Preview A QA Approach to Organic Chemistry by Michael B. Smith (Author) (2020)

A Q&A Approach to
Organic Chemistry



A Q&A Approach to
Organic Chemistry

Michael B. Smith


First edition published 2020
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Contents
Preface ...................................................................................................................................................... ix
Common Abbreviations............................................................................................................................ xi
Author .....................................................................................................................................................xiii

Part A

A Q&A Approach to Organic Chemistry

1 Orbitals and Bonding....................................................................................................................... 3
1.1
ORBITALS ............................................................................................................................. 3
1.1.1 Atomic Orbitals......................................................................................................... 3
1.1.2 Electron Confguration.............................................................................................. 5
1.1.3 Molecular Orbitals .................................................................................................... 5
1.2
BONDING .............................................................................................................................. 6
1.2.1 Ionic Bonding............................................................................................................ 6
1.2.2 Covalent Bonding...................................................................................................... 7
1.3
HYBRIDIZATION............................................................................................................... 12
1.4
RESONANCE ...................................................................................................................... 15
END OF CHAPTER PROBLEMS ...................................................................................................18

2 Structure of Molecules................................................................................................................... 19
2.1
BASIC STRUCTURE OF ORGANIC MOLECULES ........................................................ 19
2.1.1 Fundamental Structures.......................................................................................... 19
2.1.2 Structures with Other Atoms Bonded to Carbon ................................................... 22
2.2
THE VSEPR MODEL AND MOLECULAR GEOMETRY............................................... 23
2.3
DIPOLE MOMENT ............................................................................................................. 25
2.4
FUNCTIONAL GROUPS .................................................................................................... 26
2.5
FORMAL CHARGE ............................................................................................................ 28
2.6
PHYSICAL PROPERTIES................................................................................................... 28
END OF CHAPTER PROBLEMS .................................................................................................. 32
3 Acids and Bases .............................................................................................................................. 33
3.1
ACIDS AND BASES ........................................................................................................... 33
3.2
ENERGETICS ...................................................................................................................... 35
3.3
THE ACIDITY CONSTANT, Ka ......................................................................................... 38
3.4
STRUCTURAL FEATURES THAT INFLUENCE ACIDITY.......................................... 40
3.5
FACTORS THAT CONTRIBUTE TO MAKING THE ACID MORE ACIDIC ............... 45
END OF CHAPTER PROBLEMS .................................................................................................. 48
4 Alkanes, Isomers, and Nomenclature .......................................................................................... 49
4.1

DEFINITION AND BASIC NOMENCLATURE............................................................... 49
4.2
STRUCTURAL ISOMERS.................................................................................................. 50
4.3
IUPAC NOMENCLATURE ................................................................................................ 52
4.4
CYCLIC ALKANES............................................................................................................ 57
END OF CHAPTER PROBLEMS .................................................................................................. 58

v


vi

Contents

5 Conformations .................................................................................................................................61
5.1
ACYCLIC CONFORMATIONS...........................................................................................61
5.2
CONFORMATIONS OF CYCLIC MOLECULES ............................................................. 67
END OF CHAPTER PROBLEMS .................................................................................................. 75
6 Stereochemistry.............................................................................................................................. 77
6.1
CHIRALITY ........................................................................................................................ 77
6.2
SPECIFIC ROTATION ........................................................................................................ 81
6.3
SEQUENCE RULES............................................................................................................ 83
6.4

DIASTEREOMERS ............................................................................................................. 87
6.5
OPTICAL RESOLUTION ................................................................................................... 89
END OF CHAPTER PROBLEMS .................................................................................................. 90
7 Alkenes and Alkynes: Structure, Nomenclature, and Reactions .............................................. 93
7.1
STRUCTURE OF ALKENES ............................................................................................. 93
7.2
NOMENCLATURE OF ALKENES ................................................................................... 95
7.3
REACTIONS OF ALKENES .............................................................................................. 98
7.4
REACTION OF ALKENES WITH LEWIS ACID-TYPE REAGENTS ......................... 107
7.4.1 Hydroxylation........................................................................................................ 107
7.4.2 Epoxidation ............................................................................................................111
7.4.3 Dihydroxylation .....................................................................................................113
7.4.4 Halogenation ..........................................................................................................114
7.4.5 Hydroboration ........................................................................................................117
7.5
STRUCTURE AND NOMENCLATURE OF ALKYNES............................................... 122
7.6
REACTIONS OF ALKYNES ............................................................................................ 124
END OF CHAPTER PROBLEMS ................................................................................................ 129
8 Alkyl Halides and Substitution Reactions ..................................................................................133
8.1
STRUCTURE, PROPERTIES, AND NOMENCLATURE OF ALKYL HALIDES........133
8.2
SECOND-ORDER NUCLEOPHILIC SUBSTITUTION (SN2) REACTIONS ................ 134
8.3
OTHER NUCLEOPHILES IN SN2 REACTIONS .............................................................143

