senior chem - controlling reactions

8/9/2019 Senior Chem - Controlling Reactions

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USING AND CONTROLLING REACTIONS

TOPIC 4: USING AND CONTROLLING REACTIONS

The use and control of chemical reactions are important tasks undertaken by chemists. This topic looks at the

energy changes that accompany chemical reactions, as well as their rates and extents. It also examines the

ways in which chemical reactions are controlled and used to make materials and generate the energy needed

by a modern industrial society.

The increased use of energy from chemical reactions has been a major factor in the evolution of theindustrialised world. In this topic students consider the ways in which this energy is produced and begin

quantitative consideration of the energy changes that accompany chemical reactions.

The production of chemicals is the main function of the chemical industry. These chemicals allow naturally

occurring materials to be modified or replaced and previously unknown materials to be developed. The

industrialised world depends on a chemical industry for the manufacture of a diverse range of materials. In

this topic students look at how chemicals are produced and how the production can be performed most

efficiently.

Knowledge of chemistry can be applied to manipulate the reaction conditions of industrial processes in order

to determine the quantity or quality of the product.

4.1 MEASURING ENERGY CHANGES

Key Ideas Intended Student Learning

Almost all chemical reactions occur with either anabsorption or a release of heat or light energy. Otherforms of energy, such as electrical energy, can also bereleased.

Identify combustion and respiration as reactionsthat release energy and photosynthesis as areaction that absorbs energy.

Exothermic reactions release energy to thesurroundings whereas endothermic reactions absorb

energy from the surroundings.

Deduce whether a reaction is exothermic orendothermic from information provided.

The measurement of the heat change in chemicalreactions is called calorimetry; the insulatedapparatus used for the measurement is a calorimeter.

Calculate the heat produced or absorbed for areaction from experimental data, given thespecific heat capacity of water as 4.18 J g1 K1.

The heat released or absorbed in a reaction at constantpressure is called the enthalpy change for thereaction; it is given the symbol H.

Determine enthalpy changes from experimentaldata for reactions, including:

the combustion of alcohols;

the neutralisation of acids with bases;

solution processes.

Exothermic reactions have negative Hvalues.

Endothermic reactions have positive Hvalues.

Identify a reaction as exothermic or endothermic,

given a thermochemical equation or the value ofits enthalpy change.

Thermochemical equations express a quantitativerelationship between the quantities of reactants and theenthalpy change.

Write thermochemical equations that correspondto given molar enthalpies of combustion,neutralisation, and solution.

The magnitude of the heat absorbed or evolved for areaction is directly proportional to the quantities ofreactants involved.

Calculate the theoretical temperature change of aspecified mass of water or solution heated orcooled by a reaction, given molar enthalpies andquantities of reactants.

Introduction

Chemical reactions are usually accompanied by the absorption or release of heat and light energy. Electrical

and sound energy may also be produced.

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Reactions that produce heat include:

Acid-base neutralisations

Combustion reactions

Solution of a solute in a solvent

Exothermic and endothermic reactions

Exothermic reactions release heat energy into the surroundings and cause an increase in temperature.

Endothermic reactions absorb heat energy from the surroundings and cause a decrease in temperature.

Enthalpy changes for reactions

The quantity of heat energy released or absorbed when specific amounts of substances react is called the heat

of reaction. For reactions carried out at constant pressure, the heat of reaction is called the enthalpy change.

The symbol is H.

The molar enthalpy change for a reaction is the quantity of heat released or absorbed when 1.00 mole of a

specific substance reacts in a chemical reaction under constant pressure. Units - kilojoules per mole (kJ mol-

1).

Example

The molar enthalpy of combustion of butane ( Hcomb(butane)) =2874 kJ mol-1. This means that when 1 mole

of butane undergoes complete combustion, 2874 kJ of heat energy are released to the surroundings.

Calorimetry

Calorimetry is the measurement of energy changes during a chemical reaction. The vessel used for the

measurement is called a calorimeter, which is an insulated reaction vessel. eg a foam cup.

The heat energy absorbed or released by a particular reaction is calculated using the expression.

Heat = specific heat capacity of water x the temperature change x mass

Or

H = c x T x m

Where H = heat change (joules (J))

C = heat capacity (J g-1C-1) for water c = 4.18 J g-1C-1

m = mass (g)

Note -Frequently the volume of water is measured and this can easily be converted to mass since 1.0 mL

1.0 g

To convert the heat change to molar enthalpy

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where is n the number of moles of the reactant.

It is necessary to divide by 1000 to convert the J to kJ and by the number of moles (n) to give kJ mol-1

Example

Given that 3.24 g of methanol was burnt and the heat used to heat 200 mL of water from 18.2C to 38.2C,

calculate the molar enthalpy of combustion of methanol.

m = 3.24 g T = 38.2 18.2 = 20C

M(CH3OH) = 12 + 3x1 + 16 + 1 = 32 g mol-1

H = c x T x m= 4.18 x 20 x 200

= 16720 J

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m (water) = 200 mL ( 200 g)

n(CH3OH) = m/M

= 3.24/32

= 0.10125 mol

H =

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=

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H = -165 kJ mol-1

Assumptions and approximations

It is assumed that

the calorimeter absorbed no heat.

Heat losses to the surroundings can be ignored.

The specific heat of the solution is the same as distilled water.

The mass of 1 mL of solution is 1 g.

Themochemical EquationsThermochemical reactions are chemical equations that indicate:

The mole ratio of the reactants

The state of each reactant and product

The quantity of heat energy released (-) or absorbed (+) by the mole quantities of reactants or products

indicated in the equation.

Example MgCO3(s) + 2HCl(aq) MgCl2(aq) + O2(g) H = -90.4 kJ mol-1

Specific Cases of enthalpy changes

Molar enthalpy of combustion of a substance is the quantity of heat energy released when 1.00 mole of pure

element or compound is burnt completely in oxygen under constant pressure.

Example CH4(g) + 2O2 CO2(g) + 2H2O(l) H = -890 kJ mol-1

EXAMPLE

The fuel used in pocket cigarette lighters is butane, C4H10. The molar enthalpy combustion value for butane

is 2874 kJ mol-1. If 1.00 g of butane in a lighter was burnt beneath a steel can containing 1000 mL of water

initially at 18.5C, calculate the theoretical maximum temperature reached by the water. Would you expect

that the temperature is actually achieved? Discuss.

Molar enthalpy of solution of a substance is the quantity of heat energy released or absorbed when 1.00

mole of the substance dissolves in sufficient solvent so that further dilution causes no further release or

absorption of heat energy.

Example NaOH(s) + aq Na+

(aq) + OH-(aq) H = -43 kJ mol

-1

EXAMPLE

In an experiment to determine the the molar enthalpy of solution of potassium hydroxide, 4.9 g of potassium

hydroxide was added to 100 mL of distilled water in a polystyrene cup. The temperature of water increased

from 18.5C to 28.1C. Assuming that the specific heat of the solution is 4.18 J g-1C-1, calculate the molar

enthalpy of solution of potassium hydroxide.

Molar enthalpy of neutralisation is the quantity of heat energy released when 1.00 mole of hydrogen ions is

transferred from an acid to a base in an acid-base reaction occurring in aqueous solution.

Example NaOH(aq) + HCl(aq) NaCl(aq) + H2O(l) H = -57 kJ mol-1

EXAMPLE

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The experiment described below was performed to determine the molar enthalpy of neutralisation for the

following neutralisation:

NH3(aq) + HCl(aq) NH4Cl(aq)

50.0 mL of 1.0 mol L-1 ammonia solution was mixed with 50 mL of 1.0 mol L-1 hydrochloric acid is a

polystyrene cup. The temperature of both solutions was 21.4C prior to mixing. After mixing the solutions the

maximum temperature reached was 25.6C. Use this data to calculate the enthalpy of neutralization for thisreaction.

ASSIGNMENT 4.1: MEASURING ENERGY CHANGES

1. Many reactions involve heat being given out or taken in.

(a) Define an exothermic reaction.

(b) Define an endothermic reaction.

1. For each of the following reactions, write a balanced equation for the reaction and classify it as

exothermic or endothermic.

(a) The burning of ethanol in excess oxygen

(b) Photosynthesis

(c) Aerobic respiration

1. When sulfuric acid is mixed with water it generates a lot of heat.

(a) Is this reaction an exothermic or endothermic one?

(b) Sulfuric acid is also denser that water. Explain why sulfuric acid is added TO water during

dilution rather than the other way round.

(c) Draw an energy profile diagram for this reaction.

1. When sodium nitrate is added to water the resulting solution is colder than the water added.

(a) Write a balanced annotated equation for this reaction.

(b) What type of reaction is this? Ionization or dissociation?

