the_chemistry_companion.pdf

COH

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TheChemistry

Companion A C FISCHER-CRIPPS

Like the authors other companion books, The Chemistry Companion provides high quality information in unique one-page-per-topic presenta-tions that do not overburden and distract readers with excessive details. The book offers concise summaries of general chemistry concepts, easily accessible in a convenient, reader-friendly format.

Suitable as an introduction or study guide, this companion presents the minimum of what readers need to know to understand the subject. It emphasizes the physics underlying chemistry. By looking at chemistry processes from a physics point of view, readers better appreciate what is happening from the chemical perspective that is usually found in traditional chemistry books.

The author focuses on the structure of matter, chemical componentsand bonds, the periodic table, states of matter, thermodynamics, reaction rates, carbon chemistry, biochemistry, and chemical, ionic, and electronic equilibria. Each topic is covered in a single-page outline format with just enough detail to enable a good understanding of the subject.

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Provides a physical understanding of chemical concepts

Presents clear explanations of difficult material, working through any inconsistencies in understanding

Uses a convenient format for checking formulas and definitions

Includes self-contained information on each page, assuming little prior knowledge

Chemistry

K11517

ISBN: 978-1-4398-3088-8

9 781439 830888

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The Chemistry Com

panion A C

FISC

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-CR

IPPS

TheChemistry

Companion

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TheChemistry

Companion A C FISCHER-CRIPPS

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This book is dedica industria

ated to Bill Cripps al chemist

ContentsPreface 1. Structure of Matter ......................

1 1 Atoms1.1 Atoms .................................1.2 Bohr Atom ..........................1.3 Energy Levels ....................1.4 Schrdinger Equation ........1.5 The Infinite Square Well ....1.6 The Coulomb Potential ......1.7 Covalent Bond ...................1.8 Ionisation Energy ...............1.9 Electron Affinity ..................1.10 Ionic Bond ..........................1.11 Electronegativity ................1.12 Metallic Bond .....................1 13 Electronic Structure of Solid1.13 Electronic Structure of Solid

2. Chemical Components ...............2.1 Matter .................................2.2 Atomic Weight ....................2.3 Ions ....................................2.4 Molecules ...........................2.5 Mole ...................................2.6 Compounds .......................2.7 Cations ...............................2.8 Anions ................................2.9 Chemical Equation .............2 10 Stoichiometry2.10 Stoichiometry .....................2.11 Example .............................

3. The Periodic Table ......................3.1 Electron Configuration .......3.2 Periodic Law ......................3.3 Periodic Table ....................3.4 Groups ...............................3.5 Energy Levels ....................3.6 Noble Gases ......................3.7 Atomic Size ........................3.8 Covalent Radii ...................3 9 Ionic Radii3.9 Ionic Radii...........................3.10 Ionisation Energy ...............3.11 Electronegativity ................

xiii...................................................1

2

vii

...................................................2

...................................................3..................................................4..................................................5.................................................. 6..................................................7..................................................8..................................................9.................................................10.................................................11................................................ 12................................................13

ds 14ds ............................................ 14

................................................15................................................ 16................................................ 17................................................18................................................ 19................................................20................................................ 21................................................ 22.................................................23................................................ 24

25................................................25.................................................26

................................................27

................................................28

................................................29.................................................30................................................31................................................32................................................33.................................................34................................................ 35

36................................................36................................................37................................................ 38

4. Chemical Bonds ..........................4.1 Chemical Bond ...................4 2 Lewis (Electron Dot) Formu

viii

4.2 Lewis (Electron Dot) Formu4.3 Multiple Bonds ...................4.4 Lewis Single-Bonded Struct4.5 Lewis Multiple-Bonded Stru4.6 Lewis Exceptions to the Oct4.7 Oxidation Number ..............4.8 Oxidation Number Example4.9 Polar Bonds .......................4.10 Hybrid Orbitals ...................4.11 Polarisation ........................4.12 van der Waals Forces ........4.13 Hydrogen Bond ..................

5. States of Matter ...........................5.1 Changes of State ...............5.2 Changes of State of Liquids5.3 Phases of Matter ................5.4 Gases .................................5.5 Solutions ............................5.6 Aqueous Solutions .............5.7 Solubility of Solids ..............5.8 Solubility Equilibrium ..........5.9 Electrolytes ........................5.10 Osmosis .............................5 11 Solids5.11 Solids .................................5.12 Crystalline Lattice Structure5.13 Metallic Solids ....................

6. Chemical Thermodynamics .......6.1 Molecular Energy ...............6.2 Specific Heat Capacity .......6.3 Enthalpy .............................6.4 Heat of Reaction ................6.5 Heat of Reaction ................6.6 Enthalpy of Formation ........6.7 Entropy ...............................6 8 Entropy Calculations6.8 Entropy Calculations ..........6.9 Gibbs Energy .....................6.10 Spontaneous Processes ....6.11 Melting of Ice ......................6.12 Freezing of Water ..............

................................................39................................................ 40lae 41lae ..........................................41............................................... 42tures ....................................... 43ctures .................................... 44tet Rule ...................................45............................................... 46

es ............................................ 47............................................... 48............................................... 49............................................... 50............................................... 51............................................... 52

................................................53

................................................54s/Gases ...................................55................................................ 56................................................ 57................................................58............................................... 59................................................ 60................................................ 61................................................ 62............................................... 63

64................................................64es .............................................65

................................................66

................................................ 67................................................68................................................ 69................................................70................................................71................................................72................................................ 73................................................ 74

75................................................75................................................76................................................77................................................ 78................................................79

6.13 Ice/Water Equilibrium .........6.14 Chemical Equilibrium .........6 15 Statistical Entropy6.15 Statistical Entropy ..............

7. Rates of Reaction ........................7.1 Rates of Reaction ..............7.2 Collision Theory .................7.3 Reaction Mechanism .........7.4 Activation Energy ...............7.5 Nature of Reactants ...........7.6 Concentration .....................7.7 Rate Law ............................7.8 Rates of Reactions .............7.9 Determination of Order ......7 10 Determination of 1st Order R7.10 Determination of 1st Order R7.11 Half-Life Method .................7.12 Temperature ......................7.13 Catalysts ............................

8. Chemical Equilibrium .................8.1 Chemical Equilibrium .........8.2 Law of Chemical Equilibrium8.3 Equilibrium Constant ..........8.4 Le Chateliers Principle ......8.5 Summary of Le Chateliers P8.6 Equilibrium in Gaseous Sys8 7 Solubility of Solids8.7 Solubility of Solids ..............8.8 Factors Affecting Equilibrium8.9 Common Ion Effect ............8.10 Precipitation .......................8.11 Complex Ions .....................

9. Ionic Equilibrium .........................9.1 Electrolytes ........................9.2 Ionisation of Water .............9.3 H+ and OH Concentrations9.4 Acids and Bases ................9.5 BronstedLowry .................9 6 Strength of Acids9.6 Strength of Acids ................9.7 AcidBase Reactions .........9.8 Buffer Solutions ..................9.9 Indicators ...........................9.10 Neutralisation .....................

................................................ 80................................................81

82

ix

................................................82

................................................ 83................................................84................................................85................................................ 86............................................... 87................................................88................................................ 89................................................ 90................................................ 91................................................ 92Rate Law 93Rate Law ................................93................................................ 94................................................ 95................................................96

................................................ 97

................................................98m ............................................. 99...............................................100..............................................101Principle ................................102

stems .................................... 103104...............................................104

m ...........................................105..............................................106..............................................107..............................................108

...............................................109.............................................. 110..............................................111

s .............................................112..............................................113............................................ 114

115.............................................. 115............................................. 116.............................................. 117.............................................. 118..............................................119

9.11 Titration ..............................9.12 Choice of Indicator .............9 13 Hydrolysis

x

9.13 Hydrolysis ..........................9.14 Simultaneous Equilibria .....9.15 Complex Ions .....................

10. Electronic Equilibrium ..........10.1 Oxidation and Reduction ....10.2 Redox Reaction .................10.3 Single-Cell Redox Reaction10.4 Standard Hydrogen Electrod10.5 Standard Electrode Potentia10.6 Spontaneous Redox Reacti10.7 Oxidation Numbers ............10 8 Balancing Redox Half-Reac10.8 Balancing Redox Half Reac10.9 Balancing Redox Reactions10.10 Electrochemical Cell ..........10.11 EMF vs. Concentration ......10.12 Electronic Equilibrium ........10.13 Equilibrium Constant ..........10.14 Lead/Acid Battery ...............10.15 Dry Cell ..............................10.16 Corrosion ...........................10.17 Electrolysis .........................

