Chlorine is the second-lightest halogen, with the symbol Cl and atomic number 17. It has a sickly green colour and a distinctive smell, recognizable to many at low concentrations as 'the smell of pool centers' due to its compounds use as a water disinfecting agent.
Chlorine is a strong oxidizer with 7 valence electrons. Its unstable electron configuration results in high reactivity. Because of this, chlorine usually exists on earth in the form of a halide salt, and free chlorine is rare. Like fluorine, elemental chlorine forms a highly reactive diatomic gas. Chlorine, like other halogens, forms many oxoanions, negatively charged ions containing oxygen. Most notably, these are hypochlorite(ClO-), chlorite(ClO2-), chlorate(ClO3-), and perchlorate(ClO4-).
Unlike hydrochloric acid, elemental chlorine easily corrodes copper, especially in moist air.
Chlorine is a yellow-greenish gas, with a powerful odor similar to that of boiling hypochlorite solutions. It is heavier than air, and slightly soluble in water, 3.26 g/L.
While liquid (as in liquefied, and not aqueous solution) chlorine is sold by gas companies, it is hard to get hold of as it's very toxic and corrosive.
Chlorine is better produced from OTC products.
There are many methods to generating chlorine gas, due to it being such a commonly used ion.
A hypochlorite and hydrochloric acid will produce chlorine; either a solution of sodium hypochlorite or calcium hypochlorite. A violent reaction with a lot of foam may take place in the case of the latter, and starting small scale is a must to get a sense for the reaction before any large scale chlorine production is attempted.
A popular way of making chlorine on Sciencemadness is using hydrochloric acid and trichloroisocyanuric acid (TCCA). TCCA can be found as slow release chlorine tablets for swimming pools. This reaction is favorable because it not too expensive, produces a large amount of chlorine over an extended period of time (while hypochlorites tend to violently produce all the chlorine right on mixing with the acid), leaves no awful byproducts (such as MnO2) and the reaction speed at standard concentrations and temperatures is not too fast nor too slow for most applications.
Chlorine can be used to produce anhydrous metal chlorides, such as aluminum(III) chloride, iron(II) or iron(III) chloride, and many others, which cannot be made in solution, due to formation of hydrates which are irreversible, and decompose to the metal oxide and HCl gas upon heating.
- 2 Al + 3 Cl2 → 2 AlCl3
- 2 Fe + 3 Cl2 → 2 FeCl3
- Make hypochlorites
- Make sulfur dichloride
- Make disulfur dichloride
- Make interhalogen compounds
- Alkane halogenation
Elemental chlorine is extremely toxic and corrosive to many common metals. 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. It is notorious for reacting with iron at high temperatures, in a strong exothermic reaction, known as chlorine-iron fire.
Liquefied chlorine must be stored in cold places, away from any source of heat. Chlorine can be liquefied at room temperature, at a pressure of 7.4 bar.
Chlorine releasing chemicals, such as bleach and TCCA should be stored in closed bottles, usually covered with a bag or in a box, that must be opened form time to time to release the pressure.
The storage area for both chemicals should not contain any metal parts susceptible to chlorine attack.
As it is toxic and has an irritating smell, it is recommended to neutralize chlorine before disposing of it. In gaseous form, elemental chlorine can be neutralized with ammonia, reaction that produces nitrogen gas and ammonium chloride.
Aqueous chlorine however, should never be neutralized with ammonia, as it will generate toxic chloramines. Acids should also be avoided. Hydrogen peroxide will neutralize bleach and release oxygen. Ascorbic acid and its salts are also good at neutralizing chlorine. Other good neutralizing agents are certain sulfur compounds, such as bisulfites, metabisulfites, thiosulfites.
- "Handbook of Toxicology of Chemical Warfare Agents", Academic Press, 2009, p. 313.
- Handbook of Corrosion Data, by Bruce D. Craig, David S. Anderson, p. 271