(c) Explain the temperature change.

(d) Draw an energy profile diagram for this reaction.

1. Distinguish between heat of reaction and molar enthalpy of reaction.

2. Given that 2.37 g of ethanol was burnt and the heat used to heat 200 mL of water from 18.8C to

39.6C, calculate the molar enthalpy of combustion of methanol.

3. The calorimeter used in heat experiments is frequently a foam cup.

(a) Why is a foam cup used?(b) Why is the reaction mixture stirred during the reaction?

(c) Why is the final temperature measured as quickly as possible?

(d) What 4 assumptions are made during the calculation of the molar enthalpy change?

1. In industry and research laboratories for the

determination of H values is the bomb

calorimeter. The combustion of a preweighed sample

is carried out in pure oxygen inside a steel bomb

which in turn is immersed in water. Combustion is

both complete and rapid. The water and other

calorimeter parts absorb the heat energy released

and the temperature increase is recorded.

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Calibration of the bomb calorimeter is achieved by burning substances of known H and recording

the increases in temperature.

(a) Explain how the systematic errors inherent in the spirit burner method for determining H are

eliminated or minimized by the use of a bomb calorimeter.

(b) Explain how the bomb calorimeter can be calibrated.

1. Define the following terms

(a) Enthalpy of combustion

(b) Enthalpy of solution

(c) Enthalpy of neutralization.

1. Glucose is used as a source of energy in the human body. The thermochemical equation for the

combustion of glucose in the body is shown below:

C6H12O6(s) + 6O2(g) 6CO2(g) + 6H2O(l) H = -2803 kJ

(a) Calculate the amount of energy released in the body when 1.0 g of glucose undergoes

combustion.

Glucose can also undergo fermentation to produce ethanol, according to the following

equation:

C6H12O6(s) 2C2H5OH(l) + 2CO2(g)

The ethanol produced by fermentation can be burnt as a fuel, according to the following

thermochemical equation:C2H5OH(l) + 3O2(g) 2CO2(g)+3H2O(1) H = -1364 kJ

(b) Use your answer to part (a) to predict whether or not the fermentation of glucose will warm the

reaction mixture. Give a reason for your answer.

1. Methane is most commonly used as a fuel. The enthalpy of combustion of methane is 890 kJ mol-1.

(a) Write a thermochemical equation for the enthalpy of combustion of methane.

(b) 25 g of methane was burnt to heat some water on a gas stove. Calculate the amount of energy

produced.

(c) Calculate the mass of water that could be heated from 25C to 100C by the amount of energy

produced in part (b). [4.18 J raises the temperature of 1.0 g of water by 1.0C.]

(d) Suggest two reasons why the final temperature would have been less than 100C when the 25 g

of methane was burnt to heat the mass of water you have calculated in part (c).

1. (a) When 2.50 g of ammonium chloride is dissolved in 100 g of water the temperature falls by 1.7 C.

Describe, including details of the apparatus used, how this determination would have been performed

in a laboratory.

(a) From the data in (a) calculate the enthalpy of solution of ammonium chloride. State two

assumptions made in the calculation. (4.2 J changes the temperature of 1.0 g of water by 1.0

C.)

1. A pack has been developed so that small amounts of food can be heated without the need to light a fire

or burn fuel. The researchers who designed the pack found that by mixing powdered magnesium metal

with sodium chloride and iron particles the following exothermic reaction would occur rapidly when

water was added:

Mg(s) + 2H2O(l) Mg(OH)2(s) + H2(g) H = -355 kJ

When water is poured into a porous pad containing the powdered magnesium mixture the heat that is

released warms food contained in an adjacent sealed food pouch.

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Assume that the pack contains 1.0 g of powdered magnesium.

(a) Calculate the maximum amount of heat that can be released by adding water to the powdered

magnesium mixture.

(b) Calculate the minimum mass of water that must be added to release this amount of heat.

(c) The directions on the pack specify the addition of a certain volume of water. If 50 mL of water

is added, calculate its maximum rise in temperature. (4.18 J raises the temperature of 1.0 mL ofwater by 1.0 C.)

1. The enthalpy of solution of anhydrous aluminium chloride, AlC13, is -321 kJ.mol-1.

(a) Write the thermochemical equation to which this value applies.

(b) A technician prepared a solution by dissolving 2.67 g of anhydrous aluminium chloride in

enough water to make 200 mL of solution.

Calculate the final temperature of the solution if the initial temperature of the water was 20.0

C. Show your working and state any two assumptions you have made. (4.18 J changes the

temperature of 1.0 g of water by 1.0 C)

4.2 FUELS


Fossil fuels provide energy and are feedstock for thechemical industry.

Describe the advantages and disadvantages of theuse of fossil fuels as sources of heat energy,compared with their use as feedstock.

Carbon dioxide and water are produced by the completecombustion of compounds containing carbon andhydrogen.

Write balanced equations for the completecombustion of fuels in which the only productsare carbon dioxide and water.

The products of the incomplete combustion of fossil

fuels include carbon (soot) and carbon monoxide. Sootand carbon monoxide are harmful to the environment.

Describe the undesirable consequences of

incomplete combustion.

Fuels can be compared on the basis of the quantity ofheat released.

Calculate the quantities of heat evolved per mole,per gram, and per litre (for liquids) for thecomplete combustion of fuels.

Fossil Fuels

The most common fuels in use today are the fossil fuels oil, gas and coal. These materials have been formed

over millions of years by the anaerobic breakdown of plants and animals buried beneath the earths surface.

They are contain mainly hydrocarbons like methane, ethane etc and coal contains mainly carbon.

Advantages of fossil fuels

Coal abundant, cheap

Oil abundant, easily mined and transported because it is a liquid, many uses as fuel, raw material for

plastics

Gas easily extracted, purified and treated; burns cleanly, high energy output

Disadvantages of fossil fuels

They are all non-renewable and are a valuable feedstock for other industries like plastics.

Coal Burns giving off carbon dioxide (Greenhouse gas) impurities of sulfur produce sulfur dioxide (acid

rain) when coal is burnt, high temperature combustion can lead to the formation of nitrogen oxides

(acid rain, photochemical smog, hole in the ozone layer), releases particular matter which causerespiratory problems.

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Oil Produces carbon dioxide, sulfur dioxide and nitrogen oxides when burning; produces soot, unburnt

hydrocarbons contribute to photochemical smog.

Gas produces carbon dioxide, nitrogen oxides.

Combustion of fuels

Complete combustion of fuels can only occur when the supply of oxygen for the combustion is not limited.

This releases the maximum amount of energy. The products are carbon dioxide and water.

If the amount of oxygen is limited, incomplete combustion occurs. This is less efficient, and releases less

oxygen. The products are carbon monoxide, carbon (soot) and water in varying proportions depending on the

available oxygen.

Complete combustion C6H12(g) + 9O2(g)6CO2(g) + 6H2O(l) H = -3916 kJ mol-1

Incomplete combustion C6H12(g) + 6O2(g)6CO(g) + 6H2O(l) H = -2224 kJ mol-1

Incomplete combustion C6H12(g) + 3O2(g)6C(s) + 6H2O(l) H = -1558 kJ mol-1

Harmful effects of the products of incomplete combustion.

Carbon monoxide displaces oxygen from haemoglobin in the blood. Since haemoglobin is the mainmechanism for the transport of oxygen in the blood, the absorption of carbon monoxide deprives the body of

oxygen producing symptoms varying from impairment of judgement and visual perception (at 10 ppm in air)

to death (at >750 ppm).

Soot is a microcrystalline form of carbon produced by incomplete combustion. It produces visual pollution

(smoke and soot), blocks air and fuel inlets in burners and deposits on leaves causing impaired

photosynthesis. Soot in the lungs can lead to respiratory problems like bronchitis, asthma and emphysema.

Unburnt hydrocarbons contribute to photochemical smog.

Heat energy from fuels

While the energy produced by a particular combustion reaction is usually quotes in terms of the molarenthalpy of combustion, it can be quoted in terms of heat per gram (energy density) or heat per litre.

Example hexane (C6H14), H(Combustion) = -4158 kJ mol-1, molar mass = 86 g mol-1, density = 655 g L-1

Energy density =

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= 48 kJ g-1

Energy per litre =

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= 31668 kJ L-1

Fuels like methane have a high heat density, but low heat per litre, while octane has a lower heat density but

because of its greater density, it gives more heat per litre.

EXERCISE 4.2 : FUELS

1. Fossil fuels are the main energy sources used today.

(a) What is a fossil fuel?

(b) What are fossil fuels made from?

