Intermolecular Bonding - Hydrogen Bonds

INTERMOLECULAR BONDING -
HYDROGEN BONDS
This page explains the origin of hydrogen bonding - a

relatively strong form of intermolecular attraction. If you
are also interested in the other intermolecular forces (van
der Waals dispersion forces and dipole-dipole
interactions), there is a link at the bottom of the page.
The evidence for hydrogen bonding
Many elements form compounds with hydrogen. If you

plot the boiling points of the compounds of the Group 4
elements with hydrogen, you find that the boiling points
increase as you go down the group.
The increase in boiling point happens because the

molecules are getting larger with more electrons, and so
van der Waals dispersion forces become greater.
Note: If you aren't sure about van der Waals dispersion

forces, it would pay you to follow this link before you go on.
If you repeat this exercise with the compounds of the

elements in Groups 5, 6 and 7 with hydrogen, something
odd happens.
Although for the most part the trend is exactly the same
as in group 4 (for exactly the same reasons), the boiling
point of the compound of hydrogen with the first element
in each group is abnormally high.
In the cases of NH3, H2O and HF there must be some

additional intermolecular forces of attraction, requiring
significantly more heat energy to break. These relatively
powerful intermolecular forces are described as
hydrogen bonds.
The origin of hydrogen bonding
The molecules which have this extra bonding are:
Note: The solid line represents a bond in the plane of the

screen or paper. Dotted bonds are going back into the screen
or paper away from you, and wedge-shaped ones are coming
out towards you.
Notice that in each of these molecules:
The hydrogen is attached directly to one of the most

electronegative elements, causing the hydrogen to
acquire a significant amount of positive charge.
Each of the elements to which the hydrogen is

attached is not only significantly negative, but also
has at least one "active" lone pair.
Lone pairs at the 2-level have the electrons

contained in a relatively small volume of space
which therefore has a high density of negative
charge. Lone pairs at higher levels are more diffuse
and not so attractive to positive things.
Note: If you aren't happy about electronegativity, you should

follow this link before you go on.
Consider two water molecules coming close together.
The δ+ hydrogen is so strongly attracted to the lone pair

that it is almost as if you were beginning to form a co-
ordinate (dative covalent) bond. It doesn't go that far, but
the attraction is significantly stronger than an ordinary
dipole-dipole interaction.
Hydrogen bonds have about a tenth of the strength of an

average covalent bond, and are being constantly broken
and reformed in liquid water. If you liken the covalent
bond between the oxygen and hydrogen to a stable
marriage, the hydrogen bond has "just good friends"
status.
Water as a "perfect" example of hydrogen bonding
Notice that each water molecule can potentially form four

hydrogen bonds with surrounding water molecules.
There are exactly the right numbers of δ+ hydrogens and
lone pairs so that every one of them can be involved in
hydrogen bonding.
This is why the boiling point of water is higher than that of

ammonia or hydrogen fluoride.
Note: You will find more discussion on the effect of hydrogen

bonding on the properties of water in the page on molecular
structures.
In the case of ammonia, the amount of hydrogen bonding

is limited by the fact that each nitrogen only has one lone
pair. In a group of ammonia molecules, there aren't
enough lone pairs to go around to satisfy all the
hydrogens.
That means that on average each ammonia molecule

can form one hydrogen bond using its lone pair and one
involving one of its δ+ hydrogens. The other hydrogens
are wasted.
In hydrogen fluoride, the problem is a shortage of

hydrogens. On average, then, each molecule can only
form one hydrogen bond using its δ+ hydrogen and one
involving one of its lone pairs. The other lone pairs are
essentially wasted.
In water, there are exactly the right number of each.

Water could be considered as the "perfect" hydrogen
bonded system.
Warning: It has been pointed out to me that some sources

(including one of the UK A level Exam Boards) count the
number of hydrogen bonds formed by water, say, differently.
They say that water forms 2 hydrogen bonds, not 4. That is
often accompanied by a diagram of ice next to this statement
clearly showing 4 hydrogen bonds!
Reading what they say, it appears that they only count a

hydrogen bond as belonging to a particular molecule if it
comes from a hydrogen atom on that molecule. That seems
to me to be illogical. A hydrogen bond is made from two parts
- a δ+ hydrogen attached to a sufficiently electronegative
element, and an active lone pair. These interact to make a
hydrogen bond, and it is still a hydrogen bond irrespective of
which end you look at it from.
The IUPAC definitions of a hydrogen bond make no reference

at all to any of this, so there doesn't seem to be any "official"
backing for this one way or the other.
However, it is essential that you find out what your examiners

are expecting. They make the rules for the exam you will be
sitting, and you have no choice other than to play by those
rules.
More complex examples of hydrogen bonding
Hydrogen bonding in alcohols
An alcohol is an organic molecule containing an -O-H

group.
Any molecule which has a hydrogen atom attached

directly to an oxygen or a nitrogen is capable of hydrogen
bonding. Such molecules will always have higher boiling
points than similarly sized molecules which don't have an
-O-H or an -N-H group. The hydrogen bonding makes the
molecules "stickier", and more heat is necessary to
separate them.
Ethanol, CH3CH2-O-H, and methoxymethane, CH3-O-

