Rates of Reaction

Objectives• To understand that a chemical reaction

involves collisions between particles• To be able to describe the factors

which will affect the rate of a chemical reaction.

Some Everyday Chemical Reactions

• Burning wood

• Fruit ripening

• Getting a tan

• Cooking food

How Reactions Happen

• A chemical reaction involves a collision between particles.

• The particles collide and make new substances

• The particles which react are called the reactants

• The substances which are made are called the products

Reactions Happen at Different Speeds

• There are chemical reactions that occur very slowly and others that occur very quickly

RUST FORMATIONFIREWORKS

Rates of Reaction

• The rate of a reaction is how quickly a reaction proceeds

• It may be defined as the change in concentration in unit time of any one reactant or product

Rate = Change in concentration (molL-1) Time taken (s)

Calculate the rate of the following

H2O2 → H2O + ½ O2

The initial concentration of the reaction

is 5 molL-1, ten seconds later this hasdecreased 3.5 molL-1

What is the rate?

Calculate the rate of the following

2SO2 + O2 → 2SO3

The initial concentration of the reaction

is 15 molL-1, 20 seconds later this hasdecreased 12.5 molL-1

What is the rate?

Average and Instantaneous Rate

• The average rate of reaction is average rate over the course of the reaction

• The instantaneous rate of reaction is the rate at a particular point in time during the reaction

Controlling the Rates of Reaction

• Being able to control the speed (rate) of chemical reactions is important both in everyday life (cooking) and when making new materials on an industrial scale.

Factors Affecting Rate of Reaction

1. Concentration of reactants2. Temperature of reaction3. Particle size of solid reaction4. Nature of reactants5. Presence of catalyst

1. Concentration of Reactants

• The higher the concentration of reactants the higher the probability of collisions between reactant molecules

There are less red particles in the same volume so there is less chance of a collision

There are more red particles in the same volume so there is more chance of a collision so the reaction goes faster

2. Temperature of Reaction

• An increase in temperature brings about an increase in reaction rate.

• You give more energy to the system in the form of heat

3. Particle Size

• Particle size (finely divided particles react faster)

• Molecules can only collide at the surface.

• Smaller particles bigger surface area

Dust Explosions

• A dust explosion is the explosive combustion of a dust suspended in air in an enclosed location

• Any solid material that can burn in air will do so at a rate that increases with increased surface area

Grain Dust Peril• In 1998, at a series of

explosions occurred at a grain elevator facility in Haysville, Kansas

• There were seven fatalities as a result of the explosions.

• It is not the actual grain that is ignited, but the fine, thick dust which is released during the loading process when grain particles rub against each other

4. Nature of Reactants• Ionic compounds react faster than covalent

• In reactions bonds are broken and form• When an ionic compound is placed in water

it dissociates• It takes more energy to break covalent

bonds

Activation Energy (Eact)

• For a reaction to happen, energy is required.• Activation energy is the minimum energy with which particles need to collide to

cause a reaction

• This is different according to the type of bonds of the reactants.

• The activation energy may be shown on a reaction profile diagram (right)

• These diagrams show energy as a barrier that needs to be overcome by the reactants before they become products

• The difference between the energy of the reactants and the energy of the products is the heat of reaction (ΔH)

Endothermic reaction

• A reaction in which heat is taken in.

• In an endothermic reaction heat is taken in from the surroundings and the products formed have more energy than the reactants.

• It is written as + ΔH

Exothermic Reaction

• A reaction in which heat is liberated.

• In an exothermic reaction heat is lost to the surroundings and the products formed have less energy than the reactants.

• It is written as - ΔH

5. Catalysts• A catalyst is a substance that alters the

rate of a chemical reaction but is not consumed (used up) in the reaction.

• In most cases it makes the reaction go faster. Some catalysts make a reaction go slower and are called negative catalysts or inhibitors. Eg Calcium propionate added to bread to make it stay fresh longer (ie. It slows down staleness)

• The catalyst does not get used up in the reaction.

• It gives the reaction the energy to get started

General Properties of catalysts

1. Catalysts are recovered chemically unchanged at the end of a reaction

(Eg. Manganese dioxide used to speed up decomposition of Hydrogen peroxide has exactly the same chemical properties before and after the reaction)

2. Catalysts tend to be specific, even though a catalyst may catalyse one reaction it may not have any effect on a similar reaction

Enzymes in the body are examples of catalysts that are very specific

Know two examples• Protease breaks down proteins such as

blood stains on clothes and are used in washing powders

• Catalase breaks down hydrogen peroxide in the body

3. Catalysts need only be present in very small amounts

• Increasing the amount of catalyst does not greatly affect the rate of a reaction and in cases where it does it is usually something to do with the reaction itself

4. Catalysts help reactions reach equilibrium quicker but do not change what the equilibrium of a reaction is

• 5. Action of catalysts may be destroyed by catalytic poisons for example lead in petrol can destroy the catalytic converters in cars

Arsenic is a poison that inhibits the action of certain enzymes in the body

Types of Catalysis

• Chemists have discovered 3 types of catalysis

1.Homogenous Catalysis2.Heterogeneous Catalysis3.Autocatalysis

Homogenous Catalysis

• This describes when reactants and catalysts are in the same phase and there is no boundary between them

• Eg Iodine Snake where Potassium Iodide catalyses the decomposition of H2O2 to release oxygen both catalyst + Reactant are liquids

Heterogeneous Catalysis

• Catalysis where the catalyst and reactants are in different phases Eg. a liquid and a solid

• There is a boundary between the catalyst and the reactants

• Eg. MnO2 catalyses the decomposition of H2O2 to release oxygen both catalyst + Reactant are liquids

Autocatalysis

• When one of the products of the reaction catalyses the reaction

• In this type of reaction it occurs slowly at first but as the reaction proceeds it gets quicker this is because the products drive the reaction forward.

