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Physical and Chemical Properties of Some Period 3 (Na -> Cl) Compounds


The properties of the hydrides are summarized in Table 18.5. The hydrides of the less electronegative elements, such as lithium hydride and sodium hydride, are ionic. The remaining hydrides become progressively more covalent in nature as we move across the period. In water, most hydrides react to form either alkaline hydroxide solutions or acidic solutions. The remainder do not react with water but produce a neutral solution.

Table 18.5. Properties of Some Period 3 Hydrides

Reactions of Hydrides with Air

Hydrides do not react with air at room temperature. However, a few react with oxygen in the air to form their oxides and water when heated to high temperatures. Examples are shown in the equations which follow

B2H6 is unstable and very reactive. It burns with a green flame.

Carbon forms many other stable hydrides hydrocarbons such as alkanes, alkenes, alkynes and aromatic hydrocarbons which burn in air. Silicon has only alkane analogues, stable up to Si6H14 These are not as stable as the hydrocarbons and hence more reactive in air.

Reactions of Hydrides with Water

All the ionic and many of the covalent hydrides react with water to different extents. The ionic hydrides give hydroxides and hydrogen. Reactions of ionic hydrides are shown by their equations as:

B2H6 is unstable and very reactive. It burns with a green flame.

Carbon forms many other stable hydrides hydrocarbons such as alkanes, alkenes, alkynes and aromatic hydrocarbons which burn in air. Silicon has only alkane analogues, stable up to Si6H14 These are not as stable as the hydrocarbons and hence more reactive in air.

Reactions of Hydrides with Water

All the ionic and many of the covalent hydrides react with water to different extents. The ionic hydrides give hydroxides and hydrogen. Reactions of ionic hydrides are shown by their equations as:

Covalent hydrides such as aluminium hydride, boron hydride, ammonia and hydrogen fluoride react with water to give basic and acidic solutions. The reactions are as follows:

Methane does not react with water at room temperature. Silicon hydride does not react with water at room temperature except in the presence of an alkali.


The oxides of third period elements are:

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 i.e., from 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. Sodium, magnesium and aluminium oxides consist of giant structures containing metal ions and oxide ions. Magnesium oxide has a structure just like sodium chloride. The other two have more complicated arrangements of the ions beyond the scope of syllabuses at this level.

Melting and Boiling Points

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 i.e., some small and simple; others polymeric. ·

The attractive forces between these molecules are van der Waals dispersion and dipole-dipole interactions. These vary depending on the size, shape and polarity of the various molecules but will always be much weaker than the ionic or covalent bonds in a giant structure. These oxides tend to be gases, liquids or low melting 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 in the molten state.

Electrolysis of molten sodium oxide depends on whether it actually melts or sublimes or decomposes 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 lab.

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 SiO2. 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 (see Fig. 18.2).

Fig. 18.2. Crystal structure of silica.

Notice that each silicon atom is bridged to its neighbours by an oxygen atom. Do not 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 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.

In simple words silicon dioxide has giant structure, and hence the melting and boiling points are high.

Electrical Conductivity

Silicon dioxide does not have any mobile electrons or ions-so it does not 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 while others are polymeric. 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 neighbors’ 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 phosphorus uses only three of its outer electrons (the 3 unpaired p electrons) to form bonds with the oxygen atoms.


Phosphorus(V) oxide

Phosphorus(V) oxide ix 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 are polymeric.

The structure of P4O10 in a simple molecular form is given below:

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 consits of simple SO2 molecules.

The sulphur atom 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 laboratory, 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 SO3 molecules join together in long chains as shown below:

The Chlorine Oxides

Chlorine forms several oxides. Here we are just looking at two of them (i) chlorine(I) oxide, Cl2O, and (ii) 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 surprising about this molecule and its physical properties are just what one would expect for a molecule of 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 Its melting point and boiling point is higher Than chlorine (I) oxide.

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

The structures of oxides of chlorine are given below. The shape is tetrahedral around both chlorines, and V –shaped around the central oxygen

Acid-Base Behaviour of the Period 3 Oxides

The trend in acid-base behaviour is summarized below:

• On moving from left to right among the oxides we observe that metal oxides, form 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.

The oxides in highest oxidation state of element are known to be acidic in character. The acidic character decreases with decrease in oxidation state of atom.



 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  à  2NaOH

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.

Na2O + 2HCl  à  2NaCl + H2O


Magnesium Oxide

Magnesium oxide is a simple basic oxide, because it also contains oxide ions. However, it is not as strongly basic as sodium oxide because the oxide ions are not so free. In the case sodium oxide, 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 if other factors like the energy released when the positive ions develop attractions with water in the solution are taken into account, 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 does not 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.

Mg O(s) + H2O(l)   à  Mg(OH) (l)

Reaction with acids

Magnesium oxide reacts with warm dilute hydrochloric acid to give magnesium chloride solution.

Mg O(s) + 2HCl(aq)  à MgCl2 (aq)  + H2O (l)

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-A12O3 and is produced at high temperatures. Aluminium oxide is amphoteric. It has reactions as both a base and an acid.


Reaction with water

Aluminium oxide does not react in a simple way with water in the sense that sodium oxide and magnesium oxide do, and does not dissolve in it. Although it 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. Aluminium oxide reacts with hot dilute hydrochloric acid to give aluminium chloride.

Al2O3 + 6HCl  à 2AlCl3 + 3H2O


The above reaction supports the basic character of aluminium.

Reaction with bases

 Aluminium oxide is also an acidic in its nature. It shows this by reacting with bases such as sodium hydroxide solution and forms aluminate.

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

Al2O3 + 2NaOH + 3H2O  à 2NaAl (OH)4


Silicon dioxide (Silicon(IV) Oxide)

The bonds between Si and oxygen are covalent in character due to less electronegativity difference between the two.

Silicon dioxide has no basic properties i.e., it does not react with acids. Instead, it is very weakly acidic, and reacts with strong bases.

Reaction with water

 Silicon dioxide does not 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 bot and concentrated. A colourless solution of sodium silicate is formed

SiO2 + 2NaOH  à  Na2siO3 + H2O

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.

Phosphorus acid is a weak acid. Phosphorus acid has a pK a of 2.00 which makes it stronger than common organic acids like ethanoic acid (pK3 = 4.76). It is pretty unlikely that you would ever react phosphorus(ill) oxide directly with a base, but you might need to know what happens if you react the phosphorus acid with a base.


In phosphorus acid, the two hydrogen atoms in the -OH groups are acidic, but the other one is not. That means that you can get two possible reactions with, sodium hydroxide solution depending on the proportions used.


NaOH + H3PO3  à  NaH2PO3 + H2O

2NaOH + H3PO3 à 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.


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 consider one of these,

tetraoxopbosphate (V) acid H3PO4. It is also known as orthophosphoric acid.


P4O10 + 6 H2O à 4H3PO4


The pure un-ionised acid has the structure


SiO2 + CaO à CaSiO3


The Phosphorus Oxides


Phosphorus forms two 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 as phosphorus acid, orthophosphorus acid or phosphoric acid. Its reaction with hot water is much more complicated.


P4 O6 + 6H2O à 4H3PO3


The pure un-ionised acid has the structure