Bonding
HL and SL
4.1 Ionic Bonding
An ionic bond is an electrostatic attraction between ions of
opposite charge.
Positive ions (or cations) are formed when electrons are lost
from the outer energy level of, usually, a metal atom.
Negative ions (or anions) are formed when electrons are
gained to the outer energy level of a non-metal atom.
Ions usually have filled outer energy levels (except certain
transition metal ions).
Transition metals can form more than one positive ions e.g.
Fe2+ and Fe3+, Cu+ and Cu2+.
Positive ions are formed when atoms lose electrons.
e.g.
Li(2, 1)  Li+(2) + eLi(1s22s1)  Li+(1s2) + e-
HL only
Do similar to show the formation of the following ions:
Na+ Mg2+ K+
Ca2+ Al3+
Cu2+ and Fe3+ (HL only)
Think carefully about the number of electrons that need to
be lost.
Is there a link between the formula of the ion and the
position of the atom in the periodic table?
Negative ions are formed when atoms gain electrons.
e.g.
F(2, 7) + e-  F-(2, 8)
F(1s22s22p5) + e-  F-(1s22s22p6)
HL only
Do similar to show the formation of the following ions:
ClO2S2N3P3Br- (HL only)
Think carefully about the number of electrons that need to be
gained.
Is there a link between the ion formed and the position of the
element in the periodic table?
All of the ions considered so far are made up of just one
atom but we should know the formula of the following
polyatomic ions:
NO3-
nitrate(V) ion
OH-
hydroxide ion
SO42- sulfate(VI) ion
CO32- carbonate ion
PO43- phosphate(V) ion
HCO3- hydrogencarbonate ion
NH4+ ammonium ion
Brackets are used in formulae to show more than 1!
Ionic compounds are formed when electrons are transferred
from a metal atom to a non-metal atom.
A compound will be ionic if the difference in electronegativity
of the two atoms is greater than 1.8 (see later).
Look at the model for the formation of sodium chloride on the
next slide.
What is wrong with this model?
+
11 p
11 12
p n
12 n
-
chemical
reaction
a sodium ion (Na+)
a sodium atom (Na)
(2, 8)+
2, 8, 1
17 p
17 p 18 n
18 n
a chloride ion (Cl-)
a chlorine- atom (Cl)
(2, 8, 8)
2, 8, 7
The formula of sodium chloride is NaCl
2+
1212
p p
1212
n n
2-
chemical
reaction
2+)
a amagnesium
ion
(Mg
magnesium atom (Mg)
2+
(2,
8)
2, 8, 2
8p 8p
8n 8n
2-)
a
oxide
ion
(O
an oxygen atom (O)
2(2,
8)
2, 6
The formula of magnesium oxide is MgO
calcium chloride
20 p
20 n
17 p
18 n
a calcium atom (Ca)
a chlorine atom (Cl)
2, 8, 8, 2
2, 8, 7
The calcium atom has two electrons in its outer shell
which must both be lost for it to have an electronic
structure of a noble gas. The chlorine atom only has
space for one more electron in its outer shell.
What happens?
The calcium atom reacts with 2 chlorine atoms
17 p
18 n
20 p
20 n
17 p
18 n
17 p
18 n
2+
20 p
20 n
17 p
18 n
The formula of calcium
chloride is CaCl2
Compounds which contain ionic bonds are solids at room
temperature.
Ionic bonds do not exist in isolation but as part of a GIANT
IONIC LATTICE.
In a giant ionic lattice each positive ion is attracted by
negative ions which surround the positive ion in a regular
arrangement and vice versa.
This attraction is an electrostatic attraction.
e.g. sodium chloride
Or in 2-d:
sodium ion
chloride ion
Ionic compounds conduct electricity when molten or in
aqueous solution. They do not conduct when solid as the
ions are held in fixed lattice positions. When liquid or in
aqueous solution, the ions are free to move.
4.2 Covalent Bonding
Covalent bonding is an alternative way for atoms to
achieve a filled outer energy level.
