Ionic formulae

We need to use the correct formulae throughout the practical and theory work of chemistry. They are the language of chemistry and it is important to get them right.

Ionic compounds are generally formed between metals and non-metals. Most elements are metals (shown in yellow below). The non-metals can be found on the right hand side of the periodic table (in purple below). Those marked in green are metalloids with properties in between metals and non-metals.

The metal atom loses its outer electron(s) to become a positive ion or cation. The non-metal gains electrons to become a negative ion or anion. The resulting attraction between the positive and negative ions holds the compound together (opposite charges attract). A typical example is salt which is sodium chloride:

Salt has high melting and boiling points as the ions attract each other strongly and are hard to separate. As the ions are charged, if they become free to move it will conduct electricity - this happens if the salt is melted or dissolved in water. Salt dissolves in water because the polar water molecules are attracted to the charged ions. This behaviour is typical of an ionic compound. They have high melting points, conduct when molten or in solution, and tend to dissolve in polar solvents, like water, rather than non-polar solvents like petrol. Evidence for this ionic model is seen in this demonstration, where a voltage applied across a sample of potassium manganate(VII) causes the purple MnO₄^- ion to move towards the positive electrode (opposites attract).

It is very important to be able to work out the formula of an ionic compound. In order to do this you will need to learn the name and formula of the following ions (try making some flash cards):

Lithium	Li⁺	Chloride	Cl^{^-}
Sodium	Na⁺	Bromide	Br^-
Potassium	K⁺	Iodide	I^{^-}
Hydrogen	H⁺	Hydrogencarbonate
Ammonium		Hydroxide	OH^-
Silver	Ag⁺	Nitrate
Magnesium	Mg²⁺	Manganate(VII)
Calcium	Ca²⁺	Oxide	O^2-
Barium	Ba²⁺	Sulfide	S^2-
Zinc	Zn²⁺	Sulfate
Aluminium	Al³⁺	Carbonate
Fluoride	F^-	Dichromate(VI)

There are some patterns which make learning easier. All group 1 atoms form 1+ ions, all group 2 form 2+ ions, all group 3 form 3+ ions, all group 7 form 1- ions, all group 6 form 2- ions and all group 5 form 3- ions. Some atoms are able to form more than one ion, the charge of the particular ion is given in Roman numerals in brackets after the name. So iron(III) contains Fe³⁺ and iron(II) contains Fe²⁺. Some compounds are so common their formulae are worth memorizing - eg water, H₂O; carbon dioxide, CO₂; hydrochloric acid, HCl; sulfuric acid, H₂SO₄; and nitric acid, HNO₃.

In order to work out the formula of a compound, a positive ion is placed with a negative ion. The number of ions is then adjusted until the positive charge exactly balances the negative charge. For example:

Sodium chloride Na⁺ and Cl^- formula NaCl

one 1+ ion balances one 1- ion.

Calcium oxide Ca²⁺ and O^2- formula CaO

one 2+ ion balances one 2- ion.

Magnesium bromide Mg²⁺ and Br^- formula MgBr₂

one 2+ ion balances two 1- ions.

Aluminium oxide Al³⁺ and O^2- formula Al₂O₃

two 3+ ions balance three 2- ions.

Iron(III) chloride Fe³⁺ and Cl^- formula FeCl₃

one 3+ ion balance three 1- ions.

Copper(II) hydroxide Cu²⁺ and OH^- formula Cu(OH)₂

one 2+ ion balances two 1- ions. (Note the use of brackets - this shows that there are two OH^- ions. Without the brackets it would mean one O and two Hs)

Ammonium sulfate NH₄⁺ and SO₄^2- formula (NH₄)₂SO₄

two 1+ ions balances one 2- ion.

(Note NH₄⁺ means NH₄ with a 1+ charge not NH with a 4+ charge.)

Try the computer program FORMULA (if you press the # key at any time you will get a wider range of ions).