Calcium

Typical wastewaters

• lime softening plant residuals calcium carbonate and magnesium hydroxide sludge, plus the lime added; some plants regenerate lime from it
• ion exchange softener regenerant Ca²⁺ and Mg²⁺ in spent brine
• reverse osmosis concentrate and cooling water Ca²⁺ concentrated toward calcium carbonate and gypsum saturation, held off with antiscalant or acid
• lime treated acid effluent and mine water Ca²⁺ with sulfate, capped by gypsum near 2 g/L, so lime treatment leaves 1.5 to 2 g/L sulfate

1 · Identity

Symbol, number

Ca, 20

Oxidation states in water

+2 only, as the hydrated Ca²⁺ ion and its ion pairs with bicarbonate, carbonate and sulfate; the solids CaCO₃ (calcite, aragonite), CaSO₄.2H₂O (gypsum), Ca(OH)₂ (portlandite, slaked lime), Ca₅(PO₄)₃OH (hydroxyapatite) and CaF₂ (fluorite) set its solubility.

Note

The element entry covers limestone, lime burning and the narrative on solubility. This chapter is the carbonate system in practice: when calcium carbonate precipitates, how a saturation index is read, what lime does at pH 10 and 11, and what a desalination plant has to put back.

2 · Occurrence in water

Natural sources

Dissolution of limestone, chalk, dolomite and gypsum by carbon dioxide charged rainwater; calcium concentrations up to and exceeding 100 mg/L are common in natural sources, particularly groundwater, and calcium hardness usually predominates over magnesium (WHO hardness document).

Anthropogenic sources

Lime and calcium hypochlorite dosing; lime neutralisation of acid effluent and mine water; gypsum from flue gas desulfurisation and phosphoric acid (phosphogypsum) in the ledger's chemical chapter; road salt (calcium chloride); concrete leaching.

3 · Speciation

Calcium is Ca²⁺ at every pH, and its story is the solids it can form. In natural water it sits close to calcium carbonate saturation: warming the water, stripping carbon dioxide or raising pH pushes it over and calcite deposits on heaters, pipes and membranes; cooling, adding acid or carbon dioxide dissolves it. The Langelier saturation index (pH minus the pH at which the water would be saturated with CaCO₃) reads which way the water is tending, positive for scaling and negative for dissolving, and Standard Methods 2330 pairs it with the calcium carbonate precipitation potential, the mass that would actually precipitate. Sulfate rich water is capped instead by gypsum near 2 g/L. Lime softening deliberately drives the water far past saturation at pH 10 to 11.

Solubility

Calcite: log Ks₀ about minus 8.48 at 25 C (Stumm and Morgan chapter 7, from the chapter, not re-read); solubility falls with rising temperature, which is why calcium carbonate scales hot surfaces. Gypsum about 2 g/L. Calcium hydroxide about 1.7 g/L, enough to make lime water a strong base but too little to dose lime as a clear solution at scale, hence lime slurry. Fluorite and hydroxyapatite far less soluble, the basis of fluoride and phosphate removal with lime.

Hydrolysis

Negligible; Ca²⁺ is a hard, weakly hydrolysing cation. Quicklime hydrates exothermically to slaked lime, a strong base.

Complexation

Ion pairs with bicarbonate, carbonate and sulfate; strong complexes with EDTA and with phosphonate antiscalants (the basis of EDTA titration for hardness and of threshold scale inhibition); calcium binds humic substances and bridges them into floc.

Precipitates

CaCO₃ (calcite, aragonite, vaterite), CaSO₄.2H₂O gypsum, Ca(OH)₂ in lime slurry, Ca₅(PO₄)₃OH hydroxyapatite, CaF₂ fluorite, calcium arsenate, calcium silicate and ettringite in lime treated water.

4 · Role in treatment

5 · Removal and control

6 · Analytics

Sampling pitfalls

Calcium is stable in a sample, but the saturation index is not: pH and temperature must be measured in the field, and a supersaturated sample will precipitate calcite in the bottle so that calcium, alkalinity and pH all fall. Acidify samples for total calcium; do not acidify the aliquot used for alkalinity.

7 · Regulatory limits

Limits change, and many are set locally. Treat these as the published values to start from, not as your compliance target: check the standard in force at your site and the numbers written into your own permit.

8 · Health and environmental effects

Toxicity

Non toxic and essential (about 1 kg in an adult, the element entry); drinking water can be a useful contributor to calcium and magnesium intake, and WHO advises adding calcium and magnesium back to demineralised supplies to the levels the population had before (hardness background document).

Bioaccumulation

Not applicable.

Ecotoxicity

Calcium is a nutrient; its practical ecotoxicological role is that hardness reduces the toxicity of metals, which is why the US EPA freshwater criteria for chromium(III), nickel and other metals are expressed as functions of hardness.

