The Chemistry of Period 3 Elements

ATOMIC AND PHYSICAL PROPERTIES OF THE PERIOD 3 ELEMENTS

Atomic Properties

Electronic Structures

In Period 3 of the Periodic Table, the 3s and 3p orbitals are filling with electrons. The
shortened versions of the electronic structures for the eight elements are:

Na [Ne] 3s1

Mg [Ne] 3s2

Al [Ne] 3s2 3px1

Si [Ne] 3s2 3px1 3py1

P [Ne] 3s2 3px1 3py1 3pz1

S [Ne] 3s2 3px2 3py1 3pz1

Cl [Ne] 3s2 3px2 3py2 3pz1

Ar [Ne] 3s2 3px2 3py2 3pz2

In each case, [Ne] represents the complete electronic structure of a neon atom.

First ionisation energy

The first ionisation energy is the energy required to remove one mole of electrons from one
mole of gaseous atoms to produce 1 mole of gaseous ions each with a charge of 1+. This can
be represented by the equation:

X(g) X+(g) + e-

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The Chemistry of Period 3 Elements
The pattern of first ionisation energies across Period 3

The general trend is upwards, but this is broken by falls between magnesium and aluminium,
and between phosphorus and sulphur. First ionisation energy is governed by:

 the charge on the nucleus;

 the distance of the outer electron from the nucleus;
 the amount of screening by inner electrons;
 whether the electron is alone in an orbital or one of a pair.

The upward trend

In the whole of period 3, the outer electrons are in 3rd energy level orbitals which are all
roughly the same of distances from the nucleus, and are screened by the same electrons in the
first and second energy levels. The major difference is the increasing number of protons in
the nucleus as you go from sodium across to argon. That causes greater attraction between the
nucleus and the electrons and so increases the ionisation energies. In fact the increasing
nuclear charge also drags the outer electrons in closer to the nucleus. That increases
ionisation energies still more as you go across the period.

The fall at aluminium

Aluminium is expected to have a value more than the magnesium value because of the extra
proton. Offsetting that is the fact that aluminium's outer electron is in a 3p orbital rather than
a 3s. The 3p electron is slightly more distant from the nucleus than the 3s, and partially
screened by the 3s electrons as well as the inner electrons. Both of these factors offset the
effect of the extra proton.

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The Chemistry of Period 3 Elements
The fall at sulphur

As you go from phosphorus to sulphur, something extra must be offsetting the effect of the
extra proton. The screening is identical in phosphorus and sulphur (from the inner electrons
and, to some extent, from the 3s electrons), and the electron is being removed from an
identical orbital. The difference is that in the sulphur case the electron being removed is one
of the 3px2 pair. The repulsion between the two electrons in the same orbital means that the
electron is easier to remove than it would otherwise be.

Atomic radius

The trend

The diagram shows how the atomic radius changes as you go across Period 3.

The figures used to construct this diagram are based on:

 metallic radii for Na, Mg and Al;

 covalent radii for Si, P, S and Cl;
 the van der Waals radius for Ar because it doesn't form any strong bonds.

It is fair to compare metallic and covalent radii because they are both being measured in
tightly bonded circumstances. It is not fair to compare these with a van der Waals radius,
though. The general trend towards smaller atoms across the period is NOT broken at argon.
You aren't comparing like with like. The only safe thing to do is to ignore argon in the
discussion which follows.

Explaining the trend

A metallic or covalent radius is going to be a measure of the distance from the nucleus to the
bonding pair of electrons. From sodium to chlorine, the bonding electrons are all in the 3rd
energy level, being screened by the electrons in the first and second levels. The increasing
number of protons in the nucleus as you go across the period pulls the bonding electrons
more tightly to it. The amount of screening is constant for all of these elements.

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The Chemistry of Period 3 Elements
Electronegativity

Electronegativity is a measure of the tendency of an atom to attract a bonding pair of

electrons. The Pauling scale is the most commonly used. Fluorine (the most electronegative
element) is assigned a value of 4.0, and values range down to caesium and francium which
are the least electronegative at 0.7.

H2.1 Pauling’s Electronegativity Scale

Li1.0 Be1.5 B2.0 C2.5 N3.0 O3.5 F4.0

Na0.9 Mg1.2 Al1.5 Si1.8 P2.1 S2.5 Cl3.0

The trend

The trend across Period 3 looks like this:

Electronegativity is about the tendency of an atom to attract a bonding pair of electrons. Since
argon does not form covalent bonds, it obviously cannot be assigned an electronegativity
value.

