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AS Chem Unit 2.1 Shapes of molecules, polarity & intermolecular forces

determining number of lone electron pairs, molecule shape, VSEPR theory, bond + molecule polarity, intermolecular forces and affect on physical properties

J Amuah-Fuster

on 8 July 2016

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Transcript of AS Chem Unit 2.1 Shapes of molecules, polarity & intermolecular forces


lone pairs
repel to a greater degree
than bonding pairs

So this bond angle is
drastically from the tetrahedral 109.5 to >

Cl-Si-Cl bond angle ....
~109.5 degrees


Cl-P-Cl bond angle ....
~107 degrees


Cl-S-Cl bond angle ....
~180 degrees
Hydrogen bonding
a. Demonstrate an understanding of the nature of intermolecular forces resulting from interactions between permanent dipoles, instantaneous dipoles and induced dipoles (London forces) and from the formation of hydrogen bonds.
b. Relate the physical properties of materials to the types of intermolecular force present, eg:
i) The trends in boiling and melting temperatures of alkanes with increasing chain length.
ii) The effect of branching in the carbon chain on the boiling and melting temperatures of alkanes.
iii) The relatively low volatility (higher boiling temperatures) of alcohols compares to alkanes with a similar number of electrons.
iv) The trends in boiling temperatures of the hydrogen halides HF to HI.
c. Carry out experiments to study the solubility of simple molecules in different solvents.
d. Interpret given information about solvents and solubility to explain the choice of solvents in given contexts, discussing the factors that determine the solubility, including:
i) The solubility of ionic compounds in water in terms of the hydration of the ions.
ii) The water solubility of simple alcohols in terms of hydrogen bonding.
iii) The insolubility of compounds that cannot form hydrogen bonds with water molecules, eg polar molecules such as halogenoalkanes.
iv) The solubility in non-aqueous solvents of compounds which have similar intermolecular forces to those in the solvent.
Molecule Shapes
Intermolecular Forces
Calculating electron pairs
What have you learnt?
Practice questions
Molecule Shapes
Draw dot and cross diagrams for the following molecules:
carbon dioxide
Electrons, where possible, move around in pairs.
These pairs may be
bonded pairs
(involved in bonding), or
lone pairs
(electrons belonging and residing next to the atom).
The number of bonded pairs can be found by observing how many sigma-bonds (single bonds) are in the molecule.
The number of lone pairs are discovered by subtracting the number of valence electrons involved in bonding from the group number of the element, and then dividing that sub-total by two.

For example, in methane (CH4), the central carbon atom is bonded to four hydrogen atoms. Carbon is in group 4, and all of its 4 valence electrons are involved in bonding. Therefore, the number of electrons remaining is zero, which divided by two, results in zero lone pairs around the carbon atom.
# lone pairs =
(group number - # bonding electrons)
Identify the number of
bonding pairs
lone pairs
about the
central atom
in these molecules:
carbon dioxide
boron trifluoride
an ammonium

Demonstrate an understanding of the use of electron-pair repulsion theory to interpret and predict the shapes of simple molecules and ions
Recall and explain the shapes of BeCl2, BCl3, CH4, NH3, NH4+, H2O, CO2, gaseous PCl5 and SF6 and the simple organic molecules listed in Units 1 and 2
Apply the electron-pair repulsion theory to predict the shapes of molecules and ions analogous to those in 2.3b
Demonstrate an understanding of the terms bond length and bond angle and predict approximate bond angles in simple molecules and ions
VSEPR theory
V alency
S hell
E lectron
P air
R epulsion
The shapes of molecules are affected by the number of electron
around the central atom.
The electron
regions may be

made up of

bonding pairs

lone pairs
Electron pairs


each other to an

angle of maximum separation

minimum repulsion
Electron pairs of


nature will repel each other with


States that pairs of electrons arrange themselves around the central atom so that they are as far apart from each other as possible. Greater repulsion between lone pair of electrons than bonded pairs.
Let's take
as an example;
it has
2 bonding pairs
2 lone pairs
It is a
of the
shape made from 4 bonding electron regions (like in methane).
1. Silicon, phosphorus and sulphur form chlorides with molecular formulae SiCl4, PCl3, SCl2.
Draw the shapes you would expect for these molecules, suggesting a value for the bond angle in each case.
SiCl4 Cl-Si-Cl bond angle ...............
PCl3 Cl-P-Cl bond angle ...............
SCl2 Cl-S-Cl bond angle ............... (Total 3 marks)

2. (a) Explain why a water molecule does not have a linear shape. (2)
(b) State the HOH bond angle in water and explain why it has this
value. (2)

