AS Chemistry
Unit 1: Theoretical Chemistry
Part 13: Covalent Bonding
A covalent bond is made when atoms share one or more electrons to form a
molecule. Covalent bonds are usually formed between pairs of non-metallic
elements.
A single covalent bond is made when each atom donates one electron to the
bond. It is also possible to form double and triple bonds where two and three
electrons are donated. As a general rule, the number shared gives each atom
filled outer shells similar to the electronic configuration of a noble gas.
Some very simple covalent molecules
Hydrogen
Imagine if you will, two hydrogen atoms approaching each other from space.
The two atoms come close together so that their outer atomic orbitals overlap
(to form a molecular orbital). Both nuclei are attracted to the shared pair of
electrons and this attraction binds the atoms together.
or H - H
AS Chemistry
Unit 1: Theoretical Chemistry
Task 2
Use dot-cross diagrams to show the covalent bonding in the following molecules:
Cl2,
HCl,
O 2,
CH4, NH3, CO2, C2H4,
Draw your diagrams on the worksheet ‘Electron dot-cross diagrams – covalent
compounds 1’. You may discuss your answers with your partner. In each case
show only the outer electrons.
Lone pairs
Atoms in molecules frequently have pairs of electrons in their outer shells that
are not involved in covalent bonds. These non-bonding electron-pairs are called
lone-pairs.
Covalent bonding at A-level
Cases where there isn't any difference from the simple view
If you stick closely to the A-level syllabus, there is little need to move far from
the simple (IGCSE) view. The only thing which must be changed is the overreliance on the concept of noble gas structures. Most of the simple molecules
you draw do in fact have all their atoms with noble gas structures.
Even with a more complicated molecule like PCl3, there's no problem. In this
case, only the outer electrons are shown for simplicity. Each atom in this
structure has inner layers of electrons of 2,8. Again, everything present has a
noble gas structure.
Cases where the simple view throws up problems
Boron trifluoride, BF3
AS Chemistry
Unit 1: Theoretical Chemistry
A boron atom only has 3 electrons in its outer level, and there is no possibility
of it reaching a noble gas structure by simple sharing of electrons. Is this a
problem? No. The boron has formed the maximum number of bonds that it can
in the circumstances, and this is a perfectly valid structure.
Energy is released whenever a covalent bond is formed. Because energy is being
lost from the system, it becomes more stable after every covalent bond is
made. It follows, therefore, that an atom will tend to make as many covalent
bonds as possible. In the case of boron in BF3, three bonds is the maximum
possible because boron only has 3 electrons to share.
Phosphorus(V) chloride, PCl5
In the case of phosphorus 5 covalent bonds are possible - as in PCl5.
Phosphorus forms two chlorides - PCl3 and PCl5. When phosphorus burns in
chlorine both are formed - the majority product depending on how much
chlorine is available. We've already looked at the structure of PCl3.
The diagram of PCl5 (like the previous diagram of PCl3) shows only the outer
electrons.
Notice that the phosphorus now has 5 pairs of electrons in the outer level certainly not a noble gas structure. You would have been content to draw PCl3 at
GCSE, but PCl5 would have looked very worrying.
Why does phosphorus sometimes break away from a noble gas structure and
form five bonds? In order to answer that question, we need to explore territory
AS Chemistry
Unit 1: Theoretical Chemistry
beyond the limits of the A-level syllabus. Don't be put off by this! It isn't
particularly difficult, and is extremely useful if you are going to understand the
bonding in some important organic compounds.
A more sophisticated view of covalent bonding
The bonding in methane, CH4
What is wrong with the dots-and-crosses picture of bonding in methane?
We are starting with methane because it is the simplest case which illustrates
the sort of processes involved. You will remember that the dots-and-crossed
picture of methane looks like this.
There is a serious mis-match between this structure and the modern electronic
structure of carbon, 1s22s22px12py1. The modern structure shows that there are
only 2 unpaired electrons for hydrogens to share with,
instead of the 4 which the simple view requires.
You can see this more readily using the electrons-in-boxes
notation. Only the 2-level electrons are shown. The 1s2
electrons are too deep inside the atom to be involved in
bonding. The only electrons directly available for sharing
are the 2p electrons. Why then isn't methane CH2?
Promotion of an electron
When bonds are formed, energy is released and the
system becomes more stable. If carbon forms 4 bonds
rather than 2, twice as much energy is released and so the
resulting molecule becomes even more stable.
There is only a small energy gap between the 2s and 2p orbitals, and so it pays
the carbon to provide a small amount of energy to promote an electron from the
2s to the empty 2p to give 4 unpaired electrons. The extra energy released
when the bonds form more than compensates for the initial input.
AS Chemistry
Unit 1: Theoretical Chemistry
Hybridisation
The electrons rearrange themselves again in a process
called hybridisation. This reorganises the electrons into
four identical hybrid orbitals called sp3 hybrids (because
they are made from one s orbital and three p orbitals).
You should read "sp3" as "s p three" - not as "s p cubed".
sp3 hybrid orbitals look a bit like half a p orbital, and they arrange themselves in
space so that they are as far apart as possible. You can
picture the nucleus as being at the centre of a
tetrahedron (a triangularly based pyramid) with the
orbitals pointing to the corners. For clarity, the nucleus
is
drawn far larger than it really is.
What happens when the bonds are formed?
Remember that hydrogen's electron is in a 1s orbital - a spherically symmetric
region of space surrounding the nucleus where there is some fixed chance (say
95%) of finding the electron. When a covalent bond is formed, the atomic
orbitals (the orbitals in the individual atoms) merge to produce a new molecular
orbital which contains the electron pair which creates the bond.
Four molecular orbitals are formed, looking rather like the original sp3 hybrids,
but with a hydrogen nucleus embedded in each lobe. Each orbital holds the 2
electrons that we've previously drawn as a dot and a cross. The principles
involved - promotion of electrons if necessary, then hybridisation, followed by
the formation of molecular orbitals - can be applied to any covalently-bound
molecule.
The bonding in the phosphorus chlorides, PCl3 and PCl5
AS Chemistry
Unit 1: Theoretical Chemistry
What's wrong with the simple view of PCl3?
This diagram only shows the outer (bonding) electrons.
Nothing is wrong with this! (Although it doesn't account for the shape of the
molecule properly.) If you were going to take a more modern look at it, the
argument would go like this:
Phosphorus has the electronic structure 1s22s22p63s23px13py13pz1. If we look
only at the outer electrons as "electrons-in-boxes":
There are 3 unpaired electrons that can be used to form bonds with 3 chlorine
atoms. The four 3-level orbitals hybridise to produce 4 equivalent sp3 hybrids
just like in carbon - except that one of these hybrid orbitals contains a lone pair
of electrons.
Each of the 3 chlorines then forms a covalent bond by merging the atomic
orbital containing its unpaired electron with one of the phosphorus unpaired
electrons to make 3 molecular orbitals.
You might wonder whether all this is worth the bother! Probably not! It is worth
it with PCl5, though.
AS Chemistry
Unit 1: Theoretical Chemistry
What's wrong with the simple view of PCl5?
You will remember that the dots-and-crosses picture of PCl5 looks awkward
because the phosphorus doesn't end up with a noble gas structure. This diagram
also shows only the outer electrons.
In this case, a more modern view makes things look better by abandoning any
pretence of worrying about noble gas structures.
If the phosphorus is going to form PCl5 it has first to generate 5 unpaired
electrons. It does this by promoting one of the electrons in the 3s orbital to
the next available higher energy orbital.
Which higher energy orbital? It uses one of the 3d orbitals. You might have
expected it to use the 4s orbital because this is the orbital that fills before
the 3d when atoms are being built from scratch. Not so! Apart from when you
are building the atoms in the first place, the 3d always counts as the lower
energy orbital.
This leaves the phosphorus with this arrangement of its electrons:
AS Chemistry
Unit 1: Theoretical Chemistry
The 3-level electrons now rearrange (hybridise) themselves to give 5 hybrid
orbitals, all of equal energy. They would be called sp3d hybrids because that's
what they are made from.
The electrons in each of these orbitals would then share space with electrons
from five chlorines to make five new molecular orbitals - and hence five
covalent bonds. Why does phosphorus form these extra two bonds? It puts in an
amount of energy to promote an electron, which is more than paid back when the
new bonds form. Put simply, it is energetically profitable for the phosphorus to
form the extra bonds.
The advantage of thinking of it in this way is that it completely ignores the
question of whether you've got a noble gas structure, and so you don't worry
about it.
A non-existent compound - NCl5
Nitrogen is in the same Group of the Periodic Table as phosphorus, and you
might expect it to form a similar range of compounds. In fact, it doesn't. For
