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Bonding and Structure
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CHEMICAL BONDING
Stable atoms
The noble gases exist as single atoms that are very stable. All electrons are paired
and the bonding shells are full.
helium
neon
argon
Bonding
Compounds form when atoms are joined together by forces called chemical bonds.
When bonds form, unpaired electrons often pair up to form a noble gas electron
structure. This is often referred to as the ‘octet rule’.
 the bonding shells are full
 the electron structure is very stable.
Types of bonding
Bonds are classified into two main types: ionic and covalent.
As a general rule,
 Ionic bonding occurs between a metal and a non-metal.
 Covalent bonding occurs between a non-metal and a non-metal.
Questions
Predict the type of bonding in the following compounds.
(a)
sodium chloride
………………….
(d)
silver bromide
(b)
zinc oxide
………………….
(e)
nitrogen bromide ………………….
(c)
hydrogen chloride ………………….
(f)
sulphur dioxide
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………………….
………………….
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IONIC BONDS
Ionic bonds are present in a compound of a metal and a non-metal.
 Electrons are transferred from a metal atom to a non-metal atom forming ions.
 The metal ion is positive
 The non-metal ion is negative
An ionic bond is the electrical attraction
between oppositely charged ions.
Sodium chloride, NaCl
Na atom
Cl atom
Na+ ion
Cl ion
noble gas
neon
argon
The ions that are formed often have stable noble gas electron structures with full outer
electron shells.
Magnesium chloride, MgCl2
Mg atom
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Cl atoms
Mg2+ ion
Cl ions
noble gas
neon
argon
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Problems
(a)
Draw ‘dot-and-cross’ diagrams to show how atoms can form the four ionic
compounds below.
(b)
Write down the noble gas that would have the same electron structure as each
ion that you draw.
(i)
magnesium oxide,
….. atom
(ii)
….. ion
noble gas
………
………..
….. atom
….. ion
….. ion
noble gas
………
………..
….. atom
….. ion
….. ion
noble gas
………
………..
aluminium oxide.
….. atom
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….. ion
sodium oxide,
….. atom
(iii)
….. atom
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IONS AND THE PERIODIC TABLE
The charge on an ion can easily be predicted from its position in the Periodic Table.
Complete the table below.
Group
1
2
3
4
5
6
7
8
number of
outer shell
electrons
1
2
3
4
5
6
7
8
element
Li
Be
B
C
N
O
F
Ne
2,1
2,2
2,3
2,4
2,5
2,6
2,7
2,8
ion
element
Li+
F
2
2,8
Na
Mg
Al
Si
P
S
Cl
Ar
2,8,1
2,8,2
2,8,3
2,8,4
2,8,5
2,8,6
2,8,7
2,8,8
ion
Multiple charges
Sometimes elements can form ions with different charges. The charge is shown by a
Roman numeral.
 iron(II) for Fe2+ and iron(III) for Fe3+;
 copper(I) for Cu+ and copper(II) for Cu2+.
Predicting ionic formulae
calcium chloride:
equalise charge
formula
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ion
charge
ion
charge
Ca2+
2+
Cl
1–
Ca2+
2+
2 Cl
2–
CaCl2
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Predicting formulae
1.
Predict the formula of the ionic compounds below.
(a)
lithium chloride
………………….
(e)
aluminium fluoride ………………….
(b)
sodium oxide
………………….
(f)
aluminium sulphide………………….
(c)
beryllium oxide
………………….
(g)
aluminium phosphide……………….
(d)
lithium nitride
………………….
(h)
magnesium nitride………………….
2.
Predict the formula of the ionic compounds below.
(a)
copper(II) oxide
………………….
(e)
cobalt(II) oxide
(b)
copper(I) oxide
………………….
(f)
chromium(III) oxide………………..
(c)
iron(II) fluoride
………………….
(g)
manganese(VI) oxide……………….
(d)
iron(III) chloride
………………….
(h)
vanadium(V) nitride………………….
