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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:
NaCl MgCl2 AlCl3 SiCl4 PCl5 S2Cl2
PCl3
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.
Note: Follow this link if you aren't sure about intermolecular attractions such as van der Waals dispersion forces.
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Electrical conductivity
Sodium and magnesium chlorides are ionic and so will undergoelectrolysis 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.
Tho 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.
Warning: The rest of this page contains quite a lot of detail about the various chlorides, covering material from all the UK A level (or its equivalent) syllabuses. It is very unlikely that you will need all of this, and it is quite possible that your examiners will allow (or even expect) simplifications in some cases.
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.
It would also be useful to look at books written specifically for your actual syllabus. These will have been checked by your examiners, and they can hardly argue with anything you find in them. Look at the text book suggestions page to find some of the available books.
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.
Note: You will find the structure and physical properties of sodium chloride dealt with in a bit more detail (including an explanation of how to draw the last diagram) by following this link.
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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.
Note: There is a problem here, though! You would expect the attractions between magnesium ions and chloride ions to be greater than those between sodium and chloride ions due to the extra charge on the magnesium. However, magnesium chloride melts at a lower temperature than sodium chloride, and the boiling points are almost identical (to within one degree).
I have no idea what the explanation for this is, unless magnesium chloride isn't as ionic as we usually pretend. If anyone has any reliable information on this could they contact me via the address on the about this site page.
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).
Note: This is one point when you need to know exactly what your examiners want you to say about this by looking at yoursyllabus, past papers and mark schemes.
Some examiners simply say that magnesium chloride just dissolves in water. However, that wouldn't account for the slightly lowered pH. On the other hand, there is no point in learning a complicated bit of chemistry if all you need is a simplification. Be aware that it is a simplification, though.
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+.
Note: You will find the bonding in ions of this sort discussed with reference to the corresponding aluminium ion on the page about co-ordinate (dative covalent) bonding. The magnesium case is exactly the same.
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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.
Note: The reason for the colour-coding is to try to avoid confusion between the water molecules attached to the ion and those 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. If you use this form, it is essential to include the state symbols.
Note: You will find lots more about the acidity of hexaaqua ions by following this link.
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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
Al2Cl6molecules. 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.
Note: One very reliable source says that although solid aluminium chloride has zero conductivity at room temperature, it conducts just below the melting point. I haven't at the moment been able to confirm this - neither do I have any idea why it might happen.
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.
Note: The phosphorus(III) chloride molecule has a permanent dipole, which is why dipole-dipole attractions are possible. There is a discussion about polar molecules and polar bonds on the page about electronegativity. The phosphorus(III) chloride case is rather similar to CHCl3 (discussed on that page), except that there is a lone pair of electrons at the top of the molecule
rather than a hydrogen atom.
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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).
Note: Phosphorous acid is also known as orthophosphorous acid or as phosphonic acid. Notice the "-ous" ending in the first two names. That's not a spelling mistake - it's for real! It is used to distinguish it from phosphoric acid which is quite different (see below).
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.
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.
Note: Phosphorus(V) chloride does, however, undergo electrolysis in a suitable solvent which it doesn't react with. For example, Inorganic Chemistry by Heslop and Robinson quotes it as conducting electricity in solution in methyl nitrite, CH3ONO.
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:
Note: I'm not entirely happy about the conditions for these reactions. Several sources mention the need for boiling water for the second half of the reaction, but some don't quote any temperature. All of the safety data sheets available on the web talk about phosphorus oxychloride reacting strongly with water without any suggestion that it needs to be heated.
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.
Where would you like to go now?
To the Period 3 menu . . .
To the Inorganic Chemistry menu . . .
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© Jim Clark 2005
CHEMICAL REACTIONS OF THE PERIOD 3 ELEMENTS
This page describes the reactions of the Period 3 elements from sodium to argon with water, oxygen and chlorine.
Reactions with water
Sodium
Sodium has a very exothermic reaction with cold water producing hydrogen and a colourless solution of sodium hydroxide.
Magnesium
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.
Magnesium burns in steam with its typical white flame to produce white magnesium oxide and hydrogen.