example, the compound NCl3 exists, but there is no such thing as NCl5.
Nitrogen is 1s22s22px12py12pz1. The reason that NCl5 doesn't exist is that in
order to form five bonds, the nitrogen would have to promote one of its 2s
electrons. The problem is that there aren't any 2d orbitals to promote an
electron into - and the energy gap to the next level (the 3s) is far too great.
In this case, then, the energy released when the extra bonds are made isn't
enough to compensate for the energy needed to promote an electron - and so
that promotion doesn't happen.
Atoms will form as many bonds as possible provided it is energetically
profitable.
Learning Objectives:
Candidates should be able to describe, including the use of ‘dot-and-cross
diagrams, covalent bonding, as in hydrogen; oxygen; chlorine; hydrogen chloride;
carbon dioxide; methane and ethene.
AS Chemistry
Unit 1: Theoretical Chemistry
References:
A-level Chemistry: pages 33-36
Chemistry in Context: pages 86-88
AS Chemistry
Unit 1: Theoretical Chemistry
Part 14: Co-ordinate (dative covalent) bonding
Task 1
Can you complete the worksheet ‘Electron dot-cross diagrams – covalent
compounds 2’.
A covalent bond is formed by two atoms sharing a pair of electrons. The atoms
are held together because the electron pair is attracted by both of the nuclei.
In the formation of a simple covalent bond, each atom supplies one electron to
the bond - but that doesn't have to be the case.
A co-ordinate bond (also called a dative covalent bond) is a covalent bond (a
shared pair of electrons) in which both electrons come from the same atom.
The reaction between ammonia and hydrogen chloride
You will have seen the experimental set up during your IGCSE course:
As the colourless gases mix, a thick white smoke of solid ammonium chloride is
formed.
Ammonium ions, NH4+, are formed by the transfer of a hydrogen ion from the
hydrogen chloride to the lone pair of electrons on the ammonia molecule.
AS Chemistry
Unit 1: Theoretical Chemistry
When the ammonium ion, NH4+, is formed, the fourth hydrogen is attached by a
dative covalent bond, because only the hydrogen's nucleus is transferred from
the chlorine to the nitrogen. The hydrogen's electron is left behind on the
chlorine to form a negative chloride ion.
Once the ammonium ion has been formed it is impossible to tell any difference
between the dative covalent and the ordinary covalent bonds. Although the
electrons are shown differently in the diagram, there is no difference between
them in reality.
Representing co-ordinate bonds
In simple diagrams, a co-ordinate bond is shown by an arrow. The arrow points
from the atom donating the lone pair to the atom accepting it.
Dissolving hydrogen chloride in water to make hydrochloric acid
Something similar happens. A hydrogen ion (H+) is transferred from the chlorine
to one of the lone pairs on the oxygen atom.
The H3O+ ion is variously called the hydroxonium ion, the hydronium ion or, in
the case of your syllabus, the oxonium ion.
In an introductory chemistry course (such as IGCSE), whenever you have talked
about hydrogen ions in solution (for example in acids), you have actually been
talking about the oxonium ion. A raw hydrogen ion is simply a proton, and is far
too reactive to exist on its own in a test tube.
AS Chemistry
Unit 1: Theoretical Chemistry
If you write the hydrogen ion as H+(aq), the "(aq)" represents the water molecule
that the hydrogen ion is attached to. When it reacts with something (an alkali,
for example), the hydrogen ion simply becomes detached from the water
molecule again.
Note that once the co-ordinate bond has been set up, all the hydrogens
attached to the oxygen are exactly equivalent. When a hydrogen ion breaks
away again, it could be any of the three.
The reaction between ammonia and boron trifluoride, BF3
You may remember from our previous lesson that boron trifluoride is a
compound which doesn't have a noble gas structure around the boron atom. The
boron only has 3 pairs of electrons in its bonding level, whereas there would be
room for 4 pairs. BF3 is described as being electron deficient. (Although in the
words of Professor D.M.P. Mingos (Principal of St. Edmund Hall, Oxford), “There
are no such things as electron-deficient compounds, only theory-deficient
chemists.”)
Task 2
Boron trifluoride, BF3, and an ammonia molecule, can combine together to form
the molecule NH3BF3. This molecule has a dative covalent bond between the
nitrogen atom and the boron atom. Can you draw the dot-cross diagram for this
molecule?
How could this molecule be drawn more simply using displayed formula?
AS Chemistry
Unit 1: Theoretical Chemistry
The structure of aluminium chloride
Aluminium chloride sublimes (turns straight from a solid to a gas) at about
180°C. If it simply contained ions it would have a very high melting and boiling
point because of the strong attractions between the positive and negative ions.
The implication is that it when it sublimes at this relatively low
temperature, it must be covalent. The dots-and-crosses
diagram shows only the outer electrons.
AlCl3, like BF3, is electron deficient. There is likely to be a
similarity, because aluminium and boron are in the same group
of the Periodic Table, as are fluorine and chlorine.
Measurements of the relative formula mass of aluminium chloride show that its
formula in the vapour at the sublimation temperature is not AlCl3, but Al2Cl6. It
exists as a dimer (two molecules joined together). The bonding between the two
molecules is co-ordinate, using lone pairs on the chlorine atoms. Each chlorine
atom has 3 lone pairs, but only the two important ones are shown in the line
diagram.
Task 3
Can you represent the molecule above using displayed formula? Only the
important lone pairs on each chlorine atom should be shown.
AS Chemistry
Unit 1: Theoretical Chemistry
Learning Objectives:
Candidates should be able to describe, including the use of ‘dot-and-cross
diagrams, co-ordinate (dative covalent) bonding, as in the formation of the
ammonium ion and in the Al2Cl6 molecule.
References:
A-level Chemistry: pages 35-36
Chemistry in Context: pages 88-90
AS Chemistry
Unit 1: Theoretical Chemistry
Part 15: Electronegativity
What is electronegativity?
Definition
Electronegativity is a measure of the tendency of an atom to attract a bonding
pair of electrons.
Electronegativity values cannot be measured directly; they are calculated using
a number of different methods. The Pauling scale is the most commonly used.
Fluorine (the most electronegative element) is assigned a value of 4.0, and
values range down to caesium and francium which are the least electronegative
at 0.7.
What happens if two atoms of equal electronegativity bond together?
Consider a bond between two atoms, A and B. Each atom may be forming other
bonds as well as the one shown - but this makes no difference to the argument.
If the atoms are equally electronegative, both have the same tendency to
attract the bonding pair of electrons, and so it will be found on average half way
between the two atoms. To get a bond like this, A and B would usually have to be
the same atom. You will find this sort of bond in, for example, H 2 or Cl2
molecules.
This sort of bond could be thought of as being a "pure" covalent bond - where
the electrons are shared evenly between the two atoms.
What happens if B is slightly more electronegative than A?
B will attract the electron pair rather more than A does.
That means that the B end of the bond has more than its fair share of electron
density and so becomes slightly negative. At the same time, the A end (rather
short of electrons) becomes slightly positive. In the diagram, " " (read as
"delta") means "slightly" - so + means "slightly positive".
Defining polar bonds
AS Chemistry
Unit 1: Theoretical Chemistry
This is described as a polar bond. A polar bond is a covalent bond in which there
is a separation of charge between one end and the other - in other words in
which one end is slightly positive and the other slightly negative. Examples
include most covalent bonds. The hydrogen-chlorine bond in HCl or the
hydrogen-oxygen bonds in water are typical.
The polarity of a bond depends on the difference in electronegativity between
elements:
Molecule
HCl
HBr
HI
Electronegativity
difference
0.9
0.7
0.4
What happens if B is a lot more electronegative than A?
In this case, the electron pair is dragged right over to B's end of the bond. To
all intents and purposes, A has lost control of its electron, and B has complete
control over both electrons. Ions have been formed.
A "spectrum" of bonds
The implication of all this is that there is no clear-cut division between covalent
and ionic bonds. In a pure covalent bond, the electrons are held on average
exactly half way between the atoms. In a polar bond, the electrons have been
dragged slightly towards one end.
How far does this dragging have to go before the bond counts as ionic? There is
no real answer to that. You normally think of sodium chloride as being a typically
ionic solid. Lithium iodide, on the other hand, would be described as being "ionic
with some covalent character".
When a large negative ion in a lattice is adjacent to a positive ion which is small
and highly charged, the electron cloud around the negative ion is distorted so
that it is no longer spherical.
AS Chemistry
Unit 1: Theoretical Chemistry
The negative ion is distorted and is ‘polarised’. Some electron density is
concentrated between the ions and the bond begins to resemble a covalent bond
in which an electron pair is localised between two atoms. Lithium iodide, for
example, dissolves in organic solvents like ethanol - not something which ionic
substances normally do.
Summary