Explaining the trend

The trend is explained in exactly the same way as the trend in atomic radii. As you go across
the period, the bonding electrons are always in the same energy level. They are always being
screened by the same inner electrons. All that differs is the number of protons in the nucleus.
As you go from sodium to chlorine, the number of protons steadily increases and so attracts
4
the bonding pair more closely.

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The Chemistry of Period 3 Elements
Physical Properties

These properties include the electrical conductivity and the melting and boiling points of the
elements. To understand these, you first have to understand the structure of each of the
elements.

Structures of the elements

The structures of the elements change as you go across the period. The first three are metallic,
silicon is giant covalent, and the rest are simple molecules.

Three metallic structures

Sodium, magnesium and aluminium all have metallic structures. In sodium, only one electron
per atom is involved in the metallic bond. In magnesium, both of its outer electrons are
involved, and in aluminium all three. The other difference you need to be aware of is the way
the atoms are packed in the metal crystal. Sodium is 8-co-ordinated - each sodium atom is
touched by only 8 other atoms. Both magnesium and aluminium are 12-co-ordinated
(although in slightly different ways). This is a more efficient way to pack atoms, leading to
less wasted space in the metal structures and to stronger bonding in the metal.

A giant covalent structure

Silicon has a giant covalent structure just like diamond. A tiny part of the structure looks like
this:

The structure is held together by strong covalent bonds in all three dimensions.

| Notes By TP Shumba
The Chemistry of Period 3 Elements
Four simple molecular structures

The structures of phosphorus and sulphur vary depending on the type of phosphorus or
sulphur you are talking about. For phosphorus, I am assuming the common white phosphorus.
For sulphur, I am assuming one of the crystalline forms - rhombic or monoclinic sulphur.

The atoms in each of these molecules are held together by covalent bonds (apart, of course,
from argon). In the liquid or solid state, the molecules are held close to each other by van der
Waals dispersion forces.

Electrical conductivity

 Sodium, magnesium and aluminium are all good conductors of electricity.

Conductivity increases as you go from sodium to magnesium to aluminium.
 Silicon is a semiconductor.
 None of the rest conduct electricity.

The three metals, of course, conduct electricity because the delocalised electrons (the "sea of
electrons") are free to move throughout the solid or the liquid metal. In the silicon case,
explaining how semiconductors conduct electricity is beyond the scope of A level chemistry
courses. With a diamond structure, you mightn't expect it to conduct electricity, but it does.
The rest don't conduct electricity because they are simple molecular substances. There are no
electrons free to move around.

| Notes By TP Shumba
The Chemistry of Period 3 Elements
Melting and boiling points

The chart shows how the melting and boiling points of the elements change as you go across
the period. The figures are plotted in kelvin rather than °C to avoid having negative values.

It is best to think of these changes in terms of the types of structure that the elements have.

The metallic structures

Melting and boiling points rise across the three metals because of the increasing strength of
the metallic bonds. The number of electrons which each atom contributes to the delocalised
"sea of electrons" increases. The atoms also get smaller and have more protons as you go
from sodium to magnesium to aluminium. The attractions and therefore the melting and
boiling points increase because:

 The nuclei of the atoms are getting more positively charged.

 The "sea" is getting more negatively charged.
 The "sea" is getting progressively nearer to the nuclei and so more strongly attracted.

Silicon

Silicon has high melting and boiling points because it is a giant covalent structure. You have
to break strong covalent bonds before it will melt or boil. Because you are talking about a
different type of bond, it is not profitable to try to directly compare silicon's melting and
boiling points with aluminium's.

| Notes By TP Shumba
The Chemistry of Period 3 Elements
The four molecular elements

Phosphorus, sulphur, chlorine and argon are simple molecular substances with only van der
Waals attractions between the molecules. Their melting or boiling points will be lower than
those of the first four members of the period which have giant structures. The sizes of the
melting and boiling points are governed entirely by the sizes of the molecules. Remember the
structures of the molecules:

Phosphorus

Phosphorus contains P4 molecules. To melt phosphorus you don't have to break any covalent
bonds - just the much weaker van der Waals forces between the molecules.

