Chemistry 2e · Representative Metals, Metalloids, and Nonmetals

Occurrence, Preparation, and Properties of Carbonates

10 min read

Molar masses, STP molar volume (22.4 L/mol), and the carbonate/bicarbonate chemistry described are standard textbook material; verify quantitative values against current primary sources before formal citation.

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In 30 seconds
Why this matters
The college version
Eli explains
Worked example
Key takeaway
Check yourself
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In 30 seconds

Carbonates are the salts of carbonic acid, H₂CO₃, containing the CO₃^2- — a flat, trigonal-planar ion whose negative charge is delocalized across three oxygens. They are among the most abundant minerals on Earth: (CaCO₃), marble, chalk, and dolomite (CaMg(CO₃)₂) are mountains of carbonate, and coral reefs, seashells, and eggshells are biogenic calcium carbonate. The carbonate system is also the main buffer of seawater and blood.

Preparation mirrors the acid–base origin: carbonate salts form when carbon dioxide reacts with a base or when carbonic acid is neutralized. Industrially, sodium carbonate (soda ash, Na₂CO₃) comes from the , and sodium bicarbonate (baking soda, NaHCO₃) decomposes on heating to release CO₂ — the leavening action behind cakes and the chemistry of some fire extinguishers.

Carbonate chemistry is dominated by two reactions: reaction with acids, which releases CO₂ with (the "fizz" test), and thermal decomposition, which releases CO₂ on heating to leave the metal oxide. Solubility divides the family: alkali-metal carbonates (except Li₂CO₃) are soluble and basic; almost all others are insoluble. These patterns explain limestone caves, hard water, antacids, and the carbon cycle.

Why this matters

The built environment: limestone and its products — lime (CaO), slaked lime (Ca(OH)₂), and Portland cement — are the raw materials of concrete, mortar, and glass.
Everyday chemistry: baking soda (NaHCO₃) leavens baked goods, neutralizes acid spills and heartburn, and puts out small grease fires.
Geology and climate: limestone caves form because CO₂-laden water dissolves CaCO₃ (reprecipitating it as stalactites); carbonate weathering is a major term in the global carbon cycle.
Health: calcium carbonate is the active ingredient in many antacids and calcium supplements; carbonate and bicarbonate ions buffer blood pH; kidney stones are often calcium carbonate or oxalate deposits.
Exams: expect the fizz test, the limewater test for CO₂, thermal decomposition products, and solubility rules.

The college version

Core Concepts

Structure and bonding of the carbonate ion

CO₃^2- is trigonal planar (120° bond angles); the three C–O bonds are equivalent in the resonance hybrid (bond order 4/3 each). The sp² carbon's perpendicular p orbital overlaps with oxygen p orbitals, delocalizing the π electrons and negative charge across the ion — which makes CO₃²⁻ unusually stable and carbonate minerals widespread and durable.

Occurrence: carbonate minerals

Calcium carbonate (CaCO₃): limestone, marble, chalk, calcite, aragonite — the most important carbonate mineral.
Dolomite (CaMg(CO₃)₂): a double carbonate of calcium and magnesium, abundant in sedimentary rock.
Sodium carbonate: occurs naturally in arid regions (trona); historically from plant ashes, today made industrially.
Potassium carbonate (K₂CO₃, potash): from wood ashes; used in soaps, glass, and fertilizer.
Other metal carbonates (malachite, Cu₂CO₃(OH)₂; siderite, FeCO₃) are minor ores.

Preparation

From CO₂ and a base: 2NaOH + CO₂ → Na₂CO₃ + H₂O; with excess CO₂ this forms the bicarbonate: Na₂CO₃ + CO₂ + H₂O → 2NaHCO₃.
Neutralization of carbonic acid: H₂CO₃ + 2NaOH → Na₂CO₃ + 2H₂O.
The Solvay process (industrial Na₂CO₃): brine is saturated with ammonia and CO₂, precipitating the less-soluble sodium bicarbonate: NaCl + NH₃ + CO₂ + H₂O → NaHCO₃(s) + NH₄Cl. Heating gives soda ash — 2NaHCO₃ → Na₂CO₃ + CO₂ + H₂O — with the CO₂ recycled and ammonia regenerated from NH₄Cl with lime.