No electronegativity difference between two atoms leads to a pure nonpolar covalent bond.
A small electronegativity difference leads to a polar covalent bond.
A large electronegativity difference leads to an ionic bond.
Polar bonds and polar molecules
In a simple molecule like HCl, if the bond is polar, then so is the whole molecule.
What about more complicated molecules?
In CCl4, each bond is polar.
The molecule as a whole, however, isn't polar - in the sense that it doesn't have
an end (or a side) which is slightly negative and one which is slightly positive.
The whole of the outside of the molecule is somewhat negative, but there is no
overall separation of charge from top to bottom, or from left to right.
By contrast, CHCl3 is polar.
AS Chemistry
Unit 1: Theoretical Chemistry
The hydrogen at the top of the molecule is less electronegative than carbon and
so is slightly positive. This means that the molecule now has a slightly positive
"top" and a slightly negative "bottom", and so is overall a polar molecule.
Patterns of electronegativity in the Periodic Table
The most electronegative element is fluorine.
If you remember that fact, everything becomes easy, because electronegativity
must always increase towards fluorine in the Periodic Table.
Trends in electronegativity across a period
As you go across a period the electronegativity increases. The chart shows
electronegativities from sodium to chlorine - you have to ignore argon. It
doesn't have an electronegativity, because it doesn't form bonds.
Trends in electronegativity down a group
As you go down a group, electronegativity decreases. (If it increases up to
fluorine, it must decrease as you go down.) The chart shows the patterns of
electronegativity in Groups 1 and 7.
AS Chemistry
Unit 1: Theoretical Chemistry
Explaining the patterns in electronegativity
The attraction that a bonding pair of electrons feels for a particular nucleus
depends on:



the number of protons in the nucleus;
the distance from the nucleus;
the amount of screening by inner electrons.
Task 1
Using sodium and chlorine as examples, can you explain why electronegativity
increases across a period?.....................................................................................................
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Task 2
Using fluorine and chlorine as your examples, can you explain why
electronegativity decreases as you go down a group?.....................................................
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Learning Objectives:
Candidates should be able to explain the origin of polar bonds, with reference to
electronegativity differences between atoms.
References:
A-level Chemistry: pages 36-37
Chemistry in Context: pages 97-99
AS Chemistry
Unit 1: Theoretical Chemistry
Part 16: Ionic (electrovalent) Bonding
A simple view of ionic bonding
At a simple level (like IGCSE) a lot of importance is attached to the electronic
structures of noble gases like neon or argon which have eight electrons in their
outer energy levels (or two in the case of helium). These noble gas structures
are thought of as being in some way a "desirable" thing for an atom to have.
You may well have been left with the strong impression that when other atoms
react, they try to organise things such that their outer levels are either
completely full or completely empty. As we have already seen this is very much
an over-simplification.
Ionic bonding in sodium chloride
Ionic bonds are formed when there is a large electronegativity difference
between atoms. So when, for example, sodium and chlorine atoms interact the
outer electrons on both atoms start to feel a pull from both nuclei. The higher
‘core charge’ of chlorine means that an electron is transferred from the sodium
atom to the chlorine atom and ions are formed. This can happen because
chlorine (1s22s22p63s23p5) has a partially filled orbital.
The sodium has lost an electron, so it no longer has equal numbers of electrons
and protons. Because it has one more proton than electrons, it has a charge of
1+. If electrons are lost from an atom, positive ions are formed.
Positive ions are sometimes called cations.
The chlorine has gained an electron, so it now has one more electron than
protons. It therefore has a charge of 1-. If electrons are gained by an atom,
negative ions are formed.
A negative ion is sometimes called an anion.
AS Chemistry
Unit 1: Theoretical Chemistry
The nature of the bond
The sodium ions and chloride ions are held together by the strong electrostatic
attractions between the positive and negative ions.
The formula of sodium chloride
You need one sodium atom to provide the extra electron for one chlorine atom,
so they combine together in the ratio 1:1. The formula is therefore NaCl.
Below are dot and cross diagrams representing the two ions formed,
Na+, 1s22s22p6
Cl-, 1s22s22p63s23p6
Some other examples of ionic bonding
Task 1
(a)
Can you draw a similar dot-cross diagram for the bonding in magnesium
oxide? Would you expect the ionic bonding in this compound to be
stronger or weaker than that in NaCl? Can you explain why?
AS Chemistry
Unit 1: Theoretical Chemistry
(b)
Similarly, draw dot-cross diagrams for the bonding in potassium oxide,
calcium chloride and aluminium oxide. In each case give the formula of the
compound formed.
Some common ions which don't have noble gas structures
You will have come across all of the following ions. They are all common and
perfectly stable, but not one of them has a noble gas structure.
Task 2
Can you write the electronic configuration for the following ions?
Ion
Fe3+
Cu2+
Zn2+
Ag+
Electronic configuration
Ion
Electronic configuration
AS Chemistry
Unit 1: Theoretical Chemistry
Apart from some elements at the beginning of a transition series (scandium
forming Sc3+ with an argon structure, for example), all transition elements and
any metals following a transition series (like tin and lead in Group 4, for
example) will have structures like those above.
That means that the only elements to form positive ions with noble gas
structures (apart from odd ones like scandium) are those in groups 1 and 2 of
the Periodic Table and aluminium in group 3 (boron in group 3 doesn't form ions).
Negative ions are tidier! Those elements in Groups 5, 6 and 7 which form simple
negative ions all have noble gas structures.
What determines what the charge is on an ion?
Elements combine to make the compound which is as stable as possible - the one
in which the greatest amount of energy is evolved in its making. The more
charges a positive ion has, the greater the attraction towards its accompanying
negative ion. The greater the attraction, the more energy is released when the
ions come together.
That means that elements forming positive ions will tend to give away as many
electrons as possible. But there's a down-side to this.
Energy is needed to remove electrons from atoms. This is called ionisation
energy. The more electrons you remove, the greater the total ionisation energy
becomes. Eventually the total ionisation energy needed becomes so great that
the energy released when the attractions are set up between positive and
negative ions isn't large enough to cover it.
The element forms the ion which makes the compound most stable - the one in
which most energy is released over-all.
Task 3
For example, why is calcium chloride CaCl2 rather than CaCl or CaCl3?
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AS Chemistry
Unit 1: Theoretical Chemistry
Polarisation of ions
We have already seen that few molecules may be regarded as purely covalent.
Most have some ionic character, i.e. they are polar. Similarly, there are many
ionic compounds which could be described as having some covalent character.
They contain anions (negative ions) which have become polarised.
Positive cation
Negative anion
This means that the cation distorts the electron charge cloud on the anion.
Polarisation brings more electron charge between the nuclei, and thus produces
a significant degree of covalent bonding between the ions.
Task 4
Look at the diagram below. Which ion-pair is likely to have the greatest covalent
character?
Summarise your reasons in the box below:
Property
Charge
Radius
Cation is the most
powerful polarising
agent when....
Anode is most easily
polarised when...
AS Chemistry
Unit 1: Theoretical Chemistry
Learning Objectives:
Candidates should be able to describe ionic (electrovalent) bonding, as in sodium
chloride and magnesium oxide, including the use of ‘dot-and-cross’ diagrams.
References:
A-level Chemistry: pages 30-33
Chemistry in Context: pages 84-86
AS Chemistry
Unit 1: Theoretical Chemistry
Part 17: Shapes of Molecules and Ions
The (valence shell) electron pair repulsion theory (VSEPR)
The shape of a molecule or ion is governed by the arrangement of the electron
pairs around the central atom. All you need to do is to work out how many
electron pairs there are at the bonding level, and then arrange them to produce
the minimum amount of repulsion between them. You have to include both
bonding pairs and lone pairs.
Task 1
Carry out the activity on worksheet ‘Balloon Molecules’.
How to work out the number of electron pairs
For compounds which contain only single bonds, you can do this by drawing
dot-and-cross diagrams, or using the simple rules outlined below:



