AP Chemistry 5.6-5.10: Energy Profiles and Multistep Mechanisms

Activated complex (transition state)

The activated complex, or transition state, is a short-lived, high-energy arrangement formed when reactants collide with suitable orientation and sufficient energy. It can proceed toward products or revert to reactants.

Activation energy, $E_a$

Activation energy is the minimum energy required for a collision to produce products. On a reaction-energy diagram, it is the energy difference between the reactants and the transition state.

How does the size of the activation energy relative to molecular kinetic energies affect reaction rate?

If $E_a$ is large compared with typical molecular kinetic energies, only a small fraction of collisions can react and the reaction is slow. A smaller $E_a$ allows a larger fraction of collisions to react, increasing the rate.

How are activation energy and reaction enthalpy represented on a reaction-energy diagram?

$E_a$ is the energy difference from the reactants to the transition state, while $\Delta H$ is the energy difference between products and reactants. These quantities describe different energy changes.

How can an exothermic reaction still have a substantial activation energy?

A reaction can form products at a lower energy than the reactants, giving $\Delta H<0$, while still requiring energy to reach the higher-energy transition state. Thus, exothermicity does not imply a small activation energy.

Reaction mechanism

A reaction mechanism is the sequence of elementary reactions, or individual steps, by which an overall reaction occurs. The mechanism provides molecular details that are not shown by the net chemical equation.

Elementary reaction

An elementary reaction is a single mechanistic step that represents a specific molecular event, including the bonds broken and formed in that step.

How is the net reaction obtained from a proposed reaction mechanism?

Add the elementary-step equations and cancel species that are produced in one step and consumed in another. The remaining reactants and products form the balanced net reaction.

Intermediate

An intermediate is a species formed in one elementary step and consumed in a later step. It appears in the mechanism but cancels out and does not appear in the net reaction.

Rate-limiting step

The rate-limiting, or rate-determining, step is the slowest step in a multistep mechanism. Because the overall process cannot proceed faster than this step, it often determines the observed rate law.

Why must a proposed reaction mechanism agree with both the net reaction and the experimental rate law?

Adding the steps must produce the observed balanced overall equation, and the mechanism must predict the experimentally observed concentration dependence. Agreement with only the net equation is insufficient because many different mechanisms can have the same net reaction.

What does the vertical axis of a reaction-energy profile represent?

The vertical axis represents the potential energy or enthalpy of the reacting system. It allows the relative energies of reactants, products, transition states, and intermediates to be compared.

What does the horizontal axis of a reaction-energy profile represent?

The horizontal axis represents reaction progress or the sequence of structural changes as reactants are converted into products. It is not a time axis.

How are transition states and intermediates identified on an energy profile?

Transition states are located at the peaks, or local maxima, of the energy profile. Intermediates are located at the valleys, or local minima, between two transition states.

How many transition states and intermediates are shown for a mechanism with several elementary steps?

A profile generally contains one transition state for each elementary step, shown as a peak. It contains one intermediate for each species formed and consumed between steps, shown as a valley between peaks.

How is the activation energy for an individual step determined from a multistep energy profile?

For each step, measure the energy difference between the preceding minimum and the next peak. For the first step, this is the difference between the reactant energy and the first transition-state energy.

How is the overall enthalpy change determined from a multistep energy profile?

Compare the energy of the final products with the energy of the initial reactants: $\Delta H=E_{\text{products}}-E_{\text{reactants}}$. Intermediate energies and transition-state energies do not determine the overall $\Delta H$.

How can the overall reaction be exothermic even if some steps in its mechanism are endothermic?

The overall enthalpy change depends only on the energy difference between the initial reactants and final products. Individual intermediates may be higher or lower in energy, so some steps can have positive $\Delta H$ even when the overall reaction has $\Delta H<0$.

How is the rate-determining step identified on a multistep energy profile?

The rate-determining step is usually the step with the largest activation-energy barrier, measured from its preceding reactant or intermediate minimum to the following peak. The highest peak alone is not always sufficient; the height relative to the preceding minimum matters.

How does a transition state differ from an intermediate on an energy diagram?

A transition state is the high-energy configuration at the top of a barrier and does not represent a stable species. An intermediate is a relatively lower-energy species between steps, represented by a valley, and may sometimes be detected.

How does a catalyst change a reaction-energy profile?

A catalyst provides an alternate mechanism with a lower activation-energy barrier, often lowering the barriers for both the forward and reverse reactions. It does not change the energies of the initial reactants or final products, so it does not change the overall $\Delta H$.