factors affecting rate of a chemical reaction Archives

Explain the effect of concentration on the rate of reaction?

December 1, 2020 by Veerendra

Explain the effect of concentration on the rate of reaction?

Effect of concentration on the rate of reaction:

When the concentration of a reactant increases, the rate of reaction also increases.
(a) For example, two sets of experiments are carried out using the reacting conditions below:
Set I: 1 g of zinc powder and 20 cm³ of 0.4 mol dm^–³ hydrochloric acid at room temperature.
Set II: 1 g of zinc powder and 20 cm³ of 2 mol dm^–³ hydrochloric acid at room temperature.
(b) The rate of reaction of set I is higher than that of set II.
(c) This is because the concentration of hydrochloric acid used in set I is higher than that in set II.
When investigating experimentally the effect of concentration on the rate of reaction,
- the experiment is repeated a few times, each time using a different concentration of a reactant
- all the other factors/conditions are kept constant in all the experiments.

Other suitable reactions to study the effect of concentration on the rate of reaction.

The effect of concentration of a dilute acid on the rate of reaction involving the liberation of a gas such as the reaction between a reactive metal (Mg/Al/Zn/Fe) and a dilute acid (HCl/H₂SO₄) to liberate hydrogen gas, a carbonate salt and a dilute acid to liberate carbon dioxide gas can also be investigated experimentally.
- The experiment is carried out twice by changing the concentration of the dilute acid.
- The curves for the volume of gas liberated against time for both sets of the experiments are plotted on the same axes.
- The gradients of the curves are compared to deduce the difference in rates.
Different shapes of curves may be obtained depending on the volume and concentration of the dilute acid used.

Table illustrates some examples.
Explain the effect of concentration on the rate of reaction 1

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How does concentration affect the rate of reaction experiment

Aim: To investigate the effect of concentration on the rate of reaction.
Problem statement: How does the concentration of a reactant affect the rate of reaction?
Hypothesis: When the concentration of a reactant increases, the rate of reaction becomes higher.
Variables:
(a) Manipulated variable : Concentration of sodium thiosulphate solution
(b) Responding variable : Rate of reaction
(c) Controlled variables : Temperature, total volume of the reacting mixture, concentration and volume of sulphuric acid, size of conical flask
Operational definition:
Rate of reaction is inversely proportional to the time taken for the mark ‘X’ to disappear from sight.
Explain the effect of concentration on the rate of reaction 3
Materials: 0.2 mol dm^–³ sodium thiosulphate solution, 1.0 mol dm^–³ sulphuric acid, distilled water, white paper with a mark ‘X’ at the centre.
Apparatus: 150 cm³ conical flasks, 50 cm³ measuring cylinder, 10 cm³ measuring cylinder, digital stopwatch (electronically operated with an accuracy of 0.01 s).
Procedure:

45 cm³ of 0.2 mol dm^–³ sodium thiosulphate solution is measured using a 50 cm³ measuring cylinder and poured into a conical flask.
The conical flask is placed on top of a piece of white paper with a mark ‘X’ at the centre.
5 cm³ of 1.0 mol dm^–³ sulphuric acid is measured using a 10 cm³ measuring cylinder.
The sulphuric acid is then poured quickly and carefully into the conical flask and a stopwatch is started immediately.
The mixture in the conical flask is swirled a few times. The conical flask is then placed back on the white paper.
The mark ‘X’ is viewed vertically from the top through the solution, as shown in Figure.
The stopwatch is stopped immediately once the mark ‘X’ disappears from sight.
The time t required for the mark ‘X’ to disappear from sight is recorded.
The experiment is repeated four more times using different volumes of 0.2 mol dm^–³ sodium thiosulphate solution diluted with different volumes of distilled water, as shown in Table.
The results are recorded.

Results:
Explain the effect of concentration on the rate of reaction 5

Interpreting data:
Based on the results obtained, two graphs are plotted.
(a) Graph I: Graph of the concentration of sodium thiosulphate solution against time, as shown in figure.
(b) Graph II: Graph of the concentration of sodium thiosulphate solution against 1/time, as shown in figure.
Explain the effect of concentration on the rate of reaction 7

Discussion:

