Carbonates
Summary: Metal carbonates thermally decompose to metal oxide + CO₂ (except Group 1 carbonates which are stable). The limestone cycle: CaCO₃ → CaO + CO₂, CaO + H₂O → Ca(OH)₂, Ca(OH)₂ + CO₂ → CaCO₃. Tags: igcse chemistry reactions Created: 2026-07-14 Last Updated: 2026-07-14
Content
What Are Metal Carbonates?
Metal carbonates are ionic compounds containing the carbonate ion, CO₃²⁻, combined with a metal cation. The carbonate ion consists of a carbon atom covalently bonded to three oxygen atoms, carrying an overall charge of 2−. Common examples include:
| Carbonate | Formula | Common Name / Occurrence |
|---|---|---|
| Calcium carbonate | CaCO₃ | Limestone, chalk, marble |
| Sodium carbonate | Na₂CO₃ | Washing soda, soda ash |
| Potassium carbonate | K₂CO₃ | Potash |
| Copper(II) carbonate | CuCO₃ | Malachite (green mineral) |
| Zinc carbonate | ZnCO₃ | Smithsonite |
| Magnesium carbonate | MgCO₃ | Magnesite |
| Lithium carbonate | Li₂CO₃ | Used in mood-stabilising medication |
Most metal carbonates are insoluble in water. The notable exceptions are the carbonates of Group 1 elements (sodium carbonate, potassium carbonate) and ammonium carbonate, all of which dissolve readily.
Thermal Decomposition of Carbonates
Thermal decomposition is the breaking down of a compound by heating. When most metal carbonates are heated strongly, they decompose to form a metal oxide and carbon dioxide gas:
General equation: [ \text{Metal carbonate} \rightarrow \text{Metal oxide} + \text{CO}_2 ]
Examples:
- Copper(II) carbonate: CuCO₃ (s) → CuO (s) + CO₂ (g)
- Zinc carbonate: ZnCO₃ (s) → ZnO (s) + CO₂ (g)
- Calcium carbonate: CaCO₃ (s) → CaO (s) + CO₂ (g)
- Magnesium carbonate: MgCO₃ (s) → MgO (s) + CO₂ (g)
The ease of thermal decomposition depends on the position of the metal in the reactivity series. The more reactive the metal, the more stable its carbonate — meaning a higher temperature is required to decompose it. This trend is explained by the polarising power of the metal cation: smaller, more highly charged cations (such as those lower in Group 2) distort the carbonate ion more, weakening the C–O bonds and making decomposition easier at lower temperatures.
Group 1 Carbonates: Exception to Decomposition
Carbonates of Group 1 metals (except lithium) do not decompose when heated with a Bunsen burner. Sodium carbonate (Na₂CO₃) and potassium carbonate (K₂CO₃) are thermally stable up to very high temperatures — far beyond what a laboratory Bunsen flame can achieve (approximately 900–1000 °C at its hottest).
Why are Group 1 carbonates stable? Group 1 cations (Na⁺, K⁺) are large with a low charge density (+1 charge on a relatively large ion). They exert weak polarising power on the carbonate ion, so they do not significantly distort the electron cloud of CO₃²⁻. The C–O bonds within the carbonate ion remain strong, and thermal decomposition does not occur under normal heating.
Lithium carbonate (Li₂CO₃) is an exception. Li⁺ is the smallest Group 1 cation and has a higher charge density. It exerts sufficient polarising power to destabilise the carbonate ion, and Li₂CO₃ does decompose on strong heating:
[ \text{Li}_2\text{CO}_3 \rightarrow \text{Li}_2\text{O} + \text{CO}_2 ]
This behaviour is more similar to Group 2 carbonates than to other Group 1 carbonates — an example of the diagonal relationship between lithium and magnesium in the periodic table.
Copper(II) Carbonate Decomposition (Demonstration)
Heating copper(II) carbonate is a classic laboratory demonstration of thermal decomposition with a clear colour change:
- Before heating: Green powder (copper(II) carbonate, CuCO₃)
- During heating: The green powder turns black as copper(II) oxide (CuO) forms
- Gas evolved: Carbon dioxide (CO₂), which is colourless and odourless
[ \text{CuCO}_3\text{ (s, green)} \rightarrow \text{CuO (s, black)} + \text{CO}_2\text{ (g)} ]
The colour change from green to black is the key visual observation that confirms decomposition has occurred. The carbon dioxide produced can be identified by bubbling it through limewater.