11. Carbon Chemistry..................11 1 Carbon11.1 Carbon ...............................11.2 Carbon Compounds ...........11.3 Functional Groups ..............11.4 Alkanes ..............................11.5 Alkenes ..............................11.6 Alkynes ..............................11.7 Benzene .............................11.8 Alcohols .............................11.9 Ethers .................................11.10 Aldehydes ..........................11.11 Ketones ..............................11.12 Carboxylic Acids ................11 13 Esters11.13 Esters .................................11.14 Amides ...............................11.15 Amines ...............................11.16 Polymers ............................11.17 Physical and Chemical Prop

..............................................120

..............................................121122..............................................122

..............................................123..............................................124

.............................................. 125.............................................. 126..............................................127 ............................................ 128de ......................................... 129als .........................................130ons .......................................131..............................................132

ctions 133ctions .....................................133s with Oxidation Numbers ..... 134..............................................135............................................. 136..............................................137.............................................. 138.............................................. 139..............................................140..............................................141.............................................. 142

...............................................143144..............................................144

.............................................. 145

.............................................. 146..............................................147..............................................148..............................................149.............................................. 150..............................................151.............................................. 152..............................................153.............................................. 154..............................................155

156...............................................156............................................. 157............................................. 158............................................. 159

perties ...................................160

12. Biochemistry ...............................12.1 Sugars ...............................12 2 Polysaccharides12.2 Polysaccharides ................12.3 Lipids .................................12.4 Proteins ..............................12.5 Nucleic Acids .....................12.6 DNA ...................................12.7 Enzymes ............................12.8 ATP ....................................12.9 Anaerobic Metabolism .......12.10 Aerobic Metabolism ...........12.11 Cyclic Photophosphorylatio12.12 Non-Cyclic Photophosphory12.13 Metabolism ........................I

............................................. 161............................................. 162

163

xi

............................................. 163

............................................. 164

............................................. 165............................................. 166............................................. 167............................................. 168.............................................. 169..............................................170..............................................171n ...........................................172ylation ................................... 173..............................................174

This book is similar to previowhere each topic is covered in a

Preface

where each topic is covered in awith enough detail to provide asubject.

This book emphasises the phyespecially in the first chapter. Iwhat is happening from a physicmay better appreciate what is haperspective that is usually foundbook.

I am indebted to Dr. Ray STechnology Sydney an exemplTechnology, Sydney, an exemplthat anything can be learnedmanageable pieces with attainapresentations in this book haveexcellent undergraduate lectures.

I also thank Hilary Rowe for hbook into print and the editoriTaylor & Francis for their verapproach to the whole publication

Tony Fischer CrippsTony Fischer-Cripps,Killarney Heights, Australia

us Companion style booksa single page outline format

xiii

a single page outline formata good understanding of the

ysics underlying chemistry,hope that by understanding

cs point of view, the readerappening from the chemicald in a traditional chemistry

Sleet of the University oflary teacher who taught melary teacher who taught meif it is broken down intoable goals. Many of the

their origin in Dr. Sleets

er persistence in getting thisial and production team atry professional and helpfuln process.

1 Structure of Matte

1. Structure of Matter

1. Structure of Matte

Summary

222

42

8 nhqZmE

o

een

Energy levels Bohr H-like atom

22242

24 nmqZE

o

en

Energy levels Schrdinger equCoulomb potential:

Covalent bond: the co-sharing

Ionic bond: the electrostatic aafter electrons are transferred f

Metallic bond: free electrons ienergy of the system and so act

The electronegativity describatom, when it combines with anegative by more strongly celectron pair.

er

1

er

m:

uation,

of valence electrons.

attraction between ions formedfrom one atom to another.

in the valence band reduce thet to hold atoms together.

bes the relative ability of ananother atom, to become morecapturing a shared electron or

1.1 Atoms

Sixth century BC Thales of Miletoswhich all things are made is water.

2

matter is composed of earth, air, wphilosophers (Leucippus, Democritusdivide a piece of matter again and agawhere no further subdivision could takwas called the atom. About 300 BC, Aargues that matter is based upon theadds the qualities of coldness, hotneconsiderable reputation ensured that hteaching for many hundreds of years.

In the fifteenth century, new advancmade from particles in agreement wmade from particles, in agreement wBoyle taught that matter consists ofcomposed of atoms of the same type.in fixed proportions to form compproposed the atomic theory of maexperimental evidence from thecombination. He created a scale atomwere then known. Daltons referencehydrogen, which was assigned an atogiven atomic masses according to hhydrogen atom.

In 1807, Humphry Davy decomposmetals using electrolysis. In 183quantitative connection between electinto elements in electrolysis. Thesthemselves contain electric chargedemonstrated visible cathode rayscharged electrode (cathode) and tra(anode) in an evacuated tube.

In 1872, Mendeleev arranged elemeand discovered that the propertiesregular intervals When elements werregular intervals. When elements wermass going across from left to right,down, a periodic table was formedthe time, Mendeleev was able to pundiscovered elements from gaps in th

proposes that the basic element from450 BC Empedocles teaches that all

The Chemistry Companion

ater and fire. Around 400 BC, Greeks, Epicurus) proposed that if one couldain, eventually a limit would be reachedke place, this limiting amount of matterAristotle rejects the atomistic view andfour basic elements of Empedocles butss, dryness and moistness. Aristotleshis ideas became embodied in religious

ces in physics suggested that matter waswith the ancient Greek atomists Robertwith the ancient Greek atomists. Robert

different types of elements that wereDifferent elements could join together

pounds. Later, in 1803, John Daltonatter that was based on quantitativeweighing of different elements in

mic mass for the different elements thate atom was the lightest element known,omic mass of one. Other elements werehow heavy they were compared to a

sed potash into sodium and potassium32, Michael Faraday discovered atricity and the separation of compoundse observations suggested that atoms

e. Experiments by William Crookess that emanated from a negativelyavelled towards the positive electrode

ents in increasing order of atomic massof certain elements were repeated atre ordered in columns with the atomicre ordered in columns with the atomicand similar chemical properties going

whereby, using the known elements atpredict the properties of some as-yethe table.

1.2 Bohr Atom

In 1897, Thomson demonstrated thatthe cathodes of vacuum tubes were in


electrons. Thomson proposed that atsphere within which were embedded n

Rutherford found in 1911 that thedistance from a central positivelyelectrons orbited the nucleus and thnucleus and the electron was balancethe orbital motion. However, a majorthe case, then the electrons would celectromagnetic waves and very quic

In 1913, Bohr postulated two importo Rutherfords theory of atomic struc1. Electrons of mass me can orbit the

nucleus at radius r with velocity vwhat are called stationary states iwhich no emission of radiation occurs and in which the angular momentum L is constrained to havalues:

2nhvrmL e

The 2 appears bL is expressed in rather than f.

2 Electrons can make transitions fro

By summing the kinetic energy (froand the potential energy from the elean electron at a given energy level n is

2. Electrons can make transitions frothe emission or absorption of a sinbeing the absorption and emission

rvm

rq ee

o

2

2

2

41

As in the Rutherford atom, the centrifuforce is balanced by Coulomb attractio

with the addthat: vrme

g gy

222

42

8 nhqZmE

o

een

Note: atomic numthe energy of the glevels for each statmodel, the energy electron atoms can

t the rays observed to be emitted fromn fact charged particles which he called

3

toms consisted of a positively chargednegatively charged electrons.electrons were actually located at somecharged nucleus. He proposed that

he electrostatic attraction between theed by the centrifugal force arising from

problem with this was that if this wereontinuously radiate all their energy as

ckly fall into the nucleus.rtant additions Lyman

C106.1 19eqcture:

in in

ave

because terms of

y

Balmer

Paschen

n = 1n = 2

n = 3n = 4 n = 5 n = 6

Ehf

m one state to another accompanied by

om the orbital velocity) ctrostatic force, the total energy of s given by:

m one state to another accompanied by ngle photon of energy E = hf , this

spectra observed experimentally.Mechanical model of hydrogen atom

r

fugal on:dition

2nh

me

g ymber Z = 1 for the hydrogen atom where ground state is 13.6 eV. The energy te n rise as Z2. Thus, according to the Bohr level of the innermost shell for multi-

n in principle be several thousand eV.

