|Boiling Point: 239.25°K, -33.9°C,
Melting Point: 172.31°K, -100.84°C, -149.51°F
Electrons Energy Level: 2,8,7
Isotopes: 18 + 2 Stable
Heat of Vaporization: 10.2 kJ/mol
Heat of Fusion: 3.203 kJ/mol
Density: 3.214 g/L @ 273K & 1atm
Specific Heat: 0.48 J/gK
Atomic Radius: 0.97Å
Ionic Radius: 1.81Å
Electronegativity: 3.16 (Pauling); 2.83 (Allrod Rochow)
Chlorine, which is similar to Fluorine but not as reactive, was discovered
in 1774 by Swedish chemist Carl Wilhelm Scheele
when he combined the mineral pyrolusite (MnO2). It was named in 1810 by Sir Humphry
demonstrated that it was actually a distinct element. It is a greenish-yellow gas with a disagreeable odor (you can detect it near poorly
balanced swimming pools). Its name comes from the Greek word chloros, meaning
greenish-yellow. In high
concentration it is quite toxic and was used in World War I as a poison gas.
Like Fluorine and the other members of the halogen family, Chlorine is diatomic in nature, occurring as Cl2 rather than Cl. It forms -1 ions in ionic compounds with most metals. Perhaps the best known compound of that type is Sodium Chloride, common table salt (NaCl).
Small amounts of Chlorine can be produced in the lab by oxidizing HCl with MnO2. On an industrial scale, Chlorine is produced by electrolysis of brines (aqueous sodium chloride, NaCl , or even sea water. Sodium Hydroxide, NaOH, (also in high demand) is a by-product of the process.
In addition to the ionic compounds that Chlorine forms with metals, it also forms molecular compounds with non-metals such as Sulfur and Oxygen. There are four different oxides of the element. Hydrogen Chloride gas (from which we get Hydrochloric Acid, HCl) is an important industrial product.
Chlorine gas is diatomic with the formula Cl2. It combines readily with nearly all other elements, although it is not as extremely reactive as Fluorine. At 10°C one liter of water dissolves 3.10 liters gaseous Chlorine and at 30 °C only 1.77 liters.
This element is a member of the salt-forming halogen series and is extracted from Chlorides through oxidation often by electrolysis. As the Chloride Ion, Cl, it is also the most abundant dissolved species in ocean water.
Chlorine was discovered in 1774 by Swedish chemist Carl Wilhelm Scheele, who called it dephlogisticated muriatic acid and mistakenly thought it contained oxygen. Chlorine was given its current name in 1810 by Sir Humphry Davy, who insisted that it was in fact an element.
1s2 2s2p6 3s2p5
Chlorine gas, also known as bertholite, was first used as a weapon against humans in World War I by Germany on April 22, 1915 in the Second Battle of Ypres. It was pioneered by a German scientist later to be a Nobel laureate, Fritz Haber. It is alleged that his role in the use of chlorine as a deadly weapon drove his wife to suicide. After its first use, it was utilised by both sides as a chemical weapon.
In nature, chlorine is found mainly as the chloride ion (Cl-), a component of the salt that is deposited in the earth or dissolved in the oceansabout 1.9% of the mass of seawater is chloride ions. Even higher concentrations of chloride are found in the Dead Sea and in underground brine deposits. Most chloride salts are soluble in water, thus, chloride-containing minerals are usually only found in abundance in dry climates or deep underground. Common chloride minerals include halite (sodium chloride), sylvite (potassium chloride) and carnallite (potassium magnesium chloride hexahydrate).
Industrially, elemental chlorine is usually produced by the electrolysis sodium chloride dissolved in water. Along with chlorine, this chloralkali process yields hydrogen gas and sodium hydroxide, according to the chemical equation:
2 NaCl + 2H2O Cl2 + H2 + 2 NaOH
Chlorine has 9 isotopes with mass numbers ranging from 32 to 40. There are two principal stable isotopes, 35Cl (75.77%) and 37Cl (24.23%), found in the relative proportions of 3:1 respectively, giving chlorine atoms bulk an apparent atomic weight of 35.5.
