Chemical Kinetics Quiz & Flashcards

Color does not affect reaction rate, while temperature, concentration, and catalysts do.

Increasing temperature generally increases reaction rates by providing more energy to reactant molecules.

Catalysts lower activation energy, making it easier for reactions to occur without being consumed.

In zero-order reactions, rate is constant and independent of reactant concentrations.

The rate law expresses the reaction rate as a function of reactant concentrations and a rate constant.

A mechanism is the step-by-step sequence of elementary reactions that make up the overall reaction.

High activation energy means fewer molecules have sufficient energy to react, leading to a slower rate.

The initial rates method determines reaction order by comparing initial rates at different concentrations.

A catalyst does not affect the equilibrium position; it only speeds up the rate of reaching equilibrium.

Increasing pressure raises gas concentration, leading to more frequent collisions and a higher reaction rate.

The rate constant 'k' is a measure of reaction speed, indicating how quickly reactants turn into products.

A bimolecular reaction involves two reactant molecules colliding, like A + B -> C.

The Arrhenius equation does not include reaction order; it relates temperature, activation energy, and frequency factor.

In first-order reactions, the rate is directly proportional to reactant concentration, so doubling it doubles the rate.

Catalysts are not consumed; they facilitate reactions without being used up.

Half-life is the time required for the concentration of a reactant to decrease to half its initial value.

Increasing surface area allows more collisions between reactants, thus increasing the reaction rate.

Intermediates are species formed and consumed in reaction steps but do not appear in the overall reaction equation.

Solvents can either stabilize intermediates or alter collision frequency, affecting the rate positively or negatively.

Collision theory states that molecules must collide with proper orientation and sufficient energy for a reaction to occur.

A catalyst increases the reaction rate by providing an alternative pathway with a lower activation energy.

In pseudo first-order reactions, one reactant is in excess, making the rate appear first-order with respect to another reactant.

The rate-determining step is the slowest step, controlling the overall reaction rate.

The pre-exponential factor represents the frequency of collisions and the orientation of reactant molecules.

In homogeneous catalysis, the catalyst and reactants are in the same phase, usually both in solution.

The steady-state approximation assumes the concentration of intermediates remains constant during most of the reaction.

Reaction rate is typically measured by observing changes in the concentration of reactants or products over time.

Molecularity refers to the number of molecules involved in an elementary reaction step, indicating its order.

Concentration is not part of the Arrhenius equation, which relates activation energy, temperature, and the pre-exponential factor.

Studying kinetics helps in understanding reaction mechanisms, designing reactors, and developing new catalysts and materials.

A catalyst is not consumed in the reaction; it lowers activation energy and provides an alternative pathway.

In first-order reactions, the rate is directly proportional to the reactant concentration, so it doubles with increased concentration.

In heterogeneous catalysis, the catalyst is in a different phase than the reactants, often as a solid with gaseous or liquid reactants.

Increasing the surface area of a solid reactant increases the reaction rate by allowing more collisions with reactant molecules.

Increasing reactant concentration generally increases the rate of a chemical reaction due to more frequent collisions.

Temperature affects the rate constant in the Arrhenius equation, impacting how quickly reactions occur.

The rate-determining step is the slowest step in a reaction mechanism, limiting the overall reaction rate.

Activation energy is the minimum energy that reactant molecules must possess to undergo a successful chemical reaction.

An inhibitor decreases the rate of a chemical reaction by interfering with the reactants or catalysts.

Catalysts lower activation energy, increasing reaction rates, while inhibitors decrease rates by interfering with the reaction.

In second-order reactions, the rate is proportional to the square of the reactant concentration, so doubling it quadruples the rate.

An elementary reaction is a single-step process with a specific rate and molecularity, forming the basis of reaction mechanisms.

A common misconception is that catalysts are consumed in reactions, but they actually remain unchanged after facilitating reactions.

A catalyst lowers the activation energy of a reaction, making it easier for reactants to form products.

A reaction coordinate diagram represents the energy changes during the course of a chemical reaction.

An intermediate is a transient species formed during a reaction mechanism, which does not appear in the overall balanced equation.

Increasing pressure raises the concentration of gases, leading to more frequent collisions and a higher reaction rate.

A catalyst lowers the activation energy, altering the energy profile by providing an alternative pathway for the reaction.