Thursday, September 15, 2011

Exercise (6.2 Le Chatelier's Principle)

1. Heating solid sodium bicarbonate in a closed vessel establishes the following equilibrium:
        2 NaHCO3 (s) Na2CO3 (s) + H2O (g) + CO2 (g)
    What would happen to the equilibrium position if
    (a) some of the CO2 were removed from the system.
    (b) some solid Na2CO3 were added to the system.
    (c) some of the solid NaHCO3 were removed from the system.
The temperature remains constant.

2. What effect does an increase in pressure have on each of the following systems at equilibrium? The temperature is kept constant and, in each case, the reactants are in a cylinder fitted with a movable piston.
    (a) A (s) ↔ 2 B (s)
    (b) 2 A (l) B (l)
    (c) A (s) B (g)
    (d) A (g ) B (g)
    (e) A (g) 2 B (g)

3. How would you change the volume of each of the following reaction to increase the yield of the products?
    (a) CaCO3 (s) CaO (s) + CO2 (g)
    (b) S (s) + 3 F2 (g) ↔ SF6 (g)
    (c) Cl2 (g) + I2 (g) 2 ICl (g)

4. How would you change the pressure (via a volume change) the following reaction to decrease the yield of the products?
    (a) 2 SO2 (g) + O2 (g) 2 SO3 (g)
    (b) 4 NH3 (g) + 3 O2 (g) 4 NO (g) + 6 H2O (g)
    (c) CaC2O4 (g)  CaCO3 (s) + CO (g)

5. How does an increase in temperature affect the equilibrium concentration of the underlined substance and the value of K:
    (a) CaO (s) + H2O (l) Ca(OH)2 (g)        ∆H = -82 kJ
    (b) CaCO3 (g) CaO (s) + CO2 (g)        ∆H = +178 kJ
    (c) C(s) + 2 H2 (g) CH4 (g)        ∆H = -75 kJ
    (d) N2 (g) + O2 (g) 2 NO (g)        ∆H = +181 kJ
    (e) P4 (s) + 10 Cl2 (g) 4 PCl5 (g)        ∆H = -1528 KJ

6. Predict the effect of decreasing the temperature on the amounts of products in the following reactions:
    (a) C2H2 (g) + H2O (g) CH3CHO (g)       ∆H = -151 kJ
    (b) CH3CH2OH (l) + O2 (g) CH3CO2H (l) + H2O (g)        ∆H = -451 kJ
    (c) 2 C2H4 (g) + O2 (g) 2 CH3CHO (g)        (exothermic)
    (d) N2O4 (g) 2 NO2 (g)        (endothermic)

7. Predict the effect of increasing the temperature on the amounts of products in the following reactions:
    (a) CO (g) + 2 H2 (g) CH3OH (g)        ∆H = -90.7 kJ
    (b) C (s) + H2O (g) CO (g) + H2 (g)        ∆H = +131 kJ
    (c) 2 NO2 (g) 2 NO (g) + O2 (g)        (endothermic)
    (d) 2 C (s) + O2 (g) 2 CO (g)        (exothermic)

8. Consider the following equilibrium process between dinitrogen tetrafluoride (N2F4) and nitrogen difluoride (NF2):
        N2F4 (g) 2 NF2 (g)        ∆H = 38.5 kJ
    Predict the changes in equilibrium if
    (a) the reaction mixture is heated.
    (b) NF2 gas is removed.
    (c) the pressure is decreased.
    (d) inert gas, such as He, is added at constant volume.

9. Consider this equilibrium system:
        CO (g) + Fe3O4 (s) CO2 (g) + 3 FeO (s)
    How does the equilibrium position shift if
    (a) CO is added.
    (b) solid NaOH is added.
    (c) Fe3O4 is added.
    (d) Dry ice is added at constant temperature.

10. Consider the reaction
         2 SO2 (g) + O2 (g) 2 SO3 (g)       ∆H = -198.2 kJ
     Comment on the changes in the concentration of SO2, O2 and SO3 at equilibrium if
     (a) temperature is increased.
     (b) pressure is increased.
     (c) SO2 is added.
     (d) a catalyst is added.
     (e) helium gas is added at a constant volume.

11. Consider the following equilibrium process:
         PCl5 (g) PCl3 (g) + Cl2 (g)       ∆H = +92.5 kJ
     Predict the direction of the shift in equilibrium position when
     (a) the temperature is raised.
     (b) more chlorine gas is added to the reaction mixture.
     (c) some PCl3 is removed from the mixture.
     (d) the volume of the system is decreased.
     (e) a catalyst is added to the reaction mixture.

12. Copper (II) sulphate dissolves in water to yield the blue hexaaquocopper (II) ion. On the addition of concentrated hydrochloric acid, the solution turns greenish due to the formation of a tetrachlorocuprate(II) ion, which is yellow.
         Cu(H2O)62+ (aq) + 4 Cl- (aq) CuCl42- (aq) + 6 H2O (l)
                  (blue)                                (yellow)
In accordance with Le Chatelier’s principle:
     (a) Give the ways by which the equilibrium can be shifted to the left by altering the concentration of the reactants or products.
     (b) When the greenish solution is heated, the solution turns blue. Is the forward reaction exothermic or endothermic?
     (c) Predict what would be seen when a solution of potassium chloride KCl is added to the equilibrium mixture.
     (d) Predict what will be the effect on equilibrium after adding anhydrous calcium chloride, which is a hygroscopic substance (hygroscopic means that it absorbs water and removes it from the reaction medium).

13. The oxidation of SO2 to SO3 is an important industrial reactions because it is the key in sulfuric acid production:
         SO2 (g) + ½ O2 (g) SO3 (g)        ∆H = -99.2 kJ
     (a) What is the qualitative combination of T and P to maximise SO3 yield?
     (b) How does addition of O2 affect Q? K?
     (c) Suggest a reason why catalysis is used in the manufacture of H2SO4?

1. (a)
    (b) no change
    (c) no change
2. (a) no change
    (b) no change
    (d) no change
3. (a)
    (c) no change
4. (a)
5. (a) , K
    (b) , K
    (c) , K
    (d) , K
    (e) , K
6. (a)
7. (a)
8. (a)
    (d) no change
9. (a)
    (c) no change 
10. (a) [SO2] & [O2], [SO3]
     (b) [SO2] & [O2], [SO3]
     (c) [SO2] & [O2], [SO3]
     (d) no change
     (e) no change
11. (a)
     (e) no change
12. (a) Diluting with water @ Add AgNO3 to remove Cl- ion
     (b) Exothermic
     (c) Solution turns to yellow
     (d) Solution turns to yellow
13. (a) P, T
     (b) Q , K no change
     (c) speed up the reaction

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