PROPERTIES OF THE PERIOD 3 CHLORIDES

This page looks at the structures of the chlorides of the Period 3 elements (sodium to sulphur), their physical properties and their reactions with water.

Chlorine and argon are omitted - chlorine because it is meaningless to talk about "chlorine chloride", and argon because it doesn't form a chloride.

A quick summary of the trends

The chlorides

The chlorides we'll be looking at are:

There are three chlorides of sulphur, but the only one mentioned by any of the UK-based syllabuses (A level or its equivalents) is S₂Cl₂.

As you will see later, aluminium chloride exists in some circumstances as a dimer, Al₂Cl₆.

The structures

Sodium chloride and magnesium chloride are ionic and consist of giant ionic lattices at room temperature

Aluminium chloride and phosphorus(V) chloride are tricky! They change their structure from ionic to covalent when the solid turns to a liquid or vapour. There is much more about this later on this page.

The others are simple covalent molecules.

Melting and boiling points

Sodium and magnesium chlorides are solids with high melting and boiling points because of the large amount of heat which is needed to break the strong ionic attractions.

The rest are liquids or low melting point solids. Leaving aside the aluminium chloride and phosphorus(V) chloride cases where the situation is quite complicated, the attractions in the others will be much weaker intermolecular forces such as van der Waals dispersion forces. These vary depending on the size and shape of the molecule, but will always be far weaker than ionic bonds.

Electrical conductivity

Sodium and magnesium chlorides are ionic and so will undergo electrolysis when they are molten. Electricity is carried by the movement of the ions and their discharge at the electrodes.

In the aluminium chloride and phosphorus(V) chloride cases, the solid doesn't conduct electricity because the ions aren't free to move. In the liquid (where it exists - both of these sublime at ordinary pressures), they have converted into a covalent form, and so don't conduct either.

The rest of the chlorides don't conduct electricity either solid or molten because they don't have any ions or any mobile electrons.

Reactions with water

As an approximation, the simple ionic chlorides (sodium and magnesium chloride) just dissolve in water.

The other chlorides all react with water in a variety of ways described below for each individual chloride. The reaction with water is known as hydrolysis.

It is essential to know what your examiners expect. You should obviously check your syllabus, but you also need to look at past papers and mark schemes so that you know what your examiners are actually asking. If you are working towards a UK-based exam and haven't got any of these things follow this link before you go any further to find out how to get them.

The individual chlorides

Sodium chloride, NaCl

Sodium chloride is a simple ionic compound consisting of a giant array of sodium and chloride ions.

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

This is normally drawn in an exploded form as:

The strong attractions between the positive and negative ions need a lot of heat energy to break, and so sodium chloride has high melting and boiling points.

It doesn't conduct electricity in the solid state because it hasn't any mobile electrons and the ions aren't free to move. However, when it melts it undergoes electrolysis.

Sodium chloride simply dissolves in water to give a neutral solution.

Magnesium chloride, MgCl₂

Magnesium chloride is also ionic, but with a more complicated arrangement of the ions to allow for having twice as many chloride ions as magnesium ions. This structure isn't needed for UK A level purposes.

Again, lots of heat energy is needed to overcome the attractions between the ions, and so the melting and boiling points are again high.

The most likely explanation for this is that magnesium chloride is less purely ionic than we normally suggest, and shows some small degree of covalency. That means that you can't make a simple comparison between the melting and boiling points of magnesium chloride and the more purely ionic sodium chloride.

Solid magnesium chloride is a non-conductor of electricity because the ions aren't free to move. However, it undergoes electrolysis when the ions become free on melting.

Magnesium chloride dissolves in water to give a faintly acidic solution (pH = approximately 6).

When magnesium ions are broken off the solid lattice and go into solution, there is enough attraction between the 2+ ions and the water molecules to get co-ordinate (dative covalent) bonds formed between the magnesium ions and lone pairs on surrounding water molecules.

Hexaaquamagnesium ions are formed, [Mg(H₂O)₆]²⁺.

MgCl₂(s)  +  6H₂O(l)    [Mg(H₂O)₆]²⁺(aq)  +  2Cl^-(aq)

Ions of this sort are acidic - the degree of acidity depending on how much the electrons in the water molecules are pulled towards the metal at the centre of the ion. The hydrogens are made rather more positive than they would otherwise be, and more easily pulled off by a base.

