Chemical elements

The periodic table is a tabular arrangement of the chemical elements, ordered by their atomic number (number of protons), electron configurations, and recurring chemical properties. This ordering shows periodic trends, such as elements with similar behaviour in the same column. It also shows four rectangular blocks with some approximately similar chemical properties. In general, within one row (period) the elements are metals on the left, and non-metals on the right.

The rows of the table are called periods; the columns are called groups. Six groups have names as well as numbers: for example, group 17 elements are the halogens; and group 18, the noble gases. The periodic table can be used to derive relationships between the properties of the elements, and predict the properties of new elements yet to be discovered or synthesized. The periodic table provides a useful framework for analyzing chemical behaviour, and is widely used in chemistry and other sciences.

The Russian chemist Dmitri Mendeleev published the first widely recognized periodic table in 1869. He developed his table to illustrate periodic trends in the properties of the then-known elements. Mendeleev also predicted some properties of then-unknown elements that would be expected to fill gaps in this table. Most of his predictions were proved correct when the elements in question were subsequently discovered. Mendeleev’s periodic table has since been expanded and refined with the discovery or synthesis of further new elements and the development of new theoretical models to explain chemical behaviour.

All elements from atomic numbers 1 (hydrogen) to 118 (oganesson) have been discovered or synthesized, with the most recent additions (nihonium, moscovium, tennessine, and oganesson) being confirmed by the IUPAC on December 30, 2015 and officially named on November 28, 2016: they complete the first seven rows of the periodic table.[1][2] The first 94 elements exist naturally, although some are found only in trace amounts and were synthesized in laboratories before being found in nature.[n 1] Elements with atomic numbers from 95 to 118 have only been synthesized in laboratories or nuclear reactors.[3] Synthesis of elements having higher atomic numbers is being pursued. Numerous synthetic radionuclides of naturally occurring elements have also been produced in laboratories.

2.The spiral shell and bones reflect the essential presence of calcium in all living things. Calcium is a silvery-white, soft metal that tarnishes rapidly in air and reacts with water.  

 Calcium metal is used as a reducing agent in preparing other metals such as thorium and uranium. It is also used as an alloying agent for aluminium, beryllium, copper, lead and magnesium alloys.

Calcium compounds are widely used. There are vast deposits of limestone (calcium carbonate) used directly as a building stone and indirectly for cement. When limestone is heated in kilns it gives off carbon dioxide gas leaving behind quicklime (calcium oxide). This reacts vigorously with water to give slaked lime (calcium hydroxide). Slaked lime is used to make cement, as a soil conditioner and in water treatment to reduce acidity, and in the chemicals industry. It is also used in steel making to remove impurities from the molten iron ore. When mixed with sand, slaked lime takes up carbon dioxide from the air and hardens as lime plaster.

Gypsum (calcium sulfate) is used by builders as a plaster and by nurses for setting bones, as ‘plaster of Paris’.     

Biological role

Calcium is essential to all living things, particularly for the growth of healthy teeth and bones. Calcium phosphate is the main component of bone. The average human contains about 1 kilogram of calcium.

Children and pregnant women are encouraged to eat foods rich in calcium, such as milk and dairy products, leafy green vegetables, fish and nuts and seeds.    

Natural abundance

Calcium is the fifth most abundant metal in the Earth’s crust (4.1%). It is not found uncombined in nature, but occurs abundantly as limestone (calcium carbonate), gypsum (calcium sulfate), fluorite (calcium fluoride) and apatite (calcium chloro- or fluoro-phosphate). Hard water contains dissolved calcium bicarbonate. When this filters through the ground and reaches a cave, it precipitates out to form stalactites and stalagmites. Calcium metal is prepared commercially by heating lime with aluminium in a vacuum.

  1. 3. People have used calcium’s compounds for thousands of years – in cement, for example.

Limestone [calcium carbonate] was called calx by the Romans. The Romans heated calx, driving off carbon dioxide to leave calcium oxide. To make cement, all you have to do is mix calcium oxide with water. The Romans built vast amphitheaters and aqueducts using calcium oxide cement to bond stones together.