Sodium thiosulphate solution reacts with dilute sulphuric acid at a very low rate to form a yellow precipitate of sulphur. The chemical equation for the reaction is:
In this experiment, the time taken for the formation of a fixed quantity of sulphur to cover the mark ‘X’ until it disappears from sight can be used to measure the rate of reaction.
(a) Based on graph I, it can be seen that as the concentration of sodium thiosulphate solution decreases, a longer time is needed for mark ‘X’ to disappear from sight.
(b) Hence, it infers that as the concentration of sodium thiosulphate solution becomes lower, the rate of reaction also decreases.
(a) Graph II is a straight line. Thus, it implies that the concentration of sodium thiosulphate solution is directly proportional to 1/time.
(b) Since rate of reaction is directly proportional to 1/time, it can be deduced that:
Rate of reaction is directly proportional to the concentration of sodium thiosulphate solution.
(c) In other words:
When the concentration of a reactant increases, the rate of reaction becomes higher.
(a) In this experiment, conical flasks of the same shape and size are used.
(b) If the 150 cm³ conical flasks are replaced by bigger 250 cm³ conical flasks, the time taken for the mark ‘X’ to disappear from sight becomes longer.
(c) This is because the base area of the 250 cm³ conical flask is bigger and the depth of the 50 cm³ solution becomes shallower. Hence, a bigger amount of sulphur precipitate is required to cause the mark ‘X’ to disappear from sight.
(a) in this experiment, if hydrochloric acid of the same concentration is used to replace sulphuric acid, the rate of reaction will become lower.
(b) This is because hydrochloric acid is a strong monoprotic acid, whereas sulphuric acid is a strong diprotic acid. Hence, the concentration of hydrogen ions in hydrochloric acid is only half the concentration of hydrogen ions in sulphuric acid.
Total volume of the reacting mixture in all the five sets of the experiments are the same, that is, 50 cm³. This is to ensure that the quantity of sulphur required for the mark ‘X’ to disappear from sight is the same for all the five sets of the experiments.

Conclusion:
When the concentration of a reactant increases, the rate of reaction also increases. Hence, the hypothesis can be accepted.

What is the effect of a catalyst on the rate of a reaction?

December 1, 2020 by Veerendra

What is the effect of a catalyst on the rate of a reaction?

Effect of catalyst on the rate of reaction:

A catalyst is a substance which can alter the rate of a chemical reaction while itself remains chemically unchanged at the end of the reaction.
(a) Catalysts can be classified into positive catalysts and negative catalysts (inhibitors).
(b) A positive catalystis a catalyst that increases the rate of a reaction.
(c) A negative catalyst is a catalyst that decreases the rate of a reaction.

Table shows a few examples of catalysed reactions.

Type of reaction	Catalyst used	Effect of catalyst on the rate of reaction
Reaction between zinc and sulphuric acid Zn(s) + H₂SO₄(aq) → ZnSO₄(aq) + H₂(g)	Copper(II) sulphate solution	Rate of reaction increases. (Positive catalyst)
Decomposition of hydrogen peroxide solution 2H₂O₂(aq) → 2H₂O(I) + O₂(g)	Manganese(IV) oxide/ lead(II) oxide	Rate of reaction increases. (Positive catalyst)
Decomposition of sodium chlorate(l) solution 2NaOCl(aq) → 2NaCl(aq) + O₂(g)	Manganese(IV) oxide	Rate of reaction increases. (Positive catalyst)
Decomposition of hydrogen peroxide solution 2H₂O₂(aq) → 2H₂O(l) + O₂(g)	Propane-1,2,3-triol (glycerine)	Rate of reaction decreases. (Negative catalyst)

A catalyst does not change the quantity of the products formed. It only alters the rate of the reaction.
An example of the effect of catalyst on the rate of reaction is shown in table.

What is the effect of a catalyst on the rate of a reaction 1

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Characteristics of catalysts:
Catalysts exhibit the following characteristics:

Catalyst remains chemically unchanged at the end of a reaction. It means that the quantity and chemical composition of the catalyst remain unchanged at the end of the reaction.
Catalyst alters the rate of a chemical reaction. Positive catalyst increases the rate of reaction. Negative catalyst decreases the rate of reaction.
Catalyst does not change the quantity of products formed. The amount of products remains the same with or without the catalyst.
Catalyst is specific in its action. It can only catalyse a particular reaction but not other reactions.
Only a small amount of catalyst is needed to achieve a big increase in the rate of reaction.
An increase in the amount of catalyst used will cause only a very slight increase in the rate of reacton. Hence, it is not necessary to use a large amount of catalyst in a reaction.
Finely divided (powdered) catalyst is more effective than lump catalyst.
Catalyst used can be in the solid, liquid, gas or aqueous state. A solid catalyst may undergo physical changes. For example, crystals may turn into powder during the reaction.

Most catalysts are transition elements or compounds of transition elements such as iron, platinum, nickel, vanadium(V) oxide, copper(II) sulphate and manganese(IV) oxide.

Effect of pressure on the rate of‘reaction

For reactions with reactants in the form of solid, liquid and aqueous solution, the rates of the reactions do not depend on the pressure.
Only for those reactions that involve gaseous reactants, the rates of the reactions will increase when the pressure increases or vice versa.
An example is illustrated by the combustion of petrol in the car engine as explained below:
- Petrol is a liquid mixture of hydrocarbons.
- In the car engine, the mixture of petrol vapour and air is compressed to a high pressure so that the mixture can be ignited rapidly by the spark plug.
- The high pressure results in a high rate of combustion of petrol.