Sulphur

Sulphur consists of S8 rings of atoms. The molecules are bigger than phosphorus molecules,
and so the van der Waals attractions will be stronger, leading to a higher melting and boiling
point.

Chlorine

Chlorine, Cl2, is a much smaller molecule with comparatively weak van der Waals attractions,
and so chlorine will have a lower melting and boiling point than sulphur or phosphorus.

Argon

Argon molecules are just single argon atoms, Ar. The scope for van der Waals attractions
between these is very limited and so the melting and boiling points of argon are lower again.

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The Chemistry of Period 3 Elements
CHEMICAL REACTIONS OF THE PERIOD 3 ELEMENTS

Reactions with water

Sodium

Sodium has a very exothermic reaction with cold water producing hydrogen and a colourless
solution of sodium hydroxide.

2Na + 2H2O 2NaOH + 2H2

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.

Mg + 2H2O Mg(OH)2 + H2

Magnesium burns in steam with its typical white flame to produce white magnesium oxide
and hydrogen.

Mg + H2O MgO + H2

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.

2Al + 3H2O Al2O3 + 3H2

Silicon

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.

Si + 2H2O SiO2 + 2H2

9
But it is also possible to make much more reactive forms of silicon which will react with cold
water to give the same products.

| Notes By TP Shumba
The Chemistry of Period 3 Elements
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).

Cl2 + H2O HCl + HOCl

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:

2Cl2 + 2H2O 4HCl + O2

Argon

There is no reaction between argon and water.

10

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The Chemistry of Period 3 Elements
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:

4Na + O2 2Na2O

For the peroxide:

2Na + O2 Na2O2

Magnesium

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

2Mg + O2 2MgO

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.

4Al + 3O2 2Al2O3

Silicon

Silicon will burn in oxygen if heated strongly enough. Silicon dioxide is produced.

Si + O2 SiO2

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. 11

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The Chemistry of Period 3 Elements
For the phosphorus (III) oxide:

P4 + 3O2 P4O6

For the phosphorus(V) oxide:

P4 + 5O2 P4O10

Sulphur

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

S + O2 SO2

Chlorine and Argon

Despite having several oxides, chlorine won't react directly with oxygen. Argon doesn't react
either.

PHYSICAL PROPERTIES OF THE PERIOD 3 OXIDES

The oxides we'll be looking at are:

Na2O MgO Al2O3 SiO2 P4O10 SO3 Cl2O7

P4O6 SO2 Cl2O

Those oxides in the top row are known as the highest oxides of the various elements. These
are the oxides where the Period 3 elements are in their highest oxidation states. In these
oxides, all the outer electrons in the Period 3 element are being involved in the bonding, from
just the one with sodium, to all seven of chlorine's outer electrons.

The Structures

The trend in structure is from the metallic oxides containing giant structures of ions on the
left of the period via a giant covalent oxide (silicon dioxide) in the middle to molecular
oxides on the right.

Melting and Boiling Points 12

| Notes By TP Shumba
The Chemistry of Period 3 Elements
The giant structures (the metal oxides and silicon dioxide) will have high melting and boiling
points because a lot of energy is needed to break the strong bonds (ionic or covalent)
operating in three dimensions.

The oxides of phosphorus, sulphur and chlorine consist of individual molecules (some small
and simple while others are polymeric). The attractive forces between these molecules will be
van der Waals dispersion and dipole-dipole interactions. These vary in size depending on the
size, shape and polarity of the various molecules, but will always be much weaker than the
ionic or covalent bonds you need to break in a giant structure. These oxides tend to be gases,
liquids or low melting point solids.

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.

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.

Melting and boiling points

There are strong attractions between the ions in each of these oxides and these attractions
need a lot of heat energy to break. These oxides therefore have high melting and boiling
points.

Electrical conductivity

None of these conducts electricity in the solid state, but electrolysis is possible if they are
molten. They conduct electricity because of the movement and discharge of the ions present.
The only important example of this is in the electrolysis of aluminium oxide in the
manufacture of aluminium. 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. Magnesium and
aluminium oxides have melting points far too high to be able to electrolyse them in a simple 13
lab.

Silicon dioxide (silicon(IV) oxide)

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The Chemistry of Period 3 Elements
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. 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.

Melting and boiling points

Silicon dioxide has a high melting point 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.

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.

