KineticsChapter 6

6.1 Rates of Reactions• Define the term rate of reaction.

• Describe suitable experimental

procedures for measuring rates of reactions.

• Analyse data from rate experiments.

Rate

of

react

ion is

defined a

s th

e

rate

change in

co

nce

ntr

ati

on

• Rate – per unit of time• 1/time• Per second • 1/s

• s -1

• Rate of reaction – how

quickly a rxn happens• How fast reactants are

converted to products• Visa versa

• Concentration of product against

time• Concentration of reactant against

time

• Rate of rxn formuals

Negative sign shows the concentration is decreasing, but rate is expressed as a positive value

• Rate units – mol dm -3 s -1

• Change in concentration per time

• Graphs• Gradient of line is a measure of the

change in concentration per time

• On a curve gradient is not constant –

draw a tangent line to the curve &

calculate gradient of that line

• Rate of rxn is not constant

• Faster at beginning – steep line

• Slows as it continues – less steep

line• Rxns are compared at initial rates

at time zero

Measu

ring r

ate

s of

react

ion u

ses

diff

ere

nt

tech

niq

ues

dependin

g

on t

he r

eact

ion

1. Change in volume of gas

produced2. Change in mass3. Change in transmission of light:

colorimetry / spectrophotometry4. Change in concentration

measured using titration

5. Change in concentration

measured using conductivity

6. Non-continuous methods of

detecting change during a

reaction: ‘clock reactions’

1.Change in volume of

gas produced• Graphing change in volume VS change in

time• Gas syringe – tool to

collect gas• Inverted burette • Most gases are less soluble in warm water – using warm water lowers error

2. Change in mass• If the rxn gives off a gas

• Works well for light gases• Graph – decreasing

mass VS increasing time

3.Change in transmission of light:

colorimetry/spectrophotometry

• Used if reactant or product is

coloured• Different absorption in the visible

region of spectrum• Indicator can be used to make

coloured compound

• Colorimeter or spectrophotometer passes light

of a selected wavelength

through the solution measuring

the intensity of the transmitted

light

2HI(g) H2(g) + I2(g)

colourless colourless coloured

• Allows for continuous readings

• Graph of absorbance VS time

• Convert absorbance to

concentration by using a

standard curve based on

readings of known concentrations

4.Change in concentration

measured using titration• Titrating against a ‘standard’

• Cannot be done continuously

• Samples must be taken and

tested at regular time intervals

• Quenching – a substance is

introduced stopping the rxn

when the sample is taken

5.Change in concentration

measured using conductivity

• Total electrical conductivity of a

solution depends on the total

concentration of its ions

• change in conductivity indicates a

change in the concentration of

ions• Inert electrodes are immersed in

solution – calibrated using known

solutions

BrO3

-(aq) + 5Br -(aq) + 6H+(aq) 3Br2(aq) + 3H

2O(l)

• Sharp decrease in electrical conductivity due to

decrease in ions• 12 mols of ions on the reactants side & 0 moles

on products side

6.Non-continuous methods of

detecting change during a

reaction: ‘clock reactions’

• measure of the time taken to

reach a certain chosen fixed point

• Time taken for a certain mass of

Mg ribbon to react completely in

dilute acid• Color change • Limitation – only gives average

rate over a time interval

6.2 Collision Theory

• Describe the kinetic theory in terms of the

movement of particles whose average energy is

proportional to temperature in kelvins.

• Define the term activation energy, Ea.

• Describe the collision theory.

• Predict and explain, using the collision theory, the

qualitative effects of particle size, temperature,

concentration, pressure on the rate of reaction.

• Describe the effect of a catalyst on a chemical reaction.

• Sketch and explain Maxwell-Boltzmann curves for

reactions with and without catalysts.

Kin

eti

c energ

y &

te

mpera

ture

• Kinetic-molecular theory

of matter• All particles move randomly due to possessing kinetic energy• Not all particles have

the same values of kinetic energy

• Absolute temperature –

measure of the average kinetic

energy of the particles – Kelvin

scale• Increase in temperature =

increase in average kinetic

energy of the particles• Differences in the 3 phases of

matter is average kinetic

energy of the particles

The M

axw

ell-

Bolt

zman

dis

trib

uti

on c

urv

e

• Shows the range of values of kinetic energy of particles in

a gas• Shows the numbers of

particles that have a

particular value of kinetic energy• Area under the curve

represents the total number of sample particles

How

react

ions

happen

• Reactant particles collide

with each other due to

their kinetic energy• Energy from collisions may

result in bonds between

reactants• being broken• being formed• Rate of rxn depends on the

number of “successful”