(c) What are the advantages of using

(i) Coal

(ii) Oil

(iii) Gas

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(d) What are the disadvantages of using

(i) Coal

(ii) Oil

(iii) Gas

1. Write a balanced equation for the complete combustion of butane, C4H10.

2. When using a bunsen burner the flame is controlled by the amount of gas and whether the air hole is

open.

(a) When the air hole is open,

(i) Describe the flame.

(ii) What are the likely products of combustion. Explain

(iii) Write a balanced equation for the reaction.

(b) When the air hole is closed,

(i) Describe the flame.

(ii) What are the likely products of combustion. Explain

(iii) Write a balanced equation for the reaction.

1. Describe the undesirable consequences of incomplete combustion of a fossil fuel.

2. Complete the following table

Fuel H(combustion) (kJ mol-1)

Molecular

formula

Molarmass

Density(g/L)

Heat/gram

Heat/litre

Methane -890 422

Ethane -1557 546

Propane -2220 582

Butane -2875 601

Octane -5512 685

4.3 ELECTROCHEMISTRY


Electrochemical cells are conveniently divided intogalvanic cells, which produce electrical energy fromspontaneous redox reactions, and electrolytic cells,which use electrical energy from an external source to

cause a non-spontaneous chemical reaction.

Identify a cell as galvanic or electrolytic, givensufficient information.

Redox reactions can be considered as two half-reactions, one involving oxidation and the otherreduction.

Write half-equations for half-reactions, includingthose in acidic solution, given information aboutthe reactants and the products.

Galvanic and electrolytic cells involve oxidation at theanode and reduction at the cathode, with electrons

being transferred from one electrode to the otherthrough an external circuit.

Identify the anode and the cathode in a galvaniccell or an electrolytic cell, given informationabout the reactants and the products.

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Galvanic cells are commonly used as portable sourcesof electric currents.

Identify the:

charge on the electrodes;

direction of electron flow;

movement of ions in the salt bridge orelectrolyte;

given a sketch for a galvanic cell and information

about electrode reactions.

Fuel cells are galvanic cells in which the electrodereactants are available in continuous supply.

State the advantages and disadvantages of fuelcells compared with other galvanic cells.

Some galvanic cells can be recharged by using anexternal electrical supply to reverse the electrodereactions.

Describe the complementary nature of thecharging and discharging of rechargeablegalvanic cells.

Electrolytic cells are used in the production of activemetals.

Describe, with the aid of equations, theelectrolytic production of active metals.

ELECTROCHEMISTRY

Electrochemical cells

An electrochemical cell comprises the following components:

Two electrodes, one being the anode, the other the cathode. They can be made of an active metal or an

inert electrode.

An electrolyte which can be an aqueous solution containing ions or a molten ionic compound

Metal wire that connects the two electrodes

An oxidizing agent in contact with the cathode and a reducing agent in contact with the anode ( it may be

the electrode itself)

There are two types of electrochemical cell

Galvanic cells which use redox reactions to produce a direct electric current. eg dry cell batteries

Electrolytic cells which use electricity in the form of a direct current to bring about electrolysis (non

spontaneous redox reaction). These are used in electroplating and making a number of chemicals (eg salt

water chlorinators in pools)

Redox terminology

Redox is a combination of the words REDuction and OXidation. They must occur simultaneously.

Oxidation is a half reaction in which a species loses electrons. Two examples are:

Mg Mg2+ + 2

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SO2 + 2H2O

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+ 4H+ + 2

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Reduction is a half reaction in which a species gains electrons. Two examples are

Cu2+ + 2

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Cu

I2 + 2

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2

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Oxidising agents (or oxidisers) cause oxidation by accepting electrons and are reduced in the process.

Reducing agents (or reducers) cause reduction by donating electrons and are oxidised in the process.

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Electrolytes are liquids that conducts electricity. Can be a molten ionic salt or aqueous solution of a. ions.

The anode is the electrode at which oxidation occurs.

The cathode is the electrode at which reduction occurs.

Redox pairs are couplings of oxidisers (or reducers) and their reduced (or oxidised) forms. Examples:

Oxidiser Reduced form

Permanganate ion (pink)

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Manganese (II) ion(colourless)

Dichromate ion (orange)

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Chromium (III) ion

(Green)

Hydrogen peroxide H2O2(aq) H2O(l) water

Hypochlorite ion

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Chloride ion

Bromine (orange) Br 2(aq)

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Bromide ion

Iodine (yellow/brown) I2(aq)

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Iodide ion

Reducer Oxidised form

Sulfite ion

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Sulfate ion

Hydrogen sulfide H2S(g) S(s) Sulfur

Hydrogen peroxide H2O2(aq) O2(g) Oxygen (bubbles)

Sulfur dioxide SO2(g)

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Sulfate ion

Bromide

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Br2(aq) Bromine(orange)

Iodide ion

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I2(aq) Iodine (yellow/brown)

Iron (II) ion

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Iron (III) ion

Writing redox equations for reactions in neutral or acidic solutions

The procedure for writing redox equations involves writing equations for the oxidation half reaction and

reduction half reaction and then summing these equations to produce an equation for the overall reaction.

This is illustrated in the following example:

Sulfur dioxide gas is bubbled through an acidified solution containing dichromate ions; the latter is reduced

to chromium (III) ions and the sulfur dioxide is oxidised to sulfate ions. The equation for the overall reaction

is developed as follows:

1. Identification of the redox pairs

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SO2(g)

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2 Develop each redox pair into a balanced halfequation using the following steps in the order shown:

balance any element other than hydrogen or oxygen

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2

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SO2(g)

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balance the oxygen atoms by adding an appropriate number of water molecules to the left or right

hand sides

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2

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+ 7H2O

2H2O + SO2(g)

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balance the hydrogen atoms by adding an appropriate number of hydrogen ions to the left or right

hand sides

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+ 14

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2

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+ 7H2O

2H2O + SO2(g)

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+ 4

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balance the electrical charge by adding an appropriate number of negatively charged electrons to

the left or right hand sides

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+ 14

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+ 6

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2

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+ 7H2O

2H2O + SO2(g)

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+ 4

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+ 2

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3. Combining the two half reaction equations to produce an overall equation involves the following

steps:

if the number of electrons shown in the two half equations are not equal then the half equations

must be multiplied by an integer to make them equal

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+ 14

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+ 6

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2

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+ 7H2O

3x[2H2O + SO2(g)

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+ 4

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+ 2

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]

the two half equations are then added together. Electrons, water molecules and hydrogen ions

appearing on opposite sides are cancelled out

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+ 2

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+ 3SO2(g) 2

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+

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+ 2H2O

check that all atoms and charges balance

Galvanic cells

Galvanic cells are electrochemical cells which produce electrical energy in the form of a direct electriccurrent from spontaneous redox reactions. In these cells, the reducing agent and oxidising agent are not in

direct contact with each other. They are connected via a metal wire and a salt bridge.

A galvanic cell is made of two half cells:

An oxidation half cell called the anode

A reduction half cell called the cathode

The half cells are connected via an external metal wire through which electrons flow

A salt bridge connects the solutions in the two half cells. The salt bridge contains free ions which conduct

ions to complete the circuit. Positive ions migrate towards the reduction half cell, while negative ions

migrate towards the oxidation half cell. Example of the Zn| Zn2+| | Cu2+| Cu cell

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When a cell composed of two metal halfcells is operating, the more reactive of the two metals (in this case

zinc) is oxidised. The zinc electrode is eaten away:

Zn(s)

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+2

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The electrons that are lost flow through the metal wire to the cathode where they are accepted by copper ions

in the solution. The copper ions are reduced to copper metal which deposits on the surface of the electrode:

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+ 2

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Cu(s)

The equation for the overall cell reaction is obtained by adding the oxidation and reduction half equations:

Zn(s) +

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Cu(s) +

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The positive ions in the salt bridge move into the copper half cell to counter the negative charge of the

electrons moving into this half-cell via the conducting wire. The negative ions in the salt bridge move into

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the zinc half cell to replace the negative charge of the electrons moving out of this half-cell via the

conducting wire. This movement of ions constitutes an electric current that completes the electrical circuit.

A salt bridge consists of a concentrated solution of a salt, the ions of which should not react with the ions

present in the solutions in the halfcells. The ions should also not be easily oxidised or reduced. A suitable

salt for the salt bridge of the above cell is sodium chloride. In solution, sodium ions cannot be reduced and

chloride ions can only be oxidised if they are present in very high concentration. Neither ion forms aprecipitate with the ions in the halfcells.

