How Catalysts Work: An Alternative Reaction Pathway with a Lower Activation Energy
To understand catalysts, let's revisit the concept of activation energy ($E_a$).
Definition
Activation energy
Activation energy ($E_a$) is the minimum energy required for colliding particles to form an activated complex (also known as the transition state) and proceed to products.
- A catalyst speeds up a reaction by providing an alternative pathway with a lower activation energy.
- This means more particles now have enough energy to overcome the activation barrier, leading to a higher frequency of successful collisions.
Key Characteristics of Catalysts:
- Not consumed: Catalysts participate in the reaction but are regenerated at the end, so they are not used up.
- Do not alter equilibrium: A catalyst accelerates both the forward and reverse reactions equally, leaving the equilibrium position and the overall enthalpy change ($\Delta H$) unchanged.
- Specificity: Some catalysts are highly selective, working only for specific reactions (e.g., enzymes in biological systems).
A catalyst lowers the activation energy but does not affect the energy levels of the reactants or products: it only changes the pathway.
Energy Profiles: Catalyzed vs. Uncatalyzed Reactions
- Energy profiles are graphical tools that help us visualize the energy changes during a chemical reaction.
- They provide a clear picture of how catalysts reduce activation energy.
Components of an Energy Profile:
- Reactants: The starting substances, represented on the left side of the graph.
- Products: The substances formed, shown on the right side.
- Activation Energy ($E_a$): The energy barrier that reactants must overcome to reach the transition state.
- Transition State: The peak of the energy barrier, representing the unstable arrangement of atoms during the reaction.
- Enthalpy Change ($\Delta H$): The difference in energy between reactants and products.
Comparing Catalyzed and Uncatalyzed Profiles:
- In an uncatalyzed reaction, the activation energy is higher, meaning fewer particles have sufficient energy to react.
- A catalyst lowers this barrier, enabling more particles to react and speeding up the process.
The Decomposition of Hydrogen Peroxide
- The decomposition of hydrogen peroxide ($2H_2O_2 \rightarrow 2H_2O + O_2$) is slow at room temperature.
- Adding manganese dioxide ($MnO_2$) as a catalyst provides an alternative pathway with lower activation energy, significantly speeding up the reaction.
- Examine an energy profile diagram. Can you identify the activation energy for both the catalyzed and uncatalyzed reactions?
- How does the transition state differ in energy between these two pathways?
Types of Catalysts: Homogeneous and Heterogeneous
Catalysts are classified based on their phase (solid, liquid, gas) relative to the reactants.
