AP Chemistry 5.7 Introduction to Reaction Mechanisms Study Notes - New Syllabus Effective fall 2024
AP Chemistry 5.7 Introduction to Reaction Mechanisms Study Notes- New syllabus
AP Chemistry 5.7 Introduction to Reaction Mechanisms Study Notes – AP Chemistry – per latest AP Chemistry Syllabus.
LEARNING OBJECTIVE
Identify the components of a reaction mechanism.
Key Concepts:
- Reaction Mechanisms
- Reaction Mechanism & Rate Law
- Pre-Equilibrium Approximation
- Multistep Reaction Energy Profile
5.7.A.1 Reaction Mechanism: Sequence of Elementary Reactions:
1. Elementary Reactions:
An elementary reaction is a single-step process where reactants are directly transformed into products. Elementary reactions are one-step processes and depend on the molecularity of the process for the rate of reaction. In simple terms, reactants in elementary reactions are combined to produce products in a single transition state without any intermediates or multi-step processes.
i. Characteristics of Elementary Reactions:
Single Step Process: Reaction takes place in a single event or step.
No Intermediates: No intermediate species which are the products formed between the process take place. It all happens with a single reaction event.
Rate Law: An elementary rate law for the reaction is formulated directly from molecularity. Rate constant and a reactant concentration are sufficient to formulate the rate.
Molecularity: It is the number of molecules in the elementary step. Molecularity of the reaction also determines the number of molecules that should collide in order for the reaction to happen.
| Type of Reaction | Molecularity | General Form | Rate Law | Order of Reaction | Example | Characteristics |
|---|---|---|---|---|---|---|
| Unimolecular | 1 | First-order | Decomposition of N₂O₅: | – Involves only one molecule. – Rate depends on the concentration of one reactant. | ||
| Bimolecular | 2 | Second-order | Reaction between NO₂ and CO: | – Involves two molecules. – Rate depends on the concentration of two reactants. |
- Unimolecular reactions involve one molecule and are typically first-order.
- Bimolecular reactions involve two molecules and are typically second-order
2. Reaction Intermediates and Products:
Role of intermediates and separation from the reactants and products is crucial in a chemical reaction while studying the reaction mechanism.
i. Reaction Intermediates:
Definition:
Intermediates are the molecules which are produced in the process of a reaction and become used up in further steps. They are not found in the balanced equation of the overall reaction of the reaction because they never appear back in the product.
Role:
Intermediates provide the pathway from reactant to product via a multi-step mechanism (reaction mechanism). Intermediates are important for understanding how a reaction takes place at a molecular level.
Characteristics:
Short-lived Species: Intermediates are temporary and occur in certain steps of the reaction alone.
Not in Overall Balanced Equation: They get used up in subsequent steps and are not part of the overall chemical equation.
Can Be Isolated (At Times): The intermediates can be isolated and identified in certain situations if the reaction is not very rapid.
Example:
In a reaction between hydrogen and iodine to produce HI, the intermediate is the hydrogen iodide radical (HI·), which is generated and used up during the reaction.
ii. Products:
Definition:
Products are final chemicals that result after the reactants have already completed transformation after the reaction. They are included in the final balanced equation.
Role:
Products are products of a chemical reaction. They are what is left after all of the reactants are consumed and usually the stable species.
Features:
Stable Species: Products are stable and usually towards the end of the reaction.
Included in Overall Equation: Products are accounted for in the overall balanced chemical equation of the reaction.
Example:
Combustion products of methane are carbon dioxide (CO₂) and water (H₂O).
3. Catalysts:
A catalyst is a material that increases the rate of a chemical reaction without being destroyed in the reaction. It reduces the activation energy, offering an alternative route for the reaction, and is able to be used over and over again.
i. Role of Catalysts:
a. Increase Reaction Rate: Through reducing activation energy.
b. Not Consumed: They are re-formed at the end of the reaction.
c. Selective: Able to favor the production of certain products.
d. No Effect on Equilibrium: They increase the rate at which equilibrium is established but do not shift the position.
ii. Mechanism of Catalysis:
a. Adsorption: Reactants adsorb on the surface of the catalyst (in the case of heterogeneous catalysts).
b. Intermediate Formation: Reactants form transient intermediates.
c. Reaction: Reactants get transformed.
d. Desorption: Products are desorbed and the catalyst is ready for reuse
iii. Types of Catalysts:
1. Homogeneous Catalysts: Same phase as reactants (e.g., liquid acid catalysts).
2. Heterogeneous Catalysts: Of a different phase than the reactants (e.g., solid catalysts such as iron in the Haber process).
