AP Chemistry 5.6 Reaction Profiles Study Notes - New Syllabus Effective fall 2024
AP Chemistry 5.6 Reaction Profiles Study Notes- New syllabus
AP Chemistry 5.6 Reaction Profiles Study Notes – AP Chemistry – per latest AP Chemistry Syllabus.
LEARNING OBJECTIVE
Explain the relationship between the rate of an elementary reaction and the frequency, energy, and orientation of particle collisions.
Key Concepts:
- Reaction Energy Profiles
5.6.A.1 Bond Breaking and Formation in Elementary Reactions:
1. Types of Chemical Bonds:
Chemical bonds are the attractions between atoms in a compound that keep the atoms bonded together. The most important chemical bonds are:
i. Ionic Bonds:
– Formation: Ionic bonds form when an atom loses an electron to another atom, thus forming ions (charged atoms). This most frequently happens between a metal and a non-metal.
– Characteristics:
– The metal atom loses electrons and becomes a positively charged ion (cation).
– The non-metal atom gains electrons and becomes a negatively charged ion (anion).
– The electrostatic attraction between the oppositely charged ions holds them together.
– Examples: Sodium chloride (NaCl), magnesium oxide (MgO).
– Properties: Ionic compounds generally have high melting and boiling points, are generally soluble in water, and are electrically conductive when dissolved in water or melted.
ii. Covalent Bonds:
– Formation: Covalent bonds occur when two atoms share electrons, typically between non-metals. Sharing enables both atoms to have a more stable electron arrangement (typically like the closest noble gas).
– Characteristics:
– Atoms exchange pairs of electrons in a covalent bond to fill their outer electron shell (octet rule).
– Covalent bonds may be polar (in which electrons are not evenly distributed) or nonpolar (in which electrons are evenly distributed).
– Examples: Water (H₂O), oxygen (O₂), methane (CH₄).
– Properties: The melting and boiling points of covalent compounds are lower than those of ionic compounds, and covalent compounds are not electric conductors in either the solid or liquid state.
iii. Bond Dissociation Energy (BDE)
Definition: Bond dissociation energy is the energy needed to dissolve a bond in a mole of a molecule in the gas phase. It’s an indicator of the strength of the bond.
– Explanation: The greater the bond dissociation energy, the more stable is the bond. For instance, a highly dissociated bond would take more energy to break it down.
– Factors:
– Bond type: Single bonds are less energetic than double or triple bonds with respect to their dissociation.
– Bond length: More compact bonds (smaller size of the atom or multiple bonding) will exhibit higher dissociation energies.
– Electronegativity: The more electronegative a given atom is, the tighter it will grip another atom.
– Example: The O-H bond dissociation energy in water is quite high, reflecting the strength of the bond.
5.6.A.2 Reaction Coordinate and Molecular Motions:
1. Definition and Role:
A reaction coordinate is an imaginary variable that describes the evolution of a chemical reaction. It is a theoretical path taken to establish how the molecular structure of reactants changes as it evolves to products. The reaction coordinate is one method of illustrating how the bonds and atoms of a molecule change as the reaction evolves.
i. Key Points Regarding the Reaction Coordinate:
– The reaction coordinate is usually plotted on the x-axis of a reaction progress diagram.
– It does not represent any particular measurable quantity but the sequence of changes the system experiences during the course of the reaction (e.g., bond breaking, bond formation, and atom rearrangement).
ii. Role of the Reaction Coordinate:
a. Describes Reaction Progress:
– When a reaction happens, the molecules change from a starting state (reactants) to an ending state (products). The reaction coordinate is a way of explaining this change, tracing out the smooth path from reactants to products.
b. Represents Energy Changes:
– Along the reaction coordinate, the potential energy of the system varies. This may be graphed in a potential energy diagram (or energy profile), which indicates how the system’s energy varies as the reaction proceeds.
c. Activation Energy and Transition State:
– Along the reaction coordinate, the reaction will have a transition state, which is the highest point on the reactants-to-products trajectory. The difference in energy between reactants and the transition state is referred to as the activation energy.
