IONIC Bonding

A SIMPLE VIEW OF ATOMIC STRUCTURE

This page revises the simple ideas about atomic structure that you will have come across in an introductory chemistry course (for example, GCSE). You need to be confident about this before you go on to the more difficult ideas about the atom which under-pin A'level chemistry.

The sub-atomic particles

Protons, neutrons and electrons.

relative mass relative charge

proton 1 +1

neutron 1 0

electron 1/1836 -1

Beyond A'level: Protons and neutrons don't in fact have exactly the same mass - neither of them has a

mass of exactly 1 on the carbon-12 scale (the scale on which the relative masses of atoms are

measured). On the carbon-12 scale, a proton has a mass of 1.0073, and a neutron a mass of 1.0087.

The behaviour of protons, neutrons and electrons in electric fields

What happens if a beam of each of these particles is passed between two electrically charged plates - one positive and one negative? Opposites will attract.

Protons are positively charged and so would be deflected on a curving path towards the negative plate.

Electrons are negatively charged and so would be deflected on a curving path towards the positive plate.

Neutrons don't have a charge, and so would continue on in a straight line.

Exactly what happens depends on whether the beams of particles enter the electric field with the various particles having the same speeds or the same energies

If the particles have the same energy

If beams of the three sorts of particles, all with the same energy, are passed between two electrically charged plates:

Protons are deflected on a curved path towards the negative plate. Electrons are deflected on a curved path towards the positive plate.

The amount of deflection is exactly the same in the electron beam as the proton beam if the energies are the same - but, of course, it is in the opposite direction.

Neutrons continue in a straight line.

If the electric field was strong enough, then the electron and proton beams might curve enough to hit their respective plates.

If the particles have the same speeds

If beams of the three sorts of particles, all with the same speed, are passed between two electrically charged plates:

Protons are deflected on a curved path towards the negative plate. Electrons are deflected on a curved path towards the positive plate.

If the electrons and protons are travelling with the same speed, then the lighter electrons are deflected far more strongly than the heavier protons.

Neutrons continue in a straight line.

Note: This is potentially very confusing! Most chemistry sources that talk about this give either one or the

other of these two diagrams without any comment at all - they don't specifically say that they are using

constant energy or constant speed beams. But it matters!

If this is on your syllabus, it is important that you should know which version your examiners are going to expect, and they probably won't tell you in the syllabus. You should look in detail at past questions, mark schemes and examiner's reports which you can get from your examiners if you are doing a UK-based syllabus. Information about how to do this is on the syllabuses page.

If in doubt, I suggest you use the second (constant speed) version. This actually produces more useful information about both masses and charges than the constant energy version.

The nucleus

The nucleus is at the centre of the atom and contains the protons and neutrons. Protons and neutrons are collectively known as nucleons.

Virtually all the mass of the atom is concentrated in the nucleus, because the electrons weigh so little.

Working out the numbers of protons and neutrons

No of protons = ATOMIC NUMBER of the atom

The atomic number is also given the more descriptive name of proton number.

No of protons + no of neutrons = MASS NUMBER of the atom

The mass number is also called the nucleon number.

This information can be given simply in the form:

How many protons and neutrons has this atom got?

The atomic number counts the number of protons (9); the mass number counts protons + neutrons (19). If there are 9 protons, there must be 10 neutrons for the total to add up to 19.

The atomic number is tied to the position of the element in the Periodic Table and therefore the number of protons defines what sort of element you are talking about. So if an atom has 8 protons (atomic number = 8), it must be oxygen. If an atom has 12 protons (atomic number = 12), it must be magnesium.

Similarly, every chlorine atom (atomic number = 17) has 17 protons; every uranium atom (atomic number = 92) has 92 protons.

