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Group 2 of the periodic table contains the alkaline earth metals: beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), and barium (Ba). At A-Level, the focus is on magnesium to barium, since beryllium behaves anomalously due to its very small atomic radius and high charge density.
These elements share a common electron configuration ending in s², meaning each atom has two electrons in its outermost shell. This accounts for the +2 oxidation state that dominates their chemistry.
As you descend Group 2, atomic radius increases. Each successive element has an additional electron shell, so the outermost electrons are further from the nucleus. Magnesium has three shells, calcium has four, strontium has five, and barium has six.
| Element | Atomic Radius (pm) | Electron Configuration |
|---|---|---|
| Mg | 160 | 1s² 2s² 2p⁶ 3s² |
| Ca | 197 | [Ar] 4s² |
| Sr | 215 | [Kr] 5s² |
| Ba | 222 | [Xe] 6s² |
The increase in radius has important consequences for reactivity, ionisation energy, and the properties of compounds.
First ionisation energy decreases down Group 2. Although nuclear charge increases (more protons), the effect is outweighed by increased shielding from inner electron shells and greater distance of the outer electrons from the nucleus. The outermost electrons are easier to remove in barium than in magnesium.
The second ionisation energy also decreases down the group for the same reasons. Since Group 2 elements form M²⁺ ions by losing both s² electrons, the decreasing ionisation energies explain the increasing reactivity down the group.
Melting points do not follow a simple trend in Group 2. Beryllium has the highest melting point (1287 °C), then magnesium (650 °C), but calcium (842 °C) is actually higher than magnesium — an irregularity. Strontium (777 °C) and barium (727 °C) continue the general decrease. The melting points depend on metallic bond strength, which is influenced by ionic radius and crystal structure. The overall tendency is for melting point to decrease down the group as the metallic radius increases and the bond strength per unit volume decreases, but changes in crystal packing cause irregularities.
The Group 2 metals react with water to form a metal hydroxide and hydrogen gas:
M(s) + 2H₂O(l) → M(OH)₂(aq) + H₂(g)
Reactivity increases down the group because ionisation energies decrease, making it easier for atoms to lose their two outer electrons.
This increasing reactivity is a direct consequence of decreasing ionisation energy down the group.
When Group 2 metal compounds are heated in a flame, electrons are promoted to higher energy levels and then fall back, emitting light of characteristic wavelengths:
| Element | Flame Colour |
|---|---|
| Mg | No distinctive colour (white/no colour) |
| Ca | Brick red / orange-red |
| Sr | Crimson / deep red |
| Ba | Pale green |
These flame colours are used as a diagnostic test to identify Group 2 metal ions.
Two important solubility trends must be learned for Group 2 compounds:
This trend means that as you go down Group 2, the hydroxides dissolve more readily and produce more alkaline solutions. Ba(OH)₂ solution is strongly alkaline.
The insolubility of BaSO₄ is exploited in the test for sulfate ions: adding BaCl₂ solution to a sample acidified with HCl produces a white precipitate of BaSO₄ if sulfate ions are present.
Group 2 shows clear periodic trends driven by increasing atomic radius and decreasing ionisation energy. Reactivity with water increases down the group, hydroxide solubility increases, and sulfate solubility decreases. Flame tests provide a quick way to identify individual Group 2 elements.