Write an expression for the equilibrium constant of each chemical equation.

a. SbCl₅(g) ⇌ SbCl₃(g) + Cl₂(g)

b.2 BrNO(g) ⇌ 2 NO(g) + Br₂(g)

c. CH₄(g) + 2 H₂S(g) ⇌ CS₂(g) + 4 H₂(g)

d. 2 CO(g) + O₂(g) ⇌ 2 CO₂(g)

Find and fix each mistake in the equilibrium constant expressions. a. 2 H₂S(g) ⇌ 2 H₂(g) + S₂(g) K = [H₂][S₂]/[H₂S]

Find and fix each mistake in the equilibrium constant expressions. b. CO(g) + Cl₂(g) ⇌ COCl₂(g) K = [CO][Cl₂]/[COCl₂]

When this reaction comes to equilibrium, will the concentrations of the reactants or products be greater? Does the answer to this question depend on the initial concentrations of the reactants and products? A(g)+B(g) ⇌ 2 C(g) K_c = 1.4⨉10^-5

Ethene (C₂H₄) can be halogenated by this reaction: C₂H₄(g) + X₂(g) ⇌ C₂H₄X₂(g) where X₂ can be Cl₂ (green), Br₂ (brown), or I₂ (purple). Examine the three figures representing equilibrium concentrations in this reaction at the same temperature for the three different hal- ogens. Rank the equilibrium constants for the three reactions from largest to smallest.

H₂ and I₂ are combined in a flask and allowed to react according to the reaction: H₂(g) + I₂(g) ⇌ 2 HI(g) Examine the figures (sequential in time) and answer the questions: a. Which figure represents the point at which equilibrium is reached?

A chemist trying to synthesize a particular compound attempts two different synthesis reactions. The equilibrium constants for the two reactions are 23.3 and 2.2⨉10⁴ at room temperature. However, upon carrying out both reactions for 15 minutes, the chemist finds that the reaction with the smaller equilibrium constant produces more of the desired product. Explain how this might be possible.

This reaction has an equilibrium constant of K_p = 2.26⨉10⁴ at 298 K. CO(g) + 2 H₂(g) ⇌ CH₃OH(g) Calculate K_p for each reaction and predict whether reactants or products will be favored at equilibrium. a. CH₃OH(g) ⇌ CO(g) + 2 H₂(g)

This reaction has an equilibrium constant of K_p = 2.26⨉10⁴ at 298 K. CO(g) + 2 H₂(g) ⇌ CH₃OH(g) Calculate K_p for each reaction and predict whether reactants or products will be favored at equilibrium.

b. 1/2 CO(g) + H₂ (g) ⇌ 1/2 CH₃OH(g)

This reaction has an equilibrium constant of K_p = 2.26⨉10⁴ at 298 K. CO(g) + 2 H₂(g) ⇌ CH₃OH(g) Calculate K_p for each reaction and predict whether reactants or products will be favored at equilibrium.

c. 2 CH₃OH(g) ⇌ 2 CO(g) + 4 H₂(g)

This reaction has an equilibrium constant of K_p = 2.2⨉10⁶ at 298 K. 2 COF₂(g) ⇌ CO₂(g) + CF₄(g) Calculate K_p for each reaction and predict whether reactants or products will be favored at equilibrium.

a. COF₂ (g) ⇌ 1/2 CO₂(g) + 1/2 CF₄(g)

This reaction has an equilibrium constant of K_p = 2.2⨉10⁶ at 298 K. 2 COF₂(g) ⇌ CO₂(g) + CF₄(g) Calculate K_p for each reaction and predict whether reactants or products will be favored at equilibrium.

b. 6 COF₂(g) ⇌ 3 CO₂(g) + 3 CF₄(g)

This reaction has an equilibrium constant of K_p = 2.2⨉10⁶ at 298 K. 2 COF₂(g) ⇌ CO₂(g) + CF₄(g) Calculate K_p for each reaction and predict whether reactants or products will be favored at equilibrium.

c. 2 CO₂(g) + 2 CF₄(g) ⇌ 4 COF₂(g)

Consider the reactions and their respective equilibrium

constants:

NO(g) + 1/2 Br (g) ⇌ NOBr(g) K_p = 5.3

2 NO(g) ⇌ N₂(g) + O₂(g) K_p = 2.1⨉10³⁰

Use these reactions and their equilibrium constants to predict

the equilibrium constant for the following reaction: N₂(g) + O₂(g) + Br₂(g) ⇌ 2 NOBr(g)

Calculate K_c for each reaction. a. I₂(g) ⇌ 2I(g) K_p = 6.26⨉10^-22 (at 298K)

Calculate K_c for each reaction.

b. CH₄(g) + H₂O(g) ⇌ CO(g) + 3 H₂(g) K_p = 7.7x10²⁴ (at 298 K)

c. I₂(g) + Cl₂(g) ⇌ 2 ICl(g) K_p = 81.9 (at 298 K)

Calculate K_p for each reaction. a. N₂O₄(g) ⇌ 2 NO₂(g) K_c = 5.9⨉10^-3 (at 298 K)

Calculate K_p for each reaction.

b. N₂(g) + 3 H₂(g) ⇌ 2 NH₃(g) K_c = 3.7⨉10⁸ (at 298 K)

c. N₂(g) + O₂(g) ⇌ 2 NO(g) K_c = 4.10⨉10^-31 (at 298 K)

Consider the reaction: N₂(g) + 3 H₂(g) ⇌ 2 NH₃(g) Complete the table. Assume that all concentrations are equilibrium concentrations in M.

T (K) [N₂] [H₂] [NH₃] K_c

500 0.115 0.105 0.439 _

575 0.110 _ 0.128 9.6

775 0.120 0.140 _ 0.0584

Consider the following reaction: H₂(g) + I₂(g) ⇌ 2 HI(g) Complete the table. Assume that all concentrations are equilibrium concentrations in M.

T (°C) [H₂] [I₂] [HI] Kc

25 0.0355 0.0388 0.922 _

340 _ 0.0455 0.387 9.6

445 0.0485 0.0468 _ 50.2

Consider the reaction: 2 NO(g) + Br₂(g) ⇌ 2 NOBr(g) K_p = 28.4 at 298K In a reaction mixture at equilibrium, the partial pressure of NO is 108 torr and that of Br₂ is 126 torr. What is the partial pressure of NOBr in this mixture?

Consider the reaction: SO₂Cl₂(g) ⇌ SO₂(g) + Cl₂(g) K_p = 2.91⨉10³ at 298 K In a reaction at equilibrium, the partial pressure of SO₂ is 137 torr and that of Cl₂ is 285 torr. What is the partial pressure of SO₂Cl₂in this mixture?

Consider the reaction: NH₄HS(s) ⇌ NH₃(g) + H₂S(g) At a certain temperature, K_c = 8.5⨉10^-3. A reaction mixture at this temperature containing solid NH₄HS has [NH₃] = 0.166 M and [H₂S] = 0.166 M. Will more of the solid form or will some of the existing solid decompose as equilibrium is reached?

Silver sulfate dissolves in water according to the reaction: Ag₂SO₄(s) ⇌ 2 Ag⁺(aq) + SO₄^2-(aq) K_c = 1.1⨉10^-5 at 298 K. A 1.5-L solution contains 6.55 g of dissolved silver sulfate. If additional solid silver sulfate is added to the solution, will it dissolve?