THE OXIDES OF PERIOD 3 ELEMENTS

1. Formation of oxides

All the elements in Period 3 except chlorine and argon combine directly with oxygen to form oxides.

4Na(s) + O2(g) 2Na2O(s)Na2O is an ionic oxide.

Sodium

Sodium burns in oxygen with an orange flame to produce a white solid mixture of sodium oxide and sodium peroxide.

For the simple oxide:

For the peroxide:

2Mg(s) + O2(g) 2MgO(s)MgO is also an ionic oxide.

Magnesium burns in oxygen with an intense white flame to give white solid magnesium oxide.

4Al(s) + 3O2(g) 2Al2O3(s)Al2O3 is mostly ionic, but there is significant covalent character.

Aluminium will burn in oxygen if it is powdered, otherwise the strong oxide layer on the aluminium tends to inhibit the reaction. If you sprinkle aluminium powder into a Bunsen flame, you get white sparkles. White aluminium oxide is formed.

Si(s) + O2(g) SiO2(s)

SiO2 is a giant covalent oxide.Silicon will burn in oxygen if heated strongly enough. Silicon dioxide is produced.

P4(s) + 5O2(g) P4O10(s)P4O10 is a molecular covalent oxide. The oxidation number of P in this oxide is +5.

White phosphorus catches fire spontaneously in air, burning with a white flame and producing clouds of white smoke - a mixture of phosphorus(III) oxide and phosphorus(V) oxide.

The proportions of these depend on the amount of oxygen available. In an excess of oxygen, the product will be almost entirely phosphorus(V) oxide.

For the phosphorus(III) oxide:

For the phosphorus(V) oxide:

S(s) + O2(g) SO2(g)SO2 is a molecular covalent oxide.

Sulphur burns in air or oxygen on gentle heating with a pale blue flame. It produces colourless sulphur dioxide gas.

Another oxide, SO3 is formed in a reversible process when SO2 and O2 are heated with a V2O5 catalyst (the Contact Process)

2. Physical properties of oxides

The physical properties of these oxides depend on the type of bonding.

Na2O, Al2O3 and MgO are ionic oxides and hence have a high melting point. MgO and Al2O3 have a higher melting point than Na2O since the charges are higher, resulting in a stonger attraction between the ions.

SiO2 has a giant covalent structure and hence a high melting point. There are strong covalent bonds between all the atoms and thus lots of energy is required to break them.

P4O10 and SO3 are molecular covalent and so only intermolecular forces exist between the molecules. The melting points are thus much lower. P4O10 is a much bigger molecule than SO3 and so has a much higher melting point, as the van der Waal’s forces are stronger.

Element Na Mg Al Si P SFormulae of oxide

Na2O MgO Al2O3 SiO2 P4O10 SO3

Structure of oxide Ionic Ionic Mostly ionic

Giant covalent

Molecular covalent

Molecular covalent

Melting point of oxide /°C

1275 2852 2072 1703 300 -10

Electrical conductivity

None of these oxides has any free or mobile electrons. That means that none of them will conduct electricity when they are solid.

The ionic oxides can, however, undergo electrolysis when they are molten. They can conduct electricity because of the movement of the ions towards the electrodes and the discharge of the ions when they get there.

Whether you can electrolyse molten sodium oxide depends, of course, on whether it actually melts instead of subliming or decomposing under ordinary circumstances. If it sublimes, you won't get any liquid to electrolyse.

The metallic oxides

The structures

Sodium, magnesium and aluminium oxides consist of giant structures containing metal ions and oxide ions. Magnesium oxide has a structure just like sodium chloride.

Silicon dioxide (silicon(IV) oxide)

The structure

The electronegativity of the elements increases as you go across the period, and by the time you get to silicon, there isn't enough electronegativity difference between the silicon and the oxygen to form an ionic bond. Silicon dioxide is a giant covalent structure.

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 in all 3 dimensions.

