Chlorine: properties, history, production and uses

Physical and chemical characteristics

Chlorine is highly reactive: it combines with many metals and nonmetals to form salts and covalent compounds. Common stable isotopes are 35Cl and 37Cl, which give the element its average atomic mass. It is more electronegative than most elements and typically exhibits a −1 oxidation state in ionic compounds, though positive oxidation states occur in some molecular species.

History and naming

Elemental chlorine was first recognized in the 18th century; early chemists isolated a pungent gas from reactions involving hydrochloric acid and oxidizing agents. The name "chlorine," derived from the Greek word for green, was adopted in the early 19th century when the element's distinctive color became widely noted.

Production and principal uses

Industrial production is dominated by the chloralkali process, in which brine (salt water) is electrolyzed to give chlorine gas, sodium hydroxide and hydrogen. Key uses include:

• Water disinfection and sanitation (municipal and swimming pools).
• Manufacture of polyvinyl chloride (PVC) and other chlorine‑containing polymers and chemicals.
• Bleaching agents for paper and textiles, and household bleach.
• Intermediates for pharmaceuticals, agrochemicals and solvents.

Safety, environmental and notable facts

Elemental chlorine gas is toxic and can damage respiratory tissues; it was also used as a chemical weapon in World War I. Many chlorine‑containing organic compounds have significant environmental impacts: for example, chlorofluorocarbons (CFCs) released chlorine radicals that contributed to ozone layer depletion. In contrast, simple inorganic chlorides like table salt are essential to life in moderate amounts. Handling and disposal of chlorine and its compounds are strictly regulated in most countries.

For further technical details and data tables consult specialized references or databases: Cl symbol page, element overview, atomic number, atomic mass, halogen group, periodic table.