– Activation energy has great influence on the rate at which a reaction occurs; increasing activation energies makes reactions slower.
d. Reaction Intermediates:
– In the complicated reactions, there may be reaction intermediates—intermediate species encountered during the period the reaction takes place but never appear in final products. They like to be local minima in energy profiles along a reaction coordinate.
iii. Energy Profile and Reaction Coordinate Diagram:
The reaction coordinate diagram indicates the relationship between the reaction coordinate (on the x-axis) and the potential energy of the system (on the y-axis). The overall characteristics are:
– Reactants: The beginning point, of some energy.
– Transition State: The top point, of the highest energy along the reaction pathway.
– Products: The end point, of some energy.
– Activation Energy: The difference in energy between the reactants and the transition state.
– Exothermic or Endothermic Reactions:
– When products are of lower energy than the reactants, the reaction is exothermic (gives out energy).
– When the products are of higher energy than the reactants, the reaction is endothermic (takes in energy).
2. Reaction Coordinate Diagram:
A reaction coordinate diagram is a plot of the potential energy changes in a system as the reaction proceeds from reactants to products. The diagram may plot significant features of a chemical reaction like activation energy, transition states, and the energy difference between reactants and products.
i. Energy Changes Along the Reaction Path:
– During a chemical reaction, the potential energy of the system varies. The diagram usually plots the energy (y-axis) against reaction progress (x-axis).
– The x-axis is the reaction coordinate, which is a definition of the reaction progress from reactants to products (usually represented as a sequence of steps or intermediates).
– The y-axis is the potential energy of the system, which varies as bonds are broken and formed.
ii. Activation Energy (Ea):
– Activation Energy (Ea) is the energy barrier that needs to be overcome for a reaction to occur.
– It is the energy difference between the reactants and the transition state (the highest point on the reaction pathway in terms of energy).
– High Activation Energy: Takes a higher energy to initiate the reaction, and the reaction rate will be slower.
– Low Activation Energy: Takes a lower energy to initiate the reaction, and the reaction rate will be quicker.
– On the graph, activation energy is the energy between the reactants and the top of the curve (transition state).
iii. Transition State (TS):
– The transition state is the peak in energy along the reaction path, or the most unstable state of the reaction.
– It is a transient atomic arrangement where bonds are forming and breaking and the molecule contains the greatest quantity of potential energy.
– It is also known as the activated complex.
– Vital Note: The transition state cannot be separated. It exists only at the peak of the energy barrier and proceeds rapidly to produce products or return to reactants.
iv. Reaction Intermediates:
– Reaction intermediates are entities that are created in the course of the reaction but not in the products. They are usually local minima in the reaction coordinate plot.
– They are not as unstable as the transition state but more so than the reactants or products.
– All reactions do not contain intermediates, but when they do, intermediates will tend to be stable, describable species that exist for only a short lifetime during the reaction.
v. Exothermic vs. Endothermic Reactions:
a. Exothermic Reaction:
– During an exothermic reaction, the energy content of the products is less than that of the reactants.
– That is, the system releases energy (e.g., heat) to the environment.
– The diagram would be that the products contain more energy than the reactants.
b. Endothermic Reaction:
– For an endothermic reaction, the products contain more energy than the reactants.
– That is, the system absorbs energy from the environment (say, heat).
– The graph would indicate that the products are more energetic than the reactants. 
Example of a Reaction Coordinate Diagram:
– For an Exothermic Reaction:
– Reactants begin at some energy level.
– The reaction proceeds onward to arrive at the transition state (peak point).
– Once it has gone through the transition state, the energy decreases to a lower value at the products, freeing up energy in the process.
– Activation Energy is the difference in energy between the reactants and the transition state.
– For an Endothermic Reaction:
– Reactants begin at some amount of energy.
– The reaction continues, and it reaches the transition state (most peaked).
– Following the transition state, the energy level is higher, and the products possess a greater energy level than the reactants.
– Activation Energy is constant, but energy is being used when the reaction is taking place.
In general, a reaction coordinate diagram is a graphical representation of the energy change and evolution of a reaction and illustrates important characteristics such as activation energy, the transition state, and whether or not the reaction is exothermic or endothermic.
3. Effect of Catalysts:
A catalyst is a material that speeds up the rate of a chemical reaction but is not used up in the process. It does this by offering a different path for the reaction to follow that has a lower activation energy than the reaction without a catalyst. This facilitates easier passage of the reaction and also boosts the rate of the reaction.
i. Decreasing the Activation Energy:
– Activation energy refers to the energy barrier that a reaction needs to cross in order to proceed. Catalysts reduce this energy barrier by offering a lower energy pathway for the reaction to which the reactants are able to transition to the transition state.