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The Chemistry of Period 3 Elements
The Molecular Oxides

Phosphorus, sulphur and chlorine all form oxides which consist of molecules. 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.

Replace the bonds between the phosphorous atoms 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.

15
The phosphorus is using only three of its outer electrons (the 3 unpaired p electrons) to form
bonds with the oxygen atoms.

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The Chemistry of Period 3 Elements
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. 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 SO2 molecules.

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 16
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).

| Notes By TP Shumba
The Chemistry of Period 3 Elements
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, S 3O9, where
three SO3 molecules are joined up and arranged in a ring.

There are also other polymeric forms in which the SO3 molecules 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.

17

| Notes By TP Shumba
The Chemistry of Period 3 Elements
The chlorine oxides

Chlorine forms several oxides and two of them are 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.

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.

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The Chemistry of Period 3 Elements
ACID-BASE BEHAVIOUR OF THE PERIOD 3 OXIDES

The oxides we'll be looking at are:

Na2O MgO Al2O3 SiO2 P4O10 SO3 Cl2O7

P4O6 SO2 Cl2O

The trend in acid-base behaviour

The trend in acid-base behaviour is shown in various reactions, but as a simple summary:

 The trend is from strongly basic oxides on the left-hand side to strongly acidic ones
on the right, via an amphoteric oxide (aluminium oxide) in the middle. An amphoteric
oxide is one which shows both acidic and basic properties.

For this simple trend, you have to be looking only at the highest oxides of the individual
elements. Those are the ones on the top row above, and are where the element is in its highest
possible oxidation state. The pattern isn't so simple if you include the other oxides as well.
For the non-metal oxides, their acidity is usually thought of in terms of the acidic solutions
formed when they react with water - for example, sulphur trioxide reacting to give sulphuric
acid. They will, however, all react with bases such as sodium hydroxide to form salts such as
sodium sulphate.

Chemistry of the individual oxides

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.

Na2O + H2O 2NaOH

Reaction with acids

19
As a strong base, sodium oxide also reacts with acids. For example, it would react with dilute
hydrochloric acid to produce sodium chloride solution.

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The Chemistry of Period 3 Elements
Na2O + 2HCl 2NaCl + H2O

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.

MgO + H2O Mg(OH)2

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.

MgO + 2HCl MgCl2 + H2O

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-
Al2O3 and 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 20
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.

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The Chemistry of Period 3 Elements
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.

Al2O3 + 6HCl 2AlCl3 + 3H2O

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.

Al2O3 + 2NaOH + 3H2O 2NaAl(OH)4

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. 21

SiO2 + 2NaOH Na2SiO3 + H2O

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The Chemistry of Period 3 Elements
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.

SiO2 + CaO CaSiO3

The phosphorus oxides

Phosphorous has two common 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.

P4O6 + 6H2O 4H3PO3

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). 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.

NaOH + H3PO3 NaH2PO3 + H2O

2NaOH + H3PO3 Na2HPO3 + 2H2O

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. 22

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The Chemistry of Period 3 Elements
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.

4NaOH + P4O6 + 2H2O 4NaH2PO3

8NaOH + P4O6 4Na2HPO3 + 2H2O

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.

P4O10 + 6H2O 4H3PO4

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 1moldm-
3
will have a pH of about 1. 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.

NaOH + H3PO4 NaH2PO4 + H2O

2NaOH + H3PO4 Na2HPO4 + 2H2O

3NaOH + H3PO4 Na3PO4 + 3H2O

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. This is
getting ridiculous, and so I will only give one example out of the possible equations:

12NaOH + P4O10 4Na3PO4 + 6H2O 23

| Notes By TP Shumba
The Chemistry of Period 3 Elements
The sulphur oxides

We are going to be looking at sulphur dioxide, SO2, and sulphur trioxide, SO3.

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
which is 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.

SO2 + 2NaOH Na2SO3 + H2O

Na2SO3 + H2O + SO2 2NaHSO3

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.

CaO + SO2 CaSO3

Sulphur trioxide

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

SO3 + H2O H2SO4

Pure un-ionised sulphuric acid has the structure:

24
Sulphuric acid is a strong acid, and solutions will typically have pH's of around 0.

| Notes By TP Shumba
The Chemistry of Period 3 Elements
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.

H2SO4(aq) + H2O(l) H3O+(aq) + HSO4-(aq)

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:

HSO4-(aq) + H2O(l) H3O+(aq) + SO42-(aq)

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.