Trace amounts of radioactive 36Cl exist in the environment, in a ratio of about 700x10-15 to 1 with stable isotopes. 36Cl is produced in the atmosphere by spallation of 36Ar by interactions with cosmic ray protons. In the subsurface environment, 36Cl is generated primarily as a result of neutron capture by 35Cl or muon capture by 40Ca. 36Cl decays to 36S and to 36Ar, with a combined half-life of 308,000 years. The half-life of this ydrophilic nonreactive isotope makes it suitable for geologic dating in the range of 60,000 to 1 million years. Additionally, large amounts of 36Cl were produced by irradiation of seawater during atmospheric detonations of nuclear weapons between 1952 and 1958. The residence time of 36Cl in the atmosphere is about 1 week. Thus, as an event marker of 1950s water in soil and ground water, 36Cl is also useful for dating waters less than 50 years before the present. 36Cl has seen use in other areas of the geological sciences, including dating ice and sediments.
Chlorine Gas Extraction
Chlorine can be manufactured by electrolysis of a sodium chloride solution (brine). The production of chlorine results in the co-products caustic soda (sodium hydroxide, NaOH) and hydrogen gas (H2). These two products, as well as chlorine are highly reactive. There are three industrial methods for the extraction of chlorine by electrolysis.
Mercury Cell Electrolysis
Mercury cell electrolysis was the first method used to produce chlorine on an industrial scale. Titanium anodes are located above a liquid mercury cathode and a solution of sodium chloride is positioned between the electrodes. When an electrical current is applied, chloride is released at the titanium anodes and sodium dissolves into the mercury cathode forming an amalgam.
The amalgam can be regenerated into mercury by reacting it with water, producing hydrogen and sodium hydroxide. These are useful byproducts. However, this method consumes vast amounts of energy and there are also concerns about mercury emissions.
Diaphragm Cell Electrolysis
In diaphragm cell electrolysis, an asbestos diaphragm is deposited on an iron grid cathode preventing the chlorine forming at the anode and the sodium hydroxide forming at the cathode from re-mixing.
This method uses less energy than the mercury cell, but the sodium hydroxide is not as easily concentrated and precipitated into a useful substance.
Membrane Cell Electrolysis
The electrolysis cell is divided into two by a cation permeable membrane acting as an ion exchanger. Saturated sodium chloride solution is passed through the anode compartment leaving a lower concentration. Sodium hydroxide solution is circulated through the cathode compartment exiting at a higher concentration. A portion of this concentrated sodium hydroxide solution is diverted as product while the remainder is diluted with deionized water and passed through the electrolyzer again.
This method is nearly as efficient as the diaphragm cell and produces very pure sodium hydroxide but requires very pure sodium chloride solution.
Cathode: 2 H+(aq) + 2 e H2 (g)
Anode: 2 Cl Cl2 (g) + 2 e
Overall equation: 2 NaCl + 2H20 Cl2 + H2 + 2 NaOH
Other Methods of Production
Before electrolytic methods were used for chlorine production, the direct oxidation of hydrogen chloride with oxygen or air was exercised in the Deacon process:
4 HCl + O2 2 Cl2 + 2 H2O
This reaction is accomplished with the use of CuCl2 as a catalysit and is performed in 400°C. The amount of extracted chlorine is approximately at 80%. Due to the extremely corrosive reaction mixture, industrial use of this method is difficult.
Another earlier process to produce chlorine was to heat brine with acid and manganese dioxide.
2 NaCl + 2H2SO4 + MnO2 Na2SO4 + MnSO4 + 2 H2O + Cl2
Using this process, chemist Carl Wilhelm Scheele was the first to isolate chlorine in a laboratory. The manganese can be recovered by the Weldon process.
In a laboratory, small amounts of chlorine gas can be created by adding concentrated hydrochloric acid (typically about 5M) to sodium chlorate solution.
Small amounts of chlorine gas can also be made in the laboratory by putting concentrated hydrochloric acid in a flask with a side arm and rubber tubing attached. Manganese dioxide is then added and the flask stoppered. The reaction is not greatly exothermic. As chlorine is denser than air, it can be easily collected by placing the tube inside a flask where it will displace the air. Once full, the collecting flask can be stoppered.
Applications and Uses
Purification and Disinfection
Chlorine is an important chemical for some processes of water purification, in disinfectants, and in bleach. Ozone can also be used for killing bacteria, and is preferred by many municipal drinking water systems because ozone does not form organochlorine compounds and does not remain in the water after treatment. Ozone is also more reactive and will kill organisms that chlorine will not.
