1. Unit 1.4 - Bonding
• The forces holding atoms
together in compounds are known as chemical bonds
• Bonds are broken and formed in chemical reactions
• Bonding involves electrons in the outer shells of atoms
• We will be studying ionic, covalent and metallic bonding
The octet rule
The first important thing to know about bonding is that atoms will generally react and form bonds in
such a way that they end up with an outer shell full of electrons. This principle is expressed in the
octet rule:
“When elements react they tend to do so in such a way that results in an _________ shell
containing __________ electrons”
Dot- cross diagrams
• Dot-cross diagrams are an easy way of keeping track of what happens to ____________
during chemical reactions
• Only the ___________ _________ electrons are shown as others are not involved in reacting
• We will look at some examples drawing dot-cross diagrams for the formation of ionic bonds:
Sodium + chlorine
Lithium + fluorine
Magnesium + chlorine
Magnesium + oxygen
Ionic bonding
• Ions are charged particles formed when
atoms __________ or __________ electrons
• Cations are ____________ ions

2. • Anions are _____________ ions
• Metals generally form ____________
• Non- metals generally form ___________
• Anions and cations are held together by strong electrostatic forces of ____________
otherwise known as ionic bonds
Ionic radii
• An atomic or ionic radius, measured in nanometers (nm, 10-9 m), is basically the size of an
atom or ion
• The following table shows atomic and ionic radii for some common atoms and ions:
Cations Anions
Na Na+ F F-
Mg Mg2+ O O2-
K K+ Cl Cl2-
Ca Ca2+ N N3-
Trend Trend
• Cations are ______________ than the • Anions are _______________ than the
corresponding atoms corresponding atoms
• Taking electrons away means that the • Adding more electrons means that the
remaining electrons are held more tightly by __________________ force from the
the __________________ force from the _______________ is shared by more
______________ charged nucleus electrons so they are held less tightly
Remember… Think purrsitive!
Isoelectronic ions
The electronic configurations of common atoms and ions are shown below in pairs:
Na + 1s22s22p6 Cl - 1s22s22p63s23p6
Ne 2 2
1s 2s 2p 6
Ar 1s22s22p63s23p6
Because these pairs have the same electron configurations we say that they are
________________

3. Now try this…
1) Use a periodic table (p69) to determine the electron configuration of the following ions:
N 3- 1s 2s 2p
2-
O 1s 2s 2p
F- 1s 2s 2p
+
Na 1s 2s 2p
Mg 2+ 1s 2s 2p
2) What do you notice about these different ions?
3) The circles in the diagram below represent these ions. Label the circles to show the order
of ionic radii:
0.072 nm
0.171 nm 0.140 nm 0.133 nm 0.102 nm
4) Can you explain why you chose to label the ions in this order?
Hint: Think about
the electrons and
what holds them
in place around
the nucleus…
Ionic compounds
• Ionic compounds, when solid,
form crystals with characteristic
shapes and structures
• We can use X- ray diffraction to
study the structures of crystals
Lattice structures
•We call the precise crystal
structure of an ionic compound a ______________ structure
• The lattice structure of a compound is the most _____________ way of arranging the ions
• The forces of ______________ between anions and cations make lattices very strong and

4. rigid
u
n it
cell
. unit cell
Sodium chloride (NaCl ) Caesium chloride (CsCl)
Face- centred cubic structure Body- centred cubic structure
a.k.a. rock salt structure
Coordination number = 6 Coordination number = 8
• What does all this mean?
Coordination number -
Unit cell -
• Why are they different?
Hint: Think about the ionic radii (Na = 102 nm, Cs = 167 nm, Cl = 181 nm)
Properties of ionic compounds
• The physical properties of all ionic compounds are very similar
• Consider sodium chloride (NaCl):
Observation
Property Explanation
/ measurement
Melting temperature
801 °C
Boiling temperature
1413 °C
Thermal conductivity
Poor
Electrical conductivity
Solid -
Molten -
Solution -
Hardness
Malleability
NB Other ionic compounds behave similarly; you need to know their general properties and be able

