CHEMICAL KINETICS, EQUILIBRIUM AND

ENERGETICS

The Rate of Chemical Reactions

‘Copmarison between the Rates of Chemical Reactions Compare the rates of chemical reactions ‘Chemical reactions take place at different rates. Some are fast whereas others are very slow. Let us consider the following reactions: 1, Addition of sodium metal to water: 2Najy+ 2H0\)—+ 2NaOHiag + Hayy The reaction takes place immediately and violently. It is therefore o fast reaction.

• ‘The rusting of iron in the presence of air and water giving hydrated iron (Ill) oxide, P30, XHLO:

This is an extremely slow reaction These two reactions could be taken as representative examples of extremely fast and extremely slow reactions, respectively. There are, however, other reactions which proceed at rates intermediate between these two extremes. Rates of some of these reactions can be measured The rate of a chemical reaction can be measured in various ways. Let us consider the reaction between zine and sulphuric acid to produce zine sulphate and hydrogen gas

Za) + W809) -? ZaSO40)+ He When zine is added to dilute sulphuric acid in a flask, they react together. The zine slowly disappears and the gas (Hl) bubbles off: After sometime, the bubbles of a gas form less quickly.

The reaction is slowing down. Finally, no more bubbles appear. The reaction is over, because all the acid has been used up. Some zine remains behind in a beaker In this reaction both zine and sulphuric acid get used up in the reaction. At the same time, zine sulphate and hydrogen form. The raie of this reaction could be determined by measuring any of the following:

• the amount of zine used up per unit of time;
• the amount of hydrogen produced per uit of time

In general, the rate of a chemical reaction is determined by measuring the amount of reactantused up per unit of time or the amount of product produced per unit of time. Therefore, time sninutz, The hydrogen can be collected as it bubbles off and its volume ean then be measured as hydrogen (acs ee ay, pene WG 25 em? 100 ex?

f \ PSE Gee Zine + hydrochloric acid Apparatus for measuring the production of gas An experiment may be designed to messure the volume of hydrogen produced after every twenty from such an experiment.

Time (3) Volume of hydrogen gas (cm?) o ° 20 B Py 2 vw 2 100 a 120 “4

Questions from the experiment

axis) Note: rate of reaction = Change in volume Change in time

• change in intensity of colour Many chemical reactions involve a change in colour. Potassium

permanganate, for example, when it resets with sulphur dioxide it changes from purple to colourless. The rate of such a reaction couk! be determined by measuring the rate at which the colour changes.

• formation or disappearance of a precipitate: The reaction between hydrochloric acid and sodium

thiosulphate produce a yellow precipitate of sulphur. The rate at which this precipitate forms is a measure of the rate of reaction

EACTORS AFFECTING THE RATE OF CHEMICAL REACTIONS

The Effect of Concentration on the Rate of Reaction Describe the effect of concentration on the rate of a reaction 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 em’ 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/dm’) 10 20 300-4050 Time(s) forthe eross to disappear 2000 10S aS As the results show, 2 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” The results clearly indicate that, as the concentration of the reaction is mereased the time for the

disappearance of the cross decreases, which chemically means that the rate of the reaction increases. The Effect of Temperature on the Rate of a Reaction Demonstrate the effect of temperature on the rate of a reaction Increasing the temperature of the system means inereasing 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 10

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 2HChag) + Na2S:0s9) > 2NaCha) + SO + Se + HO The solid sulphur (Ss) produced in this reaction makes the colourless solution go cloudy. The reaction is usually carried out in a Mask placed on a piece of white paper, with a black cross

marked on it. An experiment is designed such that cach 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 casily. 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 eross 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, th faster isthe reaction and vice The reaction rate can be determined by the following procedure:

• 50 em’ of a solution of sodium thiosulphate (containing 40g per litre of the solution) is measured

and put into a 100 em? beaker gently warmed with a Bunsen burmer flame.

• A thick black eross is marked on a white piece of paper or cardboard

4, When the temperature reaches a little above 20°C, the beaker is taken quickly and placed on the the same time. The temperature of the mixture is noted.

