chemistry as level

7/27/2019 Chemistry AS level

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BONDING AND STRUCTURE

We are living in a world with millions of substances having properties vastly different from one another. It is the way the atoms

are held together (bonding) and how they are arranged in space (structure) that accounts for the quite distinct properties of these

different substances.

When elements form compounds, they either lose, gain or share electrons so as to achieve stable electron configurations. This

simple idea forms the basis of the electronic theory of bonding : a chemical bond is a specialarrangement of electrons between

atoms by which the resulting nuclei and electrons become morestable.

There are three main types of bonding - ionic, covalent and metallic - which involves respectively transfer, sharing and pooling of

electrons between atoms to achieve stable electronic arrangements.

Type of bonds IONIC BOND COVALENT BOND METALLIC BOND

formed between metals and non-metals

way to achieve stable

electronic arrangement

non-directional

nature of bonds

electrostatic attraction between

nuclei and shared electrons

structure adopted

example sodium chloride iodine diamond sodium

nature of forcesholding particles

together

strong metallic bondbetween ions and

sea of electrons

properties :

(1) m.p. / b.p.

(2) electrical

conductivity

(3) solubility

IONIC BONDING

A simple model of ionic bonding where electrons are being transferred from metal atoms to non-metal atoms so that the resulting

ions obtained a full outer shell of electrons :


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Bonding 2

Ionic Radii

The electron cloud of any atom or ion has no definite limit, thus the size

of an atom cannot be defined in a simple and unique manner. We could,

however, measure very precisely the distance between two nuclei by

means of electron density maps.

Class Work

The figure shows the electron density map for sodium chloride.

(a) Decide which is the sodium ion and which is chloride ion.

(b) Use a ruler and the map scale to obtain approximate values for the

ionic radii of sodium and chloride ions.

The sizes of ions can be conveniently compared by measuring their ionic radii from electron density maps.

1 Comparison of sizes of ions with theirparent atoms :

(i) cations are smaller than their parent atoms (removal of a complete outer shell leads to a contraction of electron cloud)

(ii) anions are larger than their parent atoms (affinity of incoming electrons results in an expansion of electron cloud)

cations are usually smaller than anions

2 Comparison of sizes ofisoelectronic particles :

In any isoelectronic series, all the particles carry the same number of

electrons but have a progressively increasing number of protons.

Along an isoelectronic series

thep/e ratio (no. of protons / no. of electrons) increase effective nuclear charge increase result in a contraction of the electron cloud of the ion ionic radii decrease along an isoelectronic series

Structure of Ionic Lattices

There are two common arrangements for most ionic compounds, namely the sodium chloride and caesium chloride structures.

They can be conveniently represented by unit cell, defined as the smallest portion of structure which when repeatedly stacked

together at all directions, can reproduce the entire lattice.

NaCl structure

It is made up of two inter-

penetrating face-centered

cubic lattices (f.c.c.) of

each type of ion (Na+, Cl

-)

to form a simple cubic

lattice with each type of

ion occupying alternate

corners

CsCl structure

It is made up of two

inter-penetrating simple

lattices of each type of

ion (Cs+, Cl

-) to form a

body-centered cubic

lattice (b.c.c.)

Class Work

(a) Sketch the unit cell of the two different structures.


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Bonding 3

(b) Deduce the number of each type of ion per unit cell for each structure.

NaCl structure : CsCl structure :

(c) State the respectiveco-ordination numbers of each type of ion in the two structures.

NaCl structure : CsCl structure :

Theoretically, ionic structures tend to have as higher co-ordination numbers as possible, because it allows maximum attractive forces operate

between oppositely charged ions, making the resulting lattices more stable. Sodium chloride, however, does not adopt the caesium chloride

structure because of the difference in relative ionic sizes. However, Na+

ions, being much smaller than Cs+

ions, cannot accommodate more than

6 Cl-

ions before these anions repel each other too strongly for a stable arrangement.

Class Work

Study the following two unit cells and deduce which one belong to calcium fluoride and which one to zinc sulphide.

