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Simple Explanation:
Concept of Activation Energy
Collision Theory: Chemical reactions happen due to reactant particles colliding into each other with enough energy and correct orientation.
If particles don't collide with enough kinetic energy, the reaction won't initate.
If particles don't face the right way, the reaction won't initiate.
Activation Energy: The minimal energy required to initiate a chemical reaction. Reactants must overcome this requirement to become products by methods such as:
Increasing Temperature
Raises kinetic energy of particles. → More frequent and powerful collisions, increasing the chances of overcoming the activation energy barrier.
Increasing Concentration
Particles are closer together. → More frequent collisions, boosting the likelihood of reaction.
Increasing Surface Area
More particles are exposed. → Greater interaction with other reactants, leading to more collisions.
Increasing Pressure
Compresses gas molecules, raising temperature and reducing space between particles. → More frequent collisions due to higher kinetic energy and closer proximity.
Adding catalysts lower the activation energy needed to initiate the reaction, increasing the reaction rate.
Catalysts
Catalysts are substances that speed up a reaction by lowering a reaction's activation energy once introduced to the site. Depending on the substance, they could achieve this either by:
Orienting molecular reactants in an optimal position for reacting. (Known as adsorption.)
Creating alternative elementary steps that require less activation energy to initiate by forming temporary intermediates.
Catalysts must show up at the beginning and end of the reaction, therefore not contributing to the overall reaction equation.
Homogenous Catalysis: When the catalyst used is in the same phase of matter as the reactants.
Heterogenous Catalysis: When the catalyst used is in a different phase of matter than the reactants.
Identifying Catalysts, Overall Reactants, and Intermediates in Elementary Reactions
To view the tutorial for Reaction Mechanisms, click HERE.
Demonstration with symbols:
A + B → C + X
X + Y → B + Z
Intermediate: X (Produced in the first step and then consumed in the second.)
Catalyst: B (Appears at both ends of the reaction sequence.)
To find the overall reaction equation, we cancel out the intermediates and catalysts.
Overall Reaction: A + Y → Z
Potential Energy Diagrams - Catalysis
This picture shows potential energy diagrams for two exothermic reactions. Eₐ stands for activation energy, and ΔH stands for enthalpy (measured in KJ/mol).
The peak(s) of the energy diagrams are the transition states, which highlights the point where chemical bonds break and form.
Arrhenius Equation
After wondering why reactants don't always react when put together, Arrhenius established the concept of activation energy and the Arrhenius Equation as a way to explain this concept.
The Arrhenius Equation states that the value of k (rate constant) increases exponentially when temperature increases at a linear rate.
k = Ae^(-Eₐ/RT)
k: Rate Constant
A: Pre-exponential Factor (Determines the rate of collisions with the correct orientation)
e: Euler's Number (approx. 2.72)
Eₐ: Activation Energy
R: Universal Gas Constant
T: Temperature (Kelvin)
Maxwell Boltzmann Distribution Graph
Temperature is a way of indicating the average kinetic energy of particles.
By raising the temperature, the frequency and power of impact of collisions go up, increasing the reaction rate.
Higher temperatures mean that a greater percentage of particles are able to exceed the Eₐ barrier.
In the picture, T₁ is colder than T₂. This means that the particles with a temperature of T₁ have less molecules on average that have enough kinetic energy to exceed the activation energy.
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