Glossary
Activated complex (transition state)
A short-lived, high-energy, unstable arrangement of atoms that exists momentarily at the peak of the activation energy barrier during a chemical reaction.
Example:
The activated complex is a fleeting intermediate structure that is neither reactant nor product, but rather a bridge between them.
Activation energy (Ea)
The minimum amount of energy that colliding reactant molecules must possess for a chemical reaction to occur.
Example:
Lighting a match requires overcoming a certain activation energy to initiate the combustion reaction.
Arrhenius equation
An equation that relates the rate constant (k) of a reaction to the absolute temperature, activation energy, and a pre-exponential factor.
Example:
The Arrhenius equation can be used to predict how much faster a reaction will proceed if the temperature is increased by 10 degrees Celsius.
Average rate
The change in concentration of a reactant or product over a specific time interval.
Example:
Measuring the decrease in reactant concentration from the start of a reaction to 5 minutes later gives the average rate over that period.
Catalyst
A substance that increases the rate of a chemical reaction without being consumed in the process, typically by providing an alternative reaction pathway with a lower activation energy.
Example:
Enzymes in biological systems act as catalysts, speeding up vital biochemical reactions without being used up.
Collision model
A theory that explains reaction rates based on the idea that reactant particles must collide with sufficient energy and correct orientation to form products.
Example:
The collision model helps explain why increasing temperature usually speeds up reactions: more energetic collisions occur.
Differential rate law
An equation that describes how the rate of a reaction depends on the instantaneous concentrations of reactants.
Example:
The differential rate law for A → products might be Rate = k[A]², showing how the rate changes as [A] changes.
Effective collision
A collision between reactant molecules that possesses both sufficient kinetic energy (greater than or equal to the activation energy) and the correct spatial orientation to lead to product formation.
Example:
Only a small fraction of all molecular collisions are effective collisions because most lack the necessary energy or proper alignment.
Elementary reaction
A reaction that occurs in a single step exactly as written, where the rate law can be directly determined from its stoichiometry.
Example:
The collision of two molecules, A + B → C, is an elementary reaction if it happens in one direct step.
Elementary steps
The individual, single-step reactions that make up an overall reaction mechanism.
Example:
A complex reaction might proceed through three distinct elementary steps, each involving specific molecular collisions.
First order
A reaction order where the rate of reaction is directly proportional to the concentration of a specific reactant.
Example:
Radioactive decay is a classic example of a first order process, where the decay rate depends only on the amount of radioactive material present.
Half-life (t₁/₂)
The time required for the concentration of a reactant to decrease to half of its initial value.
Example:
If a medication has a half-life of 4 hours, then after 4 hours, half of the initial dose will have been eliminated from the body.
Instantaneous rate
The rate of a reaction at a particular moment in time, determined by the slope of the tangent line to the concentration vs. time curve at that point.
Example:
To find out how fast a reaction is proceeding exactly 30 seconds after it begins, you would calculate its instantaneous rate.
Integrated rate law
An equation that relates the concentration of a reactant to time, allowing for the prediction of concentration at any given moment or the time required to reach a certain concentration.
Example:
Using the integrated rate law for a first-order reaction, you can calculate how much reactant remains after 10 minutes if you know the initial concentration and the rate constant.
Intermediate
A species that is formed in one elementary step of a reaction mechanism and consumed in a subsequent step, thus not appearing in the overall balanced chemical equation.
Example:
In the decomposition of hydrogen peroxide, the hydroxyl radical (OH•) can act as an intermediate.
Kinetics
The branch of chemistry that studies the rates of chemical reactions and the factors that influence them.
Example:
Understanding chemical kinetics helps pharmaceutical companies determine how quickly a drug will degrade over time.
Multistep reactions
Chemical reactions that occur through a sequence of two or more elementary steps rather than in a single step.
Example:
Most complex organic synthesis reactions are multistep reactions, involving several distinct transformations to reach the final product.
Rate constant (k)
A proportionality constant in the rate law that relates the rate of a reaction to the concentrations of reactants, specific to a given reaction and temperature.
Example:
A large value for the rate constant (k) indicates a fast reaction, while a small value suggests a slow reaction at that temperature.
Rate law
An experimentally determined equation that expresses the relationship between the rate of a reaction and the concentrations of its reactants.
Example:
For the decomposition of N₂O₅, the rate law is Rate = k[N₂O₅], indicating it's first order with respect to N₂O₅.
Rate of a reaction
The speed at which reactants are converted into products, typically measured as the change in concentration of a reactant or product per unit time.
Example:
If a reaction consumes 0.1 M of reactant A every 10 seconds, its rate of reaction is 0.01 M/s.
Rate-limiting step (Rate-determining step)
The slowest elementary step in a reaction mechanism, which determines the overall rate of the entire reaction.
Example:
In a multi-step manufacturing process, the rate-limiting step is the bottleneck that dictates the maximum production speed.
Reaction energy profile
A diagram that illustrates the energy changes that occur during a chemical reaction, showing the energy of reactants, products, and the transition state.
Example:
A chemist might sketch a reaction energy profile to visualize the activation energy barrier that must be overcome for a reaction to proceed.
Reaction mechanism
A series of elementary steps that describe the actual pathway by which a chemical reaction occurs at the molecular level.
Example:
The overall reaction of ozone depletion involves a complex reaction mechanism with several individual steps.
Reaction order
The exponent to which a reactant's concentration is raised in the rate law, indicating how the rate is affected by changes in that reactant's concentration.
Example:
If doubling a reactant's concentration quadruples the reaction rate, the reaction order with respect to that reactant is two.
Second order
A reaction order where the rate of reaction is proportional to the square of the concentration of a specific reactant, or to the product of the concentrations of two reactants.
Example:
Many dimerization reactions, where two identical molecules combine, often exhibit second order kinetics.
Steady-state approximation
A simplifying assumption in reaction kinetics that states the concentration of a reaction intermediate remains constant over time because its rate of formation equals its rate of consumption.
Example:
When an intermediate is highly reactive and short-lived, the steady-state approximation can be applied to simplify the derivation of the rate law.
Stoichiometric coefficients
The numbers placed in front of reactants and products in a balanced chemical equation, indicating the relative number of moles involved in the reaction.
Example:
In 2H₂ + O₂ → 2H₂O, the stoichiometric coefficients are 2 for H₂, 1 for O₂, and 2 for H₂O.
Zero order
A reaction order where the rate of reaction is independent of the concentration of a specific reactant.
Example:
In a zero order reaction, adding more of a reactant will not speed up the reaction, similar to how adding more people to a fully staffed assembly line won't increase output.