1

Bonding and Organic Chemistry

1. Covalent Bonding

In this chapter we will discuss covalent bonding. Ionic and metallic bonding will be covered later.

A covalent bond involves the sharing of one or more pairs of electrons. When the atoms involved in

the bond have the same electronegativity, as in Cl2 or O2, the electrons are shared equally. When the

electronegativities differ, as in HCl or H2O, the electrons are shared unequally and the bond is polar;

that is it has a positive end and a negative end. The more electronegative atom will bear a partial

negative charge and the less electronegative atom will have a partial positive charge. Polarity will be

discussed in greater detail at the end of this chapter.

2. Lewis Dot Structures Covalent molecules are usually drawn showing the valence electrons as dots. Since the use of such

drawings was popularized by G.N. Lewis, the British chemist and Nobel Laureate, the structures are

often called "Lewis structures." Students typically call them "dot structures."

There are two rules which govern the drawing of Lewis structures. First, the structures must use the

right number of electrons. If a structure has more electrons than do its individual atoms, it will have a

negative charge. A structure with fewer electrons will have a positive charge. Second, the electrons are

generally arranged so that each atom has the configuration of a noble gas. This will be two electrons

for hydrogen and eight electrons (an octet) for other atoms. The formation of H2 from H atoms, each of

which has a single valence electron, is:

HH + H H

Similarly the formation of Cl2from Cl atoms, each of which has 7 valence electrons, is as below:

Cl + Cl Cl Cl

For more complicated molecules, it is convenient to draw pairs of electrons as lines. Thus we have

O HH or NH H

H

Sometimes the atoms in a molecule do not have enough electrons to both form the needed bonds and to

give each atom an octet. Such a molecule, for example SO2 or N2, can form a valid structure by using

double bonds. Some examples are:

SO O N N

3. Molecular Orbitals The Lewis structure, with its emphasis on localized electron pairs, is part of a chemical bonding model

called "valence bond theory." Valence bond theory is simple to understand and easy to use. There is

another bonding model which has far more predictive power than does VBT. Unfortunately this model,

known as molecular orbital theory is far more difficult to use and understand.

3

2

Molecular orbital theory is not covered in the AP syllabus and we will discuss it no further. However, it

is used heavily in later chemistry courses. Knowing from the beginning that there is another way to

view bonding may make the eventual introduction of molecular orbital theory less of a shock.

4. Resonance

Sulfur dioxide is noteworthy in that it has two possible structures -- one with a double bond on the left

side and another with a double bond on the right side. One would expect, therefore, that each sulfur

dioxide molecule would contain two non-equivalent bonds -- a single bond and a double bond. Since

double bonds are shorter than single bonds, it is possible to determine experimentally whether this is

true. It is not. The two bonds in a sulfur dioxide molecule are both equal and are intermediate in length

between a single and a double bond. This can be explained using the concept of resonance.

Resonance requires that a molecule such as sulfur dioxide be described by drawing both possible

structures with a double-headed arrow between them. Thus:

SO O SO O

The two structures are called "resonance structures" and the actual structure is intermediate between

them.

Another example of resonance is found in sulfur trioxide, SO3. In this compound there are three

resonance structures, as shown below. Instead of a double bond and two single bonds, sulfur trioxide

contains three equivalent bonds, each of which can be considered to be a 1 ⅓ bond.

O OO OS S

O O

SO O

O

Resonance occurs when a molecule can be drawn in several different ways, each differing only by the

arrangement of electrons. Structures in which the atoms have different arrangements are not resonance

structures. While resonance is a useful concept and is widely used by chemists, it is a misleading term.

"Resonance" implies that the structures resonate. A resonating structure would exist in one form half

the time and in the other form half the time. In fact a molecule which exhibits resonance has the same

structure all the time. We just can't draw it.

Stating this differently, a molecule which exhibits resonance is like a mule. A mule is half horse and

half donkey. But that does not mean that a mule is a horse half the time and a donkey the other half of

the time. Molecular orbital theory, it should be noted, has no need for resonance structures.

5. Odd Electron Molecules - Free Radicals

There are a few cases of stable molecules with an odd number of electrons. An odd-electron molecule

will never have an octet and will always have an unpaired electron. These molecules are therefore

relatively unstable and quite reactive. Odd-electron molecules are generally called "free radicals." An

example of a stable free radical is nitric oxide, shown below. Other resonance structures are possible.

However because of the distribution of formal charge, which is discussed below, this structure is best.

