© Boardworks Ltd 2003 RATES OF REACTION. © Boardworks Ltd 2003 Rates of reaction Reactions can be...

© Boardworks Ltd 2003

RATES OF REACTION


Rates of reaction

Reactions can be very fast, like fireworks or explosives, but they can also be very slow – such as an apple turning brown.


Chemical reactions occur when particles of reactant collide with enough energy to react.

Rates of reaction


Speeding up reactions

• Anything that increases the chance of effective collision increases the rate (speed) of reaction. Factors include:

• Increased surface Area

• Increased concentration

• Increased temperature

• Use of a catalyst


Surface area

• The reactions of solids can clearly only take place at the surface of the solid.

• If we break a solid into smaller pieces we get more area and a faster reaction.

Molecules collide with the surface of the solid

Extra surface for molecules to collide with.


• If we grind up a solid to a powder we massively increase the surface area.

• We therefore massively increase the rate of any reaction

Very fast

Slow

Surface area


A B

An indigestion tablet fizzes in water – but fizzes much faster if it is crushed.Which glass has the crushed tablet?

CrushedSolid


Concentration

• Reactions in solution involve dissolved paticles that must collide before reaction is possible.

• The more crowded (concentrated) the solution, the faster the reaction.

Collisions infrequent Collisions frequent



Pressure

• Reactions involving gases are affected by the pressure of the gases present.

• If we cover one end of a bicycle pump and push in the plunger we increase the pressure.

• What we are doing is squeezing the gas molecules closer togethercloser together or making them more concentrated.

• And so - pressure speeds up gas reactions

Low pressure

High pressure


• The Haber Process, in which nitrogen reacts with hydrogen to form ammonia, is carried out at 200 atmospheres pressure.

• How and why will this affect the rate of reaction?

The particles will be 200 times closer together and so will collide much more often.

The reaction will be much faster.

compress


Temperature

• At higher temperatures molecules move faster. As a result there are more collisions per second and so a faster reaction occurs. Slow molecules are also less likely to lead to a reaction than fast ones.

Fewer collisions per second More collisions per second


• Food spoils because of chemical reactions that occur.

• Why does food remain usable for so much longer if it is kept in a freezer?

The reactions that cause the food to go off will be slower because there will be fewer and “softer” collisions between molecules at a reduced temperature.


• Before microwave ovens were common many people used pressure cookers.

• It was a pan that stopped the water boiling until it reached about 115oC.

• How would this help cooking?

The molecules move faster and collide more often and with more energy.

Cooking times were greatly reduced.


New bonds form

Old bonds start to break

Activation energy

• Chemical reactions involve the formation of bonds between atoms but often before new bonds can be formed old ones have to be broken.

• This means that there has to be enough energy (activation energyactivation energy)to start breaking the old bonds before a reaction can occur.

Reactants

Activation energy

needed to break existing

bonds

Energy is given out as new bonds

form


Most reactions are exothermic (give out heat) overall but there is still a need for energy to get the reaction started.

Ene

rgy

in c

hem

ical

s

Reaction

Activation energy

Break old bonds

Energy taken in

Form new bonds

Energy given out

Activation energy


1. Why doesn’t petrol catch fire when it is poured through air?

2. Why is just one spark enough to create a major explosion?

1. Energy is needed to break the bonds in petrol before new bonds can be formed by a reaction with oxygen.

2. Once some of the bonds in one petrol molecule have been broken the subsequent reaction with oxygen gives out enough energy to break the bonds in several other petrol molecules - and so on.



Catalysts

For chemical reactions to occur:

• Existing bonds have to begin breaking so that new ones can be formed.• The molecules have to collide in such a way that the reacting parts of the molecules are brought together.

Catalysts can help with either or both of these processes.

A catalyst is a substance that speeds up a reaction without getting used up in the process.


• In the presence of a nickel catalyst vegetable oil and hydrogen react to form margarine.

• Nickel adsorbs hydrogen gas onto its surface in such a way that the bond holding the hydrogen molecule together becomes stretched.

• This partial breaking of the bond lowers the activation energy making hydrogen more reactive.

H HH H

H HH HHH HHBond stretched

The stretching of the H-H bond lowers the activation energy and helps hydrogen react with the oil

Ni NiNi Ni NiNicatalyst

Catalysts


Other catalysts, especially enzymes, absorb molecules in a way that not only stretches bonds but also brings the reacting parts of reactants right next to each other.

