Introduction
The feasibility of a chemical reaction refers to its ability to proceed spontaneously under a given set of conditions without external intervention. Early chemists believed that exothermicity alone dictated spontaneity, but many endothermic reactions, such as the dissolution of ammonium nitrate in water, occur readily. This essay argues that a complete understanding requires the interplay of two fundamental thermodynamic concepts: energetics (enthalpy change) and entropy (disorder change) . These are unified in the Gibbs free energy equation, ΔG = ΔH – TΔS, which provides the definitive criterion for reaction feasibility. By examining both enthalpy and entropy contributions, and considering temperature dependence, this essay will demonstrate how thermodynamics explains why some reactions are spontaneous and others are not.
Enthalpy and the Role of Energy Changes
The enthalpy change (ΔH) of a reaction reflects the heat absorbed or released at constant pressure. Exothermic reactions (ΔH < 0) release energy to the surroundings, making them energetically favourable. For example, the combustion of methane releases 890 kJ mol⁻¹, a strongly negative ΔH that drives the reaction forward. Conversely, endothermic reactions (ΔH > 0) absorb energy, and on their own appear unfavourable.
However, enthalpy alone is insufficient to predict spontaneity. The second law of thermodynamics states that the total entropy of the universe must increase for a process to be spontaneous. An exothermic reaction increases the entropy of the surroundings by transferring heat, which can offset a decrease in the entropy of the system. But if the system’s entropy decreases too much, even a negative ΔH may not guarantee spontaneity. Thus, we must consider the entropy change of the system itself.
Entropy: A Measure of Disorder
Entropy (S) quantifies the degree of randomness or disorder within a system. Reactions that produce more gas molecules from solids or liquids generally have positive entropy changes (ΔS > 0). For instance, the decomposition of calcium carbonate (CaCO₃ → CaO + CO₂) yields a gas, increasing disorder: ΔS ≈ +160 J K⁻¹ mol⁻¹ at 298 K.
The second law requires ΔStotal = ΔSsystem + ΔSsurroundings > 0 for spontaneity. The entropy change of the surroundings is directly related to the enthalpy change: ΔSsurroundings = –ΔH/T. Combining these yields the Gibbs free energy equation. Therefore, even an endothermic reaction with a positive ΔH can be spontaneous if the increase in system entropy is sufficiently large to overcome the negative contribution from the surroundings (Atkins & de Paula, 2014).
Gibbs Free Energy: The Ultimate Criterion
The Gibbs free energy change (ΔG) combines ΔH and ΔS into a single value: ΔG = ΔH – TΔS. A reaction is spontaneous (feasible) at constant temperature and pressure when ΔG < 0. This equation reveals three key scenarios:
- Exothermic with increasing entropy (ΔH < 0, ΔS > 0): ΔG is always negative at all temperatures. Example: combustion of hydrogen.
- Endothermic with decreasing entropy (ΔH > 0, ΔS < 0): ΔG is always positive; the reaction is never spontaneous. Example: formation of ozone from oxygen at room temperature.
- Endothermic with increasing entropy (ΔH > 0, ΔS > 0): Spontaneity depends on temperature. At high T, the –TΔS term dominates, making ΔG negative. Example: dissolving NaCl in water (slightly endothermic but entropy-driven).
This temperature dependence explains many industrial processes. The production of ammonia in the Haber process (N₂ + 3H₂ → 2NH₃) is exothermic (ΔH < 0) but also involves a decrease in entropy (ΔS < 0) because four gas molecules become two. Lower temperatures favour exothermic reactions, but the decreased entropy makes ΔG less negative; higher temperatures increase the –TΔS term, making ΔG positive. Therefore, the Haber process operates at a compromise temperature (~400–450 °C) to balance rate and yield. This principle is closely related to the concept of equilibrium discussed in Evaluate the Significance of Equilibrium Principles in Industrial Chemical Processes.
