Question

Explain the important distinctions between each pair of terms: (a) reaction that goes to completion and reversible reaction; (b) \(K_{\mathrm{c}}\) and \(K_{\mathrm{p}} ;\) (c) reaction quotient (Q) and equilibrium constant expression ( \(K\) ); (d) homogeneous and heterogeneous reaction.

Short answer

(a) A reaction that goes to completion is one where the reactants are completely converted into products while a reversible reaction is one where the reactants form products, which react again to produce reactants. (b) \(K_{\mathrm{c}}\) is the equilibrium constant for reactions based on molar concentrations whereas \(K_{\mathrm{p}}\) refers to the equilibrium constant based on partial pressures of gases. (c) The reaction quotient (Q) is calculated at any point in the reaction whereas the equilibrium constant (\(K\)) is a measure of the concentrations of reactants and products at equilibrium. (d) A homogeneous reaction involves reactants and products in the same phase, whereas a heterogeneous reaction involves them in different phases.

Step-by-step solution

Understanding (a) reaction that goes to completion and reversible reaction

A reaction that goes to completion is one where the reactants are completely converted into products, and essentially no reactant remains at equilibrium. On the other hand, a reversible reaction is one where the reactants form products, which in turn can react together to produce the reactants again. Reversible reactions can reach a state of dynamic equilibrium, where the forward and reverse reactions continue to occur, but at the same rates, so there is no net change in the amounts of reactants and products.

Comprehending (b) \(K_{\mathrm{c}}\) and \(K_{\mathrm{p}}\)

\(K_{\mathrm{c}}\) is the equilibrium constant for reactions based on molar concentrations of reactants and products. It is dimensionless in nature. On the other hand, \(K_{\mathrm{p}}\) refers to the equilibrium constant based on partial pressures of gases in the reaction. It is also dimensionless and is often used when dealing with gaseous reactions.

Clarification of (c) reaction quotient (Q) and equilibrium constant expression (\(K\))

The reaction quotient (Q) is a measure calculated exactly like the equilibrium constant, but it can be calculated at any point in the reaction, not just at equilibrium. On the other hand, the equilibrium constant (\(K\)) is a measure of the concentrations of reactants and products at equilibrium. The value of \(Q\) gives us a prediction about which direction the reaction will proceed to reach equilibrium.

Explanation of (d) homogeneous and heterogeneous reaction

A homogeneous reaction is one in which all reactants and products are in the same phase (i.e., solid, liquid, gas, or aqueous solution). Conversely, a heterogeneous reaction is one in which the reactants and products are in different phases.

Key concepts

Reaction Completion vs Reversible Reaction

Understanding the distinction between a reaction that goes to completion and a reversible reaction is fundamental in the study of chemical reactions. When a reaction goes to completion, the reactants are entirely converted into the products, and there is no backward process. This is common in situations where the products leave the reaction mixture, such as in the formation of a precipitate or gas that escapes.
Reversible reactions, on the other hand, do not simply stop once the reactants have been turned into products. Instead, the products can themselves react, regenerating the original reactants. This creates a dynamic situation known as chemical equilibrium, where the forward reaction (reactants to products) and the reverse reaction (products to reactants) occur at the same rate. As a result, the concentrations of the reactants and products remain constant over time, not because the reactions have stopped, but because they move forward and reverse at identical rates.

Equilibrium Constants (Kc and Kp)

Equilibrium constants are crucial for quantifying the position of equilibrium in a chemical reaction. There are two main types of equilibrium constants: Kc and Kp. The equilibrium constant Kc is defined specifically for reactions in solution and looks at the molar concentrations of the reactants and products. It is a dimensionless number, giving you the ratio of product concentrations to reactant concentrations at equilibrium for a balanced chemical equation.

Kp, on the other hand, is used when dealing with gaseous reactions and is determined by the partial pressures of the gases involved. While both Kc and Kp are dimensionless and describe the position of equilibrium, they are applied to different states of matter and are calculated using concentrations for Kc and pressures for Kp. Understanding these constants helps predict the extent to which a reaction will proceed under specified conditions.

Reaction Quotient (Q)

The reaction quotient, represented as Q, plays a pivotal role in determining the direction a chemical reaction is proceeding at any given moment. Unlike the equilibrium constant (K), which is exclusive to the state of equilibrium, Q can be calculated at any point during a reaction. It is determined using the same formula as K, which includes the concentrations or partial pressures of the reactants and products involved.

However, the key difference lies in timing: Q is calculated before equilibrium is established, while K is calculated when the reaction has reached equilibrium. By comparing the value of Q to the value of K, one can predict whether a reaction will shift to the right, forming more products (if Q < K), remain unchanged (if Q = K), or shift to the left, forming more reactants (if Q > K). This predictive ability of Q is essential for understanding and controlling chemical reaction systems.

Homogeneous vs Heterogeneous Reactions

Chemical reactions can be classified depending on the physical state of the reactants and products involved. Homogeneous reactions involve substances that are in the same phase, such as all gases, all liquids, or all solids. For example, the reaction between gases or dissolved substances in a solution is considered homogeneous because the reactants and products are uniformly distributed and are in the same phase.

Heterogeneous reactions, in contrast, involve reactants and products in different phases, such as a solid reacting with a gas or a liquid reacting with a solid. These types of reactions occur at the interface between different phases, and their rates can be influenced by surface area and other factors not present in homogeneous reactions. Understanding whether a reaction is homogeneous or heterogeneous is crucial for manipulating conditions to optimize the reaction rate and yield.