International Examinations

AS Level and A Level

ChemistryBrian Ratcliff, Helen Eccles, John RaffanJohn Nicholson, David Johnson, John Newman

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ContentsIntroduction vi

1 Atomic structure (AS) 1Discovering the electron 2Discovering protons and neutrons 4Atomic and mass numbers 6Isotopes 7Electrons in atoms 7Ionisation energy 8Electronic configurations 12

2 Atoms, molecules and stoichiometry (AS) 16Counting atoms and molecules 16Determination of Ar from mass spectra 17Counting chemical substances in bulk 18Calculation of empirical and molecular formulae 21Balancing chemical equations 23Balancing ionic equations 24Calculations involving concentrations and

gas volumes 24

3 Chemical bonding and structure (AS) 29Ionic bonding 30Covalent bonding 33Bonds of intermediate character 36Shapes of simple molecules 38Metallic bonding 40Intermolecular forces 41

4 States of matter (AS) 50The three states of matter 50Differences in properties of solids, liquids

and gases 53Comparing the melting and boiling points

of substances 54Why gases liquefy, and solids melt 56Some remarkable substances 58Real and ideal gases 59The behaviour of ideal gases 59The ideal gas law 61The behaviour of real gases 62Measuring the relative molecular mass

of a volatile liquid 63Some lattice structures 65The modern uses of materials 66The effect of structure and bonding on

physical properties 68

5a Chemical energetics (AS) 71Energy transfer: exothermic and

endothermic reactions 71Energy is conserved 73Enthalpy, and enthalpy changes 73Bond making, bond breaking and

enthalpy change 76Measuring energy transfers and

enthalpy changes 77Enthalpy changes by different routes:

Hess’s law 80

5b Chemical energetics (A2) 87The Born–Haber cycle 88Trends in the lattice enthalpy 90The stability of the Group II carbonates

to heat 91

6a Electrochemistry (AS) 94Oxidation states 94Redox: oxidation and reduction 95Electrolysis 95The uses of electrolysis in industry 96

6b Electrochemistry (A2) 99Electrode potentials 100Standard electrode potentials 100Measuring a standard electrode potential 101The meaning of E! values 105Using E! values to predict cell voltages 106The chemical reaction taking place in

an electrochemical cell 107Using E! values to predict whether or not

a reaction will occur 108Limitations of the standard electrode

potential approach 110Reaction rate has a role to play too 111Commercial electrochemical cells 112Michael Faraday and electrolysis 113Determination of the Avogadro constant (L) 114What is liberated during electrolysis? 114How much change occurs during

electrolysis? 115

7a Equilibria (AS) 119Reversible reactions 119Changing conditions: Le Chatelier’s principle 121Crocodiles, blood and carbon dioxide 125Equilibria in organic reactions 126Finding the balance 127Kc and Le Chatelier’s principle 129Equilibrium constants and pressure changes 130The Haber process and calculating Kp 131Using Kc and Kp 132An equilibrium of importance: the Haber

process 133A second equilibrium of importance:

the Contact process 135Acids and their reactions 136Definitions of acids and bases 138The role of water 141Base behaviour and neutralisation 141

7b Equilibria (A2) 146Introducing Kw, the ionic product of water 146Introducing pH 147Ionic equilibria: the definition of Ka and pKa 149Calculating the pH of a weak acid 150Measuring pH 150Acids with alkalis: monitoring change 151Buffer solutions 154Solubility product and solubility 158The common ion effect 159

8a Reaction kinetics (AS) 161Speed, rates and reactions 161Rates of reaction – why bother? 162The collision theory of reactivity 166Catalysis 170Enzymes 172

8b Reaction kinetics (A2) 174More about reaction rates 175The rate equation 176Order of reaction 178Half-life and reaction rates 180Finding the order of reaction using raw data 181Rate constants and temperature changes 184Rate equations – the pay-off 184Catalysis 187Acid rain 191The oxidation of I– by S2O8

2– 192

9 Chemical periodicity (AS) 196Introduction to periodicity 197Versions of the Periodic Table 199

Periodic patterns of physical properties of elements 200

Periodic patterns of first ionisation energies 204Successive ionisation energies and the