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……..………….
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Molecular ions
common molecular ions
ion
1+
1–
2–
OH–
ammonium NH4+ hydroxide
3–
carbonate CO32– phosphate PO43–
nitrate
NO3–
sulphate
SO42–
nitrite
NO2–
sulphite
SO32–
Ionic compounds from molecular ions
sodium carbonate
equalise charge
ion
charge
ion
charge
Na+
1+
CO32–
2–
2 Na+
2+
2 CO32–
2–
formula
calcium nitrate
equalise charge
Na2CO3
ion
charge
ion
charge
Ca2+
2+
NO3–
1–
Ca2+
2+
2 NO3–
2–
formula
Ca(NO3)2
Predicting more formulae
Predict the formula of the ionic compounds below.
………………….
(a)
lithium nitrate
………………….
(e)
iron(III) sulphite
(b)
sodium carbonate ………………….
(f)
chromium(III) nitrite ……………….
(c)
aluminium sulphate…………………. (g)
ammonium phosphate …………….
(d)
calcium hydroxide ………………….
vanadium(V) sulphate……………….
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(h)
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COVALENT BONDS
A covalent bond is formed between atoms of non-metals with a similar attraction for
electrons.
A covalent bond is a shared pair of
electrons.
A covalent bond forms when two atoms attract the same pair of electrons.
H2 molecule
Cl2 molecule
Molecules with different atoms
(a)
Draw ‘dot-and-cross’ diagrams to show a molecule of the following compounds
of hydrogen.
(b)
Write down the formula of each compound.
(c)
What is the common name given to each compound?
hydrogen chloride
hydrogen oxide
formula:
………….
formula:
………….
common name:
…………………..
common name:
…………………..
hydrogen nitride
hydrogen carbide
formula:
………….
formula:
………….
common name:
…………………..
common name:
…………………..
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Multiple bonds
Double and triple bonds are possible.
O2
N2
Drawing ‘dot-and-cross’ diagrams
1
2
Draw dot-and-cross diagrams of the following molecules of (i); (ii); (iii); (iv),
F2
HF
SiF4
SCl2
Draw dot-and-cross diagrams of the following molecules, each containing at
least one multiple covalent bond.
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CO2
C2H4
HCN
H2CO
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DATIVE COVALENT BONDS
A dative covalent bond forms when the shared pair of electrons comes from just one
of the atoms.
The formation of the ammonium ion, NH4+.
NH3
H+
NH4+
More covalent questions
1.
Draw dot-and-cross diagrams to show the formation of a dative covalent bond in
the oxonium ion, H3O+.
2.
(a)
Draw a dot-and-cross diagram for a molecule of boron trifluoride.
(b)
How many electrons surround the boron atom?
...............................................................................................................................
(c)
In what way, is the electron structure different to those met previously?
...............................................................................................................................
...............................................................................................................................
...............................................................................................................................
(d)
Why do you think this structure is still stable?
...............................................................................................................................
...............................................................................................................................
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3.
When covalent bonds form, unpaired electrons often pair up to form a noble gas
electron structure. This is often referred to as the ‘octet rule’. Many molecules,
however, form electron structures that do not form an octet.
Draw ‘dot-and-cross’ diagrams for the following molecules which do not obey
the octet rule.
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AlCl3
PCl5
SF6
SO3
SO2
SO3
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SHAPES OF MOLECULES
the shape of a molecule depends upon the number of electron pairs surrounding the
central atom.
The shapes of simple molecules can be explained using electron-pair repulsion.
 The pairs of electrons that surround an atom repel one another.
 When this happens the electron pairs become as far apart as possible.