3. Phosphorus reacts with a limited amount of chlorine to produce
phosphorus trichloride, PCl3.

(a) Draw a dot and cross diagram to show the arrangement of the
electrons in phosphorus trichloride, PCl3. You need only show
the outer shell electrons. (2)

(b) Draw the phosphorus trichloride molecule, making its three-
dimensional shape clear. (1)

(c) Explain:
(i) the shape of the phosphorus trichloride molecule, and
(ii) why the Cl-P-Cl bond angle is different from the H-C-H
bond angle in methane, CH4. (3)
I understand how to use VSEPR to interpret and predict the shapes of simple molecules an ions
I can state the shapes of BeCl2, BCl3, CH4, NH3, NH4+, H2O, CO2, gaseous PCl5 and SF6, as well as simple organic molecules listed in Unit 1 and Unit 2
I can demonstrate that I understand the terms bond length and bond to predict approximate bond angles in simple molecules and ions
I can apply the VSEPR theory to predict the shapes of molecules and ions analogous to those listed in objective 2
Instantaneous dipole-dipole
Permanent dipole-dipole
polar molecules
bond strength =
10 kJ/mol (approx.)
bond strength =
1.5 kJ/mol (approx.)
occurs between all types of molecules
bond strength =
20 kJ/mol (approx.)
between sufficiently delta+ H and delta-
, or
atom (bonded to a H atom)
type of intermolecular force
Force proportional to
size of electron cloud
number of contact points
between molecules
Bond strength =
+400 kJ/mol
to bond
ionic bonding
Bond energy
+250 kJ/mol
Mg, Na, Cu
to ionic
and ionic
F = q q
(r + r )
Metallic bonding
Bond strength=
+100 kJ/mol (approx.)
H + H O H O
Non-polar molecules have a symmetrical dipole moment
Polar molecules have an
dipole moments
electronegativity increases along period & up group
molecular Force
Polar & Non-polar molecules
... small
Shapes of molecules & Intermolecular forces
carbon dioxide
sigma bonds are single covalent bonds formed by the overlap of s and/or p orbitals, so they may share a pair of electrons.
2 bonding pairs, 2 lone pairs
3 bonding pairs, 1 lone pair
Carbon dioxide:
4 bonding pairs, 0 lone pairs
Boron trifluoride:
3 bonding pairs, 0 lone pairs
4 bonding pairs, 0 lone pairs
(hint: it can help to draw the dot and cross diagram for the molecule)
End of topic summary questions
Lesson Objectives
How does a gecko stop itself from falling down???
4-nitro phenol has
H bonds, leading to stronger forces of interaction between molecules and so its BP is 279°C
Compare H O and H S
Anomalous behaviour of Water
Examples of H-Bonding contd….
Trends in Boiling points of Groups 4,5,6 & 7
Examples of H-Bonding
Types of Intermolecular forces
Kinds of molecules
Forces between ions, atoms and molecules
Biological importance of H-bonding
Propanone is miscible in water but it has no H-bonding between propanone molecules
This is because its O atoms can H-bond to water molecules.
Effect of H bonding on Organic compounds
Case 2
Dimer of Ethanoic acid
Effect of H bonding on Organic compounds
Case 1
Hydrogen Bonds in Water and Ice
Structure of Ice
Examples of H-Bonding contd….
Examples of H-Bonding contd….
Hydrogen Bonding in different compounds
Hδ+—Fδ- ---------------Hδ+—Fδ-
Hydrogen bonding
Examples to explain dipole-dipole forces
Examples to explain Van der waals’ forces
Proof for existence of intermolecular forces
A survey of the boiling temperatures of the hydrides of Group 4, 5, 6 and 7 clearly illustrates the significance of hydrogen bonding on the properties of the molecules.
Hydrogen bonds and physical properties of substances.
The molar enthalpy of vaporisation of a liquid is the enthalpy change when one mole of the liquid changes into one mole of its gas at the boiling temperature – this is a direct measure of the intermolecular forces in the liquid.
Molar enthalpy of vaporisation
2 H bonds between Adenine and Thymine
3 H bonds between Cytosine and Guanine
Hydrogen bonds in DNA Base Pairs
You may have wondered why ice floats?
This happens because more hydrogen bonds form in ice than in liquid water, and since each hydrogen bond has a bond length, the water molecules have to spread out in order for the hydrogen bonds to form. This makes ice less dense.
Furthermore, the presence of the hydrogen bonds makes ice very strong – and in large amounts, strong enough to sink a ship!
Liquid water and ice
Hydrogen bonds form because of the special nature of the hydrogen atom, which has no inner shells of electrons. As a result, the nucleus of the hydrogen atom in these molecules is left unusually exposed by the shift in electron density within the bond, making it easily accessible for strong permanent dipole – permanent dipole interactions to occur.
The reason why it is called a bond is because it is a relatively strong intermolecular force (although not as strong as a ionic or covalent bond).
Molecules that can attract each other by this form of bonding have a common feature.
They all have either Nitrogen, Oxygen or Fluorine bonded to hydrogen
Hydrogen bonding
London forces come into their own when it comes to giant molecules such as polymers.
the chains are held together by the millions of London forces.
London forces and polymers
Molecules with
permanent dipoles
can attract neighbouring molecules.
Permanent dipole - permanent dipole forces