In the magnesium case, the amount of distortion is quite small, and only a small proportion of the hydrogen atoms are removed by a base - in this case, by water molecules in the solution.

The presence of the hydroxonium ions in the solution causes it to be acidic. The fact that there aren't many of them formed (the position of equilibrium lies well to the left), means that the solution is only weakly acidic.

You may also find the last equation in a simplified form:

[Mg(H₂O)₆]²⁺(aq)    [Mg(H₂O)₅(OH)]⁺(aq)  +  H⁺(aq)

Hydrogen ions in solution are hydroxonium ions. If you use this form, it is essential to include the state symbols.

Aluminium chloride, AlCl₃

Electronegativity increases as you go across the period and, by the time you get to aluminium, there isn't enough electronegativity difference between aluminium and chlorine for there to be a simple ionic bond.

Aluminium chloride is complicated by the way its structure changes as temperature increases.

At room temperature, the aluminium in aluminium chloride is 6-coordinated. That means that each aluminium is surrounded by 6 chlorines. The structure is an ionic lattice - although with a lot of covalent character.

At ordinary atmospheric pressure, aluminium chloride sublimes (turns straight from solid to vapour) at about 180°C. If the pressure is raised to just over 2 atmospheres, it melts instead at a temperature of 192°C.

Both of these temperatures, of course, are completely wrong for an ionic compound - they are much too low. They suggest comparatively weak attractions between molecules - not strong attractions between ions.

The coordination of the aluminium changes at these temperatures. It becomes 4-coordinated - each aluminium now being surrounded by 4 chlorines rather than 6.

What happens is that the original lattice has converted into Al₂Cl₆ molecules. If you have read the page on co-ordinate bonding mentioned above, you will have seen that the structure of this is:

This conversion means, of course, that you have completely lost any ionic character - which is why the aluminium chloride vaporises or melts (depending on the pressure).

There is an equilibrium between these dimers and simple AlCl₃ molecules. As the temperature increases further, the position of equilibrium shifts more and more to the right.

Al₂Cl₆    2AlCl₃

Summary

• At room temperature, solid aluminium chloride has an ionic lattice with a lot of covalent character.

• At temperatures around 180 - 190°C (depending on the pressure), aluminium chloride coverts to a molecular form, Al₂Cl₆. This causes it to melt or vaporise because there are now only comparatively weak intermolecular attractions.

• As the temperature increases a bit more, it increasingly breaks up into simple AlCl₃ molecules.

Solid aluminium chloride doesn't conduct electricity at room temperature because the ions aren't free to move. Molten aluminium chloride (only possible at increased pressures) doesn't conduct electricity because there aren't any ions any more.

The reaction of aluminium chloride with water is dramatic. If you drop water onto solid aluminium chloride, you get a violent reaction producing clouds of steamy fumes of hydrogen chloride gas.

If you add solid aluminium chloride to an excess of water, it still splutters, but instead of hydrogen chloride gas being given off, you get an acidic solution formed. A solution of aluminium chloride of ordinary concentrations (around 1 mol dm^-3, for example) will have a pH around 2 - 3. More concentrated solutions will go lower than this.

The aluminium chloride reacts with the water rather than just dissolving in it. In the first instance, hexaaquaaluminium ions are formed together with chloride ions.

AlCl₃(s)  +  6H₂O(l)    [Al(H₂O)₆]³⁺(aq)  +  3Cl^-(aq)

You will see that this is very similar to the magnesium chloride equation given above - the only real difference is the charge on the ion.

That extra charge pulls electrons from the water molecules quite strongly towards the aluminium. That makes the hydrogens more positive and so easier to remove from the ion. In other words, this ion is much more acidic than in the corresponding magnesium case.

These equilibria (whichever you choose to write) lie further to the right, and so the solution formed is more acidic - there are more hydroxonium ions in it.

or, more simply:

[Al(H₂O)₆]³⁺(aq)    [Al(H₂O)₅(OH)]²⁺(aq)  +  H⁺(aq)

We haven't so far accounted for the burst of hydrogen chloride formed if there isn't much water present.

All that happens is that because of the heat produced in the reaction and the concentration of the solution formed, hydrogen ions and chloride ions in the mixture combine together as hydrogen chloride molecules and are given off as a gas. With a large excess of water, the temperature never gets high enough for that to happen - the ions just stay in solution.

