The rate constant k for a certain reaction is measured at two different temperatures: temperature k 1.4 x 1010 2.0 × 10⁹ Assuming the rate constant obeys the Arrhenius equation, calculate the activation energy E for this reaction. Round your answer to 2 significant digits. E = 311.0 °C 246.0 °C kJ mol 0x10 X Ś

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The rate constant \( k \) for a certain reaction is measured at two different temperatures:

| Temperature (°C) | \( k \) (s\(^{-1}\))        |
|------------------|-------------------|
| 311.0 °C          | \( 1.4 \times 10^{10} \) |
| 246.0 °C          | \( 2.0 \times 10^9 \)     |

Assuming the rate constant obeys the Arrhenius equation, calculate the activation energy \( E_a \) for this reaction.

Round your answer to 2 significant digits.

**Equation to Calculate Activation Energy:**

\[ E_a = \square \, \text{kJ/mol} \]

(To indicate where to input the answer, there is a box with functionalities such as a check box for scientific notation (\( \times 10 \)), and buttons for submitting or resetting the answer.)
Transcribed Image Text:The rate constant \( k \) for a certain reaction is measured at two different temperatures: | Temperature (°C) | \( k \) (s\(^{-1}\)) | |------------------|-------------------| | 311.0 °C | \( 1.4 \times 10^{10} \) | | 246.0 °C | \( 2.0 \times 10^9 \) | Assuming the rate constant obeys the Arrhenius equation, calculate the activation energy \( E_a \) for this reaction. Round your answer to 2 significant digits. **Equation to Calculate Activation Energy:** \[ E_a = \square \, \text{kJ/mol} \] (To indicate where to input the answer, there is a box with functionalities such as a check box for scientific notation (\( \times 10 \)), and buttons for submitting or resetting the answer.)
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