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Chemical Equilibrium (McGuire)

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

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


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


    1. The reverse of the reaction in part a


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


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


    1. 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.


    1. H2O(g) is added.


    1. The pressure is increased by adding helium gas.


    1. The temperature is increased.


    1. 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.




    Add CaO(s)



    Decrease the container volume



    Decrease the container volume

    Total Pressure


    Add CO2



    Add CO2



    Add Helium gas



    Increase the temperature



    Increase the temperature



    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 PSO3 = 2 x 10-3 atm; PSO2 = 5 x 10-3 atm; PO2 = 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.


    1. You could ___________ the volume.


    1. You could ___________ oxygen gas.


    1. 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.


    1. 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. PI2 = 0.461 atm; PH2S = 0.050 atm; PHI = 0.0 atm


    1. PI2 = 0.461 atm; PH2S = 0.050 atm; PHI = 9.0 atm



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