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Contents

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Contents

ContextsThe beginnings of chemistry . . . . . . . . . . . . . 2Where do elements come from? . . . . . . . . . . 10Elements of life . . . . . . . . . . . . . . . . . . . . . . . . 19Atmospheres . . . . . . . . . . . . . . . . . . . . . . . . . . 28Crystals . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 35Metals by design . . . . . . . . . . . . . . . . . . . . . . . 40Flavours and odours . . . . . . . . . . . . . . . . . . . 47Soil chemistry . . . . . . . . . . . . . . . . . . . . . . . . . 54The air we breathe . . . . . . . . . . . . . . . . . . . . . 62Consumer chemistry . . . . . . . . . . . . . . . . . . . 70Water quality . . . . . . . . . . . . . . . . . . . . . . . . . . 76Corrosion at a cost . . . . . . . . . . . . . . . . . . . . . 86Acids in the environment . . . . . . . . . . . . . . . 92Polymers . . . . . . . . . . . . . . . . . . . . . . . . . . . . 100Chemicals from the sea . . . . . . . . . . . . . . . . 105

Electrochemistry . . . . . . . . . . . . . . . . . . . . . 112Food analysis . . . . . . . . . . . . . . . . . . . . . . . . 120Energy supply . . . . . . . . . . . . . . . . . . . . . . . . 126Energy from food . . . . . . . . . . . . . . . . . . . . . 133Fuels . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 143Sports chemistry . . . . . . . . . . . . . . . . . . . . . . 149Wine analysis . . . . . . . . . . . . . . . . . . . . . . . . 156Quality control . . . . . . . . . . . . . . . . . . . . . . . 165Pharmaceuticals . . . . . . . . . . . . . . . . . . . . . . 172Forensic chemistry . . . . . . . . . . . . . . . . . . . . 179The chemical industry . . . . . . . . . . . . . . . . . 186Food for life . . . . . . . . . . . . . . . . . . . . . . . . . . 194Food production . . . . . . . . . . . . . . . . . . . . . . 200Biotechnology . . . . . . . . . . . . . . . . . . . . . . . . 207Marine chemistry . . . . . . . . . . . . . . . . . . . . 213

ChaptersChapter 1 Atomic structure . . . . . . . . . . . . 2211.1 The structure of atoms1.2 Elements and isotopes1.3 Atomic mass1.4 Electronic structure1.5 Electronic structure in more detail1.6 Organising information—

the periodic table

Chapter 2 Elements and compounds . . . . 2382.1 Physical and chemical properties2.2 Ionic bonding2.3 Covalent bonding2.4 Metallic bonding2.5 Organic compounds2.6 Patterns in properties

Chapter 3 Stoichiometry . . . . . . . . . . . . . . 2793.1 The mole3.2 Chemical reactions3.3 Stoichiometry of chemical reactions3.4 Limiting and excess reagents3.5 Percent yield

Chapter 4 Chemical solutions . . . . . . . . . . 2974.1 Dissolving4.2 Concentration

Chapter 5 Gases . . . . . . . . . . . . . . . . . . . . . . 3165.1 Physical properties of gases5.2 Kinetic molecular theory5.3 Measuring pressure and volume5.4 The gas laws5.5 The General Gas Equation5.6 Molar volume5.7 Stoichiometric calculations involving

gases

Chapter 6 Precipitation reactions . . . . . . . 3396.1 Types of chemical reactions6.2 Precipitation reactions6.3 Ionic equations

Chapter 7 Redox reactions . . . . . . . . . . . . . 3497.1 Oxidation and reduction7.2 Redox reactions7.3 Oxidation numbers and half equations7.4 Ease of oxidation

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Chapter 8 Acid–base reactions . . . . . . . . . 3708.1 Introducing acids and bases8.2 Reactions involving acids8.3 Defining acids and bases8.4 Concentration and strength8.5 The pH scale

Chapter 9 Polymerisation . . . . . . . . . . . . . 3899.1 The structure and properties of polymers9.2 Addition polymerisation9.3 Condensation polymerisation9.4 Resources

Chapter 10 Energy transformations . . . . . 40310.1 Enthalpy10.2 Specific heat capacity10.3 Spontaneous reactions

Chapter 11 Qualitative and quantitativetesting . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 42211.1 Testing a sample11.2 The flame test11.3 Chromatography11.4 Qualitative analysis using physical and

chemical properties11.5 Quantitative analysis

Chapter 12 Volumetric and gravimetricanalysis . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 43812.1 Gravimetric analysis12.2 Empirical and molecular formulae12.3 Volumetric analysis

Chapter 13 Instrumental and specialisedtechniques . . . . . . . . . . . . . . . . . . . . . . . . . . . 46713.1 Atomic emission spectroscopy13.2 Atomic absorption spectroscopy13.3 Colorimetry13.4 UV visible spectroscopy13.5 Infrared spectroscopy13.6 Nuclear magnetic resonance

spectroscopy13.7 Chromatography techniques

Chapter 14 Reactions rates . . . . . . . . . . . . 49214.1 Reaction speed14.2 Reaction mechanisms14.3 Order of reactions and rate laws

Chapter 15 Reactions of life . . . . . . . . . . . . 51115.1 Reactions of life15.2 The chemistry of digestion15.3 Matter cycles

Chapter 16 Reversible reactions . . . . . . . . 53416.1 Reversible reactions16.2 Le Chatelier’s Principle16.3 Equilibrium constant—a mathematical

relationship16.4 Solubility product16.5 Kw

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ContextsAtmospheresCrystalsMetals by designFlavours and odours

Chapter2 Elements and compounds

area of

Study

�Structure,bonding and

propertiesof materials

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c h e m i s t r y

Li

Na

BeSr

HeCa

RnCl2

FrBr

XeK

F2Kr

Br2Rb

NeI2

Ra

CsMg

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metals by

des ign

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�BY DESIGNMETALS⌧ 40

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�Alloys—mixtures of metalsAlloys are not pure substances. They are formed by melting a main metal,adding the required amounts of other components (usually other metals) andallowing the molten mixture to cool. Only small amounts need to be added toproduce an alloy with vastly differentproperties to the original metal. Somecommon alloys are listed in table md.1.

The addition of other componentsproduces alloys with specificproperties. By changing the latticestructure of the original metal, theregular pattern of positive metal ions isdisrupted. The alloying of metals canresult in a substance with increasedstrength and greater resistance tocorrosion. It can also produce asubstance that is brittle or has a lowermelting temperature than the puremetal. Generally alloys are harder andless malleable than their componentmetals as well as being poorerelectrical conductors.

Many alloys are the result of trialand error but an understanding of howalloying works enables the design ofmaterials with specific properties forparticular purposes.

Metals feature heavily in the homes, workplaces,industries and leisure activities of people in thedeveloped world. Copper makes great electricalwiring, aluminium foil keeps food fresh and mercuryenables us to tell the temperature. Rarely are thesemetals in their pure form. Most metals have to bemodified to make them suitable for their variousapplications. This context will outline methods used toachieve the characteristics required for the metalswe use.

Section 2.1 Physical and chemical properties describes the general propertiesof elements including metals.*

figure md.1 The discovery of bronzemoved humans into an age where metalobjects became far more prominent.

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Structure, bonding and properties

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Substitutional alloysSome alloys are made from elements with similarchemical properties and atoms of similar size. Theatoms of the metal added substitute for the mainmetal atoms in some positions of the lattice. Ions ofboth metals are attracted to the delocalised electronsso the mixture remains stable.

Bronze is an example of this kind of alloying. TheBronze Age is seen as a major step forward for thehuman race as it resulted in tools, implements andweapons that were harder and more resistant tocorrosion than the copper ones used previously. All it took was for someone to add some tin to thecopper furnace.

