1. For each of the following reactions, determine whether the value of the equilibrium constant favors the formation of reactants, products, or both sides equally.

a) Br₂(g) + Cl₂(g)  ↔ 2 BrCl(g)              K_eq = 3 x 10⁶

b) H₂(g) + Br₂(g) ↔  2 HBr(g)            K_eq = 1.15

c) I₂(g) ↔  I(g) + I(g)                          K_eq = 4.5 x 10^-7

 

2. Molecular chlorine decomposes into atoms according to the reaction:

 

Cl₂ (g)↔   2 Cl (g)

 

The equilibrium constant for the reaction at 25°C is 1.4 x 10^-38.  Would many chlorine atoms be present at this temperature?  Explain how you can determine this.

3. For the following reaction at equilibrium at 2000°C, the concentration of N2 and O2 are both 4.8 M.

 

N₂(g) + O₂(g) ↔  2 NO(g)       K_eq = 6.3 x 10^-4

 

Calculate the concentration of NO at equilibrium. Show your work; pay careful attention to exponents.

4. CaCO₃ (chalk) can produce solid CaO and CO₂ gas when heated. If 6.0 moles of carbon dioxide forms in a 4.500 L reaction vessel, what is the equilibrium constant for this reaction?

5.  Le Chatelier’s Principle states that if a stress is applied to a reversible reaction at equilibrium, the reaction will undergo a shift in order to re-establish its equilibrium. Consider the following exothermic reversible reaction at equilibrium:
2 A ↔B + C
In which direction (left or right) would the following stresses cause the system to shift?
a. decrease the concentration of A
b. increase the concentration of B
c. lower the temperature

6. Chemical engineers use Le Chatêlier’s principle to predict shifts in chemical system at equilibrium resulting from changes in reaction conditions. Predict the changes needed to maximize the yield of product in each of the following industrial chemical systems:

1. a) the production of ethene (ethylene)

C₂H₆(g) + energy ↔C₂H₄(g) + H₂(g)

1. b) the production of methanol

CO(g) + 2H₂(g) ↔ CH₃OH(g) + energy

 

7. Acetic acid, HC₂H₃O₂, is in equilibrium with its ions:

HC₂H₃O₂(aq) ↔ H+(aq) + C₂H₃O₂^–(aq)     Keq = 1.8 x 10^-5

At equilibrium, the concentration of the ions are:                  [H+] = 1.33 x 10^-3 M

[C₂H₃O₂^–] = 1.33 x10^-3 M

Calculate the concentration of the acid, HC₂H₃O₂.

 

8.Sulfuric acid is the most common commercial acid, with millions of tons produced each year. The second step in the production involves the oxidation of sulfur dioxide gas catalyzed with V₂O₅(s),

2SO₂(g) + O₂(g) ↔2 SO₃(g)      DH° = -198 kJ

a) Set-up the equilibrium law for this reaction.

b) List three stresses that could shift the equilibrium to the products side of the reaction.

c) If the equilibrium constant, K_eq, for the above reaction at a given temperature is 4 x 10³, what is the constant (K_eq) for the reverse reaction (the decomposition of sulfur trioxide)?