Bonds found in biological chemistry

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Chapter 3 Bonds found in biological chemistry

Chapter contents

The significance of different types of bond 13

Covalent bonds 13

‘Oil and water don’t mix’ 16

Ionic bonds 17

Hydrogen bonds 18

Van der Waals attraction 19

Each electron orbital has a specific allocation of spaces available for electrons to fill (see Figure 2.3, p. 9). In the case of the outer orbitals, it is unusual to find all the spaces empty or full and one atom will seek out another compatible atom so that they can share their respective electrons and fill their outer orbitals. This arrangement creates great stability for the associated atoms, which is why, whenever possible, bonds are formed, between individual atoms.

Taking the examples of hydrogen and carbon used in Chapter 2 (Figure 2.3), it is possible to see that hydrogen:

• has one electron in the outer orbital

• needs to combine with elements that will provide one electron to fill its outer orbital.

However, carbon:

• has two electrons in the first orbital, which is therefore full

• has four electrons in the second orbital

• needs four electrons to complete its outer orbital, making a total of eight electrons for a full orbital.

The final result of this union is that four hydrogen atoms attach themselves to one carbon atom, thus filling the outer orbitals of both the carbon and the hydrogens satisfying the needs of both elements (Figure 3.1).

Figure 3.1 Formation of methane, with additional diagrams showing some of the more common chemical representations of this molecule.

It is a matter of simple arithmetic to work out where this leads: a more complex molecule – ethane – can be formed from two carbon atoms and six hydrogen atoms Figure 3.2. The process can be continued to make an even longer compound called propane (Figure 3.3).

Figure 3.2 Formation of ethane and common chemical representations.

Figure 3.3 Formation of propane.

Note: as drawing chemicals out in this way can become complicated, there is a need for a method of abbreviation. Each figure, therefore, contains examples of various chemical shorthands used to represent these structures. This will enable you to become used to seeing them. Naming and representation will be covered in more detail in Chapter 5 ‘Nomenclature: representation of chemical structures and basic terminology’.

The Significance of Different Types of Bond

The way elements bond dictates the function of the particular chemical compound they create. A basic understanding of how atoms interact will help you understand the pharmacology you will encounter in your work.

Covalent Bonds

Bonds, formed by completely equal sharing of electrons, such as the previous examples, are called covalent bonds. This principle of bond formation can be applied to other elements, such as oxygen, nitrogen and sulphur, to create various organic compounds that can be structural or part of the metabolic processes of a plant or animal.

Covalent bonds have set bond lengths and angles, dictated by the electron orbitals of the atoms that are bonding together. Each atom has a distinctive arrangement of electrons due to the:

• electrical field produced by the orbiting electrons

• distance the outer orbital is from the nucleus

• shape of the orbital (Figure 3.4), as mostly electrons do not move in a perfectly circular orbit around the nucleus.

Figure 3.4 Different shapes of electron clouds.

For a greater appreciation, blow up balloons of different shapes and arrange them as you see them in Figure 3.4. The different balloons are the accepted shapes of electron orbitals and the shape of the electrostatic force created by this orbital. Notice how it is possible to bend the orbital shown in B around the axis, but that in C the movement around the axis is restricted because the sides of the balloons push against one another limiting movement.

Electron clouds are not solid but are negatively charged (like similar poles on magnets), which means that they repel one another to a particular distance and therefore into a particular shape, as in the balloon experiment. Therefore, when several of these clouds are in close proximity, which is what occurs when several atoms bond to form a molecule, the molecule acquires a specific shape.

Naturally, the above is a very simplistic explanation of what really happens, but it is important at this stage to appreciate that chemistry, and therefore pharmacology, is three dimensional, and it is very necessary to starting thinking in three dimensions to gain a proper appreciation of pharmacology, as the three-dimensional concept is very important when considering how chemicals (substrates) fit into sites that activate physiological processes.

Each substrate and site has a specific shape and in many cases only certain molecules will be allowed to fit into the site, which may be on an enzyme, the surface of a cell or at a nerve ending (see Chapter 11 ‘Amino acids and proteins’, p. 85 and Chapter 19 ‘Pharmacodynamics: how drugs elicit a physiological effect’, p. 137).

