Chlorine is the second lightest gas amongst the halogen group of gases. Amongst the distinct characteristics of chlorine, include its yellow-greenish appearance, with a choking smell at room temperature, high electron affinity, and high electronegativity, which make it a strong oxidizing agent. In terms of abundance, the gas comes second from amongst the halogen gases, after fluorine, and is the twenty first most abundant chemical element on earth. It is therefore, a common gas, with its most common compound being sodium chloride, locally known as common salt. In addition to being used as a food additive in its compound form, chlorine compounds are also used as intermediates in the production of plastics. In the periodic table of elements, chlorine is a diatomic non-metal, which appears in group 17. This group encompasses the halogen gases thus, its properties are almost similar to those of other halogen gases in the group, including fluorine, bromine and iodine. It has an electron configuration of 2.8.7 (Greenwood and Earnshaw, 1984). The seven electrons in the third shell, which is also chlorine’s outermost shell, acts as chlorine’s valence electrons. This electrons readily reacts with any element in order to attain a full octet configuration. This is why the gas is a strong oxidizing agent. The gas also has a relatively high electronegativity as compared to the gases in group 17. It is less reactive when compared to fluorine, but more reactive than the rest of the halogens. It is thus, the second most reactive halogen gas. Chlorine’s electrons are held together by weak intermolecular forces, called the van der Waals forces. This is responsible for the low melting and boiling points of chlorine, at -34 and -101 degrees Celsius respectively (Rsc.org. 2016). Very little heat is required to break the weak van der Waals forces. The darkening of the halogens color down the group from fluorine is responsible for the yellow-greenish color of chlorine. This is portrayed by fluorine, which appears first in the group, being pale yellow, followed by chlorine, which is yellow-greenish as the degree of darkness continues down the group. Gas is also a poor conductor of electricity. Chlorine is also amongst the few elements in the periodic table which exhibit isotopic characteristics. This means that chlorine has the ability to appear in more than one stable form. This gas has two isotopes, including chlorine (35) and chlorine (37). These two isotopes are the only naturally occurring forms of the element, with varying percentages of abundance. Chlorine 35 accounts for 76% abundance, while chlorine 37 represents the remaining 24% of the world’s naturally occurring chlorine (Cameron, 1973). Chlorine as an element, does not occur in nature, and this is attributed to its strong reactive property. It is, however, abundantly available in nature in the form of its chloride compounds, most of which are salts (Greenwood and Earnshaw, 1984). Its abundance is portrayed by being the 21 st most abundant element on earth. Sodium chloride, an ancient compound of chlorine, which evaporates from salty water bodies, also reveals the abundance of chlorine and its compounds. Small amounts of chlorine can also be chemically prepared in the laboratory. This can be achieved by mixing hydrochloric acid and manganese dioxide. The preparation of large amounts of chlorine, however, is done in the industry through the electrolysis of sodium chloride which has been dissolved in water. This method also known as the chloralkali process, and is responsible for industrial chlorine, which can then be used in different chemical processes amongst other applications (Greenwood and Earnshaw, 1984). The applications of chlorine gas are very risky applications because the gas is very toxic to the extent of being used as a chemical weapon. The acidic properties of chlorine enable the element to kill bacteria. Bacteria and other forms of pathogens cannot survive in chlorine. This element is thus, used in the treatment of drinking water and even, swimming pool water. This property of chlorine plays a very vital biological role with respect to the human body. In the human digestive system, there are traces of hydrochloric acid present in the stomach lining, which apart from aiding digestion, also kills germs and micro-organisms that might have been ingested together with food. The element, therefore, cleans the stomach. In the body, chlorine is also present in cell fluid and extracellular fluid. In the cell, the element appears as a negative ion, and plays a vital role in balancing the positive ions in the fluid cell, mainly potassium, while in the extracellular fluid, it balances positive sodium ions. The industrial applications of chlorine include the manufacture of plastics. Close to 20% of the total amount of chlorine produced is often aimed at the production and manufacture of PVC. PVC, once manufactured, can be used in a diverse number of applications, including in the making of vehicle interiors, the making of electrical insulators, the manufacture of water pipes, and the production of blood bags. Chlorine is also a vital element and plays an important applicable role in organic chemistry. In organic chemistry, the chemical properties of chlorine has made it a good oxidizing agent in most of substitution reactions. This is employed by pharmaceuticals in the production and manufacture of medicine. Lastly, chlorine was also used in the manufacture of chloroform, an aesthetic, but its production was stopped because of its tendency to cause liver damage. Chlorine is, therefore, a toxic element whose history dates back to 1774, when the gas was first produced by Carl Wilhelm. The chemical compound of the gas sodium chloride has also been used since time immemorial, as a food additive, commonly referred to as table salt. Despite its toxic property, chlorine elements are of importance to both chemical and biological processes.
References
Cameron, A. (1973). Abundances of the elements in the solar system. Space Science Reviews , 15(1). Available at: http://dx.doi.org/10.1007/bf00172440 [Accessed 23 Nov. 2016].
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Greenwood, N. and Earnshaw, A. (1984). Chemistry of the elements . Oxford [Oxfordshire]: Pergamon Press.
Rsc.org. (2016). Chlorine - Element information, properties and uses | Periodic Table . Available at: http://www.rsc.org/periodic-table/element/17/chlorine [Accessed 23 Nov. 2016].