Test for Carbon Dioxide: Limewater
The standard test for carbon dioxide gas uses limewater (a saturated solution of calcium hydroxide, Ca(OH)₂):
- Bubble the unknown gas through limewater, or shake a test tube of the gas with a small volume of limewater
- If CO₂ is present, the limewater turns milky or cloudy white
Chemical explanation: CO₂ reacts with dissolved calcium hydroxide to form a precipitate of insoluble calcium carbonate:
[ \text{Ca(OH)}_2\text{ (aq)} + \text{CO}_2\text{ (g)} \rightarrow \text{CaCO}_3\text{ (s)} + \text{H}_2\text{O (l)} ]
Calcium carbonate is the white solid responsible for the cloudiness. If CO₂ is bubbled through for a prolonged period, the milky suspension may clear because the excess CO₂ reacts with CaCO₃ to form soluble calcium hydrogencarbonate:
[ \text{CaCO}_3\text{ (s)} + \text{CO}_2\text{ (g)} + \text{H}_2\text{O (l)} \rightarrow \text{Ca(HCO}_3\text{)}_2\text{ (aq)} ]
This secondary reaction explains why persistent bubbling of CO₂ through limewater eventually produces a clear solution.
The Limestone Cycle
The limestone cycle describes the interconversions between calcium carbonate (limestone), calcium oxide (quicklime), and calcium hydroxide (slaked lime). It is a closed loop driven by heat and chemical reaction.
Step 1: Thermal Decomposition of Limestone
Calcium carbonate (limestone) is heated strongly to produce calcium oxide (quicklime) and carbon dioxide:
[ \text{CaCO}_3\text{ (s)} \xrightarrow{\sim 900^\circ\text{C}} \text{CaO (s)} + \text{CO}_2\text{ (g)} ]
This is carried out industrially in a lime kiln, where limestone is heated to approximately 900–1000 °C. The reaction is endothermic and requires a continuous input of heat energy. Carbon dioxide gas is released as a by-product.
Step 2: Slaking of Quicklime
Calcium oxide reacts vigorously with water in a highly exothermic reaction to produce calcium hydroxide (slaked lime):
[ \text{CaO (s)} + \text{H}_2\text{O (l)} \rightarrow \text{Ca(OH)}_2\text{ (s)} ]
This process is called slaking. It releases a large amount of heat — enough to cause the water to boil and produce steam. The resulting product, calcium hydroxide, is a white powder only slightly soluble in water. A saturated solution of Ca(OH)₂ is known as limewater.
Step 3: Carbonation (Limewater Test)
Calcium hydroxide reacts with carbon dioxide to regenerate calcium carbonate, completing the cycle:
[ \text{Ca(OH)}_2\text{ (aq)} + \text{CO}_2\text{ (g)} \rightarrow \text{CaCO}_3\text{ (s)} + \text{H}_2\text{O (l)} ]
This is the same reaction that occurs during the limewater test for CO₂. The white precipitate of calcium carbonate formed is chemically identical to the original limestone.
Summary Diagram of the Limestone Cycle
CaCO₃ ──heat (~900°C)──> CaO + CO₂
^ │
│ │ + H₂O (slaking, exothermic)
│ ▼
└──── + CO₂ ──── Ca(OH)₂ ── + more CO₂ (excess) ──> Ca(HCO₃)₂ (aq)
(limewater)
Uses of Limestone (CaCO₃)
| Use | Explanation |
|---|---|
| Building material | Limestone blocks are used directly in construction; durable and readily available |
| Cement manufacture | Heated with clay in a rotary kiln to produce cement clinker; cement is mixed with sand and aggregate to make concrete |
| Glass manufacture | Heated with sand (SiO₂) and sodium carbonate (Na₂CO₃); CaCO₃ provides CaO which acts as a stabiliser in the glass |
| Iron extraction (blast furnace) | Limestone is added to remove sandy impurities (silica, SiO₂) as molten slag: CaCO₃ → CaO + CO₂, then CaO + SiO₂ → CaSiO₃ (calcium silicate slag) |
| Neutralising acidic soils | Powdered limestone is spread on fields to raise soil pH (less effective than quicklime or slaked lime but slower-acting and gentler) |
| Toothpaste and indigestion tablets | Finely ground CaCO₃ is used as a mild abrasive and as a source of calcium |
Uses of Calcium Oxide, CaO (Quicklume)
| Use | Explanation |
|---|---|
| Drying agent | CaO reacts with water, so it can be used to dry gases that do not react with it (e.g. ammonia gas is dried by passing through a column of CaO) |