8.4
FIRST-ORDER SUBSTITUTION (SN1) REACTIONS .....................................................151
8.5
COMPETITION BETWEEN SN2 vs. SN1 REACTIONS .................................................. 156
8.6
RADICAL HALOGENATION OF ALKANES ............................................................... 158
END OF CHAPTER PROBLEMS .................................................................................................162
9 Elimination Reactions...................................................................................................................165
9.1
THE E2 REACTION...........................................................................................................165
9.2
THE E1 REACTION ...........................................................................................................172
9.3
PREPARATION OF ALKYNES ........................................................................................176
9.4
SYN ELIMINATION ..........................................................................................................178
END OF CHAPTER PROBLEMS ................................................................................................ 180
10 Organometallic Compounds ........................................................................................................183
10.1 ORGANOMETALLICS ......................................................................................................183
10.2 ORGANOMAGNESIUM COMPOUNDS .........................................................................183
10.3 ORGANOLITHIUM COMPOUNDS.................................................................................185
10.4 BASICITY ...........................................................................................................................187
10.5 REACTION WITH EPOXIDES .........................................................................................188
10.6 OTHER METALS...............................................................................................................188
END OF CHAPTER PROBLEMS ................................................................................................ 190


vii

Contents

11 Spectroscopy ..................................................................................................................................191
11.1 THE ELECTROMAGNETIC SPECTRUM .......................................................................191
11.2 MASS SPECTROMETRY ................................................................................................. 192
11.3 INFRARED SPECTROSCOPY (IR)................................................................................. 196
11.4 NUCLEAR MAGNETIC RESONANCE SPECTROSCOPY (nmr)................................. 201
END OF CHAPTER PROBLEMS .................................................................................................215
12 Aldehydes and Ketones. Acyl Addition Reactions .....................................................................219
12.1 STRUCTURE AND NOMENCLATURE OF ALDEHYDES AND KETONES.............219
12.2 REACTION OF ALDEHYDES AND KETONES WITH WEAK NUCLEOPHILES ...... 221
12.3 REACTIONS OF ALDEHYDES AND KETONES. STRONG NUCLEOPHILES ........ 230
12.4 THE WITTIG REACTION................................................................................................ 233
END OF CHAPTER PROBLEMS ................................................................................................ 235

Part B

A Q&A Approach to Organic Chemistry

13 Oxidation Reactions ..................................................................................................................... 239
13.1 OXIDATION REACTIONS OF ALKENES ..................................................................... 239
13.2 OXIDATION OF ALKENES: EPOXIDATION ................................................................ 244
13.3 OXIDATIVE CLEAVAGE: OZONOLYSIS....................................................................... 247
13.4 OXIDATIVE CLEAVAGE. PERIODIC ACID CLEAVAGE OF 1,2-DIOLS ................... 250
13.5 OXIDATION OF ALCOHOLS TO ALDEHYDES OR KETONES .................................251
END OF CHAPTER PROBLEMS ................................................................................................ 255
14 Reduction Reactions..................................................................................................................... 257
14.1 CATALYTIC HYDROGENATION................................................................................... 258
14.2 DISSOLVING METAL REDUCTION: ALKYNES......................................................... 264
14.3 HYDRIDE REDUCTION OF ALDEHYDES AND KETONES ..................................... 265
14.4 CATALYTIC HYDROGENATION AND DISSOLVING METAL REDUCTIONS.

ALDEHYDES AND KETONES....................................................................................... 269
END OF CHAPTER PROBLEMS ................................................................................................ 273
15 Carboxylic Acids, Carboxylic Acid Derivatives, and Acyl Substitution Reactions ............... 275
15.1 STRUCTURE OF CARBOXYLIC ACIDS ....................................................................... 275
15.2 PREPARATION OF CARBOXYLIC ACIDS ................................................................... 280
15.3 CARBOXYLIC ACID DERIVATIVES ............................................................................ 283
15.4 PREPARATION OF ACID DERIVATIVES..................................................................... 290
15.5 HYDROLYSIS OF CARBOXYLIC ACID DERIVATIVES ............................................ 301
15.6 REACTIONS OF CARBOXYLIC ACIDS AND ACID DERIVATIVES ........................ 305
15.7 DIBASIC CARBOXYLIC ACIDS......................................................................................310
END OF CHAPTER PROBLEMS .................................................................................................312
16 Benzene, Aromaticity, and Benzene Derivatives........................................................................315
16.1 BENZENE AND NOMENCLATURE OF AROMATIC COMPOUNDS ........................315
16.2 ELECTROPHILIC AROMATIC SUBSTITUTION ..........................................................319
16.3 SYNTHESIS VIA AROMATIC SUBSTITUTION ...........................................................335
16.4 NUCLEOPHILIC AROMATIC SUBSTITUTION ........................................................... 337
16.5 REDUCTION OF BENZENE AND BENZENE DERIVATIVES ................................... 344
16.6 POLYCYCLIC AROMATIC COMPOUNDS AND HETEROAROMATIC
COMPOUNDS ................................................................................................................... 347
END OF CHAPTER PROBLEMS .................................................................................................353


viii

Contents

17 Enolate Anions and Condensation Reactions............................................................................ 357
17.1 ALDEHYDES, KETONES, ENOLS, AND ENOLATE ANIONS .................................. 357
17.2 ENOLATE ALKYLATION ................................................................................................361
17.3 CONDENSATION REACTIONS OF ENOLATE ANIONS AND ALDEHYDES