Periodic Table 207Reactions of the Period 3 elements 208Preparation of Period 3 oxides 209Preparation of Period 3 chlorides 209The reactions of sodium and magnesium

with water 210

10a Group II (AS) 215Introduction 215General properties of the Group II elements 216Uses 217Reactions of Group II elements 218Chalk and lime chemistry 220

10b Group II (A2) 224The decomposition of the nitrates and

carbonates 224The solubilities of Group II sulphates 224

11 Group IV (A2) 227Introduction 227Bonding and expansion of the octet 228Melting points and electrical conductivities 229The tetrachlorides of the Group IV elements 230Oxides of the Group IV elements 231Oxidation states 232Some uses of Group IV elements and compounds 233

12 Group VII (AS) 237Introduction 237General properties of the Group VII elements 238The reactivity of the halogens: displacement

reactions 239Reactions of the elements with hydrogen 240Disproportionation reactions of chlorine 241Uses 242

13 The transition elements (A2) 245Electronic structures 247Properties of the elements 248Variable oxidation states 249Transition elements as catalysts 250Coloured compounds 250Precipitating transition metal hydroxides 251Complexes 251Redox behaviour 255

14 Nitrogen and sulphur (AS) 259Introduction 259Nitrogen 259Sulphur 262

15 Introduction to organic chemistry (AS) 267Types of formulae 269Functional groups 269Molecular modelling 270Naming organic compounds 271Determination of empirical formulae of

organic compounds 273Practical techniques 274Organising organic reactions 277What is a reaction mechanism? 278Isomerism 280How do we design molecules? 285Routes to new molecules 286Carrying out a synthesis 287A two-stage synthesis 289Designing molecules in industry 289

16a Hydrocarbons: alkanes (AS) 293Physical properties of alkanes 293Chemical properties of alkanes 295

16b Hydrocarbons: alkenes (AS) 299Physical properties of alkenes 300Bonding in alkenes: s and p bonds 300Cis–trans isomerism 300Characteristic reactions of alkenes 301Polymerisation of alkenes 304

16c Hydrocarbons: fuels (AS) 308Sources of hydrocarbons 308Fuels 312

16d Hydrocarbons: arenes (A2) 320Introduction 320Substitution reactions 322Addition of halogens to benzene 325

17 Halogenoalkanes (AS) 328The classification of halogenoalkanes 328Physical properties 329Nucleophilic substitution 329Elimination reactions 331The uses of halogen compounds 331Chemists and the environment 332

18a Hydroxy compounds: alcohols (AS) 335Physical properties of alcohols 335Industrial production of ethanol 336

The reactions of alcohols 336The uses of alcohols 341Structural identification using infrared

spectroscopy 341The tri-iodomethane reaction (A2 only) 343

18b Hydroxy compounds: phenols (A2) 346Phenols and their properties 346The uses of phenol 348

19 Carbonyl compounds (AS) 350Physical properties 352Redox reactions 352Characteristic tests 353

20a Carboxylic acids and derivatives (AS) 357Carboxylic acids 358Esters 360

20b Carboxylic acids and derivatives (A2) 365The acidity of carboxylic acids 365Acyl chlorides 366

21 Nitrogen compounds (A2) 368Primary amines 370Amides 373Amino acids 374Proteins and polypeptides 375Polyamides 376

22a Polymers (AS) 378The formation of polymers 380

22b Polymers (A2) 385Condensation polymerisation 385Polyamide formation 386Proteins and polypeptides 387The relation between the structure of the

monomer and the polymer 388

23 Spectroscopy 390Infrared spectroscopy 390Mass spectrometry 393Analytical applications of the mass spectrometer 394Nuclear magnetic resonance spectroscopy 394

Appendix: Periodic Table of the elements 401

Answers to self-assessment questions 402

Glossary 429

Index 435

CHAPTER 1

Atomic structure (AS)

1 recognise and describe protons, neutrons and electrons in terms of their relativecharges and relative masses;

2 describe the distribution of mass and charge within an atom;

3 describe the contribution of protons and neutrons to atomic nuclei in terms ofatomic number and mass number;

4 deduce the numbers of protons, neutrons and electrons present in both atomsand ions from given atomic and mass numbers;

5 describe the behaviour of protons, neutrons and electrons in electric fields;

6 distinguish between isotopes on the basis of different numbers of neutrons present;

7 explain the terms first ionisation energy and successive ionisation energies of an element in terms of 1 mole of gaseous atoms or ions;

8 explain that ionisation energies are influenced by nuclear charge, atomic radiusand electron shielding;

9 predict the number of electrons in each principal quantum shell of an elementfrom its successive ionisation energies;

10 describe the shapes of s and p orbitals;

11 describe the numbers and relative energies of the s, p and d orbitals for theprincipal quantum numbers 1, 2, 3 and also the 4s and 4p orbitals;

12 deduce the electronic configurations of atoms up to Z = 36 and ions, given theatomic number and charge, limited to s and p blocks up to Z = 36.