From knowledge of the relative reactivity of metals it is possible to predict the reactions that occur in any

galvanic cell consisting of metal halfcells:

the more reactive metal will be oxidised to its ions

the ions of the less reactive metal will be reduced to the metal

Having established the half reactions occurring in each half-cell, it is possible to deduce other features of the

cell:

the more reactive metal will be become the anode

the less reactive metal will become the cathode

the direction of electron flow in the external metal wire will be from anode to cathode

the direction of movement of positive ions in the salt bridge will be towards to cathode

the direction of movement of negative ions in the salt bridge will be towards to anode

Solution half cells

Galvanic cells can be constructed using inert electrodes in contact with oxidising and reducing solutions.

Inert electrodes (usually platinum or graphite) are used as the electrodes. Example below.

Electron flow

Graphite

Reducer andoxidised form

(

Fe2+

/Fe3+

) Salt bridge

Oxidiser andreduced form

(

MnO4

/H+

)/Mn2+

graphite

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Fuel Cells

In a typical fuel cell, gaseous fuels are fed continuously to an anode (negative electrode) compartment and an

oxidant (such as oxygen from the air) is fed continuously to the cathode (positive electrode) compartment.

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Oxidation and reduction half reactions take place at the anode and cathode respectively to produce an electric

current. eg

Anode : H2 2H+ + 2e-

Cathode : O2 + 2H2O + 4e- 4OH-

The electrodes are porous to increase the surface area and often impregnated with a catalyst.

In the galvanic cells studied so far, the main limitation is that when the chemicals are all used up, the cell

stops producing electricity, eg a battery becomes flat. In fuel cells, there is a continuous flow of reactantsover the electrodes and theoretically it will continue to produce electricity as long as the supply of gases

continues.

The diagram below represents an alkaline hydrogen-oxygen fuel cell, one of the first fuel cells developed.

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In this cell:

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The electrolyte is aqueous potassium hydroxide. The electrode reactions take place under alkaline

conditions. Operating temperature is in the range 50 to 200C.

Hydrogen gas flows over the anode (the negative electrode) where it undergoes oxidation as follows:

H2 + 2OH- 2H2O + 2e

-

The electrons lost flow through the anode and the external circuit to the cathode thereby creating an electric

current.

Oxygen molecules flowing over the cathode accept the electrons that have come from the anode and are

reduced. The hydroxide ions that are produced migrate through the electrolyte towards the anode:

O2 + 2H2O + 4e- 4OH-

The overall cell reaction is:

O2 + 2H2 2H2O H = -286 kJ mol-1

The total output of this cell is electrical energy, heat energy and water.

Advantages and disadvantages of fuel cells

AdvantagesFuel cells have several advantages over conventional galvanic cells and over electricity generators that use

steam driven turbines. These advantages include:

they continuously produce an electric current so long as fuel and oxidant is continuously supplied;

they have high operating efficiency (between 70 and 80%);

they offer a better mass to power output ratio than conventional galvanic cells;

they use readily available fuels and oxidants;

they do not produce pollutant gases such as sulfur dioxide and the oxides of nitrogen;

the electrodes and electrolyte are not consumed during operation;

electrode reaction products are removed as they are formed and do not remain inside the cell;

they require minimal maintenance;

they are silent during operation as there are no moving parts;

in different versions they can be used for small to large scale applications and for stationary or mobile

(particularly transport) power generation.

Disadvantages

The following disadvantages are the focal points for much of the current research being undertaken on fuel

cells:

impurities in the fuel or oxidant can 'poison' the electrode catalysts and/or contaminate the electrolyte.For example, any carbon dioxide present in the oxidant can react with a potassium hydroxide electrolyte

forming potassium carbonate solid that can clog the porous electrodes. This restricts the use of air as an

oxidant.

high purity fuels and oxidants are costly;

medium to high temperatures are needed for the cells to function most effectively;

metal electrode catalysts such as platinum and palladium are very costly;

some of the electrolytes are very corrosive.

Similarities and differences between conventional galvanic cells and fuel cells

The table below summarises the main similarities and differences between conventional galvanic cells

(Commonly called batteries) and fuel cells.

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Similarities between

conventional cells and fuel cells

Differences between conventional cells and fuel cells

Conventional galvanic cells Fuel cells

redox reactions used to produce

direct current

limited quantities of reactants

stored in cell

continuous external supply of

reactants

electrolyte exists between

electrodes

When fully discharged must be

discarded or recharged.

never discharged or run down

No emission of pollutants during

operation

Discarding presents some

environmental problems.

Recharging is time consuming.

cathode positively charged limited 'life' virtually unlimited 'life'

anode negatively charged

energy output decreases as cell

runs down

waste products are removed

continuously

Rechargeable galvanic cells

Rechargeable galvanic cells are also called storage cells. During discharge of these cells, electrode reactions

produce an electric current in an external circuit. At the same time, the oxidiser and reducer inside the cell

are both used up.

During recharge, an external power source provides an electric current to the electrodes. The discharge

reactions are reversed and the original oxidiser and reducer are regenerated. This can be repeated many

times.

The lead-acid cell

The typical lead-acid car battery consists of 6 (2V) lead-acid cells connected in series to give 12 V.

Structure Anode is spongy lead.

Cathode is lead (IV) oxide, PbO2, packed onto lead plates.

Electrolyte is 38% sulfuric acid (density 1.25 g mL-1)

Separators of polymer sheets separate the cells

During discharge

At the anode (-) : Pb(s) +

oleObject69

PbSO4(s) + 2e-

At the cathode (+) : PbO2(s) +

oleObject70

+ 4H+

+ 2e-

PbSO4(s) + 2e-

Sulfuric acid is consumed during the reaction and the density decreases (at 1.1 g mL-1 the battery is flat).

During charging the above reactions are reversed (and sulfuric acid is regenerated).

The nickel-cadmium cell

Structure

Anode is cadmium

Cathode is solid nickel (III) hydroxide, Ni(OH)3, mixed with graphite to improve conductivity.

Electrolyte is potassium hydroxide.

Separators of polymer sheets separate the cells

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During discharge

At the anode (-): Cd(s) + 2

oleObject71

Cd(OH)2(s) + 2e-

At the cathode (+) : 2Ni(OH)3(s) + 2e- 2Ni(OH)2(s) + 2

oleObject72

During recharge these reactions are reversed.

Electrolytic cells

An electrolytic cell is an electrochemical cell in which an external source of electrical energy is used to

bring about a non-spontaneous redox reaction. This process is called electrolysis.

Uses of electrolytic cells:

Producing pure metals from from their compounds;

Recharging batteries;

Electroplating; Refining metals such as copper;

Producing non-metallic elements like chlorine.

The basic components of an electrolytic cell are:

an external source of electrons such as a battery;

an electrolyte

two electrodes (can be active or inert)

an external conductor

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Production of highly reactive metals

Highly reactive metals like sodium and potassium, can only be produced by electrolysis of molten chloride

salts.

A Comparison of galvanic and electrolytic cells

Type Electrode Type of reaction Charge on electrode

Galvanic Anode Oxidation Negative (-)

Cathode Reduction Positive (+)

Electrolytic Anode Oxidation Positive (+)

Cathode Reduction Negative (-)

EXERCISE 4.3 : ELECTROCHEMISTRY

1. The diagram below represents an electrochemical cell.

19

e-e-Power

supply

Anode (+)

2Cl- Cl2

+ 2e

Cathode (-)

Na+ + e Na

Molten salt (eg NaCl)


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oleObject73

(a) Is it a galvanic or electrolytic cell?

(b) What energy change occurs in this cell?

(c) What is the role of -(V)-, (A), (B), (C)

(d) Which metal is more reactive?(e) Write half equations for the reactions at each electrode

(f) Write an equation for the overall reaction

(g) Which electrode is the anode? Why?

(h) Which electrode is positive? Why?

(i) If A contains KCl(aq), which way will the potassium ions move?

(j) Which electrode loses weight? Why?

(k) What happens to an iron nail placed in zinc sulfate solution?

1. Write equations for the half reactions and the overall reaction for (a) and (b) below. Identify each half

reaction as oxidation or reduction.

(a) An acidified solution of potassium dichromate is added to an aqueous solution of sulfite ions.

The dichromate is converted to green chromium (III) ions, Cr3+, while at the same time the

sulfite ion is converted to sulfate ions.

(a) An acidified solution of hydrogen peroxide is added to an aqueous solution of potassium iodide.

The iodide is converted to yellow/brown iodine.

1. For the following combinations of metal half-cells, predict the direction of electron flow and identify

the anode and cathode.

metal half-cell combinations direction of electron flow anode cathode

Zn|Zn2+| |Fe2+|Fe

Ag|Ag+| |Pb2+|Pb

Fe|Fe2+| | Cu2+| Cu

2. Suggest a common use for galvanic/voltaic cells.

3. Fuel cells are become popular energy sources.

(a) What is a fuel cell? Give an example of the reactants used.

(b) What are the electrodes often made from?

