Group 2

Atomic Properties

Down Group 2:

• Atomic radius increases (more electron shells)
• Ionisation energy decreases (outer electrons further from nucleus, more shielding)
• Reactivity increases (easier to lose two outer electrons)
• Electronegativity decreases

Why Reactivity Increases

Group 2 metals react by losing 2 electrons to form M²⁺: $\text{M} \rightarrow \text{M}^{2+} + 2e^-$

Down the group, the outer electrons are further from the nucleus with more shielding, so they require less energy to remove (lower IE₁ and IE₂).

2.1 With Water

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

Reactivity increases down the group.

2.2 With Oxygen

$2\text{M}(s) + \text{O}_2(g) \rightarrow 2\text{MO}(s)$

All burn with characteristic colours: Mg = brilliant white, Ca = brick red, Sr = crimson, Ba = pale green.

2.3 With Dilute Acids

$\text{M}(s) + 2\text{HCl}(aq) \rightarrow \text{MCl}_2(aq) + \text{H}_2(g)$

All react; rate increases down the group.

Oxides

Group 2 oxides are basic — they react with water to form hydroxides: $\text{MO}(s) + \text{H}_2\text{O}(l) \rightarrow \text{M(OH)}_2(aq/s)$

They also neutralise acids: $\text{MgO}(s) + 2\text{HCl}(aq) \rightarrow \text{MgCl}_2(aq) + \text{H}_2\text{O}(l)$

Hydroxides — Solubility Trend

Solubility increases down Group 2:

As solubility increases, the solutions become more alkaline (higher concentration of OH⁻).

Ca(OH)₂ (limewater) is used to test for CO₂. Ba(OH)₂ is a strong base used in titrations.

Solubility decreases down Group 2 (opposite to hydroxides):

BaSO₄ is used in the test for sulfate ions (add Ba²⁺ to acidified solution → white precipitate if sulfate present).

BaSO₄ is also used as a "barium meal" in medical X-rays — it's opaque to X-rays but safe to swallow because it's so insoluble that it passes through the body without being absorbed (despite barium being toxic).

Carbonates

$\text{MCO}_3(s) \xrightarrow{\text{heat}} \text{MO}(s) + \text{CO}_2(g)$

Thermal stability increases down the group:

• MgCO₃ decomposes easily at ~350°C
• BaCO₃ requires >1000°C

Why?

Larger cations (Ba²⁺) are less polarising — they distort the electron cloud of the carbonate ion less. The carbonate is more stable and requires more energy to decompose. Smaller cations (Mg²⁺) have higher charge density, polarise the CO₃²⁻ more, weakening the C-O bonds and making decomposition easier.

Nitrates

Group 2 nitrates decompose to form the oxide, NO₂, and O₂:

$2\text{M(NO}_3\text{)}_2(s) \rightarrow 2\text{MO}(s) + 4\text{NO}_2(g) + \text{O}_2(g)$

Same trend: thermal stability increases down the group.

(Mg and Be don't produce visible flame colours in a standard flame test.)

• Two opposite trends to learn: hydroxides (more soluble down) vs sulfates (less soluble down)
• Thermal stability explanation: link to charge density and polarisation
• Know the flame colours: Ca = orange-red, Sr = crimson, Ba = pale green
• Write both full and ionic equations for precipitation reactions

• Reactivity increases down Group 2 (easier to lose 2 electrons)
• Hydroxide solubility increases down the group → more alkaline solutions
• Sulfate solubility decreases down the group → BaSO₄ test for sulfates
• Thermal stability of carbonates/nitrates increases down the group (larger cation → less polarising)
• Flame colours: Ca = orange-red, Sr = crimson, Ba = pale green