Melting and boiling points

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 they are all around

1700°C. Very strong silicon-oxygen covalent bonds have to be broken throughout the structure before melting occurs. Silicon dioxide boils at 2230°C.

Because you are talking about a different form of bonding, it doesn't make sense to try to compare these values directly with the metallic oxides. What you can safely say is that because the metallic oxides and silicon dioxide have giant structures, the melting and boiling points are all high.

Electrical conductivity

Silicon dioxide doesn't have any mobile electrons or ions - so it doesn't conduct electricity either as a solid or a liquid.

The molecular oxides

Phosphorus, sulphur and chlorine all form oxides which consist of molecules. Some of these molecules are fairly simple - others are polymeric. We are just going to look at some of the simple ones.

Melting and boiling points of these oxides will be much lower than those of the metal oxides or silicon dioxide. The intermolecular forces holding one molecule to its neighbours will be van der Waals dispersion forces or dipole-dipole interactions. The strength of these will vary depending on the size of the molecules.

None of these oxides conducts electricity either as solids or as liquids. None of them contains ions or free electrons.

The phosphorus oxides

Phosphorus has two common oxides, phosphorus(III) oxide, P4O6, and phosphorus(V) oxide, P4O10.

Phosphorus(III) oxide

Phosphorus(III) oxide is a white solid, melting at 24°C and boiling

at 173°C.

The structure of its molecule is best worked out starting from a P4 molecule which is a little tetrahedron.

Pull this apart so that you can see the bonds . . .

. . . and then replace the bonds by new bonds linking the phosphorus atoms via oxygen atoms. These will be in a V-shape (rather like in water), but you probably wouldn't be penalised if you drew them on a straight line between the phosphorus atoms in an exam.

The phosphorus is using only three of its outer electrons (the 3 unpaired p electrons) to form bonds with the oxygens.

Phosphorus(V) oxide

Phosphorus(V) oxide is also a white solid, subliming (turning straight from solid to vapour) at 300°C. In this case, the phosphorus uses all five of its outer electrons in the bonding.

Solid phosphorus(V) oxide exists in several different forms - some of them polymeric. We are going to concentrate on a simple molecular form, and this is also present in the vapour.

This is most easily drawn starting from P4O6. The other four oxygens are attached to the four phosphorus atoms via double bonds.

The sulphur oxides

Sulphur has two common oxides, sulphur dioxide (sulphur(IV) oxide), SO2, and sulphur trioxide (sulphur(VI) oxide), SO3.

Sulphur dioxide

Sulphur dioxide is a colourless gas at room temperature with an easily recognised choking smell. It consists of simple SO2molecules.

The sulphur uses 4 of its outer electrons to form the double bonds with the oxygen, leaving the other two as a lone pair on the sulphur. The bent shape of SO2 is due to this lone pair.

Sulphur trioxide

Pure sulphur trioxide is a white solid with a low melting and boiling point. It reacts very rapidly with water vapour in the air to form sulphuric acid. That means that if you make some in the lab, you tend to see it as a white sludge which fumes dramatically in moist air (forming a fog of sulphuric acid droplets).

Gaseous sulphur trioxide consists of simple SO3 molecules in which all six of the sulphur's outer electrons are involved in the bonding.

There are various forms of solid sulphur trioxide. The simplest one is a trimer, S3O9, where three SO3 molecules are joined up and arranged in a ring.

There are also other polymeric forms in which the SO3molecules join together in long chains. For example:

The fact that the simple molecules join up in this way to make bigger structures is what makes the sulphur trioxide a solid rather than a gas.

The chlorine oxides

Chlorine forms several oxides. Here we are just looking at two of them (the only ones mentioned by any of the UK syllabuses) - chlorine(I) oxide, Cl2O, and chlorine(VII) oxide, Cl2O7.

Chlorine(I) oxide

Chlorine(I) oxide is a yellowish-red gas at room temperature. It consists of simple small molecules.

There's nothing in the least surprising about this molecule and it's physical properties are just what you would expect for a molecule this size.

