.
WRITING ORGANIC MOLECULAR STRUCTURES Hopefully, you now understand a little bit about how organic molecules bond and you understand the periodic table from previous chemistry courses. The good news is that 13
knowledge of the entire periodic table is not necessary to easily recall those things seen commonly with organic molecules. The harder part of organic chemistry is determining how to write out these molecules.
BONDING TRENDS IN ORGANIC CHEMISTRY Using the typical methods used for identifying organic molecules does not work very well, especially when the molecules are large and when there are basically strings of carbon atoms in varying arrangements. For this reason, there are things you need to know in order to write the different organic molecules. All organic molecules are carbon-based and, as you know, it has the ability to form four separate bonds. There are different ways to write this using the chemical abbreviation for carbon. Figure 9 shows what it looks like to write carbon in smaller molecules:
Figure 9.
Hydrogen is also relatively easy. It has one electron and can only have one bond and will have no formal charge. The exception is the single proton H+, and the hydride ion, H-. The hydride ion is a proton plus two electrons. These are highly reactive molecules that do not exist in nature but are important in discussing the chemistry of organic molecules. Next comes oxygen. This exists in one of three ways. It can exist in a neutral way with two bonds and two lone pairs of electrons. It can also become single-bonded with a negative-one negative charge. Lastly, it can come triple-bonded with a formal positive-
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one charge. A formal charge of negative-one is called a hydroxide ion; a formal charge of positive-one is called a hydronium ion. Nitrogen has two major patterns of bonding. It can exist as a neutral nitrogen atom with three separate bonds, a double bond and a single bond, or a triple bond. Nitrogen can also be positively-charged with four bonds to other atoms. These can be four single bonds, a double bond and two single bonds, or a triple bond and a single bond. There is no double “double-bonding situation” in nitrogen and no formal negative charge on nitrogen, except in highly unstable situations. Two other atoms should be noted as they apply to organic molecules. The first is sulfur and the second is phosphorus. Sulfur generally follows the same pattern as oxygen, while phosphorus is almost always seen in the form of the phosphate ion or PO4 with a three-negative charge. Phosphorus has five bonds, no lone pairs and a zero formal charge. Because it is a larger molecule, it has an expanded valence number with dorbitals in which the octet rule does not apply. Figure 10 shows the phosphate molecule and its charges:
Figure 10.
The halogens, like chlorine, fluorine, iodine, and bromine, are important in organic chemistry but aren’t seen in organic molecules as much as the molecules already discussed. Halogens will have a formal charge of zero and will have a single bond. Rarely, such as in the case of bromine, there will be a three-membered ring with carbon and a formal charge of plus one.
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Because it is relatively clear that carbon has specific bonding patterns and that other atoms have specific bonding patterns, you can condense the way that organic molecules are written. They can be written in a condensed form, such as CH3CH2CH2OH, as well as in simple line form, in which there is no “C” written for carbon and no “H” written for hydrogen. These are called “line structures” as they look like zig-zagged lines. Only the hydrogens attached to molecules other than carbon are listed and all other atoms are listed in other than line form. Ring structures also do not have a special listing of the atoms separately but is shown as in figure 11:
Figure 11.
These line and ring structures are pared down in a way that it makes it easier to see the bonding and the way the structures are laid out. It tends to work best for larger and more cumbersome molecules; in the case of smaller structures, it is best to write them out without doing a line structure.
CONSTITUTIONAL ISOMERS Constitutional isomers are the same as “structural isomers”. These are molecules that can be written the same way when writing the different carbon, hydrogen, and oxygen molecules but that look quite different. A typical example is fructose and glucose, which are sugar molecules. Both can be written as C4H12O6. How do you write this? Because these are two different molecules, you can imagine that they are different molecules. Figure 12 shows what these two molecules look like when drawn out:
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Figure 12.
ORGANIC MOLECULAR CHARGES You should also have an understanding of what it means to be a charged molecule. This is relatively easy to understand when it comes to ionized molecules like sodium, which becomes Na+, and chlorine, which becomes Cl- when ionized. What you need to know now is that organic molecules can also be charged. An example of this is methanol, which is CH3OH, which can be neutral. It can also lose an H+ ion to make CH3O- (an anion) or can gain a hydrogen ion to become CH3OH2+. In these cases, the positive charge or negative charge is on the oxygen atom. This can be further categorized as the “formal charge”, which is the charge specifically on the oxygen atom rather than on the molecule as a whole or on any other part of the molecule. To find the formal charge on different atoms of an organic molecule, it is necessary to add up the valence electrons. In the case of oxygen, an unbound oxygen atom has six valence electrons. When it is bound in the methanol molecule, there are two electrons used to make the bonds to carbon and hydrogen and four left over. These four electrons left over are all “owned” by the oxygen molecule. The bound electrons with hydrogen are “half-owned”. Taking the total valence electrons and subtracting the unbound electrons around the atom and half of the bound electrons, gives a net charge of 0 on a methanol molecule. Figure 13 demonstrates the formal charge of -1 on the oxygen atom in the methanol anion:
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Figure 13.
