10
1
Metallic Bonding
10.1
Metallic Bonding
10.2
Metallic Radius
10.3
Factors Affecting the Strength of
Metallic Bond
10.4
Metallic Crystals
10.5
Alloys
Nature of
Metallic Bonding
2
Not ionic : ∵ atoms of the same electronegativity
Li(g) + Li(g)
Li+ Li(g)
(1)
No favourable
Li(g) + Cl(g)
favourable
3
Li+ Cl(g)
(2)
Li(g) + Li(g)
Li+ Li(g)
(1)
Li(g) + Cl(g)
Li+ Cl(g)
(2)
Q.57 Explain, with the aid of suitable enthalpy
change cycles, why reaction (1) is
energetically less favourable than reaction (2).
(1st EA : Li = 60 kJ mol1, Cl = 349 kJ mol1)
4
Li(g) + Li(g)
H1
1st E.A. of Li
1st I.E.
of Li
+
Li+ Li(g)
H2
Li (g) + Li(g)
H1 = 1st I.E. of Li + (60 kJ mol1) + H2
Li(g) + Cl(g)
H3
1st E.A. of Cl
1st I.E.
of Li
+
Li+ Cl(g)
H4
Li (g) + Cl(g)
H3 = 1st I.E. of Li + (349 kJ mol1) + H4
5
p
e
p
e

of Li 

of Cl 
ΔH 2  
ΔH 4  
6
3
4
17
18
 0.75
 Cl is smaller than Li
 0.75
1
rLi   rLi 
1
rLi   rCl 
 H4 is more negative than H2
H4 is more negative than H2
H1 = 1st I.E. of Li + (60 kJ mol1) + H2
H3 = 1st I.E. of Li + (349 kJ mol1) + H4
H3 is more negative than H1
Li(g) + Li(g)
H1
Li+ Li(g)
(1)
No favourable
Li(g) + Cl(g)
H3
favourable
7
Li+ Cl(g)
(2)
Not covalent : ∵ efficiency of orbital overlap  as the
bonding atoms get larger
8
Molecule
Bond Enthalpy
(kJ mol1)
H2
Li2
Na2
K2
Rb2
Cs2
435
105
70
50
46
43
Explanation
1. Classical approach
Distance between shared pair and the bonding nuclei :
H2 < Li2 < Na2 < K2 < Rb2 < Cs2
Bond strength :
H–H > Li–Li > Na–Na > K–K > Rb–Rb > Cs–Cs
9
2. MO approach (Not required in AL)
1s*
LiA
2sA 
1s
10
Li2