Interstitial alloysInterstitial alloys involve the addition of atoms thatfit into the spaces between the atoms of the metallattice. In carbon steel, the smaller carbon atomsoccupy the spaces in the iron lattice. Only a smallnumber of carbon atoms are randomly fitted into thespaces between the iron ions.

Section 2.4 Metallic bonding describes the structureof metals.*

Section 2.6 Patterns in properties describes the trendsin properties such as atomic size and metallic character.*

table md.1 Some features of common alloys.

Alloy Composition Properties Uses

mild steel iron, carbon (0.5%) strong, easily worked and welded girders, car bodies

tool steel iron, carbon (1%) strong, hard, brittle tools, knives, cast iron

tungsten steel iron, tungsten (15%) stays hard when hot drills, cutting toolschromium (4%)

stainless steel iron, chromium (18%) hard, corrosion-resistant cutlery, surgicalnickel (8%) equipment

brass copper (65%), zinc (35%) corrosion-resistant, strong taps, door knobs

bronze copper (90%), tin (10%) corrosion-resistant, hard ships’ propellers, statues

cupro-nickel copper (75%), nickel (25%) corrosion-resistant coins, turbine blades

solder lead, tin melts easily joining metals

dental amalgam mercury, tin, silver, copper hardens slowly after mixing dental fillings

gold alloys gold, silver, zinc, copper harder and less expensive jewellerythan pure gold

aluminium aluminium, magnesium, light, but harder than pure aircraft, motor vehiclealloys copper, zinc aluminium fittings, boats

figure md.2 A substitutional alloy. Atoms of theadded element have taken the place of atoms of themain element in the lattice.

figure md.3 Australian silver coins are actually madefrom an alloy of 75 percent copper and 25 percentnickel—a substitutional alloy.

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Varying the composition of alloys varies theirproperties. Carbon steel containing 4% carbon is verybrittle. It is sometimes known as cast iron and isused in cast iron cookware. The brittleness of castiron makes it unsuitable for structural purposes.0.1% carbon steel is ductile and able to be drawn intoa wire for making paper clips and staples. 1% carbonsteel is stronger without being brittle and is used forstructural purposes such as bridge building.

BY DESIGNMETALS

figure md.4 Carbon steel is an interstitial alloy. The smaller carbon atoms occupy some of the spacesbetween the iron atoms.

figure md.5 Cast iron cookware is a good conductor ofheat but its brittleness makes it unsuitable for building andengineering purposes.

figure md.6 1% carbon steel is extremely strong andflexible enough for bridge building. Steel for this purposemay also contain other elements like manganese,chromium and silicon to further refine its properties.

figure md.7 Stainless steel, a corrosion resistant alloy.

Rust free stainless steelThe British chemist Harry Brearley developedstainless steel in 1913 while investigating steel witha high chromium content for use in rifle barrels. Onestep in his analysis of steel involved dissolving it inacid but Brearley found the steel would not dissolve.He also found that it stayed shiny and didn’t developa dull oxide coating.

Stainless steel develops an extremely thincoating of chromium(III) oxide (Cr2O3) that adheresclosely to the surface of the metal, protecting it fromattack by air and water while allowing the shininessof the metal to show through. If scratched, a newprotective coating quickly forms.

Stainless steel is commonly used in cutlery,commercial kitchen surfaces and surgicalequipment. While it could be used more widely andfurther reduce the expense and inconvenience ofcorrosion, stainless steel is expensive to produceand more cost effective alternatives exist.

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....................................................................................................................................................ExperimentExperiment mdmd.1�Alloys: preparation of solder

A pellet of solder is prepared by heating a mixture of lead and tin. Pellets of solder,lead and tin are then heated simultaneously. The solder melts before the othermetals.

Method1 To demonstrate the production of an alloy.2 To show how the physical properties of an alloy differ from those of their

constituent metals.

ProcedureBefore the demonstration:1 Prepare three moulds by pushing a dowel stick into a tray of damp sand. 2 Place 10 g of lead in a crucible. Using a fume cupboard, heat the crucible

until the lead melts. Pour the molten lead into one of the moulds and allow it to cool.

3 Repeat this procedure using 10 g of tin. Use a marking pen to label the twopellets produced.

For the demonstration:1 Show the class the pellets of lead and tin that have been prepared and repeat

the procedure described above using a mixture of 5 g lead and 5 g tin. 2 When the metal has cooled, drop the three pellets of lead, tin and solder

onto a hard surface from a height of 2 m. Note the relative ease with whichthey are dented.

3 This test must be performed in a fume cupboard. Place the three pellets on a tin lid so that they are equidistant from the centre. Using a tripod stand,pipeclay triangle and Bunsen burner, heat the centre of the lid as shown infigure md.8. Note which pellet melts first. Stop heating once one of the pelletshas melted.

Discussion1 Why is solder used in plumbing rather than just tin or lead?2 How do the hardness and melting point of an alloy compare with those

characteristics of its constituent metals?3 Name some uses for alloys. Explain how these uses are related to their

properties.4 Explain how the physical properties of alloys reflect their bonding.

Altering metallic structureAny solid piece of metal is made up of individual crystals that form as moltenmaterial cools down. Each crystal has a regular arrangement of ions within asea of electrons. The thin individual crystals are randomly arranged. The size ofthe crystals and their arrangement is dependent on factors such as the rate atwhich the molten metal cooled. It is the arrangement of the crystals thatdetermines properties such as malleability and ductility. Different methods canbe used to change the crystal structure and make the metal suitable forspecific uses.

● ●● ●

figure md.8

43Structure, bonding and properties

materials

• Bunsen burner andtripod

• 3 × crucibles• pipeclay triangle• tin lid• shallow metal tray

filled with damp sand• tongs• dowel rod• marking pen• 10 g lead• 10 g tin• mixture of 5 g lead

and 5 g tin• cotton thread

safety

• Wear safety glassesand a laboratory coatfor this experiment.

• Lead vapour is toxic.The lead pellets mustbe prepared in afume cupboard. Donot continue to heatthe tin lid once thesolder melts.

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BY DESIGNMETALS

Work hardeningIf you bend a paper clip once it will normally notbreak. To break it you need to bend it back and forth anumber of times. The bending of the metal in a paperclip rearranges the crystal grains making it harder andbrittle. This process is known as work hardening.Hammering (or working) a metal can be done to makeit harder as well as change its shape. Sometimesworking and heat treatments are used to shape ametal, making it harder and suitable for various uses.

Smithing is the skill used to produce horse shoesand other metal objects from iron. Jewellers alsopractice a form of smithing to shape gold and silveralloys into jewellery.

Heat treatmentThe sizes of some of the crystals in a piece of metalcan be altered by heating and cooling the metal atdifferent rates. These processes allowed ancientcivilisations to produce tools and weapons with sharpblades and continues to be used in the production ofknives and sharpened tools.

figure md.10 In these representations of a metaleach circle represents an ion. On the left a solidcrystalline metal is represented and on the right amolten metal with a random arrangement of ions.

figure md.9 This piece of boiler steel (an alloy of iron,chromium and molybdenum) has individual crystals thatformed as molten material cooled down.

figure md.11 Shaping metal byheating and hammering.

table md.2 Three methods of heat treating metals.

Treatment Process Effect on metal properties Effect on metal structure

The metal is softer withimproved ductility.

The metal is harder andbrittle.

The metal is harder butless brittle.

Larger metal crystals form.

Tiny metal crystals form.

Crystals of intermediatesize form.

A metal is heated to a moderatetemperature and allowed to coolslowly.

A metal is heated to a moderatetemperature and cooled quickly(sometimes by plunging into water).

