3.1.2 - Group 2

Electronic Configuration

As in previous sections, groups indicate the number of electrons occupying the highest electron energy level in an atom of the element. Group 2 elements have 2 electrons.

Using knowledge of subshells and orbitals, one can deduce that these 2 electrons occupy the s-orbital and subshell of the highest energy level. Therefore these atoms always have electron configurations ending in s$\small^2$. When these atoms form ions, they do not only lose one electron, as this would leave an incomplete orbital. Instead they lose both electrons in the s-orbital.

Group 2 Element Atomic Electronic Configuration Stable Ionic Electronic Configuration
Beryllium [He]2s$^2$ [He]
Magnesium [Ne]3s$^2$ [Ne]
Calcium [Ar]4s$^2$ [Ar]
Strontium [Kr]5s$^2$ [Kr]
Barium [Xe]6s$^2$ [Xe]
Radium [Rn]7s$^2$ [Rn]

Redox Reactions of Group 2 Elements

Since group 1 and 2 elements are all alkaline metals, they react very similarly to each other. For instance, both sodium and magnesium will react with water to produce a hydroxide. For this course, there are three reactions of group 2 elements of which knowledge is required.


As with all metals, group 2 elements react with oxygen to form a metal oxide. Since oxygen forms oxidative state -2, and group 2 elements have oxidative state +2 in compounds, only one mole of metal reacts with one mole of oxygen atoms to form a 1:1 ratio:

For example, (magnesium can be switched with any group 2 metal):

$$ \text{Mg} _\text{(s)} + \frac{1}{2} \text{O$_2$} _\text{(g)} \rightarrow \text{MgO} _\text{(s)} $$

In the reaction, magnesium and oxygen both start with oxidation number 0, and then deviate to +2 and -2 respectively. This is therefore a redox reaction.


As mentioned above, group 1 and 2 metals react quite violently with water to produce hydroxides and hydrogen gas, hence why they are called alkaline metals. Group 2 metals have charges of 2+, and yet hydroxide ions have a charge of 1-, so two are needed:

$$ \text{Mg} _\text{(s)} + \text{2H$_2$O} _\text{(l)} \rightarrow \text{Mg(OH)$_2$} _\text{(aq)} + \text{H$_2$} _\text{(g)} $$

The reduction in this reaction is in the hydrogen, which goes from +1 in water to 0 in the hydrogen gas (and some stay at +1 in the hydroxide). The oxidation is of the magnesium, which goes from 0 to +2.


Finally, as with all metals, group 2 elements react with acids to form a salt and hydrogen gas. Similar to before, the charges of the cation and anion must sum to equal 0, and so the number of moles of acid depends on its ‘protic’ number. Only one mole of a diprotic acid is needed, whereas two moles of a monoprotic acid:

$$ \text{Mg} _\text{(s)} + \text{H$_2$SO$_4$} _\text{(aq)} \rightarrow \text{MgSO$_4$} _\text{(aq)} + \text{H$_2$} _\text{(g)} $$

$$ \text{Mg} _\text{(s)} + \text{2HCl} _\text{(l)} \rightarrow \text{MgCl$_2$} _\text{(s)} + \text{H$_2$} _\text{(g)} $$

The reduction in this reaction is in the hydrogen, which goes from +1 in the acid to 0 in the hydrogen gas. The oxidation is of the magnesium, which goes from 0 to +2.

Reactivity of Group 2 Elements

The first and second ionisation energies of group 2 elements decreases as the period increases, as with any group. This is because the highest occupied electron energy level is further away from the nucleus, and so the inward pull from the positively charged nucleus becomes less and less powerful. This means less energy is required to remove the first and second electrons.

In a spontaneous reaction mixture, there is only limited energy, so under the same conditions, more electron transfer would occur with elements lower down in the group, since less energy is required to react. This also means that as the period increases down group 2, the readiness to react also does. In turn, this causes reactivity to increase, since the reaction occurs faster, and more gas is produced per unit time. The more gas produced, the more explosive the reaction is.

Reactions of Group 2 Compounds

Above is the reaction between magnesium, a group 2 metal, and water. It produces magnesium hydroxide and hydrogen gas.

However there is a reaction with an atom economy of 100% in producing magnesium hydroxide. The reaction between magnesium oxide and water does not produce hydrogen gas, as well as only needing one mole of water:

$$ \text{MgO} _\text{(s)} + \text{H$_2$O} _\text{(l)} \rightarrow \text{Mg(OH)$_2$} _\text{(aq)} $$

This reaction can be applied to any group 2 element, as every group 2 oxide only has one oxygen.


As the period increases down the group, the element’s hydroxide becomes more soluble in water. This is electron shielding between the positive nucleus of the group 2 cation increases down the group, and so the ionic bond between the cation and the hydroxide ion becomes less.

As explained in 2.1.4 - Acids and Bases, weak bases are ones that only partially dissociate in water, but as solubility and the subsequent level of dissociation (indicated by the system’s equilibrium constant) increases; the strength, alkalinity and pH also increases.

The relatively low solubilities of beryllium and magnesium hydroxides mean that they only partially dissociate, but since Mg(OH)$\small_2$ is more soluble, its pH is higher.

The hydroxides of calcium and downwards all fully dissociate, and so are strong bases, but pH still increases down the group.


The main use of group 2 compounds are as bases. This is because they are the easiest type of base to produce.

For example, calcium hydroxide is used the neutralise acidic soils. As well as this, magnesium hydroxide and calcium carbonate, CaCO$\small_3$, are used as antacids to treat indigestion.