STUDY GUIDE
, OpenStax Organic Chemistry: A Tenth Edition Study Guide
Preface
What enters your mind when you hear the words “organic chemistry?” Some of you may
think, “the chemistry of life,” or “the chemistry of carbon.” Other responses might include “pre-
med,” “pressure,” “difficult,” or “memorization.” Although formally the study of the
compounds of carbon, the discipline of organic chemistry encompasses many skills that are
common to other areas of study. Organic chemistry is as much a liberal art as a science, and
mastery of the concepts and techniques of organic chemistry can lead to improved competence
in other fields.
Here are several suggestions that may help you with problem solving:
1. The text is organized into chapters that describe individual functional groups. As you
study each functional group, make sure that you understand the structure and reactivity
of that group. In case your memory of a specific reaction fails you, you can rely on your
general knowledge of functional groups for help.
2. Use molecular models. It is difficult to visualize the three-dimensional structure of an
organic molecule when looking at a two-dimensional drawing. Models will help you to
appreciate the structural aspects of organic chemistry and are indispensable tools for
understanding stereochemistry.
3. Look through the appendices at the end of the Study Guide. Some of these
appendices contain tables that may help you in working problems; others present
information related to the history of organic chemistry.
While the Study Guide is written to accompany Organic Chemistry, it also contains several
unique features. Each chapter of the Study Guide begins with an outline of the text that can be
used for a concise review of the text material and can also serve as a reference. After every few
chapters a Review Unit is included. In most cases, the chapters covered in the Review Units are
related to each other, and the units are planned to appear at approximately the place in the
textbook where a test might be given. Each unit lists the vocabulary for the chapters covered,
the skills needed to solve problems, and several important points that might need reinforcing or
that restate material in the text from a slightly different point of view. Finally, the small self-test
that has been included allows you to test yourself on the material from more than one chapter.
I have tried to include many types of study aids in this Study Guide. Nevertheless, this book can
only serve as an adjunct to the larger and more complete textbook. In addition, please note that
a companion to this Study Guide, the Student Solutions Manual, contains solutions to the
problems found in Organic Chemistry: A Tenth Edition. I am pleased that this is the first time
that these are published by OpenStax and available to you for free in honor of my late son, Peter
McMurry.
Susan McMurry
Acknowledgments
Bryant Gilbert, Zane State College
Kjir Hendrickson, Arizona State University
Ike Shibley, Penn State Berks
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Table of Contents
Chapter 1: Structure and Bonding 5
Chapter 2: Polar Covalent Bonds; Acids and Bases 9
Review Unit 1 12
Chapter 3: Organic Compounds: Alkanes and Their Stereochemistry 15
Chapter 4: Organic Compounds: Cycloalkanes and Their Stereochemistry 19
Chapter 5: Stereochemistry at Tetrahedral Centers 22
Review Unit 2 25
Chapter 6: An Overview of Organic Reactions 29
Chapter 7: Alkenes: Structure and Reactivity 32
Chapter 8: Alkenes: Reactions and Synthesis 35
Review Unit 3 38
Chapter 9: Alkynes: An Introduction to Organic Synthesis 43
Chapter 10: Organohalides 45
Chapter 11: Reactions of Alkyl Halides: Nucleophilic Substitutions and Eliminations 48
Review Unit 4 53
Chapter 12: Structure Determination: Mass Spectrometry and Infrared Spectroscopy 56
Chapter 13: Structure Determination: Nuclear Magnetic Resonance Spectroscopy 60
Review Unit 5 65
Chapter 14: Conjugated Compounds and Ultraviolet Spectroscopy 70
Chapter 15: Benzene and Aromaticity 73
Chapter 16: Chemistry of Benzene: Electrophilic Aromatic Substitution 77
Review Unit 6 81
Chapter 17: Alcohols and Phenols 86
Chapter 18: Ethers and Epoxides; Thiols and Sulfides 90
Preview of Carbonyl Chemistry 93
Review Unit 7 94
Chapter 19: Aldehydes and Ketones: Nucleophilic Addition Reactions 97
Chapter 20: Carboxylic Acids and Nitriles 102
Chapter 21: Carboxylic Acid Derivatives: Nucleophilic Acyl Substitution Reactions 105
Review Unit 8 110
Chapter 22: Carbonyl Alpha-Substitution Reactions 115
Chapter 23: Carbonyl Condensation Reactions 118
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Chapter 24: Amines and Heterocycles 121
Review Unit 9 126
Chapter 25: Biomolecules: Carbohydrates 131
Chapter 26: Biomolecules: Amino Acids, Peptides, and Proteins 136
Review Unit 10 142
Chapter 27: Biomolecules: Lipids 147
Chapter 28: Biomolecules: Nucleic Acids 150
Chapter 29: The Organic Chemistry of Metabolic Pathways 154
Review Unit 11 160
Chapter 30: Orbitals and Organic Chemistry: Pericyclic Reactions 165
Chapter 31: Synthetic Polymers 168
Review Unit 12 171
Appendices
Appendix A: Functional-Group Synthesis 175
Appendix B: Functional-Group Reactions 181
Appendix C: Reagents in Organic Chemistry 186
Appendix D: Name Reactions in Organic Chemistry 194
Appendix E: Abbreviations 202
Appendix F: Infrared Absorption Frequencies 206
Appendix G: Proton NMR Chemical Shifts 209
Appendix H: Nobel Prize Winners in Chemistry 210
Appendix I: Answers to Multiple Choice Questions in Review Units 1–12 222
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Chapter 1 – Structure and Bonding n n n n n
Chapter Outline n
I. Atomic Structure (Sections 1.1–1.3).
n n n
A. Introduction to atomic structure (Section 1.1).
n n n n n
1. An atom consists
n n
2. of a dense, positively charged nucleus surrounded by negatively charged
n n n n n n n n n
electrons.
