Question

A certain first-order reaction has a rate constant of 1.65 min-1 at 20°C. What is the value of the rate constant at 60°C if
the activation energy is 75.5 kJ/mol?

The rate constant is 105 times bigger at 300 K than at 200K. Explain why reaction rate increases with temperature?

At 227°C, the reaction: SO2Cl2(g) SO2(g) + Cl2(g), has an equilibrium constant of KC = 2.99×10-7. If a reaction mixture initially contains 0.175 M SO2Cl2, what is the equilibrium concentration of Cl2 at 227°C ?

Answer 1

K_1 = 1.65 min^- 1, T_1 = 20 degree c (243 k) k_2 = 2, T_2 = 60 degree c(333k), E_a = 73.5 kJ /mol from Arrhenius equation log k_2 k_1 = E_a /2.303 R[1 /T_1 - 1 /T_2] = 75.5 times 10^13 /2.303 times 8.314 [1 /293 - 1 /333] log k_2 /k_1 = 75.5 times 10^3 times 40/2.303 times 8.314 times 293 times 333 log k_2/1.65 = 1.6165 k_ 2 /1.65 = 41.3048 k_2 = 41.3048 times 65 = 68.2 min^- 1 From Arrhenius equation, log k_2 /k_1 = E_a/2.303 [1 /T_1 - 1 /T_2] From above relation we can says that on increasing Temperature, rate constant increases SO_2 cl_2 double head arrow SO_2 (g) + cl_2 initially 0.175 m 0 0 at equilibrium (0.175 - x) x x k_c = x^2/0.175 - x = 2.99 times 10^-