2.2.2 (k,l) Types of Intermolecular Forces

Syllabus

(k) intermolecular forces based on permanent dipole–dipole interactions & induced dipole–dipole interactions

{Permanent dipole–dipole and induced dipole–dipole interactions can both be referred to as van der Waals’ forces.}

{Induced dipole–dipole interactions can also be referred to as London (dispersion) forces.}

(l) hydrogen bonding as intermolecular bonding between molecules containing N, O or F and the H atom of –NH, –OH or HF

{Including the role of lone pairs.}

What does this mean?

Ionic substances, metals and a few covalent substances exist in giant structures of billions of atoms.

But the majority of covalent substances consist of simple molecules of a few atoms.

So how does a glass of water stay liquid?

Why doesn't it immediately evaporate?

What holds one water molecule to its neighbours?

The GCSE answer is intermolecular forces.

Which is true - but now we're going to have to think about what causes them.

Permanent Dipole-Dipole Forces

You have already seen, or should have seen, how a molecule can end up polar.

By definition a polar molecule has a partly positive (δ+) and partly negative (δ-) end.

Physics tells us that positives & negatives attract electrostatically - which is equally true about δ+ & δ- charges.

So polar molecules attracted to each other by Permanent Dipole-Dipole forces.

Although there will also be other intermolecular forces acting on them.

Of course, these attractions are a lot weaker than covalent bonds.

And they extend in all directions, to all neighbouring molecules.

And the bigger the dipole the stronger the forces.

Strictly speaking, these forces are just one of van der Waals forces but textbooks often use this phrase to mean a different intermolecular force which all molecules have - even the non-polar ones.

Induced Dipole-Dipole forces.

When people talk about van der Waals forces - this is what they usually mean.

They are still electrostatic attractions between atoms or molecules but in this case between non-polar molecules.

But how can a non-polar molecule have a δ+ or δ- end?

Remember that electrons move randomly.

Some of the time there will be more electrons on one side of an atom than on the other.

This is a spontaneous dipole and it is not permanent.

Most times the electrons just even out again.

But, if a neighbouring atom is close-by (and there's not too much energy around), then the δ+ side of the atom could attract electrons in the neighbour towards itself.

Then the neighbour would have an opposite dipole - an induced dipole.

So, these induced van der Waals forces could be said to be:

"an electrostatic attraction between an atom or molecule with a spontaneous dipole and a neighbouring atom or molecule with an opposite, induced dipole".

Eventually the neighbour would induce an opposite dipole on its neighbours and they would do the same.

So, providing the temperature is low enough even a non-polar substance like Carbon Dioxide can solidify.

Note that all molecules will have van der Waals forces - even those with permanent dipole-dipole forces or Hydrogen bonds.

Note also that the larger a molecule the more electrons it will have an the stronger the van der Waals forces will be.

This is why large molecules usually have higher melting points.

Hydrogen bonds

Examiners are quite keen to penalise you for using the words bond and intermolecular forces interchangeably.

Which makes calling the strongest intermolecular force a Hydrogen bond inexcusable.

But we're stuck with it.

Most Hydrogen bonding questions focus on water.

If sked to draw Hydrogen bonds in water (or anything else), then a dotted line should connect a lone pair on the O atom to an H atom on the neighbouring molecule.

But you should know that any molecule where there is a Hydrogen atom bonded to either an Nitrogen, Oxygen, or Fluorine atom will produce Hydrogen bonding.

Notice that these three elements are the elements with the highest electronegativities and the smallest size.

(Chlorine has comparable electronegativity but is bigger - more shells - & doesn't Hydrogen bond.)

This combination makes N, O and F the most polarising elements.

They make the Hydrogens they are attached to very δ+ and these begin to attract the very compact lone-pair on the N, O or F atom.

(Lone pairs on Period 2 elements are more compact than on Chlorine - another reason why it doesn't Hydrogen bond)

So, we have a situation that sounds a lot like a dative bond where one atom (N, O or F) is donating a lone-pair to fill a space in another atom's valence shell.

But this isn't quite true because the valence shell of the H atom has been filled by the bond-pair - albeit that the bond is very polar so it has a smaller share of it than in most covalent bonds.

So, a Hydrogen bond is nowhere near as strong as a covalent bond - only around 10%.

But it is still the strongest type of intermolecular force.

