Throughout biochemistry there are many bonds without which life as it is on earth today would not be possible. One of the most important bonds of these is the hydrogen bond, a weak chemical bond that is present in essential biological molecules such as water and polypeptides. A hydrogen bond is defined by Campbell and Reece as occurring when a hydrogen atom is covalently bonded to an electronegative atom but attracted to another electronegative atom. In water molecules, there are several key reasons why hydrogen bonds can be formed and explaining them in water a good way to show the chemistry.
Firstly, the presence of covalent bond between the hydrogen and the oxygen means that the electrons in the outer shells of both atoms are shared- 1 electron from hydrogen and 1 electron from oxygen.
Since the 2 electrons are shared, they are free to move within the covalent bond to the atom that is the most electronegative. In the case of water, this is oxygen.
As a result of the electrons moving to the oxygen side of the bond, the hydrogen becomes less electron-dense and becomes a slight positive charge known as a delta-positive charge. It is this positive charge that has the ability to attract other negatively charged objects, since opposite electrostatic charged atoms attract each other. On the oxygen atom of each water molecule there is a lone pair of electrons that are negatively charged, which makes oxygen delta-negative. This means that between water molecules, the delta-positive hydrogen of one molecule is able to attract a lone pair of electrons from the delta-negative oxygen atom of another water molecule (Fig.
Fig. 1 Hydrogen bonding in water
A hydrogen bond, however, is comparatively weak to covalent or ionic bond, as much as 22 times time weaker [Libes 2009], so in order to explain why hydrogen bonds are so necessary in life it is perhaps not significant that hydrogen bonds are weak on their own, since the majority of their use within strong structures is facilitated by their strength as a large number of hydrogen bonds. For example, the fundamental strength of tendons and skin lies within the many hydrogen bonds in the collagen protein.
For formation of collagen, the strength of hydrogen bonds is required to firstly join two amino acid chains (polypeptides) together into a helix. Three helices are then bound into a triple helix by yet more hydrogen bonds. The result is a fibrous quaternary protein structure with a high tensile strength that the mammalian skeletal muscles could not function without. Tendons attach skeletal muscles to their respective bones and we would simply not be able to move without them. Other uses of hydrogen bonds in proteins include contributing to the specific conformational shape of globular proteins, called protein folding.
A precise 3D shape is required in most enzymes so that the shape of binding site (active site) is complementary to the chemical reacting with the enzyme (substrate). Hydrogen bonds are essential, along with ionic bonds, covalent bonds, disulphide bonds and hydrophobic interactions, for making secondary structures (i.e. alpha-helices and beta-pleated sheets) coil into a tertiary structure. A tertiary structure, or a quaternary structure after further protein folding, can then be utilized as a specific enzyme within organisms to carry out specific metabolic reactions.
It is the hydrogen bonding found in water, in fact, that makes the metabolic reactions in the human body so efficient. The slight increase of strength between water molecules caused by hydrogen bonds means that in comparison to other fluids without hydrogen bonds, water requires a lot of energy to raise the temperature of it. This is called high heat specific capacity and may be defined as the amount of energy required to change the temperature of 1g of a substance by 1C, an attribute that is especially useful when the body is actively maintaining the body temperature at 37C. Since it takes so much energy to change the temperature of water, the molecule estimated to take up 70% of an adult’s body, the core temperature of the body is resistant to fluctuations.
For metabolic reactions this is very useful because it means that the enzymes can work at their optimum temperature, often the same as 37C core temperature, and thus the metabolic reactions within the body are very efficient. For every 10C below optimum temperature, the rate of successive substrate-enzyme collisions decreases by 2 to 3 times [Campbell and Reece, p862]. High specific heat capacity also benefits marine environments by resisting temperature fluctuations, which is perhaps why marine food chains are often many times longer than those of terrestrial organisms.