By the end of this chapter you should be able to:

Chemistry is a science of change. Over the centuries people have heated rocks, distilledjuices and probed solids, liquids and gases

with electricity. From all this activity we havegained a great wealth of new materials – metals,medicines, plastics, dyes, ceramics, fertilisers,fuels and many more (figure 1.1). But this creationof new materials is only part of the science andtechnology of chemistry. Chemists also want tounderstand the changes, to find patterns of behaviour and to discover the innermost nature ofthe materials.

Our ‘explanations’ of the chemical behaviour ofmatter come from reasoning and model-buildingbased on the limited evidence available from

� Figure 1.1 All of these useful products, and manymore, contain chemicals that have been created byapplying chemistry to natural materials. Chemistsmust also find answers to problems caused whenpeople misuse chemicals.

2 Atomic structure (AS)

experiments. The work of chemists and physicistshas shown us the following:� All known materials, however complicated and

varied they appear, can be broken down into thefundamental substances we call elements. Theseelements cannot be broken down further intosimpler substances. So far, about 115 elementsare recognised. Most exist in combinations withother elements in compounds but some, such asgold, nitrogen, oxygen and sulphur, are alsofound in an uncombined state. Some elementswould not exist on Earth without the artificialuse of nuclear reactions. Chemists have giveneach element a symbol. This symbol is usuallythe first one or two letters of the name of theelement; some are derived from their names inLatin. Some examples are:Element Symbolcarbon Clithium Liiron Fe (from the Latin ferrum)lead Pb (from the Latin plumbum)

� Groups of elements show patterns of behaviourrelated to their atomic masses. A Russianchemist, Dmitri Mendeleev, summarised thesepatterns by arranging the elements into a‘Periodic Table’. Modern versions of the PeriodicTable are widely used in chemistry. (A PeriodicTable is shown in the appendix on page 401 andexplained, much more fully, in chapter 9.)

� All matter is composed of extremely small par-ticles (atoms). About 100 years ago, the acceptedmodel for atoms included the assumptions that(i) atoms were tiny particles, which could not bedivided further nor destroyed, and (ii) all atomsof the same element were identical. The modelhad to give way to other models, as science andtechnology produced new evidence. This evidence could only be interpreted as atomshaving other particles inside them – they havean internal structure.

Scientists now believe that there are two basictypes of particles – ‘quarks’ and ‘leptons’. Theseare the building-blocks from which everything ismade, from microbes to galaxies. For many expla-nations or predictions, however, scientists find ithelpful to use a model of atomic structure thatincludes three basic particles in any atom, the

electron, the proton and the neutron. Protonsand neutrons are made from quarks, and the electron is a member of the family of leptons.

Discovering the electronEffect of electric current in solutions(electrolysis)When electricity flows in an aqueous solution ofsilver nitrate, for example, silver metal appears atthe negative electrode (cathode). This is an example of electrolysis and the best explanationis that:� the silver exists in the solution as positively

charged particles known as ions (Ag+);� one silver ion plus one unit of electricity gives

one silver atom.The name ‘electron’ was given to this unit of electricity by the Irish scientist George JohnstoneStoney in 1891.

Study of cathode raysAt normal pressures gases are usually very poorconductors of electricity, but at low pressures theyconduct quite well. Scientists, such as WilliamCrookes, who first studied the effects of passingelectricity through gases at low pressures, sawthat the glass of the containing vessel oppositethe cathode (negative electrode) glowed when theapplied potential difference was sufficiently high.