Catalysts


Catalysts are used in the manufacture or application of a huge number of products.

Synthetic materials like polyester are made using a

catalyst.

Manufacture of fertiliser via the Haber Process

involves use of an iron catalyst.

Plastics are made using catalysts.

Biological soap powder uses

biological catalysts (enzymes)

Enzymes in pineapple help cooked ham to be more tender.

Catalysts


Inside car engines some of the nitrogen and oxygen from the air combine to form poisonous nitrogen oxide. Inside the exhaust system a catalyst encourages decomposition back into nitrogen and oxygen.

1. Copy the energy profile for the uncatalysed reaction and draw in new lines showing how the presence of a catalyst will alter the profile. 2NO2

N2 + 2O2


Measuring reaction rates

Rate implies we are measuring how things change over a period of time.

To measure the rate of a reaction we have to track the manner in which the amount of product (or reactant) changes over time.

Rate of gas formation can be measured using a syringe.

For a reaction in which sulfur is precipitated we can time how long the solution takes to go cloudy.


Slower and slower

Reactions do not proceed at a steady rate. They start fast and get slower and slower.

This is not surprising because the reactant concentration (and the chance of collision) gets lower and lower as time progresses.

Percentage completion of reaction

0

fast

100

stopped

25

slower

75

very slow

rea

cta

nts

pro

du

ct


Rates and Graphs

• These show the increasing amount of product or the decreasing amount of reactant.

Am

ount

of

prod

uct

Time

Am

ount

of

reac

tant

Time

Steep gradient Fast reaction

Shallow gradient Slow reaction

Steep gradient Fast reaction

Shallow gradient Slow reaction


Rate graphs and reactant concentrationsA

mou

nt o

f pr

oduc

t

Time

reactants

product

Reactant Concentration falls

Rate of Reaction falls

All product

All reactant

Mix of reactant

And product

Gradient of graph decreases




Some Reaction Rates Experiments

The following slides describe the four chemical reactions that are commonly used

as examples.


• Marble chips are calcium carbonate.• They react with acid to evolve a gas.

calcium carbonate

+ hydrochloric acid

calcium chloride

+ water + carbon dioxide

CaCO3(s) + 2HCl(aq) CaCl2(aq) + H2O(l) + CO2(g)

Marble chips

Glass tube

Gas syringe

Hydrochloric acid

The gas given off can be collected in a syringe and readings taken every 30 seconds or so.

Acid and marble


1. Measure the agreed mass of marble chips

2. Set up the syringe, flask and connector

3. Measure the acid / water.

4. Add the marble chips and quickly insert the bung and start stop clock.

5. Take syringe readings at 30 second intervals.

Time Reading

0 s 0 cm3

Acid and marble


Acid and marble


Mg

HCl

Acid and metal

• Reactive metals (eg. Magnesium) react with acid to evolve hydrogen gas.

magnesium + hydrochloric acid

magnesium chloride

+ hydrogen

As the gas given off leaves the flask the total mass of the flask and its contents decreases slightly.

Readings of the mass(g) can be taken. Typically at 1 minute intervals.

Mg(s) + 2HCl(aq) MgCl2(aq) + H2(g)

11.8011.7711.7411.7311.7211.7111.71


Mg

HCl

Cotton wool

1. Measure the agreed volume of acid / water into the conical flask.

2. Have a loose plug of cotton wool to prevent “spitting” of droplets of liquid.

3. Have a piece of magnesium of known mass ready.

4. Add the magnesium, place the cotton wool in the neck and start taking mass readings immediately.

Time Reading

0 s 0 cm3

60

120

11.8011.7711.7411.7311.7211.7111.71

Acid and metal


Acid and metal


Decomposition of Hydrogen Peroxide

• Hydrogen peroxide decomposes into water and oxygen.

Hydrogen peroxide water + oxygen

Oxygen gas is given off and can be measured using a gas syringe or a balance.

The reaction is catalysed by a wide range of solids.

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

Remember the catalyst NEVER produces more product - just quicker



Acid and Sodium Thiosulphate

• In this reaction sulphur is precipitated which makes the solution turn cloudy.

The effect of changing conditions such as temperature or concentrations can be studied by measuring how long it takes to produce enough sulphur to make the solution opaque (non see-through).

