Introduction to Kinetics & Equilibria

Flashcards covering key concepts in kinetics and equilibria from the Edexcel AS Chemistry syllabus.

Collision Theory

The theory that describes how the rate of a chemical reaction depends on particle collisions.

Activation Energy (E_a)

The minimum energy required for a chemical reaction to occur.

Collision Frequency

The number of collisions between reactant particles per unit time.

Effective Collision

A collision that leads to a chemical reaction, occurring with proper orientation and sufficient energy.

Ineffective Collision

A collision that does not result in a chemical reaction due to incorrect orientation or insufficient energy.

Maxwell-Boltzmann Distribution

A graph showing the distribution of energy levels among particles in a system.

Dynamic Equilibrium

A state in which the forward and reverse reaction rates are equal, and concentrations of reactants and products remain constant.

Le Chatelier's Principle

The principle stating that if a change is made to a system at equilibrium, the system will adjust to counteract that change.

Catalyst

A substance that increases the rate of a chemical reaction without being consumed, by lowering the activation energy.

Homogeneous Catalyst

A catalyst that is in the same phase as the reactants.

Heterogeneous Catalyst

A catalyst that is in a different phase than the reactants.

Exothermic Reaction

A reaction that releases energy in the form of heat.

Endothermic Reaction

A reaction that absorbs energy from its surroundings.

Rate of Reaction

The speed at which a chemical reaction occurs, often measured in mol dm^-3 s^-1.

Concentration Change Method

A method of measuring reaction rates by observing changes in concentration of reactants or products.

Mass Loss Method

A technique for measuring reaction rates using the change in mass over time due to gas production.

Gas Volume Method

Measuring the volume of gas produced in a reaction over time to determine the rate of reaction.

Titration

A laboratory method used to determine the concentration of a reactant by adding a reagent of known concentration until the reaction is complete.

Quenching

A technique used to stop a reaction at a specific time for analysis.

Reaction Profile

A graph showing the energy changes during a chemical reaction.

Catalytic Reaction Pathway

An alternative route for a reaction provided by a catalyst, which has a lower activation energy.

Equilibrium Constant (K)

A value that expresses the relationship between the concentrations of products and reactants at equilibrium.

Forward Reaction Rate

The rate at which reactants convert to products in a chemical reaction.

Backward Reaction Rate

The rate at which products convert back to reactants in a reversible reaction.

Closed System

A system where no reactants or products can escape, allowing equilibrium to be established.

Open System

A system where matter and energy can be exchanged with the surroundings.

Reaction Completion

A stage in a reaction where all reactants are converted into products.

Equilibrium Shift

A change in the concentration of reactants or products in response to a change in conditions.

Concentration Increase Shift

The direction equilibrium shifts when the concentration of reactants is increased.

Pressure Increase Shift

The direction equilibrium shifts when the pressure of a gaseous system is increased.

Temperature Increase Shift

The direction equilibrium shifts when the temperature is increased, favoring endothermic reactions.

Enthalpy Change (ΔH)

The change in heat content during a reaction, indicating whether it is exothermic or endothermic.

Slope of Rate Graph

The gradient of a plot of concentration versus time, indicating the rate of reaction.

Surface Area Increase

A method to increase the rate of reaction by increasing the surface area of solid reactants.

Temperature Effect on Kinetics

Higher temperatures increase the kinetic energy of particles, leading to more frequent and successful collisions.

Ethanol and Ethanoic Acid Reaction

An example of an equilibrium reaction used to demonstrate Le Chatelier's principle.

Dynamic vs Static Equilibrium

Dynamic equilibrium involves constant motion of reactants and products, unlike static equilibrium, where no changes occur.

Gas Collection Method

A technique for measuring gas produced in a reaction by capturing it in a graduated container.

Industrial Reaction Conditions

The specific conditions (temperature, pressure, concentration) under which industrial chemical reactions are optimized.

Reversible Reaction

A reaction where products can revert back to reactants.

Reactant Concentration

The amount of reactants present in a reaction mixture, affecting the rate and position of equilibrium.

Tolerance in Measurement

The acceptable range of values when measuring concentrations or rates.

Phase of Catalyst

Refers to whether a catalyst is in the same (homogeneous) or different (heterogeneous) phase as the reactants.

Catalyst Efficiency

The effectiveness of a catalyst in increasing reaction rates.

Zeolites

Microporous materials used as catalysts in various chemical processes.

Exothermic Activation Energy

Lower activation energy needed for an exothermic reaction compared to the reverse endothermic reaction.

Endothermic Activation Energy

Higher activation energy required for an endothermic reaction compared to exothermic.

Disappearing Cross Experiment

A method to demonstrate the rate of reaction by observing a color change.

Ammonia Yield in Haber Process

The overall amount of ammonia produced, influenced by temperature and pressure.

Financial Considerations in Industry

The balance of costs and profits in determining reaction conditions for industrial processes.