A quenched metal is heated (to alower temperature than is used forquenching) and allowed to cool.

annealing

quenching

tempering

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45Structure, bonding and properties

Since pure metals are rarely used, it is a combination of alloying and heattreatment that results in many of the objects in common use. The blanks usedfor making coins are work hardened by rolling. This also gives them a rim. Theyare then softened by heating to 850°C ready for stamping with their design.

figure md.13 A combination ofthe appropriate alloy andquenching produces the hardnessof this chisel blade, which iscapable of shaping wood.

....................................................................................................................................................Experiment mdExperiment md.2�Modifying the properties of metals

MethodTo modify the properties of iron by heating.

Procedure1 Take three needles or hair pins, keeping one as a control.2 Strongly heat the other two needles or hair pins, one at a time, in the flame of

the Bunsen burner.3 Drop one of the hot needles or hair pins directly into the beaker of cold water.

Allow the other needle or hair pin to cool very slowly by gradually moving itaway from the flame of the burner and allowing it to cool in air.

4 When the needles or hair pins are quite cool, try to bend them and to breakthem. Record your observations.

Discussion1 How does the rate of cooling affect the brittleness and hardness of the metal?

materials

• 3 × sewing needlesor hair pins

• 250 mL beaker ofcold water

• Bunsen burner• bench mat• tongs• pliers (optional)

safety

• Wear safety glassesand a laboratory coatfor this experiment.

• Remember not totouch the hot metal.

figure md.12 Body replacement parts like these knee parts are made from specialalloys that will not harm the human body or break down in such a harsh environment.

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METALSBY DESIGN

1 The melting temperatures and atomic radii of some elements are shown in table md.3.Using your knowledge of metallic structure and alloys, present a proposal for a possiblenew alloy. Design the alloy with a specific purpose in mind. Justify your choice ofelements and predict the properties the alloy would have.

Consider this

2 From your knowledge of metallic structure explain why alloying works. Why does theaddition of other elements change the properties of the main metal?

3 What do you think is best—refining metal treatments and developing new alloys orproducing new polymers and plastics from oil? Justify your response in terms of uses,cost and environmental factors.

4 Imagine iron didn’t exist. How would this have affected the development of society?Would another metal have come into dominant use?

1 Experiment md.2 makes use of simple heat treatment principles. Design your ownexperiment on the effects of smithing. You could compare different combinations ofwork hardening and heat treating.

Experimental investigations

table md.3

Element Melting Atomic Element Melting Atomictemperature radius temperature radius

(°C) (nm) (nm)

LiBeC

NaMgAlSiK

CaScTiVCr

1811278352798

649660141064

8391541166018871857

0.1520.1130.0770.1540.1600.1430.1170.2270.1970.1610.1450.1320.125

MnFeCoNiCuZnRbSrSnPbCsBa

1244153514951453108342039

76923232828729

0.1240.1240.1250.1250.1280.1330.2480.2150.1410.1750.2650.217

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Properties and the periodic tableIt was mentioned in Chapter 1 that the shape of the periodic table reflects theelectron configurations of the elements. It is the electron configuration thatdetermines how atoms bond together and so the periodic table also groupselements according to their similar physical and chemical properties.

The metallic elements can be found on the left side of the periodic table andnon-metallic elements on the right side. Hydrogen is an exception to this. It isplaced in group 1 because it has one valence electron and has many propertiessimilar to the metals in that group. At times hydrogen also behaves similarly tothe gases of group 17. Hydrogen’s properties are so exceptional that at times itis simply placed on its own.

Between the metals and non-metals of the periodic table are the metalloids.Metalloids can demonstrate properties of either metals or non-metalsdepending on the conditions under which they are observed. The semi-conductive properties of metalloids make them useful in electroniccomponents such as computer chips.

The vertical columns or groups of the periodic table tend to have elementswith similar physical and chemical properties. Table 2.1 is a summary of theproperties of some of these groups.

Sometimes it is useful to refer to the blocks of elements of the periodic table.Groups 1, 2 and 13–18 are often referred to as the main group elements. Theelements of groups 3–12 are referred to as the transition metals. Thelanthanides are the elements of atomic numbers 58–71 and the actinides haveatomic numbers 90–103. Elements are blocked like this because of theirelectron configurations but elements within each block can have a wide varietyof physical and chemical properties.

Elements and compounds

239

1

3

11

19

37Rb

55Cs

87Fr

58Ce

90Th

59Pr

91Pa

60Nd

92U

61Pm

93Np

62Sm

94Pu

63Eu

95Am

64Gd

96Cm

65Tb

97Bk

66Dy

98Cf

67Ho

99Es

68Er

100Fm

69Tm

101Md

70Yb

102No

71Lu

103Lr

4

12

20

38Sr

56Ba

88Ra

21Sc

H

Li Be

Na Mg

K Ca

39Y

57La

89Ac

22Ti

40Zr

72Hf

104Rf

23V

41Nb

73Ta

105Db

24Cr

42Mo

74W

106Sg

25Mn

43Tc

75Re

107Bh

26Fe

44Ru

76Os

108Hs

27Co

45Rh

77Ir

109Mt

28Ni

46Pd

78Pt

110Ds

29Cu

47Ag

79Au

111Rg

30Zn

48Cd

80Hg

112Uub

31Ga

49In

81Tl

114Uuq

113Uut

115Uup

116Uuh

32Ge

13Al

14Si

5B

6C

50Sn

82Pb

33As

15P

7N

51Sb

83Bi

34Se

16S

8O

52Te

84Po

35Br

17Cl

9F

53I

85At

36Kr

18Ar

10Ne

2He

54Xe

86Rn

group1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18

1

2

3

4

5

6

7

per

iod

TRANSITION ELEMENTS

metals

metalloids

non-metals

LANTHANIDES

ACTINIDES

figure 2.1 The shading in this periodic table shows that metals, non-metals and metalloids are grouped together inparticular regions of the table.

Metalloids candemonstrate propertiesof both metals and non-metals. They are alsoknown as semi-metals.

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BONDING AND PROPERTIES OF MATERIALS240

STRUCTURE,

Why do atoms bond?The noble gases—helium, neon, argon, krypton, xenon and radon—are knownfor being highly unreactive. They are extremely stable elements. Each one has afull outer shell of electrons, i.e. 2 electrons (s2) for helium and 8 electrons (s2p6)for each of the others.

The noble gases are unreactive because their electron configurations are sostable. Atoms of other elements with less stable electron configurations attemptto attain a stable configuration by bonding with other atoms. The atoms ofelements other than the noble gases attempt to do this in three ways:1 accepting electrons from other atoms or donating electrons to other atoms

to form charged particles which are attracted to each other

2 sharing electrons with other atoms in groups of two or more atoms

3 forming lattices of positive particles through which electrons can move.

These bonding types are known as ionic, covalent and metallic bonding.The next sections of this chapter will describe these methods in detail.

table 2.1 The general properties of elements of some groups of the periodic table.

Group Group Elements Physical properties Chemical propertiesnumber name

1 alkali Li, Na, K, Rb, Cs, Fr bright, shiny, metallic highly reactive with metals relatively soft water or oxygen

good conductors of electricity2 alkaline Be, Mg, Ca, Sr, Ba, metallic appearance highly reactive (less

earth Ra harder and higher melting reactive than alkali metals temperature than alkali metals metals)

good conductors of electricity17 halogens F, Cl, Br, I exist as gases at reasonably strong ability to remove

(These exist as low temperatures electrons from other diatomic molecules distinctive colours elementsF2, Cl2, Br2, I2)

18 noble He, Ne, Ar, Kr, Xe, exist as gases extremely unreactivegases Rn

figure 2.2 The noble gases have full outer shells of electrons—a very stable electronconfiguration.

ArNeHe

Nucleus

Bonding allows atoms togain the stable electronconfiguration of a noblegas, with full valenceshells—s2p6.