OR KETONES.................................................................................................................... 366
17.4 ENOLATE ANIONS FROM CARBOXYLIC ACIDS AND DERIVATIVES ................ 372
END OF CHAPTER PROBLEMS ................................................................................................ 383
18 Conjugation and Reactions of Conjugated Compounds........................................................... 385
18.1 CONJUGATED MOLECULES ......................................................................................... 385
18.2 STRUCTURE AND NOMENCLATURE OF CONJUGATED SYSTEMS .................... 387
18.3 REACTIONS OF CONJUGATED MOLECULES ............................................................391
18.4 THE DIELS–ALDER REACTION ................................................................................... 393
18.5 [3+2]-CYCLOADDITION REACTIONS .......................................................................... 401
18.6 SIGMATROPIC REARRANGEMENTS .......................................................................... 403
18.7 ULTRAVIOLET SPECTROSCOPY.................................................................................. 406
END OF CHAPTER PROBLEMS ................................................................................................ 409
19 Amines............................................................................................................................................413
19.1 STRUCTURE AND PROPERTIES....................................................................................413
19.2 PREPARATION OF AMINES ...........................................................................................416
19.3 REACTIONS OF AMINES ............................................................................................... 420
19.4 HETEROCYCLIC AMINES ............................................................................................. 424
END OF CHAPTER PROBLEMS ................................................................................................ 426
20 Amino Acids, Peptides, and Proteins ......................................................................................... 429
20.1 AMINO ACIDS.................................................................................................................. 429
20.2 SYNTHESIS OF AMINO ACIDS......................................................................................435
20.3 REACTIONS OF AMINO ACIDS .................................................................................... 437
20.4 PROTEINS ......................................................................................................................... 441
END OF CHAPTER PROBLEMS ................................................................................................ 447
21 Carbohydrates and Nucleic Acids............................................................................................... 449
21.1 CARBOHYDRATES ......................................................................................................... 449
21.2 DISACCHARIDES AND POLYSACCHARIDES............................................................ 457
21.3 SYNTHESIS OF CARBOHYDRATES ............................................................................ 459
21.4 REACTIONS OF CARBOHYDRATES............................................................................ 461
21.5 NUCLEIC ACIDS, NUCLEOTIDES, AND NUCLEOSIDES......................................... 464

END OF CHAPTER PROBLEMS .................................................................................................471
Appendix: Answers to End of Chapter Problems ............................................................................ 473
Index...................................................................................................................................................... 505


Preface
What is organic chemistry?
Organic chemistry is the science that studies molecules containing the element carbon. Carbon can
form bonds to other carbon atoms or to a variety of atoms in the periodic table. The most common
bonds observed in an organic chemistry course are C—C, C—H, C—O, C—N, C—halogen (Cl, Br, I),
C—Mg, C—B, C—Li, C—S and C—P.
This book is presented in the hope that it will provide extra practice to students taking organic chemistry for the frst time and also serve as a cogent review to those who need to refresh their knowledge
of organic chemistry. This book of questions began life as Organic Chemistry in 1993 to assist those
students taking undergraduate organic chemistry and was part of a HarperCollins Outline series that was
never completed. My book, along with those other books in the series that were completed, was sold as
a reference book rather than a textbook. In 2006, a second edition of Organic Chemistry was published
and marketed more or less the same way. The book laid fallow for several years until this version became
possible. With this book, published by CRC Press/Taylor & Francis Group, I continue the idea of teaching organic chemistry by asking leading questions.
A Q&A Approach to Organic Chemistry is intended as a supplement to virtually any organic chemistry textbook rather than a stand-alone text and it will allow a “self-guided tour” of organic chemistry.
Teaching organic chemistry with a Q&A format uses leading questions along with the answers and is
presented in a manner that allows a student to refresh and renew their working knowledge of organic
chemistry. Such an approach will also be of value to those reviewing organic chemistry for MCATs
(Medical College Admission Test); graduate record exams (a standardized test), which is an admissions
requirement for many graduate schools); the PCAT (Pharmacy College Admission Test), which identifes qualifed applicants to pharmacy colleges before commencement of pharmaceutical education; and
so on.
This Q&A format was classroom-tested here at the University of Connecticut for many years where
one of the earlier versions of this book was used as a supplement. Indeed, the book was not required for
purchase and used only on a voluntary basis by students. According to their end-of-semester evaluations,
students who wanted or needed additional homework found the book very useful and helpful. Classroom
experience and comments from students have been used for the preparation of this new student-friendly

book.
This book is organized into 21 chapters and will supplement most of the organic textbooks on the
market. In all chapters, there are leading questions to focus attention on a principle or reaction and the
answer is immediately provided. The organization of the book provides an initial review of fundamental
principles followed by reactions based on manipulation of functional groups. The intent in all cases is
to provide a focused question about a specifc principle or reaction and the answer immediately follows.
There is also a chapter on spectroscopy as well as chapters on amino acid and peptide chemistry and
carbohydrate and nucleoside chemistry. Each chapter ends with several homework questions for that
chapter, and the answers are provided in an Appendix at the end of the book.
I thank all of the organic chemistry students I taught over the years. They provided the inspiration for the book as well as innumerable suggestions that were invaluable. I thank Ms. Hilary Lafoe
and Ms. Jessica Poile, the Taylor & Francis editors for this book, and also Dr. Fiona Macdonald, the
publisher. This book would not have been possible without their interest in chemistry and their help
as the book was written. I thank Professor John D’Angelo of Alfred University who provided a very
useful and helpful review of the manuscript. I thank PerkinElmer who provided a gift of ChemDraw
Professional (Version 18.0.0.231[4318]). All the reactions and fgures were done with ChemDraw except
for those images that use molecular models and the artist-rendered drawings. All molecular models
were rendered with Spartan’18 software and I thank Warren Hehre and Sean Ohlinger of Wavefunction,
ix


x

Preface

Inc., who provided a gift of Spartan’18 software, version 1.2.0 (181121). I thank Ms. Christine Elder
(), graphics design artist, for her graphic arts expertise to render the drawings
on pages 14 (C1), 66 (C5), 93 and 122 (C7), 208 and 209 (C11), 280 (C15), 315 and 342 (C16). Finally, I
thank my wife, Sarah, for her patience and understanding while I was putting this book together.
Where there are errors, I take complete responsibility. Please contact me at
if there are questions, problems, or errors.

