# For the reaction of H2 with I2, the rate constant is 2.5 × 10-4 dm3 mol-1 s-1 at 327°C and 1.0 dm3 mol-1 s-1 at 527°C. The activation energy for the r

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For the reaction of H2 with I2, the rate constant is 2.5 × 10-4 dm3 mol-1 s-1 at 327°C and 1.0 dm3 mol-1 s-1 at 527°C. The activation energy for the reaction, in KJ mol-1 is:

(R = 8.314 JK-1 mol-1)
1. 166
2. 150
3. 72
4. 59

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Correct Answer - Option 1 : 166

Concept:

Arrhenius equation:

Arrhenius equation gives the dependence of the rate constant of a chemical reaction on the absolute temperature, a pre-exponential factor and other constants of the reaction.

$k={{e}^{\frac{-{{E}_{a}}}{RT}}}$     ----(1)

Where,

k is the rate constant,

T is the absolute temperature (in Kelvin),

Ea is the activation energy for the reaction (in the same units as RT),

R is the universal gas constant.

Calculation:

We can get the activation energy using Arrhenius equation which is given by the formula:

$k={{e}^{\frac{-{{E}_{a}}}{RT}}}$

From the question, the reaction is:

H2 (g) + I2 (g) → 2HI (g)

Now, taking logarithm on both sides on equation (1),

$\Rightarrow \ln k=-\frac{{{E}_{a}}}{RT}$

Now, the s for the given reaction is:

$\log \frac{{{k}_{2}}}{{{k}_{1}}}=\frac{{{E}_{a}}}{2.303R}\left[ \frac{1}{{{T}_{1}}}-\frac{1}{{{T}_{2}}} \right]$

T1 = 327°C + 273 = 600 K

T2 = 527°C + 273 = 800 K

$\Rightarrow \log \left( \frac{1}{2.5\times {{10}^{-4}}} \right)=\frac{{{E}_{a}}}{2.303\times 8.314}\left[ \frac{1}{600}-\frac{1}{800} \right]$

$\Rightarrow \log \left( \frac{1}{2.5\times {{10}^{-4}}} \right)=\frac{{{E}_{a}}}{2.303\times 8.314}\left[ \frac{4-3}{2400} \right]$

$\Rightarrow \log \left( \frac{1}{2.5\times {{10}^{-4}}} \right)=\frac{{{E}_{a}}}{2.303\times 8.314}\left[ \frac{1}{2400} \right]$

$\Rightarrow 3.602=\frac{{{E}_{a}}}{2.303\times 8.314}\left[ \frac{1}{2400} \right]$

⇒ Ea = 2.303 × 8.314 × 2400 × 3.602

Ea = 165.5 KJ.mol-1

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