5. to explain how they arise
Evidence for the existence of ions
Ions are obviously far too small for us to see, so how do we know that they definitely exist?
1. Electron density maps
Sodium chloride (NaCl) 4-methoxybenzoic acid
• Electron density maps are produced using X-
ray diffraction
• X-rays are fired through crystals and the interference patterns produced tell us something
about where the ___________________ are
• The big difference between electron density maps for ionic and covalent compounds is…
• This shows that…
2. Electrolysis Watch the electrolysis of copper(II) chromate(VI), draw a diagram and use
it to explain how this experiment suggests the existence of ions

6. Lattice energy “The enthalpy of formation of one mole of an ionic compound from
gaseous ions under standard conditions”
• Making bonds is an _____thermic process, in other words: energy is _______________ as
things become more stable
• Forming an ionic lattice is a very _____thermic process as the ions take up their most stable
possible configuration
• The energy released when forming an ionic lattice is known as the _______________ energy
Born- H ä b er cycles
We can determine lattice energies experimentally by measuring a number of other enthalpies and
combining them in a diagram called a Born- H ä b er cycle to give the lattice enthalpy:
Born- H ä b er cycle for the formation of NaCl:
energy

7. Definitions of terms:
∆ H θf = Standard enthalpy change of formation, measured using a bomb calorimeter
∆ H θat = Standard enthalpy change of atomisation - changing a substance into a
monatomic gas
∆ H θi1 = First ionisation energy - energy required to remove 1 electron and form a
cation,
from spectroscopic measurements
∆ H θe = First elecron affinity - energy required to add 1 electron and form and anion,
also
from spectroscopic measurements
∆ H θlat = The lattice enthalpy of the substance - the enthalpy of formation of one mole of
an
ionic compound from gaseous ions under standard conditions, calculated from
the B-
H cycle
Steps involved in drawing a B-H cycle:
1) Start with the reactants in their standard forms
2) Draw in the enthalpy of formation
(downwards)
3) Atomise the metal (upwards)
4) Ionise the metal (upwards)
5) Atomise the non-metal (upwards)
6) Add an electron to the non-metal (downwards)
7) Use the gap to calculate the lattice enthalpy:
Prep - Use the values in the table below to construct Born-H äber cycles for the formation of
potassium iodide and potassium chloride and hence calculate the lattice energies:
∆Hθf[KI(s)] = -162.8 ∆Hθat[K(s)] = 89.0 ∆Hθat[1/2 Cls(g)] = 122.0 ∆Hθe[Cl(g)] = -349.0
∆Hθf[KCl(s)] = -436.7 ∆Hθi1[K(g)] = 418.8 ∆Hθat[1/2 Is(g)] = 107.0 ∆Hθe[I(g)] = -295.2
All values given in kJ mol-1

8. What factors affect lattice energy?
Compoun Radius Radius Lattice Compound Radius Radius Lattice
d of cation of energy of of anion energy
(nm) anion (kJ mol-1) cation (nm) (kJ mol-1)
(nm) (nm)
NaF 0.102 0.133 -918 NaCl 0.102 0.180 -780
NaCl 0.102 0.180 -780 KCl 0.138 0.180 -711
NaBr 0.102 0.195 -742 MgF2 0.072 0.133 -2957
NaI 0.102 0.215 -705 CaF2 0.100 0.133 -2630
Trend in lattice enthalpies of sodium halides Trend in lattice enthalpies of group 1 & 2
halides
a) Size of cation…
b) Size of anion…
c) Charge on cation/anion…
• This can all be explained using Coulomb’s law:
F = k q1q2
r2
Where q1 and q2 are the charges on the
ions, r is the distance between them and F
is the force of attraction
In other words:
• Smaller ions can get __________ together so the force of
attraction between them is _____________
• The greater the ____________ on an ion the greater its power
to
________________
another oppositely
charged ion

