We’re being asked to **determine the rate constant** at 25°C of the reaction when k = 8.54 x 10^{-4} s ^{-1} at 45°C with a EA = 90.8 kJ

This means we need to use the ** two-point form of the Arrhenius Equation**:

$\overline{){\mathbf{ln}}\left(\frac{{\mathbf{k}}_{\mathbf{2}}}{{\mathbf{k}}_{\mathbf{1}}}\right){\mathbf{}}{\mathbf{=}}{\mathbf{}}{\mathbf{-}}\frac{\mathbf{Ea}}{\mathbf{R}}\left[\frac{\mathbf{1}}{{\mathbf{T}}_{\mathbf{2}}}\mathbf{-}\frac{\mathbf{1}}{{\mathbf{T}}_{\mathbf{1}}}\right]}$

where:

**k _{1}** = rate constant at T

**k _{2}** = rate constant at T

**E _{a}** = activation energy (in J/mol)

**R** = gas constant (8.314 J/mol•K)

**T _{1} and T_{2}** = temperature (in K).

Increasing the reaction temperature from 25°C to 30°C raises the reaction rate 7.5- fold. What is the rate constant for the at 27.5 °C? Activation energy is 3.03x10^{5} J/mol

A. 6.2 x 10^{-4}

B. 1.3 x 10^{2}

C. 4.6 x 10^{-7}

D. 11.3

E. none of these

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