2NaOH + H2SO4 Na2SO4 + 2H2O

In principle, you can also get sodium hydrogensulphate solution by using half as much
sodium hydroxide and just reacting with one of the two acidic hydrogens in the acid. Sulphur
trioxide itself will also react directly with bases to form sulphates. For example, it will react
with calcium oxide to form calcium sulphate. This is just like the reaction with sulphur
dioxide described above.

CaO + SO3 CaSO4

The chlorine oxides

Chlorine forms several oxides, but the only two chlorine (VII) oxide, Cl2O7, and chlorine (I)
oxide, Cl2O are important. Chlorine (VII) oxide is also known as dichlorine heptoxide, and
chlorine (I) oxide as dichlorine monoxide.

Chlorine (VII) oxide

Chlorine (VII) oxide is the highest oxide of chlorine where the chlorine is in its maximum
oxidation state of +7. It continues the trend of the highest oxides of the Period 3 elements
towards being stronger acids. Chlorine (VII) oxide reacts with water to give the very strong
acid, chloric (VII) acid (also known as perchloric acid) with the pH of typical solutions being
around 0.

Cl2O7 + H2O 2HClO4 25

| Notes By TP Shumba
The Chemistry of Period 3 Elements
Un-ionised chloric (VII) acid has the structure:

When the chlorate (VII) ion (perchlorate ion) forms by loss of a hydrogen ion (when it reacts
with water, for example), the charge can be delocalised over every oxygen atom in the ion.
That makes it very stable, and means that chloric (VII) acid is very strong. Chloric (VII) acid
reacts with sodium hydroxide solution to form a solution of sodium chlorate (VII).

NaOH + HClO4 NaClO4 + H2O

Chlorine (VII) oxide itself also reacts with sodium hydroxide solution to give the same
product.

2NaOH + Cl2O7 2NaClO4 + H2O

Chlorine (I) oxide

Chlorine (I) oxide is far less acidic than chlorine (VII) oxide. It reacts with water to some
extent to give chloric (I) acid, HOCl - also known as hypochlorous acid.

Cl2 + H2O 2HOCl

The structure of chloric (I) acid is exactly as shown by its formula, HOCl. It has no doubly-
bonded oxygens, and no way of delocalising the charge over the negative ion formed by loss
of the hydrogen. That means that the negative ion formed isn't very stable, and readily
reclaims its hydrogen to revert to the acid. Chloric (I) acid is very weak (pKa = 7.43). Chloric
(I) acid reacts with sodium hydroxide solution to give a solution of sodium chlorate(I)
(sodium hypochlorite).

NaOH + HOCl NaOCl + H2O

Chlorine (I) oxide also reacts directly with sodium hydroxide to give the same product.

2NaOH + Cl2O 2NaOCl + H2O

26

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The Chemistry of Period 3 Elements
PROPERTIES OF THE PERIOD 3 "HYDROXIDES"

A quick summary of the trends

Sodium and magnesium hydroxides are simple basic hydroxides while Aluminium hydroxide
is amphoteric. The other "hydroxides" all have -OH groups covalently bound to the atom
from period 3. These compounds are all acidic - ranging from the very weakly acidic silicic
acids to the very strong sulphuric or chloric (VII) acids.

Sodium and magnesium hydroxides

These are both basic because they contain hydroxide ions - a strong base. Both react with
acids to form salts. For example, with dilute hydrochloric acid, you get colourless solutions
of sodium chloride or magnesium chloride.

NaOH + HCl NaCl + H2O

Mg(OH)2 + 2HCl MgCl2 + 2H2O

Aluminium hydroxide

Aluminium hydroxide is amphoteric. It will react with acids thus showing the basic side of its
nature. With dilute hydrochloric acid, a colourless solution of aluminium chloride is formed.

Al(OH)3 + 3HCl AlCl3 + 3H2O

But aluminium hydroxide also has an acidic side to its nature. It will react with sodium 27
hydroxide solution to give a colourless solution of sodium tetrahydroxoaluminate.

Al(OH)3 + NaOH NaAl(OH)4

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The Chemistry of Period 3 Elements
The other "hydroxides"

None of these contains hydroxide ions. In each case the -OH group is covalently bound to the
Period 3 element, and in each case it is possible for the hydrogen atoms on these -OH groups
to be removed by a base. In other words, all of these compounds are acidic.