Silicon tetrachloride, SiCl₄

Silicon tetrachloride is a simple no-messing-about covalent chloride. There isn't enough electronegativity difference between the silicon and the chlorine for the two to form ionic bonds.

Silicon tetrachloride is a colourless liquid at room temperature which fumes in moist air. The only attractions between the molecules are van der Waals dispersion forces.

It doesn't conduct electricity because of the lack of ions or mobile electrons.

It fumes in moist air because it reacts with water in the air to produce hydrogen chloride. If you add water to silicon tetrachloride, there is a violent reaction to produce silicon dioxide and fumes of hydrogen chloride. In a large excess of water, the hydrogen chloride will, of course, dissolve to give a strongly acidic solution containing hydrochloric acid.

SiCl₄  +  2H₂O    SiO₂  +  4HCl

The phosphorus chlorides

There are two phosphorus chlorides - phosphorus(III) chloride, PCl₃, and phosphorus(V) chloride, PCl₅.

Phosphorus(III) chloride (phosphorus trichloride), PCl₃

This is another simple covalent chloride - again a fuming liquid at room temperature.

It is a liquid because there are only van der Waals dispersion forces and dipole-dipole attractions between the molecules.

It doesn't conduct electricity because of the lack of ions or mobile electrons.

Phosphorus(III) chloride reacts violently with water. You get phosphorous acid, H₃PO₃, and fumes of hydrogen chloride (or a solution containing hydrochloric acid if lots of water is used).

PCl₃  +  3H₂O    H₃PO₃  +  3HCl

Phosphorus(V) chloride (phosphorus pentachloride), PCl₅

Unfortunately, phosphorus(V) chloride is structurally more complicated.

Phosphorus(V) chloride is a white solid which sublimes at 163°C. The higher the temperature goes above that, the more the phosphorus(V) chloride dissociates (splits up reversibly) to give phosphorus(III) chloride and chlorine.

PCl₅    PCl₃  +  Cl₂

Solid phosphorus(V) chloride contains ions - which is why it is a solid at room temperature. The formation of the ions involves two molecules of PCl₅.

A chloride ion transfers from one of the original molecules to the other, leaving a positive ion, [PCl₄]⁺, and a negative ion, [PCl₆]^-.

At 163°C, the phosphorus(V) chloride converts to a simple molecular form containing PCl₅ molecules. Because there are only van der Waals dispersion forces between these, it then vaporises.

Solid phosphorus(V) chloride doesn't conduct electricity because the ions aren't free to move.

Phosphorus(V) chloride has a violent reaction with water producing fumes of hydrogen chloride. As with the other covalent chlorides, if there is enough water present, these will dissolve to give a solution containing hydrochloric acid.

The reaction happens in two stages. In the first, with cold water, phosphorus oxychloride, POCl₃, is produced along with HCl.

PCl₅  +  H₂O    POCl₃  +  2HCl

If the water is boiling, the phosphorus(V) chloride reacts further to give phosphoric(V) acid and more HCl. Phosphoric(V) acid is also known just as phosphoric acid or as orthophosphoric acid.

POCl₃  +  3H₂O    H₃PO₄  +  3HCl

The overall equation in boiling water is just a combination of these:

PCl₅  +  4H₂O    H₃PO₄  +  5HCl

Disulphur dichloride, S₂Cl₂

Disulphur dichloride is just one of three sulphur chlorides, but is the only one mentioned by any of the UK A level syllabuses. This is possibly because it is the one which is formed when chlorine reacts with hot sulphur.

Disulphur dichloride is a simple covalent liquid - orange and smelly!

The shape is surprisingly difficult to draw convincingly! The atoms are all joined up in a line - but twisted:

The reason for drawing the shape is to give a hint about what sort of intermolecular attractions are possible. There is no plane of symmetry in the molecule and that means that it will have an overall permanent dipole.

The liquid will have van der Waals dispersion forces and dipole-dipole attractions.

There are no ions in disulphur dichloride and no mobile electrons - so it never conducts electricity.

Disulphur dichloride reacts slowly with water to produce a complex mixture of things including hydrochloric acid, sulphur, hydrogen sulphide and various sulphur-containing acids and anions (negative ions). There is no way that you can write a single equation for this - and one would never be expected in an exam.