9. Predicting stability - Why not NaCl 2 ?
Q ~ What are the formulae of the following compounds:
i. Sodium chloride
ii. Magnesium chloride
iii. Potassium chloride
Lattice energy / kJ mol -1 iv. Calcium
Compound Born- Difference chloride
Theoretical
H ä b er
NaF -918 -912
NaCl -780 -770 Why do these compounds
NaBr -742 -735 always have the same
NaI -705 -687
AgF -958 -920 formulae? Why is sodium
AgCl -905 -833 chloride never NaCl 2 for
AgBr -891 -816 example or magnesium
AgI -889 -778
fluoride never MgF 3 ?
Dot-cross diagram for NaCl Dot-cross diagram for NaCl2
Sodium loses _____ electron Sodium loses _____ electrons
Use the data-books (p20) to find out the following information:
Sodium chloride
First ionisation energy of sodium
Second ionisation energy of sodium
NaCl2 and MgCl3 will not form Theoretical lattice energy chloride
Magnesium of NaCl -2526 kJ mol-1
because… First ionisation energy of magnesium
Second ionisation energy of magnesium
Third ionisation energy of magnesium
Theoretical lattice energy of MgCl3 -4500 kJ mol-1
Polarisation of bonds
• Born-H äber cycles allow us to calculate lattice energies from ________________ results
• Coulomb’s law allows us to calculate ________________ lattice energies
• The table below compares experimental and theoretical lattice energies:

10. Hmmm….
I think I see
a pattern
developing!
Q ~ What do you notice about the numbers in the table?
A polar bear!
But why is there a difference? We need to look closely at the bonding involved, specifically whether
there is polarisation :
(a) Totally ionic bond (b) Polarised ionic bond (c) Covalent bond
• Ionic bonds can be distorted by the attraction of the positive ______________ for the outer
electrons of the negative _____________
• This means that some ionic bonds show a degree of ________________ character
• These polarised bonds are ______________ than totally ionic bonds and give rise to more
negative lattice energies than predicted theoretically; they are ____________ stable than
expected
What causes ionic bonds to become polarised?

11. Cation Anion
• Ionic radius - Small cations have • Ionic radius - Large anions hold their outer
____________ charge density and so are electrons _________ tightly, so the larger the
more polarising anion, the more __________ it is polarised
• Charge - Cations with a large positive
charge are __________ polarising because
they have stronger attraction for the outer
electrons
Effect on physical properties
1. Because theoretical lattice enthalpies are calculated by assuming that the ions are spherical
and separate they cannot account for polarisation and so are generally lower than lattice
enthalpies calculated from experimental data using Born- H ä b er cycles .
2. The melting temperatures of silver halides are about 20% lower than those of the sodium
halides, which is in agreement with the idea that the silver compounds have a greater covalent
character to their bonding.
Covalent bonding
• Covalent bonding, unlike ionic bonding which involves the _______________ of electrons,
involves the ______________ of electrons between atoms
• Atoms share electrons in order to gain ________ ________ _________
Ionic or covalent?
Whether a compound involves ionic or covalent bonding is related to the potential lattice energy for
that compound:

12. Ionisation Lattice Ionisation Lattice
energy energy energy energy
Ionic Covalent
Dot- cross diagrams
You’ve had lots of practice so you’ll find it easy to draw dot-cross diagrams for these:
Chlorine Water
Methane Oxygen
Electron density maps
• Atoms consist of a nucleus containing __________
and ___________ which is surrounded by orbiting

13. _____________
• In covalent bonds electrons are most likely to be found in the region between the
_____________ charged nuclei
• The area of electron density between the nuclei has a _____________ charge, which
____________ the nuclei inwards and keeps them _______________ bonded together
How strong and long are covalent bonds?
The distance between two atoms involved in covalent bonding depends on two factors:
1. The attraction between the bonding _____________ and the nuclei
2. The repulsion between the like-charged _______________
How the two forces balance out can be shown in the following diagram:
Separation / nm
Energy
• Covalent bonds are of such a Bond Bond length / Bond enthalpy / kJmol -1
length that attraction and repulsion nm
H-H 0.74 436
are balanced C-C 0.154 348
• Bond enthalpy is a measure of C=C 0.134 612
C-H 0.109 412
bond strength
Dative covalent bonds
Sometimes both of the ______________ that make up a covalent bond come from the same atom,
this is called a dative covalent bond:
Ammonia
(NH4+)