Michael B. Smith
Professor Emeritus
December 2019


Common Abbreviations
Other, less common abbreviations are given in the text when the term is used.
O

Ac
AIBN
aq
AIBN
AMP
ATP
Ax
9-BBN
Bn

Acetyl
azobisisobutyronitrile
aqueous
Azobisisobutyronitrile
Adenosine monophosphate
Adenosine triphosphate
axial
9-Borabicyclo[3.3.1]nonane
Benzyl

Boc

Bu
Bz
°C
13C NMR
cat

tert-Butoxycarbonyl
n-Butyl
Benzoyl
Temperature in Degrees Celsius
Carbon Nuclear Magnetic Resonance
Catalytic

Cbz
CIP
mCPBA
DCC
DEA
DMAP

Carbobenzyloxy
Cahn–Ingold–Prelog
3-Chloroperoxybenzoic acid
1,3-Dicyclohexylcarbodiimide
Diethylamine
4-Dimethylaminopyridine

DMF
DMSO
EDTA

ee or % ee
Equiv
Et
Ether
Eq
FDNB
FMO
FVP
GC
h
1H NMR
HMPA
HOMO
hν
IP

N,N'-Dimethylformamide
Dimethyl sulfoxide
Ethylenediaminetetraacetic acid
% Enantiomeric excess
Equivalent(s)
Ethyl
diethyl ether
equatorial
Sanger’s reagent, 1-fuoro-2,4-dinitrobenzene
Frontier molecular orbitals
Flash Vacuum Pyrolysis
Gas chromatography
Hour (hours)
Proton Nuclear Magnetic Resonance

Hexamethylphosphoramide
Highest occupied molecular orbital
Irradiation with light
Ionization potential

Me

-CH2Ph

O

Ot-B
Bu

-CH2CH2CH2CH3

O
CH2 Ph
OC

c-C6H11-N=C=N-c-C6H11
HN(CH2CH3)2
O
H

Me2
NM

-CH2CH3
CH3CH2OCH2C3

xi


xii

Common Abbreviations

iPr
IR
IUPAC
K
LCAO
LDA
LUMO
mcpba
Me
min
MO
mRNA
MS
NMR
N.R.
NAD+
NBS
NCS
Ni(R)
NMO
Nu (Nuc)
PCC

PDC
\PES

Isopropyl
Infrared
International Union of Pure and Applied Chemistry
Temperature in kelvin
Linear combination of atomic orbitals
Lithium diisopropylamide
Lowest unoccupied molecular orbital
meta-Chloroperoxybenzoic acid
Methyl
minutes
Molecular orbital
Messenger ribonucleic acid
Mass spectrometry
nuclear magnetic resonance
No reaction
Nicotinamide adenine dinucleotide
N-Bromosuccinimide
N-Chlorosuccinimide
Raney nickel
N-Methylmorpholine N-oxide
Nucleophile
Pyridinium chlorochromate
Pyridinium dichromate
Photoelectron spectroscopy

Ph
PhMe

PPA
Ppm
Pr

Phenyl
Toluene
Polyphosphoric acid
Parts per million
n-Propyl

Py
RNA
rt
s
(Sia)2BH
sBuLi
SEAr
SET
SNAr
SOMO
T
t-Bu
TBHP (t-BuOOH)
TFA
ThexBH2
THF
TMEDA
Tol
Ts(Tos)
UV

VIS
VDW

Pyridine
Ribonucleic acid
Room temperature
seconds
Disiamylborane (Siamyl is sec-Isoamyl)
sec-Butyllithium
Electrophilic aromatic substitution
Single electron transfer
Nucleophilic aromatic substitution
singly occupied molecular orbital
Temperature
tert-Butyl
t-Butylhydroperoxide
Trifuoroacetic acid
Thexylborane (tert-hexylborane)
Tetrahydrofuran
Tetramethylethylenediamine
Tolyl
Tosyl = p-Toluenesulfonyl
Ultraviolet spectroscopy
visible
van der Waals

-CH(Me)2

LiN(i-Pr)2
-CH3 or Me

-CH2CH2CH3
N

CH3CH2CH(Li)CH3

-CMe3
Me3COOH
CF3COOH
Me2NCH2CH2NMe2
4-(Me)C6H4
4-(Me)C6H4SO2


Author
Professor Michael B. Smith was born in Detroit, Michigan, and moved to Madison
Heights, Virginia, in 1957. He graduated from Amherst County High School in
1964. He worked at Old Dominion Box Factory for a year and then began studies
at Ferrum Junior College in 1965. He graduated in 1967 with an AA and began
studies at Virginia Tech later that year, graduating with a BS in Chemistry in
1969. He worked as a chemist at the Newport News Shipbuilding & Dry Dock
Co., Newport News, Virginia, from 1969 until 1972. In 1972, he began studies in graduate school at
Purdue University in West Lafayette, Indiana, working with Professor Joseph Wolinsky, graduating in
1977 with a PhD in Organic Chemistry. He took a postdoctoral position at Arizona State University in
Tempe, Arizona, working on the isolation of anti-cancer agents from marine animals with Prof. Bob
Pettit. After one year, he took another postdoctoral position at MIT in Cambridge, Massachusetts, working on the synthesis of the anti-cancer drug bleomycin with Prof. Sidney Hecht.
Professor Smith began his independent career as an assistant professor in the Chemistry department
at the University of Connecticut, Storrs, Connecticut, in 1979. He received tenure in 1986, and spent six
months on sabbatical in Belgium with Professor Leon Ghosez at the Université Catholique de Louvain
in Louvain la Neuve, Belgium. He was promoted to full professor in 1994 and spent his entire career at

