|
Name: Calcium |
Boiling Point: 1757°K, 1484°C, 2703°F Melting Point: 1112°K, 839°C, 1542°F 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 @ 300°K Specific Heat: 0.632 J/g°K Atomic Radius: 2.23Å Ionic Radius: 0.99Å Electronegativity: 1 (Pauling); 1.04 (Allrod Rochow) Vapor Pressure: 254 Pa @ 839°C |
| 4 Be 9.012 |
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. |
| 12 Mg 24.30 |
|
| 20 Ca 40.07 |
|
| 38 Sr 87.62 |
|
| 56 Ba 137.3 |
|
| 88 Ra 226.0 |
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.
| 1s2 | ||
| 2s2 | 2p6 | |
| 3s2 | 3p6 | |
| 4s2 |
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.
| 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 |
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.
| Calcium Data |
Atomic Radius (Å): 2.23Å Electrochemical Equivalents: 0.7477 g/amp-hr Atomic Mass Average: 40.078 |
(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.