Physical Properties of Ionic Compounds

Compounds of ionic solids are :

1 non-volatile with high melting points and high boiling points

2 good conductors of electricity when molten, but non-conductors when solid

3 soluble in water, but insoluble in non-aqueous solvents such as tetrachloromethane (CCl4)

4 hard and brittle

Ionic crystals are hard because of the

strong attractions between opposite

charged ions. However, they are also very

brittle and may be split clearly (cleaved)

using a sharp-edged razor. When the

crystal is tapped sharply along a particular

plane it is possible to displace one layer of

ions relative to the next. As a result of this

displacement, ions of similar charge come

together. Repulsion then occurs, forcing

apart the two portions of the crystal.


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Bonding 4

COVALENT BONDING

According to valence bond theory, a covalent bond is formed by sharing the valence (outermost) electrons :

This model is inadequate because it suggests a pair of electrons to be fixed between the two nuclei, yet it is impossible to locate the

exact position of electrons. A better model using the idea of overlapping orbitals/charge clouds is adopted :

The actual charge distribution for a covalent molecule could be obtained from an

electron density map, which shows a region of negative charge (shared electrons)

located between two positive charges (nuclei).

The covalent radius of an atom is defined as half the internuclear distance between

two covalently bonded atoms in the molecule of the element.

Class Work

What is the covalent radius of hydrogen ?

Despite the limitation of valence bond theory, it is nevertheless very useful in making predictions about the electronic structures of

covalent molecules by dot-and-cross diagrams.

Class Work

Draw electronic structures for the following species :

(a) NCl3 (b) H2S (c) CHCl3

(d) C2H4 (e) N2 (f) CO2

Forpolyatomic ions, such as NH4+ and CO32-, it is important to distinguish between the ionic bonding which binds these ions toother ions, and the covalent bonding which binds the atoms within each ion.

Class Work


(a) OH- (b) CO32- (c) HCO3

-


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Bonding 5

Dative Covalent Bond

In a normal covalent bond, each atom provides one electron for the shared pair. In a few compounds, however, a bond is formed

by the sharing of an electron pair which are provided by one atom. This kind of bonding is known as dative covalentbond or

co-ordinate bond, since both electrons in the bond are donated by one atom.

Ordinary covalent bonds are represented by a short line as a shorthand for an electron pair. Dative covalent bonds are represented by an arrow

showing the direction of donation. However, once a dative bond is formed, you should not imagine them being different from ordinary bonds,

because electrons are identical irrespective of where they are coming from.

Class Work


(a) CO (b) NH4+

(c) NO3-

Limitation of Octet Rule

So far, we have considered only those covalent compounds which obey the octet rule. However, there are a number of cases

where the rule does not apply.

1 Compounds with more than eight valence electrons per atom

The octet rule always applies for elements in the first two periods of the Periodic Table : each atom cannot hold more than

eight electrons in its valence shell. Starting from period 3, however, elements could make use of their vacant d-orbitals to

accommodate more than 8 valence electrons, apparently breaking the octet rule.

Class Work


(a) PCl5 (b) SO42- (c) PO4

3-

2 Compounds with fewer than eight valence electrons per atom

When elements with fewer than four outer shell electrons form compounds they usually lose those electrons to form ions. For

small atoms, however, the relevant ionization energies may be so high that covalent bonding occurs instead. Since there are

fewer than four electrons available for sharing, there will then be fewer than eight outer shell electrons per atom in the

resulting compound. Such compounds are often called electron-deficient.

Class Work


(a) BF3 (b) BeCl2 (c) AlCl3

There is a tendency for these electron-deficient compounds to achieve 'stable octet' by accepting a lone pair of electrons from

neighbouring molecules to form a dative covalent bond.

NH3 BF3 BeCl2 (polymer) Al2Cl6(dimer)


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Bonding 6

Shapes of Molecules

The electron-pair repulsion theory allows us to make predictions about the shapes of a great many covalent molecules :

1 Each covalent bond is regarded as a pair of electrons. An electron pair may be a shared pair, lone pair, single bond, double

bond or triple bond.