1.3 Energy Levels

The stationary states or energy levelselectron shells or orbitals, and are

4

corresponding to n = 1. The numbenumber. According to the Bohr moden, but experiments show that in multiof sub-levels (evidenced by fine splittL shell n =2 has two sub-shells, 2s an

n = 3

n = 4

n =

M

NO

3s(2)3p(6)4s(2)3d(10)4p(6)4d(10)4f(14)

1.51 eV1.85 eV0.54 eV

0 eV

n = 1K

Hydrogen atom only

At each value of n the angular mvalues. The number of values is descrThe allowed values of l are 0, 1, (

n = 2L2s(2)

2p(6)

1s(2)13.6 eV

3.39 eV

, , (letter that indicates the sub-shell:

A third quantum number m deschanges in angle of the angular mopresence of an electric field. It takes th

A fourth quantum number descrelectron where the spin can be either

According to the Pauli exclprinciple, no electron in any one atohave the same combination of qunumbers. When all the electrons in anare in the lowest possible energy leveatom is said to be in its ground stateoutermost electrons in an atom arethe valence electrons.

s allowed by the Bohr model are callede labelled K, L, M, N, etc. with K


er n is called the principal quantumel, the electron energy only depends oni-electron atoms, electron shells consistting of spectral lines). For example, thed 2p.

It is convenient to assign the energy at infinity as being 0 since as an electron moves closer to the nucleus, which is positively charged, its potential to do work is less and thus the energy levels for each shell shown are negative. In hydrogen, a single-electron atom, the energies for each shell are given by:

The electron-volt is a unit of energy. 1 eV = 1.602 1019 J.

momentum can take on several distinctribed by the second quantum number l.n1). Each value of l is indicated by a

26.13

nE for hydrogen

Sometimes the splitting of principal shells into sub-shells results in some overlap (e.g. 4s is lower in energy than 3d).

) y

l = 0 sl = 1 pl = 2 dl = 3 fl = 4 gl = 5 h

scribes the allowablementum vector in thehe values l to 0 to +l.ribes the spin of an1/2 or +1/2.lusionm can

uantumn atom n = 3

l 0 1 2 ( d)

For example, the 3d sub-shell can hold up to 10 electrons:

thus:els, thee. Thecalled

l = 0, 1, 2 (s, p, d) m = -2, -1, 0, 1, 2

5 values of m times 2 for spin, thus 10 possible electrons

thus:and:

The total energy E of an electron in akinetic energies. Expressed in terms o

1.4 Schrdinger Equation


Vm

pE 2

2m, this is stated:

Thus: txVm

phf ,2

2

Let

tiE

xip

is a variable, the and value of which information about thmotion of a wave/pa

hfE since

The value both positifor differenisolated at

The solution to the Schrdinger waveis a function of x only, then the wtime-independent and time-dependent

motion of a wave/pa

itxVxm

,

2 222Thus:

The resulting solutions of these equatiwave function:

EV

xm 2

22

2 t

txtx txtx ,The wave function gives all the info

in an atom. is a complex quaninterpreted as a probability densitydetermine the probability of an electand x.

Quantum mechanics is concerned(i.e. solving the Schrdinger equafunctions such as those inside atomstime-independent wave equation ocquantised. The solutions correspond to

Solutions to the Schrdinger equatiowhich are a function of both xphenomena (e.g. the probability of trlevels in an atom) to be calculated and

an atom is the sum of the potential andof momentum, p, and mass of electron

5

form provide

he article.

of the potential function may depend on ion and time. The form of V(x,t) is different nt arrangements of atoms (e.g. a single tom, an atom in a regular array of a crystal).

e equation is the wave function . If Vwave equation can be separated intot equations that can be readily solved.

article.

t Schrdinger equation

ions, when multiplied together, give the

tEie

rmation about the motion of an electron

ntity, the magnitude of which || isfunction which in turn can be used totron being at some position between x

d with determining the wave functionation) for particular potential energy. It is found that valid solutions to theccur only when the total energy iso stationary states.on can be found for potential functionsand t. This enables time-dependent

ransitions of electrons between energyd hence the intensity of spectral lines.

1.5 The Infinite Square Well

Consider an electron that is confinedone-dimensional space L/2 and L/2

6

potential energy of the electron is a ccan be conveniently set to zero). Thisof an infinite square well potential. Tnot usually found around electronsatom, but does often represent thatelectrons in a solid and also in a chem

In the infinite square well potential,distance x = L/2 from the centre posby the infinite potential at the walls.have any value of total energy as it mo

In quantum physics the allowed s

3s2p2s

EnergyElecenerleve

mLnE 2

222

2

In quantum physics, the allowed stotal energy of the electron. That iswhose probability of being in a parstanding wave patterns from the solutthe simple case of zero potential eneare given by:

2L

2L0

1sleve

Note also that the kinetic energy odecreases as the length of the well Lgiven more room to move, then its to

It can be seen that the energy increahas a greater kinetic energy if it exisThe minimum allowable energy isenergy (in this case, 0). This is the zer

total energy is kinetic, since the potenspace is effectively a mechanism forthe system. This is important when thcome near each other during the forma

to be located in2 and where the

Energy


constant (whichs is an exampleThis potential is

in an isolatedexperienced by

mical bond., the electron cannot move more than aition because it is constrained or bound. In classical physics, the electron canoves within this space.

2L

2L0

space

stationary states indicate the possible

ctron rgy

els

n = 1,2,3,4, V(x) = 0

stationary states indicate the possibles, the electron is moving as a particlerticular position x is described by thetions to the Schrdinger equation. Forergy within the well, the energy levels

The electron can move anywhere within the confines L, but can only have

els

of the electron, for a given value of n,L increases. That is, if the electron is

otal energy is lowered (in this case, the

ases as n increases. That is, the electronts in a stationary state with a larger n.greater than the minimum potential

ro point energy.

kinetic energies allowable by the stationary states.

ntial energy was set to be 0). Availableproviding a reduction in total energy of

he valence electrons of different atomsation of chemical bonds.

1.6 The Coulomb Potential

In the infinite square well potentiainside the well was constant, and in


3s

This potential reflects the fact that the electron is attracted more strongly by the nucleus

2p2s

1s

0

Eele

-r

single-electron atom, a more realistiare bound by the Coulomb potential.

attracted more strongly by the nucleusshorter distances r and so the potentialenergy is no longer constant i.e. the potential varies with distance from thenucleus. Mathematically, the Coulombpotential is:

r

ZqrVo

e

42

where Z is the atomnumber, the numbprotons in the nucl

Note that in this potential, the energynegative as r decreases. That is, elehave a higher potential energy.

22242

24 nmqZE

o

en

The total energy levels for an electsolution to the Schrdinger equation

n = 1,2,

Here, the energy E becomes more neelectron (Z = 1) , and at n = 1, we obtCoulomb potential, as drawn abovmovement available to the electron as

When an electron is bound by a poallowable energies That is, the allowallowable energies. That is, the allowAbove this level, the electron is freindicated by the grey band in the figurenergy depends on the principle quenergy becomes less negative and is th

al, the potential energy of the electronndependent of position. In an isolated

7

s at

Electron energy evels

+r

ic scenario is the case where electronsBy convention, an electron is assigned a zero total energy when it is at rest, at an infinite distance from the nucleus. As the negatively charged electron moves towards the positively charged nucleus (opposite to the direction of the electric field), it can acquire kinetic energy and/or do work and so its electrical potential is reduced. That is, the potential energy becomess at

l

e b

mic ber of eus.

y (the potential energy) becomes moreectrons further away from the nucleus

potential energy becomes negative. This is the same convention in electric field theory, where it is usual to consider a positive test charge moving in the same direction of the electric field whereupon potential energy is reduced.

tron for this potential are given by then with V(r) as above and are expressed:

,3,4

egative as n decreases, and for a singletain the total energy E = 13.6 eV. Thee, also reflects the greater range ofr increases.tential, it is forced into having discretewable states E < 0 are bound states

This applies to an isolated single electron atom (e.g. hydrogen). Compare with Bohr model 1.2.

wable states E < 0 are bound states.ee and can have any energy. This isre above. Note that in this potential, theantum number n. As n increases, theherefore at a higher potential.

1.7 Covalent Bond0Consider a hydrogen atom.

Here we have one electron

8

orbiting the nucleus. In theground state, the electron isin one of the available 1senergy states.

Hydro

(groun

When a hydrogen atom comes intowe might be tempted to draw the ener0

Atom #1

But this would beincorrect. What happens is that the electron in atom #nucleus of atom #1 and vice versa.attracted to both electrons. The nucleutwo electrons. Therefore, the two ntogether. The co-sharing of theseattraction of the two atoms is called a

0

Nucleus for atom #1

13.6

0

Atom #1

When this attraction occurs, the siextent that the electrons are completeelectrons (by the Pauli exclusion prin

0.0

electrons (by the Pauli exclusion prinSince this is energetically favourablethen hydrogen naturally forms the mformed, energy is released (heat). Toenergy (18.0813.6 ) = 4.48 eV has

3s2p t i


3s

2p2s

1s

atomic energy levelsgen atom

nd state)

proximity with another hydrogen atom,gy levels like this:

2p2s

13.6 eV

1sAtom #2

#2 is attracted by the positively chargedOverall, the nucleus from atom #1 is

us from atom #2 is attracted to the samenuclei behave as if they were bondedvalence electrons and the resultingcovalent bond.

