Name: Calcium
Symbol: Ca
Atomic Number: 20
Atomic Weight: 40.078000
Family: Alkaline Earth Metals
CAS RN: 7440-70-2
Description: silvery, soft metal, tarnishes to grayish white after exposure to air.
State (25C): Solid
Oxidation states: +2

Molar Volume: 26.02 cm3/mole
Valence Electrons: 4s2

Boiling Point:  1757K, 1484C, 2703F
Melting Point:
1112K, 839C, 1542F
Electrons Energy Level: 2, 8, 8, 2
Isotopes: 15 + 5 Stable
Heat of Vaporization: 153.6 kJ/mol
Heat of Fusion: 8.54 kJ/mol
Density: 1.55 g/cm3 @ 300K
Specific Heat: 0.632 J/gK
Atomic Radius: 2.23
Ionic Radius: 0.99
Electronegativity: 1 (Pauling); 1.04 (Allrod Rochow)
Vapor Pressure: 254 Pa @ 839C
Calcium (Latin calx, meaning "limestone") was known as early as the first century when the Ancient romans prepared lime as Calcium Oxide.  It was not actually isolated until 1808 in England when Sir Humphry Dave electrolyzed a mixture of lime and Mercuric Oxide.   Davy was trying to isolate Calcium and when he heard that Berzelius and Pontin prepared Calcium amalgam by electrolyzing lime in Mercury, he tried it himself.  He worked with electrolysis throughout his life and also discovered/isolated Magnesium, Strontium and Barium.

Calcium is a soft silver-gray alkaline earth metal and is the fifth most abundant element in the earth's crust, widely distributed as Limestone, CaCO3, Quicklime, CaO, and Calcium Fluoride, CaF2.

The pure metal and its compounds give a characteristic brick-red color to flames. Calcium compounds are used in the manufacture of iron and steel, cements and plasters, as well as gypsum wall board. It is important biologically in the formation of bones and teeth.

Calcium metal is fairly reactive and combines with water at room temperature of produce Hydrogen gas and Calcium Hydroxide.  It slowly oxidizes in air, becoming encrusted with white CaO and CaCO3.


1s2 2s2p6 3s2p6 4s2


Calcium is a rather soft, gray, metallic element that can be extracted by electrolysis from Calcium Fluoride.  It burns with a yellow-red flame and forms a white Nitride coating when exposed to air.   It reacts with water, displacing a Hydrogen atom from the structure, then forming Calcium Hydroxide, Ca(OH)2.

2s2 2p6
3s2 3p6

Calcium is essential in muscle contraction, oocyte activation, bones and tooth structure, blood clotting, nerve impulse transmission, regulating heartbeat, and fluid balance within cells.  In the US, between about 50% and 75% of adults do not get sufficient Calcium in their diet.  Adults need between 1,000 and 1,300 mg of calcium in their daily diet.

Its electron configuration is 2 electrons in the K shell (principal quantum number 1), 8 in the L shell (principal quantum number 2), 8 in the M shell (principal quantum number 3), and 2 in the N shell (principal quantum number 4). The outer shell is the valence shell, with 2 electrons in the lone 4s orbital, the 3d orbitals being empty.


Calcium is not naturally found in its elemental state. Calcium occurs most commonly in sedimentary rocks in the minerals Calcite, Dolomite and Gypsum.  It occurs in igneous and metamorphic rocks chiefly in the silicate minerals: Plagioclase, Amphiboles, Pyroxenes and Garnets.



Calcium, combined with Phosphate to form Hydroxylapatite, is the mineral portion of human and animal bones and teeth.  The mineral portion of some corals can also be transformed into hydroxylapatite.

Calcium Oxide (lime) is used in many chemical refinery processes and is made by heating and carefully adding water to limestone.  When lime is mixed with sand, it hardens into a mortar and is turned into plaster by Carbon Dioxide uptake.  Mixed with Mixed with other compounds, lime forms an important part of Portland Cement.

When water percolates through limestone or other soluble carbonate rocks, it partially dissolves part of the rock and causes cave formation and characteristic stalactites and stalagmites and also forms hard water.  Other important Calcium compounds are Nitrate, Sulfide, Chloride, Carbide, Cyanamide, and Hypochlorite.