The electron-pairs naturally repel into the following shapes around a central atom.
molecule
number of
electron
pairs
around
central
atom
dot-and-cross
diagram
shape
bond
angle
BCl2
BF3
CH4
SF6
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Molecules with lone pairs
A lone pair of electrons repels more than a bonded pair. This is because a lone pair is
closer to the atom:
In general:
 lone-pair/lone-pair repulsion
>
 bonded-pair/lone-pair repulsion >
bonded-pair/lone-pair repulsion
bonded-pair/bonded-pair repulsion
Lone pairs will distort the shape of a molecule and reduce the bond angle:
molecule
number of dot and cross diagram
lone pairs
shape
bond
angle
CH4
NH3
H2O
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Some shape questions
For each of the following molecules, draw a dot-and-cross diagram, and predict the
shape and bond-angles.
molecule
number of dot and cross diagram
lone pairs
shape
bond
angle
BeF2
AlCl3
SiH4
H2S
PH3
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Double bonds
A double bond is treated in the same way as a bonded pair. Each double bond and
pair of electrons should be treated as a bonding ‘area’.
molecule
number of
bonding areas
‘dot and cross’ diagram
shape
bond
angle
CO2
Now try these…
molecule
number of
bonding
areas
‘dot and cross’
diagram
shape
bond
angle
CS2
C2H4
SO3
SO2
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IONIC OR COVALENT?
An ionic bond with 100% ionic character would require the complete transfer of an
electron from a metal atom to a non-metal atom.
 In practice, this never completely happens!
 Between the extremes of ionic and covalent bonding, there is a whole range of
intermediate bonds, which have both ionic and covalent contributions.
Ionic bonds with covalent character
Many ionic compounds have a degree of covalency resulting from incomplete transfer
of electrons.
This occurs in ionic compounds with
 a small + ion (cation) with a high charge density and
 a large – ion (anion) with a low charge density.
charge density = Ошибка!
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2
+
3
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C
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in
c
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a
s
e
s
 The electric field around a small cation distorts the electron shells around a
large anion.
 This effect is called polarisation.
 As a result, electrons are attracted towards the cation, giving a degree of
covalency (electron sharing).
In Al2O3, polarisation takes place between Al3+ and O2– ions.
2
-
3
+
A
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S
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ELECTRONEGATIVITY
The nuclei of the atoms in a molecule attract the electrons pair in a covalent bond.
It is this attraction which is responsible for the covalent bond.
I
n
a
h
y
d
r
o
g
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n
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t
b
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n
d
.
-+
+
H
H
Electronegativity is a measure of the attraction of an atom in
a molecule for the pair of electrons in a covalent bond.
The most electronegative atoms attract bonding electrons most strongly.
 Elements with small atoms have the most electronegative atoms.
 Highly reactive non-metallic elements (such as O, F and Cl) have the most
electronegative atoms.
 Reactive metals (such as Na and K) have the least electronegative atoms.
How is electronegativity measured?
electronegativity increases
Li
1.0
Na
0.9
K
0.8
Be
1.5
B
2.0
C
2.5
N
3.0
O
3.5
F
4.0
Cl
3.0
Br
2.8
Fluorine is the most electronegative element with an electronegativity of 4.0;
Fluorine has small atoms which attract the pair of electrons in a covalent bond more
strongly than larger atoms.
The greater the difference between electronegativities, the greater the ionic
character of the bond.
The greater the similarity in electronegativities, the greater the covalent character of
the bond
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POLAR AND NON-POLAR MOLECULES
Non-polar bonds
A covalent bond is non-polar when
 the bonded electrons are shared equally between both atoms
 the bonded atoms are the same
 the bonded atoms have similar electronegativities.
A covalent bond must be non-polar if the bonded atoms are the same.
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HH
Cl  Cl
H2 molecule
Cl2 molecule
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Polar bonds
A covalent bond is polar when
 the electrons in the bond are shared unequally making a polar bond;
 the bonded atoms are different, each has a different electronegativity.
In hydrogen chloride the H–Cl bond is mainly covalent with one pair of electrons
shared between the two atoms.
 However, the chlorine atom is more electronegative than the hydrogen atom.
 The chlorine atom attracts the bonded pair of electrons more than the hydrogen
atom making the covalent bond polar.