Isotopes

The number of neutrons in an atom can vary within small limits. For example, there are three kinds of carbon atom 12C, 13C and 14C. They all have the same number of protons, but the number of neutrons varies.

protons neutrons mass number

carbon-12 6 6 12

carbon-13 6 7 13

carbon-14 6 8 14

These different atoms of carbon are called isotopes. The fact that they have varying numbers of neutrons makes no difference whatsoever to the chemical reactions of the carbon.

Isotopes are atoms which have the same atomic number but different mass numbers. They have the same number of protons but different numbers of neutrons.

The electrons

Working out the number of electrons

Atoms are electrically neutral, and the positiveness of the protons is balanced by the negativeness of the electrons. It follows that in a neutral atom:

no of electrons = no of protons

So, if an oxygen atom (atomic number = 8) has 8 protons, it must also have 8 electrons; if a chlorine atom (atomic number = 17) has 17 protons, it must also have 17 electrons.

The arrangement of the electrons

The electrons are found at considerable distances from the nucleus in a series of levels called

energy levels. Each energy level can only hold a certain number of electrons. The first level

(nearest the nucleus) will only hold 2 electrons, the second holds 8, and the third also seems to

be full when it has 8 electrons. At GCSE you stop there because the pattern gets more

complicated after that.

These levels can be thought of as getting progressively further from the nucleus. Electrons will always go into the lowest possible energy level (nearest the nucleus) - provided there is space.

To work out the electronic arrangement of an atom

Look up the atomic number in the Periodic Table - making sure that you choose the right number if two numbers are given. The atomic number will always be the smaller one.

This tells you the number of protons, and hence the number of electrons. Arrange the electrons in levels, always filling up an inner level before you go to

an outer one.

e.g. to find the electronic arrangement in chlorine

The Periodic Table gives you the atomic number of 17. Therefore there are 17 protons and 17 electrons. The arrangement of the electrons will be 2, 8, 7 (i.e. 2 in the first level, 8 in the

second, and 7 in the third).

The electronic arrangements of the first 20 elements

After this the pattern alters as you enter the transition series in the Periodic Table.

Two important generalisations

If you look at the patterns in this table:

The number of electrons in the outer level is the same as the group number. (Except with helium which has only 2 electrons. The noble gases are also usually called group 0 - not group 8.) This pattern extends throughout the Periodic Table for the main groups (i.e. not including the transition elements).

So if you know that barium is in group 2, it has 2 electrons in its outer level; iodine (group 7) has 7 electrons in its outer level; lead (group 4) has 4 electrons in its outer level.

Noble gases have full outer levels. This generalisation will need modifying for A'level purposes.

Dots-and-crosses diagrams

In any introductory chemistry course you will have come across the electronic structures of hydrogen and carbon, for example, drawn as:

Note: There are many places where you could still make use of this model of the atom at A'level. It is,

however, a simplification and can be misleading. It gives the impression that the electrons are circling the

nucleus in orbits like planets around the sun. As you will find when you look at the A'level view of the

atom, it is impossible to know exactly how they are actually moving.

The circles show energy levels - representing increasing distances from the nucleus. You could straighten the circles out and draw the electronic structure as a simple energy diagram.

Carbon, for example, would look like this:

Thinking of the arrangement of the electrons in this way makes a useful bridge to the A'level view.

IONIC (ELECTROVALENT) BONDING

This page explains what ionic (electrovalent) bonding is. It starts with a simple picture of the formation of ions, and then modifies it slightly for A'level purposes.

A simple view of ionic bonding

The importance of noble gas structures

At a simple level (like GCSE) 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.

Note: The central role given to noble gas structures is very much an over-simplification. We shall have to

spend some time later on demolishing the concept!

Ionic bonding in sodium chloride

Sodium (2,8,1) has 1 electron more than a stable noble gas structure (2,8). If it gave away that electron it would become more stable.

Chlorine (2,8,7) has 1 electron short of a stable noble gas structure (2,8,8). If it could gain an electron from somewhere it too would become more stable.

The answer is obvious. If a sodium atom gives an electron to a chlorine atom, both become more stable.