– The smaller the activation energy, the more molecules are able to move over the barrier, so the reaction goes faster.
– The catalysts do not change the total energy of the reactants and products; they shift the energy profile along the reaction path.
ii. Alternative Reaction Pathway:
– In the absence of a catalyst, the reaction follows a particular high-energy path with a single transition state, as depicted in a reaction coordinate diagram.
– When a catalyst is introduced, it forms a new reaction pathway with a lower energy peak (lower activation energy), leading to the creation of an intermediate or another transition state.
– This alternate route facilitates the reaction to be more easily allowed, thus enhancing the rate of the reaction while not changing the end products.
On a reaction coordinate diagram:
– The uncatalyzed pathway shows a peak with high form which is indicative of a high activation energy.
– The catalyzed pathway is lower peaked, showing that there is less activation energy necessary for the reaction to occur.
iii. Catalyst and Transition State
– Transition state is the high-energy, unstable transitional structure through which the reactants have to pass to get transformed into products.
– A catalyst can stabilize the transition state, reducing its energy and facilitating it to be formed.
– Even the catalyst briefly gets attached to the reactants so that it places them in a more favorable alignment such that reaction can take place with less energy.
iv. Catalysts Are Not Used Up in the Reaction:
– Another important characteristic of catalysts is that the catalyst itself does not become consumed while undergoing the process of a reaction. The catalyst will following a reaction have its original shape recovered and will be ready to continue with subsequent reactions.
– Catalysts are in contact with the reactants but do not themselves undergo chemical change in the long term. For instance, they might be engaged in intermediate steps or in transient complexes, but return to their initial state when the reaction ends.
v. Effect on Reaction Rate:
– By reducing the activation energy, catalysts increase the number of reactant molecules that are able to effectively collide with sufficient energy to produce products.
– Catalytic reactions, therefore, occur quicker under the same pressure and temperature as compared to the uncatalytic reaction.
– Catalysts prove particularly convenient when used in industries, whereby they form faster rates of reaction without demanding higher heat or pressure, conserving cost and energy.
vi. Types of Catalysts:
– Homogeneous Catalysts: These are those catalysts having the same phase as the reactants (liquid-phase catalysts for liquid-phase reactions, e.g.). They work by normally creating temporary intermediate species with the reactants.
– Example: Acid or base catalysts of organic reactions.
– Heterogeneous Catalysts: They are in a different phase than the reactants (e.g., solid catalysts for gaseous or liquid phase reactions). The reactants adsorb on the catalyst surface, and the reaction takes place on the catalyst surface.
– Example: Catalysts employed in industrial reactions such as the Haber process (solid iron catalyst) to produce ammonia.
vii. Catalysts of Biological Reactions (Enzymes):
– Enzymes are biological catalysts that accelerate biochemical reactions.
– Enzymes are also specific in their function and act by binding substrates, lowering the activation energy for the bio-reaction.
– The active site of the enzyme offers a unique environment where the reactants are arranged in the correct position and the reaction is accelerated more efficiently.
5.6.A.3 Energy Profile and Activation Energy:
1. Energy Profile Diagram:
An energy profile diagram shows energy change in a chemical reaction, illustrating the transformation from reactants to products.
– Reactants: Starting energy level.
– Products: Final energy level following the reaction.
– Activation Energy (Ea): Barrier energy to reach the transition state, illustrated by the difference between the reactants and the peak of the curve.
– Transition State: Location of maximum energy, where bonds are breaking and forming.
– ΔE (or ΔH): Difference in energy between the reactants and products. If products are of lower energy, the reaction is exothermic; if higher, endothermic.
Diagram Overview:
– The curve rises to the transition state and then drops to the level of the product, with activation energy and overall energy change noted on the graph.
Short, the diagram shows how energy ascends to the transition state and decreases (exothermic) or increases (endothermic) as the reaction proceeds to the products.
2.Energy (Ea):
Activation energy (Ea) is the least energy required to push reactants into the transition state and proceed with a chemical reaction. Activation energy is a barrier that has to be overcome for the reaction to proceed.
– Reaction Diagram: Reactants must achieve sufficient energy to reach the transition state (the peak), after which the reaction proceeds to yield products.
