Oxidation-Reduction Reactions in Living Organisms

In this article we will discuss about oxidation-reduction reactions in living organisms.

In biological systems energy is generally released from oxidation-reduction reactions of organic compounds catalysed by enzymes. It is, therefore, necessary to understand clearly the nature of oxidation-reduction reactions occurring in living systems. The most common organic compound used as substrate for energy production is glucose.

When one mole of glucose equivalent to 180g is burnt in air under non-biological conditions, 674 Kcal of energy is liberated as heat and the products are CO₂ and H₂O. When glucose is used as substrate for aerobic respiration by an organism, the same amount of energy is liberated and the products are CO₂ and H₂O (C₆H₁₂O₆ = 6CO₂ + 6H₂O + 674 Kcal). But in respiration only a part of the liberated energy is converted to chemical energy (ATP) and the rest is lost as heat.

The question may arise wherein the glucose molecule energy is hidden. Glucose or any other molecule has the energy stored in the electrons that form the chemical bonds between the atoms constituting the molecule. Energy is required for binding the atoms together and it is stored in the chemical bonds as potential energy. When these bonds are broken, energy is liberated in usable form.

The amount of energy that can be set free from a molecule is known as its free energy, conventionally designated as G. In a hypothetical reaction, if the reactants A and B produce products C and D then the difference between the total free energy of A and B and the total free energy of C and D is called the free-energy change which is designated as ∆G.

When the reactants A and B produce C and D with liberation of energy, ∆G is conventionally given a negative sign (-∆G) which means the reaction is exergonic. Under standard conditions, i.e. when both reactants and products are present in 1 molar concentration at a temperature of 25°C under 1 atm. pressure, ∆G is represented as ∆G⁰ which is standard free energy change.

It should be made clear that a chemical or biochemical reaction can run spontaneously only when it is exergonic i.e. ∆G is negative. In other words, a reaction will run spontaneously till the free energy change (∆G) assumes the value of zero and in that stage the reaction will reach the equilibrium. At equilibrium, the concentrations of the reactants and products are such that the total free energy contents of the reactants and products are equal i.e. ∆G = 0. At this stage, the concentration of the products is higher than that of the reactants.

Standard free energy change (∆G°) of a reaction can be calculated from the equation AG°=-RT InK where R is the gas constant having a value of 1.987, T is the absolute temperature (273 + °C) and InK is the natural logarithm of the equilibrium constant of the reaction (2.303 log K). In the hypothetical reaction A+B , <===> C+D, K is calculated from the ratio of the products of concentrations of C+D and that of A + B i.e. [C] x [D] /[A]x[B] = K. If this ratio is greater than 1, ∆G° becomes negative. From the above equation, AG under physiological conditions can be calculated from the relation AG = AG° + RT InK.

Another important parameter of oxidation-reduction reactions is the redox-potential which gives a quantitative measure of the tendency of a compound or an element to lose electrons. In an oxidation- reduction reaction, electrons donated by one compound or element are accepted by another compound or element. As a result, the donor becomes oxidized (loss of electrons) and the acceptor becomes reduced (addition of electrons).

So, oxidation and reduction always go hand in hand. Such pair of reactions is generally known as redox reactions. In most biological systems, the redox reactions involve removal or addition of hydrogen atoms i.e. a proton and an electron. Such reactions are catalyzed by the enzymes called dehydrogenases.

The tendency of a compound to lose electrons is expressed in relation to the tendency of molecular hydrogen to lose electrons i.e. H₂ <===> 2H⁺ + 2e^–. For biological systems, the redox-potential is expressed as standard reduction potential or electromotive force (emf) in volts when the reactant and oxidant are present in 1.0 M concentration at pH 7.0 and 25°C. The standard redox-potential of H₂ —> 2H⁺ + 2e^– is -0.42v. The more positive the redox-potential is, greater is the tendency to lose electrons i.e. greater is the oxidizing ability.

In living organisms, the energy released by exergonic reactions is utilized to drive endergonic reactions and also for other purposes, such as locomotion. The transfer of energy from one reaction to another takes place via some common reactants which take part in both the exergonic and the endergonic reactions.

These common reactants are characterized by a high transfer potential and they are called energy-rich compounds. The most common of such compounds is ATP. The last phosphate group of ATP is bound to the rest of the molecule by an unstable bond and can be quickly separated or transferred to an acceptor with release of large amount of free energy.

Thus, ATP has a high transfer potential and can take part in many biochemical reactions.

ATP molecules are comparatively larger in size and, therefore, they are unsuitable for storing energy. For storing energy, the living systems use smaller molecules like glucose where the energy is stored in the chemical bonds. So, the potential energy of glucose has to be released through breakdown of the molecule in a stepwise manner by catabolic reactions. In living systems, glucose can be catabolized by several pathways.