ionic bonding

Foundation and higher tier

Ion formation

Metal atoms are found in the left hand-side of the periodic table. In GCSE chemistry we will be mainly concerned with the metals in groups I and II, the alkali metals and the alkaline earth metals. The alkali metals, Li, Na, K, Rb and Cs being in group 1 of the periodic table all have one electron in their outer shell. Remembering the rule: elements only react if they can achieve full last electron shells, this means that these metals will either have to gain 7 electrons to fill their last shell or simply lose one electron when they react.

As an example consider sodium, atomic number 11. Sodium has the electronic configuration 2,8,1. To completely fill its last shell it needs to gain 7 electrons or simply lose the last electron. Obviously it will be easier and require less energy to simply lose 1 electron than gain 7 electrons. So when sodium reacts it will lose its outer shell electron. The sodium atom will now have the electronic arrangement 2,8 the same as the noble gas neon. It will have 10 electrons. However the nucleus of the sodium atom will still contain 11 protons, each one having a positive charge. This means that the sodium atom has 11 positive charges (protons) but only 10 negatively charged electrons, so overall the atom has more positive charges than negative charges to cancel them out. This leaves the sodium atom with a charge of +1, we call atoms with a charges ions. Atoms are neutral because they have equal numbers of positive protons in the nucleus and negatively charged electrons in the electron shells, but in ions the numbers of protons and electrons is NOT equal, so the atom has a charge- it's now an ion. This is shown in the diagram below:


sodium atom forming a sodium ion

Ion formation and ion charges

Metal atoms always lose electrons when they react with non-metals, this means they end up with more positively charged protons than negatively charged electrons - so metal ions are always positively charged. Again remembering that group I loses one electron and forms ions with a 1+ charge, group II metals lose two electrons and so will end up forming ions with a 2+ charge. Similarly group III metals e.g. aluminium will lose 3 electrons and end up with a 3+ charge on their ions. This is summarised in the table below:

Group in periodic table number of electrons in the last shell number of electrons lost charge on ion examples
1 1 1 1+ Li+, Na+, K+
2 2 2 2+ Mg2+, Ca+
3 3 3 3+ Al3+

A similar thing happens with the reactive non-metals, they are mainly in groups 5, 6 and 7 of the periodic table. Non-metals in group 5, obviously have 5 electrons in their last shell, so they need to gain 3 electrons to form a full last shell. However if they gain 3 electrons, or 3 negative charges, then they will have 3 more negatively charged electrons than positively charged protons in their nucleus, so they will end up with forming an ion with a 3- charge. A similar thing happens with the non-metals in groups 6 and 7, e.g.

Group in periodic table number of electrons in the last shell number of electrons gained charge on ion examples
5 5 3 3- N3-, P3-
6 6 2 2- O2-, S2-
7 7 1 1- F-, Cl-, Br-

Forming ionic compounds

Since metal atoms lose electrons when they react to achieve full last shells and non-metal atoms gain electrons when they react then it would seem obvious that metals and non-metals react easily together to form compounds made of ions, that is ionic compounds. Ionic compounds consist of metal ions and non-metal ions. Consider the reaction between sodium metal, a group I alkali metal, electron arrangement 2,8,1 and chlorine, atomic number 17, a halogen in group 7, electron arrangement 2,8,7. The sodium atom needs to lose one electron to gain a full last shell and the chlorine needs to gain one electron to fill its last shell. So when sodium and chlorine react an electron is transferred from the sodium atom to the chlorine atom. We can show what is happening in terms of electron transfer in a diagram, as shown below:

dot and cross diagram for sodium chloride formation

Normally when we draw dot and cross diagram only the electrons in the last shells are drawn. The reason for this is simple, these are the electrons that are involved in the bonding. So the dot and cross diagram above can be simplified, this time showing only the outer electrons:

dot and cross diagram for sodium chloride

The positive sodium ion and the negative chloride ion produced are strongly attracted to each other, since they have opposite charges. This type of attraction or bond is called an ionic bond. It is a type of electrostatic attraction.

Example 2- Magnesium Chloride

Magnesium is an alkaline earth metal in group 2 of the periodic table, its electron arrangement is 2,8,2, it needs to lose 2 electrons to achieve a full last shell. Chlorine as above is a halogen in group 7, its electron arrangement is 2,8,7. It only needs to gain one electron. Looks like we have a problem! Magnesium needs to lose 2 electrons but chlorine only needs to gain 1 electron. The solution is simple, each magnesium reacts with 2 chlorine atoms, each chlorine atom accepts one electron from the magnesium. This is shown in the diagram below:

dot and cross diagram for magnesium chloride

As before we can simplify the dot and cross diagram so that only the electrons that actually take part in the reaction/bonding, that is the outer shell electrons are shown:

dot and cross diagram for magnesium chloride

You may notice that when the atoms react they always end up with the same electron arrangement as a Noble gas- that is full last shells.

So what exactly is an ionic bond?

Most of the discussion above has been about ions and how they are formed, however once you are confident about how ions are formed then it should be obvious that once formed postive and negatively charged particles will attract each other by electrostatic forces. Therefore an ionic bond is simply the electrostatic attraction between oppositely charged ions. Ionic compounds consist of giant structure of ions called a lattice. Click on the link below or here for more information on the structure of ionic compounds.

Key Points

Practice questions

Check your understanding - Questions on ionic bonding

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