The sodium has lost an electron, so it no longer has equal numbers of electrons and protons. Because it has one more proton than electron, 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 proton. 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.

The nature of the bond

The sodium ions and chloride ions are held together by the strong electrostatic attractions between the positive and negative charges.

The formula of sodium chloride

You need one sodium atom to provide the extra electron for one chlorine atom, so they combine together 1:1. The formula is therefore NaCl.

Some other examples of ionic bonding

magnesium oxide

Again, noble gas structures are formed, and the magnesium oxide is held together by very strong attractions between the ions. The ionic bonding is stronger than in sodium chloride because this time you have 2+ ions attracting 2- ions. The greater the charge, the greater the attraction.

The formula of magnesium oxide is MgO.

calcium chloride

This time you need two chlorines to use up the two outer electrons in the calcium. The formula of calcium chloride is therefore CaCl2.

potassium oxide

Again, noble gas structures are formed. It takes two potassiums to supply the electrons the oxygen needs. The formula of potassium oxide is K2O.

THE A'LEVEL VIEW OF IONIC BONDING

Electrons are transferred from one atom to another resulting in the formation of positive and negative ions.

The electrostatic attractions between the positive and negative ions hold the compound together.

So what's new? At heart - nothing. What needs modifying is the view that there is something magic about noble gas structures. There are far more ions which don't have noble gas structures than there are which do.

Some common ions which don't have noble gas structures

You may have come across some of the following ions in a basic course like GCSE. They are all perfectly stable , but not one of them has a noble gas structure.

Fe3+

[Ar]3d5

Cu2+

[Ar]3d9

Zn2+

[Ar]3d10

Ag+

[Kr]4d10

Pb2+

[Xe]4f145d106s2

Noble gases (apart from helium) have an outer electronic structure ns2np6.

Note: If you aren't happy about writing electronic structures using of s, p and d notation, follow this link before you go on.

Return to this page via the menus or by using the BACK button on your browser.

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.

If elements aren't aiming for noble gas structures when they form ions, what decides how many electrons are transferred? The answer lies in the energetics of the process by which the compound is made.

Warning! From here to the bottom of this page goes beyond anything you are likely to need for A'level purposes. It is included for interest only.

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.

For example, why is calcium chloride CaCl2 rather than CaCl or CaCl3?

If one mole of CaCl (containing Ca+ ions) is made from its elements, it is possible to estimate that about 171 kJ of heat is evolved.

However, making CaCl2 (containing Ca2+ ions) releases more heat. You get 795 kJ. That extra amount of heat evolved makes the compound more stable, which is why you get CaCl2 rather than CaCl.

What about CaCl3 (containing Ca3+ ions)? To make one mole of this, you can estimate that you would have to put in 1341 kJ. This makes this compound completely non-viable. Why is so much heat needed to make CaCl3? It is because the third ionisation energy (the energy needed to remove the third electron) is extremely high (4940 kJ mol-1) because the electron is being removed from the 3-level rather than the 4-level. Because it is much closer to the nucleus than the first two electrons removed, it is going to be held much more strongly.

Note: It would pay you to read about ionisation energies if you really want to understand this.

You could also go to a standard text book and investigate Born-Haber Cycles.

A similar sort of argument applies to the negative ion. For example, oxygen forms an O2- ion rather than an O- ion or an O3- ion, because compounds containing the O2- ion turn out to be the most energetically stable.

IONIC STRUCTURES

This page explains the relationship between the arrangement of the ions in a typical ionic solid like sodium chloride and its physical properties - melting point, boiling point, brittleness, solubility and electrical behaviour. It also explains why caesium chloride has a different structure from sodium chloride even though sodium and caesium are both in Group 1 of the Periodic Table.

Note: If you need to revise how ionic bonding arises, then you might like to follow this link. It isn't

important for understanding this page, however.

