Calcium (Ca)

Calcium has the ground-state configuration \([Ar]4s^2\). Losing the two 4s electrons gives the stable noble-gas core \([Ar]\), so calcium readily forms Ca²⁺ in ionic compounds such as CaCl₂ and CaCO₃.

Calcium reacts slowly with cold water to form calcium hydroxide and hydrogen gas:

\(\mathrm{Ca(s) + 2\,H_2O(l) \rightarrow Ca(OH)_2(aq/s) + H_2(g)}\)

In air it forms a surface layer of oxide/hydroxide/carbonate which dulls the metal.

Calcination: Limestone decomposes on heating:

\(\mathrm{CaCO_3(s) \xrightarrow{\Delta} CaO(s) + CO_2(g)}\)

Slaking: Quicklime reacts with water:

\(\mathrm{CaO(s) + H_2O(l) \rightarrow Ca(OH)_2(s)}\)

Carbonation (set): Slaked lime reacts with CO₂ to reform limestone:

\(\mathrm{Ca(OH)_2(s) + CO_2(g) \rightarrow CaCO_3(s) + H_2O(l)}\)

Heating excites Ca electrons to higher energy levels; when they relax, emission lines in the red/orange region appear (often described as brick-red). Flame colors help identify metal ions qualitatively.

• CaCO₃ (limestone, chalk): cement, filler, antacids.
• CaO (quicklime): steelmaking flux, soil stabilization.
• Ca(OH)₂ (slaked lime): mortar, water treatment, flue-gas desulfurization.
• CaSO₄·½H₂O (plaster of Paris): casts, building materials.
• CaCl₂: de-icing, drying agent.

Calcium bicarbonate and sulfate in water cause temporary and permanent hardness, respectively.

• Boiling removes temporary hardness by decomposing bicarbonate: \(\mathrm{Ca(HCO_3)_2 \rightarrow CaCO_3\downarrow + CO_2 + H_2O}\).
• Permanent hardness (e.g., CaSO₄) is removed by ion exchange or washing soda: \(\mathrm{CaSO_4 + Na_2CO_3 \rightarrow CaCO_3\downarrow + Na_2SO_4}\).

Bones/teeth contain hydroxyapatite:

\(\mathrm{Ca_5(PO_4)_3OH}\)

In cells, Ca²⁺ acts as a second messenger in muscle contraction, neurotransmission, and hormone signaling; extracellular Ca²⁺ is essential for blood clotting (cofactor for several steps in the cascade).

In a lime kiln, limestone (CaCO₃) is heated to ~900–1100 °C with controlled airflow to drive off CO₂ and produce CaO. Fuel choice and kiln design (shaft/rotary) affect efficiency and purity.

Calcium carbonate effervesces with acids due to CO₂ evolution. Example with HCl:

\(\mathrm{CaCO_3(s) + 2\,HCl(aq) \rightarrow CaCl_2(aq) + CO_2(g) + H_2O(l)}\)

This reaction underlies antacid action and laboratory identification of carbonates.

CaO reacts exothermically with water and is corrosive; avoid skin/eye contact and moisture ingress. Ca(OH)₂ is a strong base (alkaline, irritating). Use gloves, goggles, and dust control; add CaO to water (not the reverse) when slaking, to manage heat safely.