1. Types of intermolecular bonds
Key Learning
the relative strengths of bonds (covalent bonding, dispersion forces, dipole- dipole attraction and hydrogen bonding) and evidence and factors that determine bond strength including explanations for the floating of ice and expansion of water at higher temperatures.

2. Types of intermolecular force
Boardworks AS Chemistry
Bonding and Intermolecular Forces
Types of intermolecular force
The molecules in simple covalent substances are not entirely isolated from one another. There are forces of attraction between them. These are called intermolecular forces.
There are three main types of intermolecular force:
dispersion forces also known as van der Waals forces– for example, found between I2 molecules in iodine crystals.
permanent dipole–dipole forces – for example, found between HCl molecules in hydrogen chloride.
hydrogen bonds – for example, found between H2O molecules in water.

3. Permanent dipole–dipole forces
Boardworks AS Chemistry
Bonding and Intermolecular Forces
the relative strengths of bonds (covalent bonding, dispersion forces, dipole-dipole attraction and hydrogen bonding) and evidence and factors that determine bond strength including explanations for the floating of ice and expansion of water at higher temperatures.
Permanent dipole–dipole forces
If molecules contain bonds with a permanent dipole, the molecules may align so there is electrostatic attraction between the opposite charges on neighbouring molecules.
Permanent dipole–dipole forces (dotted lines) occur in hydrogen chloride (HCl) gas.
Teacher notes
Students could be reminded that molecules need to be polar, i.e. have an overall dipole, for permanent dipole–dipole forces to occur.
The permanent dipole–dipole forces are approximately one hundredth the strength of a covalent bond.

4. Permanent dipole–dipole or not?
Boardworks AS Chemistry
Bonding and Intermolecular Forces
Permanent dipole–dipole or not?
Teacher notes
Students could also be asked to categorize the following:
phosphine (PH3) = exhibits permanent dipole–dipole forces
hydrogen sulfide (H2S) = exhibits permanent dipole–dipole forces
chlorine (Cl2) = does not exhibit permanent dipole–dipole forces
ozone (O3) = does not exhibit permanent dipole–dipole forces

5. Hydrogen bonding Notice that in each of these molecules:
the relative strengths of bonds (covalent bonding, dispersion forces, dipole-dipole attraction and hydrogen bonding) and evidence and factors that determine bond strength including explanations for the floating of ice and expansion of water at higher temperatures.
Hydrogen bonding
Notice that in each of these molecules:
The hydrogen is attached directly to one of the most electronegative elements, causing the hydrogen to acquire a significant amount of positive charge.
Each of the elements to which the hydrogen is attached is not only significantly negative, but also has at least one lone pair.
Must be a Nitrogen, Oxygen or Fluorine bonded to Hydrogen to mean that Hydrogen bonds become possible

6. What is hydrogen bonding?
Boardworks AS Chemistry
Bonding and Intermolecular Forces
What is hydrogen bonding?
When hydrogen bonds to nitrogen, oxygen or fluorine, a larger dipole occurs than in other polar bonds.
This is because these atoms are highly electronegative due to their high core charge and small size. When these atoms bond to hydrogen, electrons are withdrawn from the H atom, making it slightly positive.
Teacher notes
Hydrogen bonds are about one tenth the strength of a covalent bond.
The H atom is very small so the positive charge is more concentrated, making it easier to link with other molecules.
Hydrogen bonds are therefore particularly strong examples of permanent dipole–dipole forces.

7. the relative strengths of bonds (covalent bonding, dispersion forces, dipole-dipole attraction and hydrogen bonding) and evidence and factors that determine bond strength including explanations for the floating of ice and expansion of water at higher temperatures.
Hydrogen bonding
The + hydrogen is so strongly attracted to the lone pair that it is almost as if you were beginning to form a covalent bond. It doesn't go that far, but the attraction is significantly stronger than an ordinary dipole-dipole interaction.
Hydrogen bonds have about a tenth of the strength of an average covalent bond and are about 10 times stronger than a dipole-dipole bond, and are being constantly broken and reformed in liquid water.
Each water molecule can potentially form four hydrogen bonds with surrounding water molecules, every one of them can be involved in hydrogen bonding.
This is why the boiling point of water is higher than that of ammonia or hydrogen fluoride.

8. Boardworks AS Chemistry Bonding and Intermolecular Forces
the relative strengths of bonds (covalent bonding, dispersion forces, dipole-dipole attraction and hydrogen bonding) and evidence and factors that determine bond strength including explanations for the floating of ice and expansion of water at higher temperatures.
Hydrogen bonding
hydrogen bond
lone pair
In molecules with OH or NH groups, a lone pair of electrons on nitrogen or oxygen is attracted to the slight positive charge on the hydrogen on a neighbouring molecule.
Teacher notes
See the ‘Alcohols’ presentation for more information about the structure and properties of alcohols.
Hydrogen bonding makes the melting and boiling points of water higher than might be expected. It also means that alcohols have much higher boiling points than alkanes of a similar size.

