Calcium
Calcium is a chemical element with symbol Ca and atomic number 20. Calcium is a soft gray alkaline earth metal, fifth-most-abundant element by mass in the Earth's crust. The ion Ca2+ is also the fifth-most-abundant dissolved ion in seawater by both molarity and mass, after sodium, chloride, magnesium, and sulfate.
Identity
- Atomic mass
- 40.078(4) u
- Category
- Alkaline earth metals
- Group
- 2
- Period
- 4
- Block
- s
- Oxidation States
- +2
- Phase (25 °C)
- Solid
Electronic Structure
- Electron configuration
- 1s2 2s2 2p6 3s2 3p6 4s2
- Short configuration
- [Ar] 4s2
- Electrons per shell
- 2, 8, 8, 2
- Electronegativity (Pauling)
- 1
- Electron affinity
- 2.37 kJ/mol
- 1st ionization energy
- 590 kJ/mol
- Ionization energies
- 590 · 1145 · 4912 · 6491 · 8153… kJ/mol
Physical properties
- Melting point
- 841.9 °C (1115 K)
- Boiling point
- 1483.8 °C (1757 K)
- Density
- 1.55 g/cm³
- Molar heat capacity
- 25.9 J/(mol·K)
Chemical behavior of calcium
Calcium has two valence electrons in the 4s orbital ([Ar] 4s²), which it easily gives up to form the divalent Ca²⁺ cation with the stable electronic configuration of the argon noble gas. It forms predominantly ionic bonds with halogens, oxygen and other non-metals. As an alkaline earth metal of period 4, it exhibits intermediate reactivity: it is more energetic and electropositive than beryllium and magnesium, but less reactive than strontium and barium. In aqueous solution it easily precipitates in the presence of anions such as carbonate, sulfate, phosphate and oxalate.
Reactivity
Ca + 2 H₂O → Ca(OH)₂ + H₂
It reacts continuously with liquid water at room temperature, releasing hydrogen bubbles and clouding the solution due to the slight solubility of the resulting calcium hydroxide.
2 Ca + O₂ → 2 CaO
In contact with air or when gently heated, it combines with oxygen to form calcium oxide (quicklime), burning with a characteristic orange-red flame.
Ca + Cl₂ → CaCl₂
It reacts vigorously with chlorine gas to form anhydrous calcium chloride, a very stable and highly hygroscopic salt.
Ca + 2 HCl → CaCl₂ + H₂
It dissolves rapidly and with effervescence in dilute mineral acids, releasing hydrogen gas with a notable heat release.
Notable compounds
| Compound | Formula | Use or importance |
|---|---|---|
| Calcium carbonate | CaCO₃ | Constituent of limestones and marbles; raw material for cement, medical antacid and mineral filler |
| Calcium oxide | CaO | Quicklime; manufacture of mortars and cement, neutralisation of iron ores and water treatment |
| Calcium hydroxide | Ca(OH)₂ | Slaked lime; lech in construction, soil acidity correction and sugar industry |
| Calcium sulfate dihydrate | CaSO₄·2H₂O | Natural gypsum; gypsum panels for bricklaying, scales and orthopaedic immobilisation moulds |
| Calcium chloride | CaCl₂ | High-performance road de-icer by exothermic dissolution and industrial desiccant |
Electron shell distribution
View in the atom builderIsotopes
Natural calcium presents six stable isotopes, an extraordinary nuclear diversity: calcium-40 (⁴⁰Ca, 96.94% with 20 neutrons), calcium-44 (⁴⁴Ca, 2.09%), calcium-42 (⁴²Ca, 0.65%), calcium-48 (⁴⁸Ca, 0.19%), calcium-43 (⁴³Ca, 0.14%) and calcium-46 (⁴⁶Ca, 0.004%). The nuclei of ⁴⁰Ca and ⁴⁸Ca are considered 'doubly magic' as they have complete closed shells of protons and neutrons, giving them an unsually nuclear resistance. Calcium-48, with a half-life of millions of times the age of the universe, is used in particle accelerators to synthesise superheavy elements.
Occurrence and production
It is the fifth most abundant element in the Earth's crust (about 4.1% by mass) and the most abundant metallic cation in the human body. It is never found free in nature due to its affinity for oxygen, carbon and sulphur. It constitutes vast sedimentary rock formations of limestone, marble, and dolomite, as well as minerals such as gypsum, anhydrite, fluorite and apatites. Elemental metal is commercially obtained by thermal reduction of calcium oxide with aluminium at high temperature and under vacuum, or by electrolysis of molten calcium chloride.
