7 Ions and ionic bonding

This section returns to bonding - the way in which atoms are joined to each other. You have already met one type of bonding involving covalent bonds, which is found in molecules. However, this is not the only bonding found in compounds. In this section you will look at ionic bonding and the ionic compounds that contain such bonding. What is the main difference between the covalent compounds you met in Section 5 and ionic compounds? Are there any similarities between these two enormous families of chemical compounds?

Glucose is a covalent compound and sodium chloride is an ionic compound. You meet glucose in solution in everyday life as it is the sugar in many sweet drinks (and is closely related to ordinary table sugar). You are certainly familiar with sodium chloride as it is the table salt used in kitchens. The formulas of both compounds tell you which atoms have combined together to make them. You can work this out for yourself in the following question.

If you bought some glucose tablets from a pharmacy or a sports shop and crushed them, you would have a fine, white, sweet powder - certainly not like the elements that form it: black carbon, gaseous hydrogen and gaseous oxygen.

Equally, sodium chloride is vastly different from the elements that combine to make it (Figure 23). Sodium chloride is a white solid that is used in food. Yet chlorine is an enormously reactive green gas, deadly poisonous in concentrated form, and an excellent disinfectant when dissolved in water, such as in swimming pools. You probably haven't seen sodium as this silvery metal catches fire in air and almost explodes in water. Once again, the compound - an ionic compound in this case - is completely different from its constituent elements.

Maximise

Figure 23 (a) Sodium metal (immersed in oil to prevent reaction with oxygen and moisture in the air); (b) chlorine gas; (c) the reaction between sodium and chlorine; (d) sodium chloride (common table salt)

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Figure 23

Figure 23 (a) Sodium metal (immersed in oil to prevent reaction with oxygen and moisture in the air); (b) chlorine gas; (c) the reaction ...

If covalent and ionic compounds are similar in that they are unlike the elements that form them, in what key way do they differ? The answer is in the nature of the bonding. To understand ionic bonding, you need to recall what you learned in Section 4 about the internal structure of atoms.

Now apply this to the elements that are in common salt. Table 3 lists the number of electrons (and hence protons) in five elements. Table 6 shows the comparable information about sodium and chlorine.

How do these two elements bond together? Picture a sphere representing a sodium atom. The sphere isn't solid: it has a tiny central nucleus (with a charge of +11) and a cloud of electrons arranged in layers (with a total charge of −11). In many chemical reactions, an atom of sodium very easily loses one electron from the outer layer of this cloud of electrons. The loss of a single electron from a sodium atom makes it into a particle bearing a net positive charge of +1. This positive particle is termed an ion (pronounced 'eye-on'). By taking away one electron from a sodium atom, a sodium ion is formed:

This equation can be easily explained. The sodium atom is electrically neutral; it comprises 11 protons (+11) and 11 electrons (−11). When it loses one electron with a charge of −1, by arithmetic, the particle that is left must have a charge of +1. Note that this is represented by the symbol for sodium with a small superscript 'plus' sign.

Now picture a chlorine atom. This is electrically neutral: it has 17 electrons and 17 protons. In its reactions a chlorine atom likes to gain one electron. Following the same line of reasoning as before, adding one electron to this atom gives one extra negative charge. Thus a chlorine ion (usually called a chloride ion) is formed. Once again, the ion is represented by the symbol for the element followed by the charge written as a superscript. Thus:

The next question covers the same ground but looks at the change from atom to ion in overall terms.

The fact that a sodium atom likes to lose an electron and a chlorine atom likes to gain an electron means that they have a great potential for satisfying each other's 'needs'. This is what lies behind the violence of the reaction between shiny sodium metal and green poisonous chlorine gas (Figure 23(c)). If you dropped a lump of sodium into a container of chlorine gas (behind armoured glass as the reaction is very violent and involves flames) electron transfer occurs. Each sodium atom loses an electron (so forming Na⁺) and each chlorine atom gains one (so forming Cl⁻). On the bottom of the container of chlorine gas, there would be a trace of white powder: the compound sodium chloride, common table salt. The Na⁺ ions and the Cl⁻ ions formed in the reaction attract each other in a similar way to a comb attracting paper. The positive ions are strongly attracted to the negative ions.

The net charge on any ionic compound is always zero. The atoms, after all, start out neutral. All that happens is that some electrons move from one place to another. The total charge on the positive ions exactly equals the total charge on the negative ions.

The answer to (e) raises an important point. In ionic compounds such as salt, it is the attraction between the negatively charged chloride ions and the positively charged sodium ions which holds the substance together. This attraction operates in all directions, unlike the bonding in covalent compounds where the linking is directly between the atoms. Figure 24 shows this in a magnified piece of salt.