LiB

2sB
Q.58
There is no gain of stability when the AOs of two
Beryllium atoms overlap.
2 e involved in bonding
Be(A)
Be2
*
2s
1s*
2sA
Overall : No e involved in bonding
2sB
2s
1s
11
Be(B)
2 e involved in antibonding
Bond order = 0
Conclusion : 1. Metals tend to form giant structures rather
than discrete molecules.
2Li  Li – Li (low bond enthalpy)
nLi  Lin
(high bond enthalpy)
Stronger bonds are formed due to extensive
delocalization of valence electrons.
12
Conclusion : 2. The electron-sea model
The valence electrons do not belong to any
specific atoms (not localized) but delocalize
throughout the whole crystal structure.
13
Conclusion : 2. The electron-sea model
Mobile es  electron sea
Stationary +ve ions
14
Conclusion : 2. The electron-sea model
The electrostatic attractive forces between the
delocalized electron cloud and the positive ions
are called the metallic bonds
15
Since metallic bonds are non-directional, they
exist in significant extent even in molten state.
The boiling points of metals are much higher
than the corresponding melting points.
E.g. Na m.p.  97.8oC ; b.p.  903.8oC
NaCl m.p. = 801C ; b.p. = 1413C
16
Conclusion : 3. MO approach : Band theory
Spacing  as the no. of molecular orbitals 
Li
Li2
Li3
17
Li4
Lin
Lin : n orbitals overlap  continuous band
Half-filled
band
18
Metallic
Radius
19
10.2 Metallic radius (SB p.262)
Metallic radius (r) is defined as half of the
internuclear distance between adjacent
atoms in a metal crystal.
20
10.2 Metallic radius (SB p.262)
Trend of metallic radius in the
Periodic Table
• Moving down a group, metallic radii increase
• Going across a period, metallic radii decrease
21
Factors
Affecting the
Strength of
Metallic Bond
22
The strength of metallic bond can be estimated by
melting point,
boiling point,
enthalpy change of fusion or
enthalpy change of vapourization.
Higher m.p./b.p./Hfusion/Hvap
 stronger metallic bond
23
10.3 Factors affecting the strength of metallic bond (SB p.262)
The metallic bond strength increases with:
1. decreasing size of the metal atom
(i.e. the metallic radius);
2. increasing number of valence electrons
of the metal atom.
24
10.3 Factors affecting the strength of metallic bond (SB p.263)
Effect of number of valence
electrons on metallic bond strength
Metal
Number of valence
electrons(s)
Sodium
1
Magnesium
2
Aluminium
3
25
Melting
point (oC)
98
650
660
10.3 Factors affecting the strength of metallic bond (SB p.263)
Effect of metallic radius on metallic
bond strength of Group IA metals
Metal
Metallic radius
(mm)
Lithium
0.152
Sodium
0.186
Potassium
0.231
Rubidium
0.244
Caesium
0.262
26
Melting point
(oC)
180
98
64
39
29
Typical properties of metals
1. High density
due to close packing of atoms in metallic
crystal (h.c.p./f.c.c. co-ordination number  12)
27
Metal
Ni
Density
(g cm3)
8.91
Cu
Ag
Pb
Hg
Au
8.94 10.49 10.66 13.53 19.30
Typical properties of metals
1. High density
Exception : Alkali metals have low densities
(< 1 for Li, Na and K )
(a) they have more open structures
(b.c.c. /co-ordination number  8)
(b) their atomic radii are the highest in their
own Periods.
E.g. Size : Na > Mg > Al
28
Typical properties of metals
2. High melting point and boiling point
Extensive delocalization of valence electrons
 stronger bonds
Bond strength : ionic bond  covalent bond  metallic bond
29
Typical properties of metals
3. High flexibility
Malleability :
The ability to be deformed under compression
Ductility :
The ability to be deformed under tension
30
Typical properties of metals
3. High flexibility
Reasons : (a)
The presence of layers in the crystal lattice
i.e. the layers can slide over one another
under strain
(b)
Metallic bonds are non-directional.
i.e. electrons can take up new positions and
reform metallic bond after the deformation
31
Typical properties of metals
32
Typical properties of metals
4. Surface lustre  Silvery and shiny
Lin
Half-filled
2s band
Since the gap between energy levels are
extremely small, radiation of any frequency in
visible region can be absorbed and emitted.
33
Typical properties of metals
Cu : 3d10, 4s1
Reddish brown
Half-filled
s band
E = 220 kJ mol1
full-filled
d band
34
Typical properties of metals
Au : 5d10, 6s1
Golden yellow
Half-filled
s band
E = 300 kJ mol1
full-filled
d band
35
Typical properties of metals
Ag : 4d10, 5s1
UV light
absorbed
Silvery
Half-filled
s band
E = 380 kJ mol1
full-filled
d band
36
Typical properties of metals
5. High Thermal and Electrical Conductivity
Due to the free movement of delocalized
electrons
37
38
Metallic
Crystals
39
10.4 Metallic crystals (SB p.263)
Closed-packed structure
Closed-packed structure is made possible with
identical particles
Two types : 1. Hexagonal closed-packed, h.c.p.
2. Cubic closed-packed, c.c.p. Or
Face-centred cubic, f.c.c.
40
10.4 Metallic crystals (SB p.263)
Hexagonal close-packed structure
abab…
41
10.4 Metallic crystals (SB p.265)
Hexagonal close-packed structure
(a) normal side view
(b) exploded view
Packing efficiency = 74 %
42
(c) a unit cell
Co-ordination no. = ?
Co-ordination
number = 12
43
10.4 Metallic crystals (SB p.265)
Cubic close-packed / Face-centred cubic
structure
abcabc…
c.c.p. or f.c.c.
rotate by 45
Co-ordination no. = 12
44
Packing efficiency = 74 %
10.4 Metallic crystals (SB p.266)
Open structure
• Structures with more empty space between
the atoms
• Most common: body-centred cubic structure
45
10.4 Metallic crystals (SB p.267)
Body-centred cubic structure
(a) normal side view
(b) exploded view
Packing efficiency = 68 %
46
(c) a unit cell
Co-ordination no. = 8
Given : Density of Cu = 8.94 g cm3
Relative atomic mass of Cu = 63.546
Atomic radius of Cu = 0.128 nm
Cu adopts f.c.c. structure
Calculate the Avogadro’s constant
2
a
+
2
a
=
a  2 2r
47
2
(4r)
a  2 2r
Density