Covalent bonds are formed between non-metal atoms or
atoms with an electronegativity difference of less than 1.8.
A covalent bond is an electrostatic attraction between a
shared pair of electrons and the positively charged nuclei of
the two atoms held by the bond.
Often shown
as H – F.
Other examples include:
Here are some hints to help draw
Lewis structures:
These diagrams are called
LEWIS STRUCTURES.
1.
Work out the number of shared
pairs required by each atom.
This is the same as the number of
extra electrons required to fill the
outer shell.
2.
Draw a diagram showing the
bonds between atoms only e.g.
Cl – Cl.
3.
Draw an outline dot-cross
diagram.
4.
Fill in the electrons in the
covalent bonds.
5.
Make sure that the outer shells
are filled.
Draw Lewis structures for:
1. Hydrogen chloride, HCl
7. Carbon dioxide, CO2
2. Oxygen, O2
8. Hydrogen cyanide, HCN
3. Tetrachloromethane, CCl4
9. Carbon disulphide, CS2
4. Hydrogen sulphide, H2S
10. Ethene, C2H4
5. Phosphorus(III) chloride,
PCl3
11. Ethyne, C2H2
6. Nitrogen, N2
12. Carbon monoxide, CO
Co-ordinate bonding
Aka dative covalency
A co-ordinate bond is formed when both electrons in a shared
pair originate from one atom. This pair is called a lone pair
on the donor atom.
Once a co-ordinate bond is formed it cannot be distinguished
from a normal covalent bond.
Co-ordinate bonds are found in transition metal complex ions
(HL see later), the ammonium ion, and the hydroxonium ion.
The ammonium ion is formed when ammonia reacts with H+
lone pair of electrons
+
H+
H+ has no electrons
H
+
co-ordinate bond
ammonium ion
H3O+ (the hydroxonium ion) is formed when water reacts
with H+. Draw similar diagrams to show the formation of
the hydroxonium ion.
Can H4O2+ be formed?
In theory yes but in practice no!
Why not?
Other examples of substance containing co-ordinate bonds
are carbon monoxide (CO) and aluminium chloride (Al2Cl6).
Draw Lewis structures for these compounds.
Bond Length and Bond Strength
Tabulate the bond length and bond strength of the following
bonds:
O–O
O=O
N–N
N=N
NN
C–C
C=C
CC
What patterns do you notice?
Find out the C to O bond lengths in either the ethanoate ion
or the carbonate ion.
How do you explain this? Try drawing Lewis structures for
these substances to help.
Other substances for which resonance hybrid structures can
be drawn include:
Sulphur dioxide, SO2
Ozone, O3
Nitrate ion, NO3Benzene, C6H6
Can you draw Lewis structures for them?
Ionic or covalent?
Generally:
Ionic if electronegativity difference between atoms > 1.8
Covalent if electronegativity difference between atoms < 1.8
However, some compounds are described as ionic with some
covalent character whilst others are described covalent with some
ionic character. Most compounds have bonding somewhere
between the two extreme models of ionic and covalent bonding.
Look at the substances properties, if they are given, to help you
decide.
For example: if it conducts when aqueous or molten it must be
ionic.
Polar Covalent bonds
Arises due to differences in electronegativity between the
two atoms held by a covalent bond.
Electronegativity is defined as the power of an atom to
withdraw electron density from a covalent bond.
Small atoms with a large number of protons attract
electron density most strongly.
So electronegativity in the periodic table increases
• from left to right
• bottom to top
H
2.1
He
Li
1.0
Be
1.5
B
2.0
C
2.5
N
3.0
O
3.5
F
4.0
Ne
Na
0.9
Mg
1.2
Al
1.5
Si
1.8
P
2.1
S
2.5
Cl
3.0
Ar
Br
2.8
Kr
See later in Periodicity
topic.
increases
When a covalent bond exists between atoms of differing
electronegativities, the shared pair is displaced towards the
more electronegative atom.
A
A
‘normal’ covalent
bond with electron
pair midway
between the two
atoms
A
B
B is more electronegative
than A so the electron pair
is displaced towards B.