n
a. The nucleus is made up of positively charged protons and uncharged neutrons.
n n n n n n n n n n n
b. The nucleus contains most of the mass of the atom.
n n n n n n n n n
c. Electrons move about the nucleus at a distance of about 2 ×10−10 m (200 pm).
n n n n n n n n n n n n n n
3. The atomic number (Z) gives the number of protons in the nucleus.
n n n n n n n n n n n
4. The mass number (A) gives the total number of protons and neutrons.
n n n n n n n n n n n
5. All atoms of a given element have the same value of Z.
n n n n n n n n n n n
a. Atoms of a given element can have different values of A.
n n n n n n n n n n
b. Atoms of the same element with different values of A are called isotopes.
n n n n n n n n n n n n
B. Orbitals (Section 1.2).
n n
1. The distribution of electrons in an atom can be described by a wave equation.
n n n n n n n n n n n n n
a. The solution to a wave equation is an orbital, represented by Ψ.
n n n n n n n n n n n
b. Ψ 2 predicts the volume of space in which an electron is likely to be found.
n n n n n n n n n n n n n n n
2. There are four different kinds of orbitals (s, p, d, f).
n n n n n n n n n n
a. The s orbitals are spherical.n n n n
b. The p orbitals are dumbbell-shaped.
n n n n
c. Four of the five d orbitals are cloverleaf-shaped.
n n n n n n n
3. An atom’s electrons are organized into electron shells.
n n n n n n n
a. The shells differ in the numbers and kinds of orbitals they contain.
n n n n n n n n n n n
b. Electrons in different orbitals have different energies. n n n n n n
c. Each orbital can hold up to a maximum of two electrons.
n n n n n n n n n n
4. The two lowest-energy electrons are in the 1s orbital.
n n n n n n n n
a. The 2s orbital is the next higher in energy.
n n n n n n n n
b. The next three orbitals are 2px, 2py and 2pz, which have the same energy.
n n n n n n n n n n n n n
i. Each p orbital has a region of zero density, called a node.
n n n n n n n n n n n
c. The lobes of a p orbital have opposite algebraic signs.
n n n n n n n n n
C. Electron Configuration (Section 1.3).
n n n
1. The ground-state electron configuration of an atom is a listing of the orbitals
n n n n n n n n n n n n
occupied by the electrons of the atom in the lowest energy configuration.
n n n n n n n n n n n n
2. Rules for predicting the ground-state electron configuration of an atom:
n n n n n n n n n
a. Orbitals with the lowest energy levels are filled first. n n n n n n n n
i. The order of filling is 1s, 2s, 2p, 3s, 3p, 4s, 3d.
n n n n n n n n n n n
b. Only two electrons can occupy each orbital, and they must be of opposite spin.
n n n n n n n n n n n n n
c. If two or more orbitals have the same energy, one electron occupies each until all
n n n n n n n n n n n n n n
are half-full (Hund’s rule). Only then does a second electron occupy one
n n n n n n n n n n n n
of the orbitals. n n
i. All of the electrons in half-filled shells have the same spin.
n n n n n n n n n n
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II. Chemical Bonding Theory (Sections 1.4–1.5).
n n n n
A. Development of chemical bonding theory (Section 1.4). n n n n n n
1. Kekulé and Couper proposed that carbon has four “affinity units”; carbon is tetravalent.
n n n n n n n n n n n n
2. Kekulé suggested that carbon can form rings and chains.
n n n n n n n n
3. Van’t Hoff and Le Bel proposed that the 4 atoms to which carbon forms bonds sit at
n n n n n n n n n n n n n n n n
the corners of a regular tetrahedron.
n n n n n n
4. In a drawing of a tetrahedral carbon, a wedged line represents a bond pointing toward
n n n n n n n n n n n n n n
the viewer, a dashed line points behind the plane of the page, and a solid line lies in
n n n n n n n n n n n n n n n n n n
n the plane of the page.
n n n n
B. Covalent bonds. n
1. Atoms bond together because the resulting compound is more stable than the
n n n n n n n n n n n
nindividual atoms. n
a. Atoms tend to achieve the electron configuration of the nearest noble gas.
n n n n n n n n n n n
b. Atoms in groups 1A, 2A and 7A either lose electrons or gain electrons to form
n n n n n n n n n n n n n n
ionic compounds.
n n
c. Atoms in the middle of the periodic table share electrons by forming covalent
n n n n n n n n n n n n
bonds.
n
d. The neutral collection of atoms held together by covalent bonds is a molecule.
n n n n n n n n n n n n
2. Covalent bonds can be represented two ways. n n n n n n
a. In electron-dot structures, bonds are represented as pairs of dots.
n n n n n n n n n
b. In line-bond structures, bonds are represented as lines drawn between two bonded
n n n n n n n n n n n
atoms.
n
3. The number of covalent bonds formed by an atom depends on the number of
n n n n n n n n n n n n n
nelectrons it has and on the number it needs to achieve an octet.
n n n n n n n n n n n n
4. Valence electrons not used for bonding are called lone-pair (nonbonding)
n n n n n n n n n
nelectrons.
a. Lone-pair electrons are often represented as dots. n n n n n n
C. Valence bond theory (Section 1.5).
n n n n
1. Covalent bonds are formed by the overlap of two atomic orbitals, each of which
n n n n n n n n n n n n n
ncontains one electron. The two electrons have opposite spins.
n n n n n n n n
2. Bonds formed by the head-on overlap of two atomic orbitals are cylindrically
n n n n n n n n n n n
nsymmetrical and are called σ bonds. n n n n n
3. Bond strength is the measure of the amount of energy needed to break a bond.
n n n n n n n n n n n n n n
4. Bond length is the optimum distance between nuclei.
n n n n n n n
5. Every bond has a characteristic bond length and bond strength.
n n n n n n n n n
III. Hybridization (Sections 1.6–1.10). n n
A. sp3 Orbitals (Sections 1.6, 1.7).
n n n n
1. Structure of methane (Section 1.6). n n n n
a. When carbon forms 4 bonds with hydrogen, one 2s orbital and three 2p orbitals
n n n n n n n n n n n n n
combine to form four equivalent atomic orbitals (sp3 hybrid orbitals).