Despite the long history of calcium’s compounds, the element itself was not discovered until electricity was available for use in experiments.

Calcium was first isolated by Sir Humphry Davy in 1808 in London. In a lecture to the Royal Society in June 1808, Davy described his experiments that year, which produced tiny amounts of metal, at best. He could not find any way to produce more calcium metal until a letter from Jöns Berzelius in Stockholm pointed him in the right direction.

Davy learned that Berzelius and Magnus Pontin had used a battery to decompose calcium oxide at a mercury electrode and they had obtained an amalgam of mercury and calcium. (Berzelius, the great Swedish chemist, exchanged a great deal of information with Davy. Berzelius had earlier learned from Davy that potassium could be dissolved in mercury to form an amalgam. Berzelius had extended the method.) Davy made a paste of slaked lime [calcium oxide, slightly moistened to form calcium hydroxide] and red oxide of mercury [mercury (II) oxide].

He made a depression in the paste and placed about 3.5 grams of mercury metal there to act as an electrode. Platinum was used as the counter electrode. Davy carried out the experiment under naptha (a liquid hydrocarbon under which he had found he could safely store potassium and sodium).

When electricity was passed through the paste, a calcium-mercury amalgam formed at the mercury electrode. Davy removed the mercury by distillation to reveal a new element: calcium.Davy used the same procedure to isolate strontium, barium, and magnesium. He named the metal calcium because of its occurrence in calx.

  1. Atomic data
  •  Atomic number (number of protons in the nucleus): 20
  • Atomic symbol (on the periodic table of the elements): Ca
  • Atomic weight (average mass of the atom): 40.078
  • Density: 1.55 grams per cubic centimeter
  • Phase at room temperature: solid
  • Melting point: 1,548 degrees Fahrenheit (842 degrees Celsius)
  • Boiling point: 2,703 F (1,484 C)
  • Number of isotopes (atoms of the same element with a different number of neutrons): 24; 5 stable
  • Most common isotopes: Ca-40 (97 percent of natural abundance); Ca-44 (2 percent of natural abundance); Ca-42 (0.6 percent of natural abundance); Ca-48 (0.2 percent of natural abundance); Ca-43 (0.1 percent of natural abundance); Ca-46 (0.004 percent of natural abundance.

 

5.

Common oxidation states2
IsotopesIsotope Atomic mass Natural abundance (%) Half life Mode of decay
40Ca39.96396.9415.92 x 1021 yEC-EC
42Ca41.9590.647
43Ca42.9590.135
44Ca43.9552.086
46Ca45.9540.004> 0.4 x 1016 yβ-β-
48Ca47.9530.1874.4 x 1019 yβ-β-
> 7.1 x 1019 yβ-
 

Calcium has five stable isotopes (40Ca, 42Ca, 43Ca, 44Ca and 46Ca), plus one more (48Ca) that has such a long half-life, it can be considered stable for many purposes. The 20% range in relative mass among naturally occurring calcium isotopes is greater than for any element other than hydrogen and helium. Calcium also has a cosmogenic isotope, radioactive 41Ca, which has a half-life of 103,000 years. Unlike cosmogenic isotopes produced in the atmosphere, 41Ca is produced by neutron activation of 40Ca, primarily in the top metre of the soil column, where the cosmogenic neutron flux is sufficiently strong. 41Ca has 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 a daughter product of 40K and 40Ar decay. While K–Ar dating is used extensively in the geological sciences, the prevalence of 40Ca in nature has impeded its use in dating. Techniques using mass spectrometry and a double spike isotope dilution are used for K-Ca age dating.

40Ca has a nucleus of 20 protons and 20 neutrons and is the heaviest stable isotope of any element that has equal numbers of protons and neutrons. In supernova explosions, calcium is formed from the reaction of carbon with various numbers of alpha particles (helium nuclei), until the most common calcium isotope (containing 10 helium nuclei) has been synthesized.