2 Electron pairs around a central atom tend to get as far apart from each other as possible, so as to minimize the electrostatic

repulsion among them.

3 Non-bonding pair of electrons exerts a greater repelling effect than bonding pair of electrons does. Therefore,

lone pair - lone pair lone pair - bond pair bond pair - bond pair

repulsion repulsion repulsion

no. of electron

pairs around

central atom

Basic skeleton Bond angle examples

2

linear

BeCl2 CO2

3

trigonal planar

BF3 SO3 SO2

4

tetrahedral

CH4 NH3 H2O

5

trigonal

bipyramidal

PCl5 ClF3 ICl2-

6

octahedral

SF6 BrF5 XeF4


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Bonding 7

Class Work

For each of the following molecules, draw a three-dimensional structure showing the positions of the bond electron pairs and

lone electron pairs (if any) of the central atom.

(i) PCl4+

(ii) PCl6-

(iii) PCl3

A few comments on bond angles :

1 CH4 molecules are perfectly symmetrical. The bond angle for H-C-H is that of a regular tetrahedron, 109.5 . Although

sharing the same molecular skeleton, NH3 and H2O have different bond angles (107 for H-N-H and 105 for H-O-H).

The region in space occupied by a non-bonding pair of electrons is smaller and closer to the nucleus of an atom than a bonding pair.

Bonded pairs of electrons are drawn out between the nuclei of the two covalently-bonded atoms. This means that a lone pair can exert a

greater repelling effect than a bonded pair and this results in a decreasing bond angle from CH 4 to NH3 to H2O.

2 NH3 and PH3 are both pyramidal with four pairs of electrons around the central atom, one of which is a lone-pair. The bond

angle for H-N-H in NH3 is 107 while that for H-P-H is 90.

N atom is smaller than P. It follows that N-H bond pairs lie much closer to the central atom than P-H bond pairs do. The repulsion amongN-H bond pairs are thus much greater than that among P-H bond pairs, resulting a decrease in bond angle down the group.

Giant Covalent Structures

The sharing of electron pairs does not necessarily result in discrete molecules. In a few cases, covalent bonding is extended

indefinitely in three dimensions. Such giant structure is well illustrated by the two allotropes (different structural forms of the

same element) of carbon, namely diamond and graphite.

The properties of diamond and graphite are in dramatic contrast. Since both solids consists of identical carbon atoms only, the

differences in properties must be entirely due to differences in structures.

DIAMOND GRAPHITE

structure

Each C atom is covalently bonded to 4 other C

atoms in a tetrahedral manner to form a 3-D

giant network

Each C atom is covalently bonded to 3 other C atoms

in trigonal planar to form a multi-layer structure.

Adjacent layers are held by van der Waals' forces

co-ordination no./ hybridization

4 / sp3 3 / sp2

3550 C 3700 Cm.p. / b.p.

strong C-C covalent bonds have to be broken in melting / boiling

strength hard; strong & directionalC-C bonds restrict

relative motion between C atoms

soft; weak van der Waals' forces allow layers

to slip over each other easily

electrical

conductivity

non-conductor of electricity;

all electrons are localizedin covalent bonds

good conductor of electricity;

each C atom has a delocalizedelectron which

can move freely along the same layer

uses jewelries, glass cutters electrodes, lubricants, pencils


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Bonding 8

INTERMEDIATE TYPE OF BONDING

In many compounds, the bonding is intermediate in character between 'pure'

ionic and 'pure' covalent. We can discuss intermediate type of bonding in two

ways, either by seeing :

1 how ordinary ionic bonds can have covalent character; and

2 how ordinary covalent bonds can have ionic character.

