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MISS CHOHANS NOTES ON BONDING UNIT 2 EDEXCEL
BONDING UNIT 2 – EDEXCEL NOTES
SHAPES OF MOLECULES
MISS CHOHANS NOTES ON BONDING UNIT 2 EDEXCEL
MISS CHOHANS NOTES ON BONDING UNIT 2 EDEXCEL
DOT AND CROSS DIAGRAM TO EXPLAIN SHAPE
Beryllium chloride, BCl2
Boron trifluoride, BF3
There are three bonding pairs (electron area in a bond);
these will spread the maximum distance apart - that is at
an angle of 120o.
This molecule is flat, that is it lies in a plane;
such a molecule is said to be planar. With
three bonds at an angle of 120o BF3 is said
to be trigonal planar.
This molecule is flat, that is it lies in a plane; such a
molecule is said to be planar. With three bonds at an
angle of 120o BF3 is said to be trigonal planar.
MISS CHOHANS NOTES ON BONDING UNIT 2 EDEXCEL
Phosphorus pentachloride, PCl5
Sulphur hexafluoride, SF6
Effect of multiple bonds on bond strength and length.
Nuclei joined by multiple (i.e. double and triple) bonds have a greater electron density between them. This causes a
greater force of attraction between the nuclei and the electrons between them, resulting in a shorter bond length and
greater bond strength
MISS CHOHANS NOTES ON BONDING UNIT 2 EDEXCEL
So in multiple bonded compounds we just ignore the double or triple bond and just take it as a single covalent
bond to work out the number of electron pairs around the central atom:
i.e.
VALENCE SHELL ELECTRON PAIR
REPULSION THEORY (VSEPR) – States that:
 Electron pairs repel as far apart as possible
 LPLP > LPBP > BPBP
Where
Lp = lone pair
Bp = Bond pair
> = more than
MISS CHOHANS NOTES ON BONDING UNIT 2 EDEXCEL
HOW TO EXPLAIN SHAPE
1. State number of bonding pairs and lone pairs of electrons.
2.
State that electron pairs repel and try to get as far apart as possible (or to a position of minimum repulsion.)
3. If there are no lone pairs state that the electron pairs repel equally
4.
If there are lone pairs of electrons, then state that lone pairs repel more than bonding pairs.
5.
State actual shape and bond angle.
Remember: lone pairs repel more than bonding pairs and so
reduce bond angles (by about 2.5o per lone pair in above examples)
Ammonia, NH3
Water, H2O
MISS CHOHANS NOTES ON BONDING UNIT 2 EDEXCEL
Bond angles in CH4, NH3, H2O
It is important to note that although there are four pairs of electrons arranged approximately tetrahedrally around
the N in NH3 and the O in H2O, the lone pairs cannot be “seen” experimentally, so the shapes of these molecules are
described by the actual positions of the atoms: ammonia is pyramidal and water is bent.
We can explain the differences in the bond angles in CH4, NH3 and H2O by noting that repulsions get less along the
series:
lone pair/lone pair > lone pair/bonding pair > bonding pair/bonding pair
This occurs because a lone pair is closer to the nucleus of the atom, and so takes up more room than a bonding pair.
When this principle is applied to CH4, NH3 and H2O, CH4 is a regular tetrahedron (angles of 109.5o); NH3 has one
lone pair, which squashes the H atoms down (angles reduced to 107o); and H2O has two lone pairs, which repel the H
atoms even more (angle now 104.5o):
O
N
H
H
H
H
H
ammonia – pyramidal
water: non-linear
bond angle 107o
bond angle 104.5o
More complicated molecules
The shapes of more complicated molecules and ions can also be “explained” by electron pair repulsion theory, or EPR.
In applying this principle, we must include both bonding and non-bonding (or “lone”) pairs, and we must count a
double or triple bond as if it were one pair (or one region of electron density).
The table shows the shapes expected for different numbers of electron pairs.
MISS CHOHANS NOTES ON BONDING UNIT 2 EDEXCEL
We can use advanced EPR theory to predict the shape of any molecule or ion.
1. Decide on the central atom
2. Count the number of bonding atoms
3. If the species is – charged
4. If the species is + charged
5. Find the total number of electron pairs.
6. Determine the shape of the species;
Pairs
Record its number of outer electrons
Add 1 e- for each atom
Add 1 e- for each charge
Subtract 1 e- for each charge
Shape
2
Linear
3
Trigonal Planar
4
Tetrahedral
6
Octahedral
5
Trigonal bipyramidal
7. Show any lone pairs.
The number of Lone pairs = Total number of pairs – bonding pairs
Example
Electrons
PH4+
Central atom
=P
5
Bonding atoms = 4 x H
4 (Forms 4 bonding pairs)
Positively charged
-1
Total = 8
4 pairs = Tetrahedral
The number of Lone pairs = Total number of pairs – bonding pairs = 4-4 = 0
MISS CHOHANS NOTES ON BONDING UNIT 2 EDEXCEL
Example
Electrons
IF4-
Central atom
=I
7
Bonding atoms = 5 x F
4 ( Forms 4 bonding pairs)
Negatively charged
+1
Total = 12
6 pairs = Octahderal
The number of Lone pairs = Total number of pairs – bonding pairs = 6-4 = 2
Carbon Allotropes
Macromolecular: diamond

