Chemical Kinetics Quiz

Chemical Kinetics

• Chemical kinetics studies reaction rates and mechanisms.
• Rate is the change in a quantity over a given time.
• Reaction rate is measured by the decrease in reactant concentration or the increase in product concentration over time.
• A negative sign is used for reactant concentration measurements to show a decrease.
• Reaction rate generally slows down as reactants are consumed.
• Reactions stop when reactants are used or equilibrium is reached.
• Rate data can be collected and used to determine the order of the reaction.
• The order of a reaction is the sum of the exponents for each reactant in the rate law.

Defining Rate

• Rate is the change in a quantity per unit of time.
• Driving a car at 60 mph is an example of rate, representing the distance it travels in one hour.
• Speed = change in distance / change in time

Defining Reaction Rate

• Reaction rate is how much reactant concentration decreases or product concentration increases per unit of time.
• For reactants, a negative sign is used.
• Rate = change in concentration / change in time
• Rate = Δ[product]/Δt or Rate = -Δ[reactant]/Δt

Reaction Rate Changes Over Time

• As time goes on, the rate of a reaction generally slows because reactant concentrations decrease.
• The reaction eventually stops because all reactants are used up or equilibrium is reached.

Hypothetical Reaction

• In a reaction between Red and Blue molecules, the total number of molecules always remains 100.
• Reaction rate can be measured as the speed of loss of Red molecules or the speed of gain of Blue molecules over time.

Reaction Rate and Stoichiometry

• In many reactions, the coefficients in the balanced chemical equation are not all the same.
• Because of this, the change in concentration of one substance is a multiple of the change in concentration of another.

Measuring Reaction Rate

• In measuring reaction rates, the concentration of at least one component in the mixture must be determined at various points in time.
• For reactions that are completed within one hour, continuous monitoring of concentration may be used.
• For reactions or processes lasting over long periods of time, the mixture may be sampled at various time intervals.

Continuous Monitoring

• Polarimetry: To measure the change in the orientation angle of a plane-polarized light, when measuring the concentration of components over time.
• Spectrophotometry: To measure the fraction of light of a particular wavelength absorbed by one component of the reaction mixture over time.
• Total pressure: in reactions involving gas mixture, the total pressure is dependent on the partial pressures of each gaseous component in the chemical reaction; all the component’s partial pressures depend on their respective stoichiometric coefficients

Sampling

• Gas Chromatography: Used to determine the different concentrations components of a mixture.
• Methods of sampling include taking aliquots of the reaction mixture at fixed time intervals, doing quantitative analysis, or determining the concentration of one component using titration or gravimetric analysis.

Methods of Measuring Reaction Rates

• Physical measurements—continuous measurements, initial rate measurements (clock reactions).
• Chemical measurements—titration.

Continuous Measurements

• The experiment is done in one take.
• The reaction rates are determined by continuously measuring a property directly proportional to the reaction mixture component concentration.

Advantages of Physical Measurements

• Suitable for fast reactions.
• Small sample size.
• More accurate than chemical method (titration).
• No interruption—continuous measurements.
• Can be automated.

Disadvantages of Physical Measurements

• More sophisticated.
• More expensive.
• More specific—only suit a limited number of reactions.

Quenching Methods

• Cooling the reaction mixture rapidly in ice.
• Introducing a cold solvent to dilute the reaction mixture.
• Removing one of the reactants or the catalyst by adding another substance.
• Example, if a reaction uses a catalyst, that catalyst can be removed by adding a reagent that will react with it.

Chemical Measurements (Titration Methods)

• Start a reaction with all conditions fixed except one.
• Withdraw and quench fixed amounts of the reaction mixture at different time intervals.
• Using titration to determine the concentration of one of the reactants or products.

