12. The following mechanism has been proposed for the gas-phase reaction of chloroform (CHCl3) and chlorine:
step 1: Cl2 (g) ⇌ 2 Cl (g). (fast)
step 2: Cl (g) + CHCl3 (g) → HCl (g) + CCl3 (g) (slow)
step 3: Cl (g) + CCl3 (g) → CCl4 (g). (fast)
a) What is the overall reaction?
b) What are the intermediates in the mechanism?
c) What is the rate law predicted by this mechanism?
Despite the potential problems with this mechanism, it is one which appears on multiple sites across the interweb.
step 1: Cl2 (g) ⇌ 2 Cl (g) ...... (fast)
step 2: Cl (g) + CHCl3 (g) → HCl (g) + CCl3 (g) ...... (slow)
step 3: Cl (g) + CCl3 (g) → CCl4 (g) ...... (fast)
1. Overall reaction.
Cl2(g) + CHCl3(g) --> CCl4(g) + HCl(g)
2. Intermediates are species which are formed in early steps of the mechanism and reactants in subsequent steps. They do not appear in the overall reaction. The intermediates are Cl• and •CCl3.
3. The rate law includes an intermediate.
rate = k'[Cl] [CHCl3]
Cl2 <==> 2Cl ............. K = ???
K = [Cl]² / [Cl2]
[Cl] = (K[Cl2])^(1/2)
rate = k'(K[Cl2])^(1/2) [CHCl3]
rate = k [Cl2]^(1/2) [CHCl3] .................. all of the constants are accumulated in k
The reaction shown is incorrect. Step 1 is slow (rate determining) and step 2 is very fast as are all free radical reactions. This is a very fast reaction overall because step 3 really is: *CCl3 + Cl2 --> CCL4 + *Cl. This is a free radical chain reaction with steps 2 and 3 oscillating back and forth to rapidly consume both reactants. This is in contrast to the reaction shown which goes one slow step at a time. Wrong.