Write down the number of electrons in the outer level of the central
atom. That will be the same as the Periodic Table group number, except in
the case of the noble gases which form compounds, when it will be 8.
Add one electron for each bond being formed. (This allows for the
electrons coming from the other atoms.)
Allow for any ion charge. For example, if the ion has a 1- charge, add one
more electron. For a 1+ charge, deduct an electron.
Now work out how many bonding pairs and lone pairs of electrons there are:


Divide by 2 to find the total number of electron pairs around the central
atom.
Work out how many of these are bonding pairs, and how many are lone
pairs. You know how many bonding pairs there are because you know how
many other atoms are joined to the central atom (assuming that only
single bonds are formed).
For example, if you have 4 pairs of electrons but only 3 bonds, there must
be 1 lone pair as well as the 3 bonding pairs.
Finally, you have to use this information to work out the shape:
AS Chemistry
Unit 1: Theoretical Chemistry

Arrange these electron pairs in space to minimise repulsions. The basic
shapes were found from the balloon molecule activity.
There is an additional factor which comes into play. Lone pairs are in orbitals
that are shorter and rounder than the orbitals that the bonding pairs occupy.
Because of this, there is more repulsion between a lone pair and a bonding pair
than there is between two bonding pairs.
That forces the bonding pairs together slightly - reducing the bond angle. It's
not much, but the examiners will expect you to know it.
Remember this:
Greatest repulsion
lone pair - lone pair
lone pair - bond pair
Least repulsion
bond pair - bond pair
Compounds which contain / triple bonds
All four electrons in a double bond (or 6 in a triple bond) count as one bonding
pair/group when considering the shapes of molecules.
E.g. Two electron pairs around the central atom
The only simple case of this is beryllium chloride, BeCl2. The electronegativity
difference between beryllium and chlorine isn't enough to allow the formation
of ions.
Beryllium has 2 outer electrons because it is in Group 2. It forms bonds to two
chlorines, each of which adds another electron to the outer level of the
beryllium. There is no ionic charge to worry about, so there are 4 electrons
altogether - 2 pairs.
It is forming 2 bonds so there are no lone pairs. The two bonding pairs arrange
themselves at 180° to each other, because that's as far apart as they can get.
The molecule is described as being linear.
AS Chemistry
Unit 1: Theoretical Chemistry
Task 2
In groups complete the table on worksheet ‘Shapes of Molecules’.
Learning Objectives
Candidates should be able to explain the shape of, and bond angles in, molecules
by using the qualitative model of electron-pair repulsion (including lone-pairs),
using as simple examples: BF3; CO2; CH4; NH3; H2O; SF6; ethane and ethene, and
analogous molecules.
References
A-level Chemistry: pages 38-39
Chemistry in Context: pages 90-92
AS Chemistry
Unit 1: Theoretical Chemistry
Some more useful examples
Three electron pairs around the central atom
The simple cases of this would be BF3 or BCl3.
Boron is in Group 3, so starts off with 3 electrons. It is forming 3 bonds, adding
another 3 electrons. There is no charge, so the total is 6 electrons - in 3 pairs.
Because it is forming 3 bonds there can be no lone pairs. The 3 pairs arrange
themselves as far apart as possible. They all lie in one plane at 120° to each
other. The arrangement is called trigonal planar.
In the diagram, the other electrons on the fluorines have been left out because
they are irrelevant.
Four electron pairs around the central atom
There are lots of examples of this. The simplest is methane, CH4 but it could
equally well apply to CCl4.
Carbon is in Group 4, and so has 4 outer electrons. It is forming 4 bonds to
hydrogens, adding another 4 electrons - 8 altogether, in 4 pairs. Because it is
forming 4 bonds, these must all be bonding pairs.
Four electron pairs arrange themselves in space in what is called a tetrahedral
arrangement. A tetrahedron is a regular triangularly-based pyramid. The carbon
atom would be at the centre and the hydrogens at the four corners. All the
bond angles are 109.5°.
AS Chemistry
Unit 1: Theoretical Chemistry
Note the use of a wedge to show atoms which are coming out of the plane of the
paper towards you, and a dotted line to show atoms which are going away from
you.
Other examples with four electron pairs around the central atom
Ammonia, NH3
Nitrogen is in Group 5 and so has 5 outer electrons. Each of the 3 hydrogen
atoms is adding another electron to the nitrogen's outer level, making a total of
8 electrons in 4 pairs. Because the nitrogen is only forming 3 bonds, one of the
pairs must be a lone pair. The electron pairs arrange themselves in a tetrahedral
fashion as in methane.
Be very careful when you describe the shape of ammonia. Although the electron
pair arrangement is tetrahedral, when you describe the shape, you only take
notice of the atoms. Ammonia is trigonal pyramidal - like a pyramid with the
three hydrogens at the base and the nitrogen at the top.
Water, H2O
Following the same logic as before, you will find that the oxygen has four pairs
of electrons, two of which are lone pairs. These will again take up a tetrahedral
arrangement. This time the bond angle closes slightly more to 104°, because of
the repulsion of the two lone pairs.
The shape isn't described as tetrahedral, because we only "see" the oxygen and
the hydrogens - not the lone pairs. Water is described as bent or V-shaped.
The ammonium ion, NH4+
The nitrogen has 5 outer electrons, plus another 4 from the four hydrogen
atoms - making a total of 9.
AS Chemistry
Unit 1: Theoretical Chemistry
But take care! This is a positive ion. It has a 1+ charge because it has lost 1
electron. That leaves a total of 8 electrons in the outer level of the nitrogen.
There are therefore 4 pairs, all of which are bonding because of the four
hydrogen atoms. The ammonium ion has exactly the same shape as methane,
because it has exactly the same electronic arrangement. NH4+ is tetrahedral.
Methane and the ammonium ion are said to be isoelectronic. Two species (atoms,
molecules or ions) are isoelectronic if they have exactly the same number and
arrangement of electrons (including the distinction between bonding pairs and
lone pairs).
Five electron pairs around the central atom
Phosphorus(V) fluoride, PF5
Phosphorus (in group 5) contributes 5 electrons, and the five fluorines 5 more,
giving 10 electrons in 5 pairs around the central atom. Since the phosphorus is
forming five bonds, there can't be any lone pairs.
The 5 electron pairs take up a shape described as a trigonal bipyramid - three
of the fluorines are in a plane at 120° to each other; the other two are at right
angles to this plane. The trigonal bipyramid therefore has two different bond
angles - 120° and 90°.
Six electron pairs around the central atom
A simple example: SF6
6 electrons in the outer level of the sulphur, plus 1 each from the six fluorines,
makes a total of 12 - in 6 pairs. Because the sulphur is forming 6 bonds, these
AS Chemistry
Unit 1: Theoretical Chemistry
are all bond pairs. They arrange themselves entirely at 90°, in a shape described
as octahedral.
Two slightly more difficult examples
XeF4
Xenon forms a range of compounds, mainly with fluorine or oxygen, and this is a
typical one. Xenon has 8 outer electrons, plus 1 from each fluorine - making 12
altogether, in 6 pairs. There will be 4 bonding pairs (because of the four
fluorines) and 2 lone pairs.
There are two possible structures, but in one of them the lone pairs would be at
90°. Instead, they go opposite each other. XeF4 is described as square planar.
ClF4Chlorine is in group 7 and so has 7 outer electrons. Plus the 4 from the four
fluorines. Plus one because it has a 1- charge. That gives a total of 12 electrons
in 6 pairs - 4 bond pairs and 2 lone pairs. The shape will be identical with that of
XeF4.
AS Chemistry
Unit 1: Theoretical Chemistry
Part 18: Metallic Bonding
What is a metallic bond?
Metallic bonding in sodium
Metals tend to have high melting points and boiling points suggesting strong
bonds between the atoms. Even a metal like sodium (melting point 97.8°C) melts
at a considerably higher temperature than the element (neon) which precedes it
in the Periodic Table.
Sodium has the electronic structure 1s22s22p63s1. When sodium atoms come
together, the electron in the 3s atomic orbital of one sodium atom shares space
with the corresponding electron on a neighbouring atom to form a molecular
orbital - in much the same sort of way that a covalent bond is formed.
The difference, however, is that each sodium atom is being touched by eight
other sodium atoms - and the sharing occurs between the central atom and the
3s orbitals on all of the eight other atoms. And each of these eight is in turn
being touched by eight sodium atoms, which in turn are touched by eight atoms and so on and so on, until you have taken in all the atoms in that lump of sodium.
All of the 3s orbitals on all of the atoms overlap to give a vast number of
molecular orbitals which extend over the whole piece of metal. There have to be
huge numbers of molecular orbitals, of course, because any orbital can only hold
two electrons.
The electrons can move freely within these molecular orbitals, and so each
electron becomes detached from its parent atom. The electrons are said to be
delocalised. The metal is held together by the strong forces of attraction
between the positive nuclei and the delocalised electrons.
This is sometimes described as "an array of positive ions (cations) in a sea
of electrons".
AS Chemistry
Unit 1: Theoretical Chemistry
If you are going to use this view, beware! Is a metal made up of atoms or ions?
It is made of atoms.
Each positive centre in the diagram represents all the rest of the atom apart
from the outer electron, but that electron hasn't been lost - it may no longer
have an attachment to a particular atom, but it's still there in the structure.
Sodium metal is therefore written as Na - not Na+.
Metallic bonding in magnesium
If you work through the same argument with magnesium, you end up with
stronger bonds and so a higher melting point.
Magnesium has the outer electronic structure
become delocalised, so the "sea" has twice the
sodium. The remaining "ions" also have twice the
this particular view of the metal bond) and so
between "ions" and "sea".
3s2. Both of these electrons
electron density as it does in
charge (if you are going to use
there will be more attraction
More realistically, each magnesium atom has one more proton in the nucleus than
a sodium atom has, and so not only will there be a greater number of delocalised
electrons, but there will also be a greater attraction for them.
Magnesium atoms have a slightly smaller radius than sodium atoms, and so the
delocalised electrons are closer to the nuclei. Each magnesium atom also has
twelve near neighbours rather than sodium's eight. Both of these factors
increase the strength of the bond still further.
Metallic bonding in transition elements
Transition metals tend to have particularly high melting points and boiling
points. The reason is that they can involve the 3d electrons in the delocalisation
as well as the 4s. The more electrons you can involve, the stronger the
attractions tend to be.
The metallic bond in molten metals
In a molten metal, the metallic bond is still present, although the ordered
structure has been broken down. The metallic bond isn't fully broken until the
metal boils. That means that boiling point is actually a better guide to the
strength of the metallic bond than melting point is. On melting, the bond is
loosened, not broken.
AS Chemistry
Unit 1: Theoretical Chemistry
The structure of metals
The arrangement of the atoms
Metals are giant structures of atoms held together by metallic bonds. "Giant"
implies that large but variable numbers of atoms are involved - depending on the
size of the bit of metal.
12-co-ordination
Most metals are close packed - that is, they fit as many atoms as possible into
the available volume. Each atom in the structure has 12 touching neighbours.
Such a metal is described as 12-co-ordinated.
Each atom has 6 other atoms touching it in each layer.
There are also 3 atoms touching any particular atom in the layer above and
another 3 in the layer underneath.
This second diagram shows the layer immediately above the first layer. There
will be a corresponding layer underneath. (There are actually two different ways
of placing the third layer in a close packed structure, but that goes beyond the
requirements of current A-level syllabuses.)
8-co-ordination
Some metals (notably those in Group 1 of the Periodic Table) are packed less
efficiently, having only 8 touching neighbours. These are 8-co-ordinated.
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Unit 1: Theoretical Chemistry
The left hand diagram shows that no atoms are touching each other within a
particular layer. They are only touched by the atoms in the layers above and
below. The right hand diagram shows the 8 atoms (4 above and 4 below)
touching the darker coloured one.
Crystal grains
It would be misleading to suppose that all the atoms in a piece of metal are
arranged in a regular way. Any piece of metal is made up of a large number of
"crystal grains", which are regions of regularity. At the grain boundaries atoms
have become misaligned.
The physical properties of metals
Melting points and boiling points
Metals tend to have high melting and boiling points because of the strength of
the metallic bond. The strength of the bond varies from metal to metal and
depends on the number of electrons which each atom delocalises into the sea of
electrons, and on the packing.
Group 1 metals like sodium and potassium have relatively low melting and boiling
points mainly because each atom only has one electron to contribute to the bond
- but there are other problems as well:


