Dr. Floyd Beckford Lyon College CHM 120 CHAPTER 16 KINETICS: Rates and Mechanisms of Chemical Reactions Dr. Floyd Beckford Lyon College

CHEMICAL KINETICS Two factors control the outcome of chemical reactions: Chemical Thermodynamics Chemical Kinetics Chemical Kinetics: study of rates of chemical reactions and mechanisms by which they occur A reaction may be spontaneous but does not occur at measurable rates

REACTION RATES Rate of reaction describes how fast reactants are used up and products are formed There are 4 basic factors that affect reaction rates Concentration Physical state Temperature Catalysts

For every reaction the particles must come into intimate contact with each other High concentrations by definition implies that particles are closer together (than dilute solutions) So rate increases with concentration The degree of intimacy of particles obviously depends on the physical nature of the particles Particles in the liquid state are closer than in the solid state

Likewise, particles in a finely divided solid will be closer than in a chunk of the solid In both situations, there is a larger surface area available for the reaction to take place This leads to an increase in rate Temperature affects rate by affecting the number and energy of collisions So an increase in temperature will have the effect of increasing reaction rate

Rate of reaction is typically measured as the change in concentration with time This change may be a decrease or an increase Likewise the concentration change may be of reactants or products

Rate has units of moles per liter per unit time - M s-1, M h-1 Consider the hypothetical reaction aA + bB  cC + dD We can write

Note the use of the negative sign - rate is defined as a positive quantity - rate of disappearance of a reactant is negative 2N2O5(g)  4NO2(g) + O2(g)

Rate may be expressed in three main ways: Average reaction rate: a measure of the change in concentration with time 2. Instantaneous rate: rate of change of concentration at any particular instant during the reaction 3. Initial rate: instantaneous rate at t = 0 - that is, when the reactants are first mixed

RATE LAW where k is called the rate constant Consider the following reaction aA + bB  products Rate of reaction changes as concentration of reactants change at constant temperature RATE LAW: equation describing the relationship between concentration of a reactant and the rate Rate = k[A]m[B]n where k is called the rate constant

m, n are called reaction orders - they indicate the sensitivity of the rate to concentration changes of each reactant NOTE: the orders have nothing to do with the stoichiometric coefficients in the balanced overall equation An exponent of 0 means the reaction is zero order in that reactant - rate does not depend on the concentration of that reactant

An exponent of 1  rate is directly proportional to the concentration of that reactant - if concentration is doubled, rate doubles - reaction is first order in that reactant An exponent of 2  rate is quadrupled if the concentration of that reactant is doubled - reaction is second order in that reactant The overall reaction order is the sum of all the orders

Rate = k[A][B]0 m = 1 and n = 0 - reaction is first order in A and zero order in B - overall order = 1 + 0 = 1 - usually written: Rate = k[A] Remember: the values of the reaction orders must be determined from experiment; they cannot be found by looking at the equation

DETERMINATION OF THE RATE LAW The method of initial rates may be used - involves measuring the initial rates as a function of the initial concentrations - avoids problems of reversible reactions - initially there are no products so they cannot affect the measured rate In this method the experiments are chosen so as to check the effect of a single reactant on the rate

THE RATE CONSTANT 1. The units of k depends on the overall order of reaction 2. The value of k is independent of concentration and time 3. The value refers to a specific temperature and changes if we change temperature 4. Its value is for a specific reaction

THE INTEGRATED RATE EQUATION This is the equation that relates concentration and time Consider a first-order reaction aA  products Rate = k[A]

The equation may be written in the form for a linear plot A plot of log [A]t vs. t is linear plot with slope = -k/2.303 Note that this plot gives a straight line ONLY if the reaction is first-order

Half-life The half-life, t1/2, is defined as the time it takes for the reactant concentration to drop to half its initial value Note: the half-life for a first order reaction does not depend on the initial concentration The value of the half-life is constant

Second order reactions Consider a reaction that is 2nd order in reactant A and 2nd overall aA  products and Rate = k[A]2 A plot of 1/[A]t vs. t gives a straight line with slope = k

RATES AND TEMPERATURE Recall that temperature is the only factor that affects the rate constant In general rates increase with temperature

This is ARRHENIUS’ EQUATION Can be arranged in the form of a straight line ln k = (-Ea/R)(1/T) + ln A Plot ln k vs. 1/T  slope = -Ea/R

Another form of Arrhenius’ equation: COLLISION THEORY: a reaction results when reactant molecules, which are properly oriented and have the appropriate energy, collide The necessary energy is the activation energy, Ea

Not all collisions leads to a reaction For effective collisions proper orientation of the molecules must be possible

TRANSITION STATE THEORY During a chemical reaction, reactants do not suddenly convert to products The formation of products is a continuous process of bonding breaking and forming At some point, a transitional species is formed containing “partial” bonds This species is called the transition state or activated complex

The transition state is the configuration of atoms at the maximum of the reaction energy diagram The activation energy is therefore the energy needed to reach the transition state Note also that the transition state can go on to form products or break apart to reform the reactants

REACTION MECHANISMS MECHANISM: the step-by-step pathway by which a reaction occurs Each step is called an elementary step NO2(g) + CO(g)  NO(g) + CO2(g) Mechanism: NO2(g) + NO2(g)  NO(g) + NO3(g) NO3(g) + CO(g)  NO2(g) + CO2(g) NO3 is a reaction intermediate

Elementary reactions are classified by the molecularity A  B + C unimolecular A + B  C + D bimolecular A + 2B  E termolecular Termolecular reactions are very unlikely For ANY SINGLE ELEMENTARY REACTION – reaction orders are equal to the coefficients for that step

Rate = kelem[A][B] The slow step is called the rate-determining step (RDS) A reaction can never occur faster than its slowest step Overall reaction = sum of all elementary steps 2. The mechanism proposed must be consistent with the rate law

There may be more than one plausible mechanism The experimentally determined reaction orders indicate the number of molecules of the reactants - in the RDS (if it occurs first) - the RDS and any fast steps before it

CATALYSIS Reaction rates are also affected by catalysts Catalyst: a substance that increases the rate of a reaction without being consumed in the reaction Catalysts work by providing alternative pathways that have lower activation energies A catalyst may be homogeneous or heterogeneous Homogeneous: catalyst and reactants are in the same phase

Heterogeneous: catalyst in a different phase Typically: a solid in a liquid An important example: catalytic converters in automobile - convert pollutants to CO2 H2O, O2, N2 - usually Pt, Pd, V2O5, Cr2O3, CuO Cars must use unleaded fuels – lead poisons the catalytic bed