Biological systems harness energy to perform the functions of life, including metabolism, growth, and reproduction, all work performed through chemical reactions. Thermodynamics addresses the flow of energy in chemical reactions, and, specifically in bioenergetics, the flow of energy in the chemical interactions that make up biological systems.
Free energy (Gibbs free energy) represents the work available from chemical reactions, capable of being harnessed to perform the functions of living organisms, and is represented in the equation ΔG = ΔH - TΔS. Taken together the aspects of this equation (change in enthalpy, temperature, and change in entropy) determine whether a reaction is thermodynamically favorable (spontaneous in the forward direction).
Negative ΔH (change in enthalpy), heat flow out of a system, brings about more stable, lower energy products from reactants.
Positive ΔS (multiplied in effect by temperature) indicates a thermodynamically favorable increase in disorder of the system.
Because entropy plays a role in the favorability of the forward reaction, changes in entropy brought about by concentration changes of reactants and products will tend to drive the reaction forwards or backwards depending on their effect. Entropy will be maximized at specific concentrations of reactants and products, a point at which the reaction is at equilibrium (equal rates forwards and backwards). These concentrations of reactants and products can then be used to calculate the temperature-specific constant Keq for the reaction.
The change in Gibbs free energy thus is largely dependent on the nature of a reaction's reactants and products and their concentrations.
Keq, the equilibrium constant, is a ratio calculated from the concentrations of reactants and products of a reaction at equilibrium and is temperature dependent. For reaction
aA + bB ⇄ xX + yY at a given temperature,
Keq = ([X]x[Y]y) / ([A]a[B]b).
As a given value of Gibbs free energy is dependent on numerous mutable factors such as temperature and reactant and product concentrations, it is useful to make reference to the value under standard-state conditions, that is standard-state free energy ΔG°. (Under standard-state, temperature is 298K and all concentrations are 1M.)
At standard state, most reactions are not at equilibrium. If ΔG° is negative, then the reaction is favorable in the forward direction, proceeding forward until reaching equilibrium, that is favoring products with a Keq > 1. If ΔG° is positive, then the reaction is unfavorable and equilibrium is to be achieved by the reverse reaction, that is favoring reactants with a Keq < 1. (A reaction at equilibrium at standard state will have a Keq = 1.)
Concentrations of reactants and products can play an influential role in determining thermodynamic favorability due to their relationship with entropy. Adding or removing reactants or products therefore affects reaction spontaneity, driving directionality.
Le Châtelier's Principle, which holds that systems at equilibrium will respond to an applied stress with shifts that reduce the stress, succinctly captures the impact changes in concentration can have on a reaction. Following this principle, adding reactant to a reaction at equilibrium (increasing reactant concentration, decreasing product concentration) can drive a reaction forward to produce more product, thereby reducing reactant and correcting the concentration stress. A similar situation occurs by removing product (drives reaction forward), or conversely, with removing reactant and adding product (drives reaction in reverse). This principle also governs changes in temperature, pressure, and volume.
Endothermic and exothermic are designations that describe the change in enthalpy for a reaction. Endothermic reactions require the transfer of heat into the system (+ΔH), while exothermic reactions transfer heat out of the system (-ΔH).
|Transfer of heat||in||out|
|Change in enthalpy||+ΔH||-ΔH|
|Spontaneity of reaction||depends!||depends!|
A reaction being endothermic or exothermic does not alone determine whether it is thermodynamically favorable. Change in entropy along with temperature also plays a role. Overall thermodynamic favorability of a reaction is represented by change in free energy and related by the terms endergonic and exergonic:
|Change in Gibbs free energy||+ΔG||-ΔG|
|Spontaneity of reaction||non-spontaneous||spontaneous|
Free energy is a concept to account for the ability to do work based on the thermodynamic favorability of changes to a local system. It relates the idea that the differences in enthalpy and entropy between reactants and products can drive chemical reactions. Conceptually, G represents the energy available to the system based on the nature and concentration of reactants/products.
Spontaneous reactions are thermodynamically favorable. At standard state, a reaction with +ΔG° is non-spontaneous, or driven in reverse, and favors reactants (at equilibrium there will be more reactants than products). Conversely, a reaction with -ΔG° is spontaneous, or driven forward, and favors products (at equilibrium there will be more products than reactants).
Phosphoryl group transfers are a class of chemical reaction that are of critical biological importance for creating the energy currency of ATP (as well as in building the backbone of nucleic acids and signal transduction).
The hydrolysis of ATP into ADP + Pi is a very thermodynamically favorable reaction (ΔG << 0). This allows ATP to act as an energy currency by coupling its hydrolysis with thermodynamically unfavorable reactions, thereby granting a combined favorability.
ATP (composed of adenine, ribose, and three phosphate groups) can be hydrolyzed at different positions on the ATP molecule, producing different options for group transfers, which in turn play roles in reaction mechanisms to improve thermodynamic favorability beyond just coupling with the -ΔG of ATP hydrolysis. Depending on nucleophilic attack on the α, β, or γ phosphate, either transfer of a single phosphoryl group, a pyrophosphoryl group, or adenylyl group (the AMP portion with inorganic pyrophosphate released) will take place, respectively.
An expansive degree of biologically important work is performed through oxidation-reduction (redox) reactions, which are a class of chemical reactions that result in a change in oxidation state for atoms of the molecules involved. An increase in an atom's oxidation state is an oxidation, and a decrease in an atom's oxidation state is a reduction. The two processes occur together with one entity (the reducing agent) becoming oxidized in the act of reducing another entity (the oxidizing agent).
Atoms more likely to be reduced are those with higher electronegativity, such as oxygen. It is the reduction of oxygen that lends to the terminology of "oxidation", referencing that an atom of lower electronegativity bound to oxygen will effectively lose electrons to uneven sharing with oxygen, producing a situation referred to as "being oxidized". Of course, oxidation applies to any increase in oxidation state, not just when bound to oxygen. An atom regularly involved in reducing is hydrogen (although its name does not lend itself to the process). In binding with hydrogen, an electron from the hydrogen atom is unevenly shared, effectively gained by the bound atom, creating a reduction.
The mnemonic OIL RIG is used to remember Oxidation is Loss, Reduction is Gain in terms of (generally in a biological context, unevenly shared) electrons.
Specifically, reduction takes place by the gain of electrons procured from the loss of electrons on the entity undergoing oxidation. Thus oxidation-reduction reactions can be broken down into two reactions to better represent these activities of loss and gain.
Each half-reaction can proceed forwards or backwards such that the half-reaction with the greatest reaction potential (strongest pull on electrons, thereby being reduced) will determine the direction of the oxidation-reduction reaction as a whole. (Some redox reactions are more complex with a series of half-reactions.)
Complex biological interactions that involve multiple oxidation-reduction reactions (e.g. the citric acid cycle in relation to the electron transport chain) are aided by soluble electron carriers, molecules that are capable of accepting or donating electrons (often via binding, releasing hydrogen) and can shuttle to/from other acceptors/donors to continue the redox process.
NAD (NAD+ when oxidized, NADH when reduced) accepts electrons (reduced) in glycolysis and the citric acid cycle and donates them (oxidized) to the electron transport chain. FAD (FAD when oxidized, FADH2 when reduced) also accepts electrons (reduced) at the citric acid cycle and donates them (oxidized) to the electron transport chain.
Flavoproteins are proteins that incorporate a co-factor that contains flavin (a riboflavin derivative of which FAD is an example), which greatly aids in enzymatic activity involving oxidation-reduction by its ability to accept one or two electrons. Involvement as a co-factor gives the molecules a regulatory role.
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