N O

3

C O

N

O

OO

-

6. Incomplete Octets – Lewis Acids

There are a few cases of stable molecules lacking one or more electron pairs. The examples you will

see will almost always be boron or aluminum compounds, for example BCl3 or AlF3.

B

F

FF Compounds with incomplete octets have the ability to accept additional electron pairs, thus making

them Lewis acids. A typical example, neglecting the unshared electron pairs, is:

B

Cl

ClCl

+ N

H

H

H

H

H

H

NB

Cl

Cl

Cl

7. Formal Charge

In many molecules some of the atoms are charged. This fact, upon which many organic reactions are

based, is often explained in terms of something called "formal charge."

Formal charge is the difference between the number of electrons possessed by a simple atom and the

number of electrons "owned" by that same atom in a molecule. In calculating formal charge, bonding

electrons are considered to be shared equally by the two atoms between which they lie. Strictly

speaking, this is not true, since shared electrons will tend to associate more with the more

electronegative atom. But it still works rather well. Note that the sum of the formal charges equals the

charge on the molecule (or ion).

In carbon monoxide, for example, carbon possesses a pair of non-bonding electrons and shares three

pairs of bonding electrons. Thus it has 5 electrons. Make sure that you see this! Since by

itself a carbon atom has 4 valence electrons, carbon has gained an electron and therefore

has a formal charge of -1. Using similar logic we find that an oxygen atom in CO has 5

electrons. Since it starts with 6 electrons, this gives it a formal charge of +1.

Formal charges in resonance structures are determined separately for each structure and then averaged

together. Consider, for example, the nitrate ion. Nitrogen, which normally has 5

electrons, now has a half share in four electron pairs. Thus it has a charge of +1. The

double-bonded oxygen has two unshared pairs and a half share in two bonding pairs.

This gives a total of 6 electrons, which is exactly what oxygen starts with. Thus its

formal charge is zero. The single bonded electrons each have 3 unshared pairs and a

half share of one bonding pair for a total of 7 electrons. Since oxygen starts with 6 valence electrons, it

now has a charge of -1. Thus each of the three resonance structures has the formal charge distribution

shown below:

Note that for each of the three resonance structures the formal charges add up to -1, which is the charge

of a nitrate ion. The actual formal charges are the averages of the charges on the individual resonance

structures. Thus the formal charge on the central nitrogen atom would remain at +1. However, the

formal charge on each oxygen atom would be -⅔.

N

O

O O N

O

OON

O

OO

- -

4

Creating these charge separations requires energy. So a structure in which none of the atoms is charged

will be in a lower energy state than a structure containing charged atom. However nitrate, like many

other molecules, does not have a valid Lewis structure with no formal charges.

As a final example, let us use formal charges to explain why the structure of N2O is NNO and not

NON. First consider the molecule NON. Its three resonance structures have the formal charge

distributions shown below:

(+2) (-2) (-1) (+2) (-1) (-2) (+2)

ON N ON N ON N

Here even the middle structure, which distributes its charge more evenly than do the other two, has a

substantial charge separation. Further, all three structures place a +2 charge on the very electronegative

oxygen atom. This, also, requires energy.

Compare this sorry state with the formal charge distribution available to NNO. Its three resonance

structures are drawn below:

(+1) (-1) (-1) (+1) (-2) (+1) (-1)

NN O NN O NN O

The two structures on the left are good. They have small charge separations and don't place a positive

charge on the oxygen. While the right-hand is energetically less desirable than the others, there is no

need to use it when two low-energy alternatives are available. This simple model does not allow us to

determine the relative importance of each of the three resonance structures. However, if we were to

assume that NNO was an equal mixture of the two left-hand structures, with the structure on the right

having relatively little importance, we would be close to right.

Thus NON, having no low-energy resonance structures, does not exist. NNO does.

8. Molecular Shapes – the VSEPR Model

One of the more elegant parts of bonding theory is the way in which it lets us use a few simple rules to

predict molecular shapes. Predicting shapes is based on the idea that the most stable structure for a

molecule is the one in which the electron pairs on the central atom are as far apart as possible. This

idea is known as the valence shell electron pair repulsion model - V.S.E.P.R.

To use VSEPR theory we must first know the shapes in which each possible number of electron pairs

will arrange themselves in order to minimize their repulsive forces. You should be aware that, for

reasons which will be discussed later in this chapter, double and triple bonds count only as a single

bonding pair.

