Chemical Equilibrium (McGuire)

1.) Write the equilibrium constant expression for each of the following reactions.

1. \(\ce{2 H2O2 (g) ⇌ 2H2O (g) + O2 (g)}\)

 

2. \(\ce{6 H2O2 (g) ⇌ 6 H2O (g) + 3 O2 (g)}\)

 

3. The reverse of the reaction in part a

 

4. \(\ce{2PbS (s) + 3O2 (g) ⇌ 2PbO (s) + 2SO2 (g)}\)

 

5. \(\ce{MgCl2(s) ⇌ Mg^2+ (aq) + 2 Cl- (aq)}\)

 

6. The reverse of the reaction in part e

 

2.) Calculate the equilibrium constant for this reaction: 

\[\ce{2 PO2Br (aq)  ⇆  2 PO2 (aq)  +  Br2 (aq)}\nonumber\]

Given:  [PO₂Br] = 0.0255M,   [PO₂] = 0.155M, and [Br₂]  = 0.00351M at equilibrium.

 

 

3.)  A solution is prepared having the following initial concentrations:

\[\ce{[Fe^3+ ]} = \ce{[Hg2^2+ ]} = \textrm{0.5000 M;} \hspace{30px} \ce{[Fe^2+ ]} = \ce{[Hg^2+ ]} = \textrm{0.03000 M}\nonumber\]

The following reaction occurs among the ions at a certain temperature.

\[\ce{2Fe^3+(aq) + Hg2^2+(aq) ⇔ 2Fe^2+(aq) + 2Hg^2+(aq)} \hspace{30px} \mathrm{K = 9.14 \times 10^{-8}}\nonumber\]

Will the concentration of each the ions be higher or lower when equilibrium is established?

 

4.) An important exothermic reaction in the commercial production of hydrogen is

\[\ce{CO(g) + H2O(g) ⇔ H2(g) CO2(g)}\nonumber\]

How will the system at equilibrium shift in each of the following cases?

1. CO₂ is removed.

 

2. H₂O(g) is added.

 

3. The pressure is increased by adding helium gas.

 

4. The temperature is increased.

 

5. The pressure is increased by decreasing volume.

 

 

 

5.) Consider the endothermic reaction

\[\ce{CaCO3(s) ⇔ CaO(s) + CO2(g)}\nonumber\]

Fill in the table below showing how the indicated change to the equilibrium system will affect the indicated quantity when a new equilibrium state is established.

Change	Quantity	Effect
Add CaO(s)	CO2
Decrease the container volume	CaO(s)
Decrease the container volume	Total Pressure
Add CO2	K
Add CO2	CaCO3
Add Helium gas	CO2
Increase the temperature	CO2
Increase the temperature	K

6.) At 1000 K, the value of K for the reaction 2SO₃(g) = 2SO₂(g) + O₂(g) is 0.338.  Calculate the value for Q, and predict the direction in which the reaction will proceed toward equilibrium if the initial pressures of reactants are P_SO3 = 2 x 10^-3 atm; P_SO2 = 5 x 10^-3 atm; P_O2 = 3 x 10^-2 atm.

 

7.) Consider this endothermic reaction:  3 O₂(g)  ⇆  2 O₃(g).  To shift this reaction towards the reactants:

1. You could ___________ the pressure.

 

2. You could ___________ the volume.

 

3. You could ___________ oxygen gas.

 

4. You could ___________ the temperature.

 

 

8.) Which of the following, if increased, will change the value of the equilibrium constant in the example above?

1. Pressure       b. Volume       c. [Product]       d. Temperature       e. [Reactant]

 

9.) For the following equilibrium

\[\ce{CH3COOH ⇔ CH3COO- + H+}\nonumber\]

If K=1.8x10^-5, at what [H⁺] concentration will there be equal  concentrations of  CH₃COOH and CH₃COO^-?

 

10.)  At 298 K, F₃SSF (g) decomposes partially to SF₂ (g). At equilibrium, the partial pressure of SF₂ (g) is 1.1 × 10^-4 atm and the partial pressure of F₃SSF is 0.0484 atm.

1. Write a balanced equilibrium equation to represent this reaction.

 

2. Compute the equilibrium constant corresponding to the equation you wrote.

 

 

11.) The equilibrium constant for the reaction

\[\ce{H2S} (\textit{g}) + \ce{I2} (\textit{g}) ⇋ \ce{2 HI} (\textit{g}) + \ce{S} (\textit{s})\nonumber\]

at 110°C is equal to 0.0023. Calculate the reaction quotient Q for each of the following conditions and determine whether solid sulfur is consumed or produced as the reaction comes to equilibrium.

1. P_I2 = 0.461 atm; P_H2S = 0.050 atm; P_HI = 0.0 atm

 

2. P_I2 = 0.461 atm; P_H2S = 0.050 atm; P_HI = 9.0 atm

 





 

Contributors and Attributions

• Molly McGuire, Bucknell University ([e-mail])
• Sourced from the Analytical Sciences Digital Library