Discussion Forum

 The activation energy of a reaction is zero, Its rate constant at 280 K is 1.6 x 10^-6 s^-1 , the rate constant at 300K is 

Answer:

Explanation:

Key Idea: Arrhenius  equation is given as:

 $\log\frac{k_{2}}{k_{1}}=\frac{E_{a}}{2.303 R}\left[\frac{1}{T_{1}}-\frac{1}{T_{2}}\right]$

 Given,

 Activation  energy  of a reaction , E_A =0

 Rate constant  , k₁ =1.6 x 10^-6 s^-1

 Temperature  , T₁ =280 K, T₂ =300K

 According to Arrhenius  equation

$\log\frac{k_{2}}{1.6\times 10^{-6}}=\frac{0}{2.303 R}\left[\frac{1}{280}-\frac{1}{300}\right]$

$\log\frac{k_{2}}{1.6\times 10^{-6}}=0$

$\frac{k_{2}}{1.6\times 10^{-6}}=antilog 0$

 $\frac{k_{2}}{1.6\times 10^{-6}}=1$

$\therefore$     $k_{2}=1.6\times 10^{-6}s^{-1}$