Two electron pairs will arrange themselves on opposite sides of the central atom, giving a linear

molecule. Three electron pairs will go to the comers of an equilateral triangle, giving a shape which can

be called either trigonal or triangular. Four electron pairs will go to the corners of a tetrahedron. Five

electron pairs will go to the corners of a trigonal bipyramid and six electron pairs will go to the corners

of an octahedron. This is summarized in Table 1.

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Table l. Shapes which Maximize Electron Separation

This would be very simple indeed if we had only to consider bonding pairs. Carbon dioxide (2 bonds)

is linear; sulfur trioxide (3 bonds) is trigonal planar (triangular); methane (4 bonds) is tetrahedral; phos-

phorus pentachloride (5 bonds) is trigonal bipyramidal; sulfur hexafluoride (6 bonds) is octahedral.

However many molecules have unshared electron pairs as well as bonding pairs.

Although the unshared pairs cannot actually be "seen," they exert influence on the

geometry of the molecule. Thus water has its four electron pairs arranged

tetrahedrally around the central atom. But since the molecular shape does not include

the two unshared pairs, the water molecule is considered to be "bent".

Before we continue to make matters more complex, let us deal with an issue of

pronunciation. The structure of PCls is a bipyramid (bi pyr' a mid'). The adjective used to describe a

bipyramid is bipyramidal (bi'pyr a'mid al). The widely used (mis)pronunciation (bi pyr a mid' al) is

wrong!

At this point we must discuss not only molecular shapes but also bond angles. A regular tetrahedron

(e.g. methane) has a bond angle 1 . . - -

1 . .However it is not; 1 . Not only do unshared pairs occupy

space around the central atom, they occupy more space than do bonding pairs. The proper way to

phrase this is to say that unshared pairs repel other electrons more than do bonding pairs, thus pushing

the bonding pairs close together. This is particularly important in molecules where the central atom has

five electron pairs, as is discussed below.

O

H

H

6

The problem with the trigonal bipyramid, the arrangement

which is assumed by five electron pairs, is that it has two sets of

non-equivalent sites -- three equatorial sites and two axial sites.

These are shown on the right.

The question which is raised by the non-equivalent sites is

illustrated by SF4, a molecule which has four bonding pairs and

one unshared pair. Where does the unshared pair go? Is it equatorial or is it axial? Although this is not

obvious, it turns out that the unshared pair(s) is always equatorial. Thus SF4, with its equatorial

unshared pair, is considered to resemble a see saw. Similarly IF3, with its two unshared pairs and three

bonding pairs is said to be "t-shaped", while XeF2 is linear.

S

F

F

F

F

I

F

F

F

Xe

F

F

See-saw t-shaped linear A complete table of possible electron arrangements along with the associated shapes is given below.

Axial

EquitorialEquitorial

EquitorialAxial

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9. Multiple Bonds – Sigma and Pi Bonding

As was mentioned previously, a multiple bond counts the same as a single bond when determining

molecular shapes. This is because the multiple bond is composed of two different types of bonds –

referred to as sigma and pi.

It was previously stated that covalent bonds result from the sharing of electron pairs. But bonds can

also be described as the result of overlapping atomic orbitals. It is the way in which the orbitals overlap

which determines whether the bond is sigma or pi. A sigma bond lies on the internuclear axis. All

single bonds and one of the bonds in a multiple bond are sigma. A pi bond lies off the internuclear axis

and is generally made from overlapping p orbitals.

Between any two adjacent atoms we can form a pi bond in the plane of the paper and another pi bond

perpendicular to the plane of the paper. That's all! So two carbon atoms can be held together by, at

most, one sigma bond and two pi bonds. The bond order can be no higher than three. Note also that the

pi bond drawn above consists of two overlapping lobes above the plane of the molecule and two lobes

overlapping below the plane. This is one bond, not two.

It also can now be said that the reason a multiple bond counts only for one in VSEPR calculations is

that the pi bonds are equally distant from all of the sigma bonds. The pi bonds don't repel one electron

pair more than another. They don't "take up space."

10. Hybridization

How is it possible for electrons in p orbitals, which are at right angles to each other, to make structures

with as many different bond angles as we have seen? The answer is that bonding does not involve

simple s and p orbitals; it involves orbitals which are mixtures --or hybrids --of s and p orbitals.

Consider methane, CH4. Carbon, the central atom, has an electron configuration of ls

2 2s

2 2p

2.

