Iron(II) ion is oxidized by chlorine in aqueous solution, the overall reaction b
ID: 491423 • Letter: I
Question
Iron(II) ion is oxidized by chlorine in aqueous solution, the overall reaction being 2 Fe^2+ + Cl_2 rightarrow 2 Fe^3+ 2Cl^- It is found experimentally that the rate of the overall reaction is decreased when either the iron(III) ion or the chloride-ion concentration is increased. Which of the following possible mechanisms is consistent with the experimental observations? Prove by finding rate law for each possible mechanism. a) Step 1 Fe^2+ Cl_2 longdoblearrow Fe^3+ Cl^- + Cl (rapid equilibrium) Step 2 Fe^2+ + Cl rightarrow Fe^3+ + Cl^- b) Step 1 Fe^2+ + Cl_2 longdoublearrow Fe^4+ 2Cl^- (rapid equilibrium) Step 2 Fe^4+ Fe^2+ rightarrow 2 Fe^3+
Explanation / Answer
mechanism 1
step 1
Keq = [Fe+3] [Cl-][Cl]/[Fe+2][Cl2]
Thus [Cl] = Keq [Fe+2][Cl2] /{Fe+3][Cl-]
Rate of reaction is wriiten only based on slow step/rate determining step
Thus rate =k2 [Fe+2] [Cl]
= k2.keq. [Fe+2] [Fe+2][Cl2] /{Fe+3][Cl-]
= k [Fe+2]2[Cl2]/[Fe+3][Cl-]
Thus from this mechanism both the products Fe+3 and Cl- are in the denominator of have negative order.
Thus increasing thier concnentrations(either Fe+3 or Cl- decreases the overall reaction rate.
thus this mechanism is consistent with the experimental results.
Mechanim 2
In the similar manner we can write the rate expression
as
Rate = k [Fe+2]2[Cl2] / [Cl-]2
which is not consistent with experimental result, as this does not contain [Fe+3] in rate expression
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