Covalent Character of Ionic Bonds

The model of ionic bonding so far assumes the complete transfer of electrons

from one atom to another giving separate spherical ions. However, the

spherical charge cloud in large anions may be easily distorted, or polarized,

by neighbouring small cations, because the nuclear charge of the anion is

insufficient to hold its electron cloud firmly. Such distortion leads to a certain

overlap of charge clouds between cations and anions, and hence a degree of

covalent character. The distortion of an electron cloud by a neighbouring

charged particle is called polarization.

Evidence of whether an ionic compound has significant covalent character comes from a comparison of the lattice energy values,

one derived from theory and the other determined from experiment.

Assumptions in deriving the theoretical lattice energy of an ionic compound :

1 ions are spherical

2 charge is uniformly distributed throughout the ions

3 forces operating between ions are only electrostatic in nature

theoretical value experimental value

underlying

principle

applying Law of Electrostatics

L.E. q q

r

1 2

applying Hess' Law

L.E. = B + C + D + E - A

NaCl

LiF

AgCl

ZnS

-777

-1021

-769

-1986

-781

-1033

-890

-2435

Cations such as Al3+, Ag+ and Zn2+ have small ionic radii and high nuclear charge, both give rise to a high charge density

(i.e. charge/volume ratio), thus distorting the electron cloud of neighbouring anions to a greater extent. These ions are said to

possess a high polarizing power.

Electrons in larger anions are further away from the nucleus and are less firmly held by the nucleus than smaller anions. Large

anions are said to have a high polarizability (the ability to become polarized).


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Bonding 9

Ionic Character of Covalent Bonds

Polar covalent bond and dipole moment

The model of covalent bonding used so far assumes the equal sharing of a pair of electrons between two atoms. Very often the

electrons are not shared equally, as in the case when a bonding pair of electrons is attracted more to one nucleus than the other (e.g.

Cl atom in H-Cl molecule is more electronegative than the H atom). This makes the centre of negative charge not coincide with

the centre of positive charge, resulting in the formation of a dipole with two equal and opposite charges (+q and -q) separated by a

distance d

This bond is said to be polarized and described as a polar covalent bond. There are four ways of representing bond polarity :

The extent of bond polarization can be measured in terms ofdipole moment, , which is given by = q x d. Dipole moments are

usually expressed in Debye (D) unit (e.g. dipole moment of HCl molecule is 1.1 D)

1 Dipole moments are vector quantities. For a molecule with more than one polar bond,

the dipole moment is given by the vector sum of the dipole moments of various polar bonds. If

the vector sum is zero, the dipole moment of the molecule is zero, and the molecule is described

as non-polar.

2 The greater the overall dipole moment, the more polar the molecule is.

3 Polarity of a liquid can be tested by studying the effect of a charged rod on a stream of

liquid from a burette. Any deflection of the stream indicates that the liquid is polar.

4 Dipole moments can provide useful information about the structure of molecules. As an example, the

zero dipole moment of CO2 shows that the molecule must be linear such that the dipole moment of

each C=O bond cancels each other out. On the other hand, the existence of a net dipole moment for

SO2 molecule indicates that the molecule contains polar bonds which are not linearly arranged.

Class Work

By drawing a 3-D structure showing the bond electron pairs and lone electron pairs (if any), state

whether or not the molecule possesses a dipole moment.

(i) BCl3 (b.p.13C) (iii) CCl4 (b.p.77C) (v) CO2 (b.p.-78C)

(ii) NCl3 (b.p.71C) (iv) CHCl3 (b.p.62C) (vi) SO2 (b.p.-10C)

Electronegativity

In polar covalent bonds, the unequal sharing of the bonded electron pairs is caused by a difference in electron attracting ability of

the bonded atoms. This electron attracting ability of the atom is described as electronegativity of the element.