UConn. Professor Smith’s research involved the synthesis of biologically interesting molecules. His most
recent work involved the preparation of functionalized indocyanine dyes for the detection of hypoxic
cancerous tumors (breast cancer), and also the synthesis of infammatory lipids derived from the dental pathogen, Porphyromonas gingivalis. He has published 26 books including Organic Chemistry: An
Acid-Base Approach, 2nd edition (Taylor & Francis), the 5th–8th edition of March’s Advanced Organic
Chemistry (Wiley), and Organic Synthesis, 4th edition (Elsevier), which is the winner of a 2018 Texty
Award. Professor Smith has published 96 peer-reviewed research papers and retired from UCONN in
January of 2017.

xiii



Part A

A Q&A Approach
to Organic Chemistry
What is organic chemistry?
Organic chemistry is the science that studies molecules containing the element carbon. Carbon can
form bonds to other carbon atoms or to a variety of atoms in the periodic table. The most common
bonds observed in an organic chemistry course are C—C, C—H, C—O, C—N, C—halogen (Cl, Br, I),
C—Mg, C—B, C—Li, C—S, and C—P.



1
Orbitals and Bonding
This chapter will introduce the carbon atom and the covalent bonds that join carbon atoms together in
organic molecules. The most fundamental properties of atoms and of covalent bonds will be introduced,
including hybridization, electronic structure, and a brief introduction to using molecular orbital theory
for bonding.

1.1 ORBITALS
1.1.1 Atomic Orbitals
What is the electronic structure of an atom?
A given atom has a fxed number of protons, neutrons, and electrons, and the protons and neutrons are
found in the nucleus. The electrons are located at discreet energy levels (quanta) away from the nucleus.
The nucleus is electrically positive, and electrons are negatively charged.
What is the Schrödinger wave equation?
The Schrödinger equation, Hψ = Eψ, is a linear partial differential equation that describes the wavefunction or state function of a quantum-mechanical system. The motion of an electron is expressed by a wave
equation, which has a series of solutions and each solution is called a wavefunction. Each electron may
be described by a wavefunction whose magnitude varies from point to point in space. The equation is a
partial differential equation that describes how the wavefunction of a physical system changes over time.
What are atomic orbitals?
An atomic orbital is a mathematical function that describes the wave-like behavior of either one electron
or a pair of electrons in an atom. If certain simplifying assumptions are made, it is possible to use the
Schrödinger equation to generate a different wavefunction for electrons with differing energies relative
to the nucleus. A particular solution to the so-called Schrödinger wave equation, for a given type of electron, is determined from the Schrödinger equation, and a solution for various values of ψ that correspond
to different energies shows the relationship between orbitals and the energy of an electron. The wavefunction is described by spatial coordinates ψ(x,y,z), and using Cartesian coordinates a point is defned
that describes the position of the electron in space.
What is a node?
A node is derived from a solution to the Schrödinger equation where the wavefunction changes phase,
and it is taken to be a point of zero electron density.
What is the Heisenberg uncertainty principle?
The Heisenberg uncertainty principle states that the position and momentum of an electron cannot be
simultaneously specifed so it is only possible to determine the probability that an electron will be found
at a particular point relative to the nucleus. The probability of fnding the electron in a unit volume of
three-dimensional space is given by |ψ(x,y,z)|2, or |ψ|2dτ, which is the probability of an electron being in
a small element of the volume dτ. This small volume can be viewed as a charge cloud if it contains an
3

4

A Q&A Approach to Organic Chemistry

electron, and the charge cloud represents the region of space where we are most likely to fnd the electron
in terms of the (x,y,z) coordinates. These charge clouds are orbitals.
What is a s-orbital?
Different orbitals are described by their distance from the nucleus, which formally corresponds to the
energy required to “hold” the electron. One solution to the Schrödinger equation is symmetrical in that
the wave does not change phase (zero nodes; a node is the point at which the wave changes its phase).
This corresponds to the frst quantum level and known as a s-atomic orbital. The 1s-orbital represents
the frst energetically favorable level where electrons can be held by the nucleus. The space in which
the electron may be found is spherically symmetrical in three-dimensional space. All spherically symmetrical orbitals are referred to as s-orbitals. The nucleus is represented by the “dot” in the middle of
the sphere.
Nucleus
s-Orbital

What is a p-orbital?
When the solution for the Schrödinger equation has one node (the wave changes phase once), electron
density is found in two regions relative to the node. When the space occupied by this electron is shown
in three-dimensional space, it is a p-orbital with a “dumbbell” shape. In the (x,y,z) coordinate system,
the single node could be in the x, the y, or the z plane and all three are equally likely. Therefore, three
identical p-orbitals must be described: px, py, and pz relative to the nucleus, as shown. Identifcation of
three identical p-orbitals means that there are three p-wavefunctions that describe three electrons that
are found at the same energy.
Nucleus