 The electrons in the bond are shared unequally making a polar bond.
+ 
H  Cl
HCl molecule
 The HCl molecule is polarised with a small positive charge + on the H atom
and a small negative charge – on the Cl atom.
 The hydrogen chloride molecule is polar with a permanent dipole.
Symmetrical and unsymmetrical molecules
In symmetrical molecules, dipoles cancel and there are no permanent dipoles.
dipoles cancel
CHCl3 - polar
CCl4 - non-polar
In CCl4
 each C–Cl bond is polar but the dipoles act in different directions
 the overall effect is for the dipoles to cancel each other
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 CCl4 is a non-polar molecule.
Some polar questions
1.
2.
Draw simple diagrams for each of the following molecules. Decide whether each
molecule is polar. If so, use the periodic table to predict the dipole present and
shows these clearly as + and  on your diagrams.
(i)
Br2
H2O
O2
HBr
NH3
CF4
Predict the shape of a molecule of BF3 and of PF3.
BF3
(ii)
PF3
Explain why BF3 is non-polar whereas PF3 is polar.
...............................................................................................................................
...............................................................................................................................
...............................................................................................................................
...............................................................................................................................
3.
(i)
Predict the shape of a molecule of H2O and of CO2.
H2O
(ii)
CO2
Explain why H2O is polar whereas CO2 is non-polar.
...............................................................................................................................
...............................................................................................................................
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...............................................................................................................................
...............................................................................................................................
INTERMOLECULAR FORCES
Strength of bonds and forces
 Ionic, covalent and metallic bonds are of comparable strength.
 Intermolecular forces are much weaker:
type of bond
covalent bond
hydrogen bond
van der Waals’
forces
bond enthalpy/ kJ mol1
200-500
5-40
2
VAN DER WAALS' FORCES
Van der Waals’ forces (induced dipole-dipole interactions) exist between all molecules
whether polar or non-polar. Without these forces, non-polar molecules could never
form a liquid or a solid.
Van der Waals’ forces are
 weak intermolecular interactions
 caused by attractions between very small dipoles in molecules.
 movement of electrons
produces an oscillating
dipole
-
+
+
-
dipole oscillates and continually
changes with time
-
- +
- +
 oscillating dipole induces
- +
- +
- +
a dipole in a
neighbouring molecule
which is induced onto
further molecules
induced dipoles attract one another
+
What affects the strength of van der Waals’ forces?
Van der Waals' forces result from interactions of electrons between molecules.
The greater the number of electrons in each molecule
 the larger the oscillating and induced dipoles
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 the greater the attractive forces between molecules
 the greater the van der Waals' forces.
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Van der Waals’ forces and electrons
Molecules with more electrons will generate larger oscillating and induced dipoles.
These produce larger attractive forces between molecules.
The table below shows the boiling points of the hydrogen compounds of group 4.
compound
CH4
SiH4
GeH4
SnH4
boiling point
/K
112
161
178
221
number of
electrons
1.
Complete the table above with the number of electrons in each compound.
2.
Explain why the boiling point increases down this group.
...............................................................................................................................
...............................................................................................................................
...............................................................................................................................
...............................................................................................................................
3.
Compare these boiling points with those for the hydrogen compounds of Group
6 (page 4). Explain the similarities and differences between the two sets of data.
...............................................................................................................................
...............................................................................................................................
...............................................................................................................................
...............................................................................................................................
...............................................................................................................................
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...............................................................................................................................
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...............................................................................................................................
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PERMANENT DIPOLE-DIPOLE INTERACTIONS
Permanent dipole-dipole interaction are weak intermolecular forces.
 They are simply attractions between weak dipole charges on different molecules.
Permanent dipole-dipole interactions are ‘non-directional’.
Intermolecular forces between HCl molecules

H
Cl
weak dipole-dipole interactions

between HCl molecules
H

H
Cl

H
Cl
Cl

H
Cl
Between HCl molecules, there will be both
 weak intermolecular forces:
 van der Waals’ forces and
 permanent dipole-dipole interactions.
permanent dipole-dipole interactions > van der Waals’ forces.