3. Enzymes: Biological catalysts within living organisms.
iv. Example:
– Haber Process: Iron catalyst to synthesize ammonia from hydrogen and nitrogen:
5.7.A.2 Elementary steps must match the overall balanced equation:
1. Elementary Steps and Reaction Mechanisms:
i. Basic Steps:
Elementary steps are individual, simple reactions where reactants transform directly into products. Each step involves collisions among molecules, and its rate law is controlled by the stoichiometry. They can be unimolecular, bimolecular, or termolecular.
ii. Reaction Mechanisms:
A reaction mechanism is a sequence of elementary steps that describe how the reactants transform into products. It includes all steps, intermediates, and the net reaction. The slowest step in a reaction is the rate-determining step (RDS) that dictates the overall rate of the reaction.
iii. Combining Steps:
The net reaction is the sum of the elementary steps. Intermediates formed in one step are consumed in another. The rate law is usually determined by the rate-determining step.
2. Balanced Chemical Equation:
For a reaction mechanism to be valid, the elementary steps must accommodate the overall balanced chemical equation through stoichiometry and conservation laws (e.g., conservation of mass and energy).
i. Key Points on Alignment:
a. Stoichiometry:
– The addition of coefficients in the elementary steps must be equal to the coefficients in the entire balanced equation.
– The products and reactants in the elementary reactions should add up to give the correct amount in the end reaction.
Example: If the overall reaction is:
The elementary steps might be:
The total of the steps should have the same reactants and products as the net reaction.
b. Conservation of Mass and Energy:
– There should be the same number of atoms of each element on both sides of the equation, maintaining mass.
– Similarly, energy must be saved in the reaction (though this is more of an abstract concept and involves activation energy and bond energies).
c. Intermediates:
– Intermediates (species that are created and consumed during the course of the mechanism) must balance out when adding the elementary steps. They must not appear in the ultimate balanced equation.
ii. Example:
For the overall reaction:
The elementary steps might be:
Here, X is an intermediate that is both used to produce and get destroyed, thus eliminating itself in the final equation.
Thus, the conservation rules and the stoichiometry ensure that the mechanism correctly describes the overall balanced equation.
3. Combining Elementary Steps:
When combining elementary steps to derive the overall reaction, several techniques are employed to ensure consistency and alignment with the balanced equation:
i. Add the Elementary Steps:
– Begin by adding the equations for all elementary steps. This will yield the overall reaction.
– Make sure that the reactants and products from each elementary step are correctly summed.
Example:
If the elementary steps are:
- Step 1:
- Step 2:
Adding them gives the overall reaction:
The final equation should be:
ii. Cancel Intermediates:
– Intermediates are substances that are generated in one elementary step and consumed in another. They do not appear in the overall balanced equation.
– When adding the elementary steps, eliminate any intermediates that show up on both sides of the reaction.
Continuing the previous example:
The intermediate
cancels out, resulting in the final balanced equation:
iii. Adjust Coefficients (if necessary):
– If any of the elementary steps need to be reversed to align with the direction of the overall reaction, reverse the equation and invert the sign of the rate constant.
– Additionally, multiply the entire step by a factor if needed to match the stoichiometric coefficients of the overall reaction.
4. Reaction Rate and Mechanism:
The rate law is the expression of the dependence of a reaction rate on the concentration of reactants. It is obtained from the rate-determining step (RDS) of the reaction mechanism, i.e., the slowest step in the sequence of elementary steps.
– The rate law agrees with the stoichiometry of the RDS, and not with the overall balanced equation.
– Easy steps are responsible for the rate law in accordance with the molecularity of the molecules that react.
– The balanced equation is the total reaction, but the rate law can be different because of intermediates and stepwise mechanism.
In short, the rate law is connected to the mechanism, i.e., the slow step, and can be different from the stoichiometric coefficients of the overall balanced equation.
