SPRING 2008 CH221 ORGANIC CHEMISTRY 1
NOTES FOR CHAPTER 1
BONDING AND STRUCTURE
What is Organic Chemistry?
Originally it was thought that compounds like amino acids, sugars, urea and
many others (called organic compounds) isolated from living things, were
fundamentally different from those in nonliving things, called inorganic
compounds. It was believed that organic compounds could not be
synthesized in the laboratory from inorganic starting materials, until Wöhler
made urea (molecular formula: CH4N2O) – a well-known organic compound –
from inorganic compounds in 1828.
It was soon discovered that all organic substances contain the element carbon
and it is now evident that all organic molecules possess at least one carbon
atom and, indeed, the huge majority possess more than one carbon atom. So,
organic chemistry can be described as the study of the chemistry of
carbon compounds. This definition holds irrespective of whether the
compound is natural or not, although a great deal of the modern emphasis of
organic chemistry is on natural products and the relationships between organic
chemistry and life sciences and medical sciences are strong ones.
Nowadays (2008), there are more than 20 million organic compounds in the
CAS registry of the American Chemical Society and thousands of new ones
are synthesized each week in laboratories all over the world. In fact, the
compounds of carbon easily outnumber the compounds of all other elements.
The reasons why carbon is able to form such a large number of compounds,
with a vast range of size and complexity, are to be found in the position of
carbon in the periodic table, which is discussed later. Carbon is a second row
(period) element: it is the first element of group 4A. As such, it is endowed
with the ability to form four strong covalent bonds, to form long chains
and rings of many sizes, and to form multiple bonds. These aspects are
fundamental to organic chemistry, as will be demonstrated many times during
this course.
Organic molecules range from the simplest, such as methane, ethanol and
trichlorofluoromethane through compounds of intermediate complexity to
extremely complex ones, such as palytoxin (see below).
H
H
C
H
C
H
H
H
C
C
H
H
C
Cl
C
H
C
H
H
H
Methane
OH
Cl
C
F
C
H
C
H
Cl
Ethanol
H
Trichlorofluoromethane
Benzene, as a shorthand
drawing
Benzene
Symbol and line drawings (similar to
Lewis diagrams)
H
O
H
HO
C
NH2
O
H
C
S
NH
H H
CH3
HO
N
CH3
O
C
H
Amoxicillin, an antibiotic
of the penicillin family
N
O
H
O
N
CH3
C
H
H H
O
HO
H
H
N3
H
AZT (azidodeoxythymidine),
anti-HIV drug
As shorthand drawings only
Modern organic chemists are extremely skilled at synthesizing new organic
chemicals in the laboratory, even those that are identical to complex natural
products, such as palytoxin:
Medicines, polymers, dyestuffs, food additives, agrochemicals, including
pesticides, and a huge range of other substances are prepared (at least
initially) in the laboratory. Organic chemistry touches the lives of everyone –
from the good (for example, medicines and polymers) through the bad (for
example, pollution) to the ugly (for example, nerve gases).
Organic chemistry as a subject interacts with many other subject areas, within
and without the general subject of chemistry: examples include physical and
computational chemistry (e.g. structure, reactivity, molecular design), inorganic
chemistry (organometallic compounds), biochemistry and medicine (e.g.
metabolic
agricultural
pathways,
drug
design,
drug
science
(e.g.
pesticides),
metabolism),
food
science
pharmacology,
and
geology
(hydrocarbons).
Some of these relationships are admirably summarized by Homer Simpson:
Before considering organic molecules in detail, it is necessary to review
important structural and bonding features that have been covered in previous
general chemistry courses. Most of Chapter 1 and some of Chapter 2 are
dedicated to this.
1.1 The Periodic Table
Atoms, Ions and Isotopes
The atomic theory of matter is well-established: a schematic of an atom of
carbon-12 (the main isotope of carbon, symbol
12 C)
6
is used to illustrate
important features.