Effect of catalyst on rate of reaction experiment 1

Aim: To investigate the effect of catalyst on the rate of reaction.
Problem statement: How does a catalyst affect the rate of reaction?
Hypothesis: The presence of a catalyst will increase the rate of reaction.
Variables:
(a) Manipulated variable : Presence or absence of a catalyst
(b) Responding variable : Rate of reaction
(c) Controlled variables : Volume and concentration of hydrogen peroxide solution, temperature
Operational definition:
The decomposition of hydrogen peroxide is fast if the wooden splint rekindles brightly and rapidly. The decomposition of hydrogen peroxide is slow if the wooden splint glows dimly and slowly.
Materials: 20-volume hydrogen peroxide solution, manganese(IV) oxide powder, filter paper, wooden splints.
Apparatus: Test tubes, retort stands and clamps, 10 cm³ measuring cylinder, filter funnel, 150 cm³ beaker, spatula, electronic balance.
Procedure:

Two test tubes are labelled I and II respectively.
5 cm³ of 20-volume hydrogen peroxide solution is measured and poured separately into test tubes I and II respectively.
1.0 g of manganese(IV) oxide powder is weighed and added to the solution in test tube II.
A glowing wooden splinter is immediately brought to the mouth of each of the test tubes.
The changes that occur are recorded in a table.
At the end of the reaction, the mixture in test tube II is filtered to separate the manganese(IV) oxide powder.
The manganese(IV) oxide powder obtained is washed with a little distilled water and pressed between sheets of filter paper.
The dry manganese(IV) oxide powder is weighed and its mass is recorded.

Results:

Test tube	Observation
I	The wooden splint glows dimly and slowly. No effervescence occurs.
II	The wooden splint rekindles brightly and rapidly. Effervescence occurs.

Mass of manganese(IV) oxide before reaction = 1.0 g
Mass of manganese(IV) oxide after reaction = 1.0 g

Inferences:

Rate of decomposition of hydrogen peroxide increases in the presence of manganese(IV) oxide powder.
Mass of manganese(IV) oxide powder remains unchanged during the reaction.
Manganese(IV) oxide acts as a catalyst.

Discussion:

(a) Oxygen gas is liberated during the decomposition of hydrogen peroxide.
2H₂O₂(aq) → 2H₂O(l) + O₂(g)
(b) The liberation of oxygen gas can be confirmed if the wooden splint rekindles brightly or continues to glow.
Manganese(IV) oxide acts as a (positive) catalyst to increase the rate of decomposition of hydrogen peroxide while itself remains chemically unchanged at the end of the reaction. Hence, its mass remains unchanged at the end of the reaction.

Conclusion:
The presence of a (positive) catalyst increases the rate of a reaction. Hence, the hypothesis can be accepted.

Effect of catalyst on rate of reaction experiment 2

Aim: To investigate the effect of the amount of catalyst on the rate of reaction.
Problem statement: How does the amount of a catalyst affect the rate of a reaction?
Hypothesis: When the amount of a catalyst used increases, the rate of reaction also increases.
Variables:
(a) Manipulated variable : Amount/mass of catalyst
(b) Responding variable : Rate of reaction
(c) Controlled variables : Temperature, volume and concentration of hydrogen peroxide solution
Operational definition:
The curve for the graph of volume of gas liberated against time with a higher gradient indicates a higher rate of reaction.
Materials: 2-volume hydrogen peroxide solution, manganese(IV) oxide powder.
Apparatus: 50 cm³ measuring cylinder, 150 cm³ conical flask, stopper fitted with a delivery tube, burette, retort stand and clamp, basin, electronic balance, stopwatch, spatula, beaker.
Procedure:

A burette is filled with water until it is full. The burette is inverted over water in a basin and clamped vertically using a retort stand.
The water level in the burette is adjusted and the initial burette reading is recorded.
50 cm³ of 2-volume hydrogen peroxide solution is measured using a measuring cylinder and poured into a conical flask.
0.2 g of manganese(IV) oxide powder is weighed using an electronic balance and poured carefully into the conical flask.
The conical flask is then closed immediately with a stopper fitted with a delivery tube directed to the burette, as shown in Figure. At the same time, a stopwatch is started immediately.
The conical flask with its contents is shaken slowly and the volume of oxygen gas collected in the burette is recorded at regular time intervals of 30 seconds for 5 minutes.
Steps 1 to 6 are repeated using 0.8 g of manganese(IV) oxide powder to replace the 0.2 g of manganese(IV) oxide powder.
The results are tabulated.

Results:
Set I: Using 0.2 g of manganese(IV) oxide powder
What is the effect of a catalyst on the rate of a reaction 4
Set II: Using 0.8 g of manganese(IV) oxide powder

Interpreting data:

Based on the results obtained, the graphs of the volume of oxygen gas liberated against time for both sets I and II are plotted on the same axes.
Based on the graphs plotted, the following inferences can be made.
- Curve I has a lower gradient, thus lower rate of reaction.
- Curve Il has a higher gradient, thus higher rate of reaction.