n n n n n n n n n n
b. These orbitals are tetrahedrally oriented. n n n n
c. Because these orbitals are unsymmetrical, they can form stronger bonds than
n n n n n n n n n n
unhybridized orbitals can.
n n n
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d. These bonds have a specific geometry and a bond angle of 109.5°.
n n n n n n n n n n n
2. Structure of ethane (Section 1.7). n n n n
a. Ethane has the same type of hybridization as occurs in methane. n n n n n n n n n n
b. The C–C bond is formed by overlap of two sp3 orbitals.
n n n n n n n n n n
c. Bond lengths, strengths and angles are very close to those of methane.
n n n n n n n n n n n
2
B. sp Orbitals (Section 1.8).
n n n
1. If one carbon 2s orbital combines with two carbon 2p orbitals, three hybrid sp2
n n n n n n n n n n n n n
orbitals are formed, and one p orbital remains unchanged.
n n n n n n n n
2. The three sp2 orbitals lie in a plane at angles of 120°, and the unhybridized p orbital is
n n n n n n n n n n n n n n n n n
perpendicular to them.
n n n
3. Two different types of bonds form between two carbons.
n n n n n n n n
a. A σ bond forms from the overlap of two sp2 orbitals.
n n n n n n n n n n
b. A π bond forms by sideways overlap of two p orbitals.
n n n n n n n n n n
c. This combination is known as a carbon–carbon double bond.
n n n n n n n n
4. Ethylene is composed of a carbon–carbon double bond and four σ bonds formed
n n n n n n n n n n n n
between the remaining four sp2 orbitals of carbon and the 1s orbitals of
n n n n n n n n n n n n n
hydrogen.
n
a. The double bond of ethylene is both shorter and stronger than the C–C bond of
n n n n n n n n n n n n n n
ethane. n
C. sp Orbitals (Section 1.10).
n n n
1. If one carbon 2s orbital combines with one carbon 2p orbital, two hybrid sp orbitals
n n n n n n n n n n n n n n
are formed, and two p orbitals are unchanged.
n n n n n n n n
2. The two sp orbitals are 180° apart, and the two p orbitals are perpendicular to them
n n n n n n n n n n n n n n n
and to each other.
n n n n
3. Two different types of bonds form.
n n n n n
a. A σ bond forms from the overlap of two sp orbitals.
n n n n n n n n n n
b. Two π bonds form by sideways overlap of four unhybridized p orbitals.
n n n n n n n n n n n
c. This combination is known as a carbon–carbon triple bond.
n n n n n n n n
4. Acetylene is composed of a carbon–carbon triple bond and two σ bonds formed
n n n n n n n n n n n n
between the remaining two sp orbitals of carbon and the 1s orbitals of hydrogen.
n n n n n n n n n n n n n n
a. The triple bond of acetylene is the strongest carbon–carbon bond.
n n n n n n n n n
D. Hybridization of nitrogen and oxygen (Section 1.10). n n n n n n
1. Covalent bonds between other elements can be described by using hybrid orbitals.
n n n n n n n n n n n
2. Both the nitrogen atom in ammonia and the oxygen atom in water form sp3 hybrid
n n n n n n n n n n n n n n
orbitals.
n
3. The bond angles between hydrogen and the central atom is often less than 109°
n n n n n n n n n n n n n
because the lone-pair electrons take up more room than the bond.
n n n n n n n n n n n n
4. Because of their positions in the third row, phosphorus and sulfur can form more than
n n n n n n n n n n n n n n
the typical number of covalent bonds.
n n n n n n
IV. Molecular orbital theory (Section 1.11).
n n n n
A. Molecular orbitals arise from a mathematical combination of atomic orbitals and
n n n n n n n n n n
belong to the entire molecule.
n n n n n
1. Two 1s orbitals can combine in two different ways.
n n n n n n n n
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a. The additive combination is a bonding MO and is lower in energy than the two
n n n n n n n n n n n n n n
hydrogen 1s atomic orbitals.
n n n n
b. The subtractive combination is an antibonding MO and is higher in energy than
n n n n n n n n n n n n
the two hydrogen 1s atomic orbitals.
n n n n n n
2. Two p orbitals in ethylene can combine to form two π MOs.
n n n n n n n n n n n
a. The bonding MO has no node; the antibonding MO has one node.
n n n n n n n n n n n
3. A node is a region between nuclei where electrons aren’t found.
n n n n n n n n n n
a. If a node occurs between two nuclei, the nuclei repel each other.
n n n n n n n n n n n
V. Chemical structures (Section 1.12).
n n n
A. Drawing chemical structures.
n n
1. Condensed structures don’t show C–H bonds and don’t show the bonds between CH3,
n n n n n n n n n n n n
CH2 and CH units.
n n n n
2. Skeletal structures are simpler still.
n n n n
a. Carbon atoms aren’t usually shown. n n n n
b. Hydrogen atoms bonded to carbon aren’t usually shown.
n n n n n n n
c. Other atoms (O, N, Cl, etc.) are shown.
n n n n n n n
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Chapter 2 – Polar Covalent Bonds; Acids and Bases n n n n n n n n
Chapter Outline n
I. Polar covalent bonds (Sections 2.1–2.3).
n n n n
A. Electronegativity (Section 2.1). n n
1. Although some bonds are totally ionic and some are totally covalent, most chemical
n n n n n n n n n n n n
bonds are polar covalent bonds.
