Question

Consider the following reactions:\(\mathrm{H}_{2}(g)+\mathrm{I}_{2}(g) \longrightarrow 2 \mathrm{HI}(g) \quad\) and \(\quad \mathrm{H}_{2}(g)+\mathrm{I}_{2}(s) \longrightarrow 2 \mathrm{HI}(g)\).List two property differences between these two reactions that relate to equilibrium.

Short answer

The two property differences relating to equilibrium for the given reactions are: 1. The equilibrium constant expressions are different, as the first reaction's equilibrium constant depends on the concentration of gaseous Iodine (\(K_1 = \frac{[HI]^2}{[H_2][I_2]^1}\)), while the second one does not, due to solid Iodine being present (\(K_2 = \frac{[HI]^2}{[H_2]}\)). 2. Pressure changes will impact the equilibrium of the first reaction (all gases), following Le Chatelier's principle, while they will not have a significant effect on the equilibrium of the second reaction (involving a solid reactant).

Step-by-step solution

Difference 1: Equilibrium Constants

The equilibrium constant K is used to express the concentration of the reactants and products at equilibrium. For the given reactions, the equilibrium constants can be written as: For Reaction 1: \(K_1 = \frac{[HI]^2}{[H_2][I_2]^1}\) For Reaction 2: \(K_2 = \frac{[HI]^2}{[H_2][I_2(S)]}\) The phase of Iodine in the second reaction is solid, whereas it is a gas in the first reaction. The concentration of a solid is not included in the equilibrium expression, so the equilibrium constant for the second reaction does not depend on the concentration of solid Iodine. This leads to a difference in the equilibrium constants between the two reactions, as the first reaction's equilibrium constant depends on the concentration of gaseous Iodine, while the second one does not.

Difference 2: Effect of Pressure Change on Equilibrium

A change in pressure affects the equilibrium of a gaseous reaction since the concentrations of the gases are directly proportional to their partial pressures. According to Le Chatelier's principle, an increase in pressure will favor the side of the reaction with fewer moles of gas, while a decrease in pressure will favor the side with more moles of gas. For Reaction 1, both reactants and products are gases. When the pressure is changed, the equilibrium will shift according to Le Chatelier's principle. However, for Reaction 2, one of the reactants is solid. Changes in pressure have negligible effects on the concentration of solids, so the equilibrium of Reaction 2 will not be significantly affected by pressure changes. In conclusion, the two property differences relating to equilibrium for the given reactions are: 1. The equilibrium constant expressions for the two reactions are different due to the difference in Iodine's phase. 2. Pressure change will affect the equilibrium of the first reaction (all gases), but it will not have a significant impact on the equilibrium of the second reaction (involving a solid reactant).

Key concepts

Equilibrium Constant

In chemical reactions, the equilibrium constant, denoted as \(K\), plays a crucial role in understanding how reactions balance between reactants and products once they reach equilibrium. Specifically, it provides a quantitative measure of the concentrations of products and reactants when a reaction has reached a state of balance. Consider two reactions:

• In Reaction 1: \(\text{H}_2(g) + \text{I}_2(g) \rightarrow 2 \text{HI}(g)\), the equilibrium constant \(K_1 = \frac{[\text{HI}]^2}{[\text{H}_2][\text{I}_2]}\).
• In Reaction 2: \(\text{H}_2(g) + \text{I}_2(s) \rightarrow 2 \text{HI}(g)\), the equilibrium constant \(K_2 = \frac{[\text{HI}]^2}{[\text{H}_2]}\) excludes \(\text{I}_2(s)\) as solids do not appear in equilibrium expressions.

The equilibrium constant is, therefore, unique to each reaction depending on the phases of the substances involved. For Reaction 1, the gaseous iodine \(\text{I}_2(g)\) influences \(K\), while in Reaction 2, solid iodine \(\text{I}_2(s)\) does not. This fundamental difference hinges on the fact that the concentration of a solid remains constant and does not influence the equilibrium constant.

Le Chatelier's Principle

Le Chatelier's principle is a fascinating guideline used in chemistry to predict how a change in conditions will affect chemical equilibrium. It states that if a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium shifts to counteract the change and re-establish equilibrium.In Reaction 1 \((\text{H}_2(g) + \text{I}_2(g) \rightarrow 2 \text{HI}(g))\), both reactants and products are in the gaseous phase. When pressure is altered, equilibrium may shift to favor the side with fewer or more molecules of gas. Increasing pressure shifts equilibrium toward fewer gas molecules and vice versa.For Reaction 2 \((\text{H}_2(g) + \text{I}_2(s) \rightarrow 2 \text{HI}(g))\), only one reactant is gaseous, so changes in pressure don't significantly affect equilibrium. In summary, Reaction 1 is susceptible to pressure changes as it involves only gases, while Reaction 2 is much less so due to the presence of a solid reactant.

Reaction Phase

Phase refers to the physical state of a substance—solid, liquid, or gas—that has a notable impact on how a reaction behaves at equilibrium. In our examples:

• Reaction 1 has both reactants \(\text{H}_2\) and \(\text{I}_2\) as gases, making it homogeneous and susceptible to changes in pressure and volume.
• Reaction 2 includes solid \(\text{I}_2(s)\), resulting in a heterogeneous reaction.

The presence of a solid affects the reaction’s equilibrium equations because solids are often considered constant in concentration. This inclusion or exclusion based on phase results in different equilibrium expressions and different responses to conditions such as pressure changes. Understanding the reaction phase helps predict how a system might behave when measures like pressure and concentration are tweaked.

Pressure Effects

Pressure has a profound impact on reactions involving gases. The relationship between pressure and equilibrium is guided by the principle that reactions favor the side with fewer gas molecules when pressure increases, and the side with more gas molecules when pressure decreases. This is especially clear in Reaction 1:For Reaction 1 \((\text{H}_2(g) + \text{I}_2(g) \rightarrow 2 \text{HI}(g))\), with all reactants and products as gases, changes in pressure directly affect equilibrium. The system will shift to minimize the effects of pressure changes.Reaction 2 contains \(\text{I}_2(s)\), and the presence of a solid means pressure variations don't majorly relieve stress within the system, as solids are not compressible like gases.Ultimately, only gases experience significant effects from pressure changes, and understanding these can aid in predicting how chemical systems behave under different pressures.

Reaction Kinetics

While equilibrium is a state of balance, reaction kinetics deals with the rate at which a reaction proceeds towards equilibrium. Kinetics and equilibrium are distinct, yet they intersect in how fast equilibrium is achieved. Factors such as temperature, concentration, and the presence of a catalyst affect the speed of both Reaction 1 and Reaction 2 but not their ultimate position at equilibrium. For both reactions, altering the concentration of reactants or adding a catalyst can change the speed at which equilibrium is met without affecting the equilibrium constant itself. Temperature changes, however, can affect both the kinetics and the position of equilibrium, highlighting the nuanced interplay between these factors. Understanding reaction kinetics is valuable for predicting how quickly a reaction will reach equilibrium and how conditions might be optimized for specific reaction rates.