collisions – not all collisions will be successful

• Two factors influence

successful collisions• 1. Energy of the collision

• 2. Geometry of the collision

• 1. Energy of collision• Activation energy (E

a)- the particles

must have a minimum value of

kinetic energy – a threshold of energy

• Necessary to overcome repulsion

between molecules / to break bonds

in reactants• An energy barrier for the reaction

• Only particles with the minimum

kinetic energy will react

• Transition state – reactants after

activation energy is supplied –

products can form• Rate of rxn depends on the

proportion of particles with kinetic

energy values higher than Ea

• 2. Geometry of collision• Collisions occur with

different orientations• For a rxn to occur both

particles must have the

correct orientation during

the collision

Fact

ors

aff

ect

ing

rate

of

react

ion 1.Temperature2.Concentration3.Particle size4.Pressure5.Catalyst

1.Temperature• Increase in temperature =

increase in average kinetic

energy of the particles =

increase in collision frequency

= increase in collisions with

enough Ea and correct

orientation = increase in rate

of rxn

2.Concentration• Increase in concentration =

increase in frequency of

collisions = increase in rate

of rxn• As rxn progresses reactants

are used up (conc. Decreases) and rxn rate

decreaces

3.Particle size• Decreasing particle size = increases

particle surface area = increase in rxn

rate• Important in heterogenous rxns –

reactants in different phases

4.Pressure• With gases, increase in pressure =

increase in rxn rate• Same as increasing concentration

5.Catalyst• Increases rxn rate without

undergoing a chemical change

• Lowers Ea – provides a different

route for the rxn• Increases the number of particles

with enough Ea to react without

raising the temperature• Equal reduction in E

a for forward

and reverse rxns

• Important to many industrial

processes• Enzyme – every biological rxn is

controlled by a catalyst

16.1 Rate Expression≈Distinguish between the terms rate

constant, overall order of reaction and

order of reaction with respect to a particular

≈Deduce the rate expression for a

reaction from experimental data.

≈Solve problems involving the rate expression.

≈Sketch, identify and analyse graphical

representations for zero-, first- and

second-order reactions

The r

ate

law

for

a

react

ion is

deri

ved f

rom

ex

peri

menta

l data

• See page 216• Chart for data collected in lab• Graph made from data

on chart.

𝑟𝑎𝑡𝑒=− Δ [𝑐60𝑂3]

Δ𝑡

• “Rate Law” = a math formula

• AKA – “rate expression”

• Reaction rate = k[conc.]• k = rate constant – fixed for

a particular rxn at a specific

temp.

• Rate is dependent on

one of the reactants’

conc.• A + B products• Rate = k[A]m [B]n• Proportional to the

concentrations of the

reactants

• ”Order with respect to…”

• A + B products• Rate = k[A]2[B]1• Rxn is 2nd order with

respect to hydrogen• Rxn is 1st order with respect to oxygen

• Overall order of the rxn is

3 rd order 2 + 1 = 3

2H2(g) + 2NO(g) 2H

2O(g) +

N2(g) is show to be 2nd order with

respect to NO and 1st order with

respect to H2.

Write the rate expression for the

above rxn.

Overall order?

Unit

s of

k va

ry

dependin

g o

n t

he

ove

rall

ord

er

of

the

react

ion

Rxn

Order with

respect to

reactant 1

Order with

respect to

reactant 2

Overall order of

rxn

H2(g)+I2(g)2HI(g)

2H2O2(aq) 2H2O(l)+O2(g)

S2O8-2(aq)+2I-(aq)

2SO4-2(aq)+I2(aq)

2N2O5(g)4NO2(g)+O2(g)

The orders of reaction do NOT necessarily correspond to their

coefficients.

Unit

s of

k va

ry

dependin

g o

n t

he

ove

rall

ord

er

of

the

rxnZero

OrderFirst Order

Second Order

Third Order

Rate = k Rate = k[A] Rate = k[A]2 Rate = k[A]3

mol dm-3 s-

1 s-1 mol-1dm3s-1 mol-2dm6s-1

Points are given for the correct units of the rate constant. You must memorize the units… they are different for each

over-all order.

A rxn has the rate expression: rate = k[A]2[B]

Calculate the value of k, including

units, for the reaction when the

conc. of A & B are 2.50 X 10 -2 mol

dm -3, and the reaction rate is 7.75 x

10 -5 mol dm -3 min -1.

Gra

phic

al

repre

senta

tions

of

rxn k

ineti

cs • Zero-order rxn• Rate = k[A]0 or rate = k

• Conc. Vs. time graph• Gradient = k• Rate Vs. conc. Graph• Horizontal line

You will have to identify reaction orders by reading a graph…

• 1st order rxn• Rate = k[A]• Conc. Vs. Time graph• Curve showing rate decreasing with conc.

• Rate Vs. Conc. Graph• Straight line passing

through the orgin• Gradient = k

• 2nd order rxn• Rate = k[A]2• Conc. VS. Time graph• Curve; steeper at start

• Rate Vs. Conc. Graph• Parabola• Gradient •proportional to the conc.

• Initially zero

• Only 1st order rxns have a

constant half-life• Conc. Vs. Time graph• Half-life – t

1/2• Time for conc. to decrease to

half of original value• Constant half-life

Dete

rmin

ati

on o

f th

e

ord

er

of

a r

eact

ion

• Initial rates method• Several separate experiments with different starting conc. of reactant A &

measuring the initial

rate of each rxn• Process repeated for

reactant B

• If changing conc. of A has no

effect on rate order of zero

with respect to A• If changes in conc. A produce

directly proportional changes in

the rate 1st order with respect

to A• If a change in conc. A leads to

an increase in the rate of rxn

equal to the square of the

change 2nd order with respect

to A

Use the following data

to work out the order

of the rxn with respect

to reactants A & B. Write the rate expression for the reaction.