(c) What are the advantages of fuel cells?

(d) What are the disadvantages of fuel cells?

1. The lead-acid battery is rechargeable.

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(a) What is the anode made of?

(b) Write an equation for the reaction occurring at the anode.

(c) What is the cathode made from?

(d) Write an equation for the reaction occurring at the cathode.

(e) Write an equation for the overall reaction.

(f) When testing a battery to see if it flat, the density is sometimes measured. Why?

(g) What happens during recharging?

1. One common use of electrolytic cells in the production of metals.

(a) What is an electrolytic cell?

(b) What are some of the other uses of electrolytic cells?

(c) Zinc is prepared by electrolysis of a zinc sulfate solution.

(i) At which electrode is zinc deposited?

(ii) Write a half reaction for the reaction where zinc is deposited.

(iii) Why does the sulfate ion not react at the other electrode?

(iv) Write a half equation for the reaction occurring at this electrode.

1. Elements like sodium and potassium cannot be produced by electrolysis of aqueous solutions of their

ions.

(a) Why?

(b) But they are produced by electrolysis. What is the electrolyte?

4.4 RATE OF REACTION


The time taken for a reaction to reach a specified pointis an indication of the rate of the reaction.

Determine the effect of varying conditions on therate of a given reaction, using experimental data.

The rates of a reaction at different times can becompared by considering the slope of a graph ofquantity (or molar concentration) of reactant or productagainst time.

Draw and interpret graphs representing changesin quantities or concentration of reactants or

products against time.

The rates of a reaction are affected by changes in the:

concentration of reactants;

temperature of the reaction mixture;

pressure of the reaction mixture (for systems involving

gases);state of subdivision of reactants;

presence of catalysts (including enzymes);

intensity of light (for photochemical reactions).

Predict and explain the effect that changes incondition have on rates of reactions in terms ofthe:

frequency of collisions between reactantparticles;

orientation of colliding particles;

energy of colliding particles;

activation energy.

The energy changes in a reaction can be represented byan energy profile diagram.

Draw and interpret energy profile diagrams thatshow the relative enthalpies of reactants and

products, the activation energy, and the enthalpychange for the reaction.

Introduction

One of the prime factors to consider when making chemicals is the speed at which chemical reactions occurand to know how these speeds can be changed. The speed at which reactions occur is known as the rate of

reaction.

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If ways can be found to increase the speed of commercially important reactions without a corresponding

increase in the energy requirements, a considerable saving in an already increasing energy bill would result.

An example of this can be found in the research for a more efficient conversion of solid fuels, e.g. coal, into

liquid or gaseous fuels.

Not only must the chemist consider how to increase the reaction rate, but also find ways of decreasing very

fast reactions to a more manageable rate. Some desirable commercial reactions are too fast to be useful.Control is important.

In chemical reactions, chemical bonds are broken and new bonds form. Studies of the speed at which

reactions occur provide a further understanding of, and theoretical support for, the chemists' ideas of how

atoms bond together.

The rates observed for different reactions vary greatly. Reactions between acids and bases in solution are so

fast that the rate is difficult to measure. Other reactions such as the rusting of iron are slow. At a further

extreme, reactions involving the weathering of rocks are so slow that change may not be observed in a

lifetime.

Measuring Reaction Rate

The rate of a reaction is measured by the change in amount of substance, either reactions or products, divided

by the time taken for that change to occur. If solutions are used, the change in concentration of a species can

be used.

Rate =

oleObject74

=

oleObject75

Rate =

oleObject76

=

oleObject77

Consequently, if one monitors the amount of a reactant over a period of time and draws a graph of the

results, the following graph may depict a typical outcome. This graph shows the decrease in the

concentration of a reactant, A, against time. The straight lines show the rate at three points during the

subsequent reaction. The slope of these tangents gives a measure of the rate

Conditions Affecting Reaction Rates

From experiment, it has been found that the rate of a chemical reaction is

determined by a number of factors. These factors are:

the nature of the reactants, the concentrations of the reactants,

the temperature of the reactants,

the pressure of the reaction mixture for gaseous systems,

the state of subdivision of the reactants, ie the particle size,

Reactions, which are affected by light, are called photochemical reactions.

the amount (intensity) of light falling on the reactants, for reactions affected by light, and

catalysts (including enzymes).

The study of factors that affect rate of reaction is called reaction kinetics.

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Collision theory of chemical reactions

The effects of these factors can be explained in terms of The Collision Theory Of Chemical Reactions. This

theory is base on the premise that, for a reaction to occur

the particles (atoms, ions or molecules) are in constant random motion.

The particles must collide.

The colliding particles must have enough energy to start a reaction.

The rate of the reaction is dependent on

The magnitude of the activation energy.

The frequency of the collisions

The energy of the particles.

The orientation of the particles.

Energy Profile diagrams

The course of a chemical reaction can be represented as an energy profile diagram which shows the relative

potential energy values for the reactants, products and the activation complex. In order to form the activation

complex, the kinetic energy of the colliding particles must exceed the activation energy or the particles just

bounce off.

oleObject78 oleObject79

Energy profile diagram for an exothermic reaction Energy profile diagram for an endothermic

reaction

Factors that affect reaction rate.

Factor Effect on rate Explanation Example

Concentration Increases withconcentration.

More particles per unit volumegives more collisions

Concentrated acids react fasterthan dilute ones

Temperature Increases withtemperature

Particles have more kineticenergy hence a collision ismore likely to produce anactivated species.

Food deteriorates faster inwarm weather, hence the needfor refrigeration

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Factor Effect on rate Explanation Example

Pressure (gases) Increases withpressure

Pressure compresses the gasand so the particles are closertogether hence more collisions.

The Haber processes forceshydrogen and nitrogenmolecules together under 200atmospheres of pressure toincrease the yield

Surface Area Increases withsurface area

Increased surface area exposesmore particles to collisions

Explosions occur in flour millsor coal mines if flames come incontact with fine dust

Catalysts Catalysts canincrease ordecrease the rateof a reactionwithout beingused up in the

process.

Catalysts can provide analternative pathway for areaction. It may provide asurface for the reaction ororientate molecules correctly.Basically it changes theactivation energy.

Platinum catalyses thedecomposition of pollutants inthe catalytic converter of a car.

Intensity of light Increased

intensity of lightincreases the rateof photochemicalreactions

Light increases the energy of

the particles.

Photosynthesis.

Time Reactions slowwith time.

The chemicals are used up andso concentration drops

Reactions slow down and stop.

EXERCISE 4.4 : REACTION RATES

1. Hydrogen peroxide solution can be used as a bleaching agent. It decomposes as shown by the

equation:

2H2O2 2H2O + O2

An experiment carried out to monitor the decomposition of hydrogen peroxide gave the following data

for the concentration of hydrogen peroxide remaining at the end of each time interval shown.

[H2O2] mol L-1 Time (min)

2.32 0

1.86 5

1.49 10

0.98 20

0.62 30

0.25 50

(a) Draw a graph of hydrogen peroxide remaining against time.

(b) Name the dependent and independent variables in this experiment

(c) Use the graph to calculate the instantaneous rate of the reaction:

(i) At 5 minutes

(ii) At 25 minutes

(d) Calculate the average rate of reaction over the first 20 minutes.

(e) On the same set of axes sketch two graphs:

(i) The graph expected if the experiment was conducted at a higher temperature.

(ii) The graph expected if a catalyst had been added before the experiment was started.

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(f) Use the collision theory to explain why the concentration of hydrogen peroxide fell more

rapidly in the first 5 minutes than it did in the 5-minute intervals between 25 and 30 minutes.

1. List the 6 factors that affect reaction rates, how changes in that variable affect the rate and explain it

in terms of collision theory. A table may be a good way to display your answer.

2. Sketch generalized energy profile diagrams for exothermic and endothermic reactions showing the

relative enthalpies of the reactants, the products, the activation energy and the enthalpy change forthe reaction.

4.5 CHEMICAL EQUILIBRIUM


All chemical reactions carried out in a closed system ata fixed temperature eventually reach a state of dynamicequilibrium in which the concentrations of all thereactants and products cease to change with time. Thetotal mass of reactants and products in a closed systemremains constant.

Describe the dynamic nature of a chemicalsystem at equilibrium.

The position of equilibrium in a chemical system at agiven temperature can be indicated by a constant, Kc,related to the concentrations of reactants and products.

Write Kc expressions that correspond to givenreaction equations, and perform calculationsinvolving Kc and equilibrium concentrations inwhich all reacting species are included in theexpression.

The changes in concentrations of reactants andproducts as a system reaches equilibrium can berepresented graphically.

Draw and interpret graphs representing changesin concentration of reactants and products againsttime.