Large amounts of chlorine are used in many industrial processes, such as in the production of paper products, plastics, dyes, textiles, medicines, antiseptics, insecticides, solvents and paints.
Chlorine is also used widely in the manufacture of many every-day items, or to purify water in various forms.
Chlorine is used extensively in organic and inorganic chemistry as an oxidizing agent and in substitution reactions because chlorine often imparts many desired properties to an organic compound when it is substituted for hydrogen (as in synthetic rubber production) because of its high electron affinity. Excess chlorine is removed from water with sulfur dioxide.
World War I
Chlorine became the first killing agent to be employed during World War I. German chemical conglomerate IG Farben had been producing chlorine as a by-product of their dye manufacturing. In cooperation with Fritz Haber of the Kaiser Wilhelm Institute for Chemistry in Berlin, they developed methods of discharging chlorine gas against an entrenched enemy.
Besides the chloride ion (Cl-), chlorine also exists in ionic form as: Cl+1, Cl+5 and Cl+7. Examples are: Hypochlorite (ClO-), Chlorite (ClO2-), Chlorate (ClO3-), and Perchlorate (ClO4-), with ionic charges of -1, +1, +5 and +7 respectively.
It is also used in the production of chlorates, chloroform, carbon tetrachloride, and in bromine extraction.
Two of the most common chlorine compounds are sodium chloride (NaCl) and hydrogen chloride (HCl). Sodium chloride, commonly known as table salt, is used to season food and in some industrial processes. Chlorine (Cl2), when mixed with water (H2O), forms hydrogen chloride or hydrochloric acid (HCl), a strong and commercially important acid. Other chlorine compounds include: chloroform (CHCl3), carbon tetrachloride (CCl4), potassium chloride (KCl), lithium chloride (LiCl), magnesium chloride (MgCl2) and chlorine dioxide (ClO2).
|Chlorine is a toxic gas that irritates the respiratory system.
Because it is heavier than air, it tends to accumulate at the bottom of poorly ventilated
spaces. Chlorine gas is a strong oxidizer, which may react with flammable materials.
Liquid chlorine will burn the skin. Concentrations of the gas as low as 3.5 parts per million will produce an odor and concentrations of 1000 parts per million can be fatal.
Atomic Radius (Å): 0.97Å
Electrochemical Equivalents: 1.3228 g/amp-hr
Atomic Mass Average: 35.4527
(Gr. chloros, greenish yellow) Discovered in 1774 by Scheele, who thought it contained oxygen; named in 1810 by Davy, who insisted it was an element. In nature it is found in the combined state only, chiefly with sodium as common salt (NaCl), carnallite, and sylvite. It is a member of the halogen (salt-forming) group of elements and is obtained from chlorides by the action of oxidizing agents and more often by electrolysis; it is a greenish-yellow gas, combining directly with nearly all elements. At 10oC one volume of water dissolves 3.10 volumes of chlorine, at 30oC only 1.77 volumes. Chlorine is widely used in making many everyday products. It is used for producing safe drinking water the world over. Even the smallest water supplies are now usually chlorinated. It is also extensively used in the production of paper products, dyestuffs, textiles, petroleum products, medicines, antiseptics, insecticides, foodstuffs, solvents, paints, plastics, and many other consumer products. Most of the chlorine produced is used in the manufacture of chlorinated compounds for sanitation, pulp bleaching, disinfectants, and textile processing. Further use is in the manufacture of chlorates, chloroform, carbon tetrachloride, and in the extraction of bromine. Organic chemistry demands much from chlorine, both as an oxidizing agent and in substitution, since it often brings many desired properties in an organic compound when substituted for hydrogen, as in one form of synthetic rubber. Chlorine is a respiratory irritant. The gas irritates the mucous membranes and the liquid burns the skin. As little as 3.5 ppm can be detected as an odor, and 1000 ppm is likely to be fatal after a few deep breaths. It was used as a war gas in 1915. Exposure to chlorine should not exceed 0.5 ppm (8-hour time-weighted average - 40 hour week.)
Source: CRC Handbook of Chemistry and Physics, 1913-1995. David R. Lide, Editor in Chief. Author: C.R. Hammond