Nucleus for atom #2

eV

3s2p

2s

1s

Molecularenergy levels

Atom #2

ides of the well are reduced so theely shared between the two atoms. Thenciple) have to have different spins .

74 nm

18.08 eV

nciple) have to have different spins .e (electrons have more space to move),molecule H2. When a covalent bond is

break the bond and separate the atoms,s to be supplied.

1.8 Ionisation Energy

In order for an electron to make a trstate in an atom, it must absorb en


sufficient energy, by heat, etc). The e(highest potential energy) electron into infinity is called the first ionisatiunits of eV. The resulting positively ch

The energy required to move the neoutermost bound electron fromground state to infinity is called tsecond ionisation energy, and so on.

Ehf Bound electronIncoming phot

3s(1)0

2p(6)2s(2)

1s(2)Sodium atomNa

The incoming photon must hasufficient energy (in this examp> 5 1 eV for Na) to lift the electron> 5.1 eV for Na) to lift the electroninfinity distance from the nucleus fthe atom to be ionised. If it does noand if the incoming photon hsufficient energy to lift the electron tohigher energy level, then the atom mbecome excited rather thionised. If the incoming photon hinsufficient energy to excite or ionian atom, then it will just pass througor be scattered by the atom.

ansition from a lower state to a highernergy (via collision with a photon of

9

energy required to move the outermostan isolated atom from its ground state

on energy and is usually expressed inharged atom is called an ion.extits

the

Sometimes ionisation energies are expressed as kJ per mole. Since 1 eV is 1.602 10-19J, then 1 eV per atom is 96.49 kJ per mole.

Free electronton

0

3s(1)

2p(6)2s(2)

1s(2)Sodium ionNa+

avele,to

Important: We cannot just assign energies to these levels by calculation like we did with the h d t Th ito

forot,

haso a

mayhanhasisegh

hydrogen atom. The energies associated with each level in a multi-electron atom depend on the size of the atom, the atomic number, the degree of screening of outer electrons by the inner electrons, whether or not the atom is bonded with another, or exists in the gaseous phase, and most importantly, whether the level is filled or not. The ionisation energies are usually measured experimentally. The formula worked OK for hydrogen because we were dealinghydrogen because we were dealing with just a single electron.

Chlorine atom

Consider a neutral chlorine atom. The

1.9 Electron Affinity

10

0

Nuclechloratom

One vacant state in the 3pvalence shell. Lower states all full with electrons.

2p2s

1s(2)

Chlorine atom

What happens when a free electronppatom? The electron may occupy a vacatom with a full outer energy shell. Tand is called a noble gas configuratinfinity into the 3p shell, its potentishow that 3.7 eV is released (perhapis called the electron affinity of the at

0

3.7 eV

The electron affinity is the energy reinfinity into a neutral atom. Likeexperimentally (although usually ind

Nuclechlor

3p valence shell now full. Lower states all full with electrons.

2p2s

1s(2)

y g yovercome the repulsion of the electrelements like Cl, the attainment of apayoff.

m

e first ionisation energy for Cl is 13 eV.


3p(5)

eus for rine

m

3s(2)p(2)(2)

)

m

n comes into contact with the neutralcancy in the 3p level, thus endowing theThis condition is energetically desirabletion. When the electron is bought fromal energy is lowered and experimentsps as heat). This energy that is releasedtom.

3p(5)

Incoming electron

leased when an electron is bought frome ionisation energy, it is measureddirectly). The incoming electron has to

eus for rine atom

3s(2)p(2)(2)

)

Chlorine ion

3p(5)

y grons already there, but, in the case ofnoble gas configuration is a sufficient

1.10 Ionic Bond

Ionic bonds usually form betweenelectrons. Consider the reaction betw


0

3s(1)

One valence electron in the 3s state. Lower states all full with electrons.

2p(6)2s(2)

1s(2)

Sodium

The ionisation energy for Na is +5.1eV. Thus, if the lone electron in the 3barrier (5.1 3.7 = 1.4 eV), then tCl chlorine ion, leaving each with a n(Coulomb) attraction is called an ionic

0

3s(1)

2p(6)

Sodium ion ++

Coattr

2s(2)3s electron transferred to Cl atom. Lower states all full with electrons.

1s(2)

When the ions are formed, the attmove towards each other and the elThe ions reach an equilibrium disattraction of their overall charge and tattraction of their overall charge and tcharged nuclei. Experiments show thenergy is lower than the 1.4 eV barrtwo atoms is energetically favourablestate of minimum energy).

elements that have unpaired valenceeen sodium and chlorine.

11

3p(5)

2p(2)

0

3s(2)

One vacant state in the 3pvalence shell. Lower states all full with electrons.

2s(2)

1s(2)

Chlorine

eV. The electron affinity for Cl is 3.7s band in Na can climb over the energythis will create an Na+ sodium ion and anet charge qe. The resulting electrostaticc bond.

0

3s(2)2p(2)

Chlorine ion

3p(5)

oulomb raction

2s(2)Vacant state filled by electron from Na atom. Lower states all full with electrons.

1s(2)

traction between them causes them toectrical potential between them drops.tance determined by the electrostaticthe repulsion offered by their positivelythe repulsion offered by their positivelyhat the bond energy is 5.5 eV. Thisrier and so an ionic bond between thesee (since atomic systems tend settle to a

1.11 Electronegativity

It is energetically favourable for anenergy shells (noble gas configurat

12

energy shells (noble gas configuratachieve this configuration becausevalence electrons are easily removedon the other hand, may have only a feprefer to gain electrons to attain a nobWhen two atoms come together to foratom to gain, or pull, an electron froionisation energies and the electron afproperty that includes these telectronegativity of the element.

Heres how it works in simplifiebi t f HCl b th f ticombine to form HCl by the formatio

is shared between the H atom and thThe ionisation energy for H is 13.6affinity of H is 0.75 eV and for Cltransferred from Cl to H, a net energyFor an electron to be transferred from9.9 eV is required. Thus, in thisbiased towards being over near the Cto transfer the electron from H to Csharing makes the bond polar (sincecharge). Although we call the bon

Two atoms with verydifferent electronegativities are

An eelect( l

character as well. In general, there iscovalent depending on the nature of th

The electronegativity describes thcombines with another atom, to becocapturing a shared electron or electromeans, but at its simplest, dependselectron affinity of the atom. ElectronFl on the right-hand side of the periothe left-hand side of the periodic table

different electronegativities areexpected to form ionic bonds.Atoms with much the sameelectronegativities are expectedto form covalent bonds.

(elecexpechlor(withexpe

atom to have completely filled outertion) Metals tend to lose electrons to


tion). Metals tend to lose electrons totheir few loosely bound outer shell

d (low ionisation energy). Non-metals,ew vacancies in their outer shells and soble gas configuration.rm a chemical bond, the ability for one

om the other atom depends on both theffinities of the two atoms. A combinedtwo characteristics is called the

d terms. Hydrogen and chlorine canf l t b d A l t ion of a covalent bond. An electron pair

he Cl atom. But is this sharing equal?eV, and for Cl is 13 eV. The electronis 3.7 eV. Thus, for an electron to bey of 13 0.75 = 12.25 eV is required.m H to Cl, a net energy of 13.6 3.7 =

covalent bond, the shared electron isCl atom because less energy is requiredCl compared to Cl to H. This unequalone end, the Cl end, has a net negative

nd covalent it does have an ionic

example is when an Na atom (with tronegativity 0.9) meets a Cl atom

i i 3 0) Th b d b h i

a gradation of bond types from ionic tohe atoms.e relative ability for an atom, when it

ome more negative by more stronglyon pair. It is measured by a variety ofon both the ionisation energy and the

negativities (no units) range from 4 (fordic table) down to

1.12 Metallic Bond

Consider what happens when we hamight be first tempted to draw the pot


This does nothappen in a solid.In a solid the electrons in atom #1 are attracted to both its own nucleus and also to some extentby the nucleus of atom #2. For exampof atom #1 is also captured to somed h P li l i i i l

0

Nucleus foatom #1

The solution tothe Schrdinger equation predicts splitting of energy levels into two sub-

due to the Pauli exclusion principlethan two electrons in the same energy#1 as seen by atom #2 are no longer pbecause there are already two electron

0

levels. Thus, all the electrons (say at the 1s level) for the pcoexist by shifting their energies a littthat their wave functions do not coincexclusion principle is satisfied. Whsplitting of many levels into fine gradthe energy barrier between the atomscarry across to meet those of neighbands become free to migrate from atthey were in a potential well. Becauthese electrons (compared to if they wthese electrons (compared to if they wtheir kinetic energy is lower (see Sequates to stability that is, the free eenergy of the system and so act to hola metallic bond.

ave two Li atoms close together. Weentials as:

13

ple, the 1s electron orbiting the nucleusextent by the nucleus of atom #2. But,

f h h

2p

2s

1s

3s

or Nucleus for atom #2

atomic energy levels

e, for each atom, we cannot have morelevel (i.e. the two 1s electrons for atomermitted to have energies at the 1s level

ns from its own atom at that level).