� Figure 1.2 Cathode rays cause a glow on thescreen opposite the cathode, and the ‘MalteseCross’ casts a shadow. The shadow will move if amagnet is brought near to the screen. This showsthat the cathode rays are deflected in a magneticfield. The term ‘cathode ray’ is still familiar today,as in ‘cathode-ray oscilloscopes’.

cathode

cross

shadow ofcross

greenglow on

screen

Atomic structure (AS) 3

A solid object, placed between the cathode andthe glow, cast a shadow (figure 1.2). They proposedthat the glow was caused by rays coming from thecathode and called these cathode rays.

For a while there was some argument aboutwhether cathode rays are waves, similar to visiblelight rays, or particles. The most important evidence is that they are strongly deflected in amagnetic field. This is best explained by assumingthat they are streams of electrically charged particles. The direction of the deflection (towardsthe positive pole) shows that the particles in cathode rays are negatively charged.

J. J. Thomson’s e/m experimentThe great leap in understanding came in 1897, atthe Cavendish Laboratory in Cambridge (figures 1.3and 1.4). J. J. Thomson measured the deflection of anarrow beam of cathode rays in both magnetic andelectric fields. His results allowed him to calculatethe charge-to-mass ratio (e/m) of the particles. Theircharge-to-mass ratio was found to be exactly thesame, whatever gas or type of electrodes were usedin the experiment. The cathode-ray particles had atiny mass, only approximately 1/2000 th of themass of a hydrogen atom. Thomson then decidedto call them electrons – the name suggested earlier

by Stoney for the ‘units of electricity’.

Millikan’s ‘oil-drop’ experimentThe electron charge was firstmeasured accurately in 1909 bythe American physicist RobertMillikan using his famous ‘oil-drop’ experiment (figure 1.5). Hefound the charge to be 1.602 ¥10–19 C (coulombs). The mass ofan electron was calculated tobe 9.109 ¥ 10–31 kg, which is1/1837 th of the mass of a hydrogen atom.

cathode

cathode rays

chargedplates(anode)

scale

magnet

� Figure 1.4 A drawing of Thomson’s apparatus. Theelectrons move from the hot cathode (negative)through slits in the anode (positive).

� Figure 1.5 Robert Millikan’s ‘oil-drop’ experiment.Millikan gave the oil drops negative charge byspraying them into air ionised by X-raybombardment. He adjusted the charge on the platesso that the upward force of attraction equalled thedownward force due to gravity, and a drop couldremain stationary. Calculations on the forcesallowed him to find the charges on the drops. Thesewere multiples of the charge on an electron.

light

earth

oilspray

X-rays

� Figure 1.3 Joseph (J. J.) Thomson (1856–1940) usinghis cathode-ray tube.

4 Atomic structure (AS)

Discovering protons andneutrons

New atomic models: ‘plum-pudding’ or ‘nuclear’ atomThe discoveries about electronsdemanded new models foratoms. If there are negativelycharged electrons in all electrically neutral atoms, theremust also be a positively chargedpart. For some time the mostfavoured atomic model was J. J.Thomson’s ‘plum-pudding’, inwhich electrons (the plums)were embedded in a ‘pudding’ ofpositive charge (figure 1.6).

Then, in 1909, came one of theexperiments that changed every-thing. Two members of ErnestRutherford’s research team inManchester University, HansGeiger and Ernest Marsden, wereinvestigating how a-particles (a is the Greek letter alpha) froma radioactive source were scattered when fired at very thinsheets of gold and other metals(figure 1.7). They detected the a-particles by the small flashes oflight (called ‘scintillations’) thatthey caused on impact with afluorescent screen. Since (inatomic terms) a-particles areheavy and energetic, Geiger andMarsden were not surprised thatmost particles passed throughthe metal with only slightdeflections in their paths. Thesedeflections could be explained bythe ‘plum-pudding’ model of theatom, as small scattering effectscaused while the positive a-parti-cles moved through the diffusemixture of positive charge andelectrons.

� Figure 1.6 J. J. Thomson’s ‘plum-pudding’ model of the atom. Theelectrons (plums) are embedded in a sphere of uniform positivecharge.

