Occurrence, Preparation, and Properties of Carbonates

Core Concepts

Structure of the carbonate ion

The carbonate ion is trigonal planar: a central carbon atom bonded to three oxygen atoms, with carbon using sp² hybrid orbitals. Because of resonance, the three C–O bonds are identical — each has a bond order of about 1.33, halfway between single and double. Carbonate is the conjugate base of bicarbonate (HCO₃⁻), which is itself the conjugate base of carbonic acid (H₂CO₃). The three species interconvert rapidly in water:

CO₂ + H₂O ⇌ H₂CO₃ ⇌ H⁺ + HCO₃^- ⇌ 2H⁺ + CO₃^2-

This equilibrium is the engine of the whole topic: add CO₂ and the balance shifts right; remove CO₂ (or add base) and carbonate forms.

Occurrence: the carbonate minerals

Most carbonate rock is biogenic — precipitated by living things. Marine organisms build shells and coral from CaCO₃; when they die, their remains compact into limestone. Heat and pressure recrystallize limestone into marble. Other common carbonates include dolomite (CaMg(CO₃)₂), magnesite (MgCO₃), and siderite (FeCO₃); sodium carbonates (trona, natron) crystallize from evaporating lakes.

Caves show the cycle in action. Rainwater absorbs CO₂ and becomes slightly acidic; it dissolves limestone as soluble calcium bicarbonate:

CaCO₃(s) + CO₂ + H₂O ⟶ Ca²⁺ + 2HCO₃^-

When the water drips into a cave, CO₂ escapes, the equilibrium reverses, and CaCO₃ reprecipitates as stalactites and stalagmites.

Preparation: from CO₂, from lime, and the Solvay process

From CO₂ and a base — the lab route. Carbon dioxide reacts with strong base to give carbonate: CO₂ + 2NaOH ⟶ Na₂CO₃ + H₂O. Bubbling CO₂ through limewater gives a white precipitate of CaCO₃ — the standard test for CO₂:

Ca(OH)₂(aq) + CO₂(g) ⟶ CaCO₃(s) + H₂O(l)

The lime cycle — the industrial heart of carbonate chemistry. Heating limestone drives off CO₂ (calcination):

CaCO₃(s) ^Δ⟶ CaO(s) + CO₂(g)

Adding water to the quicklime gives slaked lime, Ca(OH)₂, which slowly reabsorbs CO₂ from the air, returning to CaCO₃ — the same atoms recycled between rock, lime, and air.

The Solvay process — soda ash from salt and limestone. The overall reaction is 2NaCl + CaCO₃ ⟶ Na₂CO₃ + CaCl₂, achieved in steps:

1. Ammonia-saturated brine absorbs CO₂: NaCl + NH₃ + CO₂ + H₂O ⟶ NaHCO₃ + NH₄Cl.
2. Heating the bicarbonate: 2NaHCO₃ ^Δ⟶ Na₂CO₃ + CO₂ + H₂O.
3. Treating the ammonium chloride with slaked lime regenerates ammonia: Ca(OH)₂ + 2NH₄Cl ⟶ CaCl₂ + 2NH₃ + 2H₂O.

Ammonia is recycled; the only wastes are CaCl₂ and the CO₂ from step 2 (returned to step 1).

Properties: heat, acid, and water

Thermal decomposition. Carbonates of metals decompose on heating: MCO₃ ⟶ MO + CO₂. Stability increases with the size and electropositive character of the cation: MgCO₃ decomposes at a few hundred degrees Celsius, CaCO₃ near 840 °C, and Na₂CO₃ and K₂CO₃ survive even strong heating without decomposing. This trend explains why limestone (not magnesite) is the practical source of lime, and why the alkali carbonates are sold as the stable "soda ash."

Reaction with acids. All carbonates and bicarbonates react with acids, releasing CO₂ with effervescence:

CO₃^2- + 2H⁺ ⟶ CO₂(g) + H₂O

HCO₃^- + H⁺ ⟶ CO₂(g) + H₂O

This fizzing is the qualitative test for carbonate, and it is why antacids (CaCO₃, NaHCO₃) relieve excess stomach acid — while releasing harmless CO₂ gas.

Solubility and hardness. Alkali metal carbonates (except sparingly soluble Li₂CO₃) and ammonium carbonate are soluble; all other common carbonates are insoluble. Bicarbonates of the alkali metals are soluble. Dissolved calcium bicarbonate, Ca(HCO₃)₂, causes temporary hardness: boiling expels CO₂ and precipitates CaCO₃ (the scale in kettles):

Ca(HCO₃)₂ ^Δ⟶ CaCO₃(s) + CO₂(g) + H₂O

Adding washing soda (Na₂CO₃) removes both Ca²⁺ and Mg²⁺ by precipitating their carbonates — the classic water-softening reaction.

Examples: carbonate stoichiometry

Example 1 — The acid test: how much CO₂ bubbles off?

A geologist adds excess hydrochloric acid to 25.0 g of limestone (assume pure CaCO₃). Reaction: CaCO₃ + 2HCl ⟶ CaCl₂ + H₂O + CO₂. What volume of CO₂ (at STP, 22.4 L/mol) is released? Moles of carbonate first:

n(CaCO₃) = 25.0 g100.09 g/mol = 0.250 mol

The 1:1 CO₂ : CaCO₃ ratio gives:

V(CO₂) = 0.250 mol × 22.4 L1 mol = 5.59 L

5.59 L of CO₂ — enough to turn a small flask of limewater milky.

Example 2 — Lime from limestone: mass with percent purity

A kiln charges 500 g of limestone that is 95.0% CaCO₃. Using CaCO₃ ^Δ⟶ CaO + CO₂, what mass of quicklime (CaO) results? Isolate the pure carbonate:

m(CaCO₃) = 0.950 × 500 g = 475 g

n(CaCO₃) = 475 g100.09 g/mol = 4.75 mol

The 1:1 CaO : CaCO₃ ratio, then mass:

m(CaO) = 4.75 mol × 56.08 g1 mol = 266 g

About 266 g of quicklime (the rest of the mass leaves as CO₂ — a 44% mass loss that matters for cement economics).

Example 3 — The Solvay process: soda ash from salt

How much Na₂CO₃ can be made from 1.00 kg of NaCl, per the overall reaction 2NaCl + CaCO₃ ⟶ Na₂CO₃ + CaCl₂? Moles of salt:

n(NaCl) = 1000 g58.44 g/mol = 17.1 mol

Two NaCl per Na₂CO₃, then mass:

m(Na₂CO₃) = 17.1 mol NaCl × 1 mol Na₂CO₃2 mol NaCl × 105.99 g1 mol = 907 g

907 g of soda ash from a kilogram of salt — the scale that makes the Solvay process an industrial giant.

Common Confusions