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

Sodium thiosulphate

+ hydrochloric acid

sodium chloride

+ water + sulphur dioxide

+ sulphur


1. Measure the agreed volume of thiosulphate / water into the conical flask.

2. Prepare a piece of paper with a cross drawn on it.

3. Measure the required volume of acid in a measuring cylinder.

4. Add the acid to the flask, start the clock, swirl the flask.

5. Look down through the flask until the cross disappears. Note the time.

Look down here


Temp

(oC)

Time taken

(s)

25 100

30 60

35 40

40 25

45 15

50 10

Imran studied the effect of temperature upon the time it took for the flask to go cloudy.

1. Sketch a graph of the results.

2. Using the reactions at 25oC and 40oC, explain how the time taken lets you work out the relative rate of reaction.

0

20

40

60

80

100

120

20 30 40 50 60

Time(s) vs Temp(oC)

Tim

e

Temp


Some Rates Questions



A pupil performed an investigation into the rate of reaction between a

metal and an acid. The results below where obtained. Time / seconds

0 10 20 30 40 50 60 70 80

Volume of gas (cm3)

0 25 80 179 245 273 282 282 282

i) Plot a graph of gas volume (y-axis) against time (x-axis)ii) When was the rate of reaction fastest?iii) Use the graph to find the volume of gas produced after 35 seconds.iv) Use the graph to tell after how long the reaction stopped.v) On the graph sketch a line showing the experiment repeated

at a higher temperature.


Experimental Results

0

50

100

150

200

250

300

0 10 20 30 40 50 60 70 80 90

Time / seconds

Vol

ume

/ cm

3

ii) The reaction was fastest at about 25 seconds as the gradient of the line is highest at this point.iii) About 175 cm3

iv) About 55 seconds.v) Higher temperature reaction is in red.

Answer


Time / seconds

0 10 20 30 40 50 60 70 80

Volume / cm3

0 18 30 40 48 53 57 58 58

A flask was connected to a gas syringe by a glass delivery tube. 30cm3 of water and 0.5g of manganese dioxide were added to the flask. Then 5cm3 of hydrogen peroxide was added and the stopper quickly fitted. Readings of the volume of gas produced were taken every 10 seconds.

i) Plot a graph of volume of gas (y-axis) against time (x-axis). Label this curve A.ii) Without emptying the flask another 10cm3 of water and a further 5cm3 of hydrogen peroxide were added. Sketch the shape of the second experiment and label it B.


Notes: Curve B is an experiment with half the concentration of hydrogen peroxide. This should produce about half the rate as shown by a line with half the gradient of A. However, the same amount was added so 58cm3 of gas will still be produced.

0

10

20

30

40

50

60

70

0 10 20 30 40 50 60 70 80 90 100 110 120

Vol

ume

of g

as /

cm

3 A

B

Time (s)

Answer


Time / seconds 0 15 30 45 60 75 90 120

Mass loss / g 0 0.21 0.45 0.67 0.85 1.01 1.13 1.31

Time / seconds 150 180 210 240 300 360 420

Mass loss / g 1.41 1.48 1.51 1.54 1.56 1.58 1.59

i) The results above were obtained from an experiment where the loss in mass was recorded as lumps of zinc reacted with hydrochloric acid. Plot a graph of mass loss (y-axis) against time (x-axis).

ii) On the graph sketch the lines you would expect if

a) the concentration of acid was reduced, b) the temperature was increased.


0

0.5

1

1.5

2

0 200 400 600

Time / seconds

loss

in

mas

s /g

(b)

(a)

Answer


Which of these would speed up the rate at which magnesium dissolves in acid?

A. Cool the acid.

B. Cut up the magnesium.

C. Add water.

D. Coat the magnesium in oil.


Why does breaking up solids increase the rate of reaction?

A. Makes more solid.

B. Creates more energy.

C. Increases surface area.

D. Increases the concentration.


Why does temperature increase the rate of reaction?

A. Acts as a catalyst.

B. Increases the concentration.

C. Increases number of molecules.

D. Makes collisions more frequent and harder.


Why does a catalyst increase the rate of reaction?

A. Provides a route with a lower activation energy.

B. Helps provide energy for the reaction.

C.Increases the speed of reactant molecules.

D. Reduces the number of molecular collisions.


Why do most reactions start fast and get slower and slower?

A. They run out of energy.

B. They run out of catalyst.

C. The concentration of reactant molecules gets less and less.

D. The surface area increases.

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