| Neutralising acidic soils | Quicklime is more effective than limestone at raising soil pH because it is more soluble; it reacts with water in the soil to produce Ca(OH)₂ which neutralises acids |
| Steelmaking | Added to remove acidic impurities such as SiO₂ and P₄O₁₀ as slag in the basic oxygen steelmaking process |
Uses of Calcium Hydroxide, Ca(OH)₂ (Slaked Lime)
| Use | Explanation |
|---|---|
| Agriculture | Spread on fields to neutralise acidic soils (lime treatment); Ca(OH)₂ + 2H⁺ → Ca²⁺ + 2H₂O |
| Mortar | Mixed with sand and water; mortar sets as Ca(OH)₂ reacts with CO₂ from the air forming CaCO₃ crystals that bind the sand particles together |
| Limewater | Used as a laboratory reagent to test for the presence of CO₂ |
| Water treatment | Used to soften hard water by precipitating calcium and magnesium ions |
Key Concepts from Past Papers
- thermal decomposition is breaking down a compound by heating
- metal carbonate → metal oxide + carbon dioxide
- CO₂ test: limewater turns milky (Ca(OH)₂ + CO₂ → CaCO₃ + H₂O)
- Group 1 carbonates (except Li₂CO₃) are thermally stable
- limestone cycle: CaCO₃ → CaO → Ca(OH)₂ → CaCO₃
- CaO + H₂O is exothermic (slaking)
- limestone in blast furnace removes silica impurities as slag (CaSiO₃)
- uses of CaO (drying agent, soil neutralisation), Ca(OH)₂ (agriculture, mortar), CaCO₃ (building, cement, glass)
Keywords from Past Papers
carbon, dioxide, calcium, reacts, carbonate, oxide, values, irritates, monoxide, limewater, iron, formed, solution, neutralises, inclusive
Related Notes
Sources
- OpenStax, Chemistry 2e, Chapter 18: Representative Metals, Metalloids, and Nonmetals — Section 18.1, occurrence, preparation, and properties of carbonates. https://openstax.org/books/chemistry-2e/
- BBC Bitesize, GCSE Chemistry (Single Science), “Limestone” and “Thermal Decomposition” topics. https://www.bbc.co.uk/bitesize/topics/z3c3qyc
- Cambridge IGCSE Chemistry 0620 Syllabus, Topic 10: Metals — thermal decomposition of metal carbonates; Topic 13: Carbonates — manufacture and uses of lime.
- CK-12 Foundation, Chemistry, Chapter: Acids and Bases — carbonates and their reactions. https://www.ck12.org/chemistry/
Past Paper Sources
- 0620/32 Feb/March 2020: Q22(b)(ii) (1m)
- 0620/32 Feb/March 2023: Q77(a)(ii) (4m)
- 0620/33 October/November 2016: Q55(e)(iii) (0m)
- 0620/33 October/November 2020: Q44(c)(i) (1m), Q44(c)(i) (1m)
- 0971/32 May/June 2023: Q88(a)(iv) (1m)
- 0971/32 October/November 2021: Q33(b)(i) (2m), Q44(a)(i) (2m), Q44(c)(i) (2m) (+21 more)
Common Misconceptions
- All carbonates decompose on heating. Group 1 carbonates (except Li₂CO₃) are thermally stable and do not decompose with a Bunsen burner due to the low charge density of Group 1 cations.
- Limewater turns milky because CaCO₃ is produced, so the solution is now calcium carbonate. Limewater is not “turning into” calcium carbonate — CaCO₃ is an insoluble precipitate suspended in the water, which is what gives the cloudy/milky appearance.
- Quicklime (CaO) and slaked lime (Ca(OH)₂) are the same substance. They are distinct compounds: CaO is produced by thermal decomposition and reacts exothermically with water to form Ca(OH)₂. They have different uses and different chemical properties.
- The limestone cycle only works in one direction. The key insight is that it is a cycle: the same calcium atoms can move between CaCO₃, CaO, and Ca(OH)₂ repeatedly through the appropriate reactions.
- All carbonates produce the same colour change on heating. The colour change observed depends on the specific carbonate. Copper(II) carbonate changes from green to black (CuCO₃ → CuO). Calcium carbonate and zinc carbonate are white and remain white (CaCO₃ → CaO, ZnCO₃ → ZnO), so no colour change is visible — only CO₂ evolution confirms decomposition.
- Ca(OH)₂ is highly soluble because it is used in limewater. Calcium hydroxide is only sparingly soluble (about 0.16 g per 100 mL at 20 °C). Limewater is a saturated solution with a low concentration of dissolved Ca(OH)₂.