Factors Affecting Feasibility: Pressure and Concentration
While ΔG under standard conditions (ΔG° ) indicates feasibility at 1 bar and 1 M, real reactions often deviate. The Nernst equation and van’t Hoff isotherm show that ΔG = ΔG° + RT ln Q. Changing the reaction quotient Q by varying pressure or concentration can shift ΔG. For example, increasing pressure in the Haber process favours the side with fewer moles of gas, making ΔG more negative. This thermodynamic control underpins many industrial strategies, including the use of catalysts to lower activation energy without altering ΔG. Catalysis is further explored in Discuss the Role of Catalysis in Chemical Reactions and Its Importance in Modern Industry.
Real-World Examples and Applications
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Photosynthesis: 6CO₂ + 6H₂O → C₆H₁₂O₆ + 6O₂. This is highly endothermic (ΔH > 0) and decreases entropy, yet it is driven forward by the constant input of sunlight. The Gibbs free energy change is positive under standard conditions; sunlight provides the energy to overcome the thermodynamic barrier.
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Explosive decomposition of nitroglycerine: Releases a large volume of gas (huge ΔS > 0) and is strongly exothermic, making ΔG very negative.
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Melting of ice: At 0 °C, ΔH > 0 (endothermic) but ΔS > 0. At T > 273 K, ΔG becomes negative and ice melts spontaneously. This illustrates the temperature-dependent crossover.
Understanding these principles is also central to green chemistry, where the aim is to design reactions with minimal energy input and waste. The development of green chemistry principles is discussed in Discuss How the Development of Green Chemistry Principles Is Influencing Contemporary Chemical Research and Manufacturing.
Conclusion
The feasibility of chemical reactions cannot be explained by enthalpy alone. The concepts of energetics (ΔH) and entropy (ΔS) are intertwined through the Gibbs free energy equation, which provides a robust criterion for spontaneity. Exothermicity often drives reactions, but a large positive entropy change can override an unfavourable enthalpy. Temperature, pressure, and concentration further modulate ΔG, allowing chemists to control reaction conditions. Mastery of these thermodynamic principles is essential for predicting reaction outcomes and optimising industrial processes, from ammonia synthesis to biochemical pathways.
Further Resources for A Level Chemistry Students
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Frequently Asked Questions
Why is enthalpy alone not sufficient to predict spontaneity?
Enthalpy change measures heat transfer, but the second law of thermodynamics requires the total entropy of the universe to increase. An endothermic reaction can be spontaneous if its entropy increase outweighs the heat absorbed from the surroundings. Gibbs free energy combines both factors.
How does temperature affect the feasibility of reactions?
Temperature appears in the Gibbs free energy equation as the factor multiplying entropy change. For reactions where ΔS > 0, increasing temperature makes ΔG more negative, favouring spontaneity. Conversely, if ΔS < 0, higher temperatures make ΔG less negative or positive.
What is the relationship between Gibbs free energy and equilibrium?
At equilibrium, ΔG = 0 and ΔG° = –RT ln K. The equilibrium constant K indicates the extent of reaction. A negative ΔG° implies K > 1, meaning products are favoured at equilibrium. This connection is essential for industrial processes.
Can a reaction with a positive ΔG be made to occur?
Yes, by coupling it with a highly favourable reaction (e.g., ATP hydrolysis in biochemistry) or by providing external energy (e.g., electrolysis). Such non-spontaneous reactions are possible but require continuous energy input.
References
Atkins, P. and de Paula, J. (2014) Physical Chemistry: Thermodynamics, Structure, and Change. 10th edn. Oxford: Oxford University Press.
Burrows, A., Holman, J., Parsons, A., Pilling, G. and Price, G. (2017) Chemistry³: Introducing Inorganic, Organic and Physical Chemistry. 3rd edn. Oxford: Oxford University Press.
Ebbing, D.D. and Gammon, S.D. (2016) General Chemistry. 11th edn. Boston: Cengage Learning.
Whitten, K.W., Davis, R.E., Peck, M.L. and Stanley, G.G. (2013) Chemistry. 10th edn. Belmont: Brooks/Cole.