The high heat capacity of water is one of several hydrogen-bonding attributes that benefit the marine environments, unsurprisingly, with the high surface tension and the small relative density of ice also playing a large role in how aquatic organisms survive. The high surface tension of water is perhaps best explained by relating it to close proximity of adjacent water molecules in aqueous solution. This closeness is of course caused by the numerous hydrogen bonds (Fig. 2) existing between the water molecules and is named cohesion, a word that can be defined as being united as a whole.
Fig. 2 numerous hydrogen bonds
As a result of the molecules being united as a whole, the top of a body of water has a membrane that is able to withstand a small downwards force before becoming pierced. The classic example of the organism that takes advantage of this phenomenon is the pond skater insect, which is able to utilize the high surface tension of lakes and ponds by walking on the surface of the water to look for prey.
The other hydrogen-bonding attribute that benefits marine life especially but not exclusively is the small relative density of ice. While most liquids become denser when they change to a solid state, aqueous water gets less dense. This means that a volume of ice has a lighter mass than the same volume of its liquid counterpart and thus ice can float on top of water.
The reason that ice is less dense than aqueous water lies within the microstructure of the molecules. Normally, a solid is denser than a liquid because the particles within a solid are more tightly packed together and thus more particles can fit into a given space. However, within ice, the hydrogen bonds between water molecules create a lattice structure Fig. 3 that increases the distance between the molecules. This means that less water molecules per volume exist in a solid form than as a liquid form, as much as 10% less than water at 4C [Campbell and Reece, 2008].
Fig. 3 Lattice structure
Several advantages of the small relative density of ice can be observed within marine environments, such as the heat insulation that a surface sheet of ice provides and the fact that bodies of water never freeze from the bottom upwards, two essential phenomena without which many aquatic organisms would not be able to survive. It is also easy to forget that sea ice is also a habitat for sub-terrestrial organisms such as penguins and polar bears, a habitat that would not exist if ice did not float.
Moving away from how hydrogen bonds benefit marine life and towards how they benefit terrestrial life, and having previously stated that water is very good at keeping the body warm at 37C, water is controversially a very good coolant. It is for the same reason that water is able to restrict temperature fluctuations that it is able to cool off the human body and other mammals. Resisting temperature change involves water’s high specific heat capacity and is the large relative energy required to change the temperature of 1g of water by 1C.
This can be related to the high latent heat of vaporization, the phenomenon that makes cooling so effective, since latent heat is the heat energy lost from the body to evaporate 1g of sweat. In water the latent heat is particularly high because of the hydrogen bonds between the water molecules must be broken before liquid can change state into a gas, so more energy is required to evaporate the water and thus more heat is removed. Hydrogen bonds also play a large role in the evaporation of water from plants, called transpiration.
Transpiration is the loss of water from the leaves of plants by evaporation and causes water to move into the roots up vascular tubes (xylem) within in the plant stem. This means a plant can transport water around its tissues for use in respiration and other metabolic reactions. Hydrogen bonds play a large role in transpiration in the same way that they do in the cohesion of water molecules to cause high surface tension. When water molecules are moved up the xylem vessels, they move as a whole due to the hydrogen bonding cohesive forces between the molecules. The molecules also stick to the walls of the vessels by hydrogen bonds, aiding the transport of the water furthermore. The movement of water aided by cohesion and adhesion is known as mass flow, and is the same occurrence as with sucking water through a straw. Since plants form the basis of most ecosystems as producers, hydrogen bonding plays a key part of life.
Concluding, it is clear to see that without hydrogen bonds, life, as we know it today would not exist. Water makes up most of the earths surface and is perhaps the molecule that is the most essential for life- providing stable habitats for marine and terrestrial organisms as well as the transport of water in plants. The fact that most of the properties water are caused by hydrogen bonding shows how essential hydrogen bonds are for life on earth.
Libes (2009). Introduction to Marine Biogeochemistry. Elsevier Science and Technology. Pp28 Campbell and Reece (2008). Biology. 8th ed. San Francisco: Benjamin Cummings.