Flags

• The calcite solubility product, gypsum and lime solubilities, the lime softening equations and the practical softening limits are cited to textbook chapters (Stumm and Morgan 7, Snoeyink and Jenkins 6, MWH 20) from memory, not re-read.
• The fluoride residual after lime treatment is from Metcalf and Eddy chapter 6 by memory and is not quoted as a number in a source read.
• The seawater figure is a single PubChem abundance figure.
• The hardness classes are one classification quoted by WHO; other bodies use different bands.
• No calcium discharge limit exists in any document read; pH limits do the work.

Gaps

• No source read gives calcium in surface water, municipal or industrial wastewater as numbers.
• No Langelier index worked example or calcium carbonate precipitation potential figure was read; Standard Methods 2330 is cited by title.
• Lime doses, softening sludge quantities and recarbonation doses were not read; MWH chapter 20 has them.
• Detection limits for the EDTA titrations were not read.
• Remineralisation targets for desalinated water (calcium, alkalinity, index) were not read; WHO states the principle only.
• Other GCC standards were not read.
• The Langelier index itself is an index, not a reaction: it is written here as the calcite solubility and the carbonate dissociation that fix pHs; no worked example or activity correction was read.
• The hydroxyapatite, gypsum neutralisation and EDTA titration stoichiometries are cited to Metcalf and Eddy chapter 6 and Standard Methods 2340 by chapter and method number, not re-read.

Sources

Identity

Name and symbol

Calcium, Ca

Atomic number

20 protons

Position

group 2 · period 4 · s-block · alkaline earth metal

CAS number

7440-70-2

Atomic structure

Atomic mass

40.078 u

Electron configuration

1s² 2s² 2p⁶ 3s² 3p⁶ 4s²
[Ar] 4s²

Electrons per shell

2, 8, 8, 2

Valence electrons

2 outer shell

Physical properties

State at room temperature

Solid

Melting point

1,115 K (841.85 °C)

Boiling point

1,757 K (1,483.85 °C)

Density

1.54 g/cm3

Appearance

dull gray, silver; with a pale yellow tint

Thermal conductivity

201 W/(m·K)

Electrical resistivity

33.6 nΩ·m at 20 °C

Electrical conductivity

29.76 MS/m

Crystal structure

face-centered cubic

Molar heat capacity

25.929 J/(mol·K)

Chemical properties

Oxidation states

+2

Electronegativity

1 (Pauling Scale)

Ionisation energy

6.113 eV
1st 589.8, 2nd 1,145.4, 3rd 4,912.4 kJ/mol

Electron affinity

0 eV

Atomic radius

empirical 176, covalent 176, van der Waals 231 pm

Ionic radius

Ca²⁺ 100 pm

Reactivity

A typical heavy alkaline earth metal ([Ar] 4s2), reactive enough that it is never found free; it tarnishes rapidly in air, reacts with water faster than magnesium and more slowly than strontium, and is a strong reducing agent used to win other metals.

with water

Reacts steadily with cold water to calcium hydroxide and hydrogen, faster than magnesium: $\ce{Ca + 2 H2O -> Ca(OH)2 + H2}$

with oxygen, air

Forms a dark oxide-nitride coating in air; heated it is reluctant to start but then burns with an intense yellow-red flame to calcium oxide and nitride, and the fine powder burns spontaneously: $\ce{2 Ca + O2 -> 2 CaO}$

with acids

Reacts readily with dilute hydrochloric acid to the chloride and hydrogen; with sulfuric acid a coating of sparingly soluble calcium sulfate slows the reaction, and very dilute nitric acid gives hydrogen contaminated with nitrogen oxides: $\ce{Ca + 2 HCl -> CaCl2 + H2}$

with halogens

Reacts with the halogens to the ionic dihalides such as calcium chloride: $\ce{Ca + Cl2 -> CaCl2}$

Typical compounds

• CaCO₃ calcium carbonate limestone, chalk, marble; building stone and cement
• CaO calcium oxide quicklime, from roasting limestone
• Ca(OH)₂ calcium hydroxide slaked lime; mortar and plaster
• CaSO₄.2H₂O calcium sulfate dihydrate gypsum; plaster
• CaC₂ calcium carbide from lime and carbon; gives acetylene with water
• CaF₂ calcium fluoride fluorite; lowers the melting point in calcium electrolysis

Occurrence, production and use

Crustal abundance

4.15×104 milligrams per kilogram

Oceanic abundance

4.12×102 milligrams per liter

Occurrence and sources

Calcium, a metallic element, is fifth in abundance in the earth's crust, of which it forms more than 3%. It is an essential constituent of leaves, bones, teeth, and shells. Never found in nature uncombined, it occurs abundantly as limestone, gypsum, and fluorite. Apatite is the fluorophosphate or chlorophosphate of calcium.