The structure of a typical ionic solid - sodium chloride

How the ions are arranged in sodium chloride

Sodium chloride is taken as a typical ionic compound. Compounds like this consist of a giant (endlessly repeating) lattice of ions. So sodium chloride (and any other ionic compound) is described as having a giant ionic structure.

You should be clear that giant in this context doesn't just mean very large. It means that you can't state exactly how many ions there are.

There could be billions of sodium ions and chloride ions packed together, or trillions, or whatever - it simply depends how big the crystal is. That is different from, say, a water molecule which always contains exactly 2 hydrogen atoms and one oxygen atom - never more and never less.

A small representative bit of a sodium chloride lattice looks like this:

If you look at the diagram carefully, you will see that the sodium ions and chloride ions alternate with each other in each of the three dimensions.

This diagram is easy enough to draw with a computer, but extremely difficult to draw convincingly by hand. We normally draw an "exploded" version which looks like this:

Only those ions joined by lines are actually touching each other. The sodium ion in the centre is being touched by 6 chloride ions. By chance we might just as well have centred the diagram around a chloride ion - that, of course, would be touched by 6 sodium ions. Sodium chloride is described as being 6:6-co-ordinated.

You must remember that this diagram represents only a tiny part of the whole sodium chloride crystal. The pattern repeats in this way over countless ions.

How to draw this structure

Draw a perfect square:

Now draw an identical square behind this one and offset a bit. You might have to practice a bit to get the placement of the two squares right. If you get it wrong, the ions get all tangled up with each other in your final diagram.

Turn this into a perfect cube by joining the squares together:

Now the tricky bit! Subdivide this big cube into 8 small cubes by joining the mid point of each edge to the mid point of the edge opposite it. To complete the process you will also have to join the mid point of each face (easily found once you've joined the edges) to the mid point of the opposite face.

Now all you have to do is put the ions in. Use different colours or different sizes for the two different ions, and don't forget a key. It doesn't matter whether you end up with a sodium ion or a chloride ion in the centre of the cube - all that matters is that they alternate in all three dimensions.

You should be able to draw a perfectly adequate free-hand sketch of this in under two minutes - less than one minute if you're not too fussy!

Why is sodium chloride 6:6-co-ordinated?

The more attraction there is between the positive and negative ions, the more energy is released. The more energy that is released, the more energetically stable the structure becomes.

That means that to gain maximum stability, you need the maximum number of attractions. So why does each ion surround itself with 6 ions of the opposite charge?

That represents the maximum number of chloride ions that you can fit around a central sodium ion before the chloride ions start touching each other. If they start touching, you introduce repulsions into the crystal which makes it less stable.

The different structure of caesium chloride

We'll look first at the arrangement of the ions and then talk about why the structures of sodium chloride and caesium chloride are different afterwards.

Warning: Before you go on with this section, make sure that you actually need it for your syllabus. If you

don't, jump down the page to "The physical properties of sodium chloride"

How the ions are arranged in caesium chloride

Imagine a layer of chloride ions as shown below. The individual chloride ions aren't touching each other. That's really important - if they were touching, there would be repulsion.

Now let's place a similarly arranged layer of caesium ions on top of these.

Notice that the caesium ions aren't touching each other either, but that each caesium ion is resting on four chloride ions from the layer below.

Now let's put another layer of chloride ions on, exactly the same as the first layer. Again, the chloride ions in this layer are NOT touching those in the bottom layer - otherwise you are introducing repulsion. Since we are looking directly down on the structure, you can't see the bottom layer of chloride ions any more, of course.

If you now think about a caesium ion sandwiched between the two layers of chloride ions, it is touching four chloride ions in the bottom layer, and another four in the top one. Each caesium ion is touched by eight chloride ions. We say that it is 8-co-ordinated.

If we added another layer of caesium ions, you could similarly work out that each chloride ion was touching eight caesium ions. The chloride ions are also 8-co-ordinated.

Overall, then, caesium chloride is 8:8-co-ordinated.