– electron

� Figure 1.7 Geiger and Marsden’s experiment, which investigatedhow a-particles are deflected by thin metal foils.

a A drawing showing the arrangement of the apparatus.b Ernest Rutherford (right) and Hans Geiger using their apparatus for

detecting a-particle deflections. Interpretation of the results ledRutherford to propose the nuclear model for atoms.

a-particlesource

leadshield

moveable microscope forobserving scintillations

metal foilmost particles

strike here

ZnS a-particledetector

a

b

Atomic structure (AS) 5

However, Geiger and Marsden also noticed somelarge deflections. A few (about one in 20 000) wereso large that scintillations were seen on a screenplaced on the same side of the gold sheet as thesource of positively charged a-particles. This wasunexpected. Rutherford said: ‘it was almost asincredible as if you had fired a 15-inch shell at apiece of tissue paper and it came back and hit you!’

The plum-pudding model, with its diffuse positive charge, could not explain the surprisingGeiger–Marsden observations. However,Rutherford soon proposed his convincing nuclearmodel of the atom. He suggested that atoms consist largely of empty space and that the mass isconcentrated into a very small, positively charged,central core called the nucleus. The nucleus isabout 10 000 times smaller than the atom itself –similar in scale to a marble placed at the centre ofan athletics stadium.

Most a-particles will pass through the emptyspace in an atom with very little deflection. Whenan a-particle approaches on a path close to anucleus, however, the positive charges stronglyrepel each other and the a-particle is deflectedthrough a large angle (figure 1.8).

Nuclear charge and ‘atomic’ numberIn 1913, Henry Moseley, a member of Rutherford’sresearch team in Manchester, found a way of comparing the positive charges of the nuclei ofelements. The charge increases by one unit fromelement to element in the Periodic Table. Moseleyshowed that the sequence of elements in theTable is related to the nuclear charges of theiratoms, rather than to their relative atomic masses(see page 199). The size of the nuclear charge wasthen called the atomic number of the element.Atomic number defined the position of the ele-ment in the Periodic Table.

Particles in the nucleusThe protonAfter he proposed the nuclear atom, Rutherfordreasoned that there must be particles in thenucleus which are responsible for the positivenuclear charge. He and Marsden fired a-particlesthrough hydrogen, nitrogen and other materials.They detected new particles with positive chargeand the approximate mass of a hydrogen atom.Rutherford eventually called these particles protons. A proton carries a positive charge of1.602 ¥ 10–19 C, equal in size but opposite in sign

to the charge on an electron. It has amass of 1.673 ¥ 10–27 kg, about 2000times as heavy as an electron.

Each electrically neutral atom hasthe same number of electrons outsidethe nucleus as there are protonswithin the nucleus.

The neutronThe mass of an atom, which is concen-trated in its nucleus, cannot dependonly on protons; usually the protonsprovide around half of the atomicmass. Rutherford proposed that thereis a particle in the nucleus with amass equal to that of a proton butwith zero electrical charge. Hethought of this particle as a protonand an electron bound together.

Without any charge to make it ‘perform’ in electrical fields, detection

a-particlesfrom asource

deflecteda-particles

dense andmassivenucleus

nodeflection

+

� Figure 1.8 Ernest Rutherford’s interpretation of theGeiger–Marsden observations. The positively charged a-particlesare deflected by the tiny, dense, positively charged nucleus. Mostof the atom is empty space.

6 Atomic structure (AS)

of this particle was very difficult. It was not until12 years after Rutherford’s suggestion that, in1932, one of his co-workers, James Chadwick, produced sufficient evidence for the existence of anuclear particle with a mass similar to that of theproton but with no electrical charge (figure 1.9).The particle was named the neutron.

Atomic and mass numbersAtomic number (Z )The most important difference between atoms ofdifferent elements is in the number of protons inthe nucleus of each atom. The number of protonsin an atom determines the element to which theatom belongs. The atomic number of an elementshows:� the number of protons in the nucleus of an

atom of that element;� the number of electrons in a neutral atom of

that element;� the position of the element in the Periodic Table.

Mass number (A ) It is useful to have a measure for the total numberof particles in the nucleus of an atom. This iscalled the mass number. For any atom: � the mass number is the sum of the number of

protons and the number of neutrons.

Behaviour of protons, neutrons and electrons inelectric fieldsThe three particles behave differently in an elec-tric field, because of their relative masses andcharges. Protons are attracted to the negative poleand electrons are attracted to the positive pole.Because they are much lighter, electrons aredeflected more. Neutrons are not deflected as theyhave no charge (figure 1.10).