Calcium Nitrate, Ca(NO3)2
Quicklime, Calcium Oxide, CaO
Slaked Lime, Calcium Hydroxide,  Ca(OH)2
Gypsum, Calcium Sulfate, CaSO4


Calcium has five stable isotopes (40Ca and 42Ca through 44Ca), plus one more isotope, 48Ca, that has such a long half-life that for all practical purposes they can be considered stable.  It also has a cosmogenic isotope, radioactive 41Ca, which has a half-life of 103,000 years.  Unlike cosmogenic isotopes that are produced in the atmosphere, 41Ca is produced by neutron activation of 40Ca.  Most of its production is in the upper meter or so of the soil column where the cosmogenic neutron flux is still sufficiently strong.  41Ca has received much attention in stellar studies because it decays to 41K, a critical indicator of solar-system anomalies.  

The most abundant isotope, 40Ca, has a nucleus of 20 protons and 20 neutrons.  97% of naturally occurring Calcium is in the form of 40Ca.   40Ca is one of the daughter products of 40K decay, along with 40Ar. While K-Ar dating has been used extensively in the geological sciences, the prevalence of 40Ca in nature has impeded its use in dating.   Techniques using mass spectometry and a double spike isotope dilution have been used for K-Ca age dating.

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Isotope Atomic Mass Half-Life
Ca34 34.014  
Ca35 35.0048 50 ms
Ca36 35.9931 102 ms
Ca37 36.9859 181.1 ms
Ca38 37.9763 440 ms
Ca39 38.9707 859.6 ms
Ca40 39.9626 Stable
Ca41 40.9623 103,000 years
Ca42 41.9586 Stable
Ca43 42.9588 Stable
Ca44 43.9555 Stable
Ca45 44.9562 162.61 days
Ca46 45.9537 Stable
Ca47 46.9546 4.536 days
Ca48 47.9525 >6E 18 years
Ca49 48.9557 8.718 minutes
Ca50 49.9575 13.9 seconds
Ca51 50.9615 10 seconds
Ca52 51.965 4.6 seconds
Ca53 52.97 90 ms
Ca54 53.975  
Ca55 54.981  
Ca56 55.986 10 ms
Ca57 56.99  

Biological Role

Calcium is an important component of a healthy diet.  A deficit can affect bone and tooth formation, while overretention can cause kidney stones.  Vitamin D is needed to absorb calcium.  Dairy products, such as milk and cheese, are a well-known source of calcium. However, some individuals are allergic to dairy products and even more people, particularly those of non Indo-European descent, are lactose-intolerant, leaving them unable to consume dairy products. Fortunately, many other good sources of calcium exist.  These include: seaweeds such as kelp, wakame, hijiki, nuts and seeds (such as almonds and sesame): blackstrap molasses, beans, oranges, amaranth, collard greens, okra, rutabaga, broccoli, dandelion leaves, kale, and fortified products such as orange juice and soy milk.  The Calcium content of most foods can be found in the USDA National Nutrient Database.

Calcium is essential for the normal growth and maintenance of bones and teeth, and calcium requirements must be met throughout life.  Long-term Calcium deficiency can lead to osteoporosis, in which the bone deteriorates and there is an increased risk of fractures. Calcium has also been found to assist in the production of lymphatic fluids.

Recommended Adequate Intake by the IOM for Calcium:

Age Calcium (mg/day)
0 to 6 months 210
7 - 12 months 270
1 to 3 years 500
4 to 8 years 800
9 to 18 years 1300
19 to 50 years 1000
51+ years 1200

Calcium supplements are used to prevent and to treat calcium deficiencies.  There are conflicting recommendations about when to take calcium supplements.  However, most experts agree that no more than 500 mg should be taken at a time because the percent of calcium absorbed decreases as the amount of calcium in the supplement increases.   It is recommended to spread doses throughout the day, with the last dose near bedtime.  Recommended daily Calcium intake varies from 1000 to 1500 mg, depending upon the stage of life.