p-orbital

pxpypz-Orbital

What is a degenerate orbital?
Orbitals with identical energies are said to be degenerate, and the three 2p-orbitals shown in the preceding question are degenerate orbitals. The three electrons in different orbital lobes have the same energy
and have the same charge. Due to the presence of like charges, the orbital lobes repel and will assume
positions as far apart as possible in a tri-coordinate system. In other words, the three orbitals will be
directed to the x-, y-, and z-directions in a three-dimensional coordinate system as shown.
Do the electrons in a p-orbital migrate from one lobe to the other?
No! The picture of the p-orbital represents the uncertainty of where to fnd the electrons. The diagram
shows an equal probability of fnding the electrons in each of the three dumbbell-shaped orbitals, above
and below a node, which is taken as a point of zero electron density and corresponds to the position of
the nucleus in the diagram. Therefore, the electrons are found in the entire p-orbital (both lobes), and the
diagram simply indicates the uncertainty of their exact location.
How many orbitals are there in each valence shell?
Each orbital can hold two electrons. For the frst valence shell containing H and He, there is one s-orbital.
For the next valence shell (containing B, C, N, O, F), there is one 2s orbital, but three 2p-orbitals. The
2p-orbitals have different spatial orientations, correlated with the x, y and z axes of a three-dimensional
coordinate system. In other words, the three p-orbitals are px, py, and pz.


5

Orbitals and Bonding

1.1.2 Electron Configuration
What is electron confguration?
The electron confguration is the distribution of electrons of an atom or molecule in atomic or molecular
orbitals. Electrons are distributed in shells, each of which has different types of electrons: s, p, d, f. Each
orbital (energy level) occurs further from the nucleus; the electrons are held less tightly. Each orbital can
hold a maximum of two electrons and each energy level will contain different numbers of electrons (one

electron for the 1s1 and two electrons for the 1s2 orbital, as shown. There are six electrons for p-orbitals;
two each is possible for each of the three-degenerate p-orbitals. There are ten electrons for d-orbitals;
two each for the fve d-orbitals. Orbitals will fll from lowest energy to highest energy orbital, according
to the order shown in the mnemonic for the electronic flling order of orbitals.
Filling Order

H

1s1

He

1s2

1s

2s

2p

3s

3p

4s

3d

4p

5s

4d

5p 6s

4f

5d

6p

What is the Aufbau principle?
Orbitals “fll” according to the Aufbau principle. The principle states that each orbital in a sublevel s,
p, or d will contain one electron before any contains two. Orbitals containing two electrons will have
opposite spin quantum numbers (they are said to be spin paired, ↑↓). An example is helium in the preceding question.
What is the order in which the three degenerate p-orbitals fll with electrons through the 2p
level? Ignore the 1s and 2s levels.
The order for the 2p-orbitals will be 2px→2py→2pz→2px→2py→2pz:
↑ _ _ →↑↑_ → ↑↑↑ → ↑↓↑↑ → ↑↓↑↓↑ → ↑↓↑↓↑↓

1.1.3 Molecular Orbitals
What is the difference between a molecular orbital and an atomic orbital?
An atomic orbital is a mathematical function that describes the wave-like behavior of either one electron
or a pair of electrons in an atom. A molecular orbital (MO) is a mathematical function describing the
wave-like behavior of an electron in a molecule.
An atomic orbital is associated with a specifc atom. The electrons found on an individual atom of an
element are in atomic orbitals whereas the electrons found in an atom that is part of a covalent bond are
in molecular orbitals. A molecular orbital is formed once two atoms are joined in a covalent bond. Much
of the electron density is shared between the two nuclei of the two atoms rather than being exclusively on

the nuclei of the two atoms. This energy level for the electrons found in the molecular orbital is different
from electrons that are on an individual atom such as that found in an element.
What is the Linear Combination of Atomic Orbitals (LCAO) model?
The LCAO model is the superposition of atomic orbitals that constitutes a technique for calculating
molecular orbitals in quantum chemistry. The LCAO model is a mathematical model that is used to mix
the atomic orbitals of two atoms to get new orbitals for the resulting bond between those two orbitals.
In the LCAO method, the atomic orbitals of each “free” atom are mixed to form molecular orbitals. The
model requires that there can be no more or no less orbitals and no more or no less electrons in the orbitals for the new bond than are found in the atomic orbitals for the two atoms. These new orbitals must be


6

A Q&A Approach to Organic Chemistry

of a different energy than the atomic orbitals in a non-degenerate system. In other words, when mixing
two atomic orbitals, one new orbital is formed that must be lower in energy and one is formed that is
higher in energy relative to the atomic orbitals.
How does the LCAO model apply to covalent bonds in simple diatomic molecules such as
hydrogen?
Two atoms are combined to form a covalent bond, and the atomic orbitals of each atom are combined to
form a molecular orbital. Assume that the electrons in each atomic orbital are transferred from energy
levels near the atom to different energy levels that correspond to electron density between the nuclei of
the bonded atoms. A molecular orbital is an orbital associated with the molecule rather than the individual atoms, as shown below for H2. For molecules containing more electrons than hydrogen or helium,
and for those containing electrons in orbitals other than s-orbitals, the diagram is more complex and the
LCAO approach usually fails.
How are molecular orbitals formed from atomic orbitals?
Using the LCAO model, the orbitals for the molecule diatomic hydrogen (H2) can be formed by mixing the atomic orbitals of two hydrogen atoms. The orbitals formed are not atomic orbitals, but they are
associated with a molecule, in this case H2, and are called molecular orbitals. Each of the two hydrogen
1s atomic orbitals (H atomic orbital) contains one electron, and these orbitals have the same energy.
When mixed to form the molecular orbital, the molecular orbital electrons have a different energy, and

those orbitals are in a different position relative to atomic orbitals, as shown for the molecule H—H.
Therefore, if the two atomic orbitals mix, two molecular orbitals are generated, one higher in energy and
one lower than the original atomic orbitals. It is noted that this model does not work well for atoms that
have degenerate p-orbitals.
Anti-bonding molecular orbital

Increasing Energy

H

H
1s1

1s1

Bonding molecular orbital

1.2 BONDING
1.2.1 Ionic Bonding
What is a Lewis dot structure?
A Lewis electron dot formula generates a bond between two atoms by simply using dots for electrons for
the two electrons that comprise a bond. In other words, each bond is represented by two dots between the
appropriate atoms, and unshared electrons are indicated by dots (one or two) on the appropriate atom.
What is the Lewis dot structure of lithium fuoride? Add the charges!