The final equilibrium concentrations for a givenreaction depend on the:

initial concentrations of the reactants and products;

temperature;

value of Kc;

pressure (for systems involving gases).

Calculate the initial and/or equilibriumconcentrations or quantities of reactants and

products, given sufficient information about

a particular system initially and/or at equilibrium.

If a change is made to a system at equilibrium so that itis no longer at equilibrium, a net reaction will occur (if

possible) in the direction that counteracts the change.This is a statement of Le Chteliers principle.

Predict, using Le Chteliers principle, the effecton the equilibrium position of a system of achange in the:concentration of a reactant or product;

overall pressure of a gaseous mixture;

temperature of an equilibrium mixture for whichthe H value for the forward or back reaction is

specified.

CHEMICAL EQUILIBRIUM

Most chemical reactions are reversible. By convention, in a chemical equation, the reactants are placed on

the left and the reactants on the right:

Reactants products

This is called the forward reaction. When the products of a reaction react to for the original reactants, the

following representation is use:

Reactants products

This is called the back reaction.

When forward and back reactions occur at the same time, the following convention is used:

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The magnitude of the equilibrium constant, Kc, provides an indication of the yield of products.

Large Kc values (>10), indicate a high yield at equilibrium position, whereas small value (


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The equilibrium moves to the left.

Decreasing the concentration of a product The removed product is partially replaced byreactants forming more products.

The equilibrium moves to the right.

Example

Consider the following equilibrium:

N2O4(g)

oleObject90

2NO2(g)

If extra N2O4(g) is added to the system, the equilibrium shifts to the right to consume the extra N2O4(g) this

results in the formation of more NO2(g) (at constant temperature and pressure). The concentration changes

with time are illustrated in the diagram below.

oleObject91

Changes to pressure

Pressure changes can cause a shift in the position of an equilibrium if at least one of the reactants or products

is a gas. When the pressure is increased, the volume decreases and a change to the equilibrium position is

possible. In a closed container at a fixed temperature, the pressure is proportional to the total number of

molecules. A reduction of the total number of particles reduces the pressure.

External change Internal change

Increasing the pressure (by decreasing the volume) The equilibrium position is shifted to produce a

smaller number of moleculesDecreasing the pressure of a system (by increasingthe volume)

The equilibrium position is shifted to produce agreater number of molecules

Note:

If an inert gas is added to an equilibrium system to increase the pressure but without increasing the volume,

no change to the position of the equilibrium occurs (since there is no change in concentration).

If there are equal numbers of molecules on the left and right side of the reaction, the eqilibrium position is

not affected by pressure changes.

Example

Consider the following equilibrium:

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oleObject95

Catalysts

Catalysts increase the rate of attainment of equilibrium, they do not affect the equilibrium position because

they affect both forward and backward reactions equally.

EXERCISE 4.5 : EQUILIBRIUM

1. Write expression for the equilibrium constants for the following equilibria:

(a) PCl 5

oleObject96

PCl3 + Cl2

(b) N 2 + O2

oleObject97

2NO

(c) 4NH 3 + 5O2

oleObject98

4NO + 6H2O

(d) N 2 + 3H2

oleObject99

2NH3

2. A mixture of dinitrogen tetroxide, N2O4, and nitrogen dioxide was at equilibrium at 100C.

The equation for the equilibrium is: N2O4

oleObject100

2NO2

The equilibrium concentrations were measured to be[N2O4 ] = 0.485 mol L-1, [NO2 ] = 0.320 molL-1

(a) Calculate K c to this reaction.

(b) In another investigation of the same reaction, equilibrium was again established at 100C. The

concentration of NO2 was found to be 0.0540 mol L-1.

Calculate:

(i) The equilibrium concentration (in mol L-1) of N2O4.

(ii) The concentration (in mol L-1) of NO2 initially admitted to the vessel, given that there was

no N2O4 initially present.

(iii) Using the answers from (i) and (ii) draw a semi-quantitative graph of concentration

versus time from the initial time until sometime after equilibrium has been reached.

1. Changing the conditions of the reaction can change the extent of a reaction.

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(a) State le Chteliers Principle.

(b) What is the effect on the position of an equilibrium if:

(i) The concentration of the reactants is increased?

(ii) The concentration of the products is decreased?

(c) Soft drinks contain carbon dioxide dissolved in water under pressure, so equilibrium exists

between gas dissolved and gas under pressure above the water.

Explain, in terms of le Chteliers principle why an open bottle of soft drink goes flat.

(d) What is the effect on the position of an equilibrium if there is an increase in temperature:

(i) To an exothermic forward reaction.

(ii) To an endothermic backwards reaction.

1. Consider the equilibrium below:

N2(g) + O2(g)

oleObject101

2NO(g)

(a) What effect would an increase in pressure have on this equilibrium?(b) Draw a sketch graph to show the changes in concentrations with time, from the time nitrogen

and oxygen were mixed and the initial equilibrium was established, until after extra oxygen had

been added and a new equilibrium had been established.


H2 + I2

oleObject102

2HI H = +10.9 kJ mol-1

Predict and explain the effect on the position of the equilibrium of an increase in temperature.


4NH3(g) + 5O2(g)oleObject103

4NO(g) + 6H2O(g)

Predict and explain the effect on the position of the equilibrium of a decrease in pressure.


N2(g) + O2(g)

oleObject104

2NO(g)

Predict and explain the effect on the position of the equilibrium of the addition of oxygen at constantvolume and temperature.

4. Fuel-grade ethanol can be produced synthetically by a reaction between ethylene and steam. The

reaction is carried out at 300C and 60 atmospheres pressure, in the presence of a catalyst, as shownby the equation below:

C2H4(g) + H2O(g)

oleObject105

C2H5OH(g) H = -46 kJ mol-1

(a) Explain the effect of the high pressure on the yield of ethanol produced by this reaction.

(b) Methanol, another synthetic fuel, is produced industrially by a reaction between carbon

monoxide and hydrogen, as shown by the equation below:

CO(g) + 2H2(g)

oleObject106

CH3OH(g)

The graph below shows the effect of temperature on the yield of methanol produced by this reaction:

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oleObject107

(i) Write the expression for Kc for the reaction given above.

(ii) State what the graph above shows about the yield of methanol as the temperature is

increased.

(iii) Using your answer to part (ii), and giving reasons, state whether the production of

methanol is endothermic or exothermic.

(iv) T 2 is the preferred operating temperature for this industrial process. Suggest why T2 is

chosen in preference to either T1 or T3.

1. Sulfur dioxide is added to white wines as an antioxidant and bactericide. It reacts with water to give

the hydrogensulfite anion, as shown by the equations below:

SO2(aq)+ H2O

oleObject108

H2SO3(aq)

H2SO3(aq)oleObject109

oleObject110

+

oleObject111

Explain why the amount of sulfur dioxide decreases as the pH of the wine is increased.

2. Exhaust pollutants can be passed through catalytic converters that remove carbon monoxide, nitrogen

dioxide, and hydrocarbons by redox processes.

(a) Nitrogen dioxide may be converted into nitrogen by reduction with carbon monoxide. Write a

balanced equation for this reaction.

(b) Suggest why, in a catalytic converter, the catalyst is supported on a porous material.

(c) Write the equation for the reaction taking place in the catalytic converter which reduces the

nitric oxide concentration

(d) Draw energy profile diagram to show how a catalyst affects the exothermic reaction between

nitrogen dioxide and carbon monoxide.

(e) Nitric oxide (NO) is formed at high temperatures in internal-combustion engines, according to

the following equation:

N2(g) + O2(g) 2NO(g) H= +181 kJ mol-1

(i) State the effect of the high temperature on the rate of the reaction.

(ii) Write the equilibrium expression for this reaction.

(iii) Calculate the value of Kc for the reaction if an equilibrium mixture contains 0.035 mol L-1

of nitrogen, 8.9 x 10-3 mol L-1 of oxygen, and 1.4 x 10-3 mol L-1 of nitric oxide.

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(iv) Suggest why exhaust gases from internal-combustion engines are not equilibrium

mixtures.

4.6 CHEMICAL INDUSTRY


In any industrial chemical process, it is necessary to

select conditions that will give maximum yield in ashort time. This will often involve compromisesbetween conditions that produce the maximum rate,conditions that produce the maximum yield, and costs.

Explain the reaction conditions that will

maximise yield.

The steps in industrial chemical processes can beconveniently displayed in flow charts.

Interpret flow charts and use them for suchpurposes as identifying raw materials, chemicalspresent at different steps in the process, wasteproducts, and by-products.

The siting of a chemical plant will depend upon factorssuch as: access toraw material; waste products; by-

products; access to markets for the product; costs ofenergy, labour, transport, and land; regulatoryrestrictions.

