AP Chemistry 5.11 Catalysis Study Notes - New Syllabus Effective fall 2024
AP Chemistry 5.11 Catalysis Study Notes.- New syllabus
AP Chemistry 5.11 Catalysis Study Notes – AP Chemistry – per latest AP Chemistry Syllabus.
LEARNING OBJECTIVE
Explain the relationship between the effect of a catalyst on a reaction and changes in the reaction mechanism.
Key Concepts:
- Catalysts & Reaction Rate
- Catalysts in Action
5.11.A.1 Catalyst Effect on Reaction Rate:
1. Catalyst Function:
A catalyst does this by providing a reaction route with less activation energy, thus making the reaction simpler and occurring at a quicker rate. :
i. Decreasing Activation Energy: In any chemical reaction, molecules need to expend some energy (activation energy) to form bonds and react. A catalyst lowers this energy barrier so that reactants are able to become products more easily. This doesn’t change the overall energy of the reaction but only the amount of energy necessary to start the process.
ii. Providing an Alternative Pathway: A catalyst often collides with the reactants to form an intermediate complex. This intermediate complex has a lower energy transition state, meaning that the reactants are more likely to collide in such a way as to yield a successful reaction.
iii. Not Consumed in the Reaction: Another property of a catalyst is that it is not consumed during the reaction. After the reaction has occurred, the catalyst is regenerated and may be reused.
iv. Higher Collision Frequency: By stabilizing the transition state or by assisting in the orientation of reactants, catalysts can increase the likelihood that reactants will collide in the correct orientation for the reaction to proceed.
In short, catalysts speed up reactions by lowering the activation energy so more molecules of the reactants have a means to achieve the necessary level of energy in order for the reaction to occur and by providing a different more preferred path for the reaction to proceed.
2. Effective Collisions:
A catalyst plays a significant role in the energy required to initiate a reaction, not in increasing the collision frequency of reactants. Here’s what occurs:
i. Collision Frequency: Catalysts do not change the collision frequency directly between molecule colliding. The number of collisions among reactant molecules relies significantly on concentration, temperature, and pressure.
– But a catalyst can make collisions more effective by lowering the activation energy. What this means is that, out of any collection of collisions, more of them will have enough energy to overcome the lowered activation energy barrier and yield a successful reaction.
ii. Energy (Activation Energy):
– The main task of a catalyst is to lower the activation energy of the reaction. This does not change the net energy of the reactants or products but enables the reactants to acquire the necessary energy in order to react.
– By lowering the activation energy, the catalyst allows more reactant molecules to have enough energy to successfully collide and react. This increases the efficiency of the collisions, though the collision frequency remains the same.
Summary:
– Collision frequency (molecule collision frequency) isn’t significantly changed by a catalyst.
– Energy is affected by a catalyst: it lowers the activation energy, and there is more likelihood that collisions will result in a reaction.
3. Activation Energy: A catalyst lowers activation energy by providing an alternative reaction pathway with a lower energy barrier. It forms an intermediate complex with reactants, stabilizing the transition state and making it easier for the reaction to occur. This increases the likelihood of successful collisions without changing the overall energy of the reactants and products.
4. Reaction Pathway:
A catalyst provides an alternative, lower-energy reaction pathway by forming temporary complexes with reactants, which stabilizes the transition state. This lowers the activation energy required for the reaction to proceed. Essentially, the catalyst guides the reactants to react in a way that requires less energy, allowing the reaction to occur more easily and at a faster rate.
5.11.A.2 Catalyst Consumption and Regeneration in Reaction Mechanism:
1. Catalyst Role and Net Concentration:
A catalyst maintains a constant concentration because it is not consumed in the reaction. Here’s why:
i. Regeneration: The catalyst can temporarily bond with reactants in the reaction, but once the reaction is over, it is released in its initial state. This regeneration is such that the catalyst remains available to participate in other reactions without being consumed.
ii. Catalyst Cycle: The catalyst goes through a cycle in which it catalyzes the reaction and then returns to its initial state to catalyze another reaction. It continues repeating this cycle, so the amount of catalyst in the reaction remains unchanged at all times.
A catalyst thus speeds up reactions without being permanently altered, such that its concentration remains constant throughout the process.
2. Rate-Determining Step and Catalyst:
In a mechanism of a reaction, the catalyst is temporarily consumed during the course of the reaction but is regenerated at the termination of the process. Below is the step-by-step process:
i. Catalyst Consumption:
– In the initial step of the reaction mechanism, the catalyst may react with a reactant to give an intermediate complex. This step usually lowers the activation energy and facilitates the reaction.
– The catalyst enables bond breaking or making in this middle stage but becomes a part of the reaction only temporarily.
ii. Catalyst Regeneration:
– The intermediate complex breaks down in a later stage, leading to the production of end products.
