Calcium is an alkaline earth metal and has been known since prehistoric times. It was discovered in 1787 by Antoine Lavoisier and was isolated in pure form in 1808. Calcium is important nutrient for the human body.
Discovery and History
The history of calcium dates to 7000 BC, when calcium in the form of lime was used as plasters for making statues and as building material. The name calcium is derived from the word “calx” that is Latin for lime. Another form of calcium, calcium sulphate that is termed as gypsum have been found to be used in the construction of Great Egyptian Pyramid of Giza. Later in 1787, lime was considered as an oxide of a novel chemical element by Antoine Lavoisier. Calcium was first isolated by Sir Humphry Davy in 1808 in England. He electrolyzed a mixture of lime and mercuric oxide [1]. A few other scientists, Magnus Pontin and Jöns Jacob Berzelius also produce a calcium amalgam after performing electrolysis on a mixture of lime and mercury oxide.
Calcium
| Periodic Table Classification | Group 2 Period 4 |
|---|---|
| State at 20C | Solid |
| Color | Silvery-gray |
| Electron Configuration | [Ar] 4s2 |
| Electron Number | 20 |
| Proton Number | 20 |
| Electron Shell | 2, 8, 8, 2 |
| Density | 1.55 g.cm-3 at 20°C |
| Atomic number | 20 |
| Atomic Mass | 40.08 g.mol -1 |
| Electronegativity according to Pauling | 1.00 |
Occurrence
Calcium is very reactive and does not occur in free form. It occur in earth’s crust in the forms of carbonate, sulfate, fluoride, silicate and borate. The calcium carbonate occurs in marble, chalk, limestone and calcite. Calcium sulfate (CaSO4) occurs in anhydrite and gypsum, calcium fluoride in fluorspar or fluorite (CaF2) and calcium phosphate occurs in apatite. Calcium also occurs in numerous silicates and alumino silicates. Almost all natural waters, including seawater, contain either or both calcium carbonate and calcium sulfate. Many organisms concentrate calcium compounds in their shells or skeletons. For example calcium carbonate is formed in the shells of oysters and in the skeletons of coral. [2]
Physical characteristics
Its physical and chemical properties are most similar to strontium and barium. It is the fifth most abundant element in Earth’s crust and the third most abundant metal, after iron and aluminum. The atomic number of calcium is 20, and atomic weight is 40.078. The density of calcium is 1.55 grams per cubic centimeter. Its melting point is 842 °C and boiling point is 1494 °C. Calcium is harder than lead but can be cut with a knife with effort. Calcium is a poorer conductor of electricity than copper or aluminum (by volume), but it is a better conductor by mass due to its very low density. [3] It reacts with atmospheric oxygen, which makes its unfavorable to be used in the most applications, but its usage is space is being considered such in space. [4]
Chemical characteristics
The chemistry of calcium is similar to heavy alkaline earth metal. The metal reacts slowly with oxygen, water vapor, and nitrogen of the air to form a yellow coating of the oxide, hydroxide, and nitride. It burns in air or pure oxygen to form the oxide and reacts rapidly with warm water to produce hydrogen gas and calcium hydroxide. On heating, calcium reacts with hydrogen, halogens, boron, sulfur, carbon, and phosphorus. Although it compares with sodium as a reducing agent, calcium is more expensive and less reactive than the latter. In many deoxidizing, reducing, and degasifying applications, however, calcium is preferred because of its lower volatility and is used to prepare chromium, thorium, uranium, zirconium, and other metals from their oxides. [5] Calcium metal dissolves directly in liquid ammonia to give a dark blue solution. [6] Due to the large size of the Ca2+ ion, high coordination numbers are common.
Uses and Significance
- Calcium carbonate is taken as an antacid is effective for treating indigestion.
- Giving calcium gluconate intravenously (by IV) can reverse hyperkalemia, a condition in which there is too much potassium in the blood.
- Taking calcium by mouth is effective for treating and preventing hypocalcemia. It is also given intravenously (by IV) for treating very low levels of calcium in the body.
- Taking calcium carbonate or calcium acetate by mouth is effective for controlling high phosphate levels in the blood, that is present in people with kidney failure.
- Taking calcium by mouth is effective for preventing bone loss and treating osteoporosis.
- Calcium is a co-factor for many enzymes which makes it’s a vital component of the biological system.
- Calcium affects the smooth muscle that surrounds blood vessels and cause it to relax. There are various ionic channels in the membrane of living cells that are controlled by level of calcium in the body.
- It is important to note that calcium is not easily absorbed without the presence of vitamin D.
Isotopes
There are six natural isotopes of calcium, including; Ca-40 is most abundant (97 percent of natural abundance), Ca-44 (2 percent of natural abundance); Ca-42 (0.6 percent of natural abundance); Ca-48 is most stable (0.2 percent of natural abundance); Ca-43 (0.1 percent of natural abundance); Ca-46 (0.004 percent of natural abundance). It is the first element to have low density and have six natural isotopes.
References
1. https://education.jlab.org/itselemental/ele020.html
2. http://nautilus.fis.uc.pt/st2.5/scenes-e/elem/e02020.html
3. Ropp, Richard C. (31 December 2012). Encyclopedia of the Alkaline Earth Compounds. pp. 12–5. ISBN 978-0-444-59553-9.
4. Hluchan and Pomerantz, p. 484
5. https://www.britannica.com/science/calcium
6. Greenwood and Earnshaw, pp. 112–3