Now do this…..
Draw a diagram to show the dipole-dipole interactions that exist between molecules of
nitrogen monoxide, NO.
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HYDROGEN BONDS
A hydrogen bond is a special type of permanent dipole-dipole interaction found
between molecules containing the following groups:

H O

H N

H
F
-
Hydrogen bonds act between:
 a lone pair HO:, HN: or HF: and
 a hydrogen atom HO, HN or HF on a different molecule.
Hydrogen bonding occurs between molecules such as H2O.
Hydrogen bonding is important in organic compounds containing –OH or –NH bonds:
e.g. alcohols, carboxylic acids, amines, amino acids.
When water change state, the covalent bonds between the H and O atoms in an H2O
molecule are strong and do not break - the much weaker intermolecular forces break
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Some hydrogen bonding questions
1.
Put a tick below the following molecules that have hydrogen bonding:
molecule
hydrogen bonding?

2.
H2O
H2S
CH4
HOCH3
NH3
NO2
H2NCH3
Each molecule that you draw should show relevant lone pairs and dipoles.
Draw diagrams showing hydrogen bonding between
(a)
2 molecules of ammonia;
(b)
2 molecules of hydrogen fluoride;
(c)
2 molecules of ethanol, C2H5OH;
(d)
1 molecule of water and 1 molecule of ethanol.
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SPECIAL PROPERTIES OF WATER
A hydrogen bond has only about one-tenth the strength of a covalent bond.
However, hydrogen bonding is strong enough to have significant effects on physical
properties, resulting in some unexpected properties for water.
The solid (ice) is less dense than the liquid (water)
 Particles in solids are usually packed closer together than in liquids.
 Hydrogen bonds hold water molecules apart in an open lattice structure.

ice is less dense than water.
The diagram below shows how the open lattice of ice collapses on melting.
hydrogen bonds break
MELTING OF ICE
ice lattice collapses:
molecules move closer together
tetrahedral open lattice in ice
H2O has a relatively high melting point and boiling point
 There are relatively strong hydrogen bonds between H2O molecules.
 The hydrogen bonds are extra forces, over and above van der Waals’ forces.
 These extra forces result in higher melting and boiling points than would be
expected from just van der Waals’ forces.
 When the ice lattice breaks, hydrogen bonds are broken.
Other properties
The extra intermolecular bonding from hydrogen bonds also explains the relatively
high surface tension and viscosity in water.
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How do hydrogen bonds affect the properties of water?
Hydrogen bonds increase the attraction between molecules. This will result in an
increase in the boiling point of a compound. The table below shows the boiling points
of the hydrogen compounds of group 6.
H2O
boiling
point /K
373
H2S
213
H2Se
231
H2Te
270
compound
relative
molecular mass
1.
Complete the table above with the relative molecular mass of each compound.
2.
Plot a graph of boiling point against relative molecular mass. Join each point
with a line.
3.
The boiling point of water is higher than expected owing to hydrogen bonding.
Use your graph to estimate what the boiling point of water would be if there were
no hydrogen bonding. Show this clearly on your graph.
4.
Refer to Chapter 3 of the Foundation Chemistry textbook. Water has several
peculiar properties that can be explained by hydrogen bonds. The unexpected
high boiling and melting point is just one of these. Explain how hydrogen
bonding is responsible for three more unusual properties of water.
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BONDING, STRUCTURE AND PROPERTIES
The properties of a substance depend upon its bonding and structure.
 Ionic and covalent bonds are of comparable strength
- Intermolecular forces are far weaker.
 Ionic bonds break when a giant ionic lattice melts or boils
 Simple molecular structure – weak forces break - low melting point
 Covalent bonds break when a giant molecular lattice melts or boils
 Giant structure – strong forces break - high melting point
GIANT IONIC LATTICES
Ionic compounds form giant ionic lattices with each ion surrounded by ions of the
opposite charge.