Nucleus [6 protons (+ve) and 6 neutrons (neutral]
Electron cloud (comprises most of volume),
equivalent to 6 electrons (-ve)
The number of protons in the nucleus, or the number of electrons in the neutral
atom is called the atomic number (Z) and this defines the element by its position
in the periodic table.
In addition, charged species, called ions are encountered in organic chemistry:
A cation is positively charged and has fewer electrons than its neutral
form.
An anion is negatively charged and has more electrons than its neutral
form.
The number of neutrons in the nucleus of a particular element can vary.
Isotopes are two atoms of the same element having different numbers of
neutrons: they have the same atomic number (Z) but different mass numbers (A).
For many elements, the natural abundance one particular isotope (e.g.
higher than that of the others (e.g.
13 C
6
and the radioactive
number is often dropped in naming isotopes (e.g.
12C, 13C
12 C)
6
is
14 C).
6
The atomic
14C).
The atomic
and
weight of an element is the weighted average of all isotopes of the element,
reported in atomic mass units (amu), based on the scale with
12C
= 12 exactly.
This is all summarized in the table below.
Element
Isotopes
(%
natural Atomic
abundance)
weight
(to
maximum of 4 decimal
places)
Hydrogen
1H
(99.98)
deuterium,
2H
(0.02) (= 1.0079
3H
D),
(v.
small) (= tritium, T)
Carbon
12C
(98.9)
13C
(1.1),
14C
12.011
(v. small)
Nitrogen
14N
(99.63)
Oxygen
16O
(99.8)
18O
15N
(0.37)
17O
14.0067
(0.037) 15.9994
(0.20)
The elements are arranged in order of atomic number in the periodic table, which
consists of rows or periods (horizontal) and columns or groups (vertical). Each
group and row is named, but of the 100 or more known elements, only some that
belong to the first three rows are commonly found in organic molecules. This is
summarized in the diagram below (see inside front textbook cover for full periodic
table).
Two useful features of the periodic table are:
Elements in the same row are similar in size (atomic radius decreases
somewhat on traversing a row LR)
Elements
in
the
same
column
(group)
have
similar
electronic
configurations and chemical properties
Across each row of the periodic table, electrons are added to a particular shell of
orbitals around the nucleus. The shells are numbered 1, 2, 3 and so on, in the
manner of rows or periods, so that adding electrons to the first and second shells
forms the first and second rows, and so on. Electrons are first added to the shells
closest to the nucleus, where they are held most tightly: these are the core
electrons and they contribute little to chemical reactivity.
Each shell consists of a fixed number of sub-shells known as orbitals.
An orbital is most conveniently described as a region (volume) of space
that is high in electron density.
In the first and second rows (which are most important in organic chemistry), only
s orbitals (one type) and p orbitals (3 types) exist. The first row has one s orbital
(1s) only and so the electronic configurations of the first two elements (H and He)
are 1s1 and 1s2, respectively. The second row has a single 2s orbital and 3 x 2p
orbitals, making a total of 8 elements (Li, Be, B, C, N, O, F, Ne), with
configurations 1s22s1, through 1s22s2sp2 for C, to 1s22s22p6. The orbitals
associated with the second row elements are shown below.
The outermost electrons (e.g. 2s orbital of Li) are the ones that participate
in reactions: they are known as valence electrons.
1.2 Bonding
Bonding is the joining of two atoms (of the same or different elements) in a
stable arrangement: it is a favorable process because it always leads to
lowered energy and increased stability.
The bonding of two or more elements gives rise to compounds and in general
allows atoms to attain a complete outer shell of electrons, which coincides with
noble (inert) gas electronic configurations. Thus H (1s1) will achieve the He
configuration (1s2) by bonding and the second row elements (1s22sm2pn) will
achieve (with some exceptions) the configuration of He (1s 2) or Ne (1s22s22p6),
depending on their positions in the row.
There are two extreme types of bonding: ionic bonding and covalent bonding.
Ionic bonds result from the transfer of electrons from one atom to another.
Covalent bonds result from the sharing of electrons between two nuclei.