Experiment Number

Initial concentration (mol dm-3)

Initial rate of reaction (mol dm--3 s-

1)

[A] [B]

1 0.10 0.10 2.0 x 10-4

2 0.20 0.10 4.0 x 10-4

3 0.30 0.10 6.0 x 10-4

4 0.30 0.20 2.4 x 10-4

5 0.30 0.30 5.2 x 10-4

Summary of dedcutions

Change in [A]

Change in rate of

zero-order rxn

Change in rate of 1st-order rxn

Change in rate of 2nd order rxn

[A] doubles No changeRate

doubles (x 2)

Rate x 4

[A] triples No changeRate triples

(x 3)Rate x 9

[A] increases four-fold

No change

Rate increases

four-fold (x 4)

Rate x 16

16.2 Reaction Mechanism≈Explain that reactions can occur

by more than one step and that

the slowest step determines the

rate of reaction (rate determining step).

≈Describe the relationship

between reaction mechanism,

order of reaction and rate-determining step.

Most

react

ions

invo

lve a

seri

es

of

small

steps

• Reaction mechanism• The theory of what’s happening

in a rxn• Series of simple steps making a

rxn• Elementary steps – individual

steps – cannot be observed

directly • Intermediates – products of one

elementary step that are used

as reactants in the next

elementary step• The sum of the elementary

steps must equal the overall

rxn

NO2(g) + CO(g) NO(g) + CO

2(g)

Step 1: NO2(g) + NO

2(g) NO(g) + NO3(g)

Step 2: NO3(g) + CO(g) NO

2(g) + CO2(g)

NO3 – is an intermediate – produced and then consumed

• Molecularity - references an

elementary step indicate the

number of reactant species

involved• Unimolecular rxn – elementary rxn

involves a single reactant particle

• Bimolecular rxn – involves two

reactant particlesNO2(g) + CO(g) NO(g) + CO

2(g)

• Termolecular - very rare due to 3+

particles colliding at same time with

energy and orientation to react

The r

ate

-dete

rmin

ing

step is

the s

low

est

ste

p

in t

he r

eact

ion

mech

anis

m

• Rate-determining step –

slowest step

The r

ate

exp

ress

ion f

or

an o

vera

ll re

act

ion is

dete

rmin

ed b

y th

e

react

ion m

ech

anis

m

Equation for rate-

determining step

Molecularity Rate Law

A products Unimolecular Rate = k[A]

2A products Bimolecular Rate = k[A]2

A + B products

Bimolecular Rate = k[A][B]

2NO2Cl(g) 2NO(g) + Cl2(g)

Step 1: NO2Cl(g) NO

2(g) + Cl(g)slow

Step 2: NO2Cl(g) + Cl NO

2(g) + Cl2(g)

fastOverall: 2NO

2Cl(g) 2NO(g) + Cl2(g)Rate = ?Order of rxn = ?

2NO(g) + O2(g) NO

2(g)Step 1: NO(g) + NO(g) N

2O2(g) fast

Step 2: N2O2(g) + O

2(g) 2NO2(g)

slowOverall: 2NO(g) + O

2(g) NO2(g)

Rate = ?Overall order of rxn = ?

Zero order reactant –

does not take part in the

rate of the rxn

16.3 Activation Energy

≈ Describe qualitatively the relationship between

the rate constant (k) and temperature (T).

≈ Determine activation energy (Ea) values from

the Arrhenius equation by a graphical method.

The r

ate

const

ant

k is

te

mpera

ture

dependent

• Rule of Thumb: 10°C

increase = doubling of the

rate• Rate of Rxn depends on two

things:• Rate constant, k• Conc. of reactants raised to

a power• Increasing the temp has no

effect on the conc; changes

the value of the rate

constant k• k is temperature specific

• Collision Theory • Increasing temp. increases

collisions increases rxn rate

The t

em

pera

ture

dependence

of

the r

ate

const

ant

is e

xpre

ssed in

the A

rrheniu

s equati

on

• Arrhenius Equaiton – shows

that the fraction of molecules with energy > Ea

at T is proporitonal to e -Ea/RT

• k = Ae -Ea/RT• R = gas law constant; 8.31

J/K mol• T = absolute temp; K• A = Arrhenius constant;

frequency factor; pre-

exponential factor• E

a = activation energy

• Arrhenius constant• Frequency w/which successful collisions occur

• Collision geometry• Energy requirements• Constant for the rxn • Same units as k – varies

with order of rxn

• Taking the natural log of both

sides of the equation• ln k = -E

a/RT + ln A• Form of equation for straight

line• y = mx + c• Arrhenius plot – graphing ln k

VS. 1/T gives a straight line

with a gradient of -Ea/R

Determine the activation energy in kJ mol -

1 by graphical method.

Rate constant (s-1) Temperature (°C)

2.88 x 10-4 320

4.87 x 10-4 340

7.96 x 10-4 360

1.26 x 10-3 380

1.94 x 10-3 400