1. The dimerization of nitrogen dioxide occurs with the following equilibrium: 2

Question

1. The dimerization of nitrogen dioxide occurs with the following equilibrium: 2 NO2(g) <----> N2O4(g). At room temperature, the Kc for this reaction is 170. If a 2 L container holding a mixture of these substances composed of 0.040 mol NO2 and 0.030 mol N2O4, is the system at equilibrium? If not, in what direction will it proceed to achieve equilibrium?

2. For the following 65C system at equilibrium in a 1 L container, predict the qualitative change each of the following will have:

CH4(g) + 3 Cl2(g) <--------> CHCl3(g) + 3 HCl(g) Hrxn = -305.2 kJ/mol

a) raising T to 78C

b) compressing V to 0.5 L

c) lowering T to 55C

d) forcing 0.2 mol of HCl(g) into the container

3. A buffer system contains 0.55 M acetic acid and 0.47 M sodium acetate. What is its pH? The Ka for acetic acid is 1.8x10-5

Explanation / Answer

2 NO2(g) <----> N2O4(g).
2 moles of NO2 produces 1 mole of N2O4
0.04 mole sof NO2 produces 0.04/2 = 0.02 moles of N2O4
but given is 0.03 moles of N2O4 is present hence the reaction proceeds in backward direction to reach equilibrium.
pH of buffer = pKa + log(salt)/(acid)
pH = -log(1.8*10^-5) + log(0.47/0.55)
pH = 4.6764
CH4(g) + 3 Cl2(g) <--------> CHCl3(g) + 3 HCl(g) ; dH rxn = -305.2 kJ/mol
since dH rxn is negative so the reaction is exothermic reaction and if you increase the temperature then the equilibrium shift in backward direction due to large amount of heat energy the product will be dissociated and if we decrease the temperature then it will favours the forward reaction to form the product.
If we decrease the volume of the container then temperature will decrease at constant conditions hence forward reaction favoured and equilibrium shifts to right side.
If we add o.2 mols of HCl means we are increasing concentration of the product so the reaction moves to backward direction to get equilibrium.

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