• When the cross is no longer visible, the clock is stopped.
• The experiment is repeated more times at temperatures of 30°C, 40°C, 50°C, 60°C, ete, Each time

10 cm? of 2M hydrochloric acid and $0 em* of the thiosulphate solution are used. white paper or card beaker s > cross no longer visible ss J start of reaction after ashort time after a much longer time Measuring the effect of temperature on the rate of reaction The following table shows the specimen results obtained fiom such an experiment Temperature (°C) 20 30 4 506 Time (s) for cross to disappear 200 125 3 33 The results show that the higher the temperature the faster the eross disappears, which in turn,

means that the rate of the resetion inereases Therefore, the data clearly indicate that a reaction es faster when the temperature is raised, When the temperature is increased by 10°C, the rate approximately doubles The Effect of Surface Area of a Solid on the Rate of a Reaction Show the effect of surface area of a solid on the rate of a reaction In many reactions, one of the reactants is a solid. The reaction between hydrochloric acid and

calcium earbonate (marble chips) is one example. Carbon dioxide gas is produced. CaCOys;+ 2HChay) > CaCL yy) + H2Oq + COr) ‘One of the ways by which the rate of such 2 reaction can be increased is by redueing 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 inereased. Therefore, more of the marble is exposed to the acid for efficient reaction. This leads to inctease in the rate of reaction Inereased rate of

reaction results to inereased 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 (3) 0 1 2 30 4 so 6 70 80 9 104 Wholemable 0 «18-24. «32,—«R_ dS SBD 72-8 Volume (em'y 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 scen 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 rte of produetion of carbon dionide gas The Effect of Catalyst onthe Rate ofa Reaction Demonstrate the effect of catalyst on the rate of a reaction A catalyst usually increases the rate of a resction and this is called positive catalysis. In genera, « catalyst will function even if present in very small amounts. Hydrogen peroxide is a clear, colourless liquid. It ean decompose to water and oxygen

2Ha0 449) + 2H204y + Ora) The rate of decomposition (reaction) of the peroxide can be inereased tremendously by adding & very litle of manganese (IV) oxide in the reaction vessel Sy

Pct Ph

tr {it FY socmiofoxygen A 50cm? of oxygen p ae ees A 3 A ™ oO eo B hydrogen peroxidesolution manganese (IV) oxide Investigating the effect of catalyst on rate of reaction The reaction in flask A above is in fact very slow It could take several days to produce just $0 em’ of oxygen. In flask B only g of manganese (IV) oxide is added. The reaction goes very faster. 50 cm® 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. catalyst is a substance that changes the rate ofa 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 2KCIO3.) + 2KC,) + 30g Manganese (IV) oxide (MnO) Synthesis of sulphur trioxide 2SOx) + Onge.28Os) Vanadium (IV) oxide (20s)

• Synthesis of ammonia Reduced iron powder
• Decomposition of hydrogen peroxide Manganese (IV) oxide or platinum powde

Reversible and Irreversible Reactions

‘Comparison between Reversible and Irreversible Reactions Compare reversible and irreversible reactions Chemical reactions can be classified into two groups, that is, reversible and irreversible reactions.

Ineversible reactions are those reactions which go to completion. In sueh reactions, a known product 1s formed. ‘The product(s) cannot be reversed back fo the onginal reactant(s). Consider the reaction between sodium and chlorine to produce sodium chloride:2Naj + Clg) > 2NaCly, 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 + BetABSuch reactions are said to be reversible. The sign #indicates reversibility, ‘The Concept of Reversible and Irreversible Reactions Describe the concept of reversible and irreversible reactions 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.

Example 1

Examples of reversible resetions

• When you heat blue crystals of copper (II) sulphate, they breakdown into anhydrous copper (Il)

sulphate, a white powder CuSO, SHO j#CuSOys + SHO The reaction can be reversed by {ust adding water to the white powder, which quickly tums to blue erystals again. Infact, this is used as atest for water: CuSOsq) + SH2O\ie#CuSOs SHO.) 2 When you heat ammonium chloride (2 solid) in the bottom of a test tube, it breaks dawn into ammonia and hydrogen chloride (gases). The gases readily combine at the top of the tube where

itis cook: NH,Cy=NHy.) + HCl This is a reversible reaction,

Equilibrium Reactions

Difference between Equilibrium Reactions and Simple Reversible Reactions

Differentiate equilibrium reactions from simple reversible reactions 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 (Lc. entirely liquid or entirely gaseous),

2A +3BAC +20

As soon as a little of C and D ere formed, a reverse reaction will begin, At first the forward reaction will predominate, but, as Cand 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 myolves the balancing of two reactions which

are proceeding at the same time in opposite directions, itis said to be a dynamic equilibrium, that is, it 1s 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, Ina closed system, matter cannot enter or lewe, but energy can. Both matter and energy An example of a teal equilibrium reaction is the decomposition of mercury (Il) oxide. Mercury

• oxide decomposes when heated to produce mercury and oxygen gas.