Tetrahedral arrangement of carbon atoms. 4 covalent bonds per
atom

This structure, held together by strong covalent bonds, is very
difficult to break apart. So diamond has very high melting and
boiling points.

It is the hardest natural substance.

Diamonds are attractive which means they are used as jewellery.
Its hardness makes it useful for cutting instruments such as drill
tips.

Diamond cannot conduct electricity because all 4 electrons per
carbon atoms are involved in covalent bonds. They are localised
and cannot move
MISS CHOHANS NOTES ON BONDING UNIT 2 EDEXCEL
Macromolecular: Graphite

Planar arrangement of carbon atoms in layers. 3 covalent bonds per atom in each layer. 4th
outer electron per atom is delocalised. Delocalised electrons between layers.

Graphite can conduct electricity well between layers because one electron per carbon is free and
delocalised, so electrons can move easily along layers. It does not conduct electricity between
layers because the energy gap between layers is too large for easy electron transfer.

Since three bonds form, the bond angle around each carbon is 120o and a hexagonal
arrangement is set up.
These hexagons join
together in a plane forming
a sheet of hexagonally
arranged carbon atoms.

Graphite is made up of these layers held together by London forces (Van der Waals' forces).
Sometimes it is referred to as a layer structure.

The carbon atoms in graphite are held together by covalent bonds forming a giant structure, so
the melting and boiling points are high.

Since each atom has a free electron, graphite is able to conduct electricity. The layers held
together by weak intermolecular forces can slide over each other, making a soft slippery
substance.

Graphite can be used for electrodes as it is a conductor, and unlike metals it does not react
during electrolysis. It can be used as a lubricant because of its slippery nature.
MISS CHOHANS NOTES ON BONDING UNIT 2 EDEXCEL
Carbon nanotubes

Another form of carbon developed as a result of the discovery
of fullerenes is the nanotube.

The individual layers in graphite are called graphemes. A
nanotube can be regarded as a grapheme which has rolled up
to form a cylinder.

The name comes from the diameter of the cylinder. A singlewalled carbon nanotube is a one-atom thick sheet of graphite
rolled up into a seamless cylinder with diameter 1-2 nm.

These have very high tensile strength because of the strong
structure of many strong covalent bonds

Such cylindrical carbon molecules have novel properties that
make them potentially useful in many applications in
nanotechnology, electronics, optics and other fields of
materials science.

They exhibit extraordinary strength and unique electrical
properties, and are efficient conductors of heat.


Nanotubes can conduct electricity well along the tube because
one electron per carbon is free and delocalised, so electrons
can move easily along the tube.