Importance of Covalent Bonds

Covalent bonds are the strongest bonds in nature and under normal biological conditions have to be broken with the help of enzymes. This is due to the even sharing of electrons between the bonded atoms and as with anything equally shared there is no conflict to weaken the arrangement. There are two types of covalent bond:

• non-polar: possess no charge

• polar: possess a charge.

Non-polar covalent bonds (Figure 3.5(i), Figure 3.6)

• These are very strong bonds.

• The electrons in the adjacent atoms are shared equally, which means that there is no push/pull effect (dipole moment; see below).

Figure 3.5 The relative positions of the bonding pair of electrons in (i) a non-polar covalent bond, (ii) a polar covalent bond and (iii) an ionic bond.

Figure 3.6 Non-polar methane. Electrons are shared evenly between the carbon and hydrogen atoms. Note how each hydrogen is not only at the same distance from the carbon but also arranged at the same angle.

Example of relevance to pharmacology: Covalent bonds are found in protein chains where two sulphurs are joined together (disulphide bonds) (Figure 3.7). In antibody molecules, the disulphide bond firmly holds the two parts of the antibody together, allowing the rest of the structure to hinge around it. This flexibility is an important part of the function of an antibody as it allows the antibody binding sites to adapt in binding to antigens and reach more antigen sites.

Figure 3.7 The disulphide bond in an antibody maintains rigidity while also providing flexibility.

Polar covalent bonds (Figure 3.5(ii))

• A polar covalent bond is still very strong, but the electron sharing between the atoms is unequal.

• One atom is able to pull more strongly on the electrons than the other, so that the part of the molecule which pulls the electrons towards it is relatively negatively charged while the other part of the molecule which has had the electrons pulled away from it is relatively positively charged.

The following is an example of the pharmacological relevance of a polar covalent bond:

The atomic composition of oxygen compared to that of hydrogen is far more attractive to electrons than both of the hydrogen atoms and it is able to exert a much more powerful pull on the shared electrons than the hydrogen atoms. The oxygen therefore becomes more negatively charged as it pulls the bonding pair of electrons closer towards it.

The average electron density around the oxygen atom is 10×times the amount around the hydrogen atom. This charge separation creates something called a dipole moment, as the oxygen has a partial negative charge and the hydrogen a partial positive charge (Figure 3.8(i)). Water is actually a fascinating molecule, the pharmacological importance of which will become evident as the book progresses.

Figure 3.8 Dipole moments of a water molecule and of glycine.

The dipole moment described above also occurs in amino acids (Figure 3.8(ii); see also Chapter 11 ‘Amino acids and proteins’, p. 86) and becomes important in the temporary binding of chemicals to proteins, which are made up of amino acids, for example at enzyme binding sites or binding of chemicals to plasma proteins.

‘Oil and Water Don’t Mix’

Whether a molecule is charged (polar) or not charged (non-polar) has great relevance when it comes to pharmacology. As explained above, the electron bonding between two elements such as carbon and hydrogen is equally shared so the union is well balanced. Because of this equilibrium, methane (see Figure 3.1 and Figure 3.9) has no charge (i.e. it is non-polar).

Figure 3.9 The non-polar covalent bond of methane (CH₄) and the polar covalent bond of water (H₂O). (Look also at the shape of the molecule created by the bonds.)

However, as you have seen, oxygen exerts a much greater electrostatic charge than hydrogen, and so draws electrons towards itself, creating a dipole moment and therefore a charged molecule (see Figure 3.5(ii) and Figure 3.9). You now know that where there is a dipole moment the molecule becomes charged.

You might have heard the phrase ‘oil and water do not mix’; more than likely this is something you will have experienced when making a vinaigrette dressing. Shaking the components of the dressing together seems to mix them, but the effect is only temporary and the mixture soon separates out into two layers, with the oily part sitting on top. Mixing two different types of oil, however, is easy and there is no separation.

The separation occurs because water and oil have different characteristics, water being polar (charged) and oil (the backbone of which is made up of carbons) being non–polar (uncharged).

Polar compounds tend to be water loving (hydrophilic) whereas non-polar compounds tend to be water hating (hydrophobic). Some natural compounds are large enough to have a bit of both qualities in their structure, which creates some very interesting pharmacological situations.

