UTA CHEM 1442 Rogers Exam 2 (Ch.13 & 14)

Chemical Kinetics

The rates or speeds at which chemical reaction occurs.

What influences the rate of reactions?

1. Temperature

2. Reactant concentration (pressure)

3. Presence or absence of a catalyst.

4. Nature of reactions

i. aq reaction = fast

ii. Large molecules w/ strong bonds tend to be slow.

5. In a heterogeneous reaction the state of sub-division of the condensed phase.

Catalyst

substance that speeds up the rate of a chemical reaction without being consumed

Rate Constant; k

dependent only on temperature and presence or absence of a catalyst.

the bigger the k the faster the reaction.

Rate law must be determined

experimentally

The exponents in the rate law are referred to as the

order of the reaction

The sum of the exponents is called the

overall order of the reaction

rate=k[A][B]

A: 1st order

B: 1st order

Overall: 2nd order

rate=k[A]^2[B]

A: 2nd order

B: 1st order

Overall: 3rd order

first order reaction

rate=k[A]

second order reaction

Rate=k[A]^2

zero order reaction

rate=k

rate is constant, straight diagonal line when graphed

Rate constant

1. k varies w/ temperature; therefore temperature must be specified.

2. a reaction w/ a large k is FASTER than one w/ a smaller k.

3. k is independent of reactant concentration

integrated rate law first order

ln[A]t = -kt + ln[A]0

a plot of ln[A] vs. time gives a straight line with a slope of -k

The half-life of a first order reaction is

independent of the initial concentration

half-life equation of a first order reaction

t1/2 = ln2/k

The most important reason for studying chemical kinetics is in order to determine

reaction mechanism

reaction mechanism is

a particular sequence of events leading to the overall chemical change

Molecularity

how many particles are involved in an elementary reaction

Unimolecular

an elementary reaction that involves a single molecule

HBr --> H + Br

rate=k[A]

Bimolecular

an elementary step in a reaction that involves two particles

H + Br --> HBr

rate=k[A][B] or rate=k[A]2

Termolecular

simultaneous collision of three molecules (rare)

H2 + Br + Cl --> HBr + HCl

rate=k[A][B][C] or rate=k[A]2[B]

You cannot get the rate law from the _________ but you can get the rate law from _________.

overall balanced equation; an elementary reaction or an elementary step.

Multi step mechanism must satisfy two requirements

1. it must agree w/ the experimentally determined rate law.

2. the sum of the individual elementary step must equal the overall balanced equation.

reactive intermediate

formed in one step and then consumed in a subsequent step. (NO3 in example below)

NO2 + NO2 --> NO3 + NO slow, rds

NO3 + CO --> NO2 + CO2 fast

________________________________________

NO2 + CO --> NO + CO2

two observations of the collision model

first, increasing the temperature of a reaction will increase the reaction rate.

second, increasing the temperature of a reaction will increase the magnitude of the rate constant k.

In order for a reaction to occur, molecules must

1. Collide

2. Collide w/ sufficient energy

3. Collide w/ appropriate orientation in space

Endothermic reaction profile

products are higher than reactants

Exothermic reaction profile

reactants are higher than products

two important principles about Ea (activation energy)

1. The higher the Ea, the slower the reaction

2. A catalyst speeds up a reaction by lowering the Ea

How does a catalyst increase the rate of a reaction?

by lowering the Ea

Does increasing temperature change the Ea?

No

Arrhenius equation

k=Ae^(-Ea/RT)

k= rate constant at temperature (K)

A= pre-exponential term

e= base of natural log

Ea= activation energy (kJ/mol x K)

R= 8.314 J/mol x K = 8.314x10^-3 kJ/mol x K

Easier Arrhenius equation

lnk = -Ea/R(1/T)+lnA

Arrhenius plot

dynamic equilibrium

A state in which no net change occurs because there are two opposing processes occurring at the same rate.

Two characteristics of chemical equilibria

1. two opposing processes occur at the same rate

N2 + 3H2 --> 2NH3

2NH3 --> N2 + 3H2

2. at equilibrium the concentrations of reactants and products do not change with time.

Equilibrium constant for Kc

Equilibrium constant for Kp

Keq > 10^3

mostly products at equilibrium

Keq < 10^-3

mostly reactants at equilibrium

Keq describes

the extent of the reaction

Reaction quotient, Qc

Qc = undefined

1. none when 0/1

2. reverse reaction occurs when 1/0

Qc < Kc

the forward reaction predominates

Qc = Kc

the system is at equilibrium

Qc > Kc

the reverse reaction predominates

if the equation is reversed:

A --> 2B

2B --> A

Kc= [B]2/[A]

Kc'= [A]/[B]2

Kc'= 1/Kc

Heterogeneous Equilibria

ALWAYS omits any solid or liquid for and Keq equation. Because concentration (M) of a solid or liquid is an intensive (doesn't depend on how much you have) property-- like density (mass/vol)

CaCO3(s) <---> CaO(s) + CO2(g)

Kc= [CO2]

Kp= Pco2

Equilibria occurring in aq solutions

Omit H2O(l) from Keq equation

Kp=Kc(RT)^Δn

Δn = the change in the number of moles of gas in the FORWARD reaction

R= 0.08206

Px= must be in atm