CHEMISTRY: FORM THREE: Topic 7 – CHEMICAL KINETICS, EQUILIBRIUM AND ENERGETICS

Comparison between the Rates of Chemical Reactions

Compare the rates of chemical reactions

Experiments to Measure the Rates of Chemical Reactions

Perform experiments to measure the rates of chemical reactions

A chemical reaction will occur only if the particles of the reacting substances are allowed to come in contact. Increasing the concentration means increasing the density of the particles and hence the probability of particles being close together and colliding more often. Thus, a reaction can be made to go faster or slower by changing the concentration of a reactant.

The effect of concentration on the rate of reaction can be determined by mixing dilute hydrochloric acid with sodium thiosulphate solution to produce a precipitate of sulphur.Since this reaction produces a precipitate from two different colourless solutions, the intensity of the precipitate at any given moment in time represents the extent of the reaction.

The experiment to determine this effect is carried out by mixing 2M hydrochloric solution acid with 50 cm³ of the thiosulphate solution (at different concentrations as shown in the table below) and noting the time taken by the cross to disappear using a stopclock. The procedure similar to that used to investigate the effect of temperature above is used except that, in this particular experiment, the concentration of the thiosulphate is altered each time the experiment is repeated. The following table shows the results.

Concentration of thiosulphate (g/dm3)	10	20	30	40	50
Time(s) for the cross to disappear	250	120	65	32	15

As the results show, a reaction goes faster when the concentration of the thiosulphate (reactant) is increased. According to the above data, the rate of the reaction approximately doubles as the concentration of the thiosulphate is increased is increased by 10g dm^-3.

The results clearly indicate that, as the concentration of the reaction is increased the time for the disappearance of the cross decreases, which chemically means that the rate of the reaction increases.

Increasing the temperature of the system means increasing the kinetic energy and hence the speed at which the reacting particles of the substance move. Thus, the particles collide more often and combine to form new substances. Temperature also provides the energy required to break the bonds of a substance and hence enhances the decomposition or splitting of a complex substance into more simpler substances. Therefore, an increase in the temperature of the system will result to an increase in the rate of reaction.

The effect of temperature on the rate of reaction can be determined experimentally. When dilute hydrochloric acid is mixed with sodium thiosulphate solution, a fine yellow precipitate of sulphur forms.

2HCl_(aq) + Na₂S₂O_3(aq) → 2NaCl_(aq) + SO_2(g) + S_(s) + H₂O_(l)

The solid sulphur (S_(s)) produced in this reaction makes the colourless solution go cloudy. The reaction is usually carried out in a flask placed on a piece of white paper, with a black cross marked on it. An experiment is designed such that each time a different temperature is used, while keeping the concentrations of the reactants constant.

At the beginning of the reaction, the cross can be seen easily. As the reaction goes on, more and more sulphur is deposited, the flask becomes more and more cloudy and the cross gradually gets harder to see. At last, the cross can no longer be seen. It is fully covered by the precipitate of sulphur. Time taken by the cross to disappear at a given temperature indicates the rate of reaction at that temperature. The quicker the disappearance of the cross, the faster is the reaction and vice versa.

The reaction rate can be determined by the following procedure:

1. 50 cm³ of a solution of sodium thiosulphate (containing 40g per litre of the solution) is measured and put into a 100 cm³ beaker.
2. A wire gauze is placed on a tripod stand, then the beaker is put on the gauze, and the solution is gently warmed with a Bunsen burner flame.
3. A thick black cross is marked on a white piece of paper or cardboard.
4. When the temperature reaches a little above 20⁰C, the beaker is taken quickly and placed on the cross.
5. 10 cm³ of 2M hydrochloric acid solution is added quickly to the beaker and a clock is started at the same time. The temperature of the mixture is noted.
6. When the cross is no longer visible, the clock is stopped.
7. The experiment is repeated more times at temperatures of 30⁰C, 40⁰C, 50⁰C, 60⁰C, etc. Each time 10 cm³ of 2M hydrochloric acid and 50 cm³ of the thiosulphate solution are used.

The following table shows the specimen results obtained from such an experiment:

Temperature (0C)	20	30	40	50	60
Time (s) for cross to disappear	200	125	50	33	24

The results show that the higher the temperature the faster the cross disappears, which in turn, means that the rate of the reaction increases.Therefore, the data clearly indicate that a reaction goes faster when the temperature is raised. When the temperature is increased by 10⁰C, the rate approximately doubles.

In many reactions, one of the reactants is a solid. The reaction between hydrochloric acid and calcium carbonate (marble chips) is one example. Carbon dioxide gas is produced.