__ __ __ __ __ 1s 2s 2p The four orbitals available for bonding are the 2s and the three 2p orbitals. But instead of using two

different types of orbitals, these are combined to give four identical orbitals, each of which is part s and

part p. Since the orbitals are one part "s" and three parts "p," they are called sp3 hybrids (a superscript

of "1" is assumed over the "s.") The superscripts in the hybrid orbital designations come from the

number of orbitals used to make them and the sum of the superscripts gives the number of hybrid

orbitals. For the molecules which we will study, the number of sigma orbitals (sigma bonds plus

unshared electron pairs) determines the hybridization, just as it determines the arrangement of electron

pairs. This is summarized in Table III.

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Orbital Arrangement of Orbitals Hybridization

2 Linear sp

3 Trigonal sp2

4 Tetrahedral sp3

5 Trigonal bipyramidal sp3d

6 octahedral sp3d

2

11. Molecules with More than an Octet

Structures in which the central atom has 5 or even 6 pairs of electrons have been frequently discussed.

Clearly they violate the octet rule. Why does this happen and, more importantly, how can we predict

when it will happen?

As can be seen in Table III, molecules which exceed the octet rule use d orbitals. Cases in which an

atom exceeds an octet involve a fairly heavy atom in which the energy of the valence electrons is at

least close to the energy of the d orbitals. Thus sulfur, even though it has no d electrons in its ground

state, has d orbitals (3d) which are close to the energy of its outer (3p)electrons. However the outer

electrons of oxygen (2s and 2p) are much lower in energy than the 3d orbitals. Oxygen, therefore, is

unable to exceed the octet rule.

In general you will find that atoms exceed the octet rule only when it is not possible to draw a structure

which doesn't. Thus you would not use double bonds to exceed the octet rule

12. Dipole Moments

Any bond in which the two ends are different (i.e. have different electronegativities) will have a

positive end and a negative end. We don't mark the positive and negative ends with a "+" and a "-"

because that implies that the charges are +l and -1. Instead we show partial charges by marking the

positive end with a δ+ and the negative end with a δ

-. This charge separation is called a "dipole" and a

bond or molecule which has a dipole is called "polar". The size of the dipole will depend upon the

difference in electronegativity and the distribution of formal charge in the atoms associated with the

bond. However we are interested primarily in determining the polarity of molecules, since this affects

their properties, and a polar bond does not necessarily lead to a polar molecule. Consider, for example,

carbon dioxide, CO2, and carbon tetrafluoride, CF4. Each has electronegative atoms which certainly

carry a negative charge, yet neither molecule is polar. These molecules are symmetrical and the dipoles

of the individual bonds end up canceling each other. On the other hand water, H2O, and

difluoromethane, CF2H2, are polar. Since water is bent, the bond dipoles do not end up canceling each

other. Difluoromethane is somewhat symmetrical but has the negative fluorine atoms on one side and

the (relatively) positive hydrogen atoms on the other.

C OO C

F

FF

FO

H H

C

H

H FF

Carbon dioxide Tetrafuoromethane water difluoromethane (non-polar) (non-polar) (polar) (polar)

δ+

δ+ δ+

δ+ δ+ δ-

δ-

δ-

δ-

δ-

δ- δ- δ-

δ-

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ORGANIC CHEMISTRY

1. Hydrocarbons: Alkanes

There are two types of hydrocarbons: saturated, which contain only C-C single bonds, where each

carbon is bound to 4 atoms, and unsaturated, which contain C-C multiple bonds.

Saturated hydrocarbons with the formula CnH2n+2 are called alkanes. They are named according to the

number of carbons:

formula name formula name

CH4 methane C6H14 hexane

C2H6 ethane C7H16 heptane

C3H8 propane C8H18 octane

C4H10 butane C9H20 nonane

C5H12 pentane C10H22 decane

I m , ’ -chain alkane, as in

CH3–CH2–CH2–CH3 or

H

H C

H

H

C

H

H

C

H

H

H

C

H

which is named n-butane (the n stands for normal). Butane has another structural isomer, a

compound with the same number and type of atoms but different bonds and properties:

H3C CH3CH

CH3

This branched isomer is called methylpropane (as you will see below). As more carbons are added,

more structural isomers can be drawn. For pentane, C5H12, there are three:

C C C C C H3C CH2 CH2 CH2 CH2or

C C CC

C

or H3C CH2 CH3CH

CH3

C C

C

C

C

or H3C

CH3

CH3

C CH3

10

Be aware that C C C C

C

m , ’ m as #1 above. To

y v ’ , y ’v m y m .

Draw all the structural isomers for hexane (there are 5). There is no mathematical formula to calculate

m m , y ’ j ry to start writing formulas and see how many you

. J y m , y m . I ’ y – there are only 75.