Nanotubes have potentially many uses. One being the potential
to us as vehicles to deliver drugs to cells.

There are delocalized electrons in buckminsterfullerene.
Fullerenes

These molecules resembled the construction
of a building by an architect called
Buckminster Fuller.

As a result, the molecule was named
buckminsterfullerene. The molecule also
resembles a football, so it is often called a
“bucky-ball”.

Similar molecules have since been made
that contain 70 or more carbon atoms. This
family of molecules are called fullerenes.
MISS CHOHANS NOTES ON BONDING UNIT 2 EDEXCEL
Electronegativity and intermediate bonding
Electronegativity is defined as follows
Electronegativity is the ability of an atom within a covalent bond to attract the bonding
pair of electrons.
Electronegativity is measured on the Pauling scale (ranges from 0 to 4
F, O, N and Cl are the most electronegative atoms
The most electronegative element is fluorine and it is given a value of 4.0
The table below shows electronegativity values of main block elements
Factors affecting electronegativity

Electronegativity increases across a period as the number of protons increases and the atomic radius
decreases because the electrons in the same shell are pulled in more.

It decreases down a group because the distance between the nucleus and the outer electrons increases and
the shielding of inner shell electrons increases
Intermediate bonding





Ionic and covalent bonding are the extremes of a continuum of bonding type. Differences in electronegativity
between elements can determine where a compound lies on this scale
A compound containing elements of similar electronegativity and hence a small electronegativity difference
will be purely covalent
A compound containing elements of very different electronegativity and hence a very large electronegativity
difference (> 1.7) will be ionic
Formation of a permanent dipole – (polar covalent) bond A polar covalent bond forms when the elements in
the bond have different electronegativities . (Of around 0.3 to 1.7) When a bond is a polar covalent bond it
has an unequal distribution of electrons in the bond and produces a charge separation, (dipole) δ+ δ- ends
The element with the larger electronegativity in a polar compound will be the δ- end H+ – Cl-
MISS CHOHANS NOTES ON BONDING UNIT 2 EDEXCEL
MORE ON INTERMEDIATE BONDING

In Unit 1 (Ionic Bonding) it was noted that ionic compounds can be polarized which gives them a covalent
character. When a molecule has a polar bond, it gives the covalent substance an ionic character.

To regard a compound as “covalent” or “ionic” is too simplistic for understanding AS chemistry. It is more correct
to visualize type of bonding on a sliding scale where compounds can be described as “predominantly covalent” or
“predominantly ionic”.

Electronegativity values can be used to give an approximate idea of the predominant type of bonding in a binary
compound
Looking at a selection of substances
If these are then plotted on the diagram below, it can be seen the type of bonding is a continuum rather than a black and
white picture.
Polarisation of a covalent bond
A perfect covalent bond would consist of the bonding electron region being shared equally by each atom.
MISS CHOHANS NOTES ON BONDING UNIT 2 EDEXCEL

This occurs if each of the atoms have the same pull on the electron pair in the bond (equal electronegativity). It
can be found in molecules of elements such as O2, Br2 and N2. Since the atoms in each of these molecules are the
same they will have the same electronegativity.

In most compounds however one atom will have a greater electronegativity than the other, and so will have a
greater pull on the electrons, so distorting the electron region.

This process of moving away from the perfect example is called polarisation. The extent of the polarisation will
depend on the difference in electronegativity of the two atoms.

The polarisation of a covalent bond will mean that one part of the molecule is more negative (the most
electronegative atom) than the other and causing the bond to be polar.

Hydrogen chloride is an example of a molecule which contains a polar bond.

