Solubility of Gases Use Henry\'s Law (Ss = k_hP_g) to determine the pressure of

Question

Solubility of Gases

Use Henry's Law (Ss = k_hP_g) to determine the pressure of gas above the liquid required to produce a concentration of 0.010 mol/kg of each gas indicated in the table. Based on your calculations above, how does the solubility of a gas vary with the size of the Henry's Law constant? Choose one of the gases in Table 13.2 and show how increasing the pressure of the gas above the liquid results in a greater gas solubility. The dissolution of any gas in water can be assumed to be an exothermic process. A representative example is the dissolution of C02 in water, depicted below. CO_2(g) CO_2(aq) delta^H^degree = -19.4 kJ/mol define exothermic What is the effect of increasing the temperature of this equilibrium? Briefly explain in terms of LeChatlier's principle. It is sometimes necessary to remove dissolved gases from water. Would heating/boiling the liquid work? Explain.

Explanation / Answer

! st part  

a) Pg = Sg/Kh

Name og gas        Pg

He                  0.010/(3.8^10^-5) = 2.63 2

H2                 0.010/(7.8*10^-4)    =12.820

O2               0.010/(1.3*10^-3)       = 7.6923

CH4            0.010/(1.4*10^-3)       = 7.1429

CO2             0.010/(1.9*!0^-3)       =5.263

b) Solubility increases with increase in size of Kh

c) Let us take He gas

Here Pg = 2.63

solubility = 0.010

If we take Pg = 5 , solubility = Kh * Pg = 1.8*10^-4 *5 = 0*10^-4   = 0.0009

So it decreases with increase in pressure of gas

2 part

a)

An exothermic reaction is a chemical reaction that releases energy by light or heat

b) Effect of change in temperature

The effect of changing the temperature in the equilibrium can be made clear by a) incorporating heat as either a reactant or a product, and b) assuming that an increase in temperature increases the heat content of a system. When the reaction is exothermic (H is negative, puts energy out), heat is included as a product, and, when the reaction is endothermic (H is positive, takes energy in), heat is included as a reactant. Hence, whether increasing or decreasing the temperature would favor the forward or the reverse reaction can be determined by applying the same principle as with concentration changes.

Take, for example, the reversible reaction of nitrogen gas with hydrogen gas to form ammonia:

N2(g) + 3 H2(g) 2 NH3(g)    H = -92 kJ mol1

Because this reaction is exothermic, it produces heat:

N2(g) + 3 H2(g) 2 NH3(g) + heat

If the temperature was increased, the heat content of the system would increase, so the system would consume some of that heat by shifting the equilibrium to the left, thereby producing less ammonia. More ammonia would be produced if the reaction was run at a lower temperature, but a lower temperature also lowers the rate of the process, so, in practice (the Haber process) the temperature is set at a compromise value that allows ammonia to be made at a reasonable rate with an equilibrium concentration that is not too unfavorable.

In above case as temperature is increased , backward reaction will be favored by le chatelier principle that is more CO2(g) will be formed

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