Rate of reaction

It is a measure of how fast a reaction takes place. It can be the change in the amount of a
substance (reactant or product) per unit time.

Units of rate;

It can be expressed in moles/dm3/s

DETERMINING RATE OF REACTION

Rate of reactions can only be determined by experimentation.

Consider the reaction;

Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g)

Rate could be considered as the speed at which;

1) magnesium and hydrochloric acid are being depleted/used up or

2) magnesium chloride and hydrogen gas are formed.

The amount of reactants or products can be used to determine the rate of reaction.

In the above example;

i) the formation of hydrogen gas can be followed by measuring its volume at intervals, i.e at ten
seconds for a period of five minutes.

Experimental setup for measuring the volume of gas produced during the reaction

ii) the loss in mass of the reactants can be followed by measuring the mass of the reactants at
intervals, i.e at ten seconds for a period of five minutes.

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Experimental setup

a) Why does the mass of the flask and contents decrease?

………………………………………………………………………………………………………………………………[1]
b) Suggest the purpose of the cotton wool.
………………………………………………………………………………………………………………………………[1]

The results can be recorded in a table and a graph of volume against time can be plotted. The
graph can be interpreted as,

When the graph is steepest, the rate of reaction is high. This means rate of reaction is always
greatest at the start of the reaction because the amount of reactants is still high hence they
interact easily resulting in frequent successful collisions. As the reaction proceeds, the speed of
the reaction decreases.

Why?..................................................................................................................................................

For example; Slope=0 ; Reaction has finished

Small slope; reaction slowing down, lower rate

Greater slope; reaction fastest, high rate

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 COLLISION THEORY AND RATE OF REATION

Substances are made up of particles like atoms, molecules and ions. During chemical
reactions, these particles must collide with enough energy. The rate or speed of reaction
depends on the frequency of successful collisions between the particles of the reactants.

The higher the frequency of collisions between particles, the higher the rate of
reaction.

FACTORS AFFECTING RATE OF REACTION

These factors affect the frequency of successful collisions of the particles.

 Concentration ( number of particles )

 Temperature ( kinetic energy of the particles)
 Surface area (exposure for collision)
 Pressure (number of particles)
 Light intensity ( endothermic and photochemical reactions)
 Use of catalyst ( activation energy)

1) Increase in concentration

At high concentration there are many reacting particles (eg ions) in a small volume. The
particles become closer to each other and there is increased chance of collision hence rate
of reaction is high.

The reaction of sodium thiosulphate and hydrochloric acid can be used to investigate the
effect of concentration on the rate of reaction. The time taken for the appearance of
yellow precipitate of solid sulphur is noted.

Reaction equation

Na2S2O3(aq) + HCl(aq) NaCl(aq) + SO2(g) + S (s) + H2O(l)

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Experimental set up

By using a measuring cylinder, 50 cm3 of sodium thiosulphate solution was poured into a
100 cm3 beaker. The beaker was placed on a cross drawn on a piece of paper. 10 cm3
of
hydrochloric acid was added to the beaker and the timer started

The time was taken until the cross could not be seen. The time was recorded in the table.
Experiments 2, 3, 4 and 5

Experiment 1 was repeated using different volumes of sodium thiosulphate as shown in the
table. All experiments were carried out at 25 °C.

Table of results

a) Why does the cross on the paper disappear?

………………………………………………………………………………………………………

……………………………………………………………………………………………...……[1]

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b) Why was the total volume of solution kept constant?

…………………………....................................................................................................................

………………………...................................................................................................................[1]

c) (i) Plot the results on the graph paper provided(plot volume of thiousulphate vs time) , draw a
smooth line graph and label it 250C.

(ii) Sketch on the grid the graph you would expect if the experiments were repeated at 50oC.
[2]

The graph is as follows;

Low concentration, lower rate

A

volume
B Moderate concentration,low rate

High concentration, high rate

C

time

NB: Basicity can also influence rate of reaction.

For the reactions;

1. HCl + CuCO3 CuCl2 + CO2 + H2O

2. H2SO4 + CuCO3 CuSO4 + CO2 + H2O

The rate of reaction would be highest for the second reaction than the first reaction
because sulphuric acid is dibasic and hence would produce more hydrogen ions in
solution which increase the chances of successful collision.

Strong acids also react faster than weak acids because they ionize completely in solution
producing more Hydrogen ions in solution than weak acids.