The chlorine possesses a higher electronegativity, so will draw the electron pair in the covalent bond towards itself.
Polar and Non Polar molecules
Symmetric molecules
A symmetric molecule (all bonds identical and no lone pairs) will not be polar even if individual bonds within the
molecular ARE polar.
The individual dipoles on the bonds ‘cancel out’ due to the symmetrical shape of the molecule. There is no NET dipole
moment: the molecule is NON POLAR
e.g. CCl4 will be non-polar whereas CH3Cl will be polar
MISS CHOHANS NOTES ON BONDING UNIT 2 EDEXCEL
The individual dipoles on the bonds ‘cancel out’ due to the symmetrical shape of the molecule. There is no NET dipole
moment: the molecule is NON POLAR
e.g. CCl4 will be non-polar whereas CH3Cl will be polar
Experiment effect of charged rod on polar/non-polar liquids

In this experiment, a charged rod (formed by rubbing a plastic rod) is brought
close to a jet of liquid flowing from a burette.

If the liquid is polar, the jet of liquid will be attracted to the electrostatic force
of the rod. The dipoles in the polar molecules will all align and the negative
end δ- will be attracted to the positive rod (or vice versa). The stronger the
dipole the more the deflection of the jet.

Non-polar liquids will not be deflected and attracted to the charged rod
Intermolecular Forces
London Forces
London forces occur between all molecular substances and noble gases. They do not occur in ionic substances

London Forces are also called instantaneous, induced
dipole dipole interactions.

They occur between all simple covalent molecules
and the separate atoms in noble gases.

In any molecule the electrons are moving constantly
and randomly. As this happens the electron density
can fluctuate and parts of the molecule become more
or less negative i.e. small temporary or transient
dipoles form.

These temporary dipoles can cause dipoles to form in
neighbouring molecules. These are called induced
dipoles.

The induced dipole is always the opposite sign to the
original one.
MISS CHOHANS NOTES ON BONDING UNIT 2 EDEXCEL
London Forces Continued
Main factor affecting size of London Forces
 The more electrons there are in the molecule the higher the chance that temporary dipoles will
form. This makes the London forces stronger between the molecules and more energy is needed to
break them so boiling points will be greater
 The increasing boiling points of the halogens down the group 7 series can be explained by the
increasing number of electrons in the bigger molecules causing an increase in the size of the
London forces between the molecules. This is why I2 is a solid whereas Cl2 is a gas.
 The increasing boiling points of the alkane homologous series can be explained by the increasing
number of electrons in the bigger molecules causing an increase in the size of the London forces
between molecules.
 The shape of the molecule can also have an effect on the size of the London forces. Long straight
chain alkanes have a larger surface area of contact between molecules for London forces to form
than compared to spherical shaped branched alkanes and so have stronger London forces .
Boiling points of the noble gases

The boiling points of the noble gases illustrate this increase in strength of Van der Waal's forces with
molecular mass.
Noble gas
Neon
Argon
Krypton
Xenon
Radon

Molecular mass
20
40
84
131
222
Boiling point
-246
-186
-152
-108
-62
As the molecular mass of the gases increases, the atoms contain more electrons and so the size of the van der
Waal’s forces increases. As the attractive forces between the molecules increases, it becomes more difficult to
separate the molecules from each other and so the boiling points increase.
Permanent dipole-dipole bonding

Permanent dipole-dipole bonding occurs between polar molecules

It is stronger than van der waals and so the compounds have higher boiling points

Polar molecules have a permanent dipole. (commonly compounds with C-Cl, C-F, C-Br H-Cl, C=O bonds)

Polar molecules are asymmetrical and have a bond where there is a significant difference in electronegativity
between the atoms.
MISS CHOHANS NOTES ON BONDING UNIT 2 EDEXCEL

Permanent dipole bonding occurs in addition to London forces

It has been seen that in a molecule such as trichloromethane, CHCl3, the
electronic charge is pulled towards the chlorine atoms because they have a
greater electronegativity

This results in a molecule which has a partial positive charge at one end and
a partial negative charge at the other. This type of molecule is described as
polar.

The separation of charge which exists in a polar molecule is called a dipole.
In a polar material there is an attraction between the positive charge in one molecule and the negative charge in the other.
This type of force between molecules is referred to as
dipole-dipole attractions.
Hydrogen bonding

It occurs in compounds that have a hydrogen atom attached to one of the three most electronegative atoms of
nitrogen, oxygen and fluorine, which must have an available lone pair of electrons. e.g. a –O-H -N-H F- H bond.