2. Steps for Nomenclature

Naming organic compounds must be done systematically. The name must be unambiguous, so

everyone has to follow the same rules.

1. Find and name the longest carbon chain in this isomer of heptane.

H3C CH2CH

CH3

CH2 CH2 CH3 longest chain: hexane

2. Name the groups attached to the main carbon chain, which are called substituents. If the

substituent is an alkane group, drop the –ane ending and add –yl. For example, a –CH3 substituent is

called methyl, a –CH2CH3 substituent is called ethyl, etc.

3. Number the carbon chain so that the substituents have the lowest possible numbers.

H3C CH2CH

CH3

CH2 CH2 CH3 2-methylhexane

1 2 3 4 5 6

Note above that the methyl group is at the 2 position, not the 5 position. A hyphen is placed between

the number and the substituent name. If more than one methyl substituent is present, then the prefix di-

is used:

H3C CH

CH3

CH2 CH2 CH3CH

CH3

2,3-dimethylhexane

4. The substituents are listed alphabetically (neglecting the prefixes di-, tri-, etc.) followed by the root

alkane name.

H3C CH

CH3

CH2 CH2 CH3CH

CH2CH3

3-ethyl–2–methylhexane

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Here are some common substituents:

name structure

methyl –CH3

ethyl –CH2–CH3

propyl –CH2–CH2–CH3

butyl –CH2–CH2–CH2–CH3

halo (fluoro, chloro, etc.) –X (X = halogen)

Alkenes and Alkynes

Hydrocarbons that contain carbon-carbon double bonds are called alkenes and have the formula CnH2n.

For example, removing two hydrogen atoms from ethane gives ethene, H2C=CH2, or

CC

H H

HH

Hydrocarbons with carbon-carbon triple bonds are called alkynes. The simplest alkyne has the

systematic name ethyne, but it is commonly called acetylene: H–C≡C–H.

Naming alkenes and alkynes is quite similar to naming alkanes. The –ane ending of alkanes is replaced

with –ene for alkenes or –yne for alkynes:

C3H6 H2C=CH–CH3 or

CC

H CH3

HH propene

C3H4 H–C≡C–CH3 propyne

If the alkene/alkyne contains more than three carbon atoms, there is more than one possible location for

the double/triple bond. The location of the bond is then indicated by the lowest numbered carbon atom

involved in the bond:

C4H8 H2C=CH–CH2CH3 1-butene (not 2-butene)

H3C–CH=CH–CH3 2-butene

H3C C C

CH3

CH CH3 4-methyl-2-pentyne

12

Alkenes can also have another type of isomer called a stereoisomer. Stereoisomers have the same

structural formula but different spatial arrangements of atoms. This is due to the fact that double bonds

do not rotate freely. Molecules with 2 identical groups on the same side of the double bond are

designated cis and those with the groups opposite are called trans. Take a look at each alkene with

formula C4H8:

C C

H

H3C

CH3

HC C

H

H3C

H

CH3

C C

H

CH3H

CH3

C C

H H

H CH2CH3

trans–2-butene cis–2-butene methylpropene 1-butene

Only the first two molecules exhibit stereoisomerism. The second two molecules need no cis- or trans–

designation.

When the molecule contains both a double bond and one or more substituents, the cis or trans comes

right before the number designating the location of the double bond:

C C

H

H3C H

CH3CH

CH3

C C

H

H3C

H

CH3CH

CH3

4-methyl-trans-2-pentene 4-methyl-cis-2-pentene

Hydrocarbon Derivatives

Hydrocarbons molecules often contain other atoms or groups of atoms. These are called functional

groups because they exhibit different reactivity. You should be able to recognize which functional

group a molecule contains. You need not be able to name individual molecules. See the chart of

. A R y m . A R’ j

hydrocarbon fragment, which may or may not be the same as R.

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Name Functional Group General

Formula

Example

halohydrocarbon –X (X = halogen) R–X CH3Cl

chloromethane or

methyl chloride

alcohol –OH R–OH CH3OH

methanol

ether –O– R–O–R’ CH3OCH3

dimethyl ether

aldehyde O

–C–H

O

R–C–H

HCHO

methanal (formaldehyde)

ketone O

–C–

O

R–C–R’

CH3COCH3

propanone (acetone)

carboxylic acid O

–C–OH

O

R–C–OH

CH3COOH

ethanoic acid (acetic acid)

ester O

–C–O–

O

R–C–O–R’

CH3COOCH2CH3

ethyl acetate

amine –NH2 R–NH2 CH3NH2

methylamine or

aminomethane