For two measured rate constants for the same reaction at different temperatures: Temperature / °C | k / s⁻¹ - 199 | 4.50 × 10⁻⁵ - 258 | 3.2 × 10⁻³ What is the activation energy (in kJ mol⁻¹) for this reaction?

Use the Arrhenius equation in its logarithmic form: \( \ln(k_2/k_1) = -\frac{E_a}{R} \left( \frac{1}{T_2} - \frac{1}{T_1} \right) \), where \( k_1 \) and \( k_2 \) are the rate constants at temperatures \( T_1 \) and \( T_2 \) respectively, \( E_a \) is the activation energy, and \( R \) is the universal gas constant (8.314 J/mol·K).