CaCO_3(s) + 2HCl_(aq) → CaCl_2(aq) + H₂O_(l) + CO_2(g)

One of the ways by which the rate of such a reaction can be increased is by reducing the particle size of the solid substance (marble chips). If this substance is grinded to fine powder or to small pellets, the surface area of the marble is increased. Therefore, more of the marble is exposed to the acid for efficient reaction. This leads to increase in the rate of reaction. Increased rate of reaction results to increased production of carbon dioxide gas. Hence, the rate of reaction can be determined by measuring the time taken to produce a given volume of carbon dioxide by reacting equal masses of whole marble and crushed (or powdered) marble, and then comparing the two results. The data obtained is recorded in a table as shown below:

Time (s)		0	10	20	30	40	50	60	70	80	90	100
Volume (cm3)	Whole marble	0	18	24	32	38	45	53	62	72	80	80
	Grinded marble	0	34	52	67	74	76	80	80	80	80	80

The data above shows that it takes 90 seconds for whole marble chips to react to completion while for powdered marble it takes only 60 seconds. Also there is seen a small increase in volume of carbon dioxide gas per unit of time in the case of ungrounded (whole) marble as compared to the powdered sample. This absolutely proves that an increase in the surface area of marble (a solid reactant) increases the rate of reaction and hence the rate of production of carbon dioxide gas.

A catalyst usually increases the rate of a reaction and this is called positive catalysis. In general, a catalyst will function even if present in very small amounts. Hydrogen peroxide is a clear, colourless liquid. It can decompose to water and oxygen:

2H₂O_2(aq) → 2H₂O_(l) + O_2(g)

The rate of decomposition (reaction) of the peroxide can be increased tremendously by adding a very little of manganese (IV) oxide in the reaction vessel.

The reaction in flask A above is in fact very slow. It could take several days to produce just 50 cm³ of oxygen. In flask B only 1g of manganese (IV) oxide is added. The reaction goes very faster. 50 cm³ of oxygen is produced in a few minutes.

The manganese (IV) oxide speeds up the reaction without being used up itself. It is called a catalyst for the reaction.A catalyst is a substance that changes the rate of a chemical reaction but remains chemically unchanged at the end of the reaction.

Catalysts for many reactions have been discovered. They are usually transition metals or compounds of transition metals. There are also biological catalysts, called enzymes. For example, the pancreatic juice secreted by the pancreas contains enzymes that speed up digestion process.

Examples of catalysts for some common reactions are given in the following table:

Reaction	Catalyst
Heating of potassium chlorate 2KClO3(s) → 2KCl(s) + 3O2(g)	Manganese (IV) oxide (MnO2)
Synthesis of sulphur trioxide 2SO2(g) + O2(g)⇔2SO3(g)	Vanadium (IV) oxide (V2O5)
1. Synthesis of ammonia	Reduced iron powder
2. Decomposition of hydrogen peroxide	Manganese (IV) oxide or platinum powder

Chemical reactions can be classified into two groups, that is, reversible and irreversible reactions.

Irreversible reactions are those reactions which go to completion. In such reactions, a known product is formed. The product(s) cannot be reversed back to the original reactant(s). Consider the reaction between sodium and chlorine to produce sodium chloride: 2Na_(s) + Cl_2(g) → 2NaCl_(s)

You cannot turn sodium chloride back to chlorine gas and pure sodium metal by ordinary means. This is a typical irreversible reaction.

Several reactions are known to proceed in both directions, provided certain conditions are maintained.A + B⇌AB Such reactions are said to be reversible. The sign ⇌indicates reversibility.

A reversible reaction is a chemical reaction in which the products can react to re-form the reactants.In such reactions, reactants combine to form products. However, under certain conditions, the products may be converted back to reactants.The idea of “reactants” and “products” in such circumstances is really confusing.

Examples of reversible reactions

1. When you heat blue crystals of copper (II) sulphate, they breakdown into anhydrous copper (II) sulphate, a white powder: CuSO₄.5H₂O_(s)⇌CuSO_4(s) + 5H₂O_(g)The reaction can be reversed by just adding water to the white powder, which quickly turns to blue crystals again. In fact, this is used as a test for water: CuSO_4(s) + 5H₂O_(l)⇌CuSO₄.5H₂O_(s)
2. When you heat ammonium chloride (a solid) in the bottom of a test tube, it breaks down into ammonia and hydrogen chloride (gases). The gases readily combine at the top of the tube where it is cool: NH₄Cl_(s)⇌NH_3(g) + HCl_(g)This is a reversible reaction.

Suppose that 2 molecules of substance A reacts with 3 molecules of substance B to produce 1 and 2 molecules of substances C and D respectively in a homogenous system (i.e. entirely liquid or entirely gaseous).

