Kinetics and Collision Theory Flashcards

A set of vocabulary flashcards covering chemical kinetics, Maxwell-Boltzmann distributions, activation energy, and factors affecting reaction rates based on lecture notes.

Activation energy ($E_A$)

The minimum energy which particles need to collide to start a reaction.

Maxwell-Boltzmann distribution

A graphical representation showing the spread of energies that molecules of a gas or liquid have at a particular temperature.

$$E_{mp}$$

The most probable energy of the molecules in a distribution, located at the peak of the Maxwell-Boltzmann curve.

Mean energy

The average energy of all particles in a distribution, which is located to the right of the peak ($E_{mp}$) on a Maxwell-Boltzmann curve.

Rate of reaction

The change in concentration of a substance in unit time, with the usual unit being $mol\,dm^{-3}\,s^{-1}$.

Initial rate

The rate at the start of the reaction, which is the point where the reaction is fastest.

Effective collisions

Collisions that occur with sufficient energy (at least the activation energy) to result in a chemical reaction.

Concentration and Pressure effect

An increase in these factors leads to more particles per unit volume, causing a higher frequency of effective collisions.

Catalyst

A substance that increases reaction rates without getting used up by providing an alternative route or mechanism with a lower activation energy ($E_A$).

Sodium thiosulfate and hydrochloric acid experiment rate

An approximation for rate of reaction measured as $\frac{1}{t}$, where $t$ is the time taken for a cross to disappear due to sulfur cloudiness.

Surface area effect

Increasing this causes successful collisions to occur more frequently between reactant particles, increasing the reaction rate.

Collision theory

A theory stating that reactions can only occur when collisions take place between particles having sufficient energy to break relevant bonds.

Temperature effect on reaction rate

As temperature increases, particles gain kinetic energy, leading to more frequent and effective collisions, thereby increasing reaction rate.\n\n

Catalytic mechanism

The process by which a catalyst lowers the activation energy of a reaction, allowing it to proceed at a faster rate without being consumed.\n\n

Reaction order

An exponent in the rate law that indicates the dependency of the rate on the concentration of a reactant.\n\n

Dynamic equilibrium

A state in a reversible reaction where the rate of the forward reaction equals the rate of the reverse reaction, resulting in constant concentrations of reactants and products.\n\n

Enzyme catalysis

Biological catalysts that speed up reactions by lowering the activation energy, often highly specific to substrates.\n\n

Arrhenius equation

An equation that gives the dependence of reaction rates on temperature and activation energy: $k = A e^{\frac{-E_A}{RT}}$, where $k$ is rate constant, $A$ is the pre-exponential factor, $E_A$ is the activation energy, and $R$ is the gas constant.\n\n

Collision geometry

The orientation of colliding particles, which affects whether their collision will be effective in leading to a reaction.\n\n

Homogeneous catalyst

A catalyst that is in the same phase as the reactants, typically in a solution, enhancing reaction rates by uniformly distributing throughout the reactants.\n\n