Note: If you are heating the magnesium in a glass tube, the magnesium also reacts with the glass. That leaves dark grey products (including silicon and perhaps boron from the glass) as well as the white magnesium oxide.
Notice also that the oxide is produced on heating in steam. Hydroxides are
only ever produced using liquid water.
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
There is a fair amount of disagreement in the books and on the web about what silicon does with water or steam. 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.
Note: These more reactive forms are produced as powders. Cotton and Wilkinson's Advanced Inorganic Chemistry (third edition - page 316) suggests that the reactivity of one of these could be due to a very high surface area, or perhaps because the silicon exists in a graphite-like structure.
A correspondent from the silicon industry tells me that when silicon is cut into slices, the silicon dust formed reacts with water at room temperature - producing hydrogen and getting very hot. He says
"The silicon is cut in a glycol slurry [. . .] The powdered Si is protected somewhat from moisture in the glycol slurry, but when we clean the slurry in aqueous solutions the reaction with water takes off."
This is probably the effect of the high surface area of the dust produced, combined with the fact that you are exposing uncontaminated silicon to the water. One source suggests that the lack of reactivity of silicon is due to a
layer of silicon dioxide on its surface. If you expose a new surface by cutting the silicon, that layer won't, of course, exist.
Phosphorus and sulphur
These have no reaction with water.
Chlorine
Chlorine dissolves in water to some extent to give a green solution. A reversible reaction takes place to produce a mixture of hydrochloric acid and chloric(I) acid (hypochlorous acid).
Note: You may also find the chloric(I) acid written as HClO. The form I have used more accurately reflects the way the atoms are joined up. It doesn't matter which you use.
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:
Argon
There is no reaction between argon and water.
Reactions with oxygen
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:
Magnesium
Magnesium burns in oxygen with an intense white flame to give white solid magnesium oxide.
Note: If magnesium is burns in air rather than in pure oxygen, it also reacts with the nitrogen in the air. You get a mixture of magnesium oxide and magnesium nitride formed.
Aluminium
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.
Silicon
Silicon will burn in oxygen if heated strongly enough. Silicon dioxide is produced.
Note: There is disagreement between various web or textbook sources about the temperature needed to ignite the silicon, varying from 400°C to well over 1000°C. In fact, there isn't a "right" answer to this. It depends on what sort of silicon you are talking about and how finely divided it is. For example, one of the amorphous (non-crystalline powder) forms of silicon even catches fire spontaneously in air at room temperature. Other forms need higher temperatures and a richer oxygen supply.
Phosphorus
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:
Note: You may come across these oxides written as P2O3and P2O5. Don't use these forms! They are as logical as writing, say, ethene as CH2 and ethane as CH3.
Sulphur
Sulphur burns in air or oxygen on gentle heating with a pale blue flame. It produces colourless sulphur dioxide gas.
Note: Sulphur dioxide can, of course, be converted further into sulphur trioxide in the presence of oxygen, but it needs the presence of a catalyst and fairly carefully controlled conditions. If you are interested in this, see the page on theContact Process.This isn't particularly relevant to the current topic, but if you do feel the urge to follow this link, use the BACK button on your browser to return to this page.
Chlorine and argon
Despite having several oxides, chlorine won't react directly with oxygen. Argon doesn't react either.
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.
Note: You may find versions of this equation showing the aluminium chloride as Al2Cl6. In fact, this exists in the vapourat temperatures not too far above the sublimation temperature - not in the solid. The structure of aluminium chloride is discussed on the page about Period 3 chlorides.
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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 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.
Note: These equations are often given starting from P rather than P4. It depends which form of phosphorus you are talking about.
If you are talking about white phosphorus (as I am here), P4is the correct version. If you are talking about red phosphorus, then P is correct. Red phosphorus has a different (polymeric) structure, and P4 would be wrong for it.
In my experience, red phosphorus is less commonly used in labs at this level (it isn't as excitingly reactive as white phosphorus!) - which is why I am concentrating on the white form.
Sulphur
If a stream of chlorine is passed over some heated sulphur, it reacts to form an orange, evil-smelling liquid, disulphur dichloride, S2Cl2.
Chlorine and argon
It obviously doesn't make sense to talk about chlorine reacting with itself, and argon doesn't react with chlorine.