Volume
molar volume
molar mass
density

63.546 g mol
8.94 g cm
of unit cells per mole of Cu 

1
3
Molar volume of Cu
Volume
of a unit cell of Cu
63.546
48
3
molar mass
molar volume of Cu 
number
3
of a unit cell  a  (2 2 r)
cm
3
mol
1
8.94
7
(2 2  0.128  10 cm)
3
number
of Cu atoms per unit cell  8 
1
8
6
1
2
4
Since one unit cell contains 4 Cu atoms
No. of Cu atoms per mole of Cu  No. of unit cells  4
63.546

cm
3
mol
1
8.94
7
(2 2  0.128  10 cm)
= 5.99  1023 mol1
49
3
4
Alloys
50
10.5 Alloys (SB p.268)
Alloys
• Made by mixing a metal with one or more
other elements (metals or non-metals)
51
10.5 Alloys (SB p.268)
Structure and Properites of alloy
• Have structures and properties different
from that of a pure metal
• In a pure metal, all the atoms are of the
same size
52
10.5 Alloys (SB p.268)
Structure of alloy
• In an alloy, atoms of different sizes are
present
53
10.5 Alloys (SB p.268)
Structure of alloy
 Changes the regular arrangement of the
layers of atoms in the metal
 Slipping of layers of atoms becomes more
difficult
 Harder and stronger
54
10.5 Alloys (SB p.269)
Types of alloys
• 2 common types of alloys:
Substitutional alloy
Interstitial alloy
55
10.5 Alloys (SB p.269)
Substitutional alloy
• Some of the host metallic atoms are
replaced by other metallic atoms of
similar sizes
• e.g. in brass
56
10.5 Alloys (SB p.269)
Interstitial alloy
• Formed when some of the interstices
among the closely packed host metallic
atoms are occupied by atoms of smaller
atomic sizes
• e.g. in steel
57
10.5 Alloys (SB p.269)
Some common alloys - Steel
• An alloy of iron
• The presence of directional carbon-iron
bonds makes the resulting alloy harder,
stronger and less ductile than pure iron.
• Amount of carbon present affects the
properties of steel
• Mild steel: contains <0.2 % carbon, ductile,
malleable (used for nails, cables and chains)
58
10.5 Alloys (SB p.269)
Some common alloys - Steel
• Medium steel: contains 0.2 – 0.6 % carbon,
harder
used in rails and structural steel beams
• High-carbon steel: contains 0.6 – 1.5 %,
tough and hard
used for springs tools and cutlery
59
10.5 Alloys (SB p.269)
Some common alloys - Steel
Articles made from stainless steel
60
10.5 Alloys (SB p.269)
Some common alloys – Alloy Steel
• A mixed form of interstitial (carbon) and
substitutional (other metals) alloys
Example : Stainless steel (steel + Cr + Ni)
The presence of Cr and Ni greatly
increases the resistance to corrosion of
the alloy.
61
10.5 Alloys (SB p.269)
Some common alloys – Alloy Steel
Example : Tool steel (steel + W + Co)
It is very hard and has a very high m.p.
It is used for making high-speed cutting
tools
62
10.5 Alloys (SB p.270)
Some common alloys – Copper alloys
• Brass - an alloy of copper and zinc
• Attractive golden appearance
• Harder and more corrosion resistant than
copper and zinc.
• Used to make ornaments, buttons, musical
instruments, plugs and sockets, and water
taps.
63
10.5 Alloys (SB p.270)
Some common alloys – Copper alloys
• Brass - an alloy of copper and zinc
Article made from brass
64
10.5 Alloys (SB p.270)
Some common alloys – Copper alloys
• Coinage metals
65
10.5 Alloys (SB p.270)
Some common alloys – Copper alloys
• Silver coins = cupronickel (Cu + Ni)
66
10.5 Alloys (SB p.270)
Some common alloys – Copper alloys
• Copper coins (copper + tin + zinc)
67
10.5 Alloys (SB p.270)
Some common alloys – Copper alloys
Brass (Cu + Zinc)
68
10.5 Alloys (SB p.269)
Some common alloys – Duralumin
• An alloy of aluminium with Cu, Mg and Mn
It is light and is stronger and more
corrosion resistant than aluminium.
It is used for making spacecrafts and jet
fighters.
69
10.5 Alloys (SB p.270)
Some common alloys – Solder
• An alloy of lead and tin
It has a lower m.p.(about 180C) than
that of lead and tin.
It is used in joining metals together.
It melts easily to fill the gaps between
metals without melting them.
On cooling, it solidifies and completes the
circuit.
70
10.5 Alloys (SB p.270)
Some common alloys – Solder
• An alloy of lead and tin
71
Check Point
10-5
10.5 Alloys (SB p.270)
Some common alloys – Carat Gold
• An alloy of gold with silver and copper.
• Pure gold is too soft to make jewellery.
• Carat gold is harder than pure gold
• Pure gold is called 24 carat (24K) gold
18 carat (18K) gold contains