Discussion:

The amount of manganese(IV) oxide powder (catalyst) used in set II is four times the amount of manganese(IV) oxide powder used in set I.
Based on the inferences made, it can be deduced that:
As the amount of manganese(IV) oxide powder (catalyst) used increases, the rate of reaction also increases.
There is only a very slight increase in the rate of reaction when using 4 times more manganese(IV) oxide powder as a catalyst.
(a) If the reaction is allowed to proceed until all the hydrogen peroxide is completely decomposed, the graphs in Figure will be obtained.

(b) (i) The graphs show that the maximum volume of oxygen gas liberated in both sets I and II are the same, that is, V cm³.
(ii) This is because the quantities of hydrogen peroxide solution used, in mole, in both sets I and II are the same.
(iii) The manganese(IV) oxide powder as a catalyst does not affect the amount of products formed.

Conclusion:
An increase in the amount of catalyst used will increase the rate of reaction. Hence, the hypothesis can be accepted.
Note: There is only a very slight increase in the rate of reaction when the amount of catalyst increases.

How does the temperature affect the rate of a chemical reaction?

December 1, 2020December 1, 2020 by Veerendra

How does the temperature affect the rate of a chemical reaction?

Effect of temperature on the rate of reaction:

When the temperature increases, the rate of reaction also increases.
(a) For example, two sets of experiments are carried out using the reacting conditions below:
Set I: 1 g of granulated zinc and 20 cm³ of 0.2 mol dm^–³ hydrochloric acid at 60°C
Set II: 1 g of granulated zinc and 20 cm³ of 0.2 mol dm^–³ hydrochloric acid at 30°C
(b) The rate of reaction of set I is higher than that of set II.
(c) This is because the temperature (60°C) of the reacting mixture in set I is higher than the temperature (30°C) of the reacting mixutre in set II.
When investigating experimentally the effect of temperature on the rate of reaction,
- the experiment is repeated a few times, each time using a different temperature of a reactant.
- all the other factors/conditions are kept constant in all the experiments.
The following shows an example of the effect of temperature on the rate of reaction. Table shows the method to derive the shapes of curves plotted for two sets of experiments.

How does the temperature affect the rate of a chemical reaction 1

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Effect of temperature on rate of reaction experiment

Aim: To investigate the effect of temperature on the rate of reaction.
Problem statement: How does temperature affect the rate of reaction?
Hypothesis: An increase in temperature will increase the rate of reaction.
Variables:
(a) Manipulated variable : Temperature ot sodium thiosulphate solution
(b) Responding variable : Rate of reaction
(c) Controlled variables : Volume and concentration of sodium thiosulphate solution, volume and concentration of sulphuric acid, size of conical flask
Operational definition:
Rate of reaction is inversely proportional to the time taken for the mark ‘X’ to disappear from sight.
How does the temperature affect the rate of a chemical reaction 2
Materials: 0.2 mol dm^–³ sodium thiosulphate solution, 1 mol dm^–³ sulphuric acid, white paper with a mark ‘X’ at the centre.
Apparatus: 150 cm³ conical flasks, 50 cm³ measuring cylinder, 10 cm^–³ measuring cylinder, digital stopwatch (electronically operated with an accuracy of 0.01 s), thermometer, Bunsen burner, tripod stand, wire gauze.
Procedure:

50 cm³ of 0.2 mol dm^–³ sodium thiosulphate solution is measured using a measuring cylinder and poured into a conical flask.
The temperature of this solution is measured using a thermometer.
The conical flask is placed on top of a piece of white paper with a mark ‘X’ at the centre.
5 cm³ of 1 mol dm^–³ sulphuric acid is measured using a 10 cm³ measuring cylinder.
The sulphuric acid is then poured quickly and carefully into the conical flask and a stopwatch is started immediately.
The mixture in the conical flask is swirled a few times. The conical flask is then placed back on the white paper.
The mark ‘X’ is viewed vertically from the top through the solution, as shown in Figure.
The stopwatch is stopped immediately once the mark ‘X’ disappears from sight.
The time f required for the mark ‘X’ to disappear from sight is recorded.
Steps 1 to 9 are repeated using 50 cm³ of 0.2 mol dm^–³ sodium thiosulphate solution at 35°C, 40°C, 45°C and 50°C respectively by heating the solution as shown in Figure before 5 cm³ of 1 mol dm^–³ sulphuric acid is added. All other conditions remain unchanged.
The results are recorded in a table.

Results:

Interpreting data:
1. Based on the results, two graphs are plotted.
(a) Graph I: Graph of the temperature against time
(b) Graph II: Graph of the temperature against 1/time
How does the temperature affect the rate of a chemical reaction 5
2. From graph I, it can be deduced that as the temperature increases, the time taken for the mark ‘X’ to disappear from sight becomes shorter.
3. From graph II, it can be deduced that the temperature increases linearly with 1/time.