The result is a polar
covalent bond.
The displacement of electron density makes the less
electronegative atom slightly electron deficient so has a small
positive charge shown as d+.
The more electronegative atom has a slight excess of electron
density which is shown as d-.
d+
d-
A
B
If the molecule is placed between electric plates the d- end is
attracted to the positive plate and the d+ end is attracted to
the negative plate. The molecule is said to have a dipole
moment, the bigger the difference in electronegativities, the
bigger the dipole moment.
Examples of molecules with polar bonds include:
d+
d-
H
F
d+
d-
H3C
Cl
Shapes of molecules:
If a molecule contains polar bonds, it does not necessarily
follow that the molecule itself is polar.
If the molecule is symmetrical then the dipoles will cancel
each other out.
So we need to be able to deduce the shape of a molecule
before we can decide whether or not it is a polar molecule.
To do this we use the VALENCE SHELL ELECTRON PAIR
REPULSION (VSPER) theory.
For the purposes of this theory we must treat double and triple
covalent bonds as if they are single bonds.
The outer electrons of atoms are arranged in pairs.
These electron pairs can be considered as ‘clouds’ of electron
density which repel each other as far apart as possible.
Shapes of molecules depend upon number of electron pairs
around the central atom in an ion or molecule.
n.b. it is electron pairs that repel each other not the
atoms.
2 pairs
3 pairs
4 pairs
Rules to help determine number of electron pairs around
central atom:
Add together:
• number of electrons in outer shell of central atom when not
bonded,
• number of shared pairs of electrons (same as number of
surrounding atoms)
Then if it is an ion
• Add one for each negative charge, subtract one for each
positive charge.
Some examples:
BeCl2
There are 2 electrons from the outer shell of beryllium
+ 1 from each of the chlorine atoms.
Total 4 electrons
i.e. 2 pairs of electrons
Two electron pairs repel each other equally so shape is
linear.
Bond angle 180 °.
CH4
There are 4 electrons from the outer shell of carbon +
1 from each of the hydrogen atoms.
Total 8 electrons
i.e. 4 pairs of electrons
Four electron pairs repel each other equally so shape
is tetrahedral.
Bond angle 109.5 °.
Lone pairs of electrons are shorter and fatter than pairs
involved in bonding. Why?
As a result they repel more than bonding pairs leading to
a reduction in bond angles between bonding pairs.
So in ammonia, NH3
5 electrons from N + 3 from shared pairs = 8  4 pairs
But only 3 involved in bonding so there is one lone pair
which reduces the bond angle to 107 °.
So ammonia is:
This shape is
sometimes referred
to as pyramidal.
So in water, H2O
6 electrons from O + 2 from shared pairs = 8  4 pairs
But only 2 involved in bonding so there are two lone pairs
which reduces the bond angle to 105 °.
So water is:
This shape is
sometimes referred
to as v-shaped or
bent.
4.3 Intermolecular Forces
Intermolecular forces are weak forces between molecules.
There are 3 types ( in order of increasing strength):
• van der Waals forces,
• permanent dipole – permanent dipole attractions,
• hydrogen bonding.
Van der Waals forces
Aka temporary dipole – induced dipole attractions.
Found in all molecules.
About 1 % strength of a covalent bond.
Arises due to fluctuations of electron clouds.
At any instant there may be more electron density on one part
of the molecule (lhs on illustration)  small negative charge,
d-, on part of the molecule and small positive charge on
another part, d+.
d-
d+
Temporary dipole
This causes an adjacent molecule to have a dipole (an
induced dipole) as the negative charge in the temporary
dipole repels electron density in the neighbouring
molecule.
Forces only act for a
short time as the
electron density is
continually
changing. These
forces are
continually switched
on and off.
d-
d+
Induced
dipole
d-
d+
Temporary
dipole
Consequences of Van der Waals forces:
1. Increase in boiling point of alkanes with Mr. The
larger the molecule, the bigger the magnitude of van
der Waals forces. See later.
