Equilibrium Constants Quiz

PDF page 695: The numeric value of Qc for a given reaction varies; it depends on the concentrations of products and reactants present at the time when Qc is determined. When pure reactants are mixed, Qc is initially zero because there are no products present at that point. As the reaction proceeds, the value of Qc increases as the concentrations of the products increase and the concentrations of the reactants simultaneously decrease (Figure 13.6). When the reaction reaches equilibrium, the value of the reaction quotient no longer changes because the concentrations no longer change.

PDF page 697: The magnitude of an equilibrium constant is a measure of the yield of a reaction when it reaches equilibrium. A large value for Kc indicates that equilibrium is attained only after the reactants have been largely converted into products. A small value of Kc—much less than 1—indicates that equilibrium is attained when only a small proportion of the reactants have been converted into products. Once a value of Kc is known for a reaction, it can be used to predict directional shifts when compared to the value of Qc. A system that is not at equilibrium will proceed in the direction that establishes equilibrium. The data in Figure 13.7 illustrate this. When heated to a consistent temperature, 800 °C, different starting mixtures of CO, H2O, CO2, and H2 react to reach compositions adhering to the same equilibrium (the value of Qc changes until it equals the value of Kc). This value is 0.640, the equilibrium constant for the reaction under these conditions. T = 800 °C CO(g) + H O(g) ⇌ CO (g) + H (g) Kc = 0.640 2 2 2 It is important to recognize that an equilibrium can be established starting either from reactants or from products, or from a mixture of both. For example, equilibrium was established from Mixture 2 in Figure 13.7 when the products of the reaction were heated in a closed container. In fact, one technique used to determine whether a reaction is truly at equilibrium is to approach equilibrium starting with reactants in one experiment and starting with products in another. If the same value of the reaction quotient is observed when the concentrations stop changing in both experiments, then we may be certain that the system has reached equilibrium.

PDF page 697: The magnitude of an equilibrium constant is a measure of the yield of a reaction when it reaches equilibrium. A large value for Kc indicates that equilibrium is attained only after the reactants have been largely converted into products. A small value of Kc—much less than 1—indicates that equilibrium is attained when only a small proportion of the reactants have been converted into products. Once a value of Kc is known for a reaction, it can be used to predict directional shifts when compared to the value of Qc. A system that is not at equilibrium will proceed in the direction that establishes equilibrium. The data in Figure 13.7 illustrate this. When heated to a consistent temperature, 800 °C, different starting mixtures of CO, H2O, CO2, and H2 react to reach compositions adhering to the same equilibrium (the value of Qc changes until it equals the value of Kc). This value is 0.640, the equilibrium constant for the reaction under these conditions. T = 800 °C CO(g) + H O(g) ⇌ CO (g) + H (g) Kc = 0.640 2 2 2 It is important to recognize that an equilibrium can be established starting either from reactants or from products, or from a mixture of both. For example, equilibrium was established from Mixture 2 in Figure 13.7 when the products of the reaction were heated in a closed container. In fact, one technique used to determine whether a reaction is truly at equilibrium is to approach equilibrium starting with reactants in one experiment and starting with products in another. If the same value of the reaction quotient is observed when the concentrations stop changing in both experiments, then we may be certain that the system has reached equilibrium.

PDF page 695: x y [C] [D] Qc = m n [A] [B] The reaction quotient is equal to the molar concentrations of the products of the chemical equation (multiplied together) over the reactants (also multiplied together), with each concentration raised to the power of the coefficient of that substance in the balanced chemical equation. For example, the reaction quotient for the reversible reaction 2NO (g) ⇌ N O (g) is given by this expression: 2 2 4 [N O ] 2 4 Qc = 2 [NO ] 2

PDF page 701: The subscript P in the symbol KP designates an equilibrium constant derived using partial pressures instead of concentrations. The equilibrium constant, KP, is still a constant, but its numeric value may differ from the equilibrium constant found for the same reaction by using concentrations. Conversion between a value for Kc, an equilibrium constant expressed in terms of concentrations, and a value for KP, an equilibrium constant expressed in terms of pressures, is straightforward (a K or Q without a subscript could be either concentration or pressure). The equation relating Kc and KP is derived as follows. For the gas-phase reaction mA + nB ⇌ xC + yD:

PDF page 699: • Activities for pure condensed phases (solids and liquids) are equal to 1.

PDF page 699: As a consequence of this last consideration, Qc and Kc expressions do not contain terms for solids or liquids (being numerically equal to 1, these terms have no effect on the expression's value). Several examples of equilibria yielding such expressions will be encountered in this section.

PDF page 699: A homogeneous equilibrium is one in which all of the reactants and products are present in a single solution (by definition, a homogeneous mixture). In this chapter, we will concentrate on the two most common types of homogeneous equilibria: those occurring in liquid-phase solutions and those involving exclusively gaseous species. Reactions between solutes in liquid solutions belong to one type of homogeneous equilibria. The chemical species involved can be molecules, ions, or a mixture of both. Several examples are provided here.

PDF page 702: A heterogeneous equilibrium is a system in which reactants and products are found in two or more phases. The phases may be any combination of solid, liquid, or gas phases, and solutions. When dealing with these equilibria, remember that solids and pure liquids do not appear in equilibrium constant expressions (the activities of pure solids, pure liquids, and solvents are 1). Some heterogeneous equilibria involve chemical changes; for example: 2+ 2+ − − 2 PbCl (s) ⇌ Pb (aq) + 2Cl (aq) Kc = [Pb ][Cl ] 2 1 CaO(s) + CO (g) ⇌ CaCO (s) Kc = 2 3 [CO ] 2 [CS ] 2 C(s) + 2S(g) ⇌ CS (g) Kc = 2 2 [S] Other heterogeneous equilibria involve phase changes, for example, the evaporation of liquid bromine, as shown in the following equation: Br (l) ⇌ Br (g) Kc = [Br ] 2 2 2 We can write equations for reaction quotients of heterogeneous equilibria that involve gases, using partial pressures instead of concentrations. Two examples are: 1 CaO(s) + CO (g) ⇌ CaCO (s) K = P 2 3 P CO 2 P CS 2 C(s) + 2S(g) ⇌ CS (g) K = P 2 ⎛ ⎞2 ⎝P ⎠ S

PDF page 701: x+y x y (RT) [C] [D] = × m n m+n [A] [B] (RT) (x+y) − (m+n) = Kc (RT) Δn = Kc (RT) The relationship between Kc and KP is Δn K = Kc (RT) P In this equation, Δn is the difference between the sum of the coefficients of the gaseous products and the sum of the coefficients of the gaseous reactants in the reaction (the change in moles of gas between the reactants and the products). For the gas-phase reaction mA + nB ⇌ xC + yD, we have Δn = (x+y) − (m+n)

PDF page 698: Qc < Kc (0.039 < 0.64) The reaction will shift to the right. Experiment 2:

PDF page 699: Qc > Kc (140 > 0.64) The reaction will shift to the left. Experiment 3:

PDF page 697: = 1.6 × 10 . The equilibrium constant is 1.6 × 10 .

PDF page 697: equilibrium, Kc = Qc =