2p2s

1s

3s

These electrons are still attached to their

t l b t

pair of atoms cantle up and down so cide and so the hen there are many atoms present, theations creates a band of energies and ifs is low enough for the upper bands tohbouring atoms, the electrons in thesetom to atom. The electrons behave as ifuse of the large range of movement of

were still attached to their parent nuclei),

parent nucleus but the presence of the neighbour electrons causes their energy levels to shift.

were still attached to their parent nuclei),ection 1.5) and a lowering of energy

electrons in the valence band reduce theld the two atoms together. This is called

1.13 Electronic Structure of S

In a solid, the interacting potentialsspaced atoms causes atomic energy le

14

sub-levels. The energy difference betwmolecular level is considered to be vir

2p

2senergy

In the diagram here, the broadening of thsuch that electrons in this band are no loto a particular atom. These electrons are shared between all the atomic nuclei pres

1senergy bands

If the highest energy band that convalence band) in a solid is not comband can easily move around frommovement can be readily obtained bySuch solids are thermal and electrical

If the valence band in a solid is fullpositioned some distance away inelectrons within the topmost band can

h hi h b d S hto the next highest band. Such minsulators.

If the next highest available band isband, then even at room temperature,given to some electrons to be promobecomes conducting and is a semiconducting electrons is called the cvalence band is the conduction bandband (at 0K) is separated from theenergy gap.

Atoms in a solid generally form molg yin a regular pattern (crystalline sothemselves in an orderly way (amoatoms are pushed and pulled arouinteraction electron potentials reach a

Solids

of many millions of relatively closelyevels to split into a very large number of


ween each sub-level is so fine that eachrtually a continuous band of energies.

he 2p level is nger local effectively

sent.

The spaces or energy gaps between bands

Electrons in these bands are constrained by the potential well and are still bound to individual nuclei.

gaps be ee ba dsare forbidden states where no electrons can exist.

ntains electrons in the ground state (thempletely full, then electrons within that

state to state within the band. Suchy applying an electric field to the solid.conductors.

l, and the next highest available band isterms of its energy levels, then the

nnot easily move from place to place ori l h l d l i lmaterials are thermal and electrical

positioned fairly closely to the valencethere may be sufficient thermal energy

oted to this higher level. The materialiconductor. The band containing theconduction band. In a conductor, thed. In a semiconductor, the conduction

valence band (defined at 0K) by an

lecules which either arrange themselvesglids) or the molecules do not repeat

orphous solids). When a solid forms,und and settle into place when theminimum level.

2 Chemical Compo

2. Elements and Molecules

2. Chemical Compo

Summary Mass numnumber ofneutrons i

Atomic number

XAZ

Avogadros number: 6.022 1

Atoms that lose or gain electro

One atomic mass unit (amu) is 12 atom. 1 amu = 1.6602 10

Atomic numberprotons in the n

Cations (+) (electrons lost Anions () (electrons gain

Stoichiometry is the process ofatoms, molecules and compoun

A molecule is the smallest collelectrically neutral and can exiunit.

The molecular weight of a subatomic weights of its constituen

Anions ( ) (electrons gain

nents

15

nents

mber total f protons and in the nucleus

r number of

Chemical symbol

X

023 = 1 mole.

ns are called ions:

1/12th the mass of a carbon -27 kg.

r number of nucleus

t)ned)

f accounting for the masses of nds in chemical reactions.

ection of atoms that is st as a separate identifiable

bstance is the sum of the nts.

ned)

2.1 Matter

16

Elements: cannot be reduced to simchemical reactions. (Examples are

Mixtures: can be separated into cMixtures are a product of mechanicaalloys, suspensions, dispersions anduniform composition and appearancemolecular length scale while hete

elements are substances which havecertain number of protons on the nucknown, with about 95% of these occreated synthetically by nuclear proce

Elements can be broadly divided int

Compounds: combinations of two orwhere the composition of the compouhave a chemical structure, the atomsCompounds keep their chemical idenonly be reduced to their constitut(Examples are water, salt, sulphuric ac

gdistinct regions. (Examples are fruit ca


mpler units by physical processes ore oxygen, iron, carbon.) Generally,

constituent parts by physical means.al or physical processes, such as metalcolloids. Homogenous mixtures have

e (examples are air, salty water) over aerogeneous mixtures have physically

e only one kind of atom (as having aleus). There are just over 100 elementscurring naturally, the remainder being

esses.o non-metals, semi-metals and metals.

r more elements in definite proportionsund is uniform throughout. Compoundsbeing held in place by chemical bonds.ntity when altered physically, and cantive elements by chemical reactions.cid.)

g p y yake, concrete, fruit juice with pulp.)

2.2 Atomic Weight

Except in the case of nuclear transforbeing much like the Bohr atom. A

2. Chemical Components

number, a chemical symbol, and the

Ttoa(iamo

Mass number total number of protons and neutrons in the nucleus

Atomic number - number of protons in the nucleus

Chemical symbol

XAZ

Since the mass of an electron me isatom is contributed by the protons anneutrons have very nearly the same matom is given by Z A. However, itthe same element with atomic numbbecause of having a different numberof atom A of an element is called a nu

By international agreement, onemass unit (amu) is 1/12th the mass ocarbon 12 atom. The atomic werelative atomic mass) of an elemeratio of the average mass per atomelement to 1/12 of the mass of an atomNote: The atomic weight of an element is nfrom adding the masses of protons, neutronelectrons. When an atom is formed, some ois used as nuclear binding energy via the Erelationship E = mc2.

In chemistry, much of the arithmeatomic weights for the atoms that takechemical reactions is due to the numvalence electrons in the atoms ratherword that means power and inarrangement of valence electrons thaform chemical bonds with other atom

rmations, we can usually treat atoms asAtoms are characterised by an atomic

17

mass number.me = 9.1096 10-31 kgmp = 1.6726 10-27 kgmN = 1.6749 10-27 kg

The mass number is approximately equal o the atomic weight of the element. The atomic weight of an element is the mass n relative terms) of an average atom of

an element. The more precise term atomic mass is used to denote the absolute mass of a particular atom (in kg).

s very small nearly all of the mass of annd neutrons in the nucleus. Protons andmass, so the number of neutrons in ant is found in nature that many atoms ofber Z have different mass numbers Ar of neutrons in the nucleus. Each typeuclide, or isotope.

The carbon 12 isotope has 6 protons and 6 neutrons in the nucleus. The carbon 13atom has 6 protons and 7

e atomicof a singleeight (or

p ( g)

neutrons in the nucleus. Both isotopes have 6 electrons. Experiments show that carbon in nature consists of 1.11% 13C and 98.89% 12C. The atomic weight of carbon is determined to be 12.011.

g (ent is them of them of 12C.not found ns and of the mass

Einstein

etic of chemical reactions involves thee part. However, the actual nature of the

mber and arrangement of the outer-shellr than the heavy nuclei. Valence is a

some sense, it is the number andat give an atom combining power to

ms.

2.3 Ions

Atoms that lose or gain electrons aremore electrons, it is called a cation

18

electrons, it is called an anion.Examples of ions:

Atom Z (atomic number)

No. electrongained or los

H 1 -1

Na 11 -1

Ca 20 -2

F 9 +1

S 16 +2

When an ion is formed, the atomWhen an ion is formed, the atomformula of the ion signifies the neelectrons, especially unpaired outer sremoval, absorption of ionising raanother type of atom which has a stron

Metallic elements tend to lose onewith other elements and so form catioelectrons to form anions. When nammore metallic element first. In binarcation first followed by the non-mecovalent compounds, the more metalli

Sodium chloridCarbon monoxCarbon dioxideNitrogen oxide

called ions. When an atom loses one orn. When an atom gains one or more


ns st

No. electrons remaining

Formula of ion

0 H

10 Na

18 Ca2

10 F

18 S2

is no longer electrically neutral Theis no longer electrically neutral. Theet electronic charge. Atoms may loseshell or valence electrons, by physicaldiation, the electron being taken bynger attraction for it, and so on.or more electrons when they combine

ons. Non-metallic elements tend to gainming compounds, it is usual to list thery ionic compounds, we list the metaletal anion with an ide suffix. Foric element is listed first.dexidee Where there is more than one

possible compound, prefixes are used to distinguish them.