Figure 24 A magnified piece of salt. (a) Here the solid lines represent the electrical attraction between every Na⁺ and each of its neighbouring Cl⁻ ions. In fact, this is a two-dimensional representation of a more complex three-dimensional structure. (b) The ions are shown in their space-filling representation. The larger spheres are chloride ions and the smaller ones sodium ions. (c) and (d) illustrate the 'ion packing' structure within sodium chloride - that each sodium ion is surrounded by 6 chloride ions, and that each chloride ion is surrounded by 6 sodium ions, respectively

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Figure 24

Figure 24 A magnified piece of salt. (a) Here the solid lines represent the electrical attraction between every Na⁺ and each of its ...

A distinct molecule of sodium chloride cannot exist because each sodium ion is attracted to more than one chloride ion and vice versa. Thus drawing directional bonds between atoms, as you did for the covalent molecules in Section 5, would be meaningless.

So far, you have considered the Na⁺ ion and the Cl⁻ ion which have lost and gained one electron, respectively. Some atoms like to lose more than one electron or gain more than one electron. For example, atoms of the element calcium (another very reactive metal) always lose two electrons, so the calcium ion is always Ca⁺⁺ normally written as Ca²⁺ (pronounced 'see-ay-two-plus'). In this case, the superscript '2+' tells you that there are two plus charges on this ion.

Another example of a 'two electron change' is shown by atoms of the element oxygen. Recall that an oxygen atom forms two covalent bonds in some reactions (you saw this in Section 6 in H₂O and CO₂). However, in some other reactions it prefers to gain two electrons and thus forms ions. When this happens, the ion formed is O²⁻ (pronounced 'oh-two-minus'). This is called the oxide ion.

This preference for forming oxide ions is shown by oxygen in the many reactions where metals burn in oxygen. For example, calcium metal burns in oxygen to form the ionic compound calcium oxide. The formula for calcium oxide is CaO. This means that there is exactly one calcium ion (Ca²⁺) to every oxide ion (O²⁻) in a piece of calcium oxide. Thus the overall neutrality of the compound is maintained: (2+) + (2−) = 0.

Note that the formula of an ionic compound gives the ratio between the two kinds of ions. By convention, charges are not written in the formula. Thus it is NaCl and CaO and not Na⁺Cl⁻ and Ca²⁺O²⁻.

What happens in an ionic compound when an ion containing two charges is combined with ions bearing only one charge? The 'golden rule' is that electrical neutrality must be maintained.

So far, this section has described simple compounds containing only two types of element: one forms the positive ion and one forms the negative ion. However, there are several ionic compounds that are more complex than this because one or both of the ions contain more than one element. An example is the nitrate ion, the principal culprit in water pollution by agricultural fertilisers.

The nitrate ion contains a cluster of atoms: one nitrogen and three oxygen atoms covalently bonded together inside the nitrate ion. However, the cluster of atoms as a whole bears just one negative charge just like the simple Cl⁻ ion. So the chemical formula for the nitrate ions is NO₃⁻. Another ion that contains a cluster of atoms is the sulfate ion. Table 7 summarises some common ions. You may find it useful to memorise the chemical formulas for these ions.

The negative nitrate ion pairs with a positive ion, of course. Sometimes this is a sodium ion (Na⁺); alternatively, it could be potassium (K⁺). Sodium nitrate (NaNO₃) and potassium nitrate (KNO₃) are, in fact, two of the commonest fertilisers containing the nitrate ion (Figure 1b). Sodium nitrate contains a vast number of sodium ions (Na⁺) and nitrate ions (NO₃⁻). As you can tell from the formula for sodium nitrate, these ions are present in equal numbers to give overall electrical neutrality.

The attraction of the positive and negative ions in sodium nitrate, in the solid form, holds the compound together in a similar way to that shown for NaCl in Figure 23. When sodium nitrate is added to water, the water molecules interpose themselves between the two types of ion reducing the attraction between them and allowing them to separate. The result is that the solid sodium nitrate dissolves in water (Figure 1(c)).

A final point about the names of ionic compounds may be helpful. In some ionic compounds (e.g. sodium chloride and calcium oxide), the name of the negative ion ends in '-ide'. As you know from Section 5 on covalent compounds, this simply denotes 'two elements in the compound'. However, in ionic compounds where the positive ion is balanced by a negative ion containing oxygen and another element (such as in the ion NO₃⁻), the ending usually changes to '-ate'. Thus the NO₃⁻ ion is called the nitrate ion; this tells you that the ion contains nitrogen and oxygen. Some general rules about the chemical names of different substances are summarised below.

Try the following four questions to practise your understanding of chemical formulas.

Make sure you have checked your answers to Questions 21 and 22 before trying Questions 23 and 24.

In the next section you will see why ions are an important feature of the water around you.