Generally (there are many exceptions), ionic bonds are formed when elements
on the far left combine with those on the far right of the periodic table. An
example is the formation of lithium fluoride:
Li
.
Li+ +
1s22s1
..
: ..F . +
1s22s22p7
e-
1s2 ([He])
Cation
e-
Li+F- ionic compound
.. _
: ..F ..
1s22s22p8 ([Ne])
Anion
Ions are held together by strong electrostatic forces, usually in crystal lattices:
ionic compounds are commonly known as salts. See below for a good example.
The second kind of bonding, covalent bonding, most likely occurs between
elements like carbon in the middle of the periodic table or between two
elements like hydrogen and chlorine from the same side of the periodic
table.
Most covalent bonds are two-electron bonds, and a compound with covalent
bonds is called a molecule.
Hydrogen is limited to forming one covalent bond (it is monovalent):
A second-row element in a molecule has a valence that depends on its
position in the row (i.e. its outer electronic configuration), but can have no
more than 8 valence (outer) electrons – known as a complete octet.
Atoms with up to 3 valence electrons will form up to 3 bonds, but atoms with 4 or
more valence electrons will form enough bonds to give a complete octet. This
results in the following formula for predicting the valence.
These guidelines are used to summarize the usual number of bonds formed by
the common atoms in organic molecules. Note that second-row elements that
form fewer than 4 bonds will have both bonding electrons and nonbonding (lone
pair) electrons. This summary can be used to speed up the drawing of Lewis
structures (next).
1.3 Lewis Structures
G.N. Lewis
Lewis structures are electron dot representations of molecules. The three general
rules for drawing Lewis structures are
1. Draw only the valence electrons.
2. Give every second-row element an octet of electrons, if possible.
3. Give each hydrogen 2 electrons.
A simple example is given by the HF molecule.
However, drawing Lewis structures for organic molecules is a little bit different to
drawing inorganic structures, as done in a typical general chemistry program,
because, although carbon is central to the structure, there is often a large
number of hydrogen atoms and often other atoms (like O, N, S, etc – called
heteroatoms) to include in the structure.
This is illustrated by sample problem 1.2 below, showing how lone pair electrons
are accommodated into a Lewis diagram.
Multiple bonds can be incorporated easily into Lewis structures, as shown by
sample problem 1.3:
Formal Charge
Formal charge is the charge assigned to individual atoms in a Lewis structure: it
is a measure of how the number of electrons around an atom compares with its
number of valence electrons.
Formal charge =
_
Number of
valence electrons
= Group number _
Number of
electrons "owned"
Number of
nonbonded electrons
_
Half number of
bond electrons
This is illustrated by sample problem 1.4.
A shortcut method to formal charge assignment is given in the table below, for
common bonding arrangements in organic molecules.
1.4 Lewis Structures Continued
Isomers
It is usually possible to draw more than one suitable Lewis structure for a given
molecular formula, as illustrated by C2H6O: the two valid Lewis structures
represent two different well-known molecules, ethanol and dimethyl ether. They
are known as isomers.
H
H
H
C
C
H
H
H
..
O
..
H
H
C
H
..
O
..
H
Ethanol
C
H
H
Dimethyl ether
Isomers are different molecules having the same molecular formula
The examples above are constitutional isomers; their molecular structures
differ in connectivity or atom bonding arrangement. Later, many other examples
of constitutional isomerism will be seen and Chapters 4 and 5 deal with the other
main branch of isomerism: stereoisomerism.
Exceptions to the Octet Rule
Odd-electron molecules (free radicals), molecules of early second-row elements
(especially Be and B), and high valence molecules of third row elements and
beyond all provide plenty of examples where the octet rule is broken: the central
atom has either fewer or more than an octet of outer electrons. Some of these
molecules are shown below.
..
:F:
Octet-deficient molecules
H
Be
H
..
: ..F
(gas phase)
4eBeryllium hydride
S
..
CH3
10eDimethyl sulfoxide
.
:N
..