Discussion:

(a) From the graphs plotted, the following inference can be deduced:
“Temperature increases linearly with 1/time”
(b) But, rate of reaction ∝ 1/time.
(c) Hence, rate of reaction increases linearly with temperature.
In other words, as the temperature increases, the rate of reaction also increases.
Usually, the rate of a reaction approximately doubles for every 10°C rise in temperature.
The concentration and volume of both sodium thiosulphate solution and sulphuric acid are kept constant in each set of the experiment. Any change in rate of reaction is due to the difference in temperature.

Conclusion:
Rate of reaction increases when the temperature of the reaction increases. Hence, the hypothesis can be accepted.

How does the surface area affect the rate of reaction?

December 1, 2020December 1, 2020 by Veerendra

How does the size of particles affect the rate of reaction?

Effect of surface area on the rate of reaction:

When the particle size of a fixed mass of a solid reactant becomes smaller, the total exposed surface area becomes larger, the rate of reaction increases.
- For example, two sets of experiments are carried out using the reacting conditions below:
  Set I: 1 g of zinc powder and 20 cm³ of 0.2 mol dm^–³ hydrochloric acid at room conditions.
  Set II: 1 g of granulated zinc and 20 cm³ of 0.2 mol dm^–³ hydrochloric acid at room conditions.
- The rate of reaction for set I is higher than that for set II.
- This is because 1 g of zinc powder used in set I has a larger total exposed surface area than 1 g of granulated zinc.
When investigating experimentally the effect of particle size/surface area on the rate of a reaction,
- the experiment is carried out twice
- the particle size/surface area of a solid reactant used in the two experiments are varied
- all the other factors/conditions are kept constant in both the experiments.
(a) The reaction between reactive metals and dilute acids to liberate hydrogen gas can also be used to study the effect of surface area on the rate of reaction.
Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g)
Zn(s) + H₂SO₄(aq) → ZnSO₄(aq) + H₂(g)
(b) Rate of reaction of reactive metals such as magnesium, aluminium, zinc or iron with dilute acid is affected by the size of the metals.
(c) The smaller the size of a fixed mass of magnesium, aluminium, zinc or iron, the larger the total exposed surface area and the higher is the rate of reaction with dilute hydrochloric acid or dilute sulphuric acid.
An example of the effect of size of a solid reactant on the rate of reaction is shown in Table.

Experiment	Initial rate of reaction	Maximum volume of hydrogen gas
Set I: 1 g of granulated iron + 50 cm³of 0.2 mol dm^–³sulphuric acid Set II: 1 g of iron filings + 50 cm³ of 0.2 mol dm^-3sulphuric acid	Initial rate of reaction of set II is higher than that of set I because the total exposed surface area of 1 g of iron filings is larger than that of 1 g of granulated iron.	The quantities of iron and sulphuric acid used, in mole, in both sets I and II are the same. Thus, the maximum volume of hydrogen gas collected in both sets I and II are also the same.

Graph:

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How surface area affects the rate of reaction experiment

Aim: To investigate the effect of total exposed surface area of a solid reactant on the rate of reaction.
Problem statement: How does the total exposed surface area of a solid reactant affect the rate of reaction?
Hypothesis: When the total exposed surface area of marble chips increases, the rate of reaction increases.
Variables:
(a) Manipulated variable : Total exposed surface area of marble chips
(b) Responding variable : Rate of reaction
(c) Controlled variables : Mass of marble, volume and concentration of hydrochloric acid, temperature
Operational definition:

Smaller marble chips have a larger total exposed surface area than larger marble chips of the same mass.
For the graph of the volume of gas liberated against time, the curve with higher initial gradient indicates a higher initial rate of reaction.

Materials: 0.1 mol dm^–³ hydrochloric acid, large marble chips, small crushed marble chips, water.
Apparatus: 100 cm³ measuring cylinder, 150 cm³ conical flask, stopper fitted with a delivery tube, basin, burette, electronic balance, stopwatch.
Procedure:

A burette is filled with water until it is full. It is then inverted over water in a basin and clamped vertically using a retort stand.
The water level in the burette is adjusted and the initial burette reading is recorded.
40 cm³ of 0.1 mol dm^–³ hydrochloric acid is measured using a measuring cylinder and poured into a conical flask.
The apparatus as shown in Figure is set up.
2 g of large marble chips is weighed using an electronic balance and added to the acid in the conical flask. The conical flask is closed immediately with a stopper fitted with a delivery tube. At the same time, a stopwatch is started.
The carbon dioxide gas released is collected in the burette by downward displacement of water as shown in Figure.
The conical flask is shaken gently throughout the whole experiment.
The volume of gas collected in the burette is recorded at regular intervals of 30 seconds until no more gas is liberated.
Steps 1 to 8 are repeated using 2 g of small crushed marble chips to replace 2 g of large marble chips. All the other conditions remain unchanged.