There is a large electronegativity difference between the H and the O,N,F
Examples of hydrogen bonding

Always show the lone pair of electrons on the
O,F,N and the dipoles and all the δ - δ + charges

Hydrogen bonding occurs in addition to London
forces

The hydrogen bond should have a bond angle of
180o with one of the bonds in one of the
molecules
Hydrogen fluoride

The bond angle is 180O around the H atom
because there are two pairs of electrons around
the H atom involved in the hydrogen bond. These
pairs of electrons repel to a position of minimum
repulsion, as far apart as possible.
MISS CHOHANS NOTES ON BONDING UNIT 2 EDEXCEL
Hydrogen Bonding cont….

Alcohols, carboxylic acids, proteins, amides all can form hydrogen bonds

Hydrogen bonding is stronger than the other two types of intermolecular bonding.
Water

Water can form two hydrogen bonds per molecule, because the
electronegative oxygen atom has two lone pairs of electrons on it.

It can therefore form stronger hydrogen bonding and needs more
energy to break the bonds, leading to a higher boiling point.

The anomalously high boiling points of H2O,
NH3 and HF are caused by the hydrogen
bonding between these molecules in addition
to their London forces. The additional forces
require more energy to break and so have
higher boiling points

The general increase in boiling point from H2S
to H2Te or from HCl to HI is caused by
increasing London forces between molecules
due to an increasing number of electrons.
Intermolecular forces and physical properties
The greater the intermolecular force strength, the more difficult it is to separate the molecules and so the higher the
melting and boiling points.
The strengths of the intermolecular forces are summarized below.
The three molecules below illustrate the various strengths of these intermolecular forces.
Propane is composed of non-polar molecules and so only has van der Waals
forces between the molecules.
MISS CHOHANS NOTES ON BONDING UNIT 2 EDEXCEL
Methoxymethane has a similar molecular mass to propane, but
has a higher boiling point because the molecules are polar, so
there are dipole-dipole attractions as well as van der Waals
forces.
Ethanol has hydrogen bonding which is significantly stronger than
the other intermolecular forces and so a much higher boiling point
than methoxymethane even though they have the same molecular
mass.
Boiling Points of the Alkanes
A graph plotting the boiling points of the alkanes is shown below.
The boiling points of the alkanes increases with molecular mass. This
happens because the higher the molecular mass, the greater the number
of electrons and so the greater the chance of an imbalance and
formation of an instantaneous dipole. The van der Waals’ forces
(London forces) increase.
The melting point of the alkanes increases with molecular mass for the same reason, but the pattern is not so straightforward
as the different packing of molecules in the solid according to whether the number of carbon atoms is odd or even causes an
additional factor in the determination of the melting points.
The closer the molecules are able to approach each other, the greater the induction effect and so the greater the van der
Waals’ forces. When an alkane has branching present, the molecules cannot approach each other so closely and there is less
area over which contact can occur, so branched alkanes have lower boiling points.
The table below illustrates the effect of branching in alkane molecules on the boiling point.
MISS CHOHANS NOTES ON BONDING UNIT 2 EDEXCEL
Boiling Points of the alcohols
The –OH group in alcohols causes hydrogen bonding between the molecules. Therefore an alcohol will have a much higher
boiling point than an alkane with a similar number of electrons.
Boiling Points of the hydrogen halides
The boiling points of HCl, HBr and HI increase with molecular mass. This is because as the number of electrons increases
so does the chance of an electron imbalance and the formation of instantaneous dipoles and so greater van der Waals’ or
London forces.
It might be expected that HF having a smaller molecular mass than HCl would have a lower boiling point. This is not the
case. Hydrogen fluoride has the highest boiling point of this group because HF molecules form hydrogen bonds. The
graph plotting the boiling points of hydrogen halides is shown below.
MISS CHOHANS NOTES ON BONDING UNIT 2 EDEXCEL
Solvents and Solubility
Solubility of a solute in a solvent is a complicated balance of energy required to break bonds in the solute and solvent
against energy given out making new bonds between the solute and solvent.
Ionic substances dissolving in water
When an ionic lattice dissolves in water it
involves breaking up the bonds in the lattice and
forming new bonds between the metal ions and
water molecules.
The negative ions are attracted to the δ+
hydrogens on the polar water molecules and the
positive ions are attracted to the δ - oxygen on
the polar water molecules.
The higher the charge density the greater the hydration
enthalpy (e.g. smaller ions or ions with larger charges) as
the ions attract the water molecules more strongly.
In general a solvent will dissolve a substance that contains similar intermolecular forces.