Summary tableParticle name Relative mass Relative chargeelectron negligible –1proton 1 +1neutron 1 0

� Figure 1.9a Using this apparatus, James Chadwick discovered the neutron.b Drawing of the inside of the apparatus. Chadwick bombarded a block of beryllium with a-particles (42He). No

charged particles were detected on the other side of the block. However, when a block of paraffin wax (a compound containing only carbon and hydrogen) was placed near the beryllium, charged particles weredetected and identified as protons (H+). Alpha-particles had knocked neutrons out of the beryllium, and inturn these had knocked protons out of the wax.

a-particles

source ofa-particles

beryllium paraffinwax block

detector ofchargedparticles

neutrons(no charge)

protons(≈ charge)

aa b

e–

p+

n

+ –

beam of particles

� Figure 1.10 The behaviour of protons, neutronsand electrons in an electric field.

Atomic structure (AS) 7

IsotopesIn Rutherford’s model of the atom, the nucleusconsists of protons and neutrons, each with amass of one atomic unit. The relative atomicmasses of elements should then be whole numbers. It was thus a puzzle why chlorine has arelative atomic mass of 35.5.

The answer is that atoms of the same elementare not all identical. In 1913, Frederick Soddy proposed that atoms of the same element couldhave different atomic masses. He named suchatoms isotopes. The word means ‘equal place’, i.e. occupying the same place in the Periodic Tableand having the same atomic number.

The discovery of protons and neutronsexplained the existence of isotopes of an element.In isotopes of one element, the number of protonsmust be the same, but the number of neutronsmay be different.

Remember:

atomic number (Z) = number of protonsmass number (A) = number of protons

+ number of neutrons

Isotopes are atoms with the same atomic number, but different mass numbers. The symbolfor isotopes is shown as

mass number Aatomic number X or ZX

For example, hydrogen has three isotopes:Protium, Deuterium, Tritium,

11H 1

2H 13H

protons 1 1 1neutrons 0 1 2

It is also common practice to identify isotopes byname or symbol plus mass number only. For example, uranium, the heaviest naturally occurring element (Z = 92), has two particularlyimportant isotopes of mass numbers 235 and 238.They are often shown as uranium-235 and ura-nium-238, as U-235 and U-238 or as 235U and 238U.

Numbers of protons, neutrons and electrons It is easy to calculate the composition of a particular atom or ion:

number of protons = Znumber of neutrons = A – Z

number of electrons in neutral atom= Z

number of electrons in positive ion = Z – charge on ion

number of electrons in negative ion = Z + charge on ion

For example, magnesium is element 12; it is inGroup II, so it tends to form doubly charged (2+)ions. The ionised isotope magnesium-25 thus hasthe full symbol 25

12Mg2+, and

number of protons = 12number of neutrons = 13number of electrons = 10

SAQ 1.1a What is the composition (numbers of electrons,

protons and neutrons) of neutral atoms of thetwo main uranium isotopes, U-235 and U-238?

b What is the composition of the ions of potassium-40 (K+) and chlorine-37 (Cl–)?

(Use the Periodic Table, page 401, for the atomicnumbers.)

Electrons in atomsElectrons hold the key to almost the whole ofchemistry. Protons and neutrons give atoms theirmass, but electrons are the outer part of the atomand only electrons are involved in the changesthat happen during chemical reactions. If weknew everything about the arrangements of electrons in atoms and molecules, we could predict most of the ways that chemicals behave,purely from mathematics. So far this has provedvery difficult, even with the most advanced computers – but it may yet happen.

SAQ 1.2Suggest why the isotopes of an element have thesame chemical properties, though they have different relative atomic masses.

What models are currently accepted about howelectrons are arranged around the nucleus? Thefirst simple idea – that they just orbit randomly

8 Atomic structure (AS)

around the nucleus – was soon rejected.Calculations showed that any moving, electricallycharged particles, like electrons, would loseenergy and fall into the nucleus.

A model you may have used considers the electrons to be arranged in shells. These ‘shells’ correspond to different energy levels occupied bythe electrons.

Arrangements of electrons: energy levels and ‘shells’ There was a great advance in atomic theory when,in 1913, the Danish physicist Niels Bohr proposedhis ideas about arrangements of electrons in atoms.