The final diagram in this sequence takes a slightly tilted view of the structure so that you can see how the layers build up. These diagrams are quite difficult to draw without it looking as if ions of the same charge are touching each other. They aren't!

Diagrams of ionic crystals are usually simplified to show the most basic unit of the repeating pattern. For caesium chloride, you could, for example, draw a simple diagram showing the arrangement of the chloride ions around each caesium ion:

By reversing the colours (green chloride ion in the centre, and orange caesium ions surrounding it), you would have an exactly equivalent diagram for the arrangement of caesium ions around each chloride ion.

Note: These diagrams are difficult enough to draw convincingly on a computer. Trying to draw them

freehand in an exam is seriously difficult. If you are doing a syllabus which wants you to know about the

structure of caesium chloride, take a careful look at past exam papers and mark schemes to see exactly

what sort of diagrams (if any) you need to use, and then practise them so that you can draw them quickly

and well. If you haven't got any past papers and mark schemes, follow this link to the syllabuses page to

find out how to get them if you are doing a UK-based exam.

Why are the caesium chloride and sodium chloride structures different?

When attractions are set up between two ions of opposite charges, energy is released. The more energy that can be released, the more stable the system becomes. That means that the more contact there is between negative and positive ions, the more stable the crystal should become.

If you can surround a positive ion like caesium with eight chloride ions rather than just six (and vice versa for the chloride ions), then you should have a more stable crystal. So why doesn't sodium chloride do the same thing?

Look again at the last diagram:

Now imagine what would happen if you replaced the caesium ion with the smaller sodium ion. Sodium ions are, of course, smaller than caesium ions because they have fewer layers of electrons around them.

You still have to keep the chloride ions in contact with the sodium. The effect of this would be that the whole arrangement would shrink, bringing the chloride ions into contact with each other - and that introduces repulsion.

Any gain in attractions because you have eight chlorides around the sodium rather than six is more than countered by the new repulsions between the chloride ions themselves. When sodium chloride is 6:6-co-ordinated, there are no such repulsions - and so that is the best way for it to organise itself.

Which structure a simple 1:1 compound like NaCl or CsCl crystallises in depends on the radius ratio of the positive and the negative ions. If the radius of the positive ion is bigger than 73% of that of the negative ion, then 8:8-co-ordination is possible. Less than that (down to 41%) then you get 6:6-co-ordination.

In CsCl, the caesium ion is about 93% of the size of the chloride ion - so is easily within the range where 8:8-co-ordination is possible. But with NaCl, the sodium ion is only about 52% of the size of the chloride ion. That puts it in the range where you get 6:6-co-ordination.

Note: What happens below 41%? At this point the negative ions will touch each other again even with

6:6-co-ordination. A new arrangement (known as 4:4-co-ordination) then becomes necessary. This is

beyond any syllabus that I am currently tracking.

The physical properties of sodium chloride

Sodium chloride is taken as typical of ionic compounds, and is chosen rather than, say, caesium chloride, because it is found on every syllabus at this level.

Sodium chloride has a high melting and boiling point

There are strong electrostatic attractions between the positive and negative ions, and it takes a lot of heat energy to overcome them. Ionic substances all have high melting and boiling points. Differences between ionic substances will depend on things like:

The number of charges on the ions

Magnesium oxide has exactly the same structure as sodium chloride, but a much higher melting and boiling point. The 2+ and 2- ions attract each other more strongly than 1+ attracts 1-.

The sizes of the ions

If the ions are smaller they get closer together and so the electrostatic attractions are greater. Rubidium iodide, for example, melts and boils at slightly lower temperatures than sodium chloride, because both rubidium and iodide ions are bigger than sodium and chloride ions. The attractions are less between the bigger ions and so less heat energy is needed to separate them.

Sodium chloride crystals are brittle

Brittleness is again typical of ionic substances. Imagine what happens to the crystal if a stress is applied which shifts the ion layers slightly.

Ions of the same charge are brought side-by-side and so the crystal repels itself to pieces!