 

6. Interesting Facts about Calcium

  • Calcium is the most abundant of the metallic elements in the human body. The average adult body contains about 1 kg or 2 lb of calcium, 99% of which is in the bones and teeth. Only oxygen, carbon, hydrogen and nitrogen are more abundant in our bodies than calcium.
  • Calcium not only builds the structures that support our bodies, many of us also live in homes built using structural concrete or cement made with lime (calcium oxide). Snails and many shellfish use another calcium compound – calcium carbonate – to build their own homes too – their shells.
  • Modern humans were not the first people to make use of calcium to build things. Egypt’s pyramids were built using limestone blocks. Limestone is crystalline calcium carbonate. In the later pyramids, the blocks were held together with gypsum or lime based mortar. Gypsum is calcium sulfate dihydrate and lime is calcium oxide.
  • Have you ever wanted to be ‘in the limelight?’ Lime is calcium oxide, which produces a brilliant, intense light when burnt in an oxyhydrogen flame. It was used to light the stage in theaters during the 1800s until electricity took over – hence the saying.
  • Cells in animals and plants must communicate with other cells. This is called signaling. Calcium ions are the most important messengers between cells in living things and are absolutely vital for the existence of multicellular life forms.
  • When turtles (typically pet turtles) don’t have enough calcium circulating through their blood, they may get an ailment called Metabolic Bone Disease, commonly known as Soft Shell Syndrome. For a turtle to be healthy, the calcium to phosphorous ratio should be 2:1. When their calcium levels are low, however, the mineral is leached from their bones in the body’s attempt to balance things out. The result is soft bones, weakness and a soft, deformed shell, and it often results in death. The disease can be prevented with a proper diet and adequate sunlight (or other appropriate lighting for reptiles).
  • Stalactites and stalagmites, the icicle-shaped formations found in underground caverns, are formed slowly over time by the build-up of calcite residue. This occurs when water seeps through the cracks in the ceiling of a limestone cave, dissolving and carrying along traces of calcite, the building material of limestone. As the water drips from the ceiling, this calcite residue begins to build up at the site of the drip, eventually resulting in icicle-shaped stalactites hanging from the cave ceiling. This water dripping from the stalactites then forms stalagmites on the ground below.
  • Many nutritionists recommend a calcium-magnesium ratio of 2:1. But although our bodies require more calcium, we are actually more likely to become deficient in magnesium. This is because our bodies tend to store and recycle calcium, while magnesium gets used or excreted and must be replenished on a daily basis.
  • Calcium carbonate is the active ingredient in many antacids, such as Tums and Rolaids. The alkaline compound works by neutralizing the stomach acid responsible for heartburn and indigestion.

Conclusion

Calcium provides a link between tectonics, climate, and the carbon cycle. In the simplest terms, uplift of mountains exposes calcium-bearing rocks to chemical weathering and releases Ca2+ into surface water. This Ca2+ eventually is transported to the ocean where it reacts with dissolved CO2 to form limestone. Some of this limestone settles to the sea floor where it is incorporated into new rocks. Dissolved CO2, along with carbonate and bicarbonate ions, are termed “dissolved inorganic carbon” (DIC). The actual reaction is more complicated and involves the bicarbonate ion (HCO−3) that forms when CO2 reacts with water at seawater pH:

Ca2++ 2HCO−3 → CaCO3 (limestone) + CO2 + H2O

Note that at seawater pH, most of the CO2 is immediately converted back into HCO−3. The reaction results in a net transport of one molecule of CO2 from the ocean/atmosphere into the lithosphere. The result is that each Ca2+ ion released by chemical weathering ultimately removes one CO2 molecule from the surficial system (atmosphere, ocean, soils and living organisms), storing it in carbonate rocks where it is likely to stay for hundreds of millions of years. The weathering of calcium from rocks thus scrubs CO2 from the ocean and atmosphere, exerting a strong long-term effect on climate. Analogous cycles involving magnesium, and to a much smaller extent strontium and barium, have the same effect. As the weathering of limestone (CaCO3) liberates equimolar amounts of Ca2+ and CO2, it has no net effect on the CO2 content of the atmosphere and ocean. The weathering of silicate rocks like granite, on the other hand, is a net CO2 sink because it produces abundant Ca2+ but very little CO2.

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Chemical elements

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