Results:
Using large marble chips:
How does the surface area affect the rate of reaction 4
Using small crushed marble chips:

Interpreting data:
1. The graphs of volume of carbon dioxide gas liberated against time for both sets of the experiments are plotted on the same axes, as shown in Figure.
How does the surface area affect the rate of reaction 6

2. To calculate the instantaneous rate of reaction
How does the surface area affect the rate of reaction 7
(c) By comparison, the instantaneous rate of reaction at 1 minute of the experiment using small crushed marble chips is higher than that of the experiment using large marble chips.

3. To calculate the overall average rate of reaction
How does the surface area affect the rate of reaction 8
(c) By comparison, the overall average rate of reaction of the experiment using small crushed marble chips is higher than that of the experiment using large marble chips.

Discussion:

The reaction between marble chips and hydrochloric acid releases carbon dioxide gas. The chemical equation for the reaction is:
CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + CO₂(g) + H₂O(l)
- Based on the graphs plotted, the curve for the experiment using small crushed marble chips is steeper than the curve for the experiment using large marble chips.
- In other words, the initial gradient of the curve for the experiment using small crushed marble chips is higher than the initial gradient of the curve for the experiment using large marble chips.
- Thus, the inital rate of reaction for the experiment using small crushed marble chips is higher than that for the experiment using large marble chips.
- Small crushed marble chips have a larger total exposed surface area than the large marble chips with the same mass.
- Hence, it can be concluded that:
  When the size of a fixed mass of solid marble becomes smaller, the total exposed surface area becomes larger, hence the rate of reaction increases.
- The above graph shows that the maximum volume of carbon dioxide gas collected in both experiments are the same, that is, 43.0 cm³.
- This is because the quantities of hydrochloric acid and marble used, in mole, are the same in both experiments.
- Theoretical maximum volume of carbon dioxide gas that should be collected can be calculated as follows:
  Number of moles of hydrochloric acid used
- The maximum volume of carbon dioxide gas collected in the experiment is less than the theoretical volume because a small volume of the carbon dioxide has dissolved in the water when it is collected in the burette.
- A method to overcome this difference is by saturating the water with carbon dioxide gas before collecting the gas in the burette.

Conclusion:
When the total exposed surface area of a solid reactant increases, the rate of reaction also increases. Hence, the hypothesis can be accepted.

What factors affect the rate of a reaction?

December 1, 2020December 1, 2020 by Veerendra

What factors affect the rate of a reaction?

Factors Affecting the Rate of Reaction:

The rate of a chemical reaction can be altered by changing the reacting conditions.
There are five factors that affect the rate of a reaction:

Total surface area of a solid reactant or particle size of a solid reactant
Concentration of a reactant in solution
Temperature of a reaction
Use of catalyst
Pressure for a reaction involving gaseous reactants

Table shows some examples of changing the reacting conditions which affect the rates of reactions in our daily occurrences.

Example	Explanation	Factor affecting the rate
(a) It is easier to start a fire using sticks rather than logs	This is because sticks have a bigger total exposed area than the logs. Hence, the rate of burning the sticks is higher than that of the logs.	Surface area (or particle size)
(b) Blowing air over glowing charcoal causes it to ignite brightly and quickly	This is due to increase in the concentration of oxygen/air. So, the rate of burning increases.	Concentration
(c) Milk turns sour faster on hot days than on cold days	On hot days, the temperature is higher. So, the rate of the milk turns sour is higher.	Temperature
(d) The presence of amylase in the saliva breaks starch to maltose	Amylase, as a biological catalyst/enzyme, speeds up the breakdown of starch to maltose.	Catalyst
(e) The manufacture of ammonia gas is carried out at high pressure N₂(g) + 3H₂(g) → 2NH₃(g)	At high pressure of the gaseous reactants, the rate of reaction is higher.	Pressure

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Application of factors that affect the rate of reaction

Application of the factors affecting the rate of reaction in daily life:

Knowledge of the factors affecting the rate of reaction can be applied to control the rates of many reactions that occur in our daily life.
These can be done by either speeding up or slowing down the reaction to suit our needs.

Table shows a few examples of applying the factors affecting the rate of reaction to control the rates of a few reactions in our daily life.

Example	Explaining the effect on the rate of reaction	Factor affecting the rate of reaction
(a) Burning of charcoal	Food can be cooked faster when smaller pieces of charcoal are used compared to bigger pieces of charcoal. This is because smaller pieces of charcoal have a larger total exposed surface area compared to bigger pieces. Hence, smaller pieces of charcoal can burn faster to produce more heat per second and the food is cooked faster.	Surface area of charcoal that controls its rate of burning
(b) Storing of food in a refrigerator	The decaying and decomposition of food by microorganisms are chemical reactions occurring in food that cause the food to turn bad. When the food is kept in a refrigerator, the food lasts longer. This is because the low temperature in the refrigerator slows down the decaying and decomposition of food by microorganisms.	Temperature that controls the rate of decaying of food by microorganisms
(c) Cooking of food in a pressure cooker	In a pressure cooker, the high pressure raises the boiling point of water to a temperature above 100°C. Thus, cooking is carried out at a temperature higher than 100°C in the pressure cooker. The higher temperature increases the rate of cooking. Hence, the food can be cooked faster.	Temperature that controls the rate of cooking
(d) Cooking of solid food with different sizes	Solid food in big chunks such as big chunks of potato have a smaller total surface area exposed to heat during cooking. Hence, the food will be cooked at a lower speed. For the food to be cooked faster, the big chunks of solid food must be cut into smaller pieces such as potato chips to increase its total surface area exposed to heat.	Surface area that controls the rate of cooking