2.4 Molecules

Most elements bond with others tosmallest collection of atoms, that is


separate identifiable unit. The concep1811, who, on the basis of observatproposed that atoms combine to formtemperature and pressure, equal volumcontain the same number of moleculesAvogadros law provides the justification for a method of determining relative molecular weight

H2OA wateatoms

A molecule is most convenientlyformula.

atoms

In some cases, such as in a crystalline appropriate because the atoms which ma regular repeating pattern. In this case

NaCl A sodnumbearrangThe molecular weight (or

relative molar mass, or relative molethe sum of the atomic weights of its co

If we had a certain mass of aelement (or a molecule), say 12 gramof 12C ho man atoms o ld this beof 12C, how many atoms would this beThis number is called Avogadronumber NA and has the valu6.0220943 1023. NA is found bexperimental methods.

12 g of 12C = 6.0220943 10atoms; therefore the mass of one 12atom is:

kg10991

kg1002.6

012.0mass

23

23

kg1099.1 It is convenient when working with

reactions to round down Avogadros n

o form molecules. A molecule is theelectrically neutral, and can exist as a

19

pt was first proposed by Avogadro intions made by Dalton and Gay-Lussacmolecules, and at the same

mes of all gases s.

ts

er molecule consists of two H s and one O atom

described in terms of a molecular

because there was now a way to obtain equal numbers of different molecules by using different gases all at the same temperature, pressure and volume.

s and one O atom.

solid, the concept of a molecule is not make up the substance are arranged in e, we speak of the empirical formula.ium chloride crystal consists of equal ers of sodium and chlorine atoms ged in a regular array or crystal lattice.

ecular mass) Mr of a substance is onstituents.

By international agreement, one atomic mass unit (amu) is 1/12th the mass of a carbon 12 atom

anmse? the mass of a carbon 12 atom.e?sueby But!, you might say, How can

you have 6.022... atoms if atoms are indivisible? Shouldnt this be a whole number? Look at the 1023. If you write Avogadros number out in full, it is a whole number.

0232C

quantities usually involved in chemicalnumber to 6.02 1023.

2.5 Mole

Chemical reactions usually involve laris convenient to have a unit of measur

20

or kilograms, etc) of bulk chemicalsdefinition, this unit of measure is cnumber of atoms or molecules.

12 grams of 12C contains 6.02 1023 at18 grams of H2O also contains 6.02 1

The word mole is a shorthand wThat is, it is easier to say Consider oConsider one Avogadros number o1023 of sodium atoms. To find out theatoms, we simply look up theatomic weight. To find out the mass in grams of a mole of some molecules, we simply use the molecular weight.

Moles, why bother? Its convenient.the anythings. When we are talkinthat one mole has a mass equal to thmolecule because thats the way Avthe link between the relative atomic m

One atomic mass unit (amu) is 1international agreement. 1 amu is

The atomic weight (or relative aof the average mass per atom ofatom of 12C.

The total sum of the atomic wecalled the molecular weight or r

The molecular/atomic weight expmolecules/atoms. This is Avogadr

Units of chemical accounting:

The molar mass M of an elemenexpressed in g/mol. The molar matomic weights of the constituent

rge numbers of atoms and molecules. Itre that relates weight (i.e. mass in grams


s with atomic or molecular weight. Bycalled a mole, and it is the Avogadro

toms and is called 1 mole of 12C

023 molecules and is called 1 mole of H2O

way of saying Avogadros number.one mole of sodium atoms rather thanof sodium atoms or Consider 6.02 e mass in grams of a mole of some

because the atomic weight (in amu) is expressed relative to the 12C atom, and thanks to Avogadro, we know that 12 g of 12C has NA atoms.

One mole of anything is 6.02 1023 ofng about atoms or molecules, we knowhe atomic weight of the element, or thevogadros number was determined. Itsmass and the actual mass in grams.

/12th the mass of a carbon 12 atom by1.66020943 10-27 kg.

atomic mass) of an element is the ratiothe element to 1/12 of the mass of an

eights for a molecule of substance iselative molar mass Mr.

pressed in grams contains one mole ofros number 6.0220943 1023.nt is the atomic weight of the element

mass of a compound is the sum of theelements expressed in g/mol.

2.6 Compounds

Compounds consist of two or more dcombined in a definite fixed ratio (as


any proportions of ingredients).compounds:

Molecular compounds are composbonds in fixed proportion and are elethe water molecule H2O is a moleformula shows the proportions of theName Molecular formu

Methane CH4Carbon dioxide CO2Ammonia NH3Ammonia NH3Acetylene C2H2Ethylene C2H4Ethane C2H6Water H2O

Benzene C6H6Ethanol C2H5OH

Naphthalene C10H8Aspirin C9H8O4

Ionic compounds are composed ofatoms, of opposite sign that are held tcompound is formed from the joiningto be a complex ion (or a polyatomicthe positive charge in the compound icharge is the anion. The ionic formulThe ionic formula is electrically neutrName Ions formed F

Magnesium chloride Mg2+ Cl- M

Silver sulphate Ag+ SO42- A

Ammonium sulphate NH4+ SO42- (

Chromium (III) hydroxide Cr3+ OH- C

Sodium chloride Na+ Cl- N

different elements which are chemicallydistinct from mixtures, which can have

21

There are generally two types of

sed of atoms held together by chemicalectrically neutral overall. For example,ecular compound and the molecularatomic species within it.

There is a distinction to be made between the molecular formula and the empirical formula of a compound The

ula

formula of a compound. The empirical formula gives the ratio of component atoms in the lowest possible numerical terms. The empirical formula shows the ratio of component atoms as found to exist in practice. In many cases, the formulae are the same, but this cannot be always assumed. For example, the empirical formula for benzene is CH but the molecular formula is C6H6.

f charged atoms, or a charged group oftogether by ionic bonds. When an ionicg of a charged group of atoms, it is saidc ion, or a molecular ion). The ion withis the cation. The ion with the negativela gives the ratio of anions and cations.ral overall.Formula

MgCl2Ag2SO4NH4)2SO4

Cr(OH)3NaCl

2.7 Cations

+1

22

Name Formula

Hydrogen H+

Lithium Li+

Sodium Na+

Potassium K+

Rubidium Rb+

Caesium Cs+

Silver Ag+

Copper Cu+ In Mercury (I) Hg22+

Ammonium NH4+

Name Formula

Beryllium Be2+

Magnesium Mg2+

Calcium Ca2+

Strontium Sr2+

B i B 2+

+2

chit ithemefolnosecis g(e.

Barium Ba2+

Lead Pb2+

Zinc Zn2+

Cadmium Cd2+

Nickel Ni2+

Manganese Mn2+

Tin (II) Sn2+

Iron (II) Fe2+

Mercury (II) Hg2+

Cobalt (II) Co2+Cobalt (II) Co2+

Chromium (II) Cr2+

Copper (II) Cu2+

+3 , +4


Name Formula

Aluminium Al3+

Bismuth Bi3+

Iron (III) Fe3+

Cobalt (III) Co3+

Chromium (III) Cr3+

Tin (IV) Sn4+

the naming of Prefixes

1 mono

2 di

3 tri

4 tetra

5 penta

6 hexa

gemical compounds, is customary to write e name of the most etallic element first, llowed by the more n-metallic one. The cond element or ion given the ide suffix g. sodium chloride).

2.8 Anions

1


Name Formula

Fluoride F-

Chloride Cl-

Bromide Br-

Iodide I-

Hydroxide OH-

Nitrite NO2-

Nitrate NO3-

Chlorate ClO3-

Foof nunoPerchlorate ClO4-

Cyanide CN-

Permanganate MnO4-

Thiocyanate NCS-

Bicarbonate HCO3-

Bisulphate HSO4-

Dihydrogen phosphate H2PO4-

Acetate CH3CO2-

2

no

S

id

it

a

In thedisothiro

Name Formula

Oxide O2-

Peroxide O22-

Sulphide S2-

Sulphate SO42-

Sulphite SO32-

Thiosulphate S2O32-

Carbonate CO32-

Oxalate C2O42-

irochl

Chromate CrO42-

Dichromate Cr2O72-

Monohydrogen phosphate

HPO42-

3 , 4

23

Name Formula

Nitride N3-

Phosphate PO43-

Hexacyanoferrate (III) Fe(CN)63-

Hexacyanoferrate (II) Fe(CN)64-

or polyatomic anions, the assignment a suffix depends on the oxidation

umber (see Section 4.8) of the central n-metal atomn-metal atom.

Suffixes

de

e ous

te ic

some cases, the oxidation number of e metal is stated in roman numerals to stinguish compounds which would herwise have the same name (e.g. on (II) chloride and iron (III)on (II) chloride and iron (III) loride).