14. Aluminium
chloride dimer
(Al2Cl6)
• We represent dative bonds with an arrow:
Ammonia Aluminium chloride
• Dative bonds are the same length and strength as normal covalent bonds
• Another example of a compound with dative bonding is carbon monoxide (CO)
Shapes of covalent molecules
• Different molecules have different shapes, remember methane, which is tetrahedral
Draw dot-cross diagrams for the following compounds:
Compoun Electron density
Dot- cross diagram Shape
d diagram
BeCl 2
CH 4
BF 3

15. H 2O
• There is a strong force of _______________ between electron pairs
• As a result of this the bonding and non- bonding pairs of electrons will spread out to be as
far apart from each other as possible
• This gives rise to the molecular geometries shown above
Giant atomic structures
• Many covalently bonded compounds exist as molecules e.g.
• Others, like diamond (diagram right) exist as atomic
crystals or molecular crystals
• These giant atomic structures form crystal lattices
held together by ___________ bonds
• Other examples of giant atomic structures are silicon
carbide (SiC) and silicon(IV) oxide (SiO2)
• Compounds with giant atomic structures all have
similar physical properties:

16. Metallic bonding
• Most simple model of bonding in metals
• Metal crystal containing metal cations
surrounded by a delocalised ‘sea’ of the
electrons that they have lost
Properties of metals
How does the theory of
metallic bonding explain the following properties of metals? (p92)
Electrical conductivity:
High thermal conductivity:
High melting/boiling temperatures:

17. Malleability and ductility:
Predicting the properties of metals
• The theory of _______________ bonding can be used to predict the properties of metals
• The strength of metallic bonding depends on the charge density of the cloud of
delocalised _________________ and the metal ________
• Higher charge density means _______________ electrostatic attraction
• Lower charge density means ____________ electrostatic attraction
Q ~ Use the theory of metallic bonding to explain the differences shown in the following table:
Melting temp / Thermal conductivity / W/cm/
Metal
ºC K
Sodium 98 1.35
Magnesium 649 1.5
Copper 1083 3.85
a) Melting temperature…
b) Thermal conductivity…
Summary of chapter 1.4 - Bonding

18. Ionic bonding
Covalent
bonding
Metallic
bonding
Practice questions:
Ionic bonding
1. Magnesium reacts spontaneously and extremely vigorously with fluorine gas
a) Write an equation, including state symbols, for the reaction of magnesium with fluorine.
b) Draw a dot-cross diagram to show the bonding in magnesium fluoride.

19. 2. Compare and contrast the structures of sodium chloride and caesium chloride. Why do they
differ?
3. Sodium chloride is, under standard conditions, an ionic solid and is used as a flavour enhancer
a) Under what conditions does sodium chloride conduct electricity? Explain your answer.
b) For which compound would you expect the lattice enthalpy to be more negative: sodium
chloride or magnesium chloride? Explain your choice.
c) Why does sodium chloride always exist as NaCl and never NaCl2?
4. Why is the theoretical lattice enthalpy for silver iodide so different to the experimentally determined
value?
Covalent bonding
1.Define the term covalent bond.
2. Nitrogen is the most abundant gas in the Earth’s atmosphere. It occurs as a covalently bonded
molecule, N2
a) Draw an electron density diagram for nitrogen
c) Draw a dot-cross diagram to show the bonding in nitrogen
3. Beryllium chloride, BeCl2, is a chemical compound that can be said to be ‘electron-deficient’. It
forms pairs of linked molecules known as dimers when solid.
a) Draw a dot-cross diagram for a single BeCl2 molecule

20. b) Draw a dot-cross diagram for a Be2Cl4 dimer
c) What type of bonding is present in Be2Cl4?
4. Covalent bonds vary in length depending on what atoms are involved. Describe the factors which
influence the length of a covalent bond; you can include a diagram as part of your answer.
Metallic bonding
1. Sodium is shiny, grey metal with a relatively low melting temperature. It is also relatively soft and
can be cut with a sharp knife at room temperature.
a) Describe the type of bonding that is present in solid sodium metal
b) Explain why iron is much harder than sodium
c) Explain why copper conducts thermal energy more readily than sodium
2. Most metals are said to be malleable. What does the word malleable mean and how does the
theory of metallic bonding account for this property?

21. Now go away and REVISE! There is a test on this chapter just around the corner and
that’s going to be closely followed by the school exam, which will also include
questions on… guess what… BONDING!