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241

....................2.2 Ionic bondingIn ionic bonding the substances are comprised of ions. Ions are atoms orgroups of atoms that have gained or lost electrons. If they gain electrons theyhave a negative charge and if they lose electrons they have a positive charge.Salt (sodium chloride) is one example of an ionic compound. It is composed ofpositive sodium ions and negative chloride ions.

Elements in the columns on the left side of the periodic table tend to loseelectrons to become positive ions, for example lithium, sodium and potassiumlose one electron to become the ions Li+, Na+ and K+.

Elements in the columns on the right side of the periodic table tend to gainelectrons to become negative ions, for example fluorine, chlorine and brominetend to gain one electron each to become the ions F–, Cl– and Br–.

These ions form as the atoms gain or lose electrons to have the sameelectron configuration as a noble gas. It is the outer electrons, called the valenceelectrons, which are lost or gained to form ions. When a sodium atom loses oneof its 11 electrons, it has the stable electron configuration of neon which has10 electrons (see table 2.2).

When chlorine atoms gain an electron they have 18 electrons and the sameelectron configuration as argon. When an atom has gained electrons we add-ide to its name. Cl is the symbol for a chlorine atom and Cl– is the symbol for achloride ion.

Elements in Group 2 of the periodic table tend to lose 2 electrons.Magnesium atoms lose 2 electrons to become Mg2+ ions with an electronconfiguration the same as neon. Aluminium atoms (in Group 13) tend to lose 3electrons to become Al3+. Phosphorous and sulfur attain the electron

...........................Questions1 Which of the following are physical changes and

which are chemical changes?boiling water, cooking a cake, iron rusting, waxmelting, a candle burning, nail polish removerevaporating, milk going sour, concrete setting,sugar dissolving in water

2 Group these elements into pairs with similarproperties.Li, S, He, Cu, Ar, K, Mg, Au, O, Sr

3 Which element in each group below would youexpect to have different properties?a krypton, helium, hydrogen, neonb magnesium, potassium, lithium, sodiumc copper, silver, gold, platinumd magnesium, beryllium, aluminium, calcium

In ionic bonding, atomslose or gain electrons toform charged ions. Ions ofopposite charge areattracted to each otherand the forces ofattraction create a bond.

table 2.2 An atom can gain or lose electrons in order to attain the more stable electron configuration of its closest noble gas.

Element Electron Ion Electron Noble gas electronconfiguration configuration configuration the of element of ion element attains

lithium 2 1 Li+ 2 heliumsodium 2 8 1 Na+ 2 8 neonpotassium 2 8 8 1 K+ 2 8 8 argonfluorine 2 7 F– 2 8 neonchlorine 2 8 7 Cl– 2 8 8 argonbromine 2 8 18 7 Br– 2 8 18 8 krypton

The valence electrons ofan atom are the electronsin the outermost electronshell.

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BONDING AND PROPERTIES OF MATERIALS242STRUCTURE,

configuration of argon by gaining 3 and 2 electrons respectively to become P3– and S2–.

Metallic elements on the left of the periodic table tend to lose electrons tobecome positive ions; non-metallic elements on the right side of the periodictable tend to gain electrons to become negative ions. Groups of atoms can alsobe charged. They are groups of non-metallic elements. Table 2.3 lists some ofthe common ions.

Positive and negative ions may be single atoms or groups of atoms that havelost or gained electrons. Note that the transition metals have Roman numeralsto indicate their valency as it can vary.

Ionic substancesSince metals tend to lose electrons to become positive ions and non-metalstend to gain electrons to become negatively charged, most ionic substances aremade up of metallic and non-metallic elements.

The burning of magnesium is a reaction between the metal magnesium andoxygen in the air.

magnesium + oxygen → magnesium oxide

table 2.3 The names and formulae of some common ions.

Positive ions (cations)+1 +2 +3

hydrogen H+ magnesium Mg2+ aluminium Al3+

lithium Li+ calcium Ca2+ chromium(III) Cr3+

sodium Na+ barium Ba2+ iron(III) Fe3+

potassium K+ zinc Zn2+

silver Ag+ copper(II) Cu2+

copper(I) Cu+ mercury(II) Hg2+

ammonium NH4+ iron(II) Fe2+

nickel(II) Ni2+

tin(II) Sn2+

lead(II) Pb2+

Negative ions (anions)–1 –2 –3

hydroxide OH– oxide O2– nitride N3–

hydrogen sulfide HS– sulfide S2– phosphate PO43–

hydrogen sulfite HSO3– sulfite SO3

2–

hydrogen sulfate HSO4– sulfate SO4

2–

hydrogen carbonate HCO3– carbonate CO3

2–

dihydrogen phosphate H2PO4– hydrogen phosphate HPO4

2–

nitrite NO2– dichromate Cr2O7

2–

nitrate NO3–

acetate CH3COO–

fluoride F–

chloride Cl–

bromide Br–

iodide I–

permanganate MnO4–

Positive ions are alsoknown as cations andnegative ions as anions.

A substance which ismade up of two or moredifferent types of atoms iscalled a compound.

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Magnesium oxide is an ionic compound. It is formed from magnesium ionsand oxide ions. Magnesium and oxygen ions are created during the reactionbecause an exchange of electrons occurs when magnesium ribbon burns. Amagnesium atom loses two valence electrons to become an ion with a charge of+2 represented as Mg2+. Magnesium ions are said to have a valency of +2.

An oxygen atom gains two valence electrons to become an ion with a chargeof –2 represented as O2–. Oxygen has a valency of –2.

The oppositely charged ions are attracted to each other and the +2 and –2charges balance. The compound magnesium oxide is formed. The cation(positive ion) is named and written first and the anion (negative ion) second. Asone oxygen anion is needed to balance each magnesium cation, MgO is thechemical formula of magnesium oxide. The formula for an ionic substanceshows the ratio of atoms not a set number of atoms in a particular sample.

To write the chemical formula of an ionic compound follow the steps below.

1 Write the ions with their charges.

2 Swap each of the charges and make them subscripts.

3 If the subscripts are the same then they can be cancelled out. For example,MgO is the balanced formula for the compound. Check: (+2) + (–2) = 0.

Sodium ions have a charge of +1 and chloride ions have a charge of –1, sosodium chloride with form in the ratio 1 : 1.

Sometimes the ionic charges do not balance as easily. Consider magnesiumchloride—a magnesium ion has a charge of +2 and a chloride ion has a chargeof –1. Magnesium chloride will have the formula MgCl2. (Note that the 1 on theMg can be omitted.)

Groups of atoms can also have a charge. The NO3– ion has one nitrogen

atom and three oxygen atoms and overall the group has a charge of –1. (Notethat when ions have a charge of –1 the 1 can be omitted.) Table 2.3 shows thatsome of the ions result from groups of atoms as well as single atoms.

When ions composed of groups of atoms are involved, the same steps areused. Sodium hydroxide, forms from the ions Na+ and OH–, is written as NaOH.If more than one of these ions is used to balance the compound brackets areused, for example magnesium hydroxide is formed from the ions Mg2+ and OH–.It is written as Mg(OH)2.

Properties of ionic compoundsIn ionic compounds, the forces between the positive and negative ions are quitestrong. They hold the ions in a tightly ordered arrangement.

Mg2+ O2–

Mg2 O2

figure 2.3Writing chemicalformulae for ioniccompounds.

Na1+ Cl1–

Na1 Cl1figure 2.4NaCl is a balancedionic formula. Check: (+1) + (–1) = 0

Mg2+ Cl1–

Mg1 Cl2

figure 2.5MgCl2 is the balancedformula. Check: (+2) + (2 × –1) = 0

figure 2.6 In sodium chloride the ions are held in structured lattices called crystals.

+ —

—

—

sodium ion chloride ion

+ +

+

+

+

—

+

+ —

—

— —

—

+—+

The valency of an ion isthe charge it takes on toattains a noble gaselectron configuration.