All compounds containing an –OH group form hydrogen bonds.
Undoubtedly, water is the most important.
When a covalent substance is boiled
only the intermolecular forces are broken.
An increase in the number of hydrogen bonds in icy water causes the molecules to spread apart. Heating water provides enough energy to over come the hydrogen bonds allowing water molecules to evaporate.
In the exam you will be asked to draw two water molecules and show the hydrogen bonding between them. This is how your diagram should be presented. You must show the charges on the atoms, label the hydrogen bond and draw the two molecules to show that the bond angle is 180o.
The bond angle for a hydrogen bond is 180 degrees.
Hydrogen bond angle
Both these are isomers of pentane, C5H12. The shapes of the molecules determines the number of London forces that can form and hence the physical properties of the molecules.
The collective strength of London forces between molecules also depends on the
number of points of contacts
There are more points of contact between straight chain molecules than between molecules with branches.
Strength of London forces
Iodine atoms are much bigger than the halogens above it in Group 7. Therefore, the London forces between molecules is stronger and hence iodine is a solid.
The size of the London forces in a substance depends on the size of the electron clouds of the particles that are interacting.
Atoms with
large electron clouds hold onto their outer electrons less strongly
so these
electron clouds are easily deformed
This favours the creation of London forces, therefore
increasing the force of attraction between the molecules
Strength of London forces
Hydrogen bonding explains why water and alcohols such as methanol and ethanol are miscible (mix).
Examples of hydrogen bonding
Van de Waals
Fritz London
These are also known as dispersion forces, van der Waals forces or
London forces
. We shall call them London forces.
Instantaneous or temporary dipole - induced dipole forces
Why does chlorine not form hydrogen bonds?
N, O, and F are three of the most electronegative (and are small atoms) elements in the Periodic Table. When they are bonded with hydrogen, they form a polar bond.
The slightly positive hydrogen on one molecule is attracted to a slightly negative N, O or F on another molecule.
Why only N, O, and F?
Considerably stronger than other intermolecular forces.
Affects the physical properties of the compounds in which it exists.
hydrogen bond
DNA pairing occurs due to H bonds.
Secondary structure of proteins (alpha-αhelix and beta-pleated proteins).
The alpha helix is an example of INTRA MOLECULAR H-Bonding
Effect of H bonding on Organic compounds
Case 3
2-nitro phenol has
H bonds, leading to weaker forces of interaction between molecules and so has lower BP of 216°C
- - - -
- - - -
Electronegativity data sheet
Summary of bonding types
a. explain the meaning of the term electronegativity as applied to atoms in a covalent bond
b. recall that ionic and covalent bonding are the extremes of a continuum of bonding type and explain this in terms of electronegativity differences leading to bond polarity in bonds and molecules, and to ionic bonding if the electronegativity is large enough
c. distinguish between polar bonds and polar molecules and be able to predict whether or not a given molecule is likely to be polar
d. carry out experiments to determine the effect of an electrostatic force on jets of liquids and use the results to determine whether the molecules are polar or non-polar.
Polar bonds and polar molecules
Polarity Questions
Ionic Character Questions
Water is one of the most talked about polar molecules. The O-H bond is polar because of the difference in the electronegativities of the two elements.
The molecule itself is polar because the shape of the molecule gives it a dipole – the slightly negative oxygen at one end and the two slightly positive hydrogen atoms at the other end.
Polar molecules
Electronegativity Questions
The charge separation in a polar molecule makes it a dipole (there are two types or poles of charge in the molecule).
The polarity is measured as its dipole moment – the amount of charge separation multiplied by the distance between the centres of charge. The unit of dipole moments is the debye, D.
Measuring polarity
Carbon dioxide does have two polar bonds, however, the dipoles cancel out which makes the molecule non-polar.
Other examples
The direction of the arrow is always towards the more electronegative atom which attracts the bonding pairs of electrons.
Examples of polar bonds
The Greek lowercase letter  delta, means ‘small’
The difference in electronegativities is 0.9.
If there is a difference in the electronegativities of the two atoms in the bond then the bonding electrons are shared unequally.
The bonding pair of electrons will be pulled towards the more electronegative atom.
The more electronegative atom then has more electrons than protons giving it a slight negative charge while the less electronegative atom becomes slightly positive.
Polar covalent bonds
F has
small atomic radii
high nuclear charge so can attract electrons more.
Why does F have the greatest electronegativity?
The power of an atom to attract the bonding pair of electrons in a covalent bond.