Group 1 elements are also inefficiently packed (8-co-ordinated), so that
they aren't forming as many bonds as most metals.
They have relatively large atoms (meaning that the nuclei are some
distance from the delocalised electrons) which also weakens the bond.
Electrical conductivity
Metals conduct electricity. The delocalised electrons are free to move
throughout the structure in 3-dimensions. They can cross grain boundaries. Even
though the pattern may be disrupted at the boundary, as long as atoms are
touching each other, the metallic bond is still present.
AS Chemistry
Unit 1: Theoretical Chemistry
Liquid metals also conduct electricity, showing that although the metal atoms
may be free to move, the delocalisation remains in force until the metal boils.
Thermal conductivity
Metals are good conductors of heat. Heat energy is picked up by the electrons
as additional kinetic energy (it makes them move faster). The energy is
transferred throughout the rest of the metal by the moving electrons.
Strength and workability
Malleability and ductility
Metals are described as malleable (can be beaten into sheets) and ductile (can
be pulled out into wires). This is because of the ability of the atoms to roll over
each other into new positions without breaking the metallic bond.
If a small stress is put onto the metal, the layers of atoms will start to roll over
each other. If the stress is released again, they will fall back to their original
positions. Under these circumstances, the metal is said to be elastic.
If a larger stress is put on, the atoms roll over each other into a new position,
and the metal is permanently changed.
The hardness of metals
This rolling of layers of atoms over each other is hindered by grain boundaries
because the rows of atoms don't line up properly. It follows that the more grain
boundaries there are (the smaller the individual crystal grains), the harder the
metal becomes.
Offsetting this, because the grain boundaries are areas where the atoms aren't
in such good contact with each other, metals tend to fracture at grain
AS Chemistry
Unit 1: Theoretical Chemistry
boundaries. Increasing the number of grain boundaries not only makes the metal
harder, but also makes it more brittle.
Controlling the size of the crystal grains
If you have a pure piece of metal, you can control the size of the grains by heat
treatment or by working the metal.
Heating a metal tends to shake the atoms into a more regular arrangement decreasing the number of grain boundaries, and so making the metal softer.
Banging the metal around when it is cold tends to produce lots of small grains.
Cold working therefore makes a metal harder. To restore its workability, you
would need to reheat it.
You can also break up the regular arrangement of the atoms by inserting atoms
of a slightly different size into the structure. Alloys such as brass (a mixture
of copper and zinc) are harder than the original metals because the irregularity
in the structure helps to stop rows of atoms from slipping over each other.
Learning Objectives:
Candidates should be able to describe metallic bonding in terms of a lattice of
positive ions surrounded by mobile electrons.
References:
A-level Chemistry: pages 40-42
Chemistry in Context: pages 119-123
AS Chemistry
Unit 1: Theoretical Chemistry
Part 19: Intermolecular Forces
What are intermolecular attractions?
You have already seen how covalent bonds hold the atoms together within a
molecule. At room temperature, some substances made up of covalent molecules
are gases, others are liquids and the remainder are solids.
What is it that determines whether the substance is a gas, a liquid or a solid?
Task 1
Can you list three covalent molecules which are gases, three which are liquids
and three which are solids at room temperature?
Gases
Liquids
Solids
In a liquid or a solid there must be forces between the molecules causing them
to be attracted to one another, otherwise they would move apart from each
other and become a gas. These forces are called intermolecular forces.
If the temperature if lowered far enough, every substance, no matter how low
its boiling and melting point, will eventually solidify. When the solid melts and
then boils, it is the intermolecular forces that are broken. The covalent bonds
within the molecules remain intact. (Just think about ice changing into water and
then water vapour!)
Intermolecular attractions are attractions between one molecule and a
neighbouring molecule. The forces of attraction which hold an individual
molecule together (for example, the covalent bonds) are known as
intramolecular attractions. These two words are so confusingly similar that it is
safer to abandon one of them and never use it. The term "intramolecular" won't
be used again.
All molecules experience intermolecular attractions, although in some cases
those attractions are very weak. In the case of hydrogen the attractions are so
weak that the molecules have to be cooled to 21 K (-252°C) before the
AS Chemistry
Unit 1: Theoretical Chemistry
attractions are enough to condense the hydrogen as a liquid. Helium's
intermolecular attractions are even weaker - the molecules won't stick together
to form a liquid until the temperature drops to 4 K (-269°C).
How do intermolecular (or van der Waals) forces arise?
Temporary or instantaneous dipoles
Attractions are electrical in nature (i.e. they arise between positive and
negative charges). 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 lozenge-shaped diagram represents a small symmetrical molecule - H2,
perhaps, or Br2. The even shading shows that on average there is no electrical
distortion (i.e. the molecule is non-polar).
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 "sloshing 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.
If both the helium electrons happen to be on one side of the atom at the same
time, the nucleus is no longer properly covered by electrons for that instant.
AS Chemistry
Unit 1: Theoretical Chemistry
How temporary dipoles give rise to intermolecular attractions
I'm going to use the same lozenge-shaped diagram now to represent any
molecule which could, in fact, be a much more complicated shape. Shape does
matter (see later), but keeping the shape simple makes it a lot easier to both
draw the diagrams and understand what is going on.
Imagine a molecule which has a temporary dipole being approached by one which
happens to be entirely non-polar just at that moment. (A pretty unlikely event,
but it makes the diagrams much easier to draw! In reality, one of the molecules
is likely to have a greater polarity than the other at that time - and so will be
the dominant one.)
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.
There is no reason why this has to be restricted to two molecules. As long as
the molecules are close together this synchronised movement of the electrons
can occur over huge numbers of molecules.
AS Chemistry
Unit 1: Theoretical Chemistry
This diagram shows how a whole lattice of molecules could be held together in a
solid using van der Waals forces. An instant later, of course, you would have to
draw a quite different arrangement of the distribution of the electrons as they
shifted around - but always in synchronisation.
N.B. van der Waals’ forces are sometimes called (London) dispersion forces.
Task 2
Can you name a covalent molecule whose structure at room temperature
resembles the diagram above.
The strength of dispersion forces
Dispersion forces between molecules are much weaker than the covalent bonds
within molecules. It isn't possible to give any exact value, because the size of
the attraction varies considerably with the size of the molecule and its shape.
How molecular size affects the strength of the dispersion forces
The boiling points of the noble gases are:
helium
neon
-269°C
-246°C
argon
-186°C
krypton
-152°C
xenon
-108°C
radon
-62°C
All of these elements have monatomic molecules.
AS Chemistry
Unit 1: Theoretical Chemistry
Task 3
Can you explain why the boiling points increase as you go down the group?
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How molecular shape affects the strength of the dispersion forces
The shapes of the molecules also matter. Long thin molecules can lie closer
together - these attractions are at their most effective if the molecules are
really close.
For example, the hydrocarbon molecules butane and 2-methylpropane both have
a molecular formula C4H10, but the atoms are arranged differently. In butane
the carbon atoms are arranged in a single chain, but 2-methylpropane is a
shorter chain with a branch.
Butane has a higher boiling point because the intermolecular forces are greater.
The molecules are longer and can lie closer together than the shorter, fatter 2methylpropane molecules.
Task 4
1.
For each pair of chemicals given below, arrange the formulae in the order
in which the strength of the instantaneous dipole-induced dipole forces
increases. In each case, indicate which has the higher boiling point.
a) Xe and Kr
b) C8H18 and C6H14
e) Hexane and 2,3-dimethylbutane
c) CCl4 and CH4
d) Butane and methylpropane
AS Chemistry
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2.
Explain why noble gases have very low boiling points.
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3.
Draw skeletal formulae showing how two molecules of pentane can
approach close to one another. Now do the same for both of its
structural isomers. The boiling points of the three isomers are given in
the table below:
Isomer
pentane
2-methylbutane
2,2-dimethylpropane
Boiling point/K
309
301
283
a) Explain the variation in boiling points of the three isomers in terms of
the strength of the intermolecular forces present.
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b) Account for the differences in strengths of the intermolecular
forces.
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van der Waals forces: dipole-dipole interactions
AS Chemistry
Unit 1: Theoretical Chemistry
A molecule like HCl has a permanent dipole because chlorine is more
electronegative than hydrogen. These permanent, in-built dipoles will cause the
molecules to attract each other rather more than they otherwise would if they
had to rely only on just temporary dipole – induced dipole interactions.
It's important to realise that all molecules experience temporary dipole –
induced dipole interactions. Dipole-dipole interactions are not an alternative they occur in addition to them. Molecules which have permanent dipoles will
therefore have boiling points rather higher than molecules which only have
temporary fluctuating dipoles.
Surprisingly dipole-dipole attractions are fairly minor compared with temporary
dipole – induced dipole interactions, and their effect can only really be seen if
you compare two molecules with the same number of electrons and the same
size. For example, the boiling points of ethane, CH3CH3, and fluoromethane,
CH3F, are
Why choose these two molecules to compare? Both have identical numbers of
electrons, and if you made models you would find that the sizes were similar - as
you can see in the diagrams. That means that the temporary dipole – induced
dipole interactions in both molecules should be much the same.
The higher boiling point of fluoromethane is due to the large permanent dipole
on the molecule because of the high electronegativity of fluorine. However, even
given the large permanent polarity of the molecule, the boiling point has only
been increased by some 10°.
Task 5
Which molecule would you expect to have the higher boiling point - CHCl3(l) or
CCl4(l)? Explain your answer.
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Learning Objectives:
Candidates should be able to describe intermolecular forces (van der Waals’
forces), based on permanent and induced dipoles, as in CHCl3(l), Br2(l) and the
liquid noble gases.
References:
A-level Chemistry: pages 41-44
Chemistry in Context: pages 99-103
AS Chemistry
Unit 1: Theoretical Chemistry
Part 20: Hydrogen Bonding
Water is funny! (Peculiar that is – not ha! ha!)
Water is a very strange substance. If it weren’t for some of these unusual
properties life could not exist on Earth.
Many elements form compounds with hydrogen - referred to as "hydrides". If
you plot the boiling points of the hydrides of the Group 4 elements, you find
that the boiling points increase as you go down the group.
The increase in boiling point happens because the molecules are getting larger
with more electrons, and so van der Waals forces become greater.
If you repeat this exercise with the hydrides of elements in Groups 5, 6 and 7,
something odd happens.
Although for the most part the trend is exactly the same as in group 4 (for
exactly the same reasons), the boiling point of the hydride of the first element
in each group is abnormally high. In fact, for water it is so high that the
substance is a liquid (and not a gas) at room temperature.
AS Chemistry
Unit 1: Theoretical Chemistry
In the cases of NH3, H2O and HF there must be some additional intermolecular
forces of attraction, requiring significantly more heat energy to break. These
relatively powerful intermolecular forces are described as hydrogen bonds.
The origin of hydrogen bonding
Molecules which have this extra bonding include:
Hydrogen bonding is a particularly strong intermolecular force that involves
three features:



a large dipole between an H atom and the highly electronegative atoms N,
O or F;
the small H atom which can get very close to other atoms;
a lone pair of electrons on another N, O or F, with which the positively
charge H atom can line up.
Consider two water molecules coming close together.
The + hydrogen is so strongly attracted to the lone pair that it is almost as if
you were beginning to form a co-ordinate (dative covalent) bond. It doesn't go
that far, but the attraction is significantly stronger than an ordinary dipoledipole interaction.
N.B. Hydrogen bonds are often represented by a dotted line!
AS Chemistry
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Hydrogen bonds have about a tenth of the strength of an average covalent
bond, and are being constantly broken and reformed in liquid water. If you liken
the covalent bond between the oxygen and hydrogen to a stable marriage, the
hydrogen bond has "just good friends" status. On the same scale, van der Waals
attractions represent mere passing acquaintances!
Water as a "perfect" example of hydrogen bonding
Notice that each water molecule can potentially form four hydrogen bonds with
surrounding water molecules. There are exactly the right numbers of +
hydrogens and lone pairs so that 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. In the case of
ammonia, the amount of hydrogen bonding is limited by
the fact that each nitrogen only has one lone pair. In a
group of ammonia molecules, there aren't enough lone
pairs to go around to satisfy all the hydrogens.
In hydrogen fluoride, the problem is a shortage of
hydrogens. In water, there are exactly the right number of each. Water could
be considered as the "perfect" hydrogen bonded system.
Hydrogen bonding accounts for many of the other unusual properties of water
including:

its high specific heat capacity

its very high surface tension

its high viscosity and

the low density of ice compared to water
AS Chemistry
Unit 1: Theoretical Chemistry
Task
Which members of the following pairs would you expect to have the higher
boiling temperature? Give reasons for your choice.
a)
b)
c)
d)
C3H8 and CH3OCH3
CH3CH2NH2 and CH3CH2OH
CH3CH2OH and C2H6
C3H8 and (CH3)2C=O
Learning Objectives:
Candidates should be able to describe hydrogen bonding, using ammonia and
water as simple examples of molecules containing N-H and O-H groups.
References:
A-level Chemistry: pages 45-47
Chemistry in Context: pages 103-106
AS Chemistry
Unit 1: Theoretical Chemistry
Part 21: Bonding, Structure and Properties – a Summary
The properties of substances are decided by their bonding and structure.

Bonding means the way the particles are held together: ionic, covalent,
metallic or weak intermolecular bonds.