• limestone and chalk (CaCO3), gypsum (CaSO4), fluorite (CaF2), apatite (calcium fluoro and chlorophosphate) sedimentary rock worldwide; limestone and dolomite suitable for lime manufacture are very large resources in every producing country
• dissolved calcium bicarbonate hard groundwater and surface water; it precipitates as stalactites and scale when carbon dioxide is lost
• calcium phosphate the main component of bone, teeth and shells

Extraction, production

Lime burning (calcination of limestone)

$\ce{CaCO3 -> CaO + CO2}$

Limestone heated in kilns gives off carbon dioxide and leaves quicklime; balanced from the reactants and products stated by the source. World lime production about 420 million tonnes in 2024 (estimate), 73 quicklime plants and 10 hydrating plants in the United States (USGS). Large quantities of lime are regenerated by paper mills and some water treatment plants regenerate lime from softening sludge.

Slaking (hydration of quicklime)

$\ce{CaO + H2O -> Ca(OH)2}$

Quicklime reacts vigorously with a controlled amount of water in a hydrator to give hydrated (slaked) lime, a dry calcium hydroxide powder used in chemical, construction and environmental applications; balanced from the reactants and product stated by the sources.

Calcium metal by aluminothermic reduction of lime

The metal is prepared commercially by heating lime with aluminium under vacuum, or by electrolysis of fused calcium chloride with calcium fluoride added to lower the melting point. The sources name the reducing agent but not the aluminium bearing product, so no equation is written; demand for the metal is small (reducing agent for thorium, uranium and zirconium, deoxidiser and desulphuriser for alloys).

Uses

Due to its high reactivity with common materials, there is very little demand for metallic calcium. It is used in some chemical processes to refine thorium, uranium and zirconium. Calcium is also used to remove oxygen, sulfur and carbon from certain alloys. Calcium can be alloyed with aluminum, beryllium, copper, lead and magnesium. Calcium is also used in vacuum tubes as a getter, a material that combines with and removes trace gases from vacuum tubes.

Calcium carbonate (CaCO3) is one of the common compounds of calcium. It is heated to form quicklime (CaO) which is then added to water (H2O). This forms another material known as slaked lime (Ca(OH)2) which is an inexpensive base material used throughout the chemical industry. Chalk, marble and limestone are all forms of calcium carbonate. Calcium carbonate is used to make white paint, cleaning powder, toothpaste and stomach antacids, among other things. Other common compounds of calcium include: calcium sulfate (CaSO4), also known as gypsum, which is used to make dry wall and plaster of Paris, calcium nitrate (Ca(NO3)2), a naturally occurring fertilizer and calcium phosphate (Ca3(PO4)2), the main material found in bones and teeth.

The metal is used as a reducing agent in preparing other metals such as thorium, uranium, zirconium, etc., and is used as a deoxidizer, desulfurizer, or decarburizer for various ferrous and nonferrous alloys. It is also used as an alloying agent for aluminum, beryllium, copper, lead, and magnesium alloys, and serves as a "getter" for residual gases in vacuum tubes, etc.

• Steel: lime as the flux that removes impurities from molten iron and steel; steelmaking is the largest lime market in the United States steelmaking was the leading United States lime market in 2024 (USGS, ranking only)
• Food and beverage: lime and carbon dioxide for beet juice purification (carbonatation), leaving carbonatation lime cake as a residue; sugar companies burn lime for their own use; calcium salts as additives (calcium carbonate E170, propionate, lactate, phosphates)
• Chemicals: calcium ammonium nitrate fertiliser, made by mixing ammonium nitrate solution with dolomite, limestone or calcium carbonate; phosphate rock (apatite) as the calcium phosphate feed of wet process phosphoric acid and superphosphates, with phosphogypsum and calcium fluoride sludge as co product and waste stream
• Glass, pulp and paper: lime for glass, pulp and paper manufacture and for precipitated calcium carbonate; paper mills regenerate large quantities of lime
• Mining: lime, quicklime, hydrated lime and limestone for neutralisation and metal precipitation in mine water treatment, including the high density sludge lime process for base metal mine water
• Water treatment: slaked lime to reduce acidity, lime softening of hard water, flue gas desulphurisation
• Construction: cement, lime mortar and lime plaster; gypsum plaster and plasterboard

Safety, toxicity

not in sources

Discovery and name

Discovered by

Humphry Davy

Discovered

1808

First isolated

not in sources

Named by

not in sources

Origin of the name

from the Latin word for lime, calx

The metal has a silvery color, is rather hard, and is prepared by electrolysis of fused chloride and calcium fluoride (to lower the melting point).

Chemically it is one of the alkaline earth elements; it readily forms a white coating of nitride in air, reacts with water, burns with a yellow-red flame.