Sodium chloride is soluble in water

Many ionic solids are soluble in water - although not all. It depends on whether there are big enough attractions between the water molecules and the ions to overcome the attractions between the ions themselves. Positive ions are attracted to the lone pairs on water molecules and co-ordinate (dative covalent) bonds may form. Water molecules form hydrogen bonds with negative ions.

Note: The bonding in hydrated metal ions is covered in the page on co-ordinate bonding. The bonding

between negative ions like chloride ions and water molecules is covered in the page on hydrogen

bonding.

Sodium chloride is insoluble in organic solvents

This is also typical of ionic solids. The attractions between the solvent molecules and the ions aren't big enough to overcome the attractions holding the crystal together.

The electrical behaviour of sodium chloride

Solid sodium chloride doesn't conduct electricity, because there are no electrons which are free to move. When it melts, sodium chloride undergoes electrolysis, which involves conduction of electricity because of the movement and discharge of the ions. In the process, sodium and chlorine are produced. This is a chemical change rather than a physical process.

The positive sodium ions move towards the negatively charged electrode (the cathode). When they get there, each sodium ion picks up an electron from the electrode to form a sodium atom.

The movement of electrons from the cathode onto the sodium ions leaves spaces on the cathode. The power source (the battery or whatever) moves electrons along the wire in the external circuit to fill those spaces. That flow of electrons would be seen as an electric current. (The external circuit is all the rest of the circuit apart from the molten sodium chloride.)

Meanwhile, chloride ions are attracted to the positive electrode (the anode). When they get there, each chloride ion loses an electron to the anode to form an atom. These then pair up to make chlorine molecules. Overall, the change is . . .

The new electrons deposited on the anode are pumped off around the external circuit by the power source, eventually ending up on the cathode where they will be transferred to sodium ions.

Molten sodium chloride conducts electricity because of the movement of the ions in the melt, and the discharge of the ions at the electrodes. Both of these have to happen if you are to get electrons flowing in the external circuit.

In solid sodium chloride, of course, that ion movement can't happen and that stops any possibility of any current flow in the circuit.

GIANT COVALENT STRUCTURES

This page decribes the structures of giant covalent substances like diamond, graphite and silicon dioxide (silicon(IV) oxide), and relates those structures to the physical properties of the substances.

The structure of diamond

The giant covalent structure of diamond

Carbon has an electronic arrangement of 2,4. In diamond, each carbon shares electrons with four other carbon atoms - forming four single bonds.

In the diagram some carbon atoms only seem to be forming two bonds (or even one bond), but that's not really the case. We are only showing a small bit of the whole structure.

This is a giant covalent structure - it continues on and on in three dimensions. It is not a molecule, because the number of atoms joined up in a real diamond is completely variable - depending on the size of the crystal.

Note: We quoted the electronic structure of carbon as 2,4. That simple view is perfectly adequate to

explain the bonding in diamond. If you are interested in a more modern view, you could read the page on

bonding in methane and ethane in the organic section of this site. In the case of diamond, each carbon is

bonded to 4 other carbons rather than hydrogens, but that makes no essential difference.

How to draw the structure of diamond

Don't try to be too clever by trying to draw too much of the structure! Learn to draw the diagram given above. Do it in the following stages:

Practise until you can do a reasonable free-hand sketch in about 30 seconds.

The physical properties of diamond

Diamond

has a very high melting point (almost 4000°C). Very strong carbon-carbon covalent bonds have to be broken throughout the structure before melting occurs.

is very hard. This is again due to the need to break very strong covalent bonds operating in 3-dimensions.

doesn't conduct electricity. All the electrons are held tightly between the atoms, and aren't free to move.

is insoluble in water and organic solvents. There are no possible attractions which could occur between solvent molecules and carbon atoms which could outweigh the attractions between the covalently bound carbon atoms.

The structure of graphite

The giant covalent structure of graphite

Graphite has a layer structure which is quite difficult to draw convincingly in three dimensions. The diagram below shows the arrangement of the atoms in each layer, and the way the layers are spaced.

