For a reaction with an activation energy of 50 kJ/mol, what is the effect of inc

For a reaction with an activation energy of 50 kJ/mol, what is the effect of increasing the temperature from 300 K to 350 K on the rate constant?

For a reaction with an activation energy of 50 kJ/mol, what is the effect of increasing the temperature from 300 K to 350 K on the rate constant?

Step 1: Understand that the activation energy (Ea) is the energy needed for a reaction to occur. In this case, Ea is 50 kJ/mol.
Step 2: Know that temperature (T) is measured in Kelvin (K). We are increasing the temperature from 300 K to 350 K.
Step 3: Familiarize yourself with the Arrhenius equation, which is: k = A * e^(-Ea/(RT)), where k is the rate constant, A is the pre-exponential factor, R is the gas constant (8.314 J/(mol·K)), and T is the temperature in Kelvin.
Step 4: Recognize that as temperature (T) increases, the value of e^(-Ea/(RT)) increases because the exponent becomes less negative.
Step 5: Conclude that since the rate constant (k) is directly related to the exponential term, increasing the temperature from 300 K to 350 K will increase the rate constant.

Arrhenius Equation – The Arrhenius equation describes how the rate constant of a reaction depends on temperature and activation energy.
Activation Energy – Activation energy is the minimum energy required for a reaction to occur, influencing the rate constant.
Temperature Dependence – The rate constant increases with temperature due to the exponential factor in the Arrhenius equation.