topic 5 speed of chemical reactions.pdf

8/13/2019 TOPIC 5 SPEED OF CHEMICAL REACTIONS.pdf

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INTRODUCTION

If there is a running competition between a rabbit and a tortoise, which animal will win? Surely the answer will be the rabbit (if the rabbit does not fall asleep during the competition, that is). Rabbits run faster than tortoises. The tortoise will get to thefinish line eventually, but will probably reach there muchlater.This means that the rabbit runs ata greater speed than the tortoise.

In everyday life, if you put granulated sugar and fine sugar in different glasses of water with the same volume and temperature, which sugar will dissolve first?

TT oopp iicc 55 Speed Of Chemical Reactions

Yes! fine sugar will dissolvefirst. It is because fine sugarhas a larger surface area that comes in contact with water.

LEARNING OUTCOMES

By the end of this topic, you should be able to:1. Define the speed of chemical reaction;2. Calculate the speed of a chemical reaction;3. Distinguish the effects of particle size, concentration, pressure,

temperature and catalysts on the speed of chemical reaction; and

4. Evaluate the effect of activation energy on the speed of a reaction.

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TOPIC5 SPEED OF CHEMICAL REACTIONS4

5.1.2 Nature of Chemical Reactants

In order for a reaction to occur, there must be a collision between the reactants at the reactive site of the molecule with correct orientation and it has to achieve activation energy. This will lead to effective collision and chemical reaction will occur.

Figure 5.3: Particles showing the effective and ineffective collision [Source : http://2012books.lardbucket.org/books/principles of general

chemistry v1.0m/s18 07 the collision model of chemica.html ]

Particles might be atoms, molecules or ions. Before we can get a chemical reaction, particles must crash together. They must collide. This is called the collision theory .

Figure 5.4:Collisionbetween particles [Source : http://minhaji.net/classes/ 3107]

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TOPIC 5 SPEED OF CHEMICAL REACTIONS 5

5.1.3 Speed of Chemical Reaction

The area of chemistry concerned with the speed or rates at which a chemical reaction occurs is called chemical kinetics. The word kinetic suggests motion. Here, kinetics refers to the speed of a reaction, or the reaction speed , which is the change of the concentration of reactant or product with time.

Let us look at the general equation :

Reactants Products

This equation tells us that, during the course of a reaction, reactant molecules are consumed while product molecules are formed. Two obvious

changes will

occur,

namely:

i. The decrease in the quantity of a reactant with time; and

ii. The increase in the quantity of a product with time.

As a result, we can follow the progress of a reaction by monitoring:

i. Either the decrease in concentration of the reactants or the

increase in concentration of the products;

ii. Decrease in the mass of reactant or increase in the mass of

product;

iii. Increase in the volume of gas released;

iv. Formation of precipitate as a product; or

v. Change in pH, temperature or electrical conductivity.

For reactions that occur rapidly, the speed of reaction is high. Conversely,for a reaction that occurs slowly,the speed of reaction is low. The time taken for a fast reaction is short, whereas the time taken for a slow reaction is long.

How do wemeasure the speed of chemical reaction?

Speed of chemical reaction is the speed at which reactantsare converted into the products in a chemical reaction.

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The chemical equation for the reaction between marble chip (calcium carbonate, CaCO 3) and hydrochloric acid is:

CaCO 3(s) + HCl (aq) CaCl 2(aq) + CO 2(g) + H 2O (l)

Figure 5.6 shows the volume of carbon dioxide gas released measured at certain intervals plotted against time.

Figure 5.6:The volume of carbon dioxide gas liberated against time

How fast areaction progresses over an interval of time is the average speed of reaction. It is calculated as follows:

Average speed= The change in the amount of reactant or product

The time taken for the change to happen

From the graph in Figure 5.6, we can calculate the average speed of chemical reaction between marble chip and hydrochloric acid.

Volume of CO2 gas/ cm3

Time/min

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Average speed of reaction = The total volume of carbon dioxide gas released Time taken for the total carbon dioxide gas release

= 94.00cm 3 4.5 min

= 20.90cm 3min 1

Can you calculate the speed of reaction at any given time?

Let us take a look at the next example:

Based on the graph of volume of carbon dioxide gas liberated against time (Figure 5.6), you can also:

a. Calculate the average rate of reaction in the first one minute;

Figure 5.7: The average rate of reaction in the first one minute

The exact speed of reaction at any given time is called the instantaneous speed of reaction .

