The document summarizes reaction mechanisms for the reaction of molecular hydrogen with chlorine and iodine. For the hydrogen-chlorine reaction, a chain reaction mechanism is proposed involving initiation, propagation, and termination steps. For the hydrogen-iodine reaction, it was originally thought to be bimolecular but is shown to involve radical intermediates and proceed by a chain mechanism as well, though multiple mechanisms may contribute to the overall reaction.

1. Reaction between molecular hydrogen and chlorine
Photochemical reaction:
The mechanism of this reaction was proposed by Gohring in 1921
which can be represented as:
i) Cl2 +hν → 2Cl Chain initiation
ii) Cl + H2 → HCl +H Chain propagation
iii) H + Cl2→ HCl + Cl Chain propagation
iv) H + O2→ HO2
v) Cl + O2→ ClO2 Chain termination
vi) Cl + X → ClX
Here X is any substance that can remove chlorine atoms. The chain
ending step Cl+Cl→Cl2 is not taken into consideration, because the
chlorine radicals are removed more effectively by reactions v and vi.

2. The concentration of chlorine atom[Cl] in this mechanism is much
lower than that of [Br], since the activation energy of Cl is lower than
Br.
The rate of Cl formation in step (i) is 2I, where I is the intensity of the
absorbed light , because two radicals are formed for each absorbed
photon.
Applying the steady-state treatment to chlorine atoms, we get:
d[Cl] /dt = 2I-k2[Cl][H2] +k3[H][Cl2] – k5[Cl][O2] – k6[Cl][X]=0….1
The steady-state equation for H atoms is:
d[H]/dt =k2 [Cl][H2]- k3[H][Cl2] – k4[H][O2] =0 …………….2
from which it follows that
[Cl]= k3[H][Cl2] + k4[H][O2] / k2[H2] ……………..3
Introduction the above expression into eq. 1, with some
rearrangement and neglecting the small terms involving [O2]2 , gives:
[H] = 2Ik2 [H2]/ k3 k6 [X][Cl2] + [O2] k2k4 [H2] + k3k5[Cl2] + k4k6 [X]………4

3. The rate of the formation of hydrogen chloride is given by:
νHCl =d[HCl]/dt = k2[Cl][H2] + k3[H][Cl2] ………………….5
Substraction eq. 2 gives:
νHCl =2 k3[H][Cl2] + k4[H][O2] …………………..6
At low concentration of oxygen the second term can be neglected, and then,
with the introduction of eq. 4 we get:
νHCl = 2k2k3 [H2][Cl2]I/k3k6[Cl2][X] +[O2] k2k4 [H2]+k3k5[Cl2] + k4k6[X])……7
or νHCl = 2k3 /k4 I [H2][Cl2]/k3k6 / k2k4[Cl2][X] +[O2] [H2]+k3k5/k2k4 [Cl2] +
k6/k2[X])……8
Apart from the final term in the denominator, this expression is of the
form of the empirical equation i.e.
d[HCl]/dt= kI[H2][Cl2]/m [Cl2]+ [Cl2]([H2] +n[Cl2]), the constants being
related by: k= 2k3 /k4 ,m= k3k6 / k2k4[X] and n=k3k5/k2k4 .................9
The scheme of reactions therefore gives a kinetic law that is essentially
in agreement with the experimental rate equation.
For the thermal reaction between H and Cl the rate may be express as:
νHCl = d[HCl]/dt=kI[H2][Cl2]2/mI[Cl2] +[O2]([H2]+nI[Cl2]) …………10

4. Reaction between molecular hydrogen and iodine
The reaction between H2 and I2 was extensively studied by
Bodenstein at the end of the nineteenth century. For many years
it was considered to be a true bimolecular reaction. However, in
1959, Sullivan showed the presence of free radicals in the system
above 600K, leading to the proposed mechanism:
I2→2I
I+H2 → HI+H 1
H+I2→HI+I
i.e. a chain reaction. Once the intermediate HI is formed, it can
be deactivated through the reverse process of step (ii). In
addition, it is also plausible that there is the formation of a linear
transition state {I…..H…..H….I}‡, which relates I and H2I with the
products, HI. The following mechanism includes the formation of
intermediates of this type

5. i) I2↔2I (fast)
ii) I+H2 ↔ H2I (fast)
iii)H +H2I ↔2HI (slow)
where the rate determining step is step (iii) and thus:
ν = k3[I][H2I] ……………..2
Inserting the expressions for [I] and [H2I], obtain from the
equilibrium processes(i) and (ii), we get:
ν = k[I2][H2I] ……………….3
Here, k=k3k2/k1. This mechanism, together with mechanism (1), all
appear to make contributions to the overall reaction. Another
process that is almost certainly involved proceeds via the
elementary step I2+H2→2HI, which proceeds via a transition state
with a trapezoidal configuration.
This illustrates that a single reaction can often occur by more
than one mechanism proceeding simultaneously.