And it still has a fixed direction - like a bond, and unlike other intermolecular forces which attract in all directions.

One of the traditional AS questions is:

"Draw a labelled diagram of the Hydrogen bonding in water (or Ammonia or ethanol etc) showing all relevant partial charges and lone pairs. Show the Hydrogen bond as a dotted line and label it."

Well done. You just got 3 marks.

One for drawing the Hydrogen bond between the H and the O

One for putting the δ+ & δ- on the correct atoms.

One for drawing the lone pairs on the Oxygen rather than the Hydrogen atoms. (Seriously!)

In Ammonia

In Ethanol

Between Ethanol & Water

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Exam-style Questions

1. The table below shows some values of melting points and some heat energies needed for melting.

(a) Name three types of inter-molecular forces

Force 1 …………………………………………………………………………..

Force 2 …………………………………………………………………………..

Force 3 …………………………………………………………………………..

(3)

(b) (i) Describe the bonding in a crystal of Iodine

……………………………………………………………………………………………………………………………………………………………………..................................................................................................................................……...................................…..

(ii) Name the crystal type which describes an Iodine crystal

………………………………............................................................................…………………………………................………….

(iii) Explain why heat energy is required to melt an Iodine crystal.

………………………………………………………………………………………………………………………………………………………………………….....................................................................................................................................................................……

(4)

(c) In terms of intermolecular forces involved, suggest why

(i) Hydrogen Fluoride requires more heat energy for melting than does of Hydrogen Chloride

…………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………

(ii) Hydrogen Iodide requires more heat energy for melting than does Hydrogen Chloride.

…………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………….................................................................................

(5)

2. The boiling temperatures, Tb, of some Group IV and Group V hydrides are given below

(a) The polarity of a Carbon-Hydrogen bond can be shown as

(i) What does the symbol d+ above the Hydrogen atom signify?

.............................................................................................................................................................................................................................................................................................................................................

(ii) Explain briefly, in terms of its shape, why a CH4 molecule has no overall polarity.

.............................................................................................................................................................................................................................................................................................................................................

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(3)

(b) Name the type of intermolecular forces which exist between CH4 molecules in liquid methane.

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(1)

(c) Explain why the boiling temperature of PH3 is greater than that of CH4.

.....................................................................................................................................................................................................................................................

...................................................................................................................................................................................................................................................(3)

(d) Explain why the boiling temperature of NH3 is greater than that of PH3

.......................................................................................................................................................................

......................................................................................................................................................................

(2)

(e) Suggest why the strength of the C–H bond in CH4 is greater than that of the Si–H bond in SiH4. State the relationship, if any, between the strength of the covalent bond in CH4 and the boiling temperature of CH4

Reason for stronger C-H bond

.......................................................................................................................................................................

Relationship between covalent bond strength and boiling temperature

......................................................................................................................................................................

.......................................................................................................................................................................

......................................................................................................................................................................

(2)

(Total 11 marks)

Answers

1. (a) Force 1: Van der Waals’ (1)

Force 2: dipole - dipole (1)

Force 3: hydrogen bonding (1)

OR London, Dispersion, temporary dipole

3

(b) (i) covalent between atoms (1)

OR within molecule

Van der Waals’ between molecules (1)

(ii) molecular.(1)

(iii) Bonds (or forces) between molecules must be broken or loosened (1)

OR VdW forces

OR intermolecular forces

Mention of ions CE=0

4

(c) (i) H-Bonding in HF (1)

(dipole-) dipole in HCl (1)

OR V.dW

H-bonding is stronger than dipole-dipole or V.dW.(1)

OR H-bonding is a strongest intermolecular force for 3rd mark

(ii) HI bigger molecule than HCl (1)

OR Heavier, more e’s, more electron shells, bigger Mr, more polarisable

Therefore the forces between HI molecules are stronger (1)

2.

(a) (i) deficiency of electrons (1)

(or small + recharge)

(ii) tetrahedral (1) (or symmetrical)

bond polarities cancel (1)

(or d + charges cancel out)

(b) van der Waals’ (1)

(c) PH3 has dipole-dipole (1)

between molecules (1)

stronger than in CH4 (1)

(inter molecular forces in PH3 > in CH4 scores Past mark)

(d) H-bonding in NH3 (1)

stronger than intermolecular forces in PH3 (1)

(e) Shorter (1) etc

none (1)

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