In the case of methane cation, the extra electron given by hydrogen makes six bonding electrons and two nonbonding electrons. Half of six bonding electrons is three and the two electrons not involved in bonding makes 5 in total. This leads to a +1 charge. The formula for determining the formal charge on oxygen is to take the total valence electrons minus the unpaired electrons and minus half of the bound electrons. Taking six minus five, gives a plus-one charge on the oxygen. Added to the issue is the fact that there can be both negative and positive formal charges on different atoms of a larger molecule. This is the issue with amino acids, in particular. There are zwitterions that have both positive and negative charges on the molecule. The total charge on the molecule is going to be zero. Even though there is a net zero charge on the molecule, it is necessary to show the location of the positive and negative charges on each atom. Formal charges can help to determine which molecules are more stable than others. Let’s look at two possible ways to write CO2 to see which is most stable: In CO2, the C is less electronegative, so that it is the central atom. It has four valence electrons and oxygen has six valence electrons, for a total of 16 valence electrons. The binding of carbon and oxygen gives an O-C-O molecule with double bonds between the two. This would not work with a single bond because it leaves too many unpaired electrons. It can also not be written with a single bond on one side and a triple bond on the other side. It must be symmetric. Figure 14 shows the Lewis dot structure on CO2.
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Figure 14.
As you can see, the first CO2 structure is more stable because it has a zero formal charge on all atoms, while the second leaves an unacceptable charge on the two oxygen atoms. In general, the Lewis structure with the set of formal charges closest to zero will be the most stable.
RESONANCE CHEMISTRY A discussion of resonance chemistry can be had by looking at the CO3(2-). This has carbon in the middle, a double bond with one oxygen and a single bond with 2 oxygens and an overall charge of 2-. The big question is this: Which oxygen molecule gets the double bond and which get a single bond? This can be explained in terms of resonance. When more than one Lewis dot structure can be drawn for the same molecule, this is said to have “resonance”. What this means is that, for CO3(2-), the charge is rapidly shifting from one place to another in a sort of blur that gives each CO bond equal stability, so that a third of the time, the double bond will be on one of the three oxygen atoms, so that there is equal sharing of this double bond. Instead of a -1 charge on two of the atoms, there will be a 2/3 charge on each oxygen atom on average. Figure 15 shows the resonance of CO3(2-).
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Figure 15.
The doubles-sided arrow will show that these structures are equivalent. There is not a true -2/3 charge on each atom but, experimentally at least, there is no difference between structures A, B, and C and the experiment to find these structures indicates that these structures as written do not exist. You can draw resonance structures by drawing all of the possible structures and charges on a molecule; however, this can be cumbersome. The only difference between the Lewis structures of these molecules is the placement of the electrons. The atomic position is exactly the same. One can use a curved arrow in order to indicate that an electron shifts between one atom or another. In the example shown in figure 16, there is a curved arrow that indicates the shifting of an electron pair in the ethylene molecule. This is called “electron pushing”.
Figure 16.
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In drawing resonance structures, all structures must have the same number of valence electrons with no destruction or creation of electrons. The octet rule needs to be obeyed at all times. In other words, you cannot have five bonds around carbon or more than one bond associated with hydrogen. Nuclei cannot change positions in a resonance structure (only the actual electrons). Ozone is a typical resonance structure. It involves three oxygen molecules together that have a double bond between one oxygen and a single bond between the other oxygen. This, however, does not stay stable and the double bond shifts from one oxygen molecule to another. Figure 17 shows the ozone molecule in its resonance arrangement:
Figure 17.
You can try this yourself as you discover that, like ozone, molecules like NO2- will have a resonance structure as well. In short, resonance structures are a simplified way of describing molecular orbitals that extend to cover more than two atoms. The electrons will shift from one place to the next and will be “averaged out” over the course of two molecular bonds.
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KEY TAKEAWAYS •
Organic chemistry involves carbon, hydrogen, and oxygen, with several other molecules less likely seen.
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Carbon has four valence electrons, hydrogen has one valence electron, and oxygen has six valence electrons.
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Covalent bonding involves the promotion of electrons and the hybridization of orbitals so that they take on a unique molecular bonding orbital.
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The formal charge depends on the number of valence electrons, the number of bonded electrons, and the number of unbonded electrons.
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The total charge on a molecule in nature will be zero.
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The simplest way to write an organic molecule is to use line drawings for carbon atoms and all hydrogen atoms not attached to carbon atoms.
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