Chlorine(VII) oxide

In chlorine(VII) oxide, the chlorine uses all of its seven outer electrons in bonds with oxygen. This produces a much bigger molecule, and so you would expect its melting point and boiling point to be higher than chlorine(I) oxide.

Chlorine(VII) oxide is a colourless oily liquid at room temperature.

In the diagram, for simplicity I have drawn a standard structural formula. In fact, the shape is tetrahedral around both chlorines, and V-shaped around the central oxygen.

3. Acid-base character of oxides

Ionic oxides contain the O2- ion. This is a strongly basic ion which reacts with water to produce hydroxide ions:

O2-(aq) + H2O(l) 2OH-(aq)

Thus all ionic oxides are BASIC.

Covalent oxides do not contain ions, but have a strongly positive dipole on the atom which is not oxygen. This attracts the lone pair on water molecules, releasing H+ ions:

MO(s) + H2O(l) MO(OH)-(aq)+ H+(aq)

Thus all covalent oxides are ACIDIC.

Intermediate oxides can react in either of the above ways, depending on the conditions. They can thus behave as either acids or bases and are thus AMPHOTERIC.

Na2O is a basic oxide. It dissolves in water to give an alkaline solution (pH = 14). It also reacts with acids:

Na2O(s) + H2O(l) 2NaOH(aq)Na2O(s) + 2H+(aq) 2Na+(aq) + H2O(l)

MgO is a basic oxide. It is only slightly soluble in water and so the solution is only slightly alkaline (pH = 9). It reacts readily with acids:

MgO(s) + H2O(l) == Mg(OH)2(s) == Mg(OH)2(aq)MgO(s) + 2H+(aq) Mg2+(aq) + H2O(l)

Al2O3 is an amphoteric oxide. It is insoluble in water (pH = 7) but dissolves in both acids and alkalis:

Al2O3(s) + 6H+(aq) 2Al3+(aq) + 3H2O(l)Al2O3(s) + 3H2O(l) + 6OH-(aq) 2Al(OH)6

3-(aq)Al2O3(s) + 3H2O(l) + 2OH-(aq) 2Al(OH)4

-(aq)

SiO2 is an acidic oxide. It is insoluble in water (pH = 7) but dissolves in hot concentrated alkalis:

SiO2(s) + 2OH-(aq) SiO32-(aq) + H2O(l)

P4O10 is an acidic oxide. It dissolves in water to give acidic solutions and is also soluble in alkalis:

P4O10(s) + 6H2O(l) 4H3PO4(aq) pH = 3P4O10(s) + 12OH-(aq) 4PO4

3-(aq) + 6H2O(l)

SO2 and SO3 are acidic oxides. They dissolve in water to give acidic solutions, and also react with alkalis:

SO2(g) + H2O(l) == H2SO3(aq), pH = 2SO3(g) + H2O(l) H2SO4(aq), pH = 1SO2(g) + 2OH-(aq) SO3

2-(aq) + H2O(l)SO3(g) + 2OH-(aq) SO4

2-(aq) + H2O(l)

SO2 is a waste gas in many industrial processes. It is harmful because it dissolves in rain water to give acid rain. It can be removed from waste gases because it dissolves in alkali and so it is passed through an alkaline solution in waste gas outlets to minimise the amount which escapes into the atmosphere.

The acid-base properties of the oxides of Period 3 can be summarised in the following table:

Element Na Mg Al Si P SFormulae of oxides

Na2O MgO Al2O3 SiO2 P4O10 SO2

SO3

Acid-base character of oxide

Basic Basic Amphoteric Acidic Acidic Acidic

pH of solution when dissolved in water

12 - 14 8 - 9 7(insoluble)

7(insoluble)

2 - 4 2 - 4 (SO2)1 - 3 (SO3)

The oxides therefore become more acidic on moving from left to right in the periodic table.

Sodium oxide

Sodium oxide is a simple strongly basic oxide. It is basic because it contains the oxide ion, O2-, which is a very strong base with a high tendency to combine with hydrogen ions.

