Chlorine

Chlorine is a chemical element with the symbol Cl and atomic number 17. It is the second-lightest halogen, appearing between fluorine and bromine in the periodic table, and its properties are mostly intermediate between them. Chlorine is a yellow-green gas at room temperature and atmospheric pressure, with a distinctive pungent, suffocating odour. It is an extremely reactive element and a strong oxidising agent: among the elements, it has the highest electron affinity and the third-highest electronegativity, behind only oxygen and fluorine. The free element is rare on Earth and is usually a result of direct or indirect oxidation by oxygen, but the chloride ion (Cl⁻) is abundant in nature and essential to most forms of life. As a chemical building block, chlorine is used in the production of a wide range of industrial and consumer products, including polyvinyl chloride (PVC), disinfectants, solvents, pharmaceuticals, and agrochemicals.

History

The compound most connected with early chlorine chemistry is common salt (sodium chloride). In 1630, the Flemish chemist Jan Baptist van Helmont first recognised that a gas was produced when salt was treated with acid, but he did not identify it as a distinct substance. The gas was first properly studied in 1774 by the Swedish chemist Carl Wilhelm Scheele, who produced it by reacting the mineral pyrolusite (manganese dioxide) with hydrochloric acid. Scheele noted several of its properties—its greenish-yellow colour, its bleaching action on litmus and leaves, its suffocating smell—but, influenced by the phlogiston theory, he called it “dephlogisticated marine acid air” and did not consider it an element.

In 1810, the English chemist Sir Humphry Davy investigated the gas and concluded that it was a chemical element. He named it “chlorine”, from the Greek χλωρός (khlōrós), meaning “greenish-yellow”, in reference to its colour. The name was later standardised internationally. By the 1820s, chlorine was being used as a bleaching agent in the textile industry, and by the mid‑19th century its disinfectant properties were recognised, leading to its application in controlling diseases such as cholera.

Properties

Physical properties

At standard temperature and pressure, chlorine is a diatomic gas with the molecular formula Cl₂. It is about 2.5 times