2.9 Chemical Equation

When atoms or molecules (the reamolecules to form compounds (the pr

24

take place. Chemical reactions invchemical bonds.

ProduReactants

Often one can tell if a chemical reaproducts that are considerably differExamples are the formation of a prcolour, production of heat and so ospontaneously other times they have

ReactProducts

Chemical equations are writtennumber of atoms or molecules is writ

22 CO2OCO2 This equation says that two moleculmolecule of oxygen to product two mbalanced equation, because the numbnumber of atoms in each product. Thleft-hand side and two carbon atoms oatoms on the left side and two on th

spontaneously, other times they haveenergy or the lowering of energy barri

Usually, reactions proceed until therchemical equilibrium occurs whenbecomes equal to the rate of the revers

ProduReactants

atoms on the left side and two on thmoles of carbon monoxide combinemoles of carbon dioxide. Since the mis the atomic or molecular weight expbalanced in terms of both number ofequation. The net electric charge muequation.

ProduReactants By rate, we mean the rate of formation

ctants) combine with other atoms orroducts), a chemical reaction is said to


volve the breaking and formation of

ucts This is the forward reactionThis is the reverse reaction

action has occurred by observation ofrent in appearance than the reactants.

recipitate, evolution of gas, change ofon. Often, chemical reactions proceedto be forced to proceed by the input of

tants

using chemical formulae where thetten before the formula:

2

es of carbon monoxide react with onemolecules of carbon dioxide. This is aber of atoms in each reactant equals theat is, we have two carbon atoms on theon the right-hand side, and two oxygen

he right side We can also say that two

to be forced to proceed by the input ofiers.

re is chemical equilibrium. Dynamicthe rate of the forward reaction

se reaction:

ucts Dynamic ilib i

he right side. We can also say that twowith one mole of oxygen to form two

mass of a mole of an atom or a moleculepressed in grams, a balanced equation isf atoms and mass on both sides of thest also be the same on both sides of the

ucts equilibriumn of products in units of mol/s.

2.10 Stoichiometry

Stoichiometry is the process of amolecules and compounds in chemic


such accounting is the arrangemenequations. Stoichiometric calculations

1. Write a balanced equation for t2. Convert known masses of reac3. Calculate the number of moles

whose masses are unknown. 4. Calculate the masses of the unk

Balancing an equation involves botcharge. This is particularly importanspectator ions may be present and dexamples of unbalanced equationcorresponding net ionic equations are

)(aqCl2)(aqPb

KCl(aq))(aqNOPb2

23

2H)(sOHNi

)(aqHNO)(sOHNi

2

32

Practical difficulties can arise whereactants and products. For examplexcess (more present than combine wone reaction may take place at the samseries of intermediate steps. Productsthe equilibrium state of the reaction. Tmixing of the reactants and so there is

23

33

2H2HCO

N)(aqHNO)(aqNaCO

accounting for the masses of atoms,cal reactions. The basic component of

25

nt of chemical formulae in chemicals are performed in a logical sequence:

the reaction.ctants and/or products into moles.s of the reactants and/or products

known reactants and/or products.

th a balance of atoms and also electricnt in the case of ionic equations wheredo not take part in the reaction. Somens (with spectator ions) and thee:

)(sPbCl

NOK)(sPbCl

2

32

O2HNi

OH)(aqNiNO

22

23

en attempting to determine masses ofe, some reactants may be present in

with one or other reactants). More thanme time, or a reaction may proceed in as (such as gases) may escape and alterThe nature of the reactants may prevents incomplete yield of product.

22

223

COOH

COOH)(aqNaNO

2.11 Example

Acetylene gas is used together with oxprepared from a reaction of calcium ca

26

22 COHCaC Balance this equation and determine tfrom the reaction of 130 g CaC2 and

22 COH2CaC

240

13CaC2

.521

10OH.2

2

But according to the balanced equatioH2O are required. Therefore, 2(2.03) the mass of C2H2 formed would be:

.2HC 22

522

03.2

xygen in an oxy welding set. It is arbide and water:


222 HCOHCa the mass of acetylene that is produced 100 g water.

222 HCOHCa

Unbalanced eqn.

Balanced eqn.

moles03

1221.0130g 30

moles 56.16

100g 00moles03.

on, for each mole of CaC2, two moles of = 4.06 moles of H2O required, and so

moles 03.

g 8.2

1212 m

3 The Periodic Tabl

3. Periodic Table

3. The Periodic Tabl

SummaryAtomic radius decreases

adiu

s in

crea

ses

Ato

mic

ra

Ionisation energy increases

on e

nerg

y de

crea

ses

Ioni

satio

Electronegativity increases

Elec

trone

gativ

ity d

ecre

ases

E

e

27

e

2

22

ee

o

Zqmhnr

Bohr radius

He 1s2 Ne He2s22p6Ar Ne3s23p6Kr Ar3d104s24p6Xe Kr4d105s25p6Rn Xe5d106s26p6

Noble gases

Rn Xe5d 6s 6p

3.1 Electron Configuration

In a multi-electron atom, the elecCoulomb potential of the form 1/r

28

screening effect of inner-shell electrmulti-electron atom, the potential in wits energy) depends on not only n (athe second quantum number l. The arground state can be complicated, but

1. Electrons occupy the lower enehigher levels.

2. Electrons at a particular energy(Pauli exclusion principle).

3. Electrons tend to occupy an enepairing up (Hunds rule)pairing up (Hund s rule).

Hydrogen 1s1

Helium 1s2

Lithium 1s22s1

Beryllium 1s22s2

Boron 1s22s22p1

Carbon 1s22s22p2

Nitrogen 1s22s22p3

Oxygen 1s22s22p4

Fluorine 1s22s22p5

Eso

Examples ofconfiguratioin order of in

Wp

Neon 1s22s22p6

Sodium 1s22s22p63s1

Magnesium 1s22s22p63s2

Aluminium 1s22s22p63s23p1

Silicon 1s22s22p63s23p2

Phosphorus 1s22s22p63s23p3

Sulphur 1s22s22p63s23p4

Chlorine 1s22s22p63s23p5

Argon 1s22s22p63s23p6

Wenaplous

Potassium 1s22s22p63s23p64s1

Calcium 1s22s22p63s23p64s2

Scandium 1s22s22p63s23p63d14s2

ctrons do not find themselves in aas in the hydrogen atom due to the


rons on the outer-shell electrons. In awhich an electron finds itself (and hences in the single-electron atom), but alsorrangement of electrons in atoms in thefollows certain rules:

ergy levels first, before occupying

y level have opposite spin

ergy level as single electrons before

1s 2s 2p

Electrons spread out to occupy unfilled sub-shells before pairing up. For example, in oxygen, the arrangement of electrons is:

f the ground-state electron n of the first few elements arranged ncreasing atomic number Z

When we write the electron configuration ofWhen we write the electron configuration of elements, we keep the principal quantum numbers together even if they are not arranged in energy order. For example, in potassium, the 4s sub-shell actually has a ower energy than the 3d sub-shell, but we usually write the 3d along with the 3s and 3psub-shells.

Note 4s has lower energy than 3d in K.

Note 4s has lower energy than 3d in Ca.

Keep sub-shells together even if no longer in increasing energy order.

3.2 Periodic Law

In the nineteenth century, it was notorder of increasing atomic weight, ce

3. Periodic Table

to occur at periodic intervals. This isthat periodicity occurs when the elemnumber Z. Consider the first few elem

1 Hydrogen 1s1

2 Helium 1s2

3 Lithium 1s22s1

4 Beryllium 1s22s2

5 Boron 1s22s22

6 Carbon 1s22s22

i 1 22 22

Another example of periodicity. Lithium is a highly reactive

t l d 7 Nitrogen 1s22s22

8 Oxygen 1s22s22

9 Fluorine 1s22s22

10 Neon 1s22s22

11 Sodium 1s22s22

12 Magnesium 1s22s22

13 Aluminium 1s22s22

14 Silicon 1s22s22

15 Phosphorus 1s22s22

16 Sulphur 1s22s22

metal and so are sodium and potassium.

16 Sulphur 1s22s22

17 Chlorine 1s22s22

18 Argon 1s22s22

19 Potassium 1s22s22

20 Calcium 1s22s22

21 Scandium 1s22s22

Initially it seems that the chemical reaof eight in atomic number Z. Howevecompletely we see that the periarrangement of outer-shell valence elarrangement of outer shell valence elin hindsight because it is the valenformation of chemical bonds. For ex(36), xenon (54) and radon (86) all haare inert. Na and K have one valence e

ticed that if elements were arranged inertain properties of the elements tended

29

a striking observation. It is now knownments are arranged in increasing atomicments and the property of reactivity:

2p1

2p2

2 3

An example of periodicity. Helium is not very reactive (just about inert) and so are neon, and also argon.