Reaction with water

Sodium oxide reacts exothermically with cold water to produce sodium hydroxide solution. Depending on its concentration, this will have a pH around 14.

Reaction with acids

As a strong base, sodium oxide also reacts with acids. For example, it would react with dilute hydrochloric acid to produce sodium chloride solution.

Magnesium oxide

Magnesium oxide is again a simple basic oxide, because it also contains oxide ions. However, it isn't as strongly basic as sodium oxide because the oxide ions aren't so free.

In the sodium oxide case, the solid is held together by attractions between 1+ and 2- ions. In the magnesium oxide case, the attractions are between 2+ and 2-. It takes more energy to break these.

Even allowing for other factors (like the energy released when the positive ions form attractions with water in the solution formed), the net effect of this is that reactions involving magnesium oxide will always be less exothermic than those of sodium oxide.

Reaction with water

If you shake some white magnesium oxide powder with water, nothing seems to happen - it doesn't look as if it reacts. However, if you test the pH of the liquid, you find that it is somewhere around pH 9 - showing that it is slightly alkaline.

There must have been some slight reaction with the water to produce hydroxide ions in solution. Some magnesium hydroxide is formed in the reaction, but this is almost insoluble - and so not many hydroxide ions actually get into solution.

Reaction with acids

Magnesium oxide reacts with acids as you would expect any simple metal oxide to react. For example, it reacts with warm dilute hydrochloric acid to give magnesium chloride solution.

Aluminium oxide

Describing the properties of aluminium oxide can be confusing because it exists in a number of different forms. One of those forms is very unreactive. It is known chemically as alpha-Al2O3and is produced at high temperatures.

In what follows we are assuming one of the more reactive forms.

Aluminium oxide is amphoteric. It has reactions as both a base and an acid.

Reaction with water

Aluminium oxide doesn't react in a simple way with water in the sense that sodium oxide and magnesium oxide do, and doesn't dissolve in it. Although it still contains oxide ions, they are held too strongly in the solid lattice to react with the water.

Reaction with acids

Aluminium oxide contains oxide ions and so reacts with acids in the same way as sodium or magnesium oxides. That means, for example, that aluminium oxide will react with hot dilute hydrochloric acid to give aluminium chloride solution.

In this (and similar reactions with other acids), aluminium oxide is showing the basic side of its amphoteric nature.

Reaction with bases

Aluminium oxide has also got an acidic side to its nature, and it shows this by reacting with bases such as sodium hydroxide solution.

Various aluminates are formed - compounds where the aluminium is found in the negative ion. This is possible because aluminium has the ability to form covalent bonds with oxygen.

In the case of sodium, there is too much electronegativity difference between sodium and oxygen to form anything other than an ionic bond. But electronegativity increases as you go across the period - and the electronegativity difference between aluminium and oxygen is smaller. That allows the formation of covalent bonds between the two.

With hot, concentrated sodium hydroxide solution, aluminium oxide reacts to give a colourless solution of sodium tetrahydroxoaluminate.

Silicon dioxide (silicon(IV) oxide)

By the time you get to silicon as you go across the period, electronegativity has increased so much that there is no longer enough electronegativity difference between silicon and oxygen to form ionic bonds.

Silicon dioxide has no basic properties - it doesn't contain oxide ions and it doesn't react with acids. Instead, it is very weakly acidic, reacting with strong bases.

Reaction with water

Silicon dioxide doesn't react with water, because of the difficulty of breaking up the giant covalent structure.

Reaction with bases

Silicon dioxide reacts with sodium hydroxide solution, but only if it is hot and concentrated. A colourless solution of sodium silicate is formed.

You may also be familiar with one of the reactions happening in the Blast Furnace extraction of iron - in which calcium oxide (from the limestone which is one of the raw materials) reacts with silicon dioxide to produce a liquid slag, calcium silicate. This is also an example of the acidic silicon dioxide reacting with a base.