2Hg0) —*> 2H gy + Ore Again, mereury and oxygen combine t form mereury (II) oxide when heated gently 2Hig + Org) —2 > 2H20 Suppose mercury (Il) oxide is heated in closed system. Once the decomposition has begun, the reactions can proceed atthe same time. Under these conditions, the rate of the composition (re- is said to be at equilibrium. At equilibrium, mercury and oxygen will combine to form mercury amounts of mercury (Il) 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 unchanged. ‘Two Equilibrium Reactions of Industrial importance Conditions for the industrial synthesis of different substances have to be carefully chosen ifthe

• Yield—how much yield is produced?
• Rate —how fast i 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 mamufacture of lime.

‘THE HABER PROCESS FOR THE INDUSTRIAL MANUFACTURE OF AMMONIA

The Haber process involves the reaction between nitrogen and hydrogen to produce ammonia: Nag) + 3Haig#2NHaig A German Chemis, Fritz Haber, was the first to show how this reaction could be controlled to make useful amounts of ammonia. ‘The first industrial plant making ammonia by the Haber process opened in Germany in 1913. Now over 100 million tones of ammonia are produced each year by this process. Because ofits importance, the Haber process for making ammonia has been

studied over a wide range of conditions of temperature and pressure. ‘The effect of pressure Ammonia is produced from its elements by reduction of volume. Therefore, if the system is in ‘equilibrium and the pressure is raised, the equilibrium will shift to the right This is because the system will shift to favour the side of the equation that has fewer molecules Nag + 3Hagie2NHaig Therefore, high pressure will increase the yield of ammonia, Modem industrial plants use @

pressure of 200-500 atmospheres. Pressure higher than this range could be used, but high: pressure reaction vessels are expensive to build, The effect of temperature ‘The forward reaction producing ammonia is exothermic and the reverse reaction is therefore endothermic, Therefore, lowering the temperature will favour ammonia production. However, the rate at which the ammonia is produced will be so slow as to be uneconomical. So, itis

necessary to include a catalyst which will give sufficient reaction rate in spite of a relatively low pressure. In practice, a compromise or optimum temperature is used to produce enough ammonia at an acceptable rate, Modem plants use temperatures of about 450°C.

‘The effect of concentration the system is in equilibrium and more nitrogen is then added to increase its concentration in the reaction mixture, the equilibrium shifts to the right so as to tend to reduce the concentration of nitrogen. ‘That is, more ammonia will be produced to use up nitrogen. Also, if hydrogen is added the equilibrium will, similarly, shift to the right.

However, in practice, there is no particular advantage in using excess of either material (nitrogen or hydrogen) since the gases, nitrogen and hydrogen, are mixed in a ratio of 3:1 by volume Hydrogen is manufactured from partial combustion of hydrocarbons, and nitrogen is obtained from the sir. They are mixed in the ratio of 3:1 proportion by volume and dried (e.g. by silica gel). They are pre-heated by gases leaving the catalyst chamber over the catalyst at 450°C

The ammonia produced is absorbed in water or liquefied by reftigeration and the remaining nitrogen and hydrogen gases are recycled. Ifthe system was at equilibrium and then some of the ammonia was removed, more ammonia would be produced to replace that which is removed, The gas is removed from the reaction ‘chamber when the percentage of ammonia in the equilibrium mixture is 15%.

Based on the above discussions, the overall conditions used for the maximum production of ammonia are as follows

• Noand Hs are mixed in the ration of 3:1
• An optimum temperature of 450°C is chosen,
• Avery high pressure (200-500 atm) is applied.
• A catalyst of finely divided reduced iron, usually promoted by alumina (aluminium oxide) is

used. +The ammonia is condensed and removed out of the reaction mixture and the remaining Nzand Hy recycled.