Identify the factors that influence the siting of achemical plant and apply this knowledge to

justify a site for an industrial chemical process,given relevant information.

Chemical Industry

The chemical industry uses chemical reactions on a large scale to convert raw materials into useful products.

The chemical plant is a complex network of components and structures such as reaction vessels, pipes,

pumps, storage tanks, valves and heat exchangers.

The industry employs a range of personel including chemists, technicians, engineers, sales personnel,

accountants, transport workers etc.

For a particular process, the conditions of the reactions are controlled to produce the best yield of product

possible, at an economic rate of formation of products. Conditions that give the best yield may not be thefastest reactions, and so there is a need for compromise between the conflicting factors.

Chemical yield

The yield of a product from a chemical reaction is the quantity obtained.

Percentage yield =

oleObject112

The theoretical yield is the quantity of product predicted by the stoichiometry of the chemical reaction.

Haber Process

The Haber process is the industrial process used throughout the world to produce ammonia.

It involves the reaction of nitrogen and hydrogen at about 400C and 250 atmospheres in the presence of an

iron catalyst. Under these conditions a 45% yield is obtained.

N2(g) + 3H2(g)

oleObject113

2NH3(g) H = -46kJ mol-1

Using le Chteliers principles the following conditions would be expected

High pressurebecause there are more atoms on the left hand side of the equation.

Low temperatures because it is an exothermic reaction.

While the pressure use is high, the temperature used is moderate because while a low temperature would

give a higher yield, the reaction would be too slow.

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The graph below shows the effects of different temperatures and pressures on the yield.

Contact Process

Most sulfuric acid is produced via the Contact Process.

The first step involves producing sulfur dioxide by either burning sulfur or a metal sulfide:

S(s) + O2(g) SO2(g)

2ZnS(s) + 3O2(g) 2ZnO(s) + SO2(g)

The sulfur dioxide is then mixed with oxygen (in air) and passed over the catalyst, vanadium pentoxide

(V2O5), at 450C and atmospheric pressure. The sulfur dioxide is oxidised to sulfur trioxide.

2SO2(g) + O2(g)oleObject114

2SO3(g) H = -99 kJ mol

-1

The sulfur trioxide is absorbed into concentrated sufuric acid solution, forming oleum, H2S2O7, which is

diluted as required to make sufuric acid. This gives a yield of about 90%.

SO3(g) + H2SO4(l) H2S2O7(l)

H2S2O7 + H2O(l) 2H2SO4(l)

The key step is the production of sulfur trioxide. Le Chteliers principle would suggest that the highest yield

would be produced using:

High pressure - because there are less molecules on the right hand side.

Low temperature because the reaction is exothermic

The actual conditions used are:

atmospheric pressure because higher pressure only increases the yield slightly and doesnt warrant the

expense.

relatively high temperature this is a compromise between the yield of sulfur trioxide (higher at low

temperature) and the rate of formation of sulfur dioxide (faster at high temperature).

The effect of different temperatures on the yield of sulfur trioxide at different temperatures is shown in the

diagram below.

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Flow charts

A flow chart is often used to represent the movement of materials through various components of the plant.

These materials are classified as follows:

Raw materials are the materials converted into useful products.

Waste materials are those for which there are no use or market. If they are produced in large quantities

and/or they are toxic, then their disposal can present problems to the operators of the plant.

By-products are products of the process that are not the main products, but have some value or can be used

within the plant.

A flow chart is usually represented in the form of a line and block diagram.

Siting of a chemical plant

A number of factors need to be considered when selecting the site for a chemical plant:

Land cost

Availability of labour

Access to energy and raw materials

Availability of cooling water

Access to markets

Availability of equipment

Regulatory restrictions

Safety

Energy Costs

Energy costs are often a major cost in chemical processes. The can be minimised by:

Reuse of energy - heat from an exothermic reaction can be transferred to other processes

Using lower temperatures catalysts can often be used for this.

Running processes continuously this saves on heating up and cooling down costs

EXERCISE 4.6 : The Chemical Industry

1. Sulfuric acid is made by the Contact Process. A flow diagram for the process is shown below:

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An important step in the commercial manufacture of sulfuric acid is the preparation of sulfur

trioxide from the sulfur dioxide which involves the following equilibrium reaction.

2SO2(g) + O2(g) 2SO3(g); H = -198 kJ

In this industrial process a mixture of air and sulfur dioxide is passed is passed over a finely divided

catalyst bed of V2O5 (vanadium pentoxide) at 450oC to form sulfur trioxide.

(a) State and explain the effect of an increase in pressure on the equilibrium yield of sulfur trioxide.

(b) (i) State whether this formation of sulfur trioxide is an endothermic process.

(i) Explain why the equilibrium yield of sulfur trioxide increases as the temperature is

lowered.

(i) Explain why the reaction is carried out at 450oC despite the fact that the yield of SO3

increases as the temperature is lowered.

(b) The gas issuing from a reaction vessel in the industrial process is found to have the following

equilibrium composition at 450oC.

[SO2] = 3 x 10-4 molL-1 [O2] = 5 x 10-3 molL-1 [SO3] = 2 x 10-3 molL-1

With reference to the chemical equation:

2SO2(g) + O2(g)

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2SO3(g),

calculate the equilibrium constant for the reaction at 450oC.

(c) Vanadium (V) oxide is a catalyst for the industrial process described in the introduction.

(i) Explain how the catalyst increases the rate of formation of sulfur trioxide.

(ii) What effect, if any, does the presence of the catalyst have upon the value of the

equilibrium constant? Explain your answer.

(iii) Explain why it is preferable to have the catalyst in a finely divided state.

1. A flow diagram for the preparation of ammonia is shown below.

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The following thermochemical equation outlines the industrial production of ammonia.

N2(g) + 3H2(g) 2NH3(g) H = -92 kJmol-1

(a) (i) What temperature and pressure conditions would you select to maximise the yield of ammonia?

(i) State two other conditions shown in the flow diagram which aid ammonia formation

(a) What is the name given to this industrial process?

(b) Ammonia is readily separated from the equilibrium mixture by rapidly lowering the

temperature. Briefly explain why this technique is so effective.

(c) State two important uses of ammonia.

1. Ammonia is manufactured in large quantities by the Haber process. The chemical equation which

represents the synthesis of ammonia is given below:

N2(g) + 3H2(g) 2NH3(g)

The position of equilibrium is affected by both the temperature and the pressure at which the process

is carried out. The following graph shows how the percentage yield of ammonia at equilibrium varies

with temperature and pressure.

(a) (i) Use the graphical information to deduce whether the forward reaction is exo or endothermic

giving reasons.

(i) Using Le Chateliers principle, explain the effect that increasing the pressure has on the

yield of ammonia.

(a) In the chemical plant the conditions commonly used are 80-110 atmospheres pressure and a

temperature of 650-720 K. In view of the above graphical data:

(i) Explain why such a relatively high temperature is used in industry.

(ii) Suggest two reasons why a relatively low pressure is actually used.

(b) Write a balanced equation for the formation of ammonia in the Haber process.

(c) What catalyst is used in the Haber process?(d) Use the graphs to suggest what conditions produce the highest percentage of ammonia at

equilibrium when the catalyst is used effectively.

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(e) Suggest why the normal working temperature of many Haber process plants is over 350C.

1. Water gas is a mixture of hydrogen and carbon monoxide. It is made by passing steam over heated

coke.

H2O(g) + C(s) H2(g) + CO(g)

It was used for many years as a commercial fuel.

(a) This reaction was carried out at normal atmospheric pressure. Suggest and explain how

increasing the pressure affects:

(i) The time taken to reach equilibrium.

(ii) The equilibrium yield of water gas.

(b) Write the expression for the equilibrium constant, Kc, for this reaction.

(c) During the conversion of steam into water gas, it was found that 30% of the steam had been

converted. Calculate the value of Kc under these conditions.

1. Explain with examples, the meaning of the following industrial terms:

(a) raw materials

(b) waste products

(c) by-products

(d) plant yield

1. The costs of manufacturing a chemical, the environmental impact, the types of energy required, the

market, the returns and the safety of both workers and residents are some of the factors which must beconsidered before a chemical plant can be set up. Argue a case for and against setting up a small

pharmaceutical plant in your area. What is your conclusion?

4.7 METAL PRODUCTION


The likelihood that an uncombined metal will occurnaturally increases with lack of reactivity.

Predict whether a metal is likely to occur innature uncombined or combined with otherelements, given the relative position of the metalin a table of metal reactivities.

The stages in the production of metals from their oresinclude: concentration of the mineral; conversion of themineral into a compound suitable for reduction;reduction; refinement of the metal.