Li

F

7

Orbitals and Bonding
What is an ionic bond?

An ionic bond occurs when two atoms are held together by electrostatic forces, where one atom or group
assumes a positive charge and the other atom or group assumes a negative charge. Sodium chloride
(NaCl), for example, exists in the solid state as Na+Cl–.

Na Cl
Why does sodium assume a positive charge in NaCl?
If the valence electrons associated with each atom are represented as dots (one dot for each electron), the
structure for NaCl will be that shown above. Sodium chloride has an ionic bond, and in an ionic bond
all of the electrons are on chlorine and none are on sodium. Sodium (Na) is in Group 1 and has the electronic confguration 1s22s22p63s1. If one assumes that the sodium atom can react, it can either lose one
electron (ionization potential) or gain seven electrons (the ability to gain one electron is called electron
affnity) in order to achieve a “flled” shell. The loss of one electron gives the electronic confguration
1s22s22p6, which is a flled shell and very stable, and requires much less energy than gaining seven to give
another flled shell. After transfer of one electron, sodium has no electrons around it. In other words, it
is Na+, which is missing one electron relative to atomic sodium; this means it is electron defcient and so
assumes a positive charge (see formal charge in Section 2.5).
Why does chlorine assume a negative change in NaCl?
In the ionic bond for NaCl, the chlorine atom has eight electrons around it. The chlorine (Cl) atom
has the electronic confguration 1s22s22p63s23p5. If one imagines that the Cl atom reacts, and since
chlorine is in Group 17 with seven electrons in the outmost shell, it can either gain one electron or lose
seven electrons. Clearly, the loss of seven electrons will require a great deal of energy. Energetically,
it is far easier for Cl to gain an electron, leading to formation of a negatively charged atom. Therefore,
Cl gains an electron in contrast to Na, which loses an electron. With an excess of electrons, given that
electrons carry a negative charge, the Cl will take a negative charge. The strong electrostatic attraction
between the positive sodium and the negatively charged chlorine binds the two atoms together into an
ionic bond.

1.2.2 Covalent Bonding
What is a covalent bond?
A covalent bond has two electrons that are shared between two atoms. In the case of hydrogen (H2), the
covalent bond can be represented as H:H or H—H, where the (:) or the (—) indicates the presence of two
electrons. In a covalent bond, the bulk of the electron density is localized between the hydrogen nuclei.
This type of bond usually occurs when the atom cannot easily gain or lose electrons. Another way to
view this is that there is a small electronegativity difference between atoms. A model of fuorine (F—F
or F2) shows the electron density around both atoms, but signifcant electron density is clearly between
the two fuorine nuclei that represent the covalent F—F bond.


8

A Q&A Approach to Organic Chemistry

What is the Lewis dot structure of diatomic hydrogen?

H:H
What is the octet rule?
The octet rule states that every atom wants to have eight valence electrons in its outermost electron shell.
What is valence?
Valence is the number of bonds an atom can form to satisfy the octet rule and remain electrically neutral.
Valence is not to be confused with valence electrons, which are the number of electrons in the outermost
shell. In the second row from C to F, the valence is (8 – the last digit of the group number): C: 8 – 4, or
4; N: 8 – 5, or 3; O: 8 – 6, or 2; F: 8 – 7, or 1. Boron is an exception. There are only three electrons and,
therefore, boron can form no more than three covalent bonds and remain neutral. In other words, an atom
can form only as many bonds as there are electrons available to share. Note that the valence of boron is
three and it is electron defcient.
What is the Lewis dot structure of a carbon–carbon bond when drawn as a covalent bond but

ignoring all other electrons and the other valences of each carbon?

C:C
What is a Lewis acid?
A Lewis acid is any substance that can accept a pair of nonbonding electrons, so it is an electron-pair
acceptor. An example is boron, which has only three electrons in the outermost shell, can only form
three covalent bonds but is electron defcient because it requires two more electrons to satisfy the octet
rule. These two electrons are gained by reaction with an electron-rich molecule and trivalent boron compounds are Lewis acids.
What is a Lewis base?
A Lewis base is any substance that can donate a pair of nonbonding electrons, so it is an electron-pair
donor. Electron donation to a hydrogen atom is not included in the Lewis base defnition, however. In other
words, a base that donates two electrons to a hydrogen atom is a Brønsted–Lowry base not a Lewis base.
Why does carbon have a valence of four?
Carbon is in Group 14 so there are four valence electrons. Therefore, carbon forms a total of four covalent bonds by sharing the electrons with many other atoms, including another carbon atom. With carbon
(C: 1s22s22p2), there are four electrons in the highest valence shell. The gain of four electrons or the loss
of four electrons would require a prohibitively high amount of energy. In a thought experiment, assume
that a carbon atom can form bonds directly with up to four other atoms. In such an experiment, carbon,
because it is lower in energy, will form covalent bonds to share electrons with another atom rather than
“donate” or “accept” four electrons to form an ionic bond. A carbon atom with appropriate functionality
attached can undergo chemical reactions with other molecules to form covalent bonds to other carbon
atoms, to hydrogen atoms, as well as to many atoms in the periodic table.
What is covalent bond?
In a covalent bond, the electrons are mutually shared between two nuclei in that bond, so each nucleus has
a flled shell (eight in the case of carbon and two in the case of hydrogen). The most common way to show