When an ionic substance is placed in water, the water molecules, being highly polar, are attracted to the
ions.

The oxygen in the water molecule carries a partial negative charge and is attracted to cations. The
hydrogen in the water molecule carries a partial positive charge and is attracted to anions.

The process of water molecules linking to ions is called hydration of ions (as bonds are formed hydration
is always exothermic).

The water molecules are vibrating, so as they bond to the ions they shake the ions free from the lattice.
The process of dissolving is shown below.

Some ionic compounds do not dissolve in water because the electrostatic attraction between the ions, the
Lattice enthalpy, is too great for the water molecules to overcome.

To be soluble the energy produced by hydrating the ions (the hydration enthalpy) must be more negative
than the energy holding the ion together (the lattice enthalpy).
MISS CHOHANS NOTES ON BONDING UNIT 2 EDEXCEL
Solubility of simple alcohols
The smaller alcohols are soluble in water because they
can form hydrogen bonds with water. The longer the
hydrocarbon chain the less soluble the alcohol.
To be soluble in water the organic substance must be
able to form strong hydrogen bonds with the water
molecules.
Insolubility of compounds in water
Compounds that cannot form hydrogen bonds with water molecules, e.g. polar molecules such as halogenoalkanes or
non polar substances like hexane will be insoluble in water.
Solubility in non-aqueous solvents

To be soluble in a non-aqueous solvent the substance must have similar strength intermolecular forces to those
in the non-aqueous solvent.

Non-polar solvents will dissolve non-polar materials.

Solvents such as hexane will dissolve substances such as iodine. Hexane and iodine are both non-polar and so
have similar forces between their molecules (van der Waals forces). This means that the molecules are able to
interact with each other easily and allow solubility.

Non-polar solutes will dissolve in non-polar solvents. e.g. Iodine which has only London forces between its
molecules will dissolve in a non polar solvent such as hexane which also only has London forces.
 Propanone is a useful solvent because it has both polar and non polar characteristics. It can form London forces
with some non polar substances such as octane with its CH3 groups. Its polar C=O bond can also hydrogen bond
with water.
MISS CHOHANS NOTES ON BONDING UNIT 2 EDEXCEL

An ionic bond will show tendencies to have some degree of covalent character if:i)
ii)
iii)
iv)

The size of the cations is small
The charge on the cation is large
The anionic radius is large
The charge on the anion is large
E.g. LiI
Has a small
cationic radius
therefore a high
charge density

Has a large anionic radius
Li+
I-
The cation radius is very small therefore it has a high charge density so it will attract the electron
cloud of the large I- ion more towards itself rather than its own nucleus therefore the electron cloud of

the I- ion is distorted therefore this bond shows a degree of covalent character
When the electron cloud on a negative ion is distorted in this way we say that the negative ion is
POLARISED. In this case the I- ion is polarised, and that the Li+ causes this polarisation.


Hence the smaller the cationic radius the greater or larger the charge, the greater the charge density
therefore the more polarisation they cause
And the larger the anionic radius or the larger the charge on the anion the more polarisable they are.
MISS CHOHANS NOTES ON BONDING UNIT 2 EDEXCEL
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