n n n n n
a. In these bonds, electrons are attracted to one atom more than to the other atom.
n n n n n n n n n n n n n n
2. Bond polarity is due to differences in electronegativity (EN).
n n n n n n n n
a. Elements on the right side of the periodic table are more electronegative than
n n n n n n n n n n n n
elements on the left side.
n n n n n
b. Carbon has an EN of 2.5. n n n n n
c. Elements with EN > 2.5 are more electronegative than carbon. n n n n n n n n n
d. Elements with EN < 2.5 are less electronegative than carbon. n n n n n n n n n
3. The difference in EN between two elements can be used to predict the polarity of a
n n n n n n n n n n n n n n n
bond.
n
a. If ΔEN < 0.4, a bond is nonpolar covalent.
n n n n n n n n
b. If ΔEN is between 0.4 and 2.0, a bond is polar covalent.
n n n n n n n n n n n
c. If ΔEN > 2.0, a bond is ionic.
n n n n n n n
d. The symbols δ+ and δ– are used to indicate partial charges.
n n n n n n n n n n
e. A crossed arrow is used to indicate bond polarity.
n n n n n n n n
i. The tail of the arrow is electron-poor, and the head of the arrow is electron-
n n n n n n n n n n n n n n
rich. n
4. Electrostatic potential maps are also used to show electron-rich (red) and electron-
n n n n n n n n n n n
poor (blue) regions of molecules.
n n n n n
5. An inductive effect is an atom’s ability to polarize a bond.
n n n n n n n n n n
B. Dipole moment (Section 2.2).
n n n
1. Net dipole moment is the measure of a molecule’s overall polarity.
n n n n n n n n n n
2. Bond dipole moment (μ) = Q × r, where Q = charge and r = distance between charges.
n n n n n n n n n n n n n n n n n
a. Dipole moment is measured in debyes (D). n n n n n n
3. Dipole moment can be used to measure charge separation in a bond.
n n n n n n n n n n n
4. Water and ammonia have large values of D; methane and ethane have D = 0.
n n n n n n n n n n n n n n
C. Formal charge (Section 2.3).
n n n
1. Formal charge (FC) indicates electron “ownership” in a molecule.
# of bonding electrons
n n n n n n n n
2. (FC) =[# of valence electrons]− −[# nonbinding electrons]
n n n n n
n n n n n n n n n
2
II. Resonance (Sections 2.4–2.6).
n n
A. Chemical structures and resonance (Section 2.4).
n n n n n
1. Some molecules (acetate ion, for example) can be drawn as two (or more) different
n n n n n n n n n n n n n
electron-dot structures.
n n
a. These structures are called resonance structures.
n n n n n
b. The true structure of the molecule is intermediate between the
n n n n n n n n n
resonance structures.
n n
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c. The true structure is called a resonance hybrid.
n n n n n n n
2. Resonance structures differ only in the placement of π and nonbonding electrons.
n n n n n n n n n n n
a. All atoms occupy the same positions. n n n n n
3. Resonance is an important concept in organic chemistry. n n n n n n n
B. Rules for resonance forms (Section 2.5).
n n n n n
1. Individual resonance forms are imaginary, not real. n n n n n n
2. Resonance forms differ only in the placement of their π or nonbonding electrons.
n n n n n n n n n n n n
a. A curved arrow is used to indicate the movement of electrons, not atoms.
n n n n n n n n n n n n
3. Different resonance forms of a molecule don’t have to be equivalent.
n n n n n n n n n n
a. If resonance forms are nonequivalent, the structure of the actual molecule
n n n n n n n n n n
resembles the more stable resonance form(s).
n n n n n n
4. Resonance forms must obey normal rules of valency. n n n n n n n
5. The resonance hybrid is more stable than any individual resonance form.
n n n n n n n n n n
C. A useful technique for drawing resonance forms (Section 2.6).
n n n n n n n n
1. Any three-atom grouping with a multiple bond adjacent to a nonbonding p orbital has
n n n n n n n n n n n n n
two resonance forms.
n n n
2. One atom in the grouping has a lone electron pair, a vacant orbital or a single
n n n n n n n n n n n n n n n
electron.
n
3. By recognizing these three-atom pieces, resonance forms can be generated.
n n n n n n n n n
III. Acids and bases (Sections 2.7–2.11).
n n n n
A. Brønsted–Lowry definition (Section 2.7). n n n
1. A Brønsted–Lowry acid donates an H+ ion; a Brønsted–Lowry base accepts H+.
n n n n n n n n n n n
2. The product that results when a base gains H+ is the conjugate acid of the base; the
n n n n n n n n n n n n n n n n
product that results when an acid loses H+ is the conjugate base of the acid.
n n n n n n n n n n n n n n n
3. Water can act either as an acid or as a base.
n n n n n n n n n n
B. Acid and base strength (Section 2.8–2.10).
n n n n n
1. A strong acid dissociates almost completely with water (Section 2.8).
n n n n n n n n n
2. The strength of an acid in water is indicated by Ka, the acidity constant.
n n n n n n n n n n n n n
3. Strong acids have large acidity constants, and weaker acids have smaller acidity
n n n n n n n n n n n
constants.
n
4. The pKa is normally used to express acid strength.
n n n n n n n n
a. pKa = –log Ka n n n
b. A strong acid has a small pKa, and a weak acid has a large pKa.
n n n n n n n n n n n n n n
c. The conjugate base of a strong acid is a weak base, and the conjugate base of a
n n n n n n n n n n n n n n n n
weak acid is a strong base.
n n n n n n
5. Predicting acid–base reactions from pKa (Section 2.9). n n n n n n
a. An acid with a low pKa (stronger acid) reacts with the conjugate base of an acid
n n n n n n n n n n n n n n n
with a high pKa (stronger base).
n n n n n n
b. In other words, the products of an acid–base reaction are more stable than the
n n n n n n n n n n n n n
reactants. n
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Preface
What enters your mind when you hear the words “organic chemistry?” Some of you may
think, “the chemistry of life,” or “the chemistry of carbon.” Other responses might include “pre-
med,” “pressure,” “difficult,” or “memorization.” Although formally the study of the
compounds of carbon, the discipline of organic chemistry encompasses many skills that are
common to other areas of study. Organic chemistry is as much a liberal art as a science, and
mastery of the concepts and techniques of organic chemistry can lead to improved competence
in other fields.