Theories of Catalysts

• Speed up a reaction by giving the reaction a new path.

• The new path has a lower activation energy and more molecules have this energy.

• The reaction goes faster.

CATALYST

Think of a catalyst as a tunnel through a

mountain

By lowering the activation energy a catalyst makes it possible to carry out a reaction at lower temperatures (lower energy)

Mechanisms of Catalysis

• The mechanism of catalysis tells us how the catalyst works

• There are two main mechanism of catalysis for you to study

1.Intermediate Formation theory2.Surface Adsorption theory

Intermediate Formation Theory of Catalysts

• Homogeneous catalysts sometimes work by reacting with reactants to form unstable intermediate products

• The intermediate exists for a very short time and reacts with the other reactant to give the final product and regenerate the catalyst

X+A + B → C

AX

+ B

See it with you own eyes

• Oxidation of Potassium Sodium Tartrate using hydrogen peroxide

• Catalyst in this reaction is Cobalt ions which give a pink colour

• The intermediate is a green colour which appears as carbon dioxide + steam are given off

• The pink colour is restored at the end indicating the Cobalt ions have not been used up

Surface Absorption Theory of Catalysts

• Heterogenous catalysis of gas reactions by metals

• The reaction happens on the surface of the metal

• The reaction occurs at the active site of the catalyst

• A catalyst can have multiple active sites

CATALYST

CATALYST

CC H

H

H

HH H H H

C CH

H

H

HH H H H

Stages of reactions of ethene

Reactants get absorbed onto catalyst surface. Bonds are weakened

CATALYST

CATALYST

C CH

H

H

HH H H H

HHH C CH

H

H

H

H

Bonds Break

New bonds formed

CATALYST

H HC C

H

H

H

H

H

H

Second bond forms and product diffuses away from catalyst surface, leaving it absorb fresh reactants

Catalytic Poisons• Catalysts can be poisoned they can become less

efficient and sometimes they no longer work at all• In heterogeneous catalysis particles that poison the

catalyst (lead / arsenic) are absorbed more strongly onto the catalyst surface than the reactant particles

• Catalytic poisons block the active sites of enzymes

How it works (Catalytic Converter)

• The catalyst in the converter speeds up reactions that reduce atmospheric pollution

• The catalyst remains unchanged at the end of the reaction

• The catalyst is a mix of transition metals (platinum, rhodium, palladium)

Reactions Catalysed• Carbon monoxide is converted to

carbon dioxide by reaction with oxygenCO + ½ O2 → CO2

• Carbon monoxide can react with nitrogen monoxide to give carbon dioxide

2CO + 2NO → 2CO2 + N2

• Unburnt hydrocarbons are oxidised to carbon dioxide and water

C8H18 + 12½ O2 → 8CO2 + 9H2O

Environmental Benefits

• Reduction in emissions of toxic gases including unburnt hydrocarbons

• Reductions in photochemical smog

Smog in Beijing

Mandatory Experiment 14.1

Monitoring the rate of production of oxygen from hydrogen peroxide, using

manganese dioxide as a catalyst

Hydrogen peroxide→Oxygen + Water

How to layout your results

O2 Cm3

Time secs

T 0 15 30 45 60 90 120 150 180 210

Vol 0 60 110 134 150 170 184 190 197 197

15

60

oxygen production by decomposition of hydrogen peroxide

0

10

20

30

40

50

60

70

0 2 4 6 8 10 12

Time/mins

Oxy

gen

/cm

3

As the reaction proceeds more oxygen is produced

When the reaction is complete the amount of oxygen stops increasing

Average and Instantaneous Rates

• The instantaneous rate of reaction is the rate at a particular point in time during the reaction

• You use your graph to find the instantaneous rate

oxygen production by decomposition of hydrogen peroxide

0

10

20

30

40

50

60

70

0 2 4 6 8 10 12

Time/mins

Ox

yg

en

/cm

3

Draw a tangent to the curve

oxygen production by decomposition of hydrogen peroxide

0

10

20

30

40

50

60

70

0 2 4 6 8 10 12

Time/mins

Ox

yg

en

/cm

3

Connect the tangent to the X and Y axis using straight lines

V1

V2

t1 t2

CalculationsInstantaneous rate= tan θ= Δv = v2 – v1

(Slope) Δt t2 – t1

= 52-45 (cm3)

3 - 2 (min) = 7 (cm3) 1 (min)

instantaneous rate = 7cm3 min-1

Mandatory Experiment 14.2

Studying the effects on the reaction rate of (i) concentration and (ii)

temperature, using sodium thiosulfate solution and hydrochloric acid

2HCl(aq) + Na2S2O3(aq) → 2NaCl(aq) + SO2(aq) + S(s)↓ + H2O(l)

Place 100 cm3 of the sodium thiosulfate solution into a conical flask.

Add to the flask, while starting the stop clock at the same time.

What’s happening?

• Repeat the experiment using 80, 60, 40 and 20 cm3 of. sodium thiosulfate solution respectively. In each case, add water to make the volume up to 100 cm3 and mix before adding HCl.

Concentration of Thiosulphate

Reaction time(s)

Rate of Reaction (1/time)

(s-1 )

0.1 M

0.08 M

0.06 M

0.04 M

0.02 M

Draw Graph

1time

Concentration of Thiosulphate