of
Bonding
A chemical bond is a force that holds atoms together.
 Ionic Bond
 Covalent Bond
 Metallic Bonding
The atoms can be the same or different. Eg HCl, O2
The physical and chemical properties of the atoms involved in the bonding
of elements and compounds is affected.
The bonds between atoms and how they pack together can help determine
melting point trends. The arrangements of bonds in a molecule help
determine it shape.
Ionic Bonding
 Forms between a metal and a non- metal
 Happens between atoms with significant differences between their
Electronegativities
Eg Na 0.9, Cl 3.0
Oppositely charged ions are held together by a strong electrostatic force
Called an ionic bond.
Ionic compounds are electrically neutral. They consist of positive ionscations and negative ions- anions held together in a giant lattice structure.
A cation is formed when a metal loses 1 or more electrons.
Eg Mg
Mg2+
It is positively charged because it has more protons then electrons
A n anion is formed when a non-metal gains 1 or more electrons then
protons.
The lattice arrangement of an ionic compound is arrangement of ions that
maximises the attractive forces between oppositely charged ions but
minimises the repulsive forces between like charges.
The forces act equally in each direction holding the ions together tightly. The
lattice forms ionic crystals.
Why do atoms gain/lose electrons?
Full/half full subshells of electrons are stable. The high ionisation energies of
the Noble gases (full outer shell) show that they are very stable. A lot of
energy is required to remove an electron from their full outer shell.
Sodium can achieve a full outer shell by losing an electron from its 3s
subshell.
Na
(1s2 2s2 2p6 3s1)
Na+
(1s2 2s2 2p6)
This gives it the same electronic configuration as
Ne the Noble gas(1s2 2s2 2p6)
Na+ and Ne are said to be isoelectronic (the same number and arrangements
of electrons)
Chlorine can gain an electron in its 3p subshell to form the chloride ion Cl-.
The Cl- ion is isoelectronic with Argon.
Cl
Cl-
(1s2 2s2 2p6 3s2 3p5)
(1s2 2s2 2p6 3s2 3p6)
Ar
(1s2 2s2 2p6 3s2 3p6)
Na and Cl achieve Noble Gas status by transferring an electron from Na’s 3s
Subshell to Cl’s 3p subshell.
This is a complete transfer of electrons form one atom to the other and an
ionic or electrovalent bond is formed.
This transfer can be drawn showing all the electrons or as an outer shell
Electron only diagram.
Full diagram p46-7 c4u
Dot cross diagrams ed p79
Properties of Ionic Compounds
They:
1. Are made of crystals (which can be split or cleaved along certain
angles)
2. Usually have high melting points
3. Are often soluble in water
4. Conduct electricity when molten or dissolved in water, but not when
solid
Evidence
Electron density maps
If you pass x rays through a crystal the x rays are scattered or diffracted by
the electrons in the atoms or ions in the structure. This produces a
diffraction pattern which can explain the arrangement of the atoms/ions
in the crystal.
The patterns are called electron density maps. (p83)
The x rays land on a photographic film. The bigger the ion/atom the more
electrons it has and the brighter the spot it produces.
By analysing the positions and intensities of the spots experts can work
out the charge density of the electrons in the crystal.
This is defined as amount of electric charge per unit volume.
Points of equal charge density are joined up to form contour lines that
produce a ‘map’.
Migration of Ions
Ionic solids do not conduct electricity because
 They do not have delocalised electrons
 The ions are fixed in a lattice
However when the solid is dissolved or melted (molten) the lattice braks
down and the ions are free to move. The current is a flow of Ions.
The positive ions move towards the negative charged cathode where they
gain electrons and are reduced to atoms.
The negative ions move towards the positively charged anode where they
lose electrons and are oxidised to atoms.
At the cathode reduction occurs the metal gains electrons
M+ + e-
M
At the anode oxidation occurs the non metal loses electrons
NM-
NM + e-
This process is electrolysis.
The movement of ions can be seen in the electrolysis of potassium
manganate (VII).
The potassium ions are colourless
Manganate ions are purple
Transition metal ions are almost all coloured....you must learn the colours
To be able to explain the results in an electrolysis experiment (yr11)
The movement of ions in the electrolysis of copper chromate(VI).
(OHT)i
Copper chromate is then dissolved in a minimum amount of acid. The
solution is then saturated with urea, a water soluble covalent substance. Thsi
increases the density of the solution. The U tube is half filled with acid and
the copper chromate is carefully added using a long stemmed funnel so that
two separate layers are formed.
The graphite/ platinum electrodes are placed in the acid just above the
minutes.
The Cu2+ ions move to the negative cathode and the solution around the
electrode turns blue green because of the Cu ions.
The solution around the positive anode becomes yellow because of the CrO42Ions moving towards it.
Melting point and solubility
Ionic compounds have very high melting points. Only a few can be melted in
a laboratory. This is because of the strong forces of attraction between ions.