Here are several suggestions that may help you with problem solving:
1. The text is organized into chapters that describe individual functional groups. As you
study each functional group, make sure that you understand the structure and reactivity
of that group. In case your memory of a specific reaction fails you, you can rely on your
general knowledge of functional groups for help.
2. Use molecular models. It is difficult to visualize the three-dimensional structure of an
organic molecule when looking at a two-dimensional drawing. Models will help you to
appreciate the structural aspects of organic chemistry and are indispensable tools for
understanding stereochemistry.
3. Look through the appendices at the end of the Study Guide. Some of these
appendices contain tables that may help you in working problems; others present
information related to the history of organic chemistry.
While the Study Guide is written to accompany Organic Chemistry, it also contains several
unique features. Each chapter of the Study Guide begins with an outline of the text that can be
used for a concise review of the text material and can also serve as a reference. After every few
chapters a Review Unit is included. In most cases, the chapters covered in the Review Units are
related to each other, and the units are planned to appear at approximately the place in the
textbook where a test might be given. Each unit lists the vocabulary for the chapters covered,
the skills needed to solve problems, and several important points that might need reinforcing or
that restate material in the text from a slightly different point of view. Finally, the small self-test
that has been included allows you to test yourself on the material from more than one chapter.
I have tried to include many types of study aids in this Study Guide. Nevertheless, this book can
only serve as an adjunct to the larger and more complete textbook. In addition, please note that
a companion to this Study Guide, the Student Solutions Manual, contains solutions to the
problems found in Organic Chemistry: A Tenth Edition. I am pleased that this is the first time
that these are published by OpenStax and available to you for free in honor of my late son, Peter
McMurry.
Susan McMurry
Acknowledgments
Bryant Gilbert, Zane State College
Kjir Hendrickson, Arizona State University
Ike Shibley, Penn State Berks
10/27/2023 2
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Table of Contents
Chapter 1: Structure and Bonding 5
Chapter 2: Polar Covalent Bonds; Acids and Bases 9
Review Unit 1 12
Chapter 3: Organic Compounds: Alkanes and Their Stereochemistry 15
Chapter 4: Organic Compounds: Cycloalkanes and Their Stereochemistry 19
Chapter 5: Stereochemistry at Tetrahedral Centers 22
Review Unit 2 25
Chapter 6: An Overview of Organic Reactions 29
Chapter 7: Alkenes: Structure and Reactivity 32
Chapter 8: Alkenes: Reactions and Synthesis 35
Review Unit 3 38
Chapter 9: Alkynes: An Introduction to Organic Synthesis 43
Chapter 10: Organohalides 45
Chapter 11: Reactions of Alkyl Halides: Nucleophilic Substitutions and Eliminations 48
Review Unit 4 53
Chapter 12: Structure Determination: Mass Spectrometry and Infrared Spectroscopy 56
Chapter 13: Structure Determination: Nuclear Magnetic Resonance Spectroscopy 60
Review Unit 5 65
Chapter 14: Conjugated Compounds and Ultraviolet Spectroscopy 70
Chapter 15: Benzene and Aromaticity 73
Chapter 16: Chemistry of Benzene: Electrophilic Aromatic Substitution 77
Review Unit 6 81
Chapter 17: Alcohols and Phenols 86
Chapter 18: Ethers and Epoxides; Thiols and Sulfides 90
Preview of Carbonyl Chemistry 93
Review Unit 7 94
Chapter 19: Aldehydes and Ketones: Nucleophilic Addition Reactions 97
Chapter 20: Carboxylic Acids and Nitriles 102
Chapter 21: Carboxylic Acid Derivatives: Nucleophilic Acyl Substitution Reactions 105
Review Unit 8 110
Chapter 22: Carbonyl Alpha-Substitution Reactions 115
Chapter 23: Carbonyl Condensation Reactions 118
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Chapter 24: Amines and Heterocycles 121
Review Unit 9 126
Chapter 25: Biomolecules: Carbohydrates 131
Chapter 26: Biomolecules: Amino Acids, Peptides, and Proteins 136
Review Unit 10 142
Chapter 27: Biomolecules: Lipids 147
Chapter 28: Biomolecules: Nucleic Acids 150
Chapter 29: The Organic Chemistry of Metabolic Pathways 154
Review Unit 11 160
Chapter 30: Orbitals and Organic Chemistry: Pericyclic Reactions 165
Chapter 31: Synthetic Polymers 168
Review Unit 12 171
Appendices
Appendix A: Functional-Group Synthesis 175
Appendix B: Functional-Group Reactions 181
Appendix C: Reagents in Organic Chemistry 186
Appendix D: Name Reactions in Organic Chemistry 194
Appendix E: Abbreviations 202
Appendix F: Infrared Absorption Frequencies 206
Appendix G: Proton NMR Chemical Shifts 209
Appendix H: Nobel Prize Winners in Chemistry 210
Appendix I: Answers to Multiple Choice Questions in Review Units 1–12 222
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Chapter 1 – Structure and Bonding n n n n n
Chapter Outline n
I. Atomic Structure (Sections 1.1–1.3).
n n n
A. Introduction to atomic structure (Section 1.1).
n n n n n
1. An atom consists
n n
2. of a dense, positively charged nucleus surrounded by negatively charged
n n n n n n n n n
electrons.