1 The Pauling's scale is used to measure electronegativities of different elements :(i) a value of 4.0 is assigned to the mostelectronegative element, fluorine

(ii) all other elements having a lower electron attracting ability than fluorine are assigned lower electronegativity values


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Bonding 10

2 Smaller atoms with greater charge density tend to attract electrons more tightly than larger ones, i.e. electronegativity

increases across a period anddecreases down the group

3 Electronegativity differences govern the nature of chemical bonding

bondelectronegativity difference

between bonded atomstype of bond

ClCl 3.0 - 3.0 = 0 covalent

ClH 3.0 - 2.1 = 0.9 polar covalent

F- Li+ 4.0 - 1.0 = 3.0 ionic

4 Both electronegativity and electron affinity refer to the attraction by an atom for electrons. The essential difference between these terms liesin the fact that electron affinity refers to the attraction between incoming electrons (external) and isolated atoms; whereas electronegativity

refers to the attraction between bonding electrons (internal) and an atom in a covalent compound.

INTERMOLECULAR FORCES

There are three major types of interactions operating between simple covalent molecules, depending on the extent ofpolarity of

the molecules :

1 permanentdipole-dipole interactions

2 dipole-induced dipole interactions

3 temporary (induced) dipole-induced dipole interactions (or van der Waals forces in general)

Dipole - Dipole InteractionPolar molecules, or sometimes called dipoles, are attracted to one another when

the positive end of one molecule is oriented toward the negative end of its

neighbour. Such electrostatic attractions between permanent dipoles are

considerably strong, and explains why polar compounds usually have

relatively higher melting or boiling points than non-polar molecules of

similar molecular size.

Dipole - Induced Dipole Interaction

It arises when a polar molecule, in the vicinity of another molecule (which is non-polar), has the effect of polarizing the second

molecule. The induced dipole can then interact with the first molecule.

Induced Dipole - Induced Dipole Interaction

Non-polar molecules, though having no permanent dipole moment, do have temporary, constant changing dipoles due to the

continuous motion of the electrons.

1 Electrostatic attractions between temporary induced dipoles are weaker than that betweenpermanentdipoles.

2 Electron cloud of larger molecules are more readily to be polarized than smaller ones, resulting in a stronger intermolecular

attraction. This is why substances with larger molecular size generally have higher melting and boiling points.3 Temporary induced dipoles are present in BOTH polar and non-polar molecules (as long as they have polarizable electron

clouds). For large polar molecules, the attractive forces due to temporary induced dipoles may even be greater than those due

to the permanent dipoles.

-

- -

- - -

- - - -

- -


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Bonding 11

Hydrogen Bonding

Hydrogen bonding is often regarded as a special kind of permanent dipole-dipole attractions. This is formed between a lone pair

of electrons from a highly electronegative element (ie. F, O or N) and a hydrogen atom, which itself bonded to another very

electronegative element.

+ - + - + -

F H F O H N N H O

Hydrogen bond is by far the strongesttype of intermolecular forces, with strength about one-tenth of an ordinary covalent bond.

This relative strength has some important and interesting consequences :

1 Boiling points of hydrides of group V, VI and VII

In general, boiling point increases down the group, owing to the increase in

molecular size (greater v.d.W. forces)

However, the first member of the hydride from Group V, VI & VII have

exceptionally high boiling points, because of the existence of intermolecular

hydrogen bonds

Such abnormal phenomenon does not appear in Group IV hydrides, as the

first member here (i.e. carbon) is obviously not electronegative enough to

form intermolecular H-bonds

2 Structure and density of Ice

Abnormal behaviour of ice :

1 Ice floats on water while solid HF does not float on its own liquid.

2 When liquid water is heated from 0C, its density increases for a short

while (below 4C) before decreasing in the usual way.

Explanation :

H2O has two lone pairs of electrons and could form at maximum two hydrogen

bonds per molecule. Each H2O molecule is, therefore, surrounded

tetrahedrally (2 covalent bonds and 2 hydrogen bonds) to form an 'open' cage-

like structure.As ice melts, some of these hydrogen bonds are broken. The partial collapse

of such an open structure allows water molecules to pack closerto each other

density of water > density of ice

3 Hydrogen bondings in proteins and DNA

Hydrogen bonding plays a crucial role in the structures of many biochemical substances such as proteins and DNA :

Protein chains are held in close

proximity to one another by hydrogen

bonds formed between the N-H and

C=O groups of the neighbouring chains

In DNA, the two helical strands of nucleic acid chains, which controls

reproduction and inheritance, are linked together by a series of

hydrogen bonds


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Bonding 12

4 Relative viscosity

Viscosity of a liquid is related to the ease with which molecules can move past one another in the liquid.

