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Principles of Thermodynamics in Biochemistry

In this principles of thermodynamics in biochemistry post we have briefly explained about principles of thermodynamics, 1st law of thermodynamics, 2nd law of thermodynamics, gibb’s free energy concept, standard free energy change, three types of reactions.  

Principles of Thermodynamics

Thermodynamics is concerned with the flow of heat and it deals with the relationship between heat and work. Bioenergetics, or biochemical thermodynamics, is the study of the energy changes accompanying biochemical reactions. Biological systems use chemical energy to power living processes.

1st Law

The first law of thermodynamics in principles of thermodynamics states that energy can neither be created nor destroyed only altered in form.  The total energy of a system, including its surroundings, remains constant. Or, ΔE = Q – W, where Q is the heat absorbed by the system and W is the work done. 

This is also called the law of conservation of energy. If heat is transformed into work, there is proportionality between the work obtained and the heat dissipated. A system is an object or a quantity of matter, chosen for observation. All other parts of the universe, outside the boundary of the system, are called the surroundings.

2nd Law

The total entropy of a system must increase if a process is to occur spontaneously. A reaction occurs spontaneously if ΔE is negative, or if the entropy of the system increases. Entropy (S) is a measure of the degree of randomness or disorder of a system. Entropy becomes maximum in a system as it approaches true equilibrium. Enthalpy is the heat content of a system and entropy is that fraction of enthalpy which is not available to do useful work. 

A closed system approaches a state of equilibrium. Any system can spontaneously proceed from a state of low probability (ordered state) to a state of high probability (disordered state). The entropy of a system may decrease with an increase in that of the surroundings. The second law may be expressed in simple terms as Q = T x ΔS, where Q is the heat absorbed, T is the absolute temperature and ΔS is the change in entropy.

Gibb’s Free Energy

The term free energy is used to get an equation combining the first and second laws of thermodynamics. Thus, ΔG = ΔH – TΔS , where ΔG is the change in free energy, ΔH is the change in enthalpy or heat content of the system and ΔS is the change in entropy. The term free energy denotes a portion of the total energy change in a system that is available for doing work. For most biochemical reactions, it is seen that ΔH is nearly equal to ΔE. So, ΔG = ΔE – TΔS. Hence, ΔG or free energy of a system depends on the change in internal energy and change in entropy of a system.

Standard Free Energy

It is the free energy change under standard conditions. It is designated as ΔG0. The standard conditions are defined for biochemical reactions at a pH of 7 and 1 M concentration, and differentiated by a priming sign ΔG0. It is directly related to the equilibrium constant. Actual free energy changes depend on reactant and product.

Most of the reversible metabolic reactions are near equilibrium reactions and therefore their ΔG is nearly zero. The net rate of near equilibrium reactions are effectively regulated by the relative concentration of substrates and products. The metabolic reactions that function far from equilibrium are irreversible. The velocity of these reactions are altered by changes in enzyme activity. A highly exergonic reaction is irreversible and goes to completion. Such a reaction that is part of a metabolic pathway confers direction to the pathway and makes the entire pathway irreversible.

Three Types of Reactions

A reaction can occur spontaneously when ΔG is negative. Then the reaction is exergonic. If ΔG is of great magnitude, the reaction goes to completion and is essentially irreversible.

When ΔG is zero, the system is at equilibrium. The change in free energy (ΔG) is the difference between the heat released during a process and the heat released for the same process occurring in a reversible manner. If a system is at equilibrium, ΔG = 0. If the process is spontaneous, ΔG < 0.

For reactions where the ΔG is positive, an input of energy is required to drive the reaction. The reaction is termed as endergonic and those with a negative ΔG as exergonic. Similarly a reaction may be exothermic (ΔH is negative), isothermic (ΔH is zero) or endothermic (ΔH is positive).

Energetically unfavourable reaction may be driven forward by coupling it with a favourable reaction.

Glucose + Pi → Glucose-6-phosphate (reaction1)

ATP + H2O → ADP + Pi (reaction 2)

Glucose+ATP→Glucose-6-phosphate+ADP (3)

Reaction 1 cannot proceed spontaneously. But the 2nd reaction is coupled in the body, so that the reaction becomes possible. For the first reaction, ΔG0 is +13.8 kJ/mole; for the second reaction, ΔG0 is –30.5 kJ/mole. When the two reactions are coupled in the reaction 3, the ΔG0 becomes –16.7 kJ/mole, and hence the reaction becomes possible.

Reactions of catabolic pathways (degradation or oxidation of fuel molecules) are usually exergonic, whereas anabolic pathways (synthetic reactions or building up of compounds) are endergonic. Metabolism constitutes anabolic and catabolic processes that are well co-ordinated.

Further Readings

Reference

Principles of Thermodynamics in Biochemistry