Application of the factors affecting the rate of reaction in industrial processes

Knowledge of the factors affecting the rate of reaction is applied in choosing the optimum conditions to run an industrial process to achieve
(a) shorter time of production
(b) higher yield
(c) lower cost of production
It is very costly to run an industrial process at a very high temperature and pressure.
Catalysts are generally used to increase the rates of these processes.
The catalysts used enable the processes to be run in a shorter time at an optimum temperature and pressure to reduce the cost of production.

Three examples of industrial processes are:
(a) Haber Process

Haber process is an industrial process to manufacture ammonia on a large scale. In this process, nitrogen reacts with hydrogen to produce ammonia.
Table describes the effect of catalyst, temperature and pressure on the rate of production of ammonia and its yield.

Factors	Effect on the rate of reaction	Effect on the yield of ammonia	Conditions chosen
Catalyst	The use of catalyst increases the rate of production of ammonia	The quantity of ammonia (yield) produced remains the same with or without the catalyst	Iron is used as a catalyst to increase the rate of production of ammonia
Temperature	Increase in temperature will increase the rate of production of ammonia	Increase in temperature will decrease the quantity of ammonia (yield) produced	An optimum temperature of 450 to 550°C is chosen. A compromise is reached for a reasonable rate and yield.
Pressure	Increase in pressure will increase the rate of production of ammonia	Increase in pressure will increase the quantity of ammonia (yield) produced.	A high pressure of 200 to 300 atm is chosen to obtain a higher rate and yield.

So, Haber process is run in the industry as below:
In the Haber process, a mixture of nitrogen gas and hydrogen gas in the volume ratio of 1:3 is passed over iron powder (catalyst) at a temperature of 450 to 550°C and a pressure of 200 to 300 atm in the presence of molybdenum as a promoter.
These optimum conditions enable the Haber process to run at a higher rate to cut down the cost and time of production, and also to obtain a higher yield.

(b) Ostwald Process

Ostwald process is an industrial process to manufacture nitric acid on a large scale.
Ostwald process involves three stages, as shown in Table.

Stage	Reacting conditions	Chemical equation for the reaction
I	Ammonia gas is mixed with excess air and the mixture is passed over platinum (catalyst) at a temperature of 850°C and a pressure of 2 to 5 atm. Ammonia is oxidised to nitrogen monoxide. All these reacting conditions are chosen to ensure the reaction occurs at an optimum rate with an optimum yield.	4NH₃(g) + 5O₂(g) → 4NO(g) + 6H₂O(l)
II	Nitrogen monoxide gas from stage I is then reacted with excess air to produce nitrogen dioxide gas.	2NO(g) + 0₂(g) → 2N0₂(g)
III	Nitrogen dioxide gas from stage II is then mixed with excess air and dissolved in hot water at 80°C to produce nitric acid	4NO₂(g) + O₂(g) + 2H₂O(l) → 4HNO₃(aq)

(c) Contact process

Contact process is an industrial process to manufacture sulphuric acid on a large scale.
Contact process involves four stages, as shown in Table.

Stage	Reacting conditions	Chemical equation for the reaction
1	Sulphur is burnt in excess air to produce sulphur dioxide gas	S(s) + O₂(g) → SO₂(g)
II	Sulphur dioxide gas from stage 1 is mixed with excess air and the mixture is passed over vanadium(V) oxide, V₂O₅ (catalyst) at a temperature of 450°C and a pressure of 1 atmosphere. Sulphur dioxide is oxidised to sulphur trioxide. All these reacting conditions are chosen to ensure the reaction occurs at an optimum rate with an optimum yield.	2SO₂(g) + O₂(g) → SO₃(g)
III	Sulphur trioxide gas from stage II is then dissolved in concentrated sulphuric acid to form a liquid called oleum, H₂S₂O₇.	SO₃(g) + H₂SO₄(aq) → H₂S₂O₇(l)
IV	The oleum from stage III is then diluted with water to produce concentrated sulphuric acid of about 98%.	H₂S₂O₇(l) + H₂0(l) → 2H₂SO₄(aq)

