Assessing Spontaneity — Use ΔG, ΔH, and ΔS to assess spontaneity

The question of why certain chemical reactions proceed on their own while others require continuous energy input has occupied scientists for centuries. Early investigators like Antoine Lavoisier focused almost exclusively on heat transfer, assuming that exothermic reactions were inherently favored by nature. This view seemed intuitive—burning wood releases heat, iron rusts with a warm glow—but it failed spectacularly for endothermic processes that nonetheless occur spontaneously, such as the dissolution of ammonium nitrate in water, which absorbs heat yet proceeds without any external driving force. The resolution of this puzzle required an entirely new thermodynamic quantity that could account for both energy and disorder simultaneously.

The central question that Gibbs resolved was deceptively simple: given a chemical reaction at constant temperature and pressure, will it proceed in the forward direction without outside intervention? Answering this requires balancing two competing thermodynamic tendencies—the drive toward lower enthalpy and the drive toward greater entropy—within a single criterion, ΔG. The remainder of this lesson develops that criterion rigorously and shows how to apply it.

Before diving into calculations, it is essential to establish precise definitions for the three state functions that govern spontaneity. Each captures a different facet of a system's thermodynamic behavior: energy exchange with the surroundings, molecular randomness, and the net thermodynamic driving force. Together they form the foundation of the Gibbs equation, which serves as the universal criterion for spontaneity under laboratory conditions.

The sign of ΔG depends on the interplay between ΔH, ΔS, and T. Because each of ΔH and ΔS can be positive or negative, there are four distinct thermodynamic scenarios. The diagram below maps these scenarios onto a two-by-two grid, showing how the sign of ΔG changes with temperature in each case. Pay particular attention to the two temperature-dependent cases, where a crossover temperature exists at which ΔG switches sign.

The diagram makes clear that enthalpy and entropy are not independent criteria—they compete through the temperature term. In Cases 2 and 3, there exists a crossover temperature (T_crossover = ΔH / ΔS) at which ΔG transitions from negative to positive or vice versa. Below this temperature, the enthalpy term dominates; above it, the entropy term takes control. Recognizing which case applies to a given reaction is the first step in any spontaneity assessment.

The mathematical framework for assessing spontaneity begins with the fundamental Gibbs equation and extends to standard-state calculations and the relationship between ΔG° and the equilibrium constant. Each equation below is presented with its assumptions and the physical meaning of every variable.

The Gibbs equation can be viewed as a linear function of temperature: ΔG = ΔH − TΔS is analogous to y = b + mx, where ΔG plays the role of y, T is the independent variable x, −ΔS is the slope, and ΔH is the y-intercept. This perspective makes it straightforward to visualize how ΔG varies with temperature for each of the four cases, and it reveals the crossover temperature geometrically as the x-intercept of the line.

When confronted with a new reaction, the strategy is straightforward: determine the signs of ΔH and ΔS, identify the applicable case, and—if the case is temperature-dependent—calculate the crossover temperature. If the question specifies a particular temperature, substitute into ΔG = ΔH − TΔS and check the sign directly. The linear relationship ensures that these calculations are algebraically simple, but unit vigilance is essential: always convert ΔS from J/(mol·K) to kJ/(mol·K) or ΔH from kJ to J before combining terms.

Consider the thermal decomposition of calcium carbonate, a classic example of a Case 3 reaction:

CaCO₃(s) → CaO(s) + CO₂(g)

Given: ΔH° = +178.3 kJ/mol, ΔS° = +160.5 J/(mol·K). Determine (a) whether the reaction is spontaneous at 25 °C, (b) the crossover temperature, and (c) whether it is spontaneous at 1500 K.

A common misconception is that a spontaneous reaction must occur quickly. In reality, thermodynamics tells us whether a reaction can occur, not how fast it will proceed. The distinction between thermodynamic favorability and kinetic feasibility is one of the most important in all of chemistry. The table below contrasts the capabilities and limitations of the Gibbs free energy criterion with kinetic considerations.

The standard Gibbs free energy (ΔG°) applies strictly at standard-state concentrations and pressures (1 M, 1 atm, 298 K unless otherwise specified). Real systems rarely operate under these conditions, so chemistry extends the framework to non-standard states using the reaction quotient Q. The relationship ΔG = ΔG° + RT ln Q adjusts the free energy for actual concentrations, connecting the thermodynamic criterion to real-world chemical conditions. Furthermore, the equation ΔG° = −RT ln K provides a direct bridge between free energy and the equilibrium constant, enabling prediction of equilibrium position from thermodynamic data alone.

In more advanced courses—physical chemistry, biochemistry, and electrochemistry—the non-standard form ΔG = ΔG° + RT ln Q becomes indispensable. In electrochemistry, it gives rise to the Nernst equation (E = E° − (RT/nF) ln Q), which adjusts cell potentials for non-standard concentrations. In biochemistry, a modified standard state (ΔG°', pH 7) is used because biological reactions occur under physiological conditions rather than pure standard states. Understanding the standard-state Gibbs equation thoroughly is therefore not merely an end in itself but a gateway to these more specialized applications.