History and origin of the name
Although its compounds such as lime and gypsum have been used since the Neolithic and the Roman civilisation to make mortars, the pure metal was isolated in 1808 by Humphry Davy by electrolysis of a calcium oxide and mercury amalgam. The British chemist gave it the name 'calcium' from the Latin term 'calx' (calcis), the word that the ancient Romans used to designate quicklime and calcined limestone.
Applications and relevance of calcium
Uses of the element
Metallic calcium is used in ironmaking and metallurgy as a reducing agent, deoxidant and desulphurant in the refining of special steels, copper and nickel alloys. It is also incorporated as a hardener in lead alloys to manufacture automotive battery plates free of maintenance and as a degasser in high-precision vacuum tubes.
Uses of its compounds
Its compounds are pillars of global infrastructure: calcium carbonate and calcium oxide underpin the manufacture of Portland cement and mortar. Gypsum is essential in bricklaying and profiles; calcium chloride is used in winter de-icing; and calcium phosphate constitutes the base of agricultural fertilisers and biocompatible bone implants.
Health, safety and environment
Biological role
It is the most abundant mineral in the human body (about one kilogram in an adult). 99% is mineralised in bones and teeth in the form of hydroxyapatite. The remaining 1% circulates in plasma and cytoplasm acting as a key intracellular messenger in blood clotting, cardiac and skeletal muscle contraction and hormonal secretion.
Safety
Metallic calcium in contact with water or atmospheric moisture releases flammable hydrogen and generates calcium hydroxide caustic, which can cause thermal and chemical burns to skin, mucous membranes and eyes. It must be stored in airtight containers protected with inert gas or dry mineral oil.
Environment
The geochemical calcium cycle buffers the natural acidity of soils and water masses through the carbonate-bicarbonate equilibrium. The precipitation of calcium carbonate on ocean floors acts as the largest carbon sink on geological scale on Earth.
Frequently asked questions about Calcium
What is water hardness and why is it related to calcium?
Water hardness measures the concentration of dissolved polyvalent cations, fundamentally calcium and magnesium ions from limestone aquifers. Hard water makes it difficult for soap to foam by forming insoluble precipitates of calcium stearate and causes scale build-up in pipes, thermoses and washing machines. Although it is annoying for domestic and industrial installations, hard water presents no sanitary risk and provides a complementary mineral source in the human diet.
Why does calcium form in stars and why are its atoms stable?
Calcium is formed by nuclear reactions that take place in stars and stellar explosions. One of its isotopes, calcium-40, is especially stable because its nucleus contains twenty protons and twenty neutrons, a combination that favours a resistant nuclear structure. Calcium produced by these processes is scattered into space and can be incorporated into new stars, planets and other celestial bodies.
Why is vitamin D important to make use of the calcium in food?
Vitamin D helps the intestine to absorb the calcium we obtain from foods such as dairy products, some vegetables and certain fish. If vitamin D is lacking, the body absorbs less calcium and may have difficulty maintaining strong bones. That is why both nutrients are important during growth and throughout life. The amount needed depends on factors such as age, diet and sun exposure.
How does ocean acidification affect the shells of marine animals?
Many molluscs and other marine organisms build their shells with calcium carbonate. When the ocean absorbs more carbon dioxide from the atmosphere, the chemistry of the water changes and the availability of some forms of carbonate that these animals need decreases. This can make it difficult for them to form shells and skeletons, especially in early life stages. The effect varies according to species and water conditions.
How are stalagmites and stalactites formed in karst caves?
Rainwater absorbs carbon dioxide from air and soil, turning into a slightly acidic solution of carbonic acid that dissolves underground limestone rock as soluble calcium bicarbonate. As it slowly drips from a cave ceiling with lower CO₂ pressure, the gas escapes and the bicarbonate reverts to insoluble calcium carbonate. This mineral precipitates drop by drop over thousands of years, raising stalagmites on the floor and forming stalactites from the ceiling.
How does calcium help muscles move?
Each time the brain sends a signal to move a muscle, muscle cells release a wave of stored calcium ions. This calcium acts as a key that activates contractile proteins (actin and myosin), allowing them to slide against each other and shorten the muscle. On pumping calcium back, the muscle relaxes.
Why is calcium chloride used to melt snow on roads?
Unlike traditional table salt, calcium chloride releases heat when it comes into contact with water and ambient humidity. In addition to lowering the freezing point to extreme temperatures of almost thirty degrees below zero, it generates its own heat to melt ice much faster on motorways during severe snowfalls.
Main sources consulted: PubChem, Royal Society of Chemistry, NIST, IUPAC. Last revision: November 2026.