Factors affecting the rate of reaction problems with solutions

1. Table shows the reacting conditions for three sets of experiments.

Set	Reacting condition
I	5 g of excess zinc powder + 100 cm³ of 0.4 moi dm^-3 sulphuric acid + 5 cm³ of 1 mol dm^–³ copper(ll) sulphate solution at room temperature
II	5 g of excess zinc powder + 50 cm³ of 0.4 mol dm^-3 sulphuric acid at room temperature
III	5 g of excess granulated zinc + 75 cm³ of 0.4 mol dm^–³ sulphuric acid at room temperature

Sketch the graphs of the volume of hydrogen gas liberated against time for sets I, II and III on the same axes. Explain how you obtain your answers.
Solution:
1. Compare the yield
What factors affect the rate of a reaction 2
From the table, number of moles of sulphuric acid reacted in set I > set III > set II.
Thus, the maximum volume of hydrogen gas liberated in set I > set III > set II.

2. Compare the rate

Initial rate of reaction for set I > set II because of the presence of catalyst in set I.
Initial rate of reaction for set II > set III because the total surface area of zinc used in set II is larger than that in set III.
Hence, the initial rate of reaction for set I > set II > set III
This means that the initial gradient of the curve for set I > set II > set III.

3. Sketch the graph
The graphs for the three sets of the experiments are:
What factors affect the rate of a reaction 4

2. 5 g of excess marble chips is added to 80.0 cm³ of 1.0 mol dm^–³ acid R in a conical flask. The conical flask is placed on an electronic balance. The mass of the conical flask and its contents is recorded at regular time intervals. Figure shows the graph of the mass of the conical flask and its contents against time.
What factors affect the rate of a reaction 5
(a) Write the chemical equation for the reaction if acid R is ethanoic acid.
(b) Why does the curve become horizontal after 180 s?
(c) (i) Describe the change in the rate of reaction during the first 180 s. Explain how you derive your answer from the graph.
(ii) Explain the change in the rate of reaction in (c)(i).
Solution:

(a) CaCO₃(s) + 2CH₃COOH(aq) → Ca(CH₃COO)₂(aq) + CO₂(g) + H₂O(l)
(b) The reaction has stopped after 180 s because all the acid R has reacted completely with excess marble.
(c) (i) Rate of reaction decreases with time during the first 180 s. This is because the magnitude of the gradient of the curve decreases with time.
(ii) The concentration of acid R and the total surface area of marble chips decrease with time, hence the rate of reaction decreases with time.

3. 1 g of excess magnesium powder is added to 50 cm³ of 0.5 mol dm^–³ hydrochloric acid at room temperature. Figure shows the curve obtained when the volume of hydrogen gas liberated against time is plotted.
What factors affect the rate of a reaction 6
The experiment is repeated using 25 cm³ of 0.5 mol dm^–³ hydrochloric acid at 80.0°C to replace 50 cm³ of 0.5 mol dm^–³ hydrochloric acid at room temperature. Copy the graph and sketch on the same axes the curve that you would expect to obtain for the second experiment. Explain how you obtain your answer.
Solution:
1. Compare the maximum volume of hydrogen gas liberated

From the above table, the number of moles of hydrochloric acid reacted in experiment I is twice the number of moles of hydrochloric acid reacted in experiment II.
Thus, the maximum volume of hydrogen gas liberated in experiment I is twice the maximum volume of hydrogen gas liberated in experiment II.
Hence, the maximum volume of hydrogen gas liberated in experiment II
= 1/2 x 600 cm³
= 300 cm³

2. Compare the initial rate

From the above table, the temperature of experiment II is higher than that of experiment I, whereas the other two factors remain constant.
Hence, the initial rate of reaction of experiment II is higher than that of experiment I.
In other words, the initial gradient of curve II is steeper than that of curve I.

3. Sketch the curve
Based on the above analysis, the curve for experiment II is shown below.
What factors affect the rate of a reaction 9

4. Experiment I is carried out to study the rate of reaction between 1.0 g of iron foil with 60 cm³ of 0.4 mol dm^–³ hydrochloric acid. The graph in Figure is obtained.
What factors affect the rate of a reaction 10
The experiment is repeated two more times using different volumes and concentrations of hydrochloric-acid as shown in Table.

(a) Write the chemical equation for the reaction.
(b) (i) Compare the initial rates of reaction for experiments I, II and III.
(ii) Give reasons for your answers.
(c) Determine the maximum volume of hydrogen gas liberated in experiments II and III.
Given that the relative atomic mass of Fe = 56 and the molar volume of any gas is 24 dm^–³ mol^–¹ at room conditions.
(d) By using the answers in (b) and (c), sketch the graphs of the volume of hydrogen gas liberated against time for experiments II and III on the same axes in Figure.
Solution:
(a) Fe(s) + 2HCl(aq) → FeCl₂(aq) + H₂(g)
(b) (i) Initial rate of reaction for experiment III > experiment I > experiment II
(ii) This is because the concentration of hydrochloric acid in experiment III > experiment I > experiment II.
What factors affect the rate of a reaction 12