Notice that you can't really draw the side view of the layers to the same scale as the atoms in the layer without one or other part of the diagram being either very spread out or very squashed.

In that case, it is important to give some idea of the distances involved. The distance between the layers is about 2.5 times the distance between the atoms within each layer.

The layers, of course, extend over huge numbers of atoms - not just the few shown above.

You might argue that carbon has to form 4 bonds because of its 4 unpaired electrons, whereas in this diagram it only seems to be forming 3 bonds to the neighbouring carbons. This diagram is something of a simplification, and shows the arrangement of atoms rather than the bonding.

The bonding in graphite

Each carbon atom uses three of its electrons to form simple bonds to its three close neighbours. That leaves a fourth electron in the bonding level. These "spare" electrons in each carbon atom become delocalised over the whole of the sheet of atoms in one layer. They are no longer associated directly with any particular atom or pair of atoms, but are free to wander throughout the whole sheet.

If you are interested (beyond A'level): The bonding in graphite is like a vastly extended version of the

bonding in benzene. Each carbon atom undergoes sp2 hybridisation, and then the unhybridised p orbitals

on each carbon atom overlap sideways to give a massive pi system above and below the plane of the

sheet of atoms.

The important thing is that the delocalised electrons are free to move anywhere within the sheet - each electron is no longer fixed to a particular carbon atom. There is, however, no direct contact between the delocalised electrons in one sheet and those in the neighbouring sheets.

The atoms within a sheet are held together by strong covalent bonds - stronger, in fact, than in diamond because of the additional bonding caused by the delocalised electrons. So what holds the sheets together?

In graphite you have the ultimate example of van der Waals dispersion forces. As the delocalised electrons move around in the sheet, very large temporary dipoles can be set up which will induce opposite dipoles in the sheets above and below - and so on throughout the whole graphite crystal.

Note: If you aren't sure about van der Waals forces follow this link before you go on. Use the BACK

button on your browser to return to this page.

The physical properties of graphite

Graphite

has a high melting point, similar to that of diamond. In order to melt graphite, it isn't enough to loosen one sheet from another. You have to break the covalent bonding throughout the whole structure.

has a soft, slippery feel, and is used in pencils and as a dry lubricant for things like locks. You can think of graphite rather like a pack of cards - each card is strong, but the cards will slide over each other, or even fall off the pack altogether. When you use a pencil, sheets are rubbed off and stick to the paper.

has a lower density than diamond. This is because of the relatively large amount of space that is "wasted" between the sheets.

is insoluble in water and organic solvents - for the same reason that diamond is insoluble. Attractions between solvent molecules and carbon atoms will never be strong enough to overcome the strong covalent bonds in graphite.

conducts electricity. The delocalised electrons are free to move throughout the sheets. If a piece of graphite is connected into a circuit, electrons can fall off one end of the sheet and be replaced with new ones at the other end.

Note: The logic of this is that a piece of graphite ought only to conduct electricity in 2-dimensions

because electrons can only move around in the sheets - and not from one sheet to its neighbours.

In practice, a real piece of graphite isn't a perfect crystal, but a host of small crystals stuck together at all sorts of angles. Electrons will be able to find a route through the large piece of graphite in all directions by moving from one small crystal to the next.

The structure of silicon dioxide, SiO2

Silicon dioxide is also known as silicon(IV) oxide.

The giant covalent structure of silicon dioxide

There are three different crystal forms of silicon dioxide. The easiest one to remember and draw is based on the diamond structure.

Crystalline silicon has the same structure as diamond. To turn it into silicon dioxide, all you need to do is to modify the silicon structure by including some oxygen atoms.

Notice that each silicon atom is bridged to its neighbours by an oxygen atom. Don't forget that this is just a tiny part of a giant structure extending on all 3 dimensions.