Reaction with acids: the fizz test

All carbonates react with acids to release carbon dioxide:

$CaCO 3 (s) + 2 HCl (aq) \to CaCl 2 (aq) + CO 2 (g) + H 2 O (l)$

The effervescence is the standard qualitative test for carbonate minerals; the CO₂ evolved turns limewater milky: CO₂ + Ca(OH)₂ → CaCO₃(s) + H₂O. Bicarbonates react similarly but consume only one H⁺: NaHCO₃ + HCl → NaCl + CO₂ + H₂O. This is why acid rain dissolves limestone — and why antacids fizz in the stomach.

Thermal decomposition: the lime cycle

Heated strongly, metal carbonates decompose to the metal oxide and CO₂:

$CaCO 3 (s) Δ ⟶ CaO (s) + CO 2 (g)$

This is , the first step of the (CaCO₃ → CaO → Ca(OH)₂ → CaCO₃). Heat stability increases with the cation's size and electropositivity: alkali-metal carbonates (except Li₂CO₃) survive furnace heat, while most others decompose — with the required temperature rising down a group (MgCO₃ < CaCO₃ < BaCO₃). Smaller, more highly charged cations destabilize the carbonate ion — an exam favorite.

Solubility and basicity: two camps

Soluble and basic: alkali-metal carbonates (Na₂CO₃, K₂CO₃) and ammonium carbonate dissolve to give basic solutions because the carbonate ion hydrolyzes: CO₃^2- + H₂O ⇌ HCO₃^- + OH^-. Soda ash is the mild base in washing soda and water treatment.
Insoluble: carbonates of most other metals (Ca²⁺, Mg²⁺, Ba²⁺, Fe²⁺, Cu²⁺) are insoluble — the basis of the qualitative-analysis test where adding CO₃²⁻ precipitates the metal carbonate.
Bicarbonates are different: NaHCO₃ is soluble, and so is Ca(HCO₃)₂ even though CaCO₃ is not. Rainwater containing dissolved CO₂ converts insoluble limestone into soluble calcium bicarbonate, CaCO₃ + CO₂ + H₂O → Ca(HCO₃)₂, the chemistry of — boiling drives off CO₂ and reprecipitates the kettle-scale CaCO₃.

Uses at a glance

CaCO₃: construction aggregate, cement, lime, antacids and calcium supplements, paper and paint fillers.
Na₂CO₃: glassmaking, soaps and detergents, water softening, pH regulation.
NaHCO₃: baking powder, antacids, fire extinguishers (decomposes to smothering CO₂).

Common Confusions

Do Not Confuse	With	Difference
Carbonate (CO₃²⁻)	Bicarbonate (HCO₃⁻)	Carbonate needs 2 H⁺ to release CO₂ and is mostly insoluble; bicarbonate needs 1 H⁺ and is soluble (NaHCO₃ = baking soda).
CaCO₃ solubility	Ca(HCO₃)₂ solubility	CaCO₃ is insoluble; calcium bicarbonate is soluble — that's how caves dissolve and hard water forms scale.
Calcination	Roasting	Calcination decomposes a carbonate by heat, releasing CO₂; roasting converts sulfide ores to oxides in air.
Lime (CaO)	Limestone (CaCO₃)	Lime is the product of heating limestone; slaking lime with water gives Ca(OH)₂.
"Fizz test" for carbonate	Any bubbling liquid	Effervescence with acid specifically signals carbonate/bicarbonate; verify with limewater turning milky.
All carbonates decompose easily	Alkali-metal carbonates	Na₂CO₃ and K₂CO₃ survive furnace heat; only Li₂CO₃ among alkali carbonates decomposes readily.
Carbonate solutions neutral	Carbonate solutions basic	CO₃²⁻ hydrolysis produces OH⁻, so soda ash solutions are basic (pH > 7).