n
a. The nucleus is made up of positively charged protons and uncharged neutrons.
n n n n n n n n n n n
b. The nucleus contains most of the mass of the atom.
n n n n n n n n n
c. Electrons move about the nucleus at a distance of about 2 ×10−10 m (200 pm).
n n n n n n n n n n n n n n
3. The atomic number (Z) gives the number of protons in the nucleus.
n n n n n n n n n n n
4. The mass number (A) gives the total number of protons and neutrons.
n n n n n n n n n n n
5. All atoms of a given element have the same value of Z.
n n n n n n n n n n n
a. Atoms of a given element can have different values of A.
n n n n n n n n n n
b. Atoms of the same element with different values of A are called isotopes.
n n n n n n n n n n n n
B. Orbitals (Section 1.2).
n n
1. The distribution of electrons in an atom can be described by a wave equation.
n n n n n n n n n n n n n
a. The solution to a wave equation is an orbital, represented by Ψ.
n n n n n n n n n n n
b. Ψ 2 predicts the volume of space in which an electron is likely to be found.
n n n n n n n n n n n n n n n
2. There are four different kinds of orbitals (s, p, d, f).
n n n n n n n n n n
a. The s orbitals are spherical.n n n n
b. The p orbitals are dumbbell-shaped.
n n n n
c. Four of the five d orbitals are cloverleaf-shaped.
n n n n n n n
3. An atom’s electrons are organized into electron shells.
n n n n n n n
a. The shells differ in the numbers and kinds of orbitals they contain.
n n n n n n n n n n n
b. Electrons in different orbitals have different energies. n n n n n n
c. Each orbital can hold up to a maximum of two electrons.
n n n n n n n n n n
4. The two lowest-energy electrons are in the 1s orbital.
n n n n n n n n
a. The 2s orbital is the next higher in energy.
n n n n n n n n
b. The next three orbitals are 2px, 2py and 2pz, which have the same energy.
n n n n n n n n n n n n n
i. Each p orbital has a region of zero density, called a node.
n n n n n n n n n n n
c. The lobes of a p orbital have opposite algebraic signs.
n n n n n n n n n
C. Electron Configuration (Section 1.3).
n n n
1. The ground-state electron configuration of an atom is a listing of the orbitals
n n n n n n n n n n n n
occupied by the electrons of the atom in the lowest energy configuration.
n n n n n n n n n n n n
2. Rules for predicting the ground-state electron configuration of an atom:
n n n n n n n n n
a. Orbitals with the lowest energy levels are filled first. n n n n n n n n
i. The order of filling is 1s, 2s, 2p, 3s, 3p, 4s, 3d.
n n n n n n n n n n n
b. Only two electrons can occupy each orbital, and they must be of opposite spin.
n n n n n n n n n n n n n
c. If two or more orbitals have the same energy, one electron occupies each until all
n n n n n n n n n n n n n n
are half-full (Hund’s rule). Only then does a second electron occupy one
n n n n n n n n n n n n
of the orbitals. n n
i. All of the electrons in half-filled shells have the same spin.
n n n n n n n n n n
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II. Chemical Bonding Theory (Sections 1.4–1.5).
n n n n
A. Development of chemical bonding theory (Section 1.4). n n n n n n
1. Kekulé and Couper proposed that carbon has four “affinity units”; carbon is tetravalent.
n n n n n n n n n n n n
2. Kekulé suggested that carbon can form rings and chains.
n n n n n n n n
3. Van’t Hoff and Le Bel proposed that the 4 atoms to which carbon forms bonds sit at
n n n n n n n n n n n n n n n n
the corners of a regular tetrahedron.
n n n n n n
4. In a drawing of a tetrahedral carbon, a wedged line represents a bond pointing toward
n n n n n n n n n n n n n n
the viewer, a dashed line points behind the plane of the page, and a solid line lies in
n n n n n n n n n n n n n n n n n n
n the plane of the page.
n n n n
B. Covalent bonds. n
1. Atoms bond together because the resulting compound is more stable than the
n n n n n n n n n n n
nindividual atoms. n
a. Atoms tend to achieve the electron configuration of the nearest noble gas.
n n n n n n n n n n n
b. Atoms in groups 1A, 2A and 7A either lose electrons or gain electrons to form
n n n n n n n n n n n n n n
ionic compounds.
n n
c. Atoms in the middle of the periodic table share electrons by forming covalent
n n n n n n n n n n n n
bonds.
n
d. The neutral collection of atoms held together by covalent bonds is a molecule.
n n n n n n n n n n n n
2. Covalent bonds can be represented two ways. n n n n n n
a. In electron-dot structures, bonds are represented as pairs of dots.
n n n n n n n n n
b. In line-bond structures, bonds are represented as lines drawn between two bonded
n n n n n n n n n n n
atoms.
n
3. The number of covalent bonds formed by an atom depends on the number of
n n n n n n n n n n n n n
nelectrons it has and on the number it needs to achieve an octet.
n n n n n n n n n n n n
4. Valence electrons not used for bonding are called lone-pair (nonbonding)
n n n n n n n n n
nelectrons.
a. Lone-pair electrons are often represented as dots. n n n n n n
C. Valence bond theory (Section 1.5).
n n n n
1. Covalent bonds are formed by the overlap of two atomic orbitals, each of which
n n n n n n n n n n n n n
ncontains one electron. The two electrons have opposite spins.
n n n n n n n n
2. Bonds formed by the head-on overlap of two atomic orbitals are cylindrically
n n n n n n n n n n n
nsymmetrical and are called σ bonds. n n n n n
3. Bond strength is the measure of the amount of energy needed to break a bond.
n n n n n n n n n n n n n n
4. Bond length is the optimum distance between nuclei.
n n n n n n n
5. Every bond has a characteristic bond length and bond strength.
n n n n n n n n n
III. Hybridization (Sections 1.6–1.10). n n
A. sp3 Orbitals (Sections 1.6, 1.7).
n n n n
1. Structure of methane (Section 1.6). n n n n
a. When carbon forms 4 bonds with hydrogen, one 2s orbital and three 2p orbitals
n n n n n n n n n n n n n
combine to form four equivalent atomic orbitals (sp3 hybrid orbitals).