2p3

2p4

2p5

2p6

2p63s1

2p63s2

2p63s23p1

2p63s23p2

2p63s23p3

2p63s23p4

Chemical reactivity is not the only property that tends to occur periodically. Other properties such as2p63s23p4

2p63s23p5

2p63s23p6

2p63s23p64s1

2p63s23p64s2

2p63s23p63d14s2

activity of elements repeats in intervalsr, when the elements are arranged moreiodicity actually depends upon thelectrons a not surprising observation

properties such as whether an element is a metal or a non-metal, whether they form certain compounds with other elements and so on can also occur periodically.

lectrons a not surprising observationnce electrons that are involved in thexample, neon (10), argon (18), krypton

ave eight electrons in the outer shell andelectron and easily form positive ions.

3.3 Periodic Table

Periodicity in chemical (and sometimebe appreciated when elements are ar

30

the table (groups) show elements wrow is called a period. The table alswhich electron shells are occupied.

I II

1 Z1H

2 3Li

4Be

Groups Shading indicates metalssemiconductors, non-metand noble gases.

Li Be

3 11Na

12Mg

4 19K

20Ca

21Sc

22Ti

23V

24Cr

25Mn

26Fe

27Co

5 37Rb

38Sr

39Y

40Zr

41Nb

42Mo

43Tc

44Ru

45Rh

6 55Cs

56Ba

57La*

72Hf

73Ta

74W

75Re

76Os

77Ir

Transition elemen2s

3s

4s

5s

6s

3d

4d

5d

7 87Fr

88Ra

89Ac**

104Ku

105Ha

Alk

ali m

etal

s

Lanthanide (rare

* 58Ce

59Pr

60Nd

61Pm

62Sm

63Eu

The Lanthanidpresented septo make the m

6s

7s

5d

Per

iod

Alk

alin

e ea

rth m

etal

s

Actinide series

** 90Th

91Pa

92U

93Np

94Pu

95Am

96Cm

5f

es physical) properties can more readilyrranged in a table. Vertical columns in


with similar properties. Each horizontalo indicates (approximately) the way in

III IV V VI VII 0

2He

5B

6C

7N

8O

9F

10Ne

1s

s, tals

B C N O F Ne

13Al

14Si

15P

16S

17Cl

18Ar

7o

28Ni

29Cu

30Zn

31Ga

32Ge

33As

34Se

35Br

36Kr

5h

46Pd

47Ag

48Cd

49In

50Sn

51Sb

52Te

53I

54Xe

7r

78Pt

79Au

80Hg

81Tl

82Pb

83Bi

84Po

85At

86Rn

nts2p

3p

4p

5p

6p

Hal

ogen

s

Nob

le g

ases

earth) series

3u

64Gd

65Tb

66Dy

67Ho

68Er

69Tm

70Yb

71Lu

de and Actinide series are usually parately from the main table so as main table a convenient size.

6p

4f

97Bk

98Cf

99Es

100Fm

101Md

102No

103Lr

3.4 Groups

Hydrogen is placed in a separate posHydrogen has one electron, which in t

3. Periodic Table

y g ,state is in the 1s orbital. This unpairemakes hydrogen a reactive element tto form compounds with many otherincluding itself, to form a gas H2.

Group I elements are the alkali mevalence electron, they are highlyElements in this group usually form M

Group II elements are the alkalineusually form M2+ ions and are genesilicates, carbonates, sulphates and ph

The transition elements generally oand III. All these elements are metaonly a few electrons. However, unlihard, brittle and have a high melting p

Group III elements are also considsemi-metallic properties). With the exand are relatively soft.

Group IV elements range from non-to metals down the group with increaeffect of electrons). All have four elresponsible for the formation of hydrresponsible for the formation of hydrlife on Earth. Silicon, unlike carbon, tthe basis for most of the minerals of th

Group V elements range from nowith increasing atomic size and five v

Group VI elements show little meionisation potentials as we go acrosssize going down the group confers som

Group VII elements are referred totendency to complete their electron sare largely non-metallic and mostly reare largely non-metallic, and mostly re

Group 0 elements are the noble gasalthough they are able to form compcertain conditions. Argon was discove

sition in the periodic table. the ground Protium

31

H11ged electronhat is able

r elements,Deuterium (heavy hydrogen)

Tritium (radioactive)

99.9% abundance

etals. Having one weakly bound outerreactive, and metallic in character.

M+ ions which are water soluble.e earth metals. Elements in this grouperally insoluble, occurring naturally ashosphates.

1

H21

H31

occupy the positions between Groups IIals since their outermost shells containike alkali metals, transition metals arepoint (with the exception of mercury).dered metals (although boron has onlyxception of boron, they form M3+ ions

-metal, semi-metals (or semiconductors)asing size of atom (and hence screeningectrons in their outer shell. Carbon isrocarbons and derivatives the basis ofrocarbons and derivatives, the basis oftends to form bonds with oxygen and ishe Earth.

on-metallic to metallic down the groupvalence electrons.etallic character due to the increasings the periodic table. Increasing atomicme metallic properties to Se, Te and Po.

o as the halogens. They all have a highshells by forming salts. These elementseactiveeactive.ses, so called because of their inertness,pounds with oxygen and fluorine underered in 1894.

3.5 Energy Levels

The ordering of the outer electron shinterpretation of the periodic table.

32

the energy shells proceeds in the fonegative) to highest (less negative pot

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d

From 1s to 3p, it is easy to understaincreasing radial distance from the nthe energy distribution just as in thnumber determines the distance thequantum number determines the shapshell. When l = 0, we have an s sub-sWhen l = 1, we have a p sub-shell whd s b shells hich are fo r lobed geod sub-shells which are four-lobed geosignificance of this is that in the fourththe 4s sub-shell has a lower energy thathese elements, electrons prefer to be ieven though the radius of the 4s shell

For the first row of the transition3d levels become occupied only afteare filled (except for Cr). In theseelectrons in the 4s shell, being furnucleus (but having lower energyelectrons) shield those in the inner 3

The ordering of shells given aboordering of all the energy levels wonly the energy ordering of the ouexample, in K, the 4s shell is at a low

the chemical properties for these elevery similar since in each case,electrons are the 4s outer electronsnumber of inner 3d electrons that isthe outer valence electrons (withprincipal quantum number) that interatoms to form chemical bonds.

example, in K, the 4s shell is at a lowthe 3d shell (3s23p64s1). By the timethe 3d shell is at a lower energy tha(3d104s2). There is no one sequence othat applies to all elements.

hells in terms of energy is the basis forCalculations show that the ordering of


ollowing sequence from lowest (moretential):d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d

and because the electron shells are at annucleus (n quantum number dominatese one-electron atom). The n quantumelectrons are from the nucleus. The l

pe of the electronshell which is spherical. hich is lobed. l = 2 gives ometries The

l = 0 sl = 1 pl = 2 dl = 3 fometries. The

h period, in K and Ca, an the 3d sub shell. In in 4s rather than 3d in the ground state is larger than the 3d shell. elements, the

er the 4s levelselements, the

rther from thethan the 3d

3d shell and so

l = 3 fl = 4 gl = 5 h

This is a consequence of the increasing dominance of the lquantum number in the dsub-shell in determining the energy for a sub-

ove is not thewithin an atom,uter shells. Forwer energy than

ements are allthe valence

while it is thechanging. It ish the highestract with other

shell. The l quantum number is connected with the angular momentum of the electrons and so, much like the case where the angular momentum for a rotating wheel is greater if the mass is concentrated at the outer edge compared to the case when the mass is evenly distributed, the concentration of mass of electrons in d-shapedwer energy than

e we get to Zn,an the 4s shell

of energy shells

electrons in d shaped shells results in a greater kinetic energy component to the total energy compared to spherical s shells.

3.6 Noble Gases

The noble gas elements are extremcompounds with any other element

3. Periodic Table

elements is the interesting propertyshells. Consider the second period elenumber of electrons in the energyn = 2 is increasing. At neon, we havgroup of eight (octet) valence electrois very energetically stable and is calle

An atom with one more valence e

He 1s2 Ne 1s22s22p6Ar 1s22s22p63s23p6Kr 1s22s22p63s23p63d104Xe 1s22s22p63s23p63d104

An atom with one more valence ehighly reactive metal. Sodium is hvalence electron is easily removed soconfiguration in its outer shell. An atothe highly reactive gas fluorine. Fluoanother atom to form an anion to achi

Rn 1s22s22p63s23p63d104

The noble gas configuration is wheshell and the next available higher enHe where the 1s shell is filled and ththe other noble gases, we have compleenergy level is the s sub-shell for thlarge energy gap between a p sub-shegives the noble gas elements a higoccupied energy shells are filled,

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