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Chemistry LibreTexts

13.E: Chemical Equilibrium

1.  For a system at equilibrium, which of the following are true?

  1. The rate of the reaction is zero.
  2. The concentrations of reactants and products are no longer changing.
  3. The value for the equilibrium constant, K, will change when temperature is changed.
  4. The rate of the forward reaction is equal to the rate of the reverse reaction.

 

2.  The equilibrium constant for the following reaction is equal to 0.447 at 100°C:

\(\mathrm{\dfrac{1}{2} N_2O_4(g)\rightleftharpoons NO_2(g)}\)

  1. What is the value for K for the following reaction:  2 NO2(g)  ⇌  N2O4(g)?
  2. Write an expression for K, the equilibrium constant, for the reaction in (a).
  3. Find the concentration of N2O4 in a system at equilibrium (at 100°C) if the equilibrium concentration of NO2 is 0.30 mol/L.

 

3.  For each reaction below, write an expression for K and indicate what effect an increase in pressure would have on equilibrium.

  1. \(H_{2(g)}  + S_{(s)}   \rightleftharpoons  H2S_{(g)}\)         
  2. \(N_{2(g)}  +  3 H_{2(g)}  \rightleftharpoons  2 NH_{3(g)}\)
  3. \(H_{2(g)}  +  Br_{2(l)}  \rightleftharpoons  2 HBr_{(g)}\)

 

4.  For the equilibrium in Question 3b, K is 4.51 x 10-5 at 450°C.  Is a mixture containing 100 atm NH3, 30 atm N2 and 500 atm H2 at equilibrium?  If not, will the mixture shift toward product or reactants to achieve equilibrium?

 

5.  Consider the following system at equilibrium:

\[2 N_2O_{(g)} \rightleftharpoons 2 N_{2(g)}  +  O_{2(g)} \;\;\; ΔH = +163\; kJ\]

  1. For each situation below, indicate whether more product or more reactant is produced in order to re-establish equilibrium.
  1. N2 is added                                                
  2. O2 is removed
  3. the volume is increased                            
  4. the temperature is increased
  5. the pressure is increased by compressing the mixture
  1. Which of the situations above will increase yield?
  2. What effect will an increase in temperature have on the value for K?

 

6.  A mixture of 0.100 mol of NO, 0.050 mol of H2, and 0.100 mol of H2O are placed in a 1.00-liter flask. The following equilibrium is established:

\[2 NO_{(g)}  +  2 H_{2(g)}  \rightleftharpoons N_{2(g)}  +  2 H_2O_{(g)}\]

      At equilibrium, [NO] = 0.070 M.

  1. Calculate the equilibrium concentrations of H2, N2, and H2O.
  2. Write an expression for K for this reaction.
  3. Calculate K for this reaction.
  4. At equilibrium, how will the concentrations of products compare to the concentrations of reactants?

 

7.  At 1500 K, the equilibrium constant for the reaction, N2(g)  +  O2(g)   ⇌  2 NO(g), is 1.0 x 10-5. Calculate the equilibrium concentrations of N2, O2 and NO if, before any reaction, 0.500 mol of NO is placed in a 1.00-liter container.  (Ignore significant digits for NO.)

 

8.  Given the equilibrium constants for the following two reactions:

\[A_{(g)}  \rightleftharpoons B_{(g)}  \;\;\;  K =  K_1\]

 \[B_{(g)} \rightleftharpoons  C_{(g)} \;\;\;   K  =  K_2\]

  1. Derive the equilibrium constant for the reaction, \(A_{(g)} \rightleftharpoons C_{(g)}\) in terms of \(K_1\) and \(K_2\).
  2. What general statement can be made about what happens to values for \(K\) when two reactions are added?