n n n n n n n n n n
b. These orbitals are tetrahedrally oriented. n n n n
c. Because these orbitals are unsymmetrical, they can form stronger bonds than
n n n n n n n n n n
unhybridized orbitals can.
n n n
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d. These bonds have a specific geometry and a bond angle of 109.5°.
n n n n n n n n n n n
2. Structure of ethane (Section 1.7). n n n n
a. Ethane has the same type of hybridization as occurs in methane. n n n n n n n n n n
b. The C–C bond is formed by overlap of two sp3 orbitals.
n n n n n n n n n n
c. Bond lengths, strengths and angles are very close to those of methane.
n n n n n n n n n n n
2
B. sp Orbitals (Section 1.8).
n n n
1. If one carbon 2s orbital combines with two carbon 2p orbitals, three hybrid sp2
n n n n n n n n n n n n n
orbitals are formed, and one p orbital remains unchanged.
n n n n n n n n
2. The three sp2 orbitals lie in a plane at angles of 120°, and the unhybridized p orbital is
n n n n n n n n n n n n n n n n n
perpendicular to them.
n n n
3. Two different types of bonds form between two carbons.
n n n n n n n n
a. A σ bond forms from the overlap of two sp2 orbitals.
n n n n n n n n n n
b. A π bond forms by sideways overlap of two p orbitals.
n n n n n n n n n n
c. This combination is known as a carbon–carbon double bond.
n n n n n n n n
4. Ethylene is composed of a carbon–carbon double bond and four σ bonds formed
n n n n n n n n n n n n
between the remaining four sp2 orbitals of carbon and the 1s orbitals of
n n n n n n n n n n n n n
hydrogen.
n
a. The double bond of ethylene is both shorter and stronger than the C–C bond of
n n n n n n n n n n n n n n
ethane. n
C. sp Orbitals (Section 1.10).
n n n
1. If one carbon 2s orbital combines with one carbon 2p orbital, two hybrid sp orbitals
n n n n n n n n n n n n n n
are formed, and two p orbitals are unchanged.
n n n n n n n n
2. The two sp orbitals are 180° apart, and the two p orbitals are perpendicular to them
n n n n n n n n n n n n n n n
and to each other.
n n n n
3. Two different types of bonds form.
n n n n n
a. A σ bond forms from the overlap of two sp orbitals.
n n n n n n n n n n
b. Two π bonds form by sideways overlap of four unhybridized p orbitals.
n n n n n n n n n n n
c. This combination is known as a carbon–carbon triple bond.
n n n n n n n n
4. Acetylene is composed of a carbon–carbon triple bond and two σ bonds formed
n n n n n n n n n n n n
between the remaining two sp orbitals of carbon and the 1s orbitals of hydrogen.
n n n n n n n n n n n n n n
a. The triple bond of acetylene is the strongest carbon–carbon bond.
n n n n n n n n n
D. Hybridization of nitrogen and oxygen (Section 1.10). n n n n n n
1. Covalent bonds between other elements can be described by using hybrid orbitals.
n n n n n n n n n n n
2. Both the nitrogen atom in ammonia and the oxygen atom in water form sp3 hybrid
n n n n n n n n n n n n n n
orbitals.
n
3. The bond angles between hydrogen and the central atom is often less than 109°
n n n n n n n n n n n n n
because the lone-pair electrons take up more room than the bond.
n n n n n n n n n n n n
4. Because of their positions in the third row, phosphorus and sulfur can form more than
n n n n n n n n n n n n n n
the typical number of covalent bonds.
n n n n n n
IV. Molecular orbital theory (Section 1.11).
n n n n
A. Molecular orbitals arise from a mathematical combination of atomic orbitals and
n n n n n n n n n n
belong to the entire molecule.
n n n n n
1. Two 1s orbitals can combine in two different ways.
n n n n n n n n
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a. The additive combination is a bonding MO and is lower in energy than the two
n n n n n n n n n n n n n n
hydrogen 1s atomic orbitals.
n n n n
b. The subtractive combination is an antibonding MO and is higher in energy than
n n n n n n n n n n n n
the two hydrogen 1s atomic orbitals.
n n n n n n
2. Two p orbitals in ethylene can combine to form two π MOs.
n n n n n n n n n n n
a. The bonding MO has no node; the antibonding MO has one node.
n n n n n n n n n n n
3. A node is a region between nuclei where electrons aren’t found.
n n n n n n n n n n
a. If a node occurs between two nuclei, the nuclei repel each other.
n n n n n n n n n n n
V. Chemical structures (Section 1.12).
n n n
A. Drawing chemical structures.
n n
1. Condensed structures don’t show C–H bonds and don’t show the bonds between CH3,
n n n n n n n n n n n n
CH2 and CH units.
n n n n
2. Skeletal structures are simpler still.
n n n n
a. Carbon atoms aren’t usually shown. n n n n
b. Hydrogen atoms bonded to carbon aren’t usually shown.
n n n n n n n
c. Other atoms (O, N, Cl, etc.) are shown.
n n n n n n n
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Chapter 2 – Polar Covalent Bonds; Acids and Bases n n n n n n n n
Chapter Outline n
I. Polar covalent bonds (Sections 2.1–2.3).
n n n n
A. Electronegativity (Section 2.1). n n
1. Although some bonds are totally ionic and some are totally covalent, most chemical
n n n n n n n n n n n n
bonds are polar covalent bonds.
n n n n n
a. In these bonds, electrons are attracted to one atom more than to the other atom.