2A + 3B⇌C + 2D

As soon as a little of C and D are formed, a reverse reaction will begin. At first the forward reaction will predominate, but, as C and D accumulate the reverse reaction will build up until an equilibrium position is reached, with forward and reverse reactions proceeding at the same rate. The composition of the mixture will then appear constant, though it is the net result of the two opposing reactions. Since chemical equilibrium involves the balancing of two reactions which are proceeding at the same time in opposite directions, it is said to be a dynamic equilibrium, that is, it is an equilibrium involving the constant interchange of particles in motion. Equilibrium is a dynamic condition in which two opposing changes can occur at equal rates in a closed system. In a closed system, matter cannot enter or leave, but energy can. Both matter and energy can escape or enter an open system. For example, sunlight, heat from a burner, or cooling by ice can cause energy to enter or leave a system.

An example of a real equilibrium reaction is the decomposition of mercury (II) oxide. Mercury (II) oxide decomposes when heated to produce mercury and oxygen gas.

Again, mercury and oxygen combine to form mercury (II) oxide when heated gently

Suppose mercury (II) oxide is heated in a closed system. Once the decomposition has begun, the mercury and oxygen released can re-combine to form mercury (II) oxide again. Thus, both reactions can proceed at the same time. Under these conditions, the rate of the composition (re-combination) reaction will eventually equal that of the decomposition reaction. Then the reaction is said to be at equilibrium. At equilibrium, mercury and oxygen will combine to form mercury (II) oxide at the same rate that mercury (II) oxide decomposes into mercury and oxygen. The amounts of mercury (II) oxide, mercury and oxygen can then be expected to remain constant as long as these conditions persist. At this point, a state of dynamic equilibrium has been reached between the two chemical reactions. Both reactions continue, but there is no net change in the composition of the system.

A reversible reaction is in chemical equilibrium when the rate of its forward reaction equals the rate of its reverse reaction, and the concentrations of its reactants and products remain unchanged.

Conditions for the industrial synthesis of different substances have to be carefully chosen if the process is to be efficient and thus economically viable. The considerations are:

• Yield – how much yield is produced?
• Rate – how fast is it produced?
• Energy – how much energy is lost during the process?

The way in which this is achieved in practice is described briefly for three important industrial reactions, namely, the Haber Process for the manufacture of ammonia, the Contact Process for the manufacture of sulphuric acid and the thermal dissociation of calcium carbonate for the industrial manufacture of lime.

THE HABER PROCESS FOR THE INDUSTRIAL MANUFACTURE OF AMMONIA

THE CONTACT PROCESS FOR THE INDUSTRIAL MANUFACTURE OF SULPHURIC ACID

You have met many different chemical reactions so far in chemistry. But they all have one thing in common, that is, they involve an energy change. The great majority of chemical reactions are accompanied by a marked heat change.

During chemical reactions as reactants form products, there is a change in heat content. This is referred to as the enthalpy change and is always expressed in kilojoules per mole (kJmol^-1). Two types of heat change are distinguished. Those reactions that are accompanied by evolution of heat to the surroundings are termed as exothermic reactions while those that are accompanied by absorption of heat from the surroundings are endothermic reactions.

• An exothermic reaction is one during which heat is liberated to the surroundings.
• An endothermic reaction is one during which heat is absorbed from the surroundings

When magnesium is burnt in air heat is evolved.

2Mg_(s) + O_2(g) → 2MgO_(s)+ heat

The same case applies to the burning of coal in air.

C_(s)+ O_2(g) → CO_2(g) + heat

Mixing sulphur nitrate and sodium chloride solutions gives a white precipitate of silver chloride and a temperature rise.

AgNO_3(aq) + NaCl_(aq) → AgCl_(s) + NaNO_3(aq)

When ammonium nitrate is dissolved in water, there is a fall in temperature. Also adding a mixture of citric acid and sodium bicarbonate to water produces bubbles and a fall in temperature. In both reactions, the temperature of the water falls because the reactions take heat energy from it. These reactions are therefore endothermic.

The heat changes that occur during any chemical reaction represent changes in the energy content of the whole system. The energy content may increase or decrease depending upon whether heat is absorbed or evolved.

For exothermic reactions, the enthalpy changeis conventionally assigned a negative value. For example, when pellets of sodium hydroxide or concentrated sulphuric acid dissolve in water, heat is evolved and the system loses heat to the surrounding.

For endothermic reactions, the enthalpy changeis assigned a positive value. For example, when potassium iodide or ammonium chloride dissolves in water, heat is absorbed from the surroundings.