Eli explains

The same idea, in plain words

Explain it like I’m 10

Carbonates are rocks and shells made of carbon, oxygen, and a metal — like the limestone that makes up caves and the chalk in your classroom. If you pour a sour liquid (an acid) on them, they bubble and fizz because they're releasing a gas called carbon dioxide — that's how bakers make cakes rise with baking soda. And if you heat limestone really hot, it turns into quicklime, which people have used for thousands of years to make mortar and cement.

Worked example

Example 1: Lime burning — mass of quicklime from impure limestone

Limestone is calcined to quicklime:

$CaCO 3 (s) Δ ⟶ CaO (s) + CO 2 (g)$

Problem: A kiln charges 250.0 g of limestone that is 85.0% CaCO₃ by mass (the rest is inert gangue). What mass of CaO is produced? Molar masses: M(CaCO₃) = 40.08 + 12.01 + 3(16.00) = 100.09 g mol^-1, M(CaO) = 40.08 + 16.00 = 56.08 g mol^-1.

Plan and formula: isolate the pure carbonate, m_pure = m_sample × purity; then n = mM, use the 1:1 mole ratio, then m = n × M.

Substitution (dimensional analysis):

$m(CaCO 3) = 250.0 g limestone \times 85.0 g CaCO 3 100 g limestone = 212.5 g$

$n(CaCO 3) = 212.5 g \times 1 mol 100.09 g = 2.123 mol$

$n(CaO) = 2.123 mol CaCO 3 \times 1 mol CaO 1 mol CaCO 3 = 2.123 mol$

$m(CaO) = 2.123 mol \times 56.08 g 1 mol = 119.1 g$

Answer: about 119 g of quicklime. The 212.5 g of pure carbonate lost 93 g of CO₂ (about 44% of its mass) — which is why lime burning is a major source of CO₂.

Example 2: The fizz test quantified — CO₂ volume from a carbonate tablet

An antacid tablet of pure CaCO₃ reacts with excess stomach acid:

$CaCO 3 (s) + 2 HCl (aq) \to CaCl 2 (aq) + CO 2 (g) + H 2 O (l)$

Problem: What volume of CO₂ (at STP, 0 °C, 1 atm) is released when a 500.0 mg (0.5000 g) tablet reacts completely? M(CaCO₃) = 100.09 g mol^-1, V_m = 22.4 L mol^-1.

Plan and formula: n = mM, 1:1 mole ratio of CaCO₃ to CO₂, then V = n × V_m.

Substitution:

$n(CaCO 3) = 0.5000 g \times 1 mol 100.09 g = 4.996 \times 10 -3 mol$

$n(CO 2) = 4.996 \times 10 -3 mol CaCO 3 \times 1 mol CO 2 1 mol CaCO 3 = 4.996 \times 10 -3 mol$

$V(CO 2) = 4.996 \times 10 -3 mol \times 22.4 L 1 mol = 0.112 L$

Answer: about 112 mL of CO₂ — a satisfying burp's worth of gas from half a gram of carbonate, and a vivid demonstration of why carbonate antacids fizz.

Key takeaways

Carbonate ion CO₃^2-: trigonal planar, sp² carbon, resonance-delocalized charge; stable and widespread in minerals.
Major minerals: limestone/marble/chalk (CaCO₃), dolomite (CaMg(CO₃)₂); biogenic CaCO₃ in shells and coral.
Acid test: carbonate + acid → CO₂ gas (effervescence); CO₂ turns limewater milky. CaCO₃ + 2HCl → CaCl₂ + CO₂ + H₂O.
Thermal decomposition (calcination): CaCO₃ ^Δ⟶ CaO + CO₂; stability to heat increases down a group (MgCO₃ < CaCO₃ < BaCO₃); alkali carbonates (except Li₂CO₃) survive furnace heat.
Solubility: alkali-metal and ammonium carbonates soluble and basic (hydrolysis); nearly all others insoluble. Bicarbonates are more soluble than carbonates (Ca(HCO₃)₂ is soluble — temporary hardness).
Solvay process: NaCl + NH₃ + CO₂ + H₂O → NaHCO₃↓ (+ NH₄Cl), then 2NaHCO₃ → Na₂CO₃ + CO₂ + H₂O.
Industrial loop: limestone → lime (CaO) → slaked lime (Ca(OH)₂) → cement/mortar/glass; carbonate chemistry links to acid rain, cave formation, and the carbon cycle.