A giant ionic lattice is held together by strong electrostatic attraction
between positive and negative ions.
Each ion is surrounded by oppositely-charged ions, forming a giant ionic lattice.
Part of the sodium chloride lattice,
Na+
Cl
-
 Each Na+ ion surrounds 6 Cl ions
 Each Cl ion surrounds 6 Na+ ions
 Although it is convenient to look at ionic bonding between two ions only, each ion is
able to attract oppositely charged ions in all directions.
 This results in a giant ionic lattice structure with hundreds of thousands of ions
(depending upon the size of the crystal).
 This arrangement is characteristic of all ionic compounds.
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Properties of Giant ionic lattices
High melting point and boiling point
Ionic compounds are solids at room temperature.
 The attraction is strong between + and  ions.
 High temperatures are needed to break the strong electrostatic forces holding the
ions rigidly in the solid lattice.
 ionic compounds have high melting and boiling points.
Electrical conductivity
In the solid lattice,
 the ions are in a fixed position and there are no mobile charge carriers.
 an ionic compound is a non-conductor of electricity in the solid state.
When melted or dissolved in water,
 the solid lattice breaks down,
 the ions are now free to move as mobile charge carriers.
 an ionic compound is a conductor of electricity in liquid and aqueous states.
Solubility
 The ionic lattice dissolves in polar solvents (e.g. water).
 The polar water molecules break down the lattice and surround each ion in solution
as shown below for sodium chloride.
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PROPERTIES OF COVALENT COMPOUNDS
Elements and compounds with covalent bonds have either of two structures:
 a simple molecular structure,
 a giant molecular structure.
SIMPLE MOLECULAR STRUCTURES
Simple molecular structures have small molecules, such as Ne, H2, O2, N2.
 A simple molecular structure is held together by weak forces between molecules
 The atoms within each molecule are bonded strongly together by covalent bonds
Properties of simple molecular structures
Low melting point and boiling point
 Low temperatures provide sufficient energy to break the intermolecular forces.
 simple molecular structures have low melting and boiling points.
Strong covalent bonds hold
together each I2 molecule
Weak van der Waals' forces
between I2 molecules
When the simple molecular structure of I2 is broken,
 only the weak van der Waals’ forces between the I2 molecules break;
 the covalent bonds, I–I, are strong and do not break.
Electrical conductivity
 There are no free charged particles.
 simple molecular structures are non-conductors of electricity.
Solubility
 Van der Waals' forces form between a simple molecular structure and a non-polar
solvent, such as hexane.
 These interactions weaken the structure.
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 simple molecular structures are soluble in non-polar solvents (e.g. hexane).
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GIANT MOLECULAR STRUCTURES
Diamond, graphite and SiO2 are common examples of giant molecular lattices.
This type of structure is known by a variety of names: a giant molecular lattice, a giant
covalent lattice, a giant atomic lattice and a macromolecular lattice.
 A giant molecular structure is held together by strong covalent bonds between
atoms.
Properties of giant molecular structures
High melting point and boiling point
 High temperatures are needed to break the strong covalent bonds in the lattice.
 giant molecular structures have high melting and boiling points.
Electrical conductivity
 Except for graphite (see below), there are no free charged particles.
 giant
molecular structures are non-conductors of electricity.
Solubility
 The strong covalent bonds in the lattice are too strong to be broken by either polar
or non-polar solvents.
 giant
molecular structures are insoluble in polar and non-polar solvents.
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PROPERTIES OF DIAMOND AND GRAPHITE
diamond
property
graphite
tetrahedral
hexagonal layers
 symmetrical structure held
together by strong covalent
bonds throughout lattice
 strong layer structure but with
weak bonds between the layers
poor conductivity
good conductivity
 There are no delocalised
electrons
 All outer shell electrons are used
for covalent bonds.
hard
 Delocalised electrons between
layers.
