AP Chemistry 5.10 Multistep Reaction Energy Profiles Study Notes - New Syllabus Effective fall 2024
AP Chemistry 5.10 Multistep Reaction Energy Profiles Study Notes.- New syllabus
AP Chemistry 5.10 Multistep Reaction Energy Profiles Study Notes – AP Chemistry – per latest AP Chemistry Syllabus.
LEARNING OBJECTIVE
Represent the activation energy and overall energy change in a multistep reaction with a reaction energy profile.
Key Concepts:
- Reaction Mechanisms
- Reaction Mechanism & Rate Law
- Pre-Equilibrium Approximation
- Multistep Reaction Energy Profile
5.10.A.1 Energetics of Elementary Reactions:
1. Elementary Reactions and Mechanisms:
In chemical reactions, reaction intermediates and transition states are significant terms:
– Reaction Intermediates: Short-lived species that are formed during the course of the reaction and are not present in end products. They are stable enough to be detected and are engaged in subsequent steps.
Example: In the reaction between hydrogen and iodine to form HI, the iodine atom (I) can be an intermediate.
– Transition States: Unstable, high-energy atom configurations at the end of the reaction pathway. They cannot be trapped and exist for an instant.
Example: In a bimolecular reaction A + B → C, the transition state is an unstable configuration prior to the formation of C.
| Aspect | Reaction Intermediates | Transition States |
|---|---|---|
| Stability | Relatively stable, detectable. | Unstable, cannot be isolated. |
| Energy | Lower than transition states. | Highest energy point in the reaction. |
| Lifetime | Longer, can be observed. | Very short-lived. |
Mechanism Role: Intermediates provide a sequential pathway of the reaction, while the transition state determines the speed of the reaction via activation energy.
2. Activation Energy and Energy Profiles:
i. Activation Energy (Ea):Activation Energy (Ea) is the minimum energy required for a reaction to occur. It is the energy barrier that must be overcome for reactants to form products. The greater the Ea, the slower the reaction.
ii. Exothermic Reaction: Energy is released; products are lower in energy than reactants.
Endothermic Reaction: Energy is absorbed; products are more energetic than reactants.
iii. Key Points:
Activation Energy (Ea): Energy difference between reactants and transition state.
ΔH: Energy difference between reactants and products.
3. Rate-Determining Step:
The rate-determining step (RDS) is the slowest step in a reaction mechanism, governing the overall reaction rate. It is usually the step with the maximum activation energy and is the reaction bottleneck. The rate of the whole reaction is governed by the energy barrier of this step.
i. Influence on the Energy Profile:
– The rate-determining step is the highest point (transition state) of the energy profile.
– The difference in energy between reactants and transition state of rate-determining step determines the activation energy (Ea) for the reaction.
– An increased RDS activation energy results in a slower rate for the overall reaction since more energy is needed in order to reach this barrier.
ii. Summary:
– The rate-determining step determines the net reaction rate as it possesses the maximum activation energy.
– It controls the energy profile by being the maximum, controlling the total Ea for the reaction.
4. Exothermic vs Endothermic Reactions:
Exothermic Reactions:
– Energy Change: Emits energy.
– Energy Profile: Products have lower energy than reactants.
– Effect: Radiates heat, warming surroundings.
– Example: Combustion (e.g., fuel burning).
Endothermic Reactions:
– Energy Change: Absorbs energy.
– Energy Profile: Products have higher energy than reactants.
– Effect: Absorbs heat, cooling surroundings.
– Example: Photosynthesis.
| Aspect | Exothermic | Endothermic |
|---|---|---|
| Energy Change | Released | Absorbed |
| Products’ Energy | Lower than reactants | Higher than reactants |
| Effect on Surroundings | Warms | Cools |
5. Catalysis and Its Effect on Energy Profiles:
Catalysts are substances that speed up a chemical reaction without themselves being consumed during the process. They act by reducing the activation energy (Ea), allowing the reactants to reach the transition state and product with greater ease.
i. How Catalysts Lower Activation Energy:
– Catalysts possess a lower activation energy reaction pathway compared to the uncatalyzed reaction.
– This lower activation energy increases the rate of the reaction since it requires less energy to overcome the barrier.
ii. Effect on Energy Profile:
– Without Catalyst: Higher activation energy with larger peak (transition state).
– With Catalyst: Lower activation energy and the reaction follows a different, lower-energy path. The total energy change (ΔH) remains the same.
iii. Comparison of Energy Profiles:
a. Uncatalyzed Reaction:
– Higher activation energy.
– Raised peak (transition state).
b. Catalyzed Reaction:
– Reduced activation energy.
– Reduced peak (transition state).
iv. Summary:
– Catalysts reduce the activation energy and provide an alternative reaction pathway.
– This accelerates the reaction, but it doesn’t change the overall energy change (ΔH) of the reaction.
Multistep Reaction Energy Profiles
- The slowest (rate-determining) step in a mechanism will have the highest activation energy
- On graph, will have the highest peak
- Number of peaks = number of elementary steps
- Intermediates are found in minimums of graph; the transition states are found in the maximum of graphs