Volume of CO 2 gas/ cm3

Time/min

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The average rate of reaction in the first one minute

= Total volume of CO 2 collected in the first 1 minute Time taken

= 54.00cm3 1 min

= 54.00cm 3min 1

b. Calculate the average rate of reaction from 1 minute to 2 minutes; and

Figure 5.8:The average rate of reaction from 1 minute to 2 minute

The average rate of reaction from 1 minute to 2 minutes = Total volume of CO 2 collected from 1 minute to 2 minutes

Time taken

= (77.00 54.00)cm3

(2 1) min = 23.00cm3

1 min = 23.00cm3min 1


Time/min

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c. Calculate the average speed of reaction at the 2 minutespoint by drawinga tangent at the curve point.

Figure 5.9:The average speed of reaction calculatedby drawing a tangent line at the curve point

The speed of reaction at the 2nd minute = The gradient of the tangen t of the graph at the second minute

The speed of reaction at the 2nd minute = 100.00 50.00cm3

3.3 0.4 min = 50.00cm3

2.9 min = 17.24cm 3min 1


Time/min

ACTIVITY 5.3

From the graph in Figure 5.8, calculate: The average speed of reaction in 3 minutes. The average speed of reaction from 3 minutes to 4 minutes. The average speed of reaction from 2 minutes to 4.5 minutes.

Tangentline

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If you want to produce as much of a product as possible with the shortest amount of timeviaa chemical reaction, you must consider the kinetics of the reaction.

5.3.1 Effect of Particle Size of Chemical Reactants

Reaction depends on collisions. The more surface area on which collisions can occur, the faster the reaction.

You can hold a burning match to a large chunk of coal and nothing will happen. But if you take that same piece of coal, grind it up very, very fine, throw it up into the air, and strike a match, youll get an explosion because of the increased surface area of the coal.

We find that small pieces of solids, especially powders, react faster than larger pieces. It is like frying two pans of chips! One has the potato cut into small, thin chips. The other pan has bigger, thicker chips (Figure 5.10).Which chips do you think will be cooked first?Which chips have the larger surface area?

Surface area is a measure of how much surface is exposed. So for the same massof potato, small chips have a larger surface area than big chips.

Figure 5.10: Small chips with larger surface area

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Figure 5.11: (a) Bigger sized reactant; (b) Smaller sizedreactant

The smaller the size of reactant, the larger is the surface area exposed. This translates to an increase to the speed of chemical reaction.

5.3.2 Effect of Concentration of Chemical Reactants

Increasing the

number

of

collisions

will

speed

up

the

reaction

rate.

The

more

reactant molecules there are colliding, the faster the reaction will be. As the concentration becomes higher, the numberof molecules perunit volume also increases (Figure 5.12). For example, a wood splint burns moderately in the air (20 percent oxygen), but it burns much faster in pure oxygen.

(a) (b)

Pour 800ml of water in two different pots.Put 1kg of whole chicken without cutting it into the first pot and in another pot put another 1kg of chiken that had been cut in eight.Which pot of chicken will be done first? Explain why.

ACTIVITY 5.4

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Figure 5.12:(a) Low concentration; (b) High concentration of reactant

In most simple cases, increasing the concentration of the reactants increases the speed of reaction. However, if the reaction is complex and has a complex mechanism (series of steps in the reaction), this may not be the case. Determining the concentration effect on the speed of reaction can give you clues as to which reactant is involved in the rate, thus determining the step of the mechanism.

You can do this by testingthe reaction withseveral different concentrations

and observing the effect on the speed of reaction as in Experiment 5.3.

(a) (b)

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Experiment 5.3

1. Using a pencil, mark an X on a piece of white paper, as follows:

2. Using the 50cm measuring cylinder, measure 50cm of 0.2 mol dm

sodium thiosulphate solution and pour it into a conical flask. Place

the flask on the X mark on the white paper.

3. Measure 5cm of 1 mol dm sulphuric acid with a 10cm measuring

cylinder.

4. Immediately, pour the sulphuric acid into the conical flask

containing 50cm of sodium thiosulphate solution and shake the flask. At the same time, start the stopwatch.

5. Observe the yellow precipitate of sulphur at the top part of the

conical flask. Record the time when the X mark on the white

paper is no longer visible.

6. Repeat the experiment using 50cm of the 0.4 mol dm , 0.6mol

dm , 0.8 mol dm and 1.0 mol dm sodium thiosulphate

solution.The volume and concentration of the sulphuric acid used

are the same.

7. Plot two graphs:

a) Graph of concentration of sodium thiosulphate solution against

time.

b) Graph of the concentration of sodium thiosulphate solution

against 1 Time

8. Calculate the average speed of reaction for all the experiment. What

can be represented by 1 Time?

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b) Graph of concentration of sodium thiosulphate solution against time

Figure 5.15: Graph of concentration of sodium thiosulphate solution against time

From Experiment 5.3, 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 speed of reaction.

