1) Consider the following reaction at 1600°C:
Br2(g) ⇔ 2Br(g)
When 1.05 moles of Br2 are put in a 0.980-L flask, 1.2 percent of the Br2 undergoes dissociation. Calculate the equilibrium constant for Kc for the reaction.
2) Consider the heterogeneous equilibrium process:
C (s) + CO2 (g) <---> 2CO (g)
At 700o C, the total pressure of the system is found to be 4.50 atm. If the equilibrium constant KP is 1.52, calculate the equilibrium partial pressures of CO2 and CO.
3) The annual production of zinc sulfide (ZnS) is 4.0 x 10^4 tons. Estimate the number of tons of SO2 produced by roasting it to extract zinc metal.
4) On a smoggy day in a certain city the ozone concentration was 0.42 ppm by volume. Calculate the partial pressure of ozone (in atm) and the number of ozone molecules per liter of air if the temperature and pressure were 20.2 degrees Celsius and 748 mmHg, respectively.
5) The element oxygen was prepared by Joseph Priestley in 1774 by heating mercury(II) oxide.
HgO(s) Hg(l) + ½O2(g) Ho = 90.84 kJ
Use the data given below to estimate the temperature at which this reaction will become spontaneous under standard state conditions.
So(Hg) = 76.02 J/K mol
So(O2) = 205.0 J/K mol
So(HgO) = 70.29 J/K mol
6) Determine the equilibrium constant (KP) for the following reaction at 25oC:
CO(g) + H2O(g) CO2(g) + H2(g) Go = -28.5 kJ
7) Calculate Kp for the following reaction at 25° Celsius:
H2 (g) + I2 (g) -- 2HI (g) ΔG° = 2.60 kJ/mol
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This solution explains:
1) How to calculate an equilibrium constant based on percent dissociation.
2) How to calculate equilibrium partial pressures.
3) How to use enthalpy and entropy to find the temperature for a reaction to occur spontaneously.
4) How to use Gibbs free energy to calculate an equilibrium constant.