What you'll learn:

• How atomic radius, first ionisation energy, and melting point change from magnesium to barium.
• The trends in their reactivity with water and the solubility of their hydroxides and sulfates.
• Important industrial and medical applications of Group 2 metals and compounds.
• The chemical test for sulfate ions using acidified barium chloride.

Trends in physical properties (Mg to Ba)

The elements from magnesium (Mg) to barium (Ba) are known as the alkaline earth metals. Like all chemical groups, their physical properties follow distinct patterns as you move down the periodic table.

Atomic radius

Atomic radius increases as you go down Group 2. As you move from magnesium down to barium, each element has an extra electron shell compared to the one above it. This extra shell means the outer electrons are physically further from the nucleus. Although the nuclear charge (number of protons) also increases, the extra inner shells provide shielding (or screening), which cancels out the effect of the extra protons. Therefore, the overall size of the atom increases.

First ionisation energy

The first ionisation energy decreases as you go down Group 2. Although the nuclear charge increases, the atomic radius increases and there is more shielding from inner electron shells. The outermost electrons are further away from the positive nucleus and experience a weaker electrostatic attraction, making them easier (requiring less energy) to remove.

Melting point

The melting point generally decreases as you go down Group 2. Group 2 elements have a giant metallic structure. The metal atoms lose their two outer electrons to form a "sea" of delocalised electrons, leaving behind a lattice of $+2$ ions. As you go down the group, the charge on the ion remains $+2$ and the number of delocalised electrons per atom remains 2. However, the ionic radius increases. Because the ions get larger, the distance between the positive metal nucleus and the delocalised electrons increases, which weakens the electrostatic attraction. Less energy is needed to break these metallic bonds.

Chemical reactions of Group 2 elements

Reaction with water

The reactivity of Group 2 elements with water increases down the group. They undergo a redox reaction to form a metal hydroxide and hydrogen gas. For a general Group 2 metal ($M$), the equation is:

$$M_{(s)}+2H_2{O}_{(l)}\to {M(OH)}_{2(aq)}+H_{2(g)}$$

• Magnesium reacts very slowly with cold liquid water, taking days to produce a few bubbles of hydrogen and a weakly alkaline solution of sparingly soluble magnesium hydroxide.
• Calcium, Strontium, and Barium react increasingly vigorously with cold water, producing noticeable effervescence (fizzing) as hydrogen gas is released.

Because magnesium reacts so poorly with cold water, we typically react it with steam instead. This is a much faster reaction and produces magnesium oxide (a white solid) instead of magnesium hydroxide:

$${Mg}_{(s)}+H_2{O}_{(g)}\to {MgO}_{(s)}+H_{2(g)}$$

Extraction of titanium

Titanium is a highly useful metal (light, strong, anti-corrosive), but it cannot be extracted using carbon because it reacts to form titanium carbide, making the metal brittle. Instead, magnesium is used as a reducing agent to extract titanium from titanium(IV) chloride (${TiCl}_4$).

The process happens at very high temperatures in an inert argon atmosphere (to prevent magnesium reacting with oxygen):

$${TiCl}_4(g)+2{Mg}_{(l)}\to {Ti}_{(s)}+2{MgCl}_{2(l)}$$

Trends in solubility

You must know the solubility trends for two types of Group 2 compounds: hydroxides (${OH}^{-}$) and sulfates (${SO}_4^{2-}$).

Solubilities of the Hydroxides

Solubility increases down the group.

• Magnesium hydroxide (${Mg(OH)}_2$) is classed as sparingly soluble (almost insoluble).
• Barium hydroxide (${Ba(OH)}_2$) is highly soluble, forming a strongly alkaline solution.

Solubilities of the Sulfates

Solubility decreases down the group.

• Magnesium sulfate (${MgSO}_4$) is highly soluble.
• Barium sulfate (${BaSO}_4$) is completely insoluble.

Uses of Group 2 compounds

The unique solubilities of Group 2 compounds dictate how we use them in industry and medicine.

1. Magnesium hydroxide in medicine: Because it is sparingly soluble, it barely dissolves to release hydroxide ions. This makes it safe to use as an antacid (like "Milk of Magnesia") to safely neutralise excess stomach acid without damaging tissue.
2. Calcium hydroxide in agriculture: Known as "slaked lime", ${Ca(OH)}_2$ is spread on fields to neutralise acidic soils, raising the pH to an optimal level for crop growth.
3. Calcium oxide or calcium carbonate in industry: Used in "flue gas desulfurisation". When fossil fuels burn, they release toxic sulfur dioxide (${SO}_2$). A slurry of $CaO$ or ${CaCO}_3$ is sprayed into the exhaust flues to neutralise the acidic ${SO}_2$ gas, producing harmless calcium sulfite (${CaSO}_3$).
4. Barium sulfate in medicine: Used as a "Barium meal". Patients swallow a suspension of ${BaSO}_4$ before an X-ray. Barium absorbs X-rays well, outlining the gut. Barium compounds are highly toxic, but because ${BaSO}_4$ is entirely insoluble, it cannot be absorbed into the blood and safely passes through the digestive system.

Testing for sulfate ions

We can use the insolubility of barium sulfate as a definitive chemical test for sulfate ions (${SO}_4^{2-}$) in an unknown solution.

To carry out the test:

1. Add a small volume of dilute hydrochloric acid (or dilute nitric acid) to the unknown solution.
2. Add a few drops of barium chloride solution (${BaCl}_2$).
3. If sulfate ions are present, a thick white precipitate of solid barium sulfate instantly forms.

The ionic equation for the positive result is:

$${Ba}_{(aq)}^{2+}+{SO}_{4(aq)}^{2-}\to {BaSO}_{4(s)}$$

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