The bromination of acetone that occurs in acid solution is represented by this equation.
CH₃COCH_3(aq) + Br_2(aq)  $$ \to $$
     CH₃COCH₂Br(aq) + H⁺_(aq) + Br^$$-$$_(aq)
These kinetic data were obtained for given reaction concentrations.

Initial concentrations, M
[CH3COCH3		[Br2]		[H+]
0.30		0.05		0.05
0.30		0.10		0.05
0.30		0.10		0.10
0.40		0.05		0.20

Initial rate, disappearance of Br2, Ms$$-$$1
			5.7$$ \times $$10$$-$$5
			5.7$$ \times $$10$$-$$5
			1.2$$ \times $$10$$-$$4
			3.1$$ \times $$10$$-$$4

Based on these data, the rate equation is

The rate constants k₁ and k₂ for two different reactions are 10¹⁶ $$ \cdot $$ e^$$-$$2000/T and 10¹⁵ $$ \cdot $$ e^$$-$$1000/T, respectively.
The temperature at which k₁ = k₂ is

The reaction of hydrogen and iodine monochloride is given as :
H_2(g) + 2ICl_(g) $$ \to $$ 2HCl_(g) + I_2(g)
This reaction is of first order with respect to H_2(g) and ICl_(g),
following mechanisms were proposed.

Mechanism A :
     H_2(g) + 2ICl_(g) $$ \to $$ 2HCl_(g) + I_2(g)
Mechanism B :
     H_2(g) + ICl_(g) $$ \to $$ HCl_(g) + HI_(g) ; slow
     HI_(g) + ICl_(g) $$ \to $$ HCl_(g) + I_2(g) ; fast

Which of the above mechanism(s) can be consistent with the given information about the reaction?

In a first-order reaction A $$ \to $$ B, if k is rate constant and initial concentration of the reactant A is 0.5 M, then the half-life is