The formulae for ionicsubstances is also knownas the formula unit. Itshows the ratio of ions in asample of the substance.Individual particles withthat formula do not exist.

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Most ionic compounds melt at fairly high temperatures and so they aresolids at room temperature. They have high melting and boiling points. Theyare crystalline because of the ordered way the ions are arranged. If you apply aforce to an ionic crystal you are trying to push together ions of the same charge.As ions of the same charge repel each other, the crystal shatters. Most ioniccompounds dissolve in water.

For an object to conduct electricity, electrons or some other chargedparticles must be able to move. The electrons in ionic compounds are heldtightly and are not free to move, so solid ionic compounds do not conductelectricity. If you melt an ionic compound or dissolve it in water, the ions areable to move freely. If you connect either a molten ionic compound or asolution containing a dissolved ionic compound to a power supply the ions willmove through the solution carrying an electric current.

figure 2.7 The forces of repulsion between ions of the same charge cause thesodium chloride crystal to shatter.

force

crystal shatters

figure 2.8Sodium chloride insolution will conduct anelectric current becauseits ions are free to move.

...........................Questions4 Would you expect barium (atomic number 56) to

gain or lose electrons? How many?5 Iodine atoms tend to gain an electron. When this

happens, to which noble gas does it have asimilar electron arrangement?

6 Sulfur and oxygen are in the same columnbecause of their similar chemical properties.Why are they similar?

7 Write the symbols for the ions that form fromnitrogen and phosphorous atoms.

8 Explain how an atom becomes positively ornegatively charged when it gains or loseselectrons.

9 For each of the following describe the number ofprotons, number of electrons and overall chargeof the ions formed. (You will find the periodictable useful.)a a hydrogen atom loses one electronb a fluorine atom gains one electronc an oxygen atom gains two electronsd an aluminium atom loses three electrons

10 How many electrons do each of the followingions have? (Make use of the periodic table.)a K+

b Cl–

c Mg2+

d H+

e Al3+

f O2–

11 Write a formula for each of the following ioniccompounds.a copper(I) nitrateb silver oxidec chromium(III) fluorided lithium nitridee potassium carbonate

12 Name the ionic compounds with the followingformulae.a KFb LiOHc ZnCl2d AlNe (NH4)2CO3

Properties of elementsand ionic bonding

Worksheet1

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....................2.3 Covalent bondingCovalent bonding takes place between atoms of non-metallic elementsthrough the sharing of electrons. In contrast, ionic bonding takes place betweenatoms of metallic and non-metallic elements which gain or lose electrons.

The simplest example of covalent bonding is hydrogen gas. Elementalhydrogen exists as pairs of hydrogen atoms described by the formula H2.

Forming covalent bonds In covalent substances, the atoms do not gain or lose valence electrons;

instead they share electrons with the other atoms to which they are bonded.This sharing of electrons gives each atom the same number of electrons as anoble gas making it more stable. Since most noble gases have 8 electrons intheir outer shell (all except helium) we often talk about atoms bonding toachieve a stable outer shell of 8 electrons or a stable octet.

Figure 2.9 shows how hydrogen atoms share a pair of electrons in a hydrogenmolecule. Each hydrogen atom contributes one electron and the pair ofelectrons is shared equally between both hydrogen nuclei. The sharing of a pairof electrons creates a single covalent bond.

Chlorine atoms also combine with a single covalent bond to form chlorinemolecules, Cl2. Each chlorine atom has an electron configuration of 2,8,7 andthe valence shell only needs to gain one electron to obtain a stable octet. Thechlorine atoms achieve this by sharing a pair of electrons.

The pair of electrons that are shared between the two atoms is known as thebonding pair. Any pairs of electrons that are part of the outer shells but notshared between atoms are known as non-bonding pairs, or lone pairs.

Covalent bonds do not only form between atoms of the same element butalso between atoms of different non-metallic elements. Figure 2.11 illustrates

figure 2.9 The two atoms which make up a hydrogen molecule are held togetherby a covalent bond. This type of bond is the result of sharing valence electrons.Each hydrogen atom has two electrons (like the noble gas helium) making it astable molecule.

e–

1+

e–

1++

hydrogenmolecule

hydrogenatom

hydrogenatom

figure 2.10 Chlorine molecules areformed when a pair of electrons areshared between two chlorine atoms in asingle covalent bond. Only the outer shell(valence) electrons are shown here.

Cl Cl

Bonding electrons(one pair)

Non-bonding electrons(three pairs on each atom)

Bonding pairs of electronsare those that are sharedbetween atoms, forming abond. Non-bonding pairs ofelectrons are not shared.

When an atom has 8electrons in its outer shell,it is described as having astable octet. This is ahighly stable configurationcomprising of a full sorbital (2 electrons) and afull p (orbital 6 electrons).

A molecule is a group oftwo or more atoms heldtogether by covalent bonds.

Covalent bonding is thesharing of electronsbetween atoms of non-metallic elements.

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the covalent bonding between a chlorine and hydrogen atom to form ahydrogen chloride molecule. Each atom contributes an electron to a bondingpair to form a covalent bond.

As well as single covalent bonds, double and triple covalent bonds can beformed. A double covalent bond is the sharing of two pairs of electrons betweentwo atoms. Oxygen molecules are two oxygen atoms held together by a doublecovalent bond. Nitrogen molecules are two nitrogen atoms bonded by thesharing of three pairs of electrons in a triple covalent bond.

The same bonding can also result in larger molecules. Water consists of twohydrogen atoms and an oxygen atom covalently bonded. Two single covalentbonds are formed as each hydrogen atom shares a pair of electrons with theoxygen atom. The oxygen has an electron configuration of 2,6 and by sharing apair of electrons with each hydrogen atom the oxygen attains a stable outershell of eight electrons. The hydrogen atoms also have two electrons in theirouter shell, making the molecule very stable.

Carbon atoms have four valence electrons and can form four covalentbonds. Methane consists of a carbon atom covalently bonded to four hydrogenatoms. It has the formula CH4. The sharing of valence electrons between carbonand hydrogen is illustrated in figure 2.14.

There are many possible combinations of elements in covalent substances.Many others will be introduced in following sections and chapters.

Molecules and molecular substancesAll of the examples of covalent bonding used so far in this section are molecularsubstances. A molecule is a set number of atoms joined together. As discussedabove, the elements hydrogen, oxygen, nitrogen and chlorine exist as diatomicmolecules, that is molecules consisting of two atoms joined together. Somemolecules are made up of one type of atom (homonuclear); other examplessuch as hydrogen chloride, water and methane consist of atoms of differentelements covalently bonded together (heteronuclear). Some molecules are farmore complex and consist of many atoms. Examples include plastics, proteinsand the DNA (deoxyribonucleic acid) found in all living things.

figure 2.11 A stablearrangement of outer shellelectrons can also beattained by the covalentbonding of different elements.

ClH Non-bondingelectrons

Bonding electrons

figure 2.12 A representation of the double covalent bond in an oxygen molecule andthe triple covalent bond in a nitrogen molecule.

O O

(a) (b)

N N

O

H H

figure 2.13 The outer shell electronsof a water molecule.

H H

H

H

C

figure 2.14 The covalent bondsin methane.

Diatomic moleculesconsist of two atomsbonded together. They canbe atoms of the sameelement (homonuclear) or different elements(heteronuclear).

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table 2.5 Molecular shapes.

Shape Description Example

linear

V-shaped(also knownas angularor bent)

pyramidal

tetrahedral

The two atoms of a diatomic molecule must bein a straight line.

The double bonds of triatomic molecules suchas carbon dioxide repel each other—a straightline is as far apart as the electrons can get.