5.7.A.3 Reaction Intermediate:
1. Definition and Role:
Intermediates are species that form and are consumed during a chemical reaction but do not show up in the overall balanced equation. They exist only temporarily during one or more elementary steps of a reaction mechanism.
i. Formation and Consumption of Intermediates:
a. Formation:
Intermediates arise in one elementary step. For instance, during a reaction, an unstable compound may be produced as a product of one step but is not stable enough to persist and instead reacts in the following step.
b. Consumption:
Intermediates are utilized in later elementary steps, resulting in the creation of the final products. They do not feature in the overall balanced equation because they are not present at the beginning or end of the reaction.
c. Example:
In the reaction mechanism for the breakdown of hydrogen peroxide (H₂O₂):
– Step 1: H₂O₂ → H₂O + O* (oxygen atom, an intermediate).
– Step 2: O* + H₂O₂ → H₂O + O₂ (oxygen molecule, final product).
In this case, O* (the oxygen atom) is an intermediate formed in step 1 and consumed in step 2.
ii. Key Points:
– Intermediates are short-lived and exist only between the steps of the mechanism.
– They do not appear in the overall balanced equation because they are consumed before the reaction concludes.
2. Elementary Steps and Mechanism:
Intermediates are species that are produced in one elementary step and manipulated in another. They are essential to the mechanism of the reaction but are not present in the balanced equation as a whole since they are transient.
i. How Intermediates are Produced and Consumed:
a. Production of Intermediates:
– All reactants are consumed in a single step when they collide to produce an intermediate. This is a product of the reaction but too unstable to be considered as one of the end products.
– Example: During the reaction of hydrogen and iodine (H₂ + I₂ → 2HI), an intermediate H-I (hydrogen-iodine complex) might be produced before it breaks down into the end products.
b. Consumption of Intermediates:
– The intermediate is not present for very long in the reaction; it is consumed during the subsequent elementary step.
– Example: In a two-stage reaction, an intermediate can react with another reactant to give the end product(s).
ii. Example: A Two-Step Reaction Mechanism
For hydrogen peroxide decomposition (H₂O₂):
– Step 1: H₂O₂ → H₂O + O* (oxygen atom is an intermediate).
– Step 2: O* + H₂O₂ → H₂O + O₂ (the product is the oxygen molecule).
Here in this mechanism:
– The O* (oxygen atom) is formed in Step 1 as an intermediate.
– O* is destroyed in Step 2 to form the product O₂.
iii. Summary:
– Intermediates are formed in one elementary step.
– They are destroyed in another simple step, such that they do not appear in the overall balanced equation.
– Intermediates are created and used in a way that is central in linking the elementary steps of a reaction mechanism.
3. Identification and Example:
Intermediates are usually transient species present within the mechanism of the reaction but not within the overall balanced equation. You can identify intermediates in a mechanism by searching for species which:
– Are formed in one elementary step.
– Are destroyed in another following elementary step.
– Do not show up within the overall reaction.
i. Typical Examples of Intermediates
a. Carbocations (carbon species with a positive charge):
– They are created when a molecule has undergone heterolytic fission of a bond, leaving behind a positively charged carbon atom.
– Example: In the SN1 reaction between tert-butyl chloride (C₄H₉Cl) and water:
– Step 1 (Formation of carbocation): C₄H₉Cl → C₄H₈⁺ + Cl⁻ (tertiary carbocation).
– Step 2 (Nucleophilic attack): C₄H₈⁺ + H₂O → C₄H₉OH₂⁺ (tert-butyl alcohol).
– In this instance, the tertiary carbocation is both an intermediate formed and one used in multiple stages of the mechanism.
b. Free Radicals (species with an unpaired electron):
– Free radicals are formed when a bond homolytically breaks (each atom captures an electron from the bond).
– Example: In the halogenation of alkanes (chlorination of methane):
– Step 1 (Initiation): Cl₂ → 2Cl· (chlorine radicals).
– Step 2 (Propagation): CH₄ + Cl· → CH₃Cl + H· (hydrogen atom and methyl radical produced).
– Step 3 (Propagation): CH₃Cl + Cl· → CH₂Cl₂ + HCl.
