For a reaction with an activation energy of 50 kJ/mol, what is the effect of a 1

Practice Questions

For a reaction with an activation energy of 50 kJ/mol, what is the effect of a 10 kJ/mol increase in activation energy on the rate constant at a constant temperature?

Rate constant increases
Rate constant decreases
Rate constant remains the same
Rate constant becomes zero

Questions & Step-by-Step Solutions

For a reaction with an activation energy of 50 kJ/mol, what is the effect of a 10 kJ/mol increase in activation energy on the rate constant at a constant temperature?

Step 1: Understand what activation energy (Ea) is. It is the minimum energy required for a reaction to occur.
Step 2: Know that the rate constant (k) is a number that tells us how fast a reaction happens.
Step 3: Familiarize yourself with the Arrhenius equation: k = A * e^(-Ea/RT). Here, A is a constant, e is the base of natural logarithms, Ea is the activation energy, R is the gas constant, and T is the temperature in Kelvin.
Step 4: Recognize that if the activation energy (Ea) increases, the exponent (-Ea/RT) becomes more negative.
Step 5: Understand that a more negative exponent means that e^(-Ea/RT) becomes smaller.
Step 6: Conclude that if the value of k decreases, the reaction rate slows down when activation energy increases.

Arrhenius Equation – The Arrhenius equation relates the rate constant (k) of a reaction to its activation energy (Ea), temperature (T), and a pre-exponential factor (A).
Effect of Activation Energy on Rate Constant – An increase in activation energy leads to a decrease in the rate constant, indicating that the reaction becomes slower.

For a reaction with an activation energy of 50 kJ/mol, what is the effect of a 1

Practice Questions

Questions & Step-by-Step Solutions

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