Identify the stages in the production of a metalfrom its ore, and explain why not all stages arenecessary in the production of some metals.

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The stages in the electrolytic production of zinc fromits ore are: concentration of the zinc mineral;conversion of the zinc mineral into a form suitable forreduction; electrolytic reduction.

Describe, with the aid of equations, theproduction of zinc from its ore.

Electrolysis of molten electrolyte is used in the

reduction stage for the production of more reactivemetals.

Explain why the production of aluminium

requires a molten non-aqueous electrolyte.

Reduction of the oxide using carbon can be used forthe production of less active metals.

Explain why zinc and iron can be obtained byreduction using carbon whereas this is not

possible for aluminium.

The method used in the reduction stage in theproduction of a metal is related to the reactivity of themetal.

Predict the likely method of reduction of a metalcompound to the metal, given the position of themetal in the activity series of metals.

Energy cost is a factor taken into account in theproduction of all metals.

Explain why reduction using electrolysis of anaqueous solution is preferable to electrolysis of a

melt.Metal reactivity

Reactive metals can easily be oxidised (losing electrons) eg

Ca Ca2+ + 2e

To make metals from their ores (metal compounds), they have to reduced (by accepting electrons). However

metals that are easily oxidised, are hard to reduce back to their metallic state.

Metals can be ranked in order of reactivity according to the ease with which they displace hydrogen from

water or acids or to displace one another from an aqueous solution.

The reactivity series below lists metals in order of decreased reactivity.

Metal Reaction with acid Reaction with water Metal ion

Most

reactive

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Least

reactive

K

Ca

Na

react violently with

dilute acids to produce

metal salts and

hydrogen

react vigorously with

liquid water to produce

metal hydroxide and

hydrogen

K+

Ca2+

Na

Most

difficult to

reduce

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Ions easilyreduced

Mg

Al

ZnFe

Ni

Sn

Pb

react with dilute acidsto produce metal salts

and hydrogen

react readily with steam toproduce metal oxides and

hydrogen

Mg2+

A13+

Zn2+

Fe3+

Ni2+

Sn2+

Pb2+

Cu

Hg

Ag

Au

do not react with dilute

acids to produce

hydrogen

do not react with liquid

water or steam.

Cu2+

Hg2+

Ag+

Au3+

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The natural occurrence of metals

Metals occur in the earths crust as compounds called minerals. Some common economically important

minerals are listed below.

metal mineral forms metal mineral forms

K KCl Ni NiSCa CaCO3, CaSO4, CaF2 Sn SnO2

Na NaCl, Na 2CO3 Pb PbS, PbSO4

Mg MgCO3, MgSO4 Cu CuFeS2, Cu2S

Al Al2O3 Hg HgS

Zn ZnS, ZnCO3 Ag Ag, Ag2S

Fe Fe2O3, Fe3O4 Au Au

Metals ores

A metallic ore is a mineral that can be mined at a profitThe production of metals from their ores

The first stage in the process is mining. This can be by open cut or underground mining. This produces

an ore which is a mixture of the metal mineral and waste minerals called gangue.

After mining it is usual to concentrate the ore. Techniques used to produce the concentrate include:

Crushing

Grinding to powder

Froth flotation or chemical leaching

The concentrate often needs to be converted to a substance suitable for treatment.

Reduction of the metal by smelting or electrolysis

Refining of the metal

Not all these steps need be followed however. eg Some ores are quite concentrated and need no further

treatment.

Zinc

Zinc is mined as a sulfide, sphalerite (ZnS). One place where this is commonly done is Broken Hill.

Crushing and Grinding

The first stage in the treatment of the ore is to crush it (to about 2 cm) and then grind it to particles that are

small enough (200 microns) to be separated by froth flotation.

Froth flotation

The powder is mixed with water, frothing agents (oil or detergent) and collecting agents are added, then air is

blown into the mixture. The zinc sticks to the air bubbles and floats to the top where it can be skimmed off.

Roasting the zinc sulfite

The concentrate is roasted in air to make zinc oxide.

ZnS(s) + 3O2(g) 2ZnO(s) + SO2(g)

Leaching of the zinc oxide

The sulfur dioxide from the roasting step can be used to make sulfuric acid which can then be used toconvert the zinc oxide to zinc sulfate solution.

ZnO(s) + H2SO4(aq) ZnSO4(aq) + H2O(l)

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Zinc powder is added to the solution to displace any least reactive metal ion impurities like Ag+, Cd2+ or

Cu2+.

Zn(s) +

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+ Cd(s)

Electrolysis of the zinc sulfate solution

Zinc is extracted from the final solution by electrolysis. The anode is lead the cathode is aluminium.

Anode reaction: 2H2O(aq) O2(g) + 2H+ + 4e

Cathode reaction:

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+ 2e Zn(s)

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Aluminium

Aluminium is mined as the mineral bauxite (Al2O3) in northern Australia eg Weipa.

Bauxite contains iron oxide and silica as impurities and these need to be removed before refining.

The bauxite is mixed with hot concentrated sodium hydroxide solution.

The aluminium oxide dissolves (Note: aluminium is amphoteric)

Al2O3(s) + 2

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+ 2H2O 2

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The silica also dissolves

SiO2(s) + 2

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2

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+ 2H2O

The basic iron oxide does not dissolve and is filtered off.

The aluminium oxide is then crystallised out of the solution with seed crystals.

Because aluminium is so reactive, it cannot be extracted by electrolysis of an aqueous solution. Water would

be reduced first.

The final stage in the industrial production of aluminium involves electrolysis of molten alumina, Al2O3. As

the melting point of alumina is very high (2030C), it is mixed with cryolite (Na3AlF6) and other ionic

compounds to form a mixture of lower melting point (just less than 1000C). This still consumes a lot ofenergy, a lot more than would be used if an aqueous solution could be used.

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Electrolysis is carried out in a carbon-lined steel cathode vessel using carbon anodes. Fresh alumina is added

continuously.

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Anode reaction: Oxide ions are oxidised, but at this temperaure, they burn the carbon electrodes

away and need to be replaced periopdically.

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O2(g) + 4e

then C(s) + O2(g) CO2(g)

Cathode reaction: Aluminium ions are reduced to aluminium metal

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+ 3e Al(l)

Chemical reduction of metals oxides to metals

Metals below aluminium on the reactivity series can be produced industrially by the reaction of the metal

oxides with carbon. Carbon is a strong enough reducer to take oxygen from the ore.

3C(s) + 2Fe2O3(s) 4Fe(l) + 3CO2(g)

C(s) + 2ZnO(s) 2Zn(l) + CO2(g)

Carbon monoxide also reduces metal oxides.

3CO(g) + Fe2O3(s) 2Fe(l) + 3CO2(g)

CO(g) + ZnO(s) Zn(l) + CO2(g)

Summary of reduction methods

Metal Most common method of reduction used to

obtain the metal

Metal ion

Most

reactive

K

Ca

Na

Mg

Al

Electrolysis of molten chloride or molten

oxide

KCl

CaCl2

NaCl

MgCl2

Al2O3

Most

difficult to

reduce

Zn Reduction of metal oxide by carbon or carbon

monoxide.

ZnO

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Least

reactive

Fe

Ni

Sn

Pb

(note about 50% of zinc is produced this way)

Fe2O3

NiO

SnO2

PbO

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Ions easily

reduced

Cu

HgRoasting metal sulfide in oxygen

HgS + O2 Hg + SO2

Cu2S

HgS

Ag

Au

Native metals no reduction required

EXERCISE 4.7 : Metal Production

1. The reactivity of metals can be determined by the reaction of metals with water, dilute acids and by

displacement reactions. Write balanced annotated equations for the following reactions.

(a) Sodium and water

(b) Calcium and hydrochloric acid

(c) Zinc and steam

(d) Iron and copper sulfate

(e) Copper and silver nitrate

2. The metals opposite are arranged in a metal reactivity series.

(a) List the metal(s) likely to occur in nature in the uncombined state.

(b) Explain the difference between the terms ore and mineral.

(c) Complete a table for the metals (except Hg) listed in activity series with the following

headings, Metal, Metal Ion, Method of Production of Metal

(d) Which metal ion would require the most energy to be reduced to the metal?

(e) In chemical reactions what do metals act as?

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Ca

Al

Zn

Pb


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3. This table shows the stages involved in production of a metal from its ore.

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(a) Explain why not all stages are necessary in the production of some metals.

(b) Draw a similar diagram to show the production of zinc from its ore. Include equations for the

chemical reactions.

1. Copper occurs naturally as a mineral, malachite, CuCO3.Cu(OH)2, iron as haematite, Fe2O3 and

aluminium as bauxite, Al2O3.

(a) What is the oxidation stat

senior chem - controlling reactions

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