Orbitals and Bonding

9

mutual sharing of electrons for two carbon atoms is to draw a single line between the two atoms (C—C) rather
than using the Lewis dot structure, C:C. The two electrons are equally distributed between the two carbon
atoms and the resultant bond has a symmetrical distribution of electron density between the two atoms.
What does a covalent bond between two hydrogen atoms look like in the molecule H2?
The two hydrogen atoms are identical, the mutual sharing of electrons leads to a symmetrical distribution
of the electron density between the two hydrogen nuclei, as shown in the accompanying molecular model
(an electronic potential map)

What does a covalent bond between two carbon atoms with identical atoms attached?
If the two carbon atoms are identical, the mutual sharing of electrons leads to a symmetrical distribution
of the electron density between the two carbon nuclei.
What is electronegativity?
Electronegativity is a measure of the attraction that an atom has for the bonding pair of electrons in a
covalent bond. A more electronegative atom will attract more electron density toward itself than a less
electronegative atom.
What is a polar covalent bond?
If two atoms are part of a covalent bond, and one atom is more electronegative than the other, the shared
electron density is distorted toward the more electronegative atom, as shown in the molecular model of
H—F, rather than the symmetrical distribution found in H—H. The shaded area on the far right indicates
higher electron density, which is clearly on the more electronegative fuorine atom.

When a bond is formed between two atoms that are not identical, the electrons do not have to be equally
shared. If one atom is more electronegative (electronegativity is the ability of an atom to attract electrons
to itself; the electronegativity is higher), it will pull a greater share of electrons from the covalent bond
toward itself. The larger the difference in electronegativity, the greater the distortion of electron density
in the covalent bond toward the more electronegative atom.


10

A Q&A Approach to Organic Chemistry

What is an example of a molecule with a polar covalent bond?
An example is H—F (see the molecular model in the preceding question), where the fuorine is signifcantly
more electronegative than the hydrogen atom. This difference in electronegativity will lead to electron distortion in the covalent bond toward the fuorine, away from hydrogen, and an unsymmetrical covalent bond
will form that has less electron density between the nuclei. In other words, it is a weaker bond. The model
shown indicates this electron distortion, with the shaded area on the far right (higher electron density)
toward the fuorine and the shaded area on the far left (lower electron density) toward the hydrogen atom.
How can HF be drawn to represent the polarized covalent bond?
In a polarized bond such as that found in HF, fuorine is more electronegative, and the bond density is
distorted such that there is more electron density on fuorine relative to the hydrogen. The more electronegative atom will be “more negative” and the other will be “more positive.” The molecule is neutral so
there are no ionic charges, but the distortion of electron density in the polarized covalent bond is represented by a “partial charge” (δ+) at the atom with the least electron density and (δ-) at the atom with the
most electron density. Therefore, the polarization of HF is represented as δ+H—Fδ–.
Can the dipole of a polarized covalent bond be represented by an appropriate symbol?
A common way to represent this distortion of electrons is with the symbol +⟶, with the + representing
the positive atom and ⟶ representing the direction of electron fow. As noted in the previous question,
it is perhaps more common to use a δ+ at the atom with the least electron density and δ– at the atom with
the most electron density, as shown for H—F. Such a covalent bond is polarized, and this disparity in
electron density leads to a dipole moment. Any covalent bond between two atoms where one is less electronegative, and the other is more electronegative, will be a polar covalent bond.
What is dipole moment?
Bond dipole moment is a measure of the polarity of a chemical bond, generally induced by differences
in electronegativity of the two atoms in that bond. The bond dipole symbol is μ and the unit of measurement is the Debye (D).
Is the C—H bond considered to be polarized?
No! Although C and H have different electronegativities (H = 2.1 and C = 2.5 on the Pauling electronegativity scale), the C—H bond is not considered to be polarized. This assumption is based on the polarity
of molecules containing only C—H bonds, but the chemical reactivity of molecules that contain only
C—H bond will support this view.
What is a heteroatom?
A heteroatom is defned as any atom other than carbon or hydrogen. Examples are O, N, S, P, Cl, Br, F,
Mg, Na, etc.
Which of the following are polar covalent bonds? For the polarized bonds, identify the negative

and positive poles: (a) C—O (b) C—C (c) O—H (d) H—H (e) Br—Br (f) C—N (g) N—N (h)
O—O (i) H—Br (j) NaCl
Only those bonds between dissimilar atoms will be polarized, therefore (a), (c), (f), (i), and (j) are polarized covalent. NaCl (j) is an ionic bond. The negative poles will be oxygen in (a) and (c), nitrogen in (f),
and bromine in (i). The positive poles are carbon in (a) and (f), hydrogen in (c) and (i).
What is van der Waals attraction?
When there are no polarizing atoms in the molecules, the only attraction between molecules results from
the electrons of one molecule being attracted to the positive nuclei of atoms in another molecule. This
interaction is known as van der Waals attraction (sometimes called London forces).