Physics & Chemistry

Calcium

Calcium is the chemical element with atomic number 20; it has an atomic mass of 40.078 atomic mass units (amu). The chemical symbol for calcium is Ca. Calcium is a soft gray alkaline earth metal, and is the fifth most abundant element by mass in the Earth's crust; moreover, it is the fifth most abundant dissolved ion in seawater both in terms of number of atoms and mass, after sodium, chloride, magnesium, and sulfate.[1]

Calcium an is essential nutrient for almost all living organisms, with vital roles in cellular metabolism, especially with regard to movement of the calcium ion Ca++ into and out of the cytoplasm functions as a signal for many cellular processes. As a chief component needed in mineralization of bones and shells, calcium is the most abundant metal by mass in a large number of faunal species, especially vertebrates, testudines and mollusca.

Previous Element: Potassium

Next Element: Scandium
20

Ca

40.08
Physical Properties
Color Silvery-white
Phase at Room Temp. solid
Density (g/cm3) 1.55
Hardness (Mohs) 1.5
Melting Point (K) 1112.2
Boiling Point (K) 1767
Heat of Fusion (kJ/mol) 9.3
Heat of Vaporization (kJ/mol) 151
Heat of Atomization (kJ/mol) 178
Thermal Conductivity (J/m sec K) 201
Electrical Conductivity (1/mohm cm) 255.754
Source Calcite (oxide)
Atomic Properties
Electron Configuration [Ar]4s2
Number of Isotopes 22 (6 natural)
Electron Affinity (kJ/mol) ---
First Ionization Energy (kJ/mol) 589.8
Second Ionization Energy (kJ/mol) 1145.4
Third Ionization Energy (kJ/mol) 4911.8
Electronegativity 1
Polarizability (Å3) 22.8
Atomic Weight 40.08
Atomic Volume (cm3/mol) 25.9
Ionic Radius2- (pm) ---
Ionic Radius1- (pm) ---
Atomic Radius (pm) 197
Ionic Radius1+ (pm) ---
Ionic Radius2+ (pm) 114
Ionic Radius3+ (pm) ---
Common Oxidation Numbers +2
Other Oxid. Numbers ---
Abundance
In Earth's Crust (mg/kg) 4.15×104
In Earth's Ocean (mg/L) 4.12×102
In Human Body (%) 1.43%
Regulatory / Health
CAS Number 7440-70-2
OSHA Permissible Exposure Limit (PEL) No limits
OSHA PEL Vacated 1989 No limits
NIOSH Recommended Exposure Limit (REL) No limits
Sources:
Mineral Information Institute
Jefferson Accelerator Laboratory
EnvironmentalChemistry.com
 

Natural occurrence

caption Karst limestone formation at the edge of the Anjajavy Forest along the Indian Ocean, Madagascar. Source: C.Michael Hogan

Elemental calcium is not found in the natural environment on Earth. As a solid calcium occurs chiefly in sedimentary rock, especially in limestone and in the specific minerals calcite, dolomite and gypsum. Some of the most dramatic occurrences are in limestones, where stalactites and stalagmites are formed in underground caverns under favorable conditions of pH, temperature and ionic concentration of calcium. Calcium is additionally found in certain igneous and metamorphic rocks: principally in the silicates:[2] plagioclase, amphiboles, pyroxenes, garnet and some serpentine rocks.

Biogeochemical cycles

Calcium is a key link between tectonics, climate and the carbon cycle. Since mountain block uplift exposes calcium bearing strata to chemical weathering, significant release occurs of Ca++ into surface waters. Much of this dissolved calcium ultimately is transported to the oceans, where it reacts with dissolved carbon ioxide to produce limestone. Certain quantities of this limestone precipitates to the sea floor where it is incorporated into new rock strata. Dissolved carbon dioxide, together with carbonate and bicarbonate ions, are termed dissolved inorganic carbon.

The actual reaction is more complex, involving bicarbonate ion (HCO3-), which is produced when carbon dioxide ionizes in pH levels found in oceans.

Ca++  + 2 HCO3- →  CaCO3 (limestone) +  CO2 +  H2O

Note that at pH levels of the oceans, most of the CO2 produced in this reaction is immediately converted back into HCO3-. The reaction thus results in a net transport of carbon dioxide from the ocean or atmosphere into the lithosphere.[3] This cycling is a significant sink for carbon dioxide, and provides an important negative feedback loop for carbon dioxide additions to the atmosphere, and thus a key mitigation measure to potential global warming.

Isotopes

caption Creation of the newest element ununseptium by collision of 48Ca and 249Bk.

There are four stable calcium isotopes (40Ca,  42Ca,  43Ca, 44Ca); in addition,[4] 46Ca and 48Ca are radioactive, but possess extremely long half-lives. Furthermore radioactive 41Ca also is produced in the uppermost meter of the Earth crust by neutron activation of 40Ca; moreover, 41Ca has the long half-life of 103,000 years.  41Ca is received much attention in stellar studies because it decays to 41K, a critical indicator of solar-system anomalies.

Ninety seven percent of naturally occurring calcium is in the form of 40Ca, which is one of the daughter products of 40K decay, along with the daughter 40Ar. While potassium-argon dating has been widely employed extensively in the earth sciences field, the  utility of this dating technique is limited due to the minute quantities of argon gas that are required to be measured; in practical terms the technique is useful where long geological time scales are involved (e.g. the order of 100,000 years);[5] moreover, as measurement techniques steadily improve, the utility is rising.