In a triatomic molecule, non-bonding pairs repelthe bonding pairs of electrons. As the bondingpairs also repel each other and try to achieve alinear shape, a V shape results.

A molecule consisting of a central atom andthree bonded atoms might be expected to forma flat triangular shape. As these central atomswill usually have a non-bonded pair that repelsthe three bonded pairs, the triangular shape isturned into a pyramid.

A molecule consisting of a central atom and fourbonded atoms has four sets of bonding pairstrying to get as far apart as possible. Theresulting shape is described as tetrahedral.

Cl Cl

Chlorine

CO O

Carbon dioxide

O

HH

Water(showing bonding and non-bonding electron pairs)

O

HH

Water(drawn without the non-bonding electron pairs,

the V-shape of a water molecule is more obvious)

N

HH

Ammonia

H

C

H

HHH

Methane

Two different atoms bonded together may not attract the shared electrons tothe same degree—one of the atoms may have a stronger pull on the electrons.Electronegativity is a measure of the attraction of a bonded element for theelectrons in the bond.

Table 2.6 shows the electronegativities of some elements. Theelectronegativity of elements tends to increase from left to right across theperiodic table and decrease down a group.

A measure of the electronattracting ability of anelement is known as itselectronegativity.

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When two elements are bonded together, the bonding electrons are moreattracted to the element with the higher electronegativity. This means that theelectrons spend more time around that atom and the region becomes slightlynegatively charged. The other atom, with lower electronegativity, becomesslightly positively charged.

The hydrogen atom of the hydrogen chloride molecule has anelectronegativity of 2.1 and the chlorine atom has an electronegativity of 3.2.The chlorine atom will attract the electrons in the bonding pair more stronglyand have a partial negative charge. The hydrogen atom will have a partialpositive charge. This is shown in figure 2.17.

The hydrogen chloride molecule has a polarised bond. It can also bedescribed as being a polar molecule or a dipole.

Since each hydrogen chloride molecule has regions of positive and negativecharge, the positive area of one molecule can be attracted to the negative areaof another. These intermolecular forces are much weaker than covalent or ionicbonding. This explains why hydrogen chloride exists as a gas at roomtemperature.

Formaldehyde has three bonds, two C–H bonds and a C=O bond. Carbonhas a higher electronegativity than hydrogen and so a polar bond existsbetween these atoms. Oxygen has a higher electronegativity than carbon and soanother polar bond exists in the formaldehyde molecule. These dipolescombine and overall the hydrogen end of the formaldehyde has a partialpositive charge and the oxygen end has a partial negative charge. Formaldehydeis a polar molecule. It is not just the electronegativity of the atoms present thatdetermines whether a molecule is polar.

Carbon dioxide has two polar bonds due to the different electronegativitiesof carbon and oxygen. However, unlike formaldehyde, carbon dioxide issymmetrical across the polar bonds and the polarity cancels out. The carbondioxide molecule is not a dipole.

table 2.6 The electronegativities of some elements (arranged according to their positions in the periodic table).

H2.1

Li Be B C N O F1.0 1.6 2.0 2.5 3.0 3.5 4.0Na Mg Al Si P S Cl0.9 1.3 1.6 1.9 2.2 2.6 3.2

figure 2.17 A hydrogen chloridemolecule has a polarised bond due to the difference in electronegativities of hydrogen and chlorine atoms.

H Clδ –δ +

δ + represents a small amountof positive charge.

δ – represents a smallamount of negative charge.

C O

δ+

δ+

δ+

δ –H

H

C Oδ –δ –

O

Formaldehyde

Carbon dioxide

figure 2.18Formaldehyde is anasymmetrical moleculeacross its polar bondswhile carbon dioxide issymmetrical and not adipole.

In a polarised bond theelectrons are not sharedequally between the twoatoms.

Notice that theelectronegativity differenceof HCl is 3.2 – 2.1 = 1.1.In NaCl, the difference is3.2 – 0.9 = 2.3. In NaClthe electrons are not justunevenly shared—eachsodium atoms totally loses an electron and each chlorine atom gainsone forming ions ofopposite charge.

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STRUCTURE,

Tetrafluoromethane (CF4) is also a symmetrical molecule and is tetrahedralin shape (figure 2.19). The bonds between carbon and fluorine are polariseddue to the electronegativity difference of the atoms but the symmetry of theatom means that the molecule is not a dipole. It does not have a negative andpositive end.

With an understanding of electronegativity and polar molecules the types ofintermolecular forces can be explained. There are three types of intermolecularforces: dipole–dipole attractions, hydrogen bonding and dispersion forces.Their features are summarised in table 2.7.

F

C

δ –

Fδ –

Fδ –F

δ –

table 2.7 The types of intermolecular forces.

Type ofintermolecular Explanation Examplesforce

dispersionforces (alsocalled van deWaal’s forces)

The weakest type of intermolecular forces.Due to a brief and temporary electrostaticcharge on an atom as electrons move about anatom or molecule causing a momentaryimbalance of charge.The more electrons in an atom or molecule thestronger the dispersion forces.Dispersion forces become insignificant whenmolecules are influenced by strongerintermolecular forces. They are onlyconsidered when stronger forces do not exist.

dipole–dipoleattraction

Stronger than dispersion forces (but weakerthan hydrogen bonding).The partial positive charge of one molecule is attracted to the partial negative charge of another.This type of force only exists when themolecules are dipoles.

hydrogenbonding (aspecial formof dipole–dipoleattraction)

The strongest type ofintermolecular forces.Hydrogen bonds only occurwhen a molecule contains ahydrogen atom bonded to oneof the highly electronegativeatoms—nitrogen, oxygen or fluorine.The hydrogen atom has arelatively high partial positivecharge due to a large differencein electronegativity.The hydrogen atom is stronglyattracted to a non-bonding pairof electrons of a highlyelectronegative atom of anearby molecule.

Dispersion forces result as the electronsmomentarily spend more time around one atomthan the other. Symmetrical molecules such ashydrogen, chlorine and tetrafluoromethane arenon-polar. The forces of attraction between suchmolecules are due to dispersion forces.Noble gases rarely combine with other elements.They can be liquefied and solidified (at lowtemperatures) due to weak dispersion forces.

Hydrogen and chlorine molecules.

HH

ClCl

Molecules in solid hydrogen chloride arrangethemselves so that their positive and negativeends are aligned.

δ+ δ –H Cl

δ+ δ –H Cl

δ+ δ –H Cl

δ+ δ –H Cl

δ+ δ –H Cl

δ+ δ –H Cl

δ+ δ –H Cl

δ+ δ –H Cl

δ+ δ –H Cl

Hydrogen bonding of (a) ammonia, (b) water and (c) hydrogen fluoride.

δ–N

δ+H

δ+Hδ+

H

δ+H

δ+H

δ–F

δ+H

δ–F represents

hydrogen bond

δ–N

δ+H

δ+H

δ–O

δ+Hδ+

H

δ+H

δ–O

δ+H

(a)

(c)

(b)

figure 2.19Tetrafluoromethanehas polarised bondsbut is not a dipole.

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....................................................................................................................................................Polar bonding within water molecules

A stream of water bends when charged rods are brought near it. A stream ofhexane is almost unaffected by the presence of the rods.

PurposeTo demonstrate the dipolar nature of water molecules.

Method1 Fill one burette with water and another with hexane. Place a beaker under

each burette.2 Open the burette taps so that fine continuous streams of liquid run from

the burettes into beakers from a height of about 40 cm. (A tall goosenecktap, such as is often found on a front bench, may be used as analternative source of the water stream.)

3 Bring the positively charged glass rod close to, but not touching, each ofthe streams of liquid. The water stream should be attracted markedlytowards the charged rod.