Structure means the way the particles are arranged relative to one
another. You have already met the major types of structure at IGCSE.
Task 1
The quantity of energy needed to break a particular covalent bond in a molecule
is called the bond enthalpy. Average bond enthalpies are given in the table
below. What general trends do you observe?
Bond
Average
bond
enthalpy/kJmol-1
Bond length/nm
C–C
C=C
C≡C
C–H
O–H
C–O
C=O
+347
+612
+838
+413
+464
+358
+805
0.154
0.134
0.120
0.108
0.096
0.143
0.116
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AS Chemistry
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Task 2
Can you use pages 65 – 66 of your textbook and your own scientific knowledge
to complete the table overleaf?
Learning Objectives:
Candidates should be able to



describe, interpret and/or predict the effect of different types of
bonding on the physical properties of substances.
describe, in simple terms, the lattice structure of a crystalline solid
which is ionic, simple molecular, giant molecular, hydrogen-bonded and
metallic.
suggest from quoted physical data the type of structure and bonding
present in a substance.
References:
A-level Chemistry: pages 48 an 65 – 66
Chemistry in Context: pages 119 – 132
AS Chemistry
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AS Chemistry
Unit 1: Theoretical Chemistry
Ionic
What
substances
have this type of
structure?
Examples
What
type
of
particle
does
it
contain?
How are the particles
bonded together?
What are the typical
properties?
M. pt and b.pt.
Hardness
Electrical conductivity
Solubility in water
Solubility in non-polar
solvents
(e.g.
hexane)
Compounds of metals
with non-metals.
GIANT LATTICE
Covalent network
Some
elements
in
Group 4 and some of
their compounds.
Metallic
Metals
COVALENT MOLECULAR
Simple molecular
Macromolecular
Some
non-metal
elements and usually
some non-metal/nonmetal compounds.
Polymers
AS Chemistry
Unit 1: Theoretical Chemistry
Part 22: The modern use of materials
Modern day man has learned to employ an enormous range of materials: metals
and their alloys, ceramics, glass and polymers.
Metals
Q.1
Can
you
name
the
five
most
commonly
used
metals
or
alloys?..........................................................................................................................................
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Q.2
Aluminium and its alloys have a number of uses. Which physical property
or properties make it ideal for the following uses?
a)
Bodywork of aeroplanes, trains, buses etc............................................................
b)
Overhead electrical cables........................................................................................
c)
Food packaging..............................................................................................................
Q.3.
Copper has high electrical and thermal conductivity, is malleable and
ductile and is resistant to corrosion. Can you list some of its main uses?.................
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Q.4
Copper has two main alloys. Can you name them and say which other
metals they contain? For each alloy give its main properties and uses.
Copper Alloy
Other metal it
contains
Main properties
Uses
AS Chemistry
Unit 1: Theoretical Chemistry
Ceramics
Ceramics are hard, brittle, heat- and corrosion-resistant materials made by
firing minerals (both natural, purified and synthetic) at high temperature. They
typically contain metallic elements, or silicon, combined with oxygen (or with
carbon, nitrogen, or sulphur). They all have giant structures and include
magnesium oxide, aluminium oxide and silicon (IV) oxide.
Material
Property
Uses
Magnesium/aluminium
Very high melting points.
Furnace linings
oxide
Silicon (IV) oxide
Very high melting point, Electrical insulators for
strong,
electrical
hard,
and
rigid, overhead power lines.
thermal
insulator
Manufacture of glass and
crockery.
Heat shields for Space
Shuttle.
Recycling
Read the passage below and give a definition for each of the terms in bold.
Raw materials extracted from the Earth cannot last forever. Although some
materials are more abundant than others, they are all finite resources.
Increasing demand for raw materials, coupled with ever growing problems of
waste disposal, have led to considerable interest in recycling waste.
Recycling has a number of possible advantages:

It leads to reduced demand for new raw materials;

It leads to a reduction in environmental damage;

It reduces the demand for landfill sites to dump waste;

It reduces the cost of waste disposal;

It may reduce energy costs.
AS Chemistry
Unit 1: Theoretical Chemistry
Learning Objectives:
Candidates should be able to




Explain the strength, high melting point and insulating properties of
ceramics in terms of their giant molecular structure.
Relate the uses of ceramics to their properties.
Describe and interpret the uses of the metals aluminium and copper (and
their alloys) in terms of their physical properties.
Understand that materials are a finite resource and the importance of
recycling processes.
References:
A-level Chemistry: pages 66 - 68
AS Chemistry
Unit 1: Theoretical Chemistry
Part 23: The Kinetic – Molecular Model of Liquids
The three states of matter are solid, liquid and gas. Whether a substance
exists as a solid, liquid or gas mainly depends on two things:


Kinetic energy – which increases as a substance is heated and brings
disorder to the movement of particles, and
Intermolecular forces – which tend to bring order to the movement of
particles.
At any given temperature, a substance will exist as a solid, liquid or gas
depending on the balance between these two opposing influences.
Task 1
Can you complete the table below?
Solid
Arrangement
of particles
Movement
of
particles
Proximity
of
particles
Compressibility
of substance
Conduction
heat
of
Liquid
Gas
AS Chemistry
Unit 1: Theoretical Chemistry
Task 2
Can you complete the notes below?
Solids have particles in fixed positions within a __________. The particles can
__________ but not move about. Hence a solid has a fixed __________. The
lattice may be held together by __________ attractions, covalent bonds,
metallic bonding, hydrogen bonding, dipole-dipole forces or __________. If
sufficient energy is supplied the particles can begin to move around each other;
this is melting. The particles in a solid may be atoms, ions, or molecules.
Liquids do not have a fixed __________ because the particles can move about.
However, they remain very __________ together. This shows that the interparticle forces have not been __________ broken. If sufficient __________
is supplied, the particles overcome the inter-particle forces almost completely
and __________ from the liquid. This is called __________ or boiling. The
energy required to boil a liquid is always __________ than that required to
melt the same substance and is a better __________ of the strength of interparticle forces.
Gases are made up of particles which are widely __________ and move in rapid
__________ motion. The forces between particles in the gas phase are
__________.
A key term: Vapour pressure
Vapour pressure is the pressure of a vapour over a liquid at equilibrium.
Imagine an empty closed box of several litres in size. I inject some liquid into
the box, but the box is not full. What will happen to the liquid?
Each molecule in the liquid has energy, but not the same amount. The energy is
distributed according to the Maxwell-Boltzmann distribution.
AS Chemistry
Unit 1: Theoretical Chemistry
This means some particles have a fairly large amount of energy compared to the
average.
Now, if some of these high energy particles happen to be sitting at the surface
of the water then they might have enough energy to escape the inter-particle
forces and become a vapour. We are making some vapour pressure. This
happens to another and another and another particle.
But wait! The vapour pressure stops going up and winds up staying at some fixed
value. What’s going on?
As more and more molecules leave the surface, what do some start to do?
That’s right; some return to the surface and resume their former life as a liquid
molecule. Soon the number of molecules in the vapour phase is constant because
the rate of returning equals the rate of leaving and so the pressure remains
constant.
We can measure the vapour pressure using a system like that shown below:-
AS Chemistry
Unit 1: Theoretical Chemistry
Learning Objectives:
Candidates should be able to describe using a kinetic-molecular model, the liquid
state; melting; vaporisation and vapour pressure.
References:
A-level Chemistry: pages 50-58
AS Chemistry
Unit 1: Theoretical Chemistry
Part 24: More on Ideal Gases
In an ideal gas:
 The molecules have mass but negligible size;
 There are no intermolecular forces.
We have seen from our previous lesson that a gas is most like a gas and least
like a liquid i.e. most ideal when a) the temperature is high and b) the pressure
is low. As we change these conditions different gases begin to deviate from
ideality to different extents.
Task - Can you answer the following past paper question?
Explain your answer in the space below:
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AS Chemistry
Unit 1: Theoretical Chemistry
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Learning Objectives:
Candidates should be able to explain qualitatively in terms of intermolecular
forces and molecular size the limitations of ideality at very high pressures and
very low temperatures.
References:
A-level Chemistry: pages 50-58