INTERMOLECULAR BONDING - HYDROGEN BONDS

The evidence for hydrogen bonding

Many elements form compounds with hydrogen - referred to as "hydrides". If you plot the boiling points of the hydrides of the Group 4 elements, 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.

If you repeat this exercise with the hydrides of elements in Groups 5, 6 and 7, 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 hydride of 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:

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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.

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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.

Consider two water molecules coming close together.

The + 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. On the same scale, van der Waals attractions represent mere passing acquaintances!

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 + 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. 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.

In hydrogen fluoride, the problem is a shortage of hydrogens. In water, there are exactly the right numbers of each. Water could be considered as the "perfect" hydrogen bonded system.

More complex examples of hydrogen bonding

The hydration of negative ions

When an ionic substance dissolves in water, water molecules cluster around the separated ions. This process is called hydration.

Water frequently attaches to positive ions by co-ordinate (dative covalent) bonds. It bonds to negative ions using hydrogen bonds.

The diagram shows the potential hydrogen bonds formed to a chloride ion, Cl^-. Although the lone pairs in the chloride ion are at the 3-level and wouldn't normally be active enough to form hydrogen bonds, in this case they are made more attractive by the full negative charge on the chlorine.

However complicated the negative ion, there will always be lone pairs that the hydrogen atoms from the water molecules can hydrogen bond to.

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.

They have the same number of electrons, and a similar length to the molecule. The 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 + charge.

In methoxymethane, the lone pairs on the oxygen are still there, but the hydrogens aren't sufficiently + 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 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.

THE STRUCTURE OF PROTEINS

Drawing the amino acids

In chemistry, if you were to draw the structure of a general 2-amino acid, you would probably draw it like this:

However, for drawing the structures of proteins, we usually twist it so that the "R" group sticks out at the side. It is much easier to see what is happening if you do that.

That means that the two simplest amino acids, glycine and alanine, would be shown as:

Peptides and polypeptides

Glycine and alanine can combine together with the elimination of a molecule of water to produce a dipeptide. It is possible for this to happen in one of two different ways - so you might get two different dipeptides.

Either:

Or:

In each case, the linkage shown in blue in the structure of the dipeptide is known as a peptide link. In chemistry, this would also be known as an amide link, but since we are now in the realms of biochemistry and biology, we'll use their terms.

If you joined three amino acids together, you would get a tripeptide. If you joined lots and lots together (as in a protein chain), you get a polypeptide.

A protein chain will have somewhere in the range of 50 to 2000 amino acid residues. You have to use this term because strictly speaking a peptide chain isn't made up of amino acids. When the amino acids combine together, a water molecule is lost. The peptide chain is made up from what is left after the water is lost - in other words, is made up of amino acid residues.

By convention, when you are drawing peptide chains, the -NH₂ group which hasn't been converted into a peptide link is written at the left-hand end. The unchanged -COOH group is written at the right-hand end.

The end of the peptide chain with the -NH₂ group is known as the N-terminal, and the end with the -COOH group is the C-terminal.

A protein chain (with the N-terminal on the left) will therefore look like this:

The "R" groups come from the 20 amino acids which occur in proteins. The peptide chain is known as the backbone, and the "R" groups are known as side chains.

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Note:  In the case where the "R" group comes from the amino acid proline, the pattern is broken. In this case, the hydrogen on the nitrogen nearest the "R" group is missing, and the "R" group loops around and is attached to that nitrogen as well as to the carbon atom in the chain.

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The primary structure of proteins

Now there's a problem! The term "primary structure" is used in two different ways.

At its simplest, the term is used to describe the order of the amino acids joined together to make the protein. In other words, if you replaced the "R" groups in the last diagram by real groups you would have the primary structure of a particular protein.

This primary structure is usually shown using abbreviations for the amino acid residues. These abbreviations commonly consist of three letters or one letter.

Using three letter abbreviations, a bit of a protein chain might be represented by, for example:

If you look carefully, you will spot the abbreviations for glycine (Gly) and alanine (Ala) amongst the others.

If you followed the protein chain all the way to its left-hand end, you would find an amino acid residue with an unattached -NH₂ group. The N-terminal is always written on the left of a diagram for a protein's primary structure - whether you draw it in full or use these abbreviations.

The wider definition of primary structure includes all the features of a protein which are a result of covalent bonds. Obviously, all the peptide links are made of covalent bonds, so that isn't a problem.

But there is an additional feature in proteins which is also covalently bound. It involves the amino acid cysteine.

If two cysteine side chains end up next to each other because of folding in the peptide chain, they can react to form a sulphur bridge. This is another covalent link and so some people count it as a part of the primary structure of the protein.

Because of the way sulphur bridges affect the way the protein folds, other people count this as a part of the tertiary structure (see below). This is obviously a potential source of confusion!

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The secondary structure of proteins

Within the long protein chains there are regions in which the chains are organised into regular structures known as alpha-helices (alpha-helixes) and beta-pleated sheets. These are the secondary structures in proteins.