4 Repeat with a negatively charged rod.

TheoryOxygen is more electronegative than hydrogen. As a result it attracts theshared bonding electrons in the water molecule more strongly than thehydrogen atom. The oxygen atom therefore behaves as if it has a smallnegative charge, and the hydrogen atom behaves as if it has a small positivecharge. The molecule is described as dipolar.

When charged rods are placed near a fine stream of water, the chargedends of the molecule become attracted to the charged rod, causing the waterto deviate from its normal flow.

The orientation of atoms in the hexane molecule, and the similarelectronegatives of carbon and hydrogen, cause hexane to be non-polar. (Models may help make these ideas easier to understand.) As aresult, hexane is attracted to a charged rod to a much lesser extentthan water. However, the charged rod does induce dipoles in theliquid, and slight attraction does occur.

Discussion1 Explain the principles behind charging one article by rubbing it

with a piece of material. 2 Why is the stream of water attracted to both the positively and

negatively charged rods?3 Sketch a diagram of a water molecule.4 Why is hexane less strongly attracted to the charged rods?

ExperimentExperiment 2.1materials

• glass or celluloseacetate rod rubbed withsilk (positively charged)

• polystyrene rod or aballoon rubbed with wool(negatively charged)

• 2 × burettes• 2 × 250 mL beakers• 2 × stands, bossheads

and clamps• de-ionised water• 50 mL hexane (C6H14)

safety

• Wear safety glasses anda laboratory coat forthis experiment.

• Hexane is extremelyflammable. The vapouris irritating to the skin,eyes and respiratorysystem. Ensure that thesolvent does not remainopen in the laboratoryfor extended periods.

Properties of molecular substancesSince the forces between molecules in molecular substances are weak,molecular substances tend to have low melting and boiling points, and many ofthem exist as gases or liquids at room temperature. Hydrogen, oxygen, chlorine,carbon dioxide and methane exist as gases at room temperature, and water,ethanol (CH3CH2OH), formaldehyde and tetrachloromethane (or carbontetra-chloride, CCl4) exist as liquids at room temperature. Notice that larger

figure 2.20

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molecules tend to be liquids and involve the stronger types of intermolecularforces. Only quite large molecules, with stronger intermolecular forces, exist assolids. These solids are usually soft and pliable. The molecules in the solid canmove around each other to some extent, allowing it to bend. Waxes and plasticsare some examples of solid molecular substances.

The electrons in covalent bonds are held tightly and do not move to carry anelectric current. Molecular substances do not contain ions either. This meansthat molecular substances are non-conductors of electricity.

Some molecules are attracted to water molecules and can be dissolved.Ethanol for example is soluble in water, while baby oil is insoluble.

Names and formulae of molecularsubstancesSome molecular substances are elements. For example oxygen, nitrogen andchlorine gases are made up of molecules consisting of two atoms. This is shownby their formulae—O2, N2, Cl2. When we talk about oxygen gas we mean O2

molecules not atoms. Molecular elements are known by the name of theelement.

Molecular compounds consist of more than one type of atom. Prefixes areused to indicate how many atoms of each type are in a molecule (see table 2.8).The first element in the formula is given its full name and the second elementhas its name shortened and -ide is added as a suffix. For example, the namecarbon dioxide tells us that the molecule consists of one carbon atom and twooxygen atoms. Note the prefix mono is not used on the first element.

Some molecular compounds are given common (or trivial) names. Carbontetrahydride (CH4) is more commonly known as methane (found in natural gas).Table 2.9 shows the name and formula for some common molecular substances.

Covalent latticesNot all substances made up of covalently bonded atoms are molecules.A covalent network lattice is a vast lattice of non-metallic atoms covalentlybonded together. In case you are tempted to consider covalent bonds to beweaker or inferior to ionic bonds you may find it interesting to learn that oneexample of a covalent network lattice is the hardest naturally occurringsubstance on earth—diamond.

Diamond consists of a lattice of carbon atoms. Each carbon atom iscovalently bonded to four other carbon atoms. The four surrounding atoms arepositioned at the corners of a tetrahedron. This three dimensional structure,held together by strong covalent bonds, forms an extremely stable three-dimensional lattice as illustrated by figure 2.21.

table 2.8The prefixes used to namemolecular compounds.

mono 1di 2tri 3tetra 4penta 5hexa 6hepta 7octa 8nona 9deca 10

table 2.9 The names of some common molecular compounds and their formula.

Name Formula

sulfur dioxide SO2

carbon monoxide COcarbon tetrachloride CCl4

dinitrogen dioxide N2O2

A covalent networklattice consists of manyatoms held in fixedpositions by covalentbonds.

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figure 2.22 Thelayered structure of agraphite lattice.

Strong forceswithin layer

Weak forcesbetweenlayers

The strength of the carbon lattice makes diamond extremely hard. It takes alot of energy to break down the diamond lattice. At extremely hightemperatures, the covalent bonds within the lattice break and, rather thanforming a liquid, the carbon atoms form individual atoms of a gas. This processis called sublimation; diamond is said to sublime.

The valence electrons of a diamond lattice are held tightly in covalent bonds.They are said to be localised. Without free electrons to carry an electric current,diamond does not conduct electricity.

Other covalent network lattice substances have similar properties todiamond. Some of these are described in table 2.10.

Graphite is another covalent lattice of carbon (figure 2.22). In graphite, eachcarbon atom is attached to three other carbon atoms by strong covalent bondsin a flat layer structure called a covalent layer lattice. Graphite has only three ofcarbon’s four valence electrons involved in covalent bonding. The fourthvalence electron is delocalised. The delocalised electrons are free to movebetween the layers of the covalent lattice.

Graphite, like diamond, has a high melting temperature due to the largeamount of energy required to break the strong covalent bonding of the lattice.Unlike diamond, graphite is able to conduct an electric current as thedelocalised electrons are free to move. While bonding within the layers ofgraphite is strong, the forces between layers are weak and they readily slide overeach other. This gives graphite a soft greasy texture and makes it useful as alubricant. Lead pencils actually contain graphite and during writing layers areleft behind on the paper as the pencil slides over it.

figure 2.21 The covalent networklattice of diamond. Each carbonatom is covalently bonded to fourother carbon atoms.

C

CC

CC

C

C

CC

C

CC

C C

C

C

C

C

CC

CC

CC

C

CC

C

C

C

C

C

table 2.10 The properties of some covalent network substances.

Material Formula Melting Boilingtemperature (°C) temperature (°C)

diamond* C >3550 –silicon Si 1410 2680silicon carbide* SiC 2200 –silicon dioxide SiO2 1700 2230*Diamond and silicon carbide do not have boiling temperatures as they undergosublimation at normal pressures.

A substance is said tosublime if it changesfrom a solid to a gaswithout passing througha liquid phase.

Localised electrons belongto a particular atom.Delocalised electrons areshared by many atoms atthe same time.

Terms and definitions

Worksheet3

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STRUCTURE,

...........................Questions13 Draw electron dot and structural formulae for the

following molecules. a fluorine (F2)b hydrogen fluoride (HF)c methane (CH4)d tetrachloromethane (CCl4)e ethane (C2H6)f ammonia (NH3)g carbon dioxide (CO2).

14 How many lone pairs would you expect atoms ofthe following elements to have when they formcovalent bonds with other non-metal atoms? a Hb Fc Cd Ne Ar

15 Draw electron dot formulae for the followingmolecules and identify the number of bondingand non-bonding electrons in each molecule.a NCl3b CH4c H2S

16 Consider the table of electronegativities(table 2.6). Covalent bonds can form betweenthe following pairs of elements in a variety ofcompounds. Use the electronegativity valuesgiven in the table to identify the atom in eachpair that would have the larger share ofbonding electrons.a S and Ob C and Hc C and Nd N and He F and Of P and F

17 Are the following molecules polar or non-polar?(Draw structural formulae or make models tohelp you to answer this question.)a CS2b Cl2Oc SiH4

18 Between which of the following moleculesare there

i dipole–dipole interactions?ii hydrogen bonds?

a NH3b CHCl3c CH3Cld F2Oe HBrf H2Sg HFh H2Oi H2

19 The melting temperatures of four of the halogensare given in table 2.11 below. Describe andexplain the trend in melting temperatures ofthese elements. (Refer to the periodic table toestablish where the halogens occur in the table.)