High solubility also provides evidence for ionic bonding....when the ions
become hydrated the energy released compensates for the energy required
to break the bonds between the positive and negative ions in the lattice. The
positive ions are surrounded by δ – oxygen atoms of water molecules and
the negative ions are surrounded by the δ+ hydrogen.
Both high mp and solubility are evidence of ionic compounds (as covalent
compounds can have high mp too).
Strength of Ionic Bonds
Determined by
 The charges on the ions
The force of attraction depends on the product of the charges eg the
attraction between a 2+ and 1- is twice the attraction between a 1+
and a 1- ion
 The radii of the ions
Positive ions are smaller than the parent atom
Negative ions are larger than the parent atom
Na 0.191nm
Na+ 0.102nm
F 0.071nm
F- 0.133nm
The more e- an atom gains or loses the greater the difference between the
atomic and ionic radii.
Na 0.191nm
Na+ 0.102nm
Mg 0.160nm
Mg2+ 0.072nm
F
F-
0.071nm
O 0.073 nm
0.133nm
O2- 0.140nm
In positive ions the remaining e- are more tightly bound to the positive
nucleus.
In negative ions the additional e- means that all the e- are less tightly
bound to the nucleus.
As we go down a group the ions have more electron shells therefore the
ions get larger therefore the radii of group 7 increases in the order
F-<Cl-<Br-<IExtent of covalency in an ionic compound!
The positive cation (metal) exerts an attraction on the electrons in the
negative anion (non metal).
If the e- are significantly pulled towards the cation, the anion is polarised.
 Cations with a small radius and high charge have a high polarising
power
 Mg ion is more polarising than Ca both have a charge of 2+
 Mg ion is more polarising than Na ion, because it is 2+ and it is smaller




Polarising power is measured by charge density of the cation
Anions with high charge and large radii are easily polarised
I- is more easily polarisable than the smaller ClSulphide ion S2- is more easily polarisable than a chloride ion because
it has a greater charge and larger radius
Cations with a large charge and small radius are highly polarising .
Anions with a large charge and large radii are highly polarisable.
Substance
NaCl
K2S
MgI2
Extent of Covalency Explanation
Almost none
Significant
Considerable
Which shows more covalent character calcium chloride or potassium
chloride.
Explain why?
Trends in isoelectronic ions
N3-, O2-, F-, Na+, Mg2+ and Al3+ all have the same electronic configuration 2,8.
They are said to be isoelectronic –identical electronic configuration. Their
radii decreases from N3- ,which is the largest , to Al3+ which is the smallest.
This is because the charge on the nucleus increases from 7 in N3- to
13 in Al 3+. The greater nuclear charge is experienced by the same number of
electrons, so the attraction between the nuclei and the orbiting electrons
goes up, making the ions progressively smaller.
Lattice energy
Like bond enthalpies, enthalpy changes involved in the
formation of a compound are useful in helping us predict
whether a reaction is exothermic or endothermic, in the
same way the formation of ionic substances is associated
with energy changes too!
We can estimate the strength of the attraction between a
positive and negative ion in an ionic lattice from the lattice
enthalpy.
The enthalpy of formation of one mole of an ionic compound
from gaseous ions under standard conditions.
The formation of ions in the gaseous state from the elements
in the standard state is endothermic, whilst the formation of
the lattice involves a release of energy-exothermic!
Therefore the lattice energy is always negative.
The higher the lattice enthalpy the greater the attraction
between the positive and negative ions.
Na+ (g) +Cl- (g)
NaCl (s)
ΔHθle
The lattice energy of each ionic compound is different and
depends on the degree of attraction between the positive
and negative ions.
Born Haber Cycles
Experimentally it is impossible to use a direct method to
determine the lattice energy of an ionic solid. An indirect
method involving Hess’s Law is used.
Hess’s Law states that if a change can be brought about by
more than one route than the enthalpy change for each
route must be the same, providing that the starting and
finishing conditions are the same for each route.
Enthalpy of formation of NaCl (s)
This is the enthalpy change for the reaction of solid sodium
with gaseous chlorine molecules.
The enthalpy change for the one step reaction, by Hess’s
law, should be equal to the sum of changes involved if the
reaction were to take place in several steps. First
suggested by Born and Haber! The steps involved in this
process are:




Turning solid sodium into gaseous sodium
Splitting chlorine into atoms
Removing an electron from each gaseous sodium
Adding an electron to each chlorine atom
 Bringing the ions together into an ionic lattice.
See oht/fig1.4.11 p84 AS ed
The value for the lattice energy gives you an idea of the attractive
forces between the ions of NaCl. This explains the high mp and bp of
ionic substance.
Copy table 1.4.2 p85. Explain the trends seen in the table.
Calculate the lattice energy for Calcium Oxide given the following
data:
ΔHf (CaO) = -635Kj /mol
ΔHat (Ca)= +178kj/mol
ΔHie first and second (Ca)= +1735kj/mol
ΔHat(O2) = +249kj/mol
ΔHea first and second (O) = = +657KJ/mol
Factors that affect lattice energy p109 facer/p85 as ed
Charge and size of ions