Earlier the German physicist Max Planck hadproposed, in his ‘Quantum Theory’ of 1901, thatenergy, like matter, is ‘atomic’. It can only betransferred in packets of energy he called quanta;a single packet of energy is a quantum. Bohrapplied this idea to the energy of electrons. Hesuggested that, as electrons could only possessenergy in quanta, they would not exist in a stableway, anywhere outside the nucleus, unless theywere in fixed or ‘quantised’ energy levels. If anelectron gained or lost energy, it could move tohigher or lower energy levels, but not somewherein between. It is a bit like climbing a ladder; youcan only stay in a stable state on one of the rungs.You will find that, as you read more widely, there

are several names given to these energy levels. Themost common name is shells.

Shells are numbered 1, 2, 3, 4, etc. These num-bers are known as principal quantum numbers(symbol n). Such numbers correspond to the num-bers of rows (or Periods) in the Periodic Table.

We can now write the simple electronicconfigurations as shown in table 1.1. Remember thatthe atomic number tells us the number of electronspresent in an atom of the element. For a given element, electrons are added to the shells as follows:� up to 2 electrons in shell 1;� up to 8 electrons in shell 2;� up to 18 electrons in shell 3.Some of the best evidence for the existence ofelectron shells comes from ionisation energies.

Ionisation energyWhen an atom loses an electron it becomes a posi-tive ion. We say that it has been ionised. Energy isneeded to remove electrons and this is generallycalled ionisation energy. More precisely, the firstionisation energy of an element is the amount ofenergy needed to remove one electron from eachatom in a mole of atoms of an element in thegaseous state.

The general symbol for ionisation energy is DHi

and for a first ionisation energy it is DHi1. Theprocess may be shown by the example of calciumas:

Ca(g) Æ Ca+(g) + e–; DHi1 = +590 kJ mol–1

(If the symbols seem unfamiliar at this stage, seepage 73.)

The energy needed to remove a second electronfrom each ion in a mole of gaseous ions is the second ionisation energy. For calcium:

Ca+(g) Æ Ca2+(g) + e–; DHi2 = +1150 kJ mol–1

Note that the second ionisation energy is muchlarger than the first. The reasons for this are discussed on page 9.

We can continue removing electrons until onlythe nucleus of an atom is left. The sequence of first,second, third, fourth, etc. ionisation energies (orsuccessive ionisation energies) for the first 11 elements in the Periodic Table are shown in table 1.2.

� Table 1.1 Electronic configurations of the first 11elements in the Periodic Table.

Number of electrons in shell

Atomic number n = 1 n = 2 n = 3

H 1 1

He 2 2

Li 3 2 1

Be 4 2 2

B 5 2 3

C 6 2 4

N 7 2 5

O 8 2 6

F 9 2 7

Ne 10 2 8

Na 11 2 8 1

Atomic structure (AS) 9

We see that the following hold for any one element:� The ionisation energies increase. As each

electron is removed from an atom, the remaining ion becomes more positivelycharged. Moving the next electron away fromthe increased positive charge is more difficultand the next ionisation energy is even larger.

� There are one or more particularly large riseswithin the set of ionisation energies of each element (except hydrogen and helium).

Ionisation energies of elements are measuredmainly by two techniques:� calculating the energy of the radiation causing

particular lines in the emission spectrum of theelement;

� using electron bombardment of gaseous elements in discharge tubes.

We now know the ionisation energies of all of theelements.

These data may be interpreted in terms of theatomic numbers of elements and their simpleelectronic configurations.

Before doing so, we must consider the factorswhich influence ionisation energies.

Factors influencing ionisation energiesThe three strongest influences on ionisation energies of elements are the following: � The size of the positive nuclear charge

This charge affects all the electrons in an atom.

The increase in nuclear charge with atomicnumber will tend to cause an increase in ionisation energies.

� The distance of the electron from the nucleusIt has been found that, if F is the force of attraction between two objects and d is the distance between them, then

F is proportional to 1/d2

(the ‘inverse square law’)

This distance effect means that all forces ofattraction decrease rapidly as the distancebetween the attracted bodies increases. Thusthe attractions between a nucleus and electronsdecrease as the quantum numbers of the shellsincrease. The further the shell is from thenucleus, the lower are the ionisation energiesfor electrons in that shell.

