reactions of aluminum chloride in o level chemistry

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.

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.
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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.