Speed of reaction is directly proportional to:

1 time taken for the mark X to disappear from sight

Concentrationof sodiumthiosulphatesolution(mol/dm 3)

1Time (s-1)

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5.3.3 Effect of Pressure of Gaseous Reactants

The pressure of gaseous reactants has basically the same effect as concentration. The higher the reactant pressure, the higher the reaction speed. This is due to the increased number of collisions (Figure 5.16).

Figure 5.16: (a) Low pressure ; (b) High pressure

5.3.4 Effect of Temperature

Increasing the temperature causes molecules to move faster, so there is an increased chance of them colliding with each other and reacting. But increasing the temperature also increases the average kinetic energy of the molecules.

Figure 5.17shows an example of how increasing the temperature affects the kinetic energy of the reactants and increases the reaction speed.

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the minimum amount of kinetic energy needed, which means a lot more collisions will be energetic enough to lead to reaction.

Increasing the temperature not only increases the number of collisions but also increases the number of collisions that are effective that transfer enough energy to cause a reaction to take place (Figure 5.18).

Figure 5.18: Effect of temperature on the reaction between particle A and particle B

Design and carry out an experiment to study the effect of temperature on the rate of reaction. The various temperatures that are suggested for this experiment are 30C, 35C, 40C, 45C and 50C. The materials and apparatus supplied are as shown in the following:

Materials: 1 mol dm sulphuric acid, H 2SO4 , 0.2 mol dm sodium thiosulphate solution, Na 2S2O4 , white paper.

Apparatus: 100cm conical flask, 50cm measuring cylinders,

stopwatch, Bunsen burner, wire gauze, tripod stand, thermometer.

Calculate the rate of reaction at the third minute.

ACTIVITY 5.5

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The catalyst may react to form an intermediate, but it is regenerated in a subsequent step of the reaction. In the laboratory preparation of molecular oxygen, a sample of potassium chlorate is heated, as shown in Figure 5.19, andthe reaction is noted as follows:

2KCIO (s) 2KCI (s) + 3O (g)

However, this thermal decomposition will occurvery slowly in the absence of a catalyst. The rate of decomposition can be increased dramatically by adding a small amount of the catalyst manganese (MnO ), a powdery black

substance.All of the MnO can be recovered at the end of the reaction, just as all of the iodine ions,I ,remain following H O decomposition.

Regardless of its nature, a catalyst speeds up a reaction by providing a set of elementary steps with a more favourable kinetics than those that exist in its absence. The smaller the activation energy, E , the greater the rate. In many cases, a catalyst increases the rate by lowering the activation energy for the reaction.

Let us carry out Experiment 5.4 to study the effect of catalyst on the speed of

reaction.

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Experiment 5.4

1. Fill the basin and small measuring cylinder with water. Invert the

measuring cylinder into the basin that is filled with water (Figure

5.20).

2. Measure 50cm of 20 volume of hydrogen peroxide solution using

a measuring cylinder and pour it into the conical flask.

3. Put a weighing bottle containing a half spatulamanganese (IV)

oxide powder into the hydrogen peroxide solution.

4. Immediately cover the conical flask with the rubber stopper and

shake the flask slowly. Start the stopwatch at the same time.

5. Record the volume of oxygen released every 30 seconds for 300

seconds (5 minutes).

6. Repeat the experiment by adding a spatula of manganese (IV)

oxide powder.

7. The volume and concentration of the hydrogen peroxide solution

used are the same. 8. Then, plot two graphs of the volume of gas against time with

different amount of catalyst, on the same graph paper.

9. Calculate the average rate of reaction for each experiment. Does

the amount of catalyst increase the rate of reaction?

Figure 5.20: Set up of the apparatus for Experiment 5.4

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THE EFFECT OF ACTIVATION ENERGY ON THE SPEED OF A

REACTION

All molecules possess a certain minimum amount of energy. The energy can be in the form of kinetic energy and/or potential energy. When molecules collide, the kinetic energy of the molecules can be used to stretch, bend and ultimately break the bonds, leading to chemical reactions.

If molecules are moving too slowly with little kinetic energy, or collided with an improper orientation, they will not react and simply bounce off each other. However, if the molecules are moving at a fast enough velocity with a proper collision orientation, such as the kinetic energy upon collision is greater than the minimum energy barrier, then a reaction will occur. The minimum energy barrier that must be met for a chemical reaction to happen is called the activation energy, Ea. It can be represented by trying to push a stone to the other side as shown in Figure 5.22.

Figure 5.23: The man is trying to push the stone from point A to point B [Source : http://sites.tenafly.k12.nj.us/~shilfstein/demo_notes.htm ]

The reaction pathway can be observed in Figure 5.23. In order to get the product to react, the reactant has to overcome the activation energy, or a new product cannot be achieved if it does not have the same amount of energy.

5.4

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