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Question description

The following reaction is a step in the commercial production of sulfuric acid.

2SO2(g) + O2(g)  2SO3(g)

The equilibrium constant is very high at room temperature, but the reaction is very slow. It must be run at high temperatures to achieve a reasonable rate of reaction. However the equilibrium constant decreases as temperature increases. Calculate Kc at 897 °C from the following concentration data:
[SO3] = 0.0077 M
[SO2] = 0.017 M
[O2] = 0.019 M
(Kc does not have units).

Sometimes in such a gaseous reaction it is more convenient to measure the partial pressures of the reactants and products rather than the molar concentration. In that case Kp values are obtained. Calculate the Kp for the above reaction at 702 °C from the following pressure data:
PSO3 = 3.1 atm
PSO2 = 2.3 atm
PO2= 0.27 atm
(Kp does not have units).

Iodine is sparingly soluble in pure water. However, it does `dissolve' in solutions containing excess iodide ion because of the following reaction:

 I-(aq) + I2(aq) I3-(aq) K = 710

3-> to [I2].

5.00×10-2 mol of I2 is added to 1.00 L of 5.00×10-1 M KI solution.

The solution above is diluted to 12.00 L.

Given the following equilibrium constants at 427°C:

 O2(g) + 4Na(l) > 2Na2O(s) K1 = 2.50×1049

 O2(g) + 2Na(l) > 2NaO(g) K2 = 2.50×109

 O2(g) + 2Na(l)> Na2O2(s) K3 = 2.00×1028

 O2(g) + Na(l) > NaO2(s) K4 = 3.33×1013

What would be the value of the equilibrium constant for each of the following reactions, at 427°C?

1.
 Na2O2(s) + Na(l) > NaO(g) + Na2O(s)
2.
 2NaO2(s) + 2Na(l) > 2NaO(g) + Na2O2(s)
At a particular temperature, Kp = 0.260 for the reaction
 N2O4(g)"> 2NO2(g)

A flask containing only N2O4(g) at an initial pressure of 3.60 atm is allowed to reach equilibrium. Calculate the total pressure in this flask at equilibrium.

4.052 atm

With no change in the amount of material in the flask, the volume of the container in question is increased to 5.000 times the original. Assuming constant temperature, calculate the (new) total pressure, at equilibrium.

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