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Exercise

Sketch the graphs of the following on the space below;

A= 50 cm3 of 2 mol/dm3 H2O2 (Hydrogen peroxide) decomposing to form oxygen gas for 5
minutes.

B= 100 cm3 of 1 mol/dm3 H2O2 (Hydrogen peroxide) decomposing to form oxygen gas for 5
minutes.

C= 100 cm3 of 2 mol/dm3 H2O2 (Hydrogen peroxide) decomposing to form oxygen gas for 5
minutes.

volume

time

2) Increase in temperature

At high temperature the reacting particles have more kinetic energy hence they move
very fast and collide frequently with each other leading to an increase in the rate of
reaction. At low temperature, particles have less energy hence, the reaction is slower.

3) Increase in Surface area (solids)

Surface area is the area of contact between particles of solid reactants. Powdered
reactants have greatest surface area than lumps, granules and ribbons because more
particles of a solid are exposed hence can collide frequently and react faster.

Exercise

Same masses of calcium carbonate, CaCO3, of different forms were reacted with hydrochloric,
HCl, acid of the same concentration. The forms are;

A- large lumps of CaCO3 B- Medium size particles of CaCO3 C- powder of CaCO3

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-Write a balanced reaction equation.

……………………………………………………………………………………………………

-Label the graphs A,B,C, for the experiments to compare the rates of reactions.

Q. Why do all experiments produce equal volumes of carbon dioxide gas at the end of the
reaction?

………………………………………………………………………………………………………

PROBLEMS CAUSED BY SURFACE AREA

It may cause fire accidents in coal mine and flour mills.

i) Coal mines-When coal is mined, a lot of coal dust particles are produced. The particles
mix with oxygen from air forming a very explosive mixture with more surface area. This
mixture may explode and cause fire accidents if ignited with flame.

ii) Flour mills-Flour dust with more surface area is also produced hence, may also cause
fire accidents.

NB: Warning sign like “NO SMOKING!”or “NO OPEN FIRE!” is found.

d) Pressure (affects gases reacting)

At high pressure, the concentration of gas molecules is very high because molecules
occupy a smaller volume(the molecules are much closer). This enables them to collide
more frequently with each other resulting in a higher rate of reaction.

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 Eg. H2(g) + 3N2(g) 2 NH3(g)

More ammonia is produced when pressure is increased because of increased frequency of

collision between hydrogen and nitrogen gases. What happens at low pressure?

e) Light intensity

It affects endothermic and photochemical reactions. These are reactions which have to
absorb light energy to proceed like photosynthesis. The reactions are faster when there is
too much light.
Silver salts in photography also absorb light and become dark.

2AgBr light 2Ag + Br2

……

The reaction above if faster when there is too much light.

f) Use of catalyst

A catalyst is a substance that speeds up a chemical reaction but remains chemically

unchanged at the end. In the presence of catalyst, collisions between reacting particles
need less energy to be successful.
Catalysts increase the rate of reaction by lowering the activation energy (Ea)
required by a given reaction. In the presence of a catalyst, reversible reactions reach the
equilibrium faster.

Examples of catalysts;

catalyst Type of reaction catalyzed

Iron metal Haber process (manufac. Of ammonia )

Vanadium (v) oxide Contact process ( manufacture of sulphuric acid

Platinum/rhodiu Ostwald process ( manufacture of nitric acid

Manganese (iv) oxide Preparation of………………………………

nickel Manufacture of margerine

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 Energy level diagram showing effect of catalyst.

BIOLOGICAL CATALYST/organic catalysts

They are enzymes produced by living organisms.eg

Salivary amylase digest ……………………………..

Pepsin digest ………………………………

DIFFERENCES BETWEEN ENZYMES AND INORGANIC CATALYSTS

Enzymes Inorganic catalysts

 Are proteins  Are elements and compounds of
transition elements.
 Are denatured by high temperatures  Functions under a wide range of
and inactivated by low temperatures.
temperatures
 Are substrate specific  Used for different purpose.

 They are pH sensitive.  Not usually affected by pH

INDUSTRIAL USES OF ENZYMES

They are used in Biotechnology in the following industries;

 In biological washing powders

 In fermentation i.e baking and brewing alcohol
 In the manufacture of cheese, youghut and baby foods.

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Everyday methods of speeding up and slowing down reactions.