2. Variation in b.p. of isomers of pentane. See later
309 K
185 K
283 K
Dipole – Dipole attractions
Molecules with permanent dipoles attract each other as shown
below:
d+
d-
d-
d+
H – Cl
d+
d-
H – Cl
Cl – H
= weak electrostatic attractions
Hydrogen Bonding
The big daddy of them all but still only maximum of 1/10th
strength of a covalent bond.
It is a special case of a dipole – dipole attraction force which
exists between a lone pair of electrons on a N, O or F atom
and a hydrogen atom with a strong partial positive charge
(d+) as it is attached to an atom with a large electronegativity
(N, O or F)
The electronegative atom pulls electron density from the
hydrogen atom to such degree that the hydrogen almost
appears like an unshielded proton.
e.g in water
180 °
hydrogen bond
Note that the nucleus of the hydrogen bonded hydrogen
atom is always in line with the nuclei of the two
electronegative atoms on either side.
Hydrogen bonding is also found in:
HF, hydrogen fluoride
NH3, ammonia
C2H5OH, ethanol
CH3COOH, ethanoic acid
In each of these cases the molecules have relatively
high boiling points for such small molecules.
What is the trend in boiling points of the hydrides of
groups 4, 5, 6 and 7?
i.e for group 6: H2O, H2S, H2Se, H2Te and,
for group 7: HF, HCl, HBr, HI
Plot a graph of period number against boiling point for
each of these groups on one set of axes.
Boiling points should increase as molecules increase in
size down the group so the van der Waals forces
increase.
But ……….
H2O and HF have higher
boiling points than expected
due to hydrogen bonding
Why is hydrogen bonding essential to life?
4.2 Covalent Bonding (more!)
Giant covalent structures
Carbon exhibits allotropy (it exists in more than 1 form in
the solid state).
The allotropes of carbon are DIAMOND, GRAPHITE and
the FULLERENES.
Diamond
Graphite
Bucky ball
In diamond each C atom is bonded covalently to 4 others in a
giant tetrahedral structure.
Diamond is extremely hard as a consequence of the strong
covalent bonds within the molecule.
To melt diamond requires a lot of energy as all of the covalent
bonds have to be broken.
There are no free electrons so diamond does not conduct
electricity.
Silicon and silicon dioxide also have a diamond structure
In graphite each C atom is bonded covalently to 3 others in a
trigonal planar arrangement forming hexagonal rings in layers.
The forces between the layers are weak as they are formed by
delocalised electrons between the layers (the 4th electron from
the outer shell of C).
Graphite conducts electricity due to these delocalised electrons.
As the forces are weak between the layers graphite feels slippery.
The layer can move over each other so graphite can be used as a
lubricant.
It also has a high melting point as there are lots of strong
covalent bonds to break.
The fullerenes were discovered jointly by Harold Kroto and
Richard Smalley in 1996. In 1997 they were awarded the
Nobel Chemistry Prize.
The basic fullerene, C60, is generally known as buckminster
fullerene (or a bucky ball). It consists of 60 carbon atoms
joined in a combination of hexagonal and pentagonal rings (like
a football).
4.4 Metallic Bonding
Metallic bonds do not exist in isolation. They form part
of a giant metallic lattice.
A giant metallic lattice consists of close-packed metal
ions surrounded by delocalised electrons. These
delocalised electrons are free to move through the lattice.
Where do they come from?
Metallic Bonding
Metallic bonds do not exist in isolation. They form part
of a giant metallic lattice.
A giant metallic lattice consists of close-packed metal
ions surrounded by delocalised electrons. These
delocalised electrons are free to move through the lattice.
Where do they come from?
4.5 Physical Properties
Type of
bonding
Structure
Covalent Small molecule with van der
Waals forces
Covalent Small molecule with dipole –
dipole forces
Covalent Small molecule with hydrogen
bonding
Covalent Giant lattice of atoms
Ionic
Giant lattice of positive and
negative ions
Metallic
Giant lattice of metal ions
surrounded by delocalised
electrons
Example
Properties