These secondary structures are held together by hydrogen bonds. These form as shown in the diagram between one of the lone pairs on an oxygen atom and the hydrogen attached to a nitrogen atom:

The alpha-helix

In an alpha-helix, the protein chain is coiled like a loosely-coiled spring. The "alpha" means that if you look down the length of the spring, the coiling is happening in a clockwise direction as it goes away from you.

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Note:  If your visual imagination is as hopeless as mine, the only way to really understand this is to get a bit of wire and coil it into a spring shape. The lead on your computer mouse is fine for doing this!

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The next diagram shows how the alpha-helix is held together by hydrogen bonds. This is a very simplified diagram, missing out lots of atoms. We'll talk it through in some detail after you have had a look at it.

What's wrong with the diagram? Two things:

First of all, only the atoms on the parts of the coils facing you are shown. If you try to show all the atoms, the whole thing gets so complicated that it is virtually impossible to understand what is going on.

Secondly, I have made no attempt whatsoever to get the bond angles right. I have deliberately drawn all of the bonds in the backbone of the chain as if they lie along the spiral. In truth they stick out all over the place. Again, if you draw it properly it is virtually impossible to see the spiral.

So, what do you need to notice?

Notice that all the "R" groups are sticking out sideways from the main helix.

Notice the regular arrangement of the hydrogen bonds. All the N-H groups are pointing upwards, and all the C=O groups pointing downwards. Each of them is involved in a hydrogen bond.

And finally, although you can't see it from this incomplete diagram, each complete turn of the spiral has 3.6 (approximately) amino acid residues in it.

If you had a whole number of amino acid residues per turn, each group would have an identical group underneath it on the turn below. Hydrogen bonding can't happen under those circumstances.

Each turn has 3 complete amino acid residues and two atoms from the next one. That means that each turn is offset from the ones above and below, such that the N-H and C=O groups are brought into line with each other.

Beta-pleated sheets

In a beta-pleated sheet, the chains are folded so that they lie alongside each other. The next diagram shows what is known as an "anti-parallel" sheet. All that means is that next-door chains are heading in opposite directions. Given the way this particular folding happens, that would seem to be inevitable.

It isn't, in fact, inevitable! It is possible to have some much more complicated folding so that next-door chains are actually heading in the same direction. We are getting well beyond the demands of UK A level chemistry (and its equivalents) now.

The folded chains are again held together by hydrogen bonds involving exactly the same groups as in the alpha-helix.

The tertiary structure of proteins

What is tertiary structure?

The tertiary structure of a protein is a description of the way the whole chain (including the secondary structures) folds itself into its final 3-dimensional shape. This is often simplified into models like the following one for the enzyme dihydrofolate reductase (DHFR). Enzymes are, of course, based on proteins.

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Note:  This diagram was obtained from the RCSB Protein Data Bank. If you want to find more information about dihydrofolate reductase, their reference number for it is 7DFR.

There is nothing particularly special about this enzyme in terms of structure. I chose it because it contained only a single protein chain and had examples of both types of secondary structure in it.

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The model shows the alpha-helices in the secondary structure as coils of "ribbon". The beta-pleated sheets are shown as flat bits of ribbon ending in an arrow head. The bits of the protein chain which are just random coils and loops are shown as bits of "string".

The colour coding in the model helps you to track your way around the structure - going through the spectrum from dark blue to end up at red.

You will also notice that this particular model has two other molecules locked into it (shown as ordinary molecular models). These are the two molecules whose reaction this enzyme catalyses.

What holds a protein into its tertiary structure?

The tertiary structure of a protein is held together by interactions between the the side chains - the "R" groups. There are several ways this can happen.

Ionic interactions

Some amino acids (such as aspartic acid and glutamic acid) contain an extra -COOH group. Some amino acids (such as lysine) contain an extra -NH₂ group.

You can get a transfer of a hydrogen ion from the -COOH to the -NH₂ group to form zwitterions just as in simple amino acids.

You could obviously get an ionic bond between the negative and the positive group if the chains folded in such a way that they were close to each other.

Hydrogen bonds

Notice that we are now talking about hydrogen bonds between side groups - not between groups actually in the backbone of the chain.

Lots of amino acids contain groups in the side chains which have a hydrogen atom attached to either an oxygen or a nitrogen atom. This is a classic situation where hydrogen bonding can occur.

For example, the amino acid serine contains an -OH group in the side chain. You could have a hydrogen bond set up between two serine residues in different parts of a folded chain.

You could easily imagine similar hydrogen bonding involving -OH groups, or -COOH groups, or -CONH₂ groups, or -NH₂ groups in various combinations - although you would have to be careful to remember that a -COOH group and an -NH₂ group would form a zwitterion and produce stronger ionic bonding instead of hydrogen bonds.

van der Waals dispersion forces

Several amino acids have quite large hydrocarbon groups in their side chains. A few examples are shown below. Temporary fluctuating dipoles in one of these groups could induce opposite dipoles in another group on a nearby folded chain.

The dispersion forces set up would be enough to hold the folded structure together.

Sulphur bridges

Sulphur bridges which form between two cysteine residues have already been discussed under primary structures.