But they vary considerably in strength:

 Orthosilicic acid is very weak indeed.

 Phosphoric(V) acid is a weak acid - although somewhat stronger than simple organic
acids like ethanoic acid.
 Sulphuric acid and chloric(VII) acids are both very strong acids.

The main factor in determining the strength of the acid is how stable the anion (the negative
ion) is once the hydrogen has been removed. This in turn depends on how much the negative
charge can be spread around the rest of the ion. If the negative charge stays entirely on the
oxygen atom left behind from the -OH group, it will be very attractive to hydrogen ions. The
lost hydrogen ion will be easily recaptured and the acid will be weak. On the other hand, if
the charge can be spread out (delocalised) over the whole of the ion, it will be so "dilute" that
it won't attract the hydrogen back very easily. The acid will then be strong. Wherever
possible, the negative charge is delocalised by interacting with doubly-bonded oxygens. For
example, in chloric(VII) acid, the ion produced is the chlorate(VII) ion (also known as the
perchlorate ion), ClO4-. The structure of the ion doesn't stay like this:

Instead, the negative charge is

delocalised over the whole ion,
and all four chlorine-oxygen
bonds are identical.

When sulphuric acid loses a hydrogen ion to form the hydrogensulphate ion, HSO4-, the
charge can be spread over three oxygens (the original one with the negative charge, and the
two sulphur-oxygen double bonds. That's still an effective delocalisation, and sulphuric acid
is almost as strong as chloric(VII) acid.

Phosphoric(V) acid is much weaker than sulphuric acid because it only has one phosphorus-
oxygen double bond which it can use to help delocalise the charge on the ion formed by
losing one hydrogen ion - so the charge on that ion is delocalised less effectively.

In orthosilicic acid, there aren't any silicon-oxygen double bonds to delocalise the charge.
28
That means the ion formed by loss of a hydrogen ion isn't at all stable, and easily recovers its
hydrogen.

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The Chemistry of Period 3 Elements
Reactions with chlorine

Sodium

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

2Na + Cl2 2NaCl

Magnesium

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

Mg + Cl2 MgCl2

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.

2Al + 3Cl2 2AlCl3

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.

Si + 2Cl2 SiCl4

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.

P4 + 6Cl2 4PCl3

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

29
P4 + 10Cl2 4PCl5

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The Chemistry of Period 3 Elements
Sulphur

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

2S + Cl2 S2Cl2

Chlorine and Argon

It obviously doesn't make sense to talk about chlorine reacting with itself, and argon doesn't
react with chlorine.

PROPERTIES OF THE PERIOD 3 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 of importance is S2Cl2 while 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 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 30
bonds.

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The Chemistry of Period 3 Elements
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.

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
31
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.
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The Chemistry of Period 3 Elements
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 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 = 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.

[Mg(H2O)6]2+ + H2O [Mg(H2O)5(OH)]+ + H3O+

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 atom 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 32
about 180°C. If the pressure is raised to just over 2 atmospheres, it melts instead at a
temperature of 192°C.

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The Chemistry of Period 3 Elements
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.

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 1moldm-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 33
together with chloride ions.

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

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The Chemistry of Period 3 Elements
That extra charge pulls electrons from the water molecules quite strongly towards the
aluminium. That makes the hydrogen atoms 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.

[Al(H2O)6]3+ + H2O [Al(H2O)5(OH)]2+ + H3O+

or, more simply:

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

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 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 important phosphorus chlorides namely phosphorus (III) chloride, PCl3, and
phosphorus (V) chloride, PCl5.

Phosphorus (III) chloride (phosphorus trichloride), PCl3

This is a simple covalent chloride, 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) 34
chloride reacts violently with water to give phosphorous acid, H3PO3, and fumes of hydrogen
chloride (or a solution containing hydrochloric acid if lots of water is used).

| Notes By TP Shumba
The Chemistry of Period 3 Elements
PCl3 + 3H2O H3PO3 + 3HCl

Phosphorus (V) chloride (phosphorus pentachloride), PCl5

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, [PCl 4]+, 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 35

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

| Notes By TP Shumba
The Chemistry of Period 3 Elements
chlorine reacts with hot sulphur. Disulphur dichloride is a simple covalent liquid which is
orange in colour and smelly.

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.

36

| Notes By TP Shumba