In July 2006, a report citing research from Fred Hutchinson Cancer Research Center in Seattle, Washington claimed that women in their 50s gained 5 pounds less in a period of 10 years by taking more than 500 mg of Calcium supplements than those who did not.   However, the doctor in charge of the study, Dr. Alejandro J. Gonzalez also noted it would be stretching it to suggest Calcium supplements as a weight-limiting aid.

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Calcium Data


Atomic Structure

Atomic Radius (): 2.23
Atomic Volume cm3/mol : 29.9cm3/mol
Covalent Radius: 1.74
Crystal Structure: Cubic face centered
Ionic Radius: 0.99

Chemical Properties

Electrochemical Equivalents: 0.7477 g/amp-hr
Electron Work Function: 2.87eV
Electronegativity: 1 (Pauling); 1.04 (Allrod Rochow)
Heat of Fusion: 8.54 kJ/mol
Incompatibilities: Water, oxidizers, acids, air, chlorine, chlorine tri-fluoride, fluorine, oxygen, silicon, sulfur
First Ionization Potential: 6.113
Second Ionization Potential: 11.871
Third Ionization Potential: 50.908
Valence Electron Potential(-eV): 29
Ionization Energy (eV): 6.113 eV

Physical Properties

Atomic Mass Average: 40.078
Boiling Point: 1757K, 1484C, 2703F
Melting Point: 1112K, 839C, 1542F
Heat of Vaporization: 153.6 kJ/mol
Coefficient of Lineal Thermal Expansion/K-1: 22E-6
Electrical Conductivity: 0.298 106/cm
Thermal Conductivity: 2.01 W/cmK
Density: 1.55 g/cm3 @ 300K
Elastic Modulus (Bulk): 17/GPa
Elastic Modulus (Rigidity): 7.4/GPa
Elastic Modulus Youngs: 20/GPa
Enthalpy of Atomization: 184 kJ/mole @ 25C
Enthalpy of Fusion: 8.54 kJ/mole
Enthalpy of Vaporization: 150 kJ/mole
Hardness Scale (Brinell): 167 MN m-2
Hardness Scale (Mohs): 1.75
Hardness Scale (Vickers): unknown
Flammability Class: Flammable solid
Molar Volume: 26.02 cm3/mole
Optical Reflectivity: unknown
Optical Refractive Index: unknown
Relative Gas Density (Air=1): unknown
Specific Heat: 0.632 J/gK
Vapor Pressure: 254 Pa @ 839C
Estimated Crustal Abundance: 4.15104 milligrams per kilogram
Estimated Oceanic Abundance:
4.12102 milligrams per liter

(L. calx, lime) Though lime was prepared by the Romans in the first century under the name calx, the metal was not discovered until 1808. After learning that Berzelius and Pontin prepared calcium amalgam by electrolyzing lime in mercury, Davy was able to isolate the impure metal. Calcium is a metallic element, fifth in abundance in the earth's crust, of which if 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. The metal has a silvery color, is rather hard, and is prepared by electrolysis of the fused chloride to which calcium fluoride is added 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, forming largely the nitride. 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. Its natural and prepared compounds are widely used. Quicklime (CaO), made by heating limestone and changed into slaked lime by the careful addition of water, is the great cheap base of chemical refinery with countless uses. Mixed with sand it hardens as mortar and plaster by taking up carbon dioxide from the air. Calcium from limestone is an important element in Portland cement. The solubility of the carbonate in water containing carbon dioxide causes the formation of caves with stalactites and stalagmites and is responsible for hardness in water. Other important compounds are the carbide, chloride, cyanamide, hypochlorite, nitrate, and sulfide.

Source: CRC Handbook of Chemistry and Physics, 1913-1995. David R. Lide, Editor in Chief. Author: C.R. Hammond

Although calcium is the fifth most abundant element in the Earth's crust, it is never found free in nature since it easily forms compounds by reacting with oxygen and water. Metallic calcium was first isolated by Sir Humphry Davy in 1808 through the electrolysis of a mixture of lime (CaO) and mercuric oxide (HgO). Today, metallic calcium is obtained by displacing calcium atoms in lime with atoms of aluminum in hot, low-pressure containers. About 4.2% of the Earth's crust is composed of calcium.

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.