The importance of polar and non-polar compounds will soon become clear, and a basic understanding of this type of bond, and of hydrogen bonds, will make basic pharmacology much easier to understand.

Ionic Bonds (Figure 3.5(iii))

As previously discussed, all atoms like to have full outer orbitals (see Chapter 2 ‘The atom: the smallest unit of pharmacology’, p. 7). Elements such as chlorine and sodium are very unstable on their own because:

• chlorine is short of one electron in its outer shell

• sodium has only one electron in its outer shell.

Both elements are desperately seeking a way of creating stability, in this case by chlorine acquiring another electron in its outer shell and by sodium losing its electron:

• Chlorine + 1 electron = stability

• Sodium − 1 electron = stability

When the nucleus is relatively large and only one electron is involved, the reactivity of the element is very vigorous as it achieves stability. This is why, when sodium and chlorine atoms come into contact, a great deal of energy is released (this can be seen in the safe environment of a laboratory).

The chlorine snatches the electron from the sodium atom, making its outer orbital stable, but leaving sodium with a positive charge (Figure 3.10). The sodium atom has thus become a positively charged sodium ion (Na⁺). This is because the protons in the nucleus normally exactly equal the negative electrons in the outer shell (see Figure 2.1, p. 8). Sodium is now short of an electron and so, overall, gains a positive charge:

• Before ionic bonding = 11 protons + 11 electrons

• After ionic bonding = 11 protons + 10 electrons

• Therefore overall positive charge is +1

Figure 3.10 The formation of ions.

Using the same principle, the chlorine has an extra electron, which gives it an overall negative charge, so it is now a negatively charged chloride ion:

• Before ionic bonding = 11 protons + 11 electrons

• After ionic bonding = 11 protons + 10 electrons

• Therefore overall negative charge −1

As always, the reality is not as simple as this, but for the purposes of the book it is fine to make this assumption.

As ions of opposite charge are attracted to one another, like opposite poles of a magnet, they form a type of bond called an ionic bond. Thus a salt crystal (of the variety you use on your food) is a lattice of ionic bonds (Figure 3.11).

Figure 3.11 The basic lattice of a sodium chloride salt crystal and its three-dimensional model. These units link together to form visible crystals.

Hydrogen Bonds

Hydrogen bonds are very special as they are, strictly speaking, not true bonds but forces of attraction between the hydrogen atoms in one molecule and atoms of high electronegativity in other molecules. This situation is due to the uneven sharing of the bonding electron pair in a polar covalent bond (see Figure 3.5(ii)).

The strength of the hydrogen bond is only one-tenth that of a covalent bond, but when a large number of hydrogen bonds are formed between two molecules or areas of the same molecule the overall result is very strong.

The section on polar bonds above shows that the hydrogen attached to nitrogen or oxygen becomes relatively positively charged, as the nitrogen and oxygen atoms, pulling the electrons towards them, become relatively negatively charged (see Figure 3.8).

Hydrogen Bonding in Water

The relatively negatively charged oxygen atom in a water molecule is able to attract positively charged hydrogen atoms attached to other water molecules. Simple mathematics then dictates that each water molecule has the potential to form four hydrogen bonds with surrounding water molecules (Figure 3.12). This enables water to form a type of lattice even as a liquid, a phenomenon that accounts for its relatively high boiling point compared with molecules of a similar size, which are gaseous at the same temperature. It also explains why ice floats and water freezes from the top downwards.

Figure 3.12 One water molecule can form four hydrogen bonds.

There is a constant exchange of hydrogen between water molecules but the overall integrity stays the same providing the temperature does not get too high and agitates the molecules beyond a level where they can stick together.

A molecule that has a tightly and securely bound non-polar covalent bond cannot form a hydrogen bond. This form of hydrogen bonding is important in biological organisms in the following ways:

• It holds separate strands of amino acids (polypeptides) together to create three-dimensional structures (Figure 3.13). This enables protein to have a structural function (e.g. collagen) but also enables the protein comprising the structure of enzymes to have a specific shape, creating distinctive sites for substrates to fit into.

• It ensures the two strands of the DNA double helix stay together, but at the same time is weak enough to allow the two strands to separate when duplication of the DNA occurs (see Chapter 12 ‘Purines and pyrimidines’, p. 91).

• It helps antibodies to bind to their antigen.