– The chlorine radical (Cl·) and methyl radical (CH₃·) are intermediates which are formed and consumed in the propagation steps.
c. Carbanions (negatively charged carbon species):
– Carbanions are nucleophilic substitution or elimination reaction intermediates in which the carbon is a negative charge bearer.
– Example: In the E2 elimination reaction of bromoethane:
– Step 1 (Formation of carbanion): CH₃CH₂Br + OH⁻ → CH₃CH₂⁻ + H₂O.
– Step 2 (Elimination): CH₃CH₂⁻ + CH₃CH₂Br → CH₂=CH₂ + HBr.
– In this, the ethyl anion (CH₃CH₂⁻) is the consumed and produced intermediate.
ii. Summary:
– Intermediates are carbocations, free radicals,carbanions, or other products consumed and generated in a reaction mechanism.
– They do not appear in the overall balanced equation but serve as a key connection to couple the elementary steps.
By knowing how intermediates are consumed and generated, we can more readily predict and rationalize reaction mechanisms.
5.7.A.4 Experimental Detection of Reaction Intermediates:
1. Role of Reaction Intermediates:
i. Definition:
A reaction intermediate is a chemical species that is produced in one step and is also used up in the subsequent step in a chemical reaction. It never occurs in the overall balanced reaction because it’s temporary, lasting only for the duration of the reaction process.
ii. Importance in Reaction Mechanisms:
a. Bridge Between Reactants and Products:
Intermediates are accountable for the stepwise process of a reaction. They are formed and utilized as reactants are gradually being transformed into products, bridging elementary steps.
b. Facilitate Reaction Pathways Description:
The presence of intermediates lends sense to the type of molecular events occurring during the progress of a reaction. By means of them, chemists can map the complex mechanism resulting in final products.
c. Influence Reaction Rate:
Intermediates usually play a crucial role in the identification of the rate-determining step (RDS), which determines the rate of the overall process. Intermediates may influence the rate at which the reaction proceeds.
d. Impact on Reaction Conditions:
Knowledge of intermediates can help in control or optimizing reaction conditions, such as temperature or catalysts, in order to guide specific pathways or products.
Briefly, intermediates are critical in the mechanism of understanding how reactants are transformed into products and play a role in predicting and controlling reaction rates and mechanisms.
2. Experimental Detection Methods:
i. Spectroscopy:
a. UV-Vis Spectroscopy:
Records the absorption of ultraviolet or visible light by a molecule, something that may be evidence for the presence of intermediates with distinctive electronic transitions or absorption patterns.
b. Infrared (IR) Spectroscopy:
Recognizes molecular vibrations and functional groups. Intermediates will possess unique IR absorption peaks, which disclose their structure at the time of the reaction.
c. Nuclear Magnetic Resonance (NMR) Spectroscopy:
Provides detailed information on the structure of intermediates by detecting the magnetic properties of nuclei (e.g., hydrogen or carbon). Changes in the NMR spectra over the course of a reaction might reveal the formation and breakdown of intermediates.
d. Electron Paramagnetic Resonance (EPR) Spectroscopy:
Utilized for the identification of radicals and paramagnetic species (such as free radicals, transition metal complexes), typically fleeting reaction intermediates.
ii. Mass Spectrometry (MS):
– Mass Spectrometry assists in identifying the molecular weight and structure of the intermediates through measuring the mass-to-charge ratio (m/z). The structure of the intermediate may be obtained from the pattern of fragmentation.
– TOF MS can be utilized for monitoring the formation and vanishing of intermediates as a function of time in a reaction.
iii. Trap and Quench Techniques:
– Reaction trapping is where reaction intermediates are quickly trapped by freezing the reaction mixture at one end or the other. The trapped intermediates are then examined using techniques such as NMR or MS.
– Quenching of the reaction (e.g., by addition of a reagent that terminates the reaction) can “freeze” an intermediate at a specific time, and it can subsequently be found and analyzed.
iv. Laser Flash Photolysis:
– An intermediate or a molecule is photolyzed with a laser, and the decay or transformation of the intermediate can be followed in real-time. It is especially useful in the study of short-lived intermediates, like free radicals.
Methods such as spectroscopy (UV-Vis, IR, NMR, EPR) and mass spectrometry are very important for identifying and characterizing reaction intermediates. These yield information on the structure, concentration, and properties of the intermediates and assist in the determination of reaction mechanisms.