20 a Write chemical formulae for these molecularsubstances.

i hydrogen gasii dinitrogen monoxide

iii phosphorous trihydrideiv carbon tetrachloridev dihydrogen sulfide

b Name the following molecular substances.i F2

ii NO2iii CH4iv NF3v N2H4

21 Explain the following properties ofa diamond andb graphitein terms of their respective structures.

i high melting temperatureii hardness or softness

iii ability or inability to conduct electricity22 Explain the following uses in terms of the

structures of graphite and diamond.a Graphite is used as a lubricant.b Diamond is often used as an edge on saws

and a tip on drills.

table 2.11

Halogen Melting temperature (°C)

fluorine (F2) –220chlorine (Cl2) –101bromine (Br2) –7iodine (I2) 114

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....................2.4 Metallic bondingThe properties of metals are a result of the way metal atoms bond together. Thischaracteristic form of bonding is called metallic bonding.

Metals consist of positive ions in a regular, three-dimensional lattice withdelocalised valence electrons, which are free to move throughout the lattice.Compare this to ionic substances in which electrons belong to particular ionsor molecular substances in which electrons are shared between particularatoms. However, metallic bonding achieves a similar result to the other forms ofbonding. By setting free the valence electrons into a common pool or sea ofelectrons, the metal ions achieve a stable electron configuration similar to thenoble gases.

The metallic bonds are the attraction between positive ions and the freemoving electrons. The strength of these bonds can be seen in the fact thatmetals (except mercury) are solids at room temperature.

Metals can be hammered into shape or drawn into wires because the atomscan be forced to move without breaking up the lattice. The positive ions andelectrons just move around to take up the new shape.

If a piece of metal is connected into an electric circuit it will conductelectricity. The delocalised electrons are free to move within the metal, allowingan electric current to flow.

figure 2.23 The structure of a metal. Thepositive ions are arranged in a fixed, three-dimensional lattice. The valence electrons arefree to move and can be thought of as a seaof negative charge.

++++

+ +

++ + +

+ +

positive ions

‘sea’ of electrons

figure 2.24 As a force is applied to a metal, the positive ions move and the electronsadjust to the new shape.

Metal is hit here

A layer of positive ions has been displaced,but the delocalised electrons move tocompensate. The essential features of themetallic structure and bonding are unchanged.

++++

++++

++++++++

++++

++++

Metallic bonding is theforces of attractionbetween positive metalions in a fixed lattice andtheir delocalised electrons.

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figure 2.25 Metals conduct electricity because the electrons are able to move freelythrough the lattice.

flow of electrons

+ + + + + + + + +

++++++++

This end of thecircuit is connectedto the negativeterminal of the battery.

When the circuit iscomplete, electronswill flow throughthe metal.

This end of the circuit isconnected to the positiveterminal of the battery.

....................................................................................................................................................Growing metal crystals

PurposeTo grow crystals of a range of metals and to observe their shapes.

Method1 Add 0.5 g of agar to 40 mL of de-ionised water in a beaker and warm

gently, stirring until the agar is dissolved. Remove from the heat. 2 Add 0.3 g of one of the available solids to the agar solution, stirring

until it dissolves.3 Pour the agar solution into a Petri dish and place the strip of zinc in the

centre. Do not move the Petri dish until the agar has set.4 Once the agar has set, place a lid on the Petri dish to prevent the agar

from drying out.5 Observe the formation of metal crystals over the next 1–2 days.6 Repeat the above procedure using each of the metal salts listed.7 Sketch the appearance of the different metal crystals that have been

grown in your class.

TheoryZinc is a more reactive metal than silver, lead, copper or tin. When a pieceof zinc is placed in solutions that contain the positive ions of these othermetals, the zinc metal reacts to form ions and the ion of the less reactivemetal reacts to form its metal. This process can be summarised by the word equation:

zinc + silver nitrate → zinc nitrate + silver

Each of the other compounds react in a similar way. These reactions areknown as displacement reactions.

When a metal is formed quickly, as would be the case if thedisplacement reaction occurred in aqueous solution, the metal crystals tendto be small. It is difficult to see the details of their shape. If, however, theformation of the metal is slowed down, as when the reaction occurs in agar,larger crystals form and it is easier to examine their shape.

If your crystals have a regular, uniform shape, you can conclude that theparticles in the metal are arranged in a regular manner in a crystal lattice.

ExperimentExperiment 2.22.2

STRUCTURE,

BONDING AND PROPERTIES OF MATERIALS

materials

• 0.5 g agar• 40 mL de-ionised water• approximately 0.3 g of

one of the followingsolids: AgNO3,Pb(NO3)2, CuSO4, SnCl2

• 1 cm × 4 cm strip ofclean zinc sheet

• 250 mL beaker• Petri dish• hotplate or Bunsen

burner, tripod and gauze mat

• glass stirring rod

safety

• Wear gloves, safetyglasses and a laboratorycoat for this experiment.

• AgNO3 causes burnsand stains skin andclothing.

• Pb(NO3)2 is poisonous.• CuSO4 irritates the skin

and eyes.• SnCl2 can irritate the

skin and eyes.

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The properties of metalsMetals tend to have a shiny appearance (lustre), and high melting and boilingpoints. Mercury is a liquid but all other metals are solids at room temperature.

Metals are good conductors of electricity. Some metals conduct an electriccurrent better than others. They are also good conductors of heat.

The ability of a metal to be bent or hammered into shape is called itsmalleability. Some metals can be drawn into a wire. They are said to be ductile.

table 2.12 The physical properties of metals and what they indicate about metallic structure.

Property

Metals conduct electricity in the solidstate.Metals are lustrous or reflective.Metals are malleable and ductile.

Metals tend to have high melting andboiling temperatures.Metals generally have high densities.

What this tells us about structure

Metals have charged particles thatare free to move.Metals can reflect light.The forces between the particlesmust be able to adjust when theparticles are moved.The forces between particles in ametal are strong.The particles are closely packed in a metal.

...........................Questions23 Which properties listed in table 2.12 could be

used to tell the difference between a piece ofnickel and a piece of tin?

24 What is the difference between malleabilityand ductility?

25 If two metals have similar melting points,electrical conductivity and lustre, what other

feature could be used to separate the twometals?

26 Explain why mercury is a liquid at roomtemperature while iron is a solid.

27 The bonds in an ionic solid can be described as directional while those in a metal are non-directional. Why are these terms used?

2.5 Organic compoundsCarbon atoms have four valence electrons and can form four covalent

bonds. Hydrocarbons are molecules consisting only of hydrogen and carbonatoms. The smallest hydrocarbon is methane consisting of one carbon and fourhydrogen atoms but there are an infinite number of possible combinations ofhydrogen and carbon, some of which are quite large. Figure 2.26 shows someexamples of fairly simple hydrocarbon molecules.

Discussion1 Explain the principles behind charging one article by rubbing it with a

piece of material. 2 Why is the stream of water attracted to both the positively and negatively

charged rods?3 Sketch a diagram of a water molecule.4 Why is hexane less strongly attracted to the charged rods?

Malleable substancesare able to be beateninto shape.

Ductile substances canbe drawn into a wire.

Hydrocarbons consistonly of hydrogen andcarbon atoms.

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· pdf file11.4 qualitative analysis using physical and ... some common alloys are listed in...

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