F:
..
6eBoron trifluoride
Octet-excessive molecules
..
:O
CH3
B
7eNitric oxide
(a free radical)
..
: O: :OH : O :
..
O:
..
HO
..
S
..
O:
..
OH
..
:O
.. 12e
Sulfuric acid
..
HO
..
P
C
P
..
OH
..
CH2 :OH
10e- : OH
..
..
..
CH2CH2NH2
Alendronic acid
1.5 Resonance
Many molecules cannot be adequately represented by a single Lewis structure.
Their true structures (those that are in accord with physical and chemical data)
are composites of the various Lewis structures.
This phenomenon is known as resonance, the Lewis structures are known
as resonance structures and the real (composite) structure is called a
resonance hybrid. The resonance symbol is .
The Lewis structures imply electron localization, whereas resonance allows
electron pairs to be delocalized over the molecule: a situation that leads to
increased stability (lower energy): molecules with two or more resonance
structures are said to be resonance stabilized.
An Introduction to Resonance Theory
The basic principles of resonance theory are given below.
For example A and B (below) are resonance structures, whereas C and D are
isomers.
One H atom in different location
..
..
: OH
O:
One electron pair is in
a different location
CH3
C
+
O:
and
CH3
A
+
C
..
O:
B
Resonance structures
CH3
C
CH3
and
CH2
C
CH3
D
C
Isomers
Drawing Resonance Structures
To draw resonance structures, use the following 3 rules:
Rule 1. Two resonance structures differ in the position of multiple bonds and
nonbonded electrons. The placement of atoms and single bonds are identical.
E.g. The formamide anion
Rule 2. The resonance structures must have the same number of unpaired
electrons.
The two structures above comply with this rule (they both have zero unpaired
electrons), but the one below does not, so it is not a valid resonance structure of
the formamide anion.
..
. O:
H
_
..
NH
..
C
.
This structure has 2 unpaired electrons
Rule 3. Resonance structures must be valid Lewis structures: hydrogen is limited
to two electrons and second-row atoms are limited to 8 electrons in their outer
shells.
The structure below breaks this rule, so is not a valid resonance structure of the
formamide anion.
..
O:
_
H C
NH
..
10 electrons around C:
not a valid Lewis structure
Curved arrow notation often helps to show the electron distribution
relationship between resonance structures. A curved arrow always has its
tail at an electron pair (bonded or nonbonded) and its head points the
“new” position.
E.g.
..
O:
H
_
..
NH
..
C
.. _
: O:
H
..
NH
C
CH2
C
+
CH2
+
CH2
H
CH3
..
O
..
+
CH2
CH3
+
O
..
CH2
_
..
CH2
C
O
.. :
C
CH2
H
CH3
CH2
C
_
:O
.. :
CH3
The Resonance Hybrid
The resonance hybrid is the composite of all possible resonance structures: it
better accounts for physical and chemical properties than any of the resonance
structures. However, it cannot be represented by a conventional Lewis structure:
instead, either we have to use an approximate structure (with partial bonds
and/or charges) to account for the inevitable electron delocalization in the hybrid,
or give the major resonance scheme. Some examples are shown below.
The resonance hybrid is more stable than its resonance structures, because
electron density is delocalized over a larger volume. It looks most like the
resonance structure that gives the biggest contribution; the major contributor.
Which resonance structures are major and which are minor depends on several
structural factors, but the most important one is the number of formal charges
and bonds.
The lower the number of formal charges and the greater the number of
bonds in a resonance structure, the greater is its contribution.
When resonance structures are identical in this respect, they contribute equally,
as in the case of the acetate ion, above. Otherwise, there will be unequal
contribution, as in the example below.
..
:O
CH3
C
.. _
:O:
CH3
major
CH3
C CH3
+
minor
The resonance hybrid
looks more like this
Also, resonance structures with like charges on adjacent atoms are unfavourable
(they contribute little) and formal + or – charges are generally preferred on O or
N, rather than C (see problem 1.15a).