Calcium - Ca
Calcium (Ca) is a silvery-white alkaline earth metal in group 2. In natural waters it occurs mainly as dissolved calcium ions (Ca2+), which contribute to hardness alongside magnesium. Calcium metal and calcium compounds have different properties and uses.
Properties · Uses and compounds · Health and safety · Environment and water
| Atomic number | 20 |
|---|---|
| Standard atomic weight | 40.078(4) |
| Electronegativity (Pauling scale) | 1.00 |
| Density at 20 °C | Approximately 1.54 g cm−3 |
| Melting point | 842 °C |
| Boiling point | 1484 °C |
| Non-bonded atomic radius | 0.231 nm |
| Common oxidation state | +2 |
| Naturally occurring isotopes | 40Ca, 42Ca, 43Ca, 44Ca, 46Ca and 48Ca |
| Electron configuration | [Ar] 4s2 |
| First ionisation energy | 589.83 kJ mol−1 |
| Second ionisation energy | 1145.447 kJ mol−1 |
| Isolation of metal | Humphry Davy, 1808 |
Property sources: Royal Society of Chemistry and CIAAW. Standard atomic weight is dimensionless; the corresponding molar mass is approximately 40.078 g mol−1. Parentheses indicate uncertainty in the final digits. Naturally occurring isotopes should not be confused with the total number of known isotopes or described collectively as stable.
Properties and occurrence of calcium
Humphry Davy first isolated calcium metal in 1808. Natural calcium is dominated by 40Ca; CIAAW lists six isotopes contributing to its usual natural isotopic composition. Calcium occurs widely in minerals and compounds, rather than as free metal under ordinary environmental conditions.
Important calcium-bearing minerals include calcite and aragonite, both forms of calcium carbonate, as well as gypsum, fluorite and apatite. Weathering and dissolution of minerals introduce calcium into soil water, groundwater and surface water.
Uses and important calcium compounds
Calcium metal is used in selected metallurgical and reducing applications. Calcium compounds have much wider uses in construction, chemical processing, agriculture and water treatment. Their formulas and hydration states matter:
- Calcium carbonate (CaCO3): a major constituent of limestone and chalk, used in construction and as a source of lime. Suitable calcite media can also be used to condition low-mineral water.
- Calcium oxide (CaO): quicklime, produced by heating limestone. Its reaction with water releases heat and forms calcium hydroxide.
- Calcium hydroxide [Ca(OH)2]: hydrated or slaked lime, used for pH adjustment and chemical softening. It is a hydroxide, not an oxide.
- Calcium sulfate: gypsum is the dihydrate, CaSO4·2H2O; anhydrous calcium sulfate is CaSO4. Plaster of Paris is principally the hemihydrate, CaSO4·½H2O.
Health effects and safe handling
Calcium as a nutrient
Calcium supports bones and teeth, muscle function, nerve signalling and blood clotting. Vitamin D helps its absorption. Sources include dairy products, fortified foods and selected vegetables. Drinking water can also contribute, depending on its mineral composition.
US NIH dietary reference information gives recommended adult intakes of 1,000–1,200 mg/day, depending on age and sex. Requirements and upper intake limits vary; these figures are nutritional reference values, not drinking-water limits. Calcium and magnesium have separate intake recommendations.
Osteoporosis is influenced by age, hormones, physical activity, nutrition and other factors. Calcium alone does not guarantee prevention. High supplemental intake may cause adverse effects and can interact with medicines. People considering supplements, particularly those with kidney disease or a history of kidney stones, should obtain individual healthcare advice.
Metal, powder and lime
Calcium metal and reactive calcium powder require specific safety controls. Contact with water can release flammable hydrogen and form alkaline calcium hydroxide. Pyrophoric powder can ignite in air and cause severe burns or eye damage. These hazards differ from those of dissolved calcium in drinking water.
Quicklime reacts strongly with water and releases heat; hydrated lime is alkaline and can harm eyes and skin. Handling and emergency procedures should follow the specific product's safety data sheet and the task's exposure assessment. Burning calcium powder needs specialist response; water and carbon dioxide are unsuitable extinguishing agents.
Calcium in the environment and water
Related page: Calcium in water (Ca + H₂O) — occurrence, reactions, water quality and treatment.
Calcium is an essential nutrient and a widespread constituent of rocks, soils and natural waters. The effects of a calcium-containing release depend on the compound, dose and receiving environment. Lime discharges, for example, can alter pH and create solids; the behaviour of one calcium compound does not define every form of the element.
Hardness and scale formation
Calcium and magnesium are the principal hardness ions. 1 mg/L of calcium corresponds to approximately 2.50 mg/L as CaCO3 when expressing calcium hardness. This is an equivalent reporting basis; it does not mean the water contains that mass of solid calcium carbonate.
Dissolved calcium does not automatically form scale. Calcium-carbonate precipitation depends on pH, alkalinity, temperature, concentration and gas exchange, among other conditions. Other calcium salts can also limit industrial operation. Assess the complete water chemistry rather than using calcium concentration alone.
WHO does not propose a health-based guideline value for hardness at concentrations normally found in drinking water. Hardness treatment is often driven by operational and consumer-acceptability considerations. This does not establish a single ideal mineral composition for every supply.
Removing calcium
- Ion-exchange softening: commonly replaces calcium and magnesium with sodium. It reduces hardness without necessarily reducing total dissolved salts.
- Precipitative softening: controlled lime or lime–soda treatment converts selected dissolved hardness constituents into solids, followed by separation and pH conditioning.
- Membrane treatment: nanofiltration and reverse osmosis can reduce calcium. Selection and removal depend on membrane properties, feed chemistry and operating conditions.
Ordinary sediment filtration does not remove dissolved calcium ions. Treatment produces residuals such as regeneration wastewater, sludge or membrane concentrate, which need appropriate management. Read our water softener FAQs for more detail.
Adding calcium after purification
Low-mineral desalinated water may need conditioning before distribution. Calcite contactors or controlled chemical dosing can add calcium and adjust alkalinity and pH. Remineralisation should meet the intended water-quality objectives; it does not, by itself, establish microbiological safety or guarantee corrosion control in every pipe material.
Read more about calcium in water and dietary reference information.
Calcium page last reviewed: by Lenntech.
Back to the periodic table of elements.
Sources and further reading
- RSC: calcium properties, applications and history.
- CIAAW: calcium atomic weight and natural isotopic composition.
- NIH: calcium nutrition, supplements and health.
- ILO/WHO: International Chemical Safety Card 1192 — pyrophoric calcium powder.
- NIOSH: calcium sulfate, gypsum and plaster of Paris.
- WHO: hardness in drinking water (PDF).
- EPA: ion exchange, nanofiltration and reverse osmosis.
- EPA: precipitative softening and residuals management (PDF).
- Texas Water Development Board: calcite contactors for remineralisation (PDF).