CH3, both have the same molecular formula, C2H6O.
Note: If you haven't done any organic chemistry yet, don't

worry about the names.
They have the same number of electrons, and a similar

length to the molecule. The other van der Waals
attractions (both dispersion forces and dipole-dipole
attractions) in each will be much the same.
However, ethanol has a hydrogen atom attached directly

to an oxygen - and that oxygen still has exactly the same
two lone pairs as in a water molecule. Hydrogen bonding
can occur between ethanol molecules, although not as
effectively as in water. The hydrogen bonding is limited
by the fact that there is only one hydrogen in each
ethanol molecule with sufficient δ+ charge.
In methoxymethane, the lone pairs on the oxygen are still

there, but the hydrogens aren't sufficiently δ+ for
hydrogen bonds to form. Except in some rather unusual
cases, the hydrogen atom has to be attached directly to
the very electronegative element for hydrogen bonding to
occur.
The boiling points of ethanol and methoxymethane show

the dramatic effect that the hydrogen bonding has on the
stickiness of the ethanol molecules:
ethanol (with hydrogen bonding) 78.5°C

methoxymethane (without hydrogen bonding) -24.8°C
The hydrogen bonding in the ethanol has lifted its boiling

point about 100°C.
It is important to realise that hydrogen bonding exists in

addition to other van der Waals attractions. For example,
all the following molecules contain the same number of
electrons, and the first two are much the same length.
The higher boiling point of the butan-1-ol is due to the
additional hydrogen bonding.
Comparing the two alcohols (containing -OH groups),

both boiling points are high because of the additional
hydrogen bonding due to the hydrogen attached directly
to the oxygen - but they aren't the same.
The boiling point of the 2-methylpropan-1-ol isn't as high

as the butan-1-ol because the branching in the molecule
makes the van der Waals attractions less effective than
in the longer butan-1-ol.
Hydrogen bonding in organic molecules containing

nitrogen
Hydrogen bonding also occurs in organic molecules

containing N-H groups - in the same sort of way that it
occurs in ammonia. Examples range from simple
molecules like CH3NH2 (methylamine) to large
molecules like proteins and DNA.
The two strands of the famous double helix in DNA are

held together by hydrogen bonds between hydrogen
atoms attached to nitrogen on one strand, and lone pairs
on another nitrogen or an oxygen on the other one.
Questions to test your understanding

If this is the first set of questions you have done,
please read the introductory page before you start. You
will need to use the BACK BUTTON on your browser
to come back here afterwards.
questions on hydrogen bonding
answers
Where would you like to go now?
To look at van der Waals forces . . .
To the bonding menu . . .
To the atomic structure and bonding menu . . .
To Main Menu . . .
© Jim Clark 2000 (last modified January 2019)

Intermolecular Bonding - Hydrogen Bonds

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INTERMOLECULAR BONDING -

This page explains the origin of hydrogen bonding - a

The evidence for hydrogen bonding

Many elements form compounds with hydrogen. If you

The increase in boiling point happens because the

Note: If you aren't sure about van der Waals dispersion

If you repeat this exercise with the compounds of the

In the cases of NH3, H2O and HF there must be some

The origin of hydrogen bonding

The molecules which have this extra bonding are:

Note: The solid line represents a bond in the plane of the

Notice that in each of these molecules:

The hydrogen is attached directly to one of the most

Each of the elements to which the hydrogen is

Lone pairs at the 2-level have the electrons

Note: If you aren't happy about electronegativity, you should

Consider two water molecules coming close together.

The δ+ hydrogen is so strongly attracted to the lone pair

Hydrogen bonds have about a tenth of the strength of an

Water as a "perfect" example of hydrogen bonding

Notice that each water molecule can potentially form four

This is why the boiling point of water is higher than that of

Note: You will find more discussion on the effect of hydrogen

In the case of ammonia, the amount of hydrogen bonding

That means that on average each ammonia molecule

In hydrogen fluoride, the problem is a shortage of

In water, there are exactly the right number of each.

Warning: It has been pointed out to me that some sources

Reading what they say, it appears that they only count a

The IUPAC definitions of a hydrogen bond make no reference

However, it is essential that you find out what your examiners

More complex examples of hydrogen bonding

Hydrogen bonding in alcohols

An alcohol is an organic molecule containing an -O-H

Any molecule which has a hydrogen atom attached

Ethanol, CH3CH2-O-H, and methoxymethane, CH3-O-

Note: If you haven't done any organic chemistry yet, don't

They have the same number of electrons, and a similar

However, ethanol has a hydrogen atom attached directly

In methoxymethane, the lone pairs on the oxygen are still

The boiling points of ethanol and methoxymethane show

ethanol (with hydrogen bonding) 78.5°C

The hydrogen bonding in the ethanol has lifted its boiling

It is important to realise that hydrogen bonding exists in

Comparing the two alcohols (containing -OH groups),

The boiling point of the 2-methylpropan-1-ol isn't as high

Hydrogen bonding in organic molecules containing

Hydrogen bonding also occurs in organic molecules

The two strands of the famous double helix in DNA are

Questions to test your understanding

Where would you like to go now?

To look at van der Waals forces . . .

To the bonding menu . . .

To the atomic structure and bonding menu . . .

© Jim Clark 2000 (last modified January 2019)

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