 Electrons are free to move
parallel to the layers when a
voltage is applied.
soft
 Tetrahedral shape enables
external forces to be spread
throughout the lattice.
 Strong bonding within each layer
 weak forces between layers
easily allow layers to slide.
structure
electrical
conductivity
hardness
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BONDING AND STRUCTURE IN METALS
Metallic bonding
A metallic bond holds atoms together in a solid metal or alloy.
In solid metals, the atoms are ionised.
 The positive ions occupy fixed positions in a lattice;
 The outer shell electrons are delocalised
– they are spread throughout the metallic structure and are able to move freely
throughout the lattice.
A metallic bond is the electrostatic attraction between the
positive metal ions and delocalised electrons.
 In the metallic lattice, each metal atom exists as + ions by releasing its outer shell
electrons to the ‘sea of electrons’.
+ - + - +
- - + - + - + - +
- - + - + - +
The ‘sea of electrons’
Delocalised and localised electrons
In a metallic bond,
 the delocalised electrons in metals are spread throughout the metal structure;
 the delocalised electrons are able to move throughout the structure;
 it is impossible to assign any electron to a particular positive ion.
In a covalent bond,
 the localised pair of electrons is always positioned between the two atoms
involved in the bond;
 the electron charge is concentrated between the bonded atoms.
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PROPERTIES OF GIANT METALLIC LATTICES
All metals form giant metallic lattices in the solid state
 A giant metallic lattice is held together by strong electrostatic attractions between
positive ions and negative electrons.
 A metallic bond acts between the positive metal ions and delocalised electrons.
High melting point and boiling point
 Generally high temperatures are needed to separate the ions from their rigid
positions within the lattice.
 most metal have high melting and boiling points.
Good thermal and electrical conductivity
 The existence of mobile, delocalised electrons allows metals to conduct heat and
electricity well, even is the solid state.
-
+
+
-
+
+
- + - +
+
- + - +
+
+
+
+
+
+
drift of delocalised electrons
across potential difference
 The electrons are free to flow between positive ions.
 The positive ions do not move.
 When a metal conducts electricity, only the electrons move.
Solubility
 Metals are insoluble.
 Solvents such as water are unable to form strong enough forces with the ions and
electrons to pull the lattice apart.
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COMPARISON OF STRUCTURE, BONDING AND
PROPERTIES
The different types of bonds and forces
 Covalent bonds act between atoms.
 Ionic bonds act between ions.
 Metallic bonds act between positive ions and electrons.
 Hydrogen bonds act between polar molecules.
 Dipole-dipole interactions act between polar molecules.
 Van der Waals' forces act between induced dipoles of molecules.
structure
m pt/
b pt
reason
electrical
conductivity
reason
solubility
reason
giant
ionic
simple
molecular
giant
molecular
hydrogen
bonded
giant
metallic
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Properties questions
1.
Magnesium oxide forms a similar ionic lattice to that of sodium chloride.
(a)
Draw a clear diagram of the magnesium oxide lattice.
(b)
Predict, with reasons, the following properties of magnesium oxide:
(i)
melting and boiling points
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(ii)
electrical conductivity
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(iii) solubility.
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(c)
Suggest, with reasons, why magnesium oxide has a higher melting point
than sodium chloride,
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2.
(a)
Describe what is meant by hydrogen bonding,
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(b)
Describe how hydrogen bonding influences the properties of water.
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3.
The data below gives some properties of four substances, A, B, C and D.
compound
solubility in
water
solubility in
hexane
A
B
C
D
good
poor
poor
poor
insoluble
good
poor
poor
(a)
electrical
conductivity
aqueous
solid
solution
poor
good
poor
poor
poor
good
-
boiling
point /K
1738
456
2503
3160
Explain how these data suggest different structures for the four
substances shown.
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(b)
Suggest an identity for each substance.
A .................................................................................................................
B .................................................................................................................
C .................................................................................................................
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D .................................................................................................................
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