chemical kinetics

the area of chemistry concerned with the speeds, or rates, at which a chemical reaction occursthe speed of a reaction is not neccessarily related to the energy change

reaction rate

the change in the concentration of a reactant or a product with timeunits of M/salways highest at the beginning of a reaction2A --> Brate = - 1/2 ^[A] / ^t = ^[B] / ^ t (rate expression)

activation energy

Ea - the minimum amount of energy required to initiate a chemical reaction-must have a collision in order to have a reaction -collisions must occur with enough kinetic energy to break the existing bondsif energy is too small the molecules will merely bounce off each other intact

rate law

expresses the realtionship of the rate of a reaction to the rate constant and the concentrations of the reactants raised to some powersfor rxn aA +bB --> cCthe rate law is: rate = k[A]^x[B]^yx and y must be experimentally determinedsum of the exponents gives "order"of reaction

rate constant

k rate =k[A]every time a reaction takes place, at a constant temperature, the same value for k is obtained

first order reactions

rate = k[A] = - ^[A] / ^ tIn [A]0 / [A]t = kt[A] at time = 0, [A] at time = tfor gas-phase reactions we can replace the concentration terms with the pressures of the gaseous reactant

half-life

t 1/2 - the time required for the concentration of a reactant to decrease to half of its initial concentrationat t 1/2, [A] = 1/2[A]0t 1/2 = 0.693/kin 1st order reaction it is independant of the [A]

second-order reactions

A --> productsrate = - ^[A] / ^ t = k[A]2 (squared)1/[A]t = kt + 1/[A]0half life = 1 / k[A]0initial concentration does matter in 2nd order rxns

zero order reactions

rate = k[A]0 rate = khalf life = [A]0 / 2k

collision theory of chemical kinetics

molecules must collide in order to reactrate increases with the # of collisions per second

activated complex

aka - transition statea temporary species formed by the reactant molecules as a result of the collision before they form the product

effect temperature has on reaction rate

more molecules have the kinetic energy necessary to overcome the acivation energymolecules move faster so they have more collisionscollisions are more energetic

The Arrhenius Equation

k = Ae ^(-Ea / RT) OR In k = In A - Ea / RTEa = activation energy in kJ/molAe = frequency factor (we usually don't know)

Arrhenius Equation with two temperatures and respective constants

In k1/k2 = Ea / R ( 1/ T2 - 1/ T1)for any order reaction

elementary steps

a series of simple reactions that represent the progress of the overall reaction on the molecular level

reaction mechanism

the sequence of elementary steps that leads to product formationdetails how the reaction is thought to take placereactions are hypotheses (always subject to revision)

intermediates

appear in the mechanism of the reaction (elementary steps) but not in the overall balanced equationthey are canceled as you add the elementary steps

rate-determining step

the slowest step in the sequence of steps leading to product formation

unimolecular reaction

an elementary step in which only one reacting molecule participates

bimolecular reaction

an elementary step involving two molecules

termolecular reaction

an elementary step involving 3 moleculesthey are very rarerequire 3 molecules to "slam" into each other at once

molecularity of a reaction

the number of molecules reacting in an elementary step

catalyst

a substance that increases the rate of a chemical reaction without itself being consumedreacts within an elementary step of the mechanism but is regenerated in a subsequent stepit works by allowing an alternative route which lowers the activation energy

heterogeneous catalyst

catalytic converter - gas vs solid

homogeneous catalyst

same phase

enzyme catalysis

biological catalyst

catalyst vs intermediate

catalyst appears as a reactant and later gets spit back outintermediate appears as a product and gets sucked back in