in proteins, as well as this ionic interaction, hydrogen bonding (which is discussed later in this chapter) also plays a role in linking the two amino acids together. It is this combination of ionic bonds and hydrogen bonds that forms the complete link. This combination of ionic and hydrogen bond interaction is referred to by biochemists as a “salt bridge.”
3.7 Covalent Bonding
In ionic bonding, the central feature of the bond is the transfer of an electron from one atom to another to generate two ions. In covalent bonds, an electron is shared. Covalent bonds take place between atoms that are generally close to each other in the Periodic Table and have a small difference in electronegativity (the tendency of an atom to attract electrons). Similarly to ionic bonds, the sharing of an electron allows the noble gas electron configuration to be attained by both atoms.
The simplest example of a covalent bond is found in the hydrogen molecule, shown in Figure 3.7. To return to the electron configurations that we discussed earlier, a hydrogen atom has an electron configuration of 1s1: one electron in the first shell, the s subshell. It would like to have two in this subshell so that the subshell is full. In the hydrogen molecule, the sharing of each electron between two hydrogen atoms allows for two electrons in this subshell.
Figure 3.7 The covalent bond in the hydrogen molecule. The two electrons are shared between both hydrogen atoms.
Covalent bonds are very strong. An example is diamond, a covalent network of carbon. Taking the same logic that we used for ionic bonds, the energy per bond is ∼6 × 10−19 J. That's equivalent to 150 times the thermal energy at room temperature. In other words, the thermal energy in the bonds of diamond is much lower than the energy needed to break carbon–carbon bonds. Diamond is very stable at room temperature.
3.7.1 Covalent Bonds and Life
Covalent bonds play a central role in biology because they are the bonds that hold atoms together in the vast array of molecules from which life is constructed.
Covalent bonds are also found in specific situations where biologically important molecular structure is required. In analogy to the role of ionic bonds in holding charged amino acids together in protein chains, covalent bonds can form between sulfur-containing amino acids (cysteine and methionine). The sulfur atoms within the amino acids join together and form a disulfide bridge. These bonds anchor the structure of the protein, making sure that its three-dimensional shape is maintained. Figure 3.8 shows a schematic example of disulfide bridges made of covalent bonds holding a protein chain together in a loop.
Figure 3.8 As well as holding atoms together in molecules, covalent bonds link within molecules to provide structure. Here covalent bonds within two disulfide bridges hold a protein chain made of amino acids into a loop. The covalent bonds in this structure, as in many similar diagrams, are shown as solid black lines.
As disulfide bridges can add rigidity to a protein, life can use them to enhance the stability of proteins. For example, some microbes, such as those that live in hot springs in Yellowstone National Park (Figure 3.9), must cope with high temperatures, often well over 60 °C. Some proteins have been found to contain extra covalent bonds to enhance their thermostability. We return to examine adaptations to extreme environments later in the book. For now, the point to learn is that changes in bonding in molecules can be used by life as one of a repertoire of adaptations to extreme physical conditions.
Figure 3.9 An enhanced number of covalent bonds in proteins is used in some microbes to stabilize proteins against high temperatures. Enhanced stability of biological molecules is needed in such environments, for example in these volcanic pools in Yellowstone National Park, USA. In this image, the microbes live within the browns and yellows in the spring, colors caused by microbial pigments and minerals such as iron.
Source: Reproduced with permission of Public-domain-photos.com; Jon Sullivan.
3.8 Metallic Bonding
Metallic bonding is the type of bonding found in metals such as sodium, copper, magnesium, and iron. Consider sodium. We have already seen how it can form ionic bonds, such as in NaCl. We saw that it has one spare electron which it can transfer to an atom such as chlorine to become a sodium ion and achieve a stable noble gas configuration. Another way it could lose this electron is for the electron to dissociate from the atom and form a “sea” of electrons around other sodium ions that are behaving in an identical way. The positively charged sodium ions do not fly apart from electrostatic repulsion because of this intervening sea of negatively charged electrons. This is depicted in Figure 3.10. In this way, the atoms, by delocalizing the electrons (which is essentially like losing them), achieve a stable noble gas configuration.
Figure 3.10 Metallic bonding showing a “sea” of delocalized electrons around positively charged metal ions.
Metals form strong bonds. Again, using the reasoning discussed earlier, consider potassium. Its energy per bond is ∼5 × 10−19 J, which is equivalent to approximately 125 times the thermal energy in the bond at room temperature. Many metals are therefore stable at room temperature (e.g. iron and silver).
Interestingly, metallic bonding is not relevant for life. Life certainly does make use of metal ions in many diverse ways that will crop up throughout this book. Metal ions are particularly prominent in enzymes, in which they play a role in mediating the catalysis of chemical reactions as cofactors. These are found in many situations in cells where electron transfer is needed, such as iron–sulfur clusters used in energy-gathering processes. Examples of these ions include copper, iron, and magnesium. However, in all these examples, the metal ions are generally present as individual ions, not as bulk metallic deposits. So where are all the metal structures in life? You might like to read the Discussion Point.
Discussion Point: Why Doesn't Terrestrial Biochemistry Use Metal Structures?
Although terrestrial biochemistry makes extensive use of metal ions, such as iron and copper as cofactors in enzymes, life does not fabricate metal structures. This is surprising given the vast use of metals and metallic structures in human engineering. Why aren't bones made from steel for example, or horns in animals made from titanium? There are a number of possible reasons that may not be mutually exclusive: (i) It may be energetically too expensive. Most metals are not pure in the environment and considerable energy is used in human manufacturing processes to purify them. Locked up in sulfides, carbonates, and other compounds, they are difficult to purify biologically. (ii) Pure metals have large densities. For example, compare the density of steels (typically ∼8000 kg m−3) to calcium phosphate bone (typically ∼3800 kg m−3) and one can see that metal structures would come with a considerable mass burden. (iii) Metals are generally at low abundance in the environment. It should be pointed out that this argument holds for some metals such as vanadium, but it clearly does not for metals such as iron and aluminum, which constitute 8.1% and 5.0% of Earth's crust, respectively. The preceding three arguments assume that nature could have attempted to use metals but for various reasons rejected them. However, we have no evidence of extensive experiments in metal construction in biology. It could be that pathways to other high-strength compounds are