Acid-base reactions: Kinetics and Thermodynamics

Let's restate that kinetics and thermodynamics are the two crucial elements in any chemical reaction. The rate of the reaction is governed by kinetics, and its viability is determined by thermodynamics.

Kinetics

To put it more precisely, when A + B --- > C, two reactions are always in action: one is the forward reaction that creates the product C, and the other is the reverse reaction that converts C into A + B at two different rate constants, K1 and K2. When K1 > K2, there is a forward reaction. Let us expand this important point a little more.

When the rate constant of the forward reaction (K1) is greater than the rate constant of the reverse reaction (K2), the forward reaction dominates, and the reaction proceeds predominantly in the forward direction. Conversely, when the rate constant of the reverse reaction (K2) is greater than the rate constant of the forward reaction (K1), the reverse reaction dominates, and the reaction proceeds predominantly in the reverse direction. When the rate constants of the forward and reverse reactions are equal (K1 = K2), the system is at equilibrium, and both forward and reverse reactions occur at the same rate.

Thermodynamics

According to thermodynamics, for a chemical reaction to occur spontaneously, internal energy must be consumed, as shown by the formula dG = dU - TdS. The reaction's free energy, denoted by the symbol dG, must be negative. This means that the reaction's term TdS must be greater than the change in internal energy, dU. There are no exceptions to acid-base reactions. This post's objective is to explain in more detail how these factors affect acid-base reactions.

To summarize, neutralization reactions involving weak acids and bases are primarily thermodynamically controlled, meaning that the equilibrium position of the reaction is determined by the thermodynamic properties of the reactants and products. The pH of the solution at equilibrium is determined by the dissociation constants of the weak acid and base.

On the other hand, neutralization reactions involving strong acids and bases are primarily kinetics controlled. The reaction rate is influenced by the concentrations of the H+ and OH- ions in the system, as these ions are responsible for the reaction.

The specifics of acid-base reactions

Weak acid-base reactions are more fundamental in nature because they do not go to completion. Unlike strong acid-base reactions where the reactants are completely converted into products, weak acid-base reactions only partially convert the reactants into products, resulting in an equilibrium state. This equilibrium allows for the coexistence of both reactants and products, and the reaction can be influenced by factors such as concentration and temperature. Understanding these equilibrium-based reactions is important as they have widespread applications in processes. To summarize, in strong acid-base reactions, the equilibrium strongly favours the formation of products, while in weak acid-base reactions, the equilibrium lies more towards the reactants. This is because weak acid-base reactions involve weak acids or bases, which only partially dissociate in solution.

Remember acid-base reactions are not simple reactions

If a weak acid or weak base has infiltrated your process, it may not be possible to neutralize them completely to pH 7. The presence of these compounds can contribute to corrosion. In order to prevent or eliminate corrosion, it may be necessary to physically remove the weak acid or weak base from the system.

Acid-base reactions, although they may seem simple, are influenced by both kinetics and thermodynamics. Understanding the kinetics and thermodynamics of these reactions is crucial in understanding their effects on corrosion processes and other chemical reactions.

The reason why it may be difficult to neutralize a weak acid or weak base to pH 7 is because their dissociation constants (Ka or Kb) are relatively weak compared to strong acids or bases. Weak acids and bases only partially dissociate in water, resulting in a limited number of hydrogen (H+) or hydroxide (OH-) ions being produced.

When attempting to neutralize a weak acid with a strong base or vice versa, the addition of a strong base to a weak acid or a strong acid to a weak base may result in excess H+ or OH- ions remaining in the solution. This residual concentration of H+ or OH- ions, although relatively small, can still affect the overall pH of the solution, preventing it from reaching a pH of exactly 7.

Let us take a few specific cases

Thermodynamics controlled nonequilibrium reactions

[1] A gas phase reaction NH3 [g] + HCl [g] < = > NH4Cl [s]

This is a typical reaction that can occur in a distillation column. This reaction would never go to completion and favor a reverse reaction because of the gain in entropy.

In this case, the formation of a solid compound can indeed contribute to decreasing the overall entropy and potentially driving the reaction in the reverse direction and making the Gibbs free energy go positive.

[2] A gas -liquid phase reaction CO2 [g] + NaOH{aq] <= > NaHCO3 [aq]

?This is yet another thermodynamically non-favorable reaction. in a non-equilibrium situation where the entropy decreases [ 2 reactants form 1 product] The Gibbs free energy becomes positive. To reiterate, in cases where the entropy decreases and the Gibbs free energy becomes positive, it signifies that the reaction is not favored energetically and will not proceed in the forward direction. The reaction may require external intervention, such as changes in temperature, pressure, or the addition of catalysts, to drive it in the desired direction.

?Kinetics controlled reactions

[3] A liquid-liquid equilibrium reaction NH4OH [aq] + HCl [aq] < = > NH4Cl [aq]

This is a neutralization reaction in the aqueous phase between NH4OH and HCl that does not go to completion. This is a more kinetics-controlled reaction. The equilibrium constant (K) relates the concentrations of the reactants and products at equilibrium. The values of K1 and K2 represent the forward and reverse rate constants respectively.

If K2 is greater than K1, at equilibrium it implies that the reverse reaction is favored over the forward reaction at equilibrium. This indicates that the reaction does not proceed to completion, and the system is not at equilibrium. In such cases, the reaction is indeed more kinetics controlled, as the rate of the reverse reaction is higher compared to the rate of the forward reaction.