The phosphorus oxides

We are going to be looking at two phosphorus oxides, phosphorus(III) oxide, P4O6, and phosphorus(V) oxide, P4O10.

Phosphorus(III) oxide

Phosphorus(III) oxide reacts with cold water to give a solution of the weak acid, H3PO3 - known variously as phosphorous acid, orthophosphorous acid or phosphonic acid. Its reaction with hot water is much more complicated.

The pure un-ionised acid has the structure:

The hydrogens aren't released as ions until you add water to the acid, and even then not many are released because phosphorous acid is only a weak acid.

Phosphorous acid has a pKa of 2.00 which makes it stronger than common organic acids like ethanoic acid (pKa = 4.76).

It is unlikely that you would ever react phosphorus(III) oxide directly with a base, but you might need to know what happens if you react the phosphorous acid formed with a base.

In phosphorous acid, the two hydrogen atoms in the -OH groups are acidic, but the other one isn't. That means that you can get two possible reactions with, for example, sodium hydroxide solution depending on the proportions used.

In the first case, only one of the acidic hydrogens has reacted with the hydroxide ions from the base. In the second case (using twice as much sodium hydroxide), both have reacted.

If you were to react phosphorus(III) oxide directly with sodium hydroxide solution rather than making the acid first, you would end up with the same possible salts.

Phosphorus(V) oxide

Phosphorus(V) oxide reacts violently with water to give a solution containing a mixture of acids, the nature of which depends on the conditions. We usually just consider one of these, phosphoric(V) acid, H3PO4 - also known just as phosphoric acid or as orthophosphoric acid.

This time the pure un-ionised acid has the structure:

Phosphoric(V) acid is also a weak acid with a pKa of 2.15. That makes it fractionally weaker than phosphorous acid. Solutions of both of these acids of concentrations around 1 mol dm-3 will have a pH of about 1.

Once again, you are unlikely ever to react this oxide with a base, but you may well be expected to know how phosphoric(V) acid reacts with something like sodium hydroxide solution.

If you look back at the structure, you will see that it has three -OH groups, and each of these has an acidic hydrogen atom. You can get a reaction with sodium hydroxide in three stages, with one after another of these hydrogens reacting with the hydroxide ions.

Again, if you were to react phosphorus(V) oxide directly with sodium hydroxide solution rather than making the acid first, you would end up with the same possible salts.

One example out of the possible equations:

Sulphur dioxide

Sulphur dioxide is fairly soluble in water, reacting with it to give a solution known as sulphurous acid, and traditionally given the formula H2SO3. However, the main species in the solution is simply hydrated sulphur dioxide - SO2, xH2O. It is debatable whether any H2SO3 as such exists at all in the solution.

Sulphurous acid is also a weak acid with a pKa of around 1.8 - very slightly stronger than the two phosphorus-containing acids above. A reasonably concentrated solution of sulphurous acid will again have a pH of about 1.

Sulphur dioxide will also react directly with bases such as sodium hydroxide solution. If sulphur dioxide is bubbled through sodium hydroxide solution, sodium sulphite solution is formed first followed by sodium hydrogensulphite solution when the sulphur dioxide is in excess.

Another important reaction of sulphur dioxide is with the base calcium oxide to form calcium sulphite (calcium sulphate(IV)). This is at the heart of one of the methods of removing sulphur dioxide from flue gases in power stations.

Sulphur trioxide

Sulphur trioxide reacts violently with water to produce a fog of concentrated sulphuric acid droplets.

Pure un-ionised sulphuric acid has the structure:

Sulphuric acid is a strong acid, and solutions will typically have pH's of around 0.

The acid reacts with water to give a hydroxonium ion (a hydrogen ion in solution, if you like) and a hydrogensulphate ion. This reaction is virtually 100% complete.

