Rate Constant Calculator
Find a reaction rate constant with Arrhenius.
Results update as you type.
How it works
The Arrhenius equation links a reaction's rate constant to temperature and activation energy, and this calculator applies it. You enter the pre-exponential factor (also called the frequency factor), the activation energy in kilojoules per mole, and the absolute temperature in kelvin, and it returns the rate constant and the value of the exponent. The equation multiplies the pre-exponential factor by e raised to the negative activation energy divided by the gas constant times temperature. Physically, the activation energy is the barrier that reacting molecules must overcome, and the exponential term represents the fraction of molecular collisions with enough energy to react. Because temperature sits in that exponential, the rate constant is extremely sensitive to it — a modest temperature increase can multiply the rate several times, which is why reactions generally speed up sharply when heated. A lower activation energy, such as one provided by a catalyst, also raises the rate constant dramatically. The pre-exponential factor accounts for collision frequency and orientation. This calculator uses the gas constant of 8.314 joules per mole-kelvin and converts your activation energy from kilojoules. Enter the pre-exponential factor, activation energy, and temperature to find the rate constant.
k = A × e^(−Ea / (R·T))
The Arrhenius equation gives the rate constant as the pre-exponential factor A times e raised to the power of negative activation energy divided by the gas constant times absolute temperature. Higher temperature or lower activation energy makes the exponent less negative, raising the rate constant and speeding the reaction.
Source:Arrhenius equation (chemical kinetics)
We review all our tools from time to time to ensure accuracy.
Worked example
How to use this calculator
- 1
Enter your values
Enter pre-exponential factor (a), activation energy (ea) and temperature.
- 2
Adjust the options
Fine-tune any options to match your situation.
- 3
Read your result
Rate constant (k) updates instantly as you type — no button to press.
Frequently asked questions
What is the Arrhenius equation?
The Arrhenius equation, k = A·e^(−Ea/RT), relates a reaction's rate constant to temperature and activation energy. A is the pre-exponential factor, Ea the activation energy, R the gas constant, and T the absolute temperature. It quantifies how the rate constant, and thus reaction speed, rises with temperature and falls with a higher energy barrier.
What is activation energy?
Activation energy is the minimum energy that colliding molecules must have for a reaction to occur, the barrier separating reactants from products. A higher activation energy means fewer collisions succeed, so the reaction is slower. Lowering it, for example with a catalyst, allows more collisions to react and speeds the reaction, raising the rate constant substantially.
Why does temperature affect the rate constant so strongly?
Because temperature appears in the exponential term of the Arrhenius equation, small increases produce large changes in the rate constant. Raising the temperature increases the fraction of molecules with enough energy to overcome the activation barrier. As a rule of thumb, many reaction rates roughly double for every 10-degree Celsius rise, reflecting this exponential sensitivity.
What is the pre-exponential factor?
The pre-exponential factor, or frequency factor A, represents how often molecules collide with the correct orientation to react, independent of energy. It sets the maximum possible rate constant if every collision had enough energy. Combined with the exponential energy term, it gives the actual rate constant. Its value depends on the specific reaction and is often determined experimentally.
How does a catalyst change the rate constant?
A catalyst provides an alternative reaction pathway with a lower activation energy. Since activation energy sits in the exponent, even a modest reduction sharply increases the rate constant and speeds the reaction, without the catalyst being consumed. This is why catalysts are so powerful and widely used in industry and biology to accelerate reactions efficiently.