n n n n n n n n n n n n n n
2. Bond polarity is due to differences in electronegativity (EN).
n n n n n n n n
a. Elements on the right side of the periodic table are more electronegative than
n n n n n n n n n n n n
elements on the left side.
n n n n n
b. Carbon has an EN of 2.5. n n n n n
c. Elements with EN > 2.5 are more electronegative than carbon. n n n n n n n n n
d. Elements with EN < 2.5 are less electronegative than carbon. n n n n n n n n n
3. The difference in EN between two elements can be used to predict the polarity of a
n n n n n n n n n n n n n n n
bond.
n
a. If ΔEN < 0.4, a bond is nonpolar covalent.
n n n n n n n n
b. If ΔEN is between 0.4 and 2.0, a bond is polar covalent.
n n n n n n n n n n n
c. If ΔEN > 2.0, a bond is ionic.
n n n n n n n
d. The symbols δ+ and δ– are used to indicate partial charges.
n n n n n n n n n n
e. A crossed arrow is used to indicate bond polarity.
n n n n n n n n
i. The tail of the arrow is electron-poor, and the head of the arrow is electron-
n n n n n n n n n n n n n n
rich. n
4. Electrostatic potential maps are also used to show electron-rich (red) and electron-
n n n n n n n n n n n
poor (blue) regions of molecules.
n n n n n
5. An inductive effect is an atom’s ability to polarize a bond.
n n n n n n n n n n
B. Dipole moment (Section 2.2).
n n n
1. Net dipole moment is the measure of a molecule’s overall polarity.
n n n n n n n n n n
2. Bond dipole moment (μ) = Q × r, where Q = charge and r = distance between charges.
n n n n n n n n n n n n n n n n n
a. Dipole moment is measured in debyes (D). n n n n n n
3. Dipole moment can be used to measure charge separation in a bond.
n n n n n n n n n n n
4. Water and ammonia have large values of D; methane and ethane have D = 0.
n n n n n n n n n n n n n n
C. Formal charge (Section 2.3).
n n n
1. Formal charge (FC) indicates electron “ownership” in a molecule.
# of bonding electrons
n n n n n n n n
2. (FC) =[# of valence electrons]− −[# nonbinding electrons]
n n n n n
n n n n n n n n n
2
II. Resonance (Sections 2.4–2.6).
n n
A. Chemical structures and resonance (Section 2.4).
n n n n n
1. Some molecules (acetate ion, for example) can be drawn as two (or more) different
n n n n n n n n n n n n n
electron-dot structures.
n n
a. These structures are called resonance structures.
n n n n n
b. The true structure of the molecule is intermediate between the
n n n n n n n n n
resonance structures.
n n
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c. The true structure is called a resonance hybrid.
n n n n n n n
2. Resonance structures differ only in the placement of π and nonbonding electrons.
n n n n n n n n n n n
a. All atoms occupy the same positions. n n n n n
3. Resonance is an important concept in organic chemistry. n n n n n n n
B. Rules for resonance forms (Section 2.5).
n n n n n
1. Individual resonance forms are imaginary, not real. n n n n n n
2. Resonance forms differ only in the placement of their π or nonbonding electrons.
n n n n n n n n n n n n
a. A curved arrow is used to indicate the movement of electrons, not atoms.
n n n n n n n n n n n n
3. Different resonance forms of a molecule don’t have to be equivalent.
n n n n n n n n n n
a. If resonance forms are nonequivalent, the structure of the actual molecule
n n n n n n n n n n
resembles the more stable resonance form(s).
n n n n n n
4. Resonance forms must obey normal rules of valency. n n n n n n n
5. The resonance hybrid is more stable than any individual resonance form.
n n n n n n n n n n
C. A useful technique for drawing resonance forms (Section 2.6).
n n n n n n n n
1. Any three-atom grouping with a multiple bond adjacent to a nonbonding p orbital has
n n n n n n n n n n n n n
two resonance forms.
n n n
2. One atom in the grouping has a lone electron pair, a vacant orbital or a single
n n n n n n n n n n n n n n n
electron.
n
3. By recognizing these three-atom pieces, resonance forms can be generated.
n n n n n n n n n
III. Acids and bases (Sections 2.7–2.11).
n n n n
A. Brønsted–Lowry definition (Section 2.7). n n n
1. A Brønsted–Lowry acid donates an H+ ion; a Brønsted–Lowry base accepts H+.
n n n n n n n n n n n
2. The product that results when a base gains H+ is the conjugate acid of the base; the
n n n n n n n n n n n n n n n n
product that results when an acid loses H+ is the conjugate base of the acid.
n n n n n n n n n n n n n n n
3. Water can act either as an acid or as a base.
n n n n n n n n n n
B. Acid and base strength (Section 2.8–2.10).
n n n n n
1. A strong acid dissociates almost completely with water (Section 2.8).
n n n n n n n n n
2. The strength of an acid in water is indicated by Ka, the acidity constant.
n n n n n n n n n n n n n
3. Strong acids have large acidity constants, and weaker acids have smaller acidity
n n n n n n n n n n n
constants.
n
4. The pKa is normally used to express acid strength.
n n n n n n n n
a. pKa = –log Ka n n n
b. A strong acid has a small pKa, and a weak acid has a large pKa.
n n n n n n n n n n n n n n
c. The conjugate base of a strong acid is a weak base, and the conjugate base of a
n n n n n n n n n n n n n n n n
weak acid is a strong base.
n n n n n n
5. Predicting acid–base reactions from pKa (Section 2.9). n n n n n n
a. An acid with a low pKa (stronger acid) reacts with the conjugate base of an acid
n n n n n n n n n n n n n n n
with a high pKa (stronger base).
n n n n n n
b. In other words, the products of an acid–base reaction are more stable than the
n n n n n n n n n n n n n
reactants. n
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