– The catalyst is released back to its original form, which can again be used for repeated cycles of the reaction in the future.
iii. Catalyst Mechanism Example:
– During an oxidation reaction, a catalyst like platinum can initially bind to a reactant molecule such that it will be more inclined to react. After the product is formed, the catalyst is regenerated into its original form and remains unchanged, enabling it to continue catalyzing further reactions.
iv. Role in Rate-Determining Step:
– The rate-determining step is the slowest step in a reaction mechanism. Though the catalyst speeds up the reaction overall by lowering activation energy, it affects the rate-determining step significantly, making it occur faster.
Briefly speaking, the catalyst is used during the reaction as an intermediate but regenerated when the reaction ends so that it remains in a constant concentration and induces the reaction repeatedly.
5.11.A.3 Catalyst Binding and Reaction Acceleration:
1. Catalyst-Reactant Binding:
A catalyst binds to reactants via transitory interactions (bonds or weak attractions such as van der Waals) that stabilize the transition state and reduce the activation energy. This is how it goes:
i. Binding to Reactants: The catalyst is bound to the reactants, most often by the creation of a transitory complex. This is achieved by physical adsorption or chemical bonding (such as hydrogen bonding, coordination bonds, etc.), depending on the type of catalyst.
ii. Correct Orientation: The catalyst places the reactants in the correct orientation for the reaction to occur more effectively by binding onto the reactants. This improves the likelihood of successful collisions among reactant molecules.
iii. Stabilization of Transition State: The catalyst can decrease the activation energy by stabilizing the energetic transition state of the reaction. This makes it require less energy for the reactants to get to the transition state so that there can be more easy reaction with less energy input.
iv. Lower Activation Energy: The catalyst merely lowers the energy required for breaking or forming bonds by facilitating a lower-energy pathway for the reaction. This creates a faster rate of reaction without the catalyst’s consumption.
Overall, the catalyst binds reactants to allow favorable conditions for reaction, lowers activation energy through transition state stabilization, and enables easier reaction to proceed.
2. Reaction Intermediates:
Reaction intermediates are temporary and unstable species that arise during a reaction but are not the final products. They exist only briefly and are quickly converted into the final products.
1. How Catalysts Form Reaction Intermediates:
i. Catalyst-Reactant Interaction: Initially, the catalyst interacts with the reactants, often binding to them to create a complex. This interaction may involve temporary bonds or other weak forces, such as hydrogen bonds, van der Waals forces, or coordination bonds.
ii. Intermediate Formation: Throughout the reaction mechanism, the catalyst can transfer energy or electrons to the reactants or form a transient complex. This process results in the creation of an intermediate, which has a lower energy state than the transition state but is not yet a product.
iii. Intermediate Conversion: The intermediate is then transformed into the final products in subsequent steps. The catalyst aids this transformation by stabilizing the intermediate or providing a pathway with lower activation energy for the conversion.
iv. Regeneration of the Catalyst: Once the intermediate is formed and converted into the final products, the catalyst is released in its original form, ready to catalyze additional reactions.
v. Example:
– In heterogeneous catalysis, such as with a metal catalyst, the catalyst may form a complex with the reactant, breaking or forming bonds to generate an intermediate. This intermediate will then undergo further transformations, ultimately yielding the products, while the catalyst remains unchanged.
vi Summary:
– Reaction intermediates are short-lived species produced during the reaction process, typically facilitated by the catalyst.
– The catalyst generates these intermediates by interacting with reactants and reducing the activation energy for their formation, thereby accelerating the overall reaction.
5.11.A.4 Covalent Bonding in Catalysis: Acid-Base:
1. Covalent Bonding and Acid-Base Catalysis:
Covalent bonding and proton transfer are essential components of acid-base catalysis and play a significant role in the mechanisms of various reactions involving catalysts. Here’s how they influence the reaction process:
i. Covalent Bonding in Catalysis:
– Covalent catalysis refers to the creation of a temporary covalent bond between the catalyst and the reactant. This interaction stabilizes the transition state and reduces the activation energy required for the reaction.
– The catalyst establishes a covalent intermediate with the reactant, which facilitates the breaking or forming of bonds more efficiently. Once the reaction is complete, the catalyst regenerates and can be reused.
– Example: In certain enzyme-catalyzed reactions, a nucleophile on the enzyme forms a covalent bond with the substrate, resulting in a reaction intermediate. This intermediate then transforms, leading to product formation, while the enzyme (catalyst) is restored to its original state.
ii. Proton Transfer in Acid-Base Catalysis:
– Acid-base catalysis involves the transfer of protons (H⁺), which can either activate or deactivate the reactants, making the reaction more favorable.
– Proton donation (acid catalysis): An acidic catalyst donates a proton to a reactant, increasing its electrophilicity and thus its reactivity (which aids in bond cleavage or rearrangement).