Factors Affecting Reaction Rates

• Nature of reactants: reactivity of molecules varies greatly based on the types of particles involved, and the physical state (gas, liquid, solid); smaller molecules tend to react faster than larger ones. The more surface area available for interaction, the higher the reaction rate; powdered solids generally react faster than large, solid blocks.
• Temperature: increasing temperature speeds up reactions; a rule of thumb is that reaction rates double for each 10°C rise in temperature.
• Catalysts: catalysts speed up reactions by providing an alternative pathway with a lower activation energy. Catalysts are not consumed in the reaction.
• Concentration: increasing the concentration of reactants generally speeds up the reaction. The frequency of collisions between reacting particles is proportional to the concentration.

Rate Law

• The rate law of a reaction shows the relationship between the rate of the reaction and the reactant concentrations.
• It has the form: Rate = k[A]ⁿ[B]ᵐ where k is the rate constant, and n and m are the orders of reactions for reactants A and B, respectively.

Reaction Rate Order

• The sum of the exponents in the rate law equation is called the overall order of the reaction.

Methods of Determining the Rate Law

• Plot [A] versus time. If the plot is a straight line, the reaction is zero order. The slope of the line will equal "-k."
• Plot ln[A] versus time. If the plot is a straight line, the reaction is first order. The slope of the line will equal "-k."
• Plot 1/[A] versus time. If the plot is a straight line, the reaction is second order. The slope of the line will equal "k."

Zero-Order Reactions

• The rate of a reaction is constant (independent of the reactant concentration).
• Rate = k
• [A] = -kt + [A]₀
• The graph of [A] versus time is a straight line. The slope equals "-k" and the y-intercept equals [A]₀

First-Order Reactions

• The rate is directly proportional to the concentration of the reactant.
• Rate = k[A]
• ln[A] = -kt + ln[A]₀
• The graph of ln[A] versus time is a straight line. The slope equals "-k," and the y-intercept equals ln[A]₀

Half-Life

• The time it takes for the concentration of a reactant to decrease to one-half of its initial value.
• t₁⁄₂ =0.693/k (for first-order reactions)

Second-Order Reactions

• The rate is proportional to the square of the concentration of a reactant.
• 1/[A] = kt + 1/[A]₀
• The graph of 1/[A] versus time is a straight line. The slope equals "k."

Activation Energy

• The minimum amount of energy needed for a reaction to occur.
• The higher the activation energy, the slower the reaction.
• Exothermic reactions release energy that may be used to compensate for the activation energy to start the reaction.

Arrhenius Equation

• A mathematical relationship between the rate constant, activation energy, frequency factor, and temperature.
• In the form: k = Ae⁻Ea/RT

Collision Theory

• A theory that explains reaction rates using collision theory.
• The collision frequency is the number of collisions that occur per unit of time.
• Effective collisions: reacting particles must have enough energy and correct orientation.

Effective Collisions

• Collisions in which reacting molecules have sufficient energy to overcome the activation energy barrier and proper orientation for bond changes.
• The higher the effective collisions, the faster the reaction.

Orientation Factor

• The likelihood that colliding molecules have the correct orientation required for the bond changes to take place.

Transition State Theory

• A theory that focuses on what happens to the reaction mixture after the colliding molecules come together.
• Explains why specific reaction pathways take place.
• It can be used to determine the activation energy and pre-exponential factor (A).

Reaction Mechanisms

• The series of elementary steps that make up a complex reaction.
• Elementary steps are reactions that cannot be simplified into smaller steps.
• The steps occur in a specific order, and one or more steps determine the reaction.

Intermediates

• Materials made as a product of one step, but then utilized as a reactant in a later step.

###Molecularity

• Unimolecular steps involve one reactant.

• Bimolecular elementary steps involve two reactants.

• Termolecular steps, involving three reactants, are rare.

Rate Laws for Elementary Steps

• The rate law of an elementary reaction can be determined from the reaction's stoichiometry (the number of reacting particles).

Rate-Determining Step

• In multi-step reaction mechanisms, the slowest step determines the overall speed of the reaction.
• The slowest step in a mechanism often has the highest activation energy.
• The rate law for the rate-determining step dictates the rate law for the overall reaction.