Viscosity of alcohols increases in the following order : propan-1-ol < propane-1,2-diol < propane-1,2,3-triol

Each alcohol molecule forms hydrogen bonds with its neighbours through hydroxyl (-OH) groups. As the number of hydroxyl

groups increases, so does the extent of hydrogen bonding in the liquid; making it more viscous.

5 Pressure-temperature diagram

vapour pressure of a liquid at a particular temp. is the pressure exerted by the vapour molecules, which are in equilibrium

with the liquid molecules at that temp.

graphical representation :

-- at each temp., the vapour pressure of the more

volatile substance is always higher than the less

volatile one

-- boiling point is defined as the temp. at which the

vapour pressure of the substance is equal to the

atmospheric pressure

-- note that liquids with lowervapour pressures (ie.

less volatile) will have correspondingly higherb.p.

the curve represents the variation of boiling points with differentpressure and temp. (notice that the y-axis now indicates external

pressure rather than vapourpressure of the substance)

each area included by the curve represents the conditions of temp.and pressure under which a particular phase is stable (or dominant)

every point on the curve represents the conditions of temp. &pressure at which the two phases co-exist in equilibrium

the concept can be further elaborated to represent equilibria

involving all the 3 phases :

interpretation of phase diagram

AB :

BC :

BD :

B : __________ point (at which all the three phases co-exist in

equilibrium)

C : __________ point (above which the vapour cannot be liquefied

no matter how high the applied pressure will be)

Class Work

Determine each of the following by referring to the above phase diagram of a substance,X:

(a) the m.p. and b.p. ofX at atmospheric pressure ;

(b) the state ofXat (I) 200C and 2 atm; (ii) 0C and 1 atm ;

(c) the condition under whichXwill sublime.

va

our

ressure

temperature

external

ressure

temperature


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Bonding 13

comparison of phase diagrams between CO2 and H2O

features : (i) triple pointfor CO2 is above 1 atm while that for water is below 1 atm

(ii) slope of solid-liquid curve for CO2 is +ve while that for water is -ve

consequences : (i) liquid CO2 is not stable under ordinary conditions

solid carbon dioxide sublimes rather than melts

(ii) an increase in P will lower m.p. for H2O

delay freezing or melting occurs earlier than before water freezes with expansion in volume or ice melts with contraction in volume

water is denser than ice (formation ofopen cage structure in ice)

Molecular Crystals

Molecular crystals consist of molecules held in simple cubic or face-centred cubic lattice by weak intermolecular forces such as

van der Waals' forces (e.g. iodine) or hydrogen bonds (e.g. ice) :

Molecular solids are :

1 volatile with low m.p. / b.p.

2 poor conductor of electricity

3 usually soft and of low densities

Buckministerfullerene, with formula C60, was discovered in 1985 as the third form of carbon

allotrope, which is made from interlocking hexagonal and pentagonal rings of C atoms. It is the

parent of a new family of structures called fullerenes, including C70, C82 and C100, mostly

produced from the sooty flames from burning benzene.

Fullerenes can conduct electricity due to the presence of delocalized electrons on their surface.

By housing various metal atoms (e.g. K), inside the hollow cage, a whole new range of

applications have been found, including catalysts, superconductors, rocket fuels, laser materials

and drugs.

Class Work

Deduce the number of carbon atoms in a unit cell of C60, if it is known to have a face-centred cubic structure.

P

T

P

T


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Bonding 14

METALLIC BONDING

The structure of a metal can be conveniently illustrated by a model in which a lattice of regularly packed cations is surrounded by

a sea of delocalized electrons.