Linus Pauling produced a relative scale from 0 to 4 – the higher the number, the greater the electronegativity.
The same would be true if polar molecules were placed in an electric field.
A simple test for polarity, is to produce a stream of liquid and hold a charged rod or balloon next to it.
If the stream bends then the liquid is made of polar molecules.
Testing for polarity
Why would a molecule with polar bonds not necessarily be polar itself?
You have already learned about polar bonds, produced as a result of the difference in the electronegativities of the two atoms in the bond.
A polar molecule is one which has a region of slightly negative charge and a region of slightly positive charge (a charge separation called a dipole).
A polarity of a molecule is obviously a direct result of having polar bonds, but not all molecules with polar bonds are polar molecules.
Polarity of molecules
The equal sharing of electrons only occurs if the two atoms have the same electronegativities., i.e., they are the atoms of the same element.
The electron clouds around the nuclei are distributed equally.
In an ideal covalent bond the bonding pair of electrons should be shared equally between the two atoms.
An ideal covalent bond
Does carbon dioxide have polar bonds, and if so, is it a polar molecule?
A polar molecule must have a defined dipole.
If the polar bonds in a molecule cancel – usually when there is symmetry in the arrangement of the polar bonds, then the molecule will not be polar.
Polar molecules
Summary of bonding types
Learning objectives
1. HF has the most ionic character

3. Lithium Iodide has 45% ionic character. Compared to potassium fluoride (93% ionic character), lithium iodide will have a lower melting and boiling point, and a lower solubility in water, but it will have a higher lattice enthalpy.
1. The next three most electronegative elements after fluorine are oxygen, nitrogen and chlorine.

2. The trend in electronegativity from Sc to Zn shows a general increase with the exception of drops found at Mn and Zn.

3. The slightly higher electron density around the carbon atom indicates that the C-H bond is mainly
very slight


, with carbon as delta negative and H as delta positive.
This is due to very small differences in electronegativities of carbon and hydrogen
1. a) Non-polar
b) Polar
c) Non-polar
d) Polar

2. A molecule with polar bonds may not be a polar if the
dipole moment of the polar bond is cancelled out
by other polar bonds with the same polarity. This arises due to symmetrical dipole moments, e.g. carbon dioxide has two polar C=O bonds, but the
symmetrically opposite nature
of these bonds means the dipoles are cancelled out, leaving a non-polar molecule.

3. All three compounds (chloromethane, dichloromethane & tetrachloromethane) have polar bonds, which should imply polarity in the molecule. However, in the example of tetrachloromethane, the
polarity of the C-Cl bond is cancelled out due to overall symmetry of dipole moments
. The first two examples have
asymmetric dipole moments

and therefore maintain polarity
Electronegativity data sheet
to ionic






metal ion
delocalised electrons
Energy produced affects the
of a substance
atom and
2 2
Permanent dipole-dipole
London forces (van der Waals)
Instantaneous dipole-dipole
Dispersion forces
Hydrogen bonding
(i) Phosphorus trichloride has a triganol pyramidal shape due to 3 bonding electron pairs and 1 lone pair of electrons. The forces between the electron pairs creates a shape that allows the maximum angle between the electron pairs while causing minimum repulsion.

(ii) The Cl-P-Cl bond angle in PCl3 is smaller than (107 degrees) the H-C-H bond angle (109.5) in methane, due to the greater level of repulsion between lone pairs and bonding pairs than that between bonding pairs alone. As a result, the bonding electron pairs are forced closer together, reducing the typical bond angle.
Question 3

Question 3
(a) The oxygen atom in the water molecule has 2 lone pairs of electrons as well as 2 bonding pairs. Due to the repulsion between these 4 electron pairs, the shape of water is a v-shape; based on a tetrahedral structure.

(b) The H-O-H bond angle in water is about 104.5 degrees, this is due to the greater repulsion between nearby lone pairs and more distant bonding pairs, than the repulsion between bonding pairs. Hence the bond angle is reduced from a tetrahedral angle of 109.5 to 104.5.
Question 2
Question 1
An example
As a rule of thumb,
'.. for
each lone pair
that replaces a bonding pair,
the standard bond
angle by 2.5
According to
lone pairs exert a greater degree of repulsion than bonding pairs
, because they are
closer to the nucleus
of the central atom.
Making Molecule shapes
Bond energy

H -Cl
I use Pritt Stick innit!
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