� The ‘shielding’ effect by electrons in filled inner shellsAll electrons are negatively charged and repeleach other. Electrons in the filled inner shellsrepel electrons in the outer shell and reducethe effect of the positive nuclear charge. This iscalled the shielding effect. The greater theshielding effect upon an electron, the lower isthe energy required to remove it and thus thelower the ionisation energy.

Consider the example of the successive ionisationenergies of lithium. We see a low first ionisationenergy, followed by much larger second and thirdionisation energies. This confirms that lithium

� Table 1.2 Successive ionisation energies for the first 11 elements in the Periodic Table (to nearest 10 kJ mol–1).

Electrons removed1 2 3 4 5 6 7 8 9 10 11

1 H 1310

2 He 2370 5250

3 Li 520 7300 11800

4 Be 900 1760 14 850 21000

5 B 800 2420 3660 25 000 32 800

6 C 1090 2350 4620 6220 37 800 47 300

7 N 1400 2860 4580 7480 9450 53 300 64 400

8 O 1310 3390 5320 7450 11000 13 300 71300 84100

9 F 1680 3470 6040 8410 11000 15 200 17 900 92 000 106 000

10 Ne 2080 3950 6120 9370 12 200 15 200 – – – 131400

11 Na 510 4560 6940 9540 13 400 16 600 20 100 25 500 28 900 141000 158 700

10 Atomic structure (AS)

log 1

0 io

nis

atio

n e

ner

gy (k

J m

ol–1

)

5

1 2 4 6 9 10 11

Number of electrons removed

3 5 7 82

4

3

has one electron in its outer shell n = 2, which iseasier to remove than either of the two electronsin the inner shell n = 1. The large increase in ionisation energy indicates where there is achange from shell n = 2 to shell n = 1.

The pattern is seen even more clearly if we plota graph of ionisation energies (y axis) againstnumber of electrons removed (x axis). As the ionisation energies are so large, we must use logarithm to base 10 (log10) to make the numbersfit on a reasonable scale. The graph for sodium isshown in figure 1.11.

SAQ 1.3a In figure 1.11 why are there large increases

between the first and second ionisation energiesand again between the ninth and tenth ionisationenergies?

b How does this graph confirm the suggested sim-ple electronic configuration for sodium of (2,8,1)?

Successive ionisation energies are thus helpful forpredicting or confirming the simple electronicconfigurations of elements. In particular, theyconfirm the number of electrons in the outershell. This leads also to confirmation of the position of the element in the Periodic Table.

Elements with one electron in their outer shellare in Group I, elements with two electrons intheir outer shell are in Group II, and so on.

SAQ 1.4The first four ionisation energies of an element are,in kJ mol–1: 590, 1150, 4940 and 6480. Suggest theGroup in the Periodic Table to which this elementbelongs.

Need for a more complex modelElectronic configurations are not quite so simpleas the pattern shown in table 1.1. You will see inchapter 9 (page 206) that the first ionisation energies of the elements 3 (lithium) to 10 (neon)do not increase evenly. This, and other, variationsshow the need for a more complex model of electron configurations than the Bohr model.

The newer models depend upon an understand-ing of the mathematics of quantum mechanicsand, in particular, the Schrödinger equation andHeisenberg’s uncertainty principle. Explanationsof these will not be attempted in this book, but anoutline of some implications for the chemist’sview of atoms is given.

It is now thought that the following hold:� The energy levels (shells) of principal quantum

numbers n = 1, 2, 3, 4, etc. do not have preciseenergy values. Instead, they each consist of a setof subshells, which contain orbitals with different energy values.

� The subshells are of different types labelled s, p,d and f. An s subshell contains one orbital; a psubshell contains three orbitals; a d subshellcontains five orbitals; and an f subshell contains seven orbitals.

� An electron orbital represents a region of spacearound the nucleus of an atom, within whichthere is a high chance of finding that particularelectron.

� Each orbital has its own approximate, three-dimensional shape. It is not possible to drawthe shape of orbitals precisely. They do not haveexact boundaries but are fuzzy, like clouds;indeed, they are often called ‘charge-clouds’.

Approximate representations of orbitals areshown in figure 1.12. Some regions, where there is

� Figure 1.11 Graph of logarithm (log10) ofionisation energy of sodium against the number ofelectrons removed.