– Influence on Reaction Rate: Faster reactions due to high activation energy and slow reactions due to fewer molecules possessing enough energy. Slow reactions with low activation energy.
– Factors that Influence Ea:
– Temperature: When temperature is high, molecular energy is raised, resulting in quicker reactions.
– Catalysts: Reducing the activation energy leads to faster reactions.
– Concentration: Adding more reactants increases the likelihood of the reaction.
In short, activation energy is the “barrier energy” reactants need to overcome to transform into products.
3.Transition State:
– Energy Peak: It is the highest energy point along the reaction pathway, where bonds are being broken and formed to some degree.
– Unstable: The transition state is extremely unstable, and upon its formation, the system quickly proceeds to form products or revert back to reactants.
– Short-lived: It exists for an extremely brief period, either prior to the reaction continuing to form products or the molecules returning to their original state.
That is, it’s a “saddle point” in the reaction, where the system is neither fully reactant nor fully product but at its highest energy level in the course of conversion.
5.6.A.4 Temperature Dependence of Reaction Rate and the Arrhenius Equation:
1. Temperature and Reaction Rate:
Temperature is one of the factors that influence the rate of an elementary reaction by regulating the kinetic energy of the molecules. Here’s how it functions:
i. How Temperature Influences Reaction Rate:
a. Higher Molecular Kinetic Energy:
– As temperature increases, the average kinetic energy of the molecules is greater. This indicates that molecules travel faster and bump into other molecules more often.
b. More Energetic Collisions:
– In order for a reaction to take place, molecules need to collide with sufficient energy to surpass the activation energy barrier. There will be an increase in collisions with sufficient energy at higher temperatures to surpass this barrier, resulting in more successful reactions.
c. Increased Collision Frequency:
– Molecules not only collide with greater energy at higher temperatures, but also collide more frequently, further raising the chances of successful reactions.
ii. Summary:
– More energetic and frequent collisions = more temperature.
– It causes an enhancement of the rate of reaction as there are more molecules able to cross the activation energy, i.e., reactions occurring within the same period.
Rising the temperature will make the reaction occur more rapidly through more molecular collision energy and frequency.
2. Activation Energy (Ea):
Activation Energy (Ea) is the least amount of energy reactants must have to be converted to products in a chemical reaction. It is the energy barrier to be surpassed in order for the reaction to proceed.
i. Effect of Temperature on Activation Energy:
a. Temperature and Molecular Energy:
– As temperature rises, kinetic energy of molecules rises. That is, molecules travel faster and collide with other molecules forcefully.
b. More Molecules Cross Ea:
– As a result of higher kinetic energy at high temperatures, more molecules will have sufficient energy to cross the activation energy barrier (Ea). More molecules are thus able to proceed to the transition state and continue forming products.
c. Higher Rate of Reaction:
– Efficient collisions lead to a greater rate of reaction. Therefore, as the temperature rises, more molecules can overcome the Ea and speed up the reaction.
ii. Key Takeaway: Though Ea itself is independent of temperature, temperature influences the amount of molecules having sufficient energy to cross this barrier. Increased temperature implies increased speed of reactions as more molecules can attain or surpass the activation energy required for reaction.
3. Arrhenius Equation: The Arrhenius Equation illustrates how the rate constant ( k ) of a chemical reaction relates to the activation energy ( Ea ), temperature ( T ), and a constant ( A ), which is referred to as the pre-exponential factor or frequency factor.
The equation is expressed as:
Where:
– ( k ) = rate constant of the reaction
– ( A ) = pre-exponential factor (linked to the frequency of collisions and the orientation of reactants)
– ( Ea ) = activation energy
– ( R ) = universal gas constant 8.314J/mol.K
– ( T ) = temperature in Kelvin (K)
– ( e ) = base of the natural logarithm (approximately 2.718)
i. Key Points:
a. Rate Constant and Temperature: The rate constant ( k ) tends to increase with temperature since higher temperatures enable more molecules to possess sufficient energy to surpass the activation energy barrier.
ii. Effect of Activation Energy: A higher activation energy ( Ea) results in a lower rate constant at a specific temperature, causing the reaction to proceed more slowly. In contrast, a lower activation energy facilitates a quicker reaction.
iii. Exponential Dependence: The rate constant ( k ) is exponentially influenced by both the activation energy ( Ea ) and temperature ( T ). Even minor temperature variations can cause substantial changes in the rate constant.