Note: If you want to be fussy, the Si-O-Si bond angles are wrong in this diagram. In reality the "bridge"

from one silicon atom to its neighbour isn't in a straight line, but via a "V" shape (similar to the shape

around the oxygen atom in a water molecule). It's extremely difficult to draw that convincingly and tidily in

a diagram involving this number of atoms. The simplification is perfectly acceptable.

The physical properties of silicon dioxide

Silicon dioxide

has a high melting point - varying depending on what the particular structure is (remember that the structure given is only one of three possible structures), but around 1700°C. Very strong silicon-oxygen covalent bonds have to be broken throughout the structure before melting occurs.

is hard. This is due to the need to break the very strong covalent bonds. doesn't conduct electricity. There aren't any delocalised electrons. All the

electrons are held tightly between the atoms, and aren't free to move. is insoluble in water and organic solvents. There are no possible attractions

which could occur between solvent molecules and the silicon or oxygen atoms which could overcome the covalent bonds in the giant structure.

COVALENT BONDING - SINGLE BONDS

This page explains what covalent bonding is. It starts with a simple picture of the single covalent bond, and then modifies it slightly for A'level purposes. It also takes a more sophisticated view (beyond A'level) if you are interested. You will find a link to a page on double covalent bonds at the bottom of the page.

A simple view of covalent bonding The importance of noble gas structures At a simple level (like GCSE) 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 achieve noble gas structures.

As well as achieving noble gas structures by transferring electrons from one atom to another as in ionic bonding, it is also possible for atoms to reach these stable structures by sharing electrons to give covalent bonds.

Some very simple covalent molecules Chlorine For example, two chlorine atoms could both achieve stable structures by sharing

their single unpaired electron as in the diagram.

The fact that one chlorine has been drawn with electrons marked as crosses and

the other as dots is simply to show where all the electrons come from. In reality there is no difference between them.

The two chlorine atoms are said to be joined by a covalent bond. The reason that the two chlorine atoms stick together is that the shared pair of electrons is attracted to the nucleus of both chlorine atoms.

Hydrogen

Hydrogen atoms only need two electrons in their outer level to reach the noble

gas structure of helium. Once again, the covalent bond holds the two atoms together because the pair of electrons is attracted to both nuclei.

Hydrogen chloride

The hydrogen has a helium structure, and the chlorine an argon structure.

Covalent bonding at A'level Cases where there isn't any difference from the simple view If you stick closely to modern A'level syllabuses, there is little need to move far

from the simple (GCSE) view. The only thing which must be changed is the over-reliance 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.

For example:

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

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.

Note: You might perhaps wonder why boron doesn't form ionic bonds with fluorine instead.

Boron doesn't form ions because the total energy needed to remove three electrons to form a B3+

ion is simply too great to be recoverable when attractions are set up between the boron and

fluoride ions.

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 beyond the limits of current A'level syllabuses. 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

Warning! If you aren't happy with describing electron arrangements in s and p notation, and with

the shapes of s and p orbitals, you need to read about orbitals before you go on.

Use the BACK button on your browser to return quickly to this point.

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, 1s22s22px12py

1. 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.

The carbon atom is now said to be in an excited state.

Note: People sometimes worry that the promoted electron is drawn as an up-arrow, whereas it

started as a down-arrow. The reason for this is actually fairly complicated - well beyond the level

we are working at. Just get in the habit of writing it like this because it makes the diagrams look

tidy!

Now that we've got 4 unpaired electrons ready for bonding, another problem arises. In methane all the carbon-hydrogen bonds are identical, but our electrons are in two different kinds of orbitals. You aren't going to get four identical bonds

unless you start from four identical orbitals. 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.

Note: You will find this bit on methane repeated in the organic section of this site. That article on

methane goes on to look at the formation of carbon-carbon single bonds in ethane.

The bonding in the phosphorus chlorides, PCl3 and PCl5 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 1s22s22p63s23px13py

13pz1. 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.

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:

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 1s22s22px12py

12pz1. 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.