Speeding up Slowing down

 Warming  Freezing
 Crushing  Refrigeration
 Stirring or shaking  Drying e.t.c
 Using more soap,
 Adding yeast e.t.c

Experiments used for measuring rate of reaction

i) Measuring loss in mass of substances at time intervals

For the following reactions, overall mass decreases at time intervals because gases are
released to the surroundings during the reaction.

__________________________________________________________________

____________________________________________________________________

The experiment may be carried out on a weighing balance and mass may be recorded
at time intervals.

EXPERIMENTAL SET UP

Q. Why does the scale reads negative?

………………………………………………………………………………………
What is the purpose of the cotton wool plug?
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 …………………………………………………………………………………………

Exercise
A conical flask with cotton wool plugged and dilute hydrochloric acid weighs 123.0g.
100g of marble chips was added to the conical flask then the total mass was measured
at time intervals and recorded below.

Time 0 1 2 3 4 5 6 7
(mins)
Mass (g) 223.0 204.9 192.8 184.1 180.1 179.0 179.0 179.0
Loss in
mass (g)

Plot a graph of loss in mass against time then answer the following questions.

Why was the graph flat after 4 minutes?

…………………………………………………………………………………

Does the loss in mass decrease or increase with time?

…………………………………………………………………………………

When was the rate of reaction greatest?

…………………………………………………………………………………

Use your graph to determine loss in mass after 3.5 minutes.

…………………………………………………………………………………
How long did the reaction last?

…………………………………………………………………………………

Why is it important to cover the reaction vessel with a cotton wool?

…………………………………………………………………………………
ii) Measuring volume of a gas produced at time intervals.

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 For the reaction;

___________________________________________________________________
Hydrogen gas is produced hence its volume can be measure at time intervals and used to determine rate
of reaction. Volume of gas can be measured using the following;
Using gas syringe-suitable for any gas.
Exp. Set up

Q. What can be done to start the above reaction?

Using graduated measuring cylinder or gas tubes-suitable for insoluble gases only.
Exp. Set up

Q. The following results were obtained when calcium carbonate was reacted with dilute hydrochloric
acid.

Time (min) 0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0
Volume of CO2 0 15 24 28 31 33 35 35 35
(cm3)
Plot a graph of volume of carbon dioxide produced against time then answer the following;

Write a balanced reaction equation

…………………………………………………………………………………

How long did the reaction last?

…………………………………………………………………………………

How long did it take for the reaction to produce 30 cm3 of carbon dioxide gas?
…………………………………………………………………………………

Draw a diagram to show how the experiment could be performed.

Rate of reaction can also be determined by monitoring time taken for;
Bubbles/effervescence to stop
Solid to react completely or disappear
A reaction to produce a given volume of gas (eg 10 cm3)
Colour of products to appear or reactants to disappear
e.g Time taken for colour of products to appear
Effect of concentration on the rate of reaction of sodium thiosulphate and hydrochloric acid is used.
Time taken for appearance of yellow precipitate of solid sulphur is noted.
Reaction equation
____________________________________________________________________________
Experimental set up

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COURSEWORK

SKILL 4

EXPERIMENT WORKCARD

TITTLE: Effect of Surface area of a solid reactant on the rate of reaction

AIM:……………………………………………………………………………………………………….

Apparatus and reagents

 0.1 g ribbon, powder and small pieces.

 Measuring cylinder ( 50 cm3 )
 2 mol/dm3 HCl
 250 cm3 beaker
 Stop clock

Experimental steps/plan/procedure

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 1. Put 0.1 g magnesium powder into a 250 cm3 beaker.
2. Measure out 50 cm3 of hydrochloric acid of concentration 2 mol/dm3 using a measuring cylinder.
3. Pour this acid into a beaker containing magnesium powder and start the stop clock. swirl the
reaction contents of the beaker and when the reaction stops ( the “fizzing” in the beaker will stop
and no solid magnesium can be seen) stop the clock.
4. Record the time taken on the table
5. Rinse out the beaker and repeat steps 1 -4 now using 0.1 g magnesium ribbon and smaller pieces.

Discussion of results

Look carefully at your result table and answer the following questions.

 The variable tested was ‘surface area of magnesium’, name two other variables that might have
affected the experiment and how you handled them.
 How can you improve your experiment in the future?
 Write one or two sentences to critic the method used, procedure.
 Write a one sentence conclusion to your test

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