Check yourself

5 review questions from the chapter. Try each one, then open the answer.

Write the reaction of CaCO₃ with HCl, and describe how you would confirm that the gas evolved is CO₂.
Show answer
CaCO₃ + 2HCl → CaCl₂ + CO₂ + H₂O. Bubbling suggests a carbonate; passing the gas through limewater (aqueous Ca(OH)₂) makes it turn milky — the classic CO₂ test.
Why does a carbonate's decomposition temperature decrease as the cation gets smaller and more highly charged (e.g., MgCO₃ decomposes more easily than BaCO₃)?
Show answer
A small, highly charged cation (like Mg²⁺) strongly polarizes the carbonate ion, pulling electron density away from the C–O bonds and destabilizing CO₃²⁻ — so less thermal energy is needed to break it apart. Larger, less polarizing cations (Ba²⁺) destabilize the ion less, so their carbonates need higher temperatures.
Sodium carbonate solution has a pH above 7. Explain why using an equation.
Show answer
CO₃^2- + H₂O ⇌ HCO₃^- + OH^-. The carbonate ion hydrolyzes, generating hydroxide ions, which makes the solution basic.
How many grams of Na₂CO₃ form when 168.0 g of NaHCO₃ is heated? (2NaHCO₃ → Na₂CO₃ + CO₂ + H₂O; M(NaHCO₃) = 84.01 g mol^-1, M(Na₂CO₃) = 105.99 g mol^-1.)
Show answer
n(NaHCO₃) = 168.0/84.01 = 2.000 mol; the 2:1 mole ratio gives n(Na₂CO₃) = 1.000 mol; m = 1.000 × 105.99 = 106.0 g.
What is temporary hardness, what causes it, and how is it removed by boiling?
Show answer
Temporary hardness is caused by dissolved calcium (or magnesium) bicarbonate, Ca(HCO₃)₂. Boiling drives off CO₂ and reverses the dissolution reaction, reprecipitating insoluble CaCO₃ — the scale that coats kettles — leaving softened water.

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Study tools & related lessonsKey vocabulary · Related

Key vocabulary

Carbonate ion: CO₃^2-, trigonal-planar ion with delocalized charge
Bicarbonate (hydrogen carbonate): HCO₃^-, carbonate with one H⁺ added back
Limestone: Sedimentary rock composed mainly of CaCO₃
Calcination: Heating a carbonate to drive off CO₂, leaving the oxide
Lime cycle: CaCO₃ → CaO → Ca(OH)₂ → CaCO₃
Effervescence: Bubbling caused by escaping gas
Temporary hardness: Hardness from dissolved Ca(HCO₃)₂
Solvay process: Industrial route to Na₂CO₃ from NaCl, NH₃, and CO₂
Hydrolysis of CO₃²⁻: Carbonate reacting with water to give OH⁻

Sources & references

openstax.org — Chemistry 2e

This lesson was adapted from the open educational references above; their licenses and attributions are preserved. See Copyright & Licensing.

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In 30 seconds

Why this matters

The college version

Core Concepts

Structure and bonding of the carbonate ion

Occurrence: carbonate minerals

Preparation

Reaction with acids: the fizz test

Thermal decomposition: the lime cycle

Solubility and basicity: two camps

Uses at a glance

Common Confusions

The same idea, in plain words

Worked example

Example 1: Lime burning — mass of quicklime from impure limestone

Example 2: The fizz test quantified — CO₂ volume from a carbonate tablet

Key takeaways

Check yourself

Keep learning

Key vocabulary

Related lessons

Sources & references