Group 2

Four metals, one pattern

The syllabus uses four Group 2 metals: magnesium, calcium, strontium and barium. They are reactive, silvery metals. Each one reacts with a given reagent in the same kind of reaction, giving the same kind of product. What changes down the group is mainly the rate. So you learn one general equation for each reaction (M stands for any of the four metals), then learn how the rate and the observations change from Mg to Ba.

Group 2 compounds are ionic and contain the $\text{M}^{2+}$ ion, so the formulae are always $\text{MO}$, $\text{M(OH)}_2$, $\text{MCl}_2$, $\text{MSO}_4$, $\text{MCO}_3$ and $\text{M(NO}_3)_2$.

With oxygen: the oxide, MO

Every Group 2 metal burns in oxygen to form its oxide, a white solid:

What you see:

• Magnesium burns with a bright white flame (light) and leaves a white solid (some of it rises as white smoke).
• Calcium, strontium and barium burn with coloured flames (calcium brick-red, strontium red, barium green, or white with a green tinge) and also leave a white solid.

Exam questions often say "do not refer to temperature changes", so give the flame or the white solid, not "it gets hot".

With cold water: the hydroxide and hydrogen

This one equation covers all four metals with cold water. (For magnesium, write $\text{Mg(OH)}_2\text{(s)}$, because it hardly dissolves.) The rate is very different at the two ends of the group:

• Magnesium reacts very slowly with cold water: a few bubbles of gas form slowly on the metal. A thin layer of oxide on the surface, and the barely soluble $\text{Mg(OH)}_2$ that forms, both slow it down.
• Calcium reacts steadily: effervescence (fizzing), the calcium gets smaller, and the liquid goes cloudy white because some $\text{Ca(OH)}_2$ does not dissolve.
• Strontium and barium react faster still: rapid effervescence, the metal disappears quickly, and a colourless solution is left, because $\text{Sr(OH)}_2$ and $\text{Ba(OH)}_2$ are more soluble.

In every case the solution formed is alkaline (it contains $\text{OH}^-$ ions).

To compare two metals, describe the difference in what you see: the metal lower in the group gives more bubbles per second, disappears more quickly, and stops fizzing first.

Magnesium with steam: the oxide, not the hydroxide

Heated magnesium reacts quickly with steam, glowing with a bright white light. With steam the product is the oxide, not the hydroxide:

So magnesium gives a different product with steam, not just a faster version of the cold-water reaction.

Why reactivity increases down the group

You met this reasoning in the AS Atomic Structure note. Each element down Group 2 has one more shell of electrons. The two outer electrons are further from the nucleus and more shielded by the inner shells, so the nucleus attracts them less, even though it has more protons. So the first and second ionisation energies decrease down the group.

In every reaction in this section the metal atom loses its two outer electrons to become $\text{M}^{2+}$. The lower the ionisation energy, the more easily this happens, so the reaction is faster.

The metals are reducing agents

Look at the oxidation numbers in the reaction with water:

• The metal goes from $0$ to $+2$: it loses electrons, so it is oxidised.
• Hydrogen in water goes from $+1$ to $0$: water gains electrons, so water is reduced.

The metal gives electrons to the other substance, so the metal is the reducing agent. The same is true with oxygen and with dilute acids. Because the metals lose electrons more easily down the group, their strength as reducing agents increases down the group: barium is the strongest reducing agent of the four, magnesium the weakest.

With dilute acids: a salt and hydrogen

Every Group 2 metal, including magnesium, reacts quickly with dilute acid: effervescence and the metal dissolves (disappears), leaving a colourless solution. The rate still increases down the group.

With hydrochloric acid every chloride is soluble, so the metal reacts completely.

With sulfuric acid the sulfate matters. Magnesium sulfate is soluble, so write $\text{MgSO}_4\text{(aq)}$ and the magnesium dissolves completely. Calcium sulfate is only slightly soluble, and strontium and barium sulfates are insoluble (you meet this trend in “Solubility of the hydroxides and sulfates”). For these metals the sulfate forms as a solid, $\text{MSO}_4\text{(s)}$, coats the metal, and the reaction soon stops, leaving most of the metal unreacted.