THE CONTACT PROCESS FOR THE INDUSTRIAL MANUFACTURE OF

SULPHURIC ACID

The first step in the production of sulphuric acid is the conversion of sulphur dioxide to sulphur trioxide: The process involves the reaction between sulphur dioxide and oxygen. The reaction is exothermic and reversible 28034) + Oxin%*2SOxs) To make sulphuric acid, the sulphur trioxide gas produced is dissolved in 98% concentrated sulphuric aeid, forming “oleum’” which is then diluted with the correct amount of water to give

ordinary concentrated sulphuric acid

• The effect of pressure: There are fewer gos molecules on the right of thc equation

Therefore, increasing the pressure would favour the production of sulphur trioxide. In fact, the process is run at atmospheric pressure because the conversion of sulphur dioxide to sulphur trioxide is about 96% complete under these conditions. In practice, it is found that use of high pressure produces a very small gain in yield and involves extra cost, It is for this reason that ordinary atmospheric pressure is used.

• The effect of temperature:The reaction to produce sulphur trioxide is exothermic so, if the

temperature is lowered the equilibrium shifts to the right. Hence, more sulphur trioxide will be produced. This means that sulphur trioxide production would be favoured by low temperatures, However, 100 low temperatures reduce the rate of reaction so increasing the time requited for the production of the sulphur trioxide, This would mean an increase in production cost. For this reason, a catalyst must be introduced, A catalyst of vanadium (V) oxide is used to increase the

rate of reaction, An optimum temperature of about 450°C is used. This gives sufficient sulphur trioxide at an economic rate. In general, low temperatures give an equilibrium favourable to an exothermic reaction, but catalysis is needed to give a favourable reaetion rate.

• Effeet of concentration of reactants:Suppose that in the contact process reaction, an

‘equilibrium has been reached in certain conditions and oxygen is added to the system, According to Le Chatelier’s principle, the reaction would shift to the right so as to oppose this change, that is, to reduce the concentration of the oxygen added towards its former level. This can only be achieved by combining it with sulphur dioxide to form sulphur trioxide. So, increased concentration of oxygen favours conversion of more sulphur dioxide to trioxide. Likewise, a

similar case can occur when sulphur dioxide is added to the system. In a similar manner, increased concentration of sulphur trioxide would favour the conversion of oxygen to sulphur trioxide and hence increased formation of the trioxide.In general, the conditions used in the contact process are as follows: An optimum temperature of about 450°C iy chosen; A catalyst vanadium (V) oxide, is used to speed up the reaction; An operating pressure of 1 atmosphere is

applied. A higher pressure, which would theoretically increase the yield of SOs, is not used as i is uneconomical; Yield of SOs would be increased by increasing either SO; or Oz, The cheaper reactant is Os

Endothermic and Exothermie Reaction

‘The Concept of Endothermic and Exothermie Reactions Explain the concept of endothermic and exothermic reactions You have met many different chemical reactions so far in chemistry. But they all have one thing in common, thet is, they involve an energy change. The great majority of chemical reactions are accompanied by a marked heat change.

Dunng chemical reactions as reactants form products, there is a change in heat content. ‘his is referred to as the enthalpy changeand is always expressed in kilojoules per mole (kJmot"). 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

• Anexothermic 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 burt in air heat is evolved 2Mges + Oxy) —> 2Me0jat heat The same ease applies to the burning of coal in air. Cot Oxg) 7 COn + heat Mixing sulphur nitrate and sodium chloride solutions gives a white precipitate of silver chloride and a temperature rise ABNOsay) + NaClag 7 AgClo) + NNOxaq) When ammonium nitrate is dissolved im water, there is a fall in temperature. Also adding © mixture of citrie acid and sodium bicarbonate w water produces bubbles anda fall in

temperature, In both reactions, the temperature of the water falls because the reactions take heat energy fiom 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 deerease depending upon whether heat is absorbed or evolved.

Energy Level Diagrams for Exothermic and Endothermie Reactions Dravr energy level diagrams for exothermic and endothermic reactions 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

A t / \ Eq= activation . energy reactants _/ \ energy -———<— ~-2 AH negative 4 products

Energy level diagram for exothermic reaction For endothermic reactions, the enthalpy changeis assigned a positive value, For example, when potassium iodide or ammonium chloride dissolves in water, beat is absorbed fiom the surroundings

4 ey \ Ea= activation energy \_ product g Se AH reactant Positive a t i