The second hydrogen is more difficult to remove. In fact the hydrogensulphate ion is a relatively weak acid - similar in strength to the acids we have already discussed on this page. This time you get an equilibrium:

Sulphuric acid, of course, has all the reactions of a strong acid that you are familiar with from introductory chemistry courses. For example, the normal reaction with sodium hydroxide solution is to form sodium sulphate solution - in which both of the acidic hydrogens react with hydroxide ions.

THE REACTION OF PERIOD 3 ELEMENTS WITH WATER

Na, Mg, Al and Si are more electropositive than H and can reduce the water to hydrogen gas:

Na reacts vigorously with water to give the hydroxide and hydrogen:2Na(s) +2H2O(l) 2NaOH(aq) + H2(g)The resulting solution is strongly alkaline, and will have a pH of 14.Sodium has a very exothermic reaction with cold water producing hydrogen and a colourless solution of sodium hydroxide

Magnesium has a very slight reaction with cold water, but burns in steam.

A very clean coil of magnesium dropped into cold water eventually gets covered in small bubbles of hydrogen which float it to the surface.

Magnesium hydroxide is formed as a very thin layer on the magnesium and this tends to stop the reaction.

Mg reacts with steam to give the oxide and hydrogen: white flameMg(s) + H2O(g) MgO(s) + H2(g)The resulting solution is weakly alkaline, since the oxide is slightly basic (pH = 9).

Al and Si also react with steam under certain conditions.

Aluminium

Aluminium powder heated in steam produces hydrogen and aluminium oxide. The reaction is relatively slow because of the existing strong aluminium oxide layer on the metal, and the build-up of even more oxide during the reaction.

Silicon

. The truth seems to depend on the precise form of silicon you are using.

The common shiny grey lumps of silicon with a rather metal-like appearance are fairly unreactive. Most sources suggest that this form of silicon will react with steam at red heat to produce silicon dioxide and hydrogen.

But it is also possible to make much more reactive forms of silicon which will react with cold water to give the same products.

P, S and Cl2 do not reduce water to hydrogen gas. Phosphorus and sulphur do not react with water but chlorine will disproportionate to give an acidic solution: (green)Cl2(g) + H2O(l) HClO(aq) + HCl(aq)The resulting solution contains HCl(aq) and is thus acidic (pH = 2).

In the presence of sunlight, the chloric(I) acid slowly decomposes to produce more hydrochloric acid, releasing oxygen gas, and you may come across an equation showing the overall change:

The reactivity of the elements of period 3 towards water thus decreases from Na to Si, and then increases from P to Cl. The resulting solutions become increasingly acidic.

Reactions with chlorine

Sodium

Sodium burns in chlorine with a bright orange flame. White solid sodium chloride is produced.

Magnesium

Magnesium burns with its usual intense white flame to give white magnesium chloride.

Aluminium

Aluminium is often reacted with chlorine by passing dry chlorine over aluminium foil heated in a long tube. The aluminium burns in the stream of chlorine to produce very pale yellow aluminium chloride. This sublimes (turns straight from solid to vapour and back again) and collects further down the tube where it is cooler.

Silicon

If chlorine is passed over silicon powder heated in a tube, it reacts to produce silicon tetrachloride. This is a colourless liquid which vaporises and can be condensed further along the apparatus.

Phosphorus

White phosphorus burns spontaneously in chlorine to produce a mixture of two chlorides, phosphorus(III) chloride and phosphorus(V) chloride (phosphorus trichloride and phosphorus pentachloride).

Phosphorus(III) chloride is a colourless fuming liquid.

Phosphorus(V) chloride is an off-white (going towards yellow) solid.

Sulphur

If a stream of chlorine is passed over some heated sulphur, it reacts to form an orange, evil-smelling liquid, disulphur dichloride, S2Cl2.

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.

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.

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.

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

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+.

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:

Hydrogen ions in solution are hydroxonium ions.

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 AlCl3molecules. As the temperature increases further, the position of equilibrium shifts more and more to the right.

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.

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:

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.

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).

Phosphorus(V) chloride (phosphorus pentachloride), PCl5

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.

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.

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.

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

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).