The second law of thermodynamics says that entropy doesn't decrease in adiabatic or isolated systems. We use this rule to find out if a process can happen on its own. However, most of the bio systems we encounter are not isolated. The external environment should be added in our system to make an isolated one in order to apply this law. For example, the photons absorbed by plants and heat radiation emitted from reaction centers need to be included as part of such enlarged system.
There are several challenges in this approach. A boundary must be drawn so as to create an isolated system each time. Some details will be omitted during analysis of a complicated process and entropy is too abstract for understanding. Our purpose is to substitute entropy with a familiar physical concept without creating isolation every time. We introduce Gibbs free energy criterion rules to determine whether a reaction can proceed under constant temperature and pressure. It is essential to learn the enthalpy before discussing this.
Enthalpy: Exothermic and Endothermic Reactions
Enthalpy (H) represents internal energy (U) plus product of external pressure and volume (H = U + PV). Like potential energy, enthalpy is also state function which doesn’t depend on path taken by process. We are mainly interested in change rather than absolute value itself(ΔH = ΔU + Δ(PV)). Chemical reactions generally occur at constant pressure within biological systems where heat given off equals ΔH(if pressure changes then ΔQ≠ΔH). Here U stands for system’s internal energy while ΔQ denotes released heat; W refers work done by system ; H=U+PV implies enthalpy of system.
ΔU = ΔQ + ΔW = ΔQ – PΔV ⇒ ΔU + PΔV = ΔQ ⇒ ΔH = ΔQ
The total enthalpy change (ΔH) in exothermic reactions is negative, hence reaction releases thermal energy to surroundings. It is often observed as an increase in temperature or emit of light. An example is fuel combustion like gasoline or wood. The positive ΔH in endothermic reactions shows that it absorbs heat from the surroundings. Ammonium nitrate dissolving in water is a classic example. The ancient people used this process to make cold beverages during summer.
Gibbs energy determines whether a process is spontaneous
We introduce concept called Gibbs free energy (also known as Gibbs function, Gibbs energy). It was named after American physicist Josiah Willard Gibbs who developed this concept (G = H – TS). This function can be applied anywhere but Gibbs criterion only works for isothermal and isobaric systems. Its derivation is very simple.
ΔS≥(ΔQ)/T ⇒ TΔS≥ΔH ⇒ ΔH-TΔS≤0 ⇒ Δ(H-TS)≤0
We obtain ( G = H - TS ) from the new state function ( ΔG = ΔH - TΔS). In this case, ΔG is the change in Gibbs free energy, ΔH is the change in enthalpy, T represents temperature in Kelvin and ΔS refers to the change in entropy. The inequality ΔS≥(ΔQ)/T holds true for any situation and not just isolated systems. We thereby convert the entropy criterion into energy criterion so that it can be applied directly to our system without creating an isolated system each time.
ΔG < 0: The process is spontaneous. The system releases energy and the reaction is exergonic.
ΔG > 0: The process is non-spontaneous. An input of energy is required by the system and this is the endergonic reaction.
ΔG = 0: The system is at equilibrium. There exists no net transfer of energy because zero free energy change.
Discussion on Gibbs Free Energy with Applications and Examples
If ΔH < 0，ΔS > 0, the process is spontaneous. If ΔH > 0，ΔS < 0, the process is not spontaneous. For other cases, temperature must be considered to determine whether a chemical reaction occurs spontaneously. If ΔH < 0、ΔS < 0, reaction tends to occur at low temperatures. If ΔH ＞ 0, ΔS ＞ 0, it tends to occur at high temperatures.
ΔH < 0，ΔS > 0
This includes most reactions such as combustion of solids or liquids. For example, when carbohydrates burn in air, carbon dioxide and water are produced(C₆H₁₂O₆+6O₂→6CO₂+6H₂O). Heat and light indicate it’s an exothermic reaction (ΔH < 0). The products include water and gas, while the reactants are gas and solid, hence ΔS＞0.
ΔH > 0，ΔS < 0
These reactions are not spontaneous but can occur under certain conditions. It happens when they are coupled with a reaction that greatly decreases Gibbs free energy. Water doesn’t decompose into oxygen and hydrogen at room temperature. However, this reaction will take place with the help of electric current, because chemical energy from battery is consumed (the Gibbs free energy of the battery decreases largely). Many non-spontaneous reactions in biology require ATP consumption, as ATP hydrolysis is a reaction that significantly decreases free energy.
ΔH < 0，ΔS < 0
This group includes gases that condense into liquids or burn to make liquids. Burning hydrogen in oxygen is exothermic reaction with a blue flame or explosion, so ΔH < 0. Hydrogen gas turns into liquid, so ΔS < 0. These processes tend to occur at low temperatures.
ΔH > 0，ΔS > 0
An endothermic reaction that happens spontaneously belongs to this category (to satisfy ΔG < 0, ΔS must be positive). One of the instances is ammonium nitrate dissolving in water. Such reactions will probably happen at higher temperatures. For example, calcium carbonate doesn’t break down at room temperature, but it decomposes into carbon dioxide spontaneously when heated up to 1000°C.