– Proton abstraction (base catalysis): A basic catalyst accepts a proton from a reactant, enhancing its nucleophilicity and likelihood of participating in a reaction.
– Both acid and base catalysis contribute to lowering the activation energy by stabilizing transition states or intermediates through proton transfer.
– Example: In the esterification reaction, an acid catalyst (like H₂SO₄) donates a proton to the reactant, boosting its capacity to bond with another reactant (such as alcohol), resulting in the formation of an ester.
2. Reaction Intermediates:
Reaction intermediates are transient species that are formed in a reaction, e.g., covalent bonds, carbocations, or enzyme-substrate complexes. They are more reactive than the reactants and are transformed into products.
Elementary reactions are elementary steps of the reaction mechanism, in which intermediates are formed or consumed, often involving bond formation, bond cleavage, or proton transfer. Examples include proton transfers in acid-base catalysis or enzyme-substrate binding in biological catalysis.
5.11.A.5 Surface Catalysis and Reaction Intermediates:
1. Surface Binding and Covalent Bonding:
Reactants or intermediates are bound to the surface of a catalyst by surface binding , which can be by covalent bonding or non-covalent interactions. The following is how they happen:
i. Surface Binding through Non-Covalent Interactions:
– Physical adsorption: Intermediates or reactants can adsorb onto the catalyst surface via weak forces like van der Waals forces or hydrogen bonding. These interactions are reversible and allow the orientation of the reactants in a way that enables the reaction to proceed.
– Example: During heterogeneous catalysis, reactants adsorb onto metal surfaces, taking up a configuration that minimizes activation energy for the formation or breaking of bonds.
ii. Surface Binding through Covalent Bonding:
– Covalent bonding: Occasionally, reactants or intermediates form stronger bonds with the catalyst surface. This may involve the transfer of electrons from the reactant to the surface to form a covalent bond.
– Example: In enzyme catalysis, the active site of the enzyme can covalently bind to the substrate and create a reaction intermediate. The intermediate is converted, and the product is released, and the catalyst is restored.
iii. Summary:
– Non-covalent binding (e.g., van der Waals forces) allows reactants to bind to the surface temporarily, positioning them for the reaction.
– Covalent bonding leads to stronger and more stable interactions that may lead to the formation of reaction intermediates, which are then transformed into products.
2. Reaction Intermediates and Elementary Reactions:
Bound intermediates create new elementary reactions through stabilizing the reactant and lowering the activation energy for the reaction to take place. This is how it works:
i. Formation of Bound Intermediate:
– Reactants or intermediates are attached to the catalyst surface or active site (in enzymes) by non-covalent or covalent bonds. This attachment is what positions the molecules in optimal orientation for the reaction.
– The bound intermediate is more reactive than the free reactant because the catalyst stabilizes the intermediate and provides a favorable environment for the reaction.
ii. New Elementary Reactions:
– The bound intermediate can form new elementary reactions that were not possible prior to binding. These reactions are likely to involve bond-breaking or bond-forming steps catalyzed by the catalyst.
– The role of the catalyst is to lower the activation energy of such elementary steps, which speeds up the reaction.
iii. Example:
– In enzyme catalysis, the reactant (substrate) is drawn towards the active site of the enzyme and comes together to give an intermediate complex. This intermediate exists at a lower state of energy and hence makes it easy for the reaction to proceed through elementary processes like transfer of protons or bond rearrangement. The product is removed when the intermediate has reacted and the enzyme remains free to catalyze another reaction.
Catalysis
- Catalyst: chemical agent that speeds up a reaction without being consumed by the reaction → can be used over and over again
- Do not affect the free energy/enthalpy of a reaction!
- Reactions with a catalyst will typically require at least 2 steps
- Catalysts allow reactions to occur with a lower AE → lower activation energy means that more collisions will have enough energy to overcome AE and form a product → reaction rate increased
- Homogeneous catalyst: is present in the same phase as the reacting molecules
- Heterogeneous catalyst: exists in a different phase (usually as a solid; a catalytic converter is one type)
- 4 steps of heterogeneous catalysis
- Adsorption (collection of one substance on the surface of another substance) and activation of the react
ants - Migration of the adsorbed reactants on the surface
- Reaction of the adsorbed substances
- Escape (desorption) of the products
- Adsorption (collection of one substance on the surface of another substance) and activation of the react
How Catalysts Speed Up Reactions
- A catalyst increases the rate constant and lowers Ea Barrier by…
- Forming a more stable activated complex
- Increased collision frequency
- Improved orientation effects
- Allow chemical reactions to occur at lower temperatures
- Speeds up natural reactions ≠ cause them
- Reactions can occur without catalysts, but would be slower & cost a lot more energy
Activation Energy Barrier
|