Notable properties

Calcium possesses the lightest density of all the alkali earth metals. It is chemically reactive and is so soft that it can be cut with a knife, although it is somewhat harder than  lead. As a metal it is silvery in appearance; to produce the metallic form, one usually utilizes electrolysis upon a fused salt like calcium chloride. When exposed to the air in its metallic form, a coating is quickly formed as a whitish gray oxide and nitride coating. Unlike magnesium, it is difficult to ignite; however, upon combustion it manifests a brilliant brick-red high intensity luminence. As a metal, calcium reacts with water, rapidly producing hydrogen gas, but exuding little heat. If the metal has been powdered, the reaction with water is more robust, due to the considerably greater surface area.

caption Bright red flame of calcium being combusted

Calcium has a higher electrical resistivity than copper or aluminium. Calcium actually has an electrical conductivity about forty percent that of copper; however, its conductivity under normal atmospheric conditions is usually limited by its chemical reactivity with air.

Ionic solutions of calcium salts are generally colorless. Many calcium salts are not soluble in water. When in solution, the calcium ion to the human taste varies with its anion, being reported variously as moderately salty, sour or mineral like. It has been demonstrated that many fauna can taste calcium, and use this sense to detect the mineral in salt licks.[6]

Example compounds

Calcium phosphate leads to the formation of calcium hydroxylapatite, the mineral portion of human and animal bones and dentin.[7] The mineral portion of some corals can also be transformed into hydroxylapatite.

Calcium hydroxide (slaked lime), employed in numerous chemical refinery processes, is produced from applying heat to limestone at temperatures above 825 °C, following by hydration. When lime is mixed with sand, it hardens into a mortar and is converted into plaster after carbon dioxide uptake. Mixed with other compounds, lime forms an  component of Portland cement. Calcium carbonate, one of the common compounds of calcium, can be heated to form quicklime (CaO). After hydration, quicklime is converted to  slaked lime (Ca(OH)2), which is a cheap commodity utilized widely in the chemical manufacturing industry.[8] Chalk and marble are other alternative forms of calcium carbonate.

Other notable calcium compounds are calcium sulfate, calcium chloride, calcium sulfide, calcium carbide and calcium hypochlorite.

Faunal metabolism

caption Normal vs. osteoporotic bone viewed by electron micrograph. Only with extreme magnification can one see the massive damage from calcium deficiency. Source: U.S. CDC

Calcium is the most abundant mineral not only in the human body but for many other faunal species. A human adult body has about one kilogram of calcium, 99 percent of which is stored in bones as calcium phosphate. Human extracellular fluid holds around 22.5 mmol, almost half of which is within  blood plasma. About 500 mmol of calcium is exchanged between bone and the extracellular fluid within a complete daily cycle.[9] Hypocalcaemia in humans is the condition of  deficient calcium concentration in blood plasma; this deficit circumstance is one in which  less than 90 milligramsl/liter or ionized calcium concentration less than 45 milligrams/liter). This condition is an electrolyte imbalance that can lead to osteoporosis;[10] it is important to note that it is the ionized calcium which is key to metabolic regulation. Conversely, elevated blood calcium levels, known as hypercalcaemia, can lead to adverse health effects including heart arhythmia, depression and peptic ulcers.

Hypocalcaemia is also studied in depth for dairy cows, other domestic animals and fauna in the wild. In the case of dairy cows the condition is linked to inferior milk production and other disease issues.[11]

References

  1. ^ A.G.Dickson and C. Goyet. 1994. Handbook of method for the analysis of the various parameters of the carbon dioxide system in sea water. Publisher:  ORNL/CDIAC
  2. ^ Harvey Blatt, Robert J.Tracy and Brent E.Owens. 2006. Petrology: igneous, sedimentary, and metamorphic. Macmillan. 530 pages
  3. ^ Richard Zeebe. 2006. Marine carbonate chemistry. National Council for Science and the Environment
  4. ^ H.G. Attendorn and Robert Bowen. 1998. Isotopes in the Earth Sciences, Springer. 668 pages
  5. ^ Kevin Greene. 1996. Archaeology, an introduction: the history, principles, and methods. Routledge. 208 pages
  6. ^ M.G.Tordoff. 2001. Calcium: Taste, Intake, and Appetite. Physiological Reviews. vol. 81, issue 4, page 1567
  7. ^ Joon Bu Park and Roderic S. Lakes. 2007. Biomaterials: an introduction, Springer. 561 pages
  8. ^ Daniel D. Walker. 1992. Innovations and uses for lime. ASYM International. 141 pages
  9. ^ W.J.Marshall. 1995. Clinical Chemistry, third edition. Mosby Publishers, London.UK
  10. ^ Maurice Edward Shils and Moshe Shike. 2006. Modern nutrition in health and disease. Lippincott Williams & Wilkins. 2069 pages
  11. ^ U.S. National Research Council, Subcommittee on Dairy Cattle Nutrition. 2001. Nutrient requirements of dairy cattle. National Academies Press. 381 pages
Glossary

Citation

Hogan, C. (2013). Calcium. Retrieved from http://www.eoearth.org/view/article/150845

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