1 The delocalized nature of electrons accounts for their high electrical and thermal conductivity, whereas the flexibility of

packing the cationic lattice accounts for their malleable (pressed easily) and ductile (pulled easily) behaviour

2 Strength of metallic bonding is governed by two factors :

(a) no. of valence electrons metallic bond strength increases with the number of outermost (valence) electrons m.p./b.p.increases across the period (i.e. Na < Mg < Al)

(b) metallic radii ionic radius increases down the group, resulting in a decrease in charge density and, therefore, a

reduction of electrostatic attraction between cations and the valence electrons m.p./b.p. decreases down the group (i.e.

Li > Na > K)

3 When a metal is melted, only a small portion of metallic bonds are broken, resulting in a slight distortion of the entire structure.

On the other hand, all metallic bonds are broken in boiling, making the melting points of metals are usually significantly

lower than their boiling points. (e.g. m.p. & b.p. for Na are 98 and 883C respectively)

Metallic Crystals

The problem of packing cation together in a lattice is not so very different from the problem of packing ping-pong balls into a box.

There are two different ways of packing spheres within a single layer :

close packing loose packing

1 In building up metal spheres in a 3-D close-packedstructure :

a In a close-packed layer, each atom is in contact with six others, resulting a hexagonal arrangement of atoms.

b In the second layer, atoms pack as closely as possible to those in the first layer by 'sitting'in the holes between atoms in

the first layer. (Notice that there are now two different kinds of holes, namely tetrahedral and octahedral, depending on

the shape formed by the neighbouring spheres surrounding the hole)

c The third layer of atoms can now be added in two quite distinct ways so that :

(i) each sphere in the layer is sitting on the tetrahedral holes formed by the previous two layers, resulting in an overall

pattern that repeats itself every alternate layer (denoted by abab... etc.)hexagonal close packing (h.c.p.)

(ii) each sphere in the layer is sitting on the octahedral holes formed by the previous two layers, resulting in an overall

pattern that repeats repeats itself every three layers (denoted by abcabc... etc.)cubic close packing (c.c.p.)

(the word cubic is derived from the fact that a c.c.p. pattern could be visualized as a face-centred cubic (f.c.c.)

structure from a different point of view)


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Bonding 15

2 In building up metal spheres in a three-dimensional open (not close-packed) manner, a second layer can be formed by placing

a layer of spheres in the holes formed by the first layer, and a third layer formed in a similar way. This produces a

body-centred cubic (b.c.c.) arrangement in which each sphere is in contact with 4 spheres in the layer above and 4 in the

layer below, but in contact with no spheres within its own layer.

Different metals adopt different types of structures depending on their metallic radii.

h.c.p. c.c.p. / f.c.c. b.c.c.

arrangement for atoms

in different layersABAB ABCABC

unit cell

representation

co-ordination number

packing efficiency closely packed loosely packed

Alloys

An alloy is a mixture of two or more metals (or of a metal with a non-metal) such that the resulting properties are generally moredesirable than the pure metals alone. While some alloys (e.g. Pb-Sn alloy solder) exist in heterogeneous mixtures with separate

solid phases, most alloys are homogeneous solid solutions and can be further classified into:

1 substitutional alloy: some of the metal atoms being substituted by another metal (e.g. 18-carat gold alloy with silver); atomic

radii of the metals in the alloy should match with each other

2 interstitial alloy: the smaller atoms occupy holes among lattice formed by larger atoms (e.g. 1% carbon among iron in steel)

Compared with their constituents, the properties of alloys can be modified in the following ways:

Alloys are harder and stronger than pure metals. In the case of pure atoms, layers of atoms with identical radius can easily

slide over each other. However, the addition of a small amount of one metal with different atomic radii disrupts the orderly

arrangement of atoms, making it more difficult to slide over each other. (Degree of hardness increase as the relative

motion between particles become more restricted.) some alloys are more resistant to corrosion than pure metals (e.g. stainless steel, Al-Mg alloy duralumin)

many alloys form mixtures with much lower melting points than pure metals (e.g. solder to join metals together)

chemistry as level

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