9. Hydrogen bonding and boiling points
Boardworks AS Chemistry
Bonding and Intermolecular Forces
Hydrogen bonding and boiling points

10. Boiling points of the hydrogen halides
Boardworks AS Chemistry
Bonding and Intermolecular Forces 20 40
-20
-40
-60
-80
-100
boiling point (°C) HF
HCl
HBr HI
The boiling point of hydrogen fluoride is much higher than that of other hydrogen halides, due to fluorine’s high electronegativity.
The means that hydrogen bonding between molecules of hydrogen fluoride is much stronger than the permanent dipole–dipole forces between molecules of other hydrogen halides. More energy is therefore required to separate the molecules of hydrogen fluoride.

11. Permanent dipole–dipole forces
Boardworks AS Chemistry
Bonding and Intermolecular Forces
Permanent dipole–dipole forces

12. How do we explain the existence of intermolecular forces in non polar substances? We know that intermolecular forces are present in non-polar substance because we see non-polar substances form liquids and solids. Without intermolecular forces there would be nothing to hold the molecules together and non-polar substances would only exist as gases. Lots of non- polar substances are liquids at room temperature (vegetable oil) and even non-polar solids (candle wax). And all non polar substances can form liquids or solids if cooled to low enough temperatures. So how to we explain these forces??

13. Boardworks AS Chemistry Bonding and Intermolecular Forces
Dispersion forces

14. Dispersion forces
Attractions are electrical in nature. In a symmetrical molecule like hydrogen, however, there doesn't seem to be any electrical distortion to produce positive or negative parts. But that's only true on average.
The diagram represents a small symmetrical molecule - H2, perhaps, or Br2. The even shading shows that on average there is no electrical distortion.
But the electrons are mobile, and at any one instant they might find themselves towards one end of the molecule, making that end -. The other end will be temporarily short of electrons and so becomes +.
An instant later the electrons may well have moved up to the other end, reversing the polarity of the molecule.
This constant moving around of the electrons in the molecule causes rapidly fluctuating dipoles even in the most symmetrical molecule. It even happens in monatomic molecules - molecules of noble gases, like helium, which consist of a single atom.

15. How do Dispersion Forces Work?
Imagine a molecule that has a temporary dipole alreday set up, is being approached by a molecule with no polarity.
As the right hand molecule approaches, its electrons will tend to be attracted by the slightly positive end of the left hand one.
This sets up an induced dipole in the approaching molecule, which is orientated in such a way that the + end of one is attracted to the - end of the other.
An instant later the electrons in the left hand molecule may well have moved up the other end. In doing so, they will repel the electrons in the right hand one.
The polarity of both molecules reverses, but you still have + attracting -. As long as the molecules stay close to each other the polarities will continue to fluctuate in synchronisation so that the attraction is always maintained.
This diagram shows how a whole lattice of molecules could be held together in a solid using van der Waals forces in one instant.

16. Strength of dispersion forces
Boardworks AS Chemistry
Bonding and Intermolecular Forces
the relative strengths of bonds (covalent bonding, dispersion forces, dipole-dipole attraction and hydrogen bonding) and evidence and factors that determine bond strength including explanations for the floating of ice and expansion of water at higher temperatures.
Strength of dispersion forces
The strength of dispersion forces increases as molecular size increases.
Atomic radius increases down the group, so the outer electrons become further from the nucleus. They are attracted less strongly by the nucleus and so temporary dipoles are easier to induce. 50
100
150
200
-50
-100
-150
-200
boiling point (°C)
Br2
This is illustrated by the boiling points of group 7 elements. F2
Cl2 I2
element

17. Strength of dispersion forces
Boardworks AS Chemistry
Bonding and Intermolecular Forces
the relative strengths of bonds (covalent bonding, dispersion forces, dipole-dipole attraction and hydrogen bonding) and evidence and factors that determine bond strength including explanations for the floating of ice and expansion of water at higher temperatures.
Strength of dispersion forces
The points of contact between molecules also affects the strength of dispersion forces.
butane (C4H10)
boiling point = 272 K
2-methylpropane (C4H10)
boiling point = 261 K
Straight chain alkanes can pack closer together than branched alkanes, creating more points of contact between molecules. This results in stronger van der Waals forces.

18. Boiling points of alkanes
Boardworks AS Chemistry
Bonding and Intermolecular Forces
Boiling points of alkanes