18/24 or 75% gold.
72
The END
73
10.3 Factors affecting the strength of metallic bond (SB p.263)
Back
It is said that bonding in most metals is strong but nondirectional. Can you think of some facts to support the
above statement?
Answer
Metals are durable and have high melting (and boiling)
points. These indicate that metallic bonds are strong. On
the other hand, metals can be pulled into wires or
hammered into sheets (I.e. it is relatively easy to change
the shape of most metals). This shows that metal atoms
can slide over each other which is a consequence of the
non-directional nature of the metallic bond.
74
10.4 Metallic crystals (SB p.266)
Back
How many tetrahedral holes and octahedral holes are
there adjacent to each sphere in cubic close-packed
structure?
Answer
In cubic close-packed structure, there are 6
octahedral holes and 8 octahedral holes adjacent to
each sphere.
75
10.4 Metallic crystals (SB p.267)
X-ray crystallography shows that aluminium and
potassium have f.c.c. and b.c.c. structures respectively.
Calculate the number of atoms in a unit cell of
(a) aluminium; and
(b) potassium
Answer
76
10.4 Metallic crystals (SB p.267)
(a)
For the face-centred cubic structure of aluminium, an atom
1 so
on each face of the unit cell is shared by two cell and
2
of
the atom belongs to the unit1cell; an atom at each corner is
shared by eight cells and so 8
of the atom belongs to the
1
1
unit cell.
6
Number of aluminium atoms in a unit cell =
77
2
 8
=4
8
10.4 Metallic crystals (SB p.267)
Back
(b)
For the body-centred cubic structure of potassium, an atom
at the centre of the unit cell is not shared with other cells
and totally belongs to the unit cell; an atom at each corner is
1
shared by eight cells and
so
of the atom belongs to
8
1
the unit cell.
1 8 
Number of potassium atoms in a unit cell =
78
=2
8
10.4 Metallic crystals (SB p.268)
(a) X-ray crystallography shows that copper has the
cubic close-packed structure. Calculate the number
of atoms in a unit cell of copper.
Answer
79
10.4 Metallic crystals (SB p.268)
(a) For the cubic close-packed structure of copper, an atom on each
1
face of the unit cell is shared by two cells and
of the
2
atom belongs to the unit cell; an atom at each
corner is shared by
1 so
8 cells and
of the atom belongs to the unit cell.
8
Number of Cu atoms in a unit cell =6 
80
1
2
 8
=4
1
8
10.4 Metallic crystals (SB p.268)
(b) It is a known that sodium metal has a body-centred
cubic structure.
(i) Draw a unit cell of sodium.
(ii) Is this structure a close-packed structure?
Explain this in terms of the coordination number
of sodium.
Answer
81
10.4 Metallic crystals (SB p.268)
Back
(b) (i) A unit cell of sodium is drawn as follows:
(ii) Refer to the unit cell drawn in (b)(i), one atom is at each of
the
eight corners of a cube, and one atom is at the centre
touching
these eight atoms, so the coordination number of the
central
atom is 8. Thus, the structure is not a close-packed
structure.
82
10.5 Alloys (SB p.271)
(a) (i) Give two advantages of steel compared to the pure
iron.
(ii) Why is tungsten added to certain types of alloy
steels?
Answer
(a) (i) Steel is harder and stronger than iron. It is also less
ductile.
(ii) The addition of metal tungsten to certain types of alloy
steels make them become hard and strong with a very
high melting point. These materials are ideal for making
high-speed cutting tools.
83
10.5 Alloys (SB p.271)
(b) Cupronickel replaced earlier silver coins which
contained silver.
Give two reasons for the replacement.
Answer
(b) The main reason for the replacement was due to the
relatively high cost of silver, as the cost of making a pure
silver coin was higher than the value of the coin. Besides,
cupronickel is much harder and more durable than pure
silver.
84
10.5 Alloys (SB p.271)
Back
(c) (i) Why the low melting point of solder makes it
useful in joining metals together?
(ii) Explain how soldering joins up metals.
Answer
(c) (i) Due to the low melting point of solder, it needs not to
be heated up to a high temperature. As a result, there is
no risk for the metals to be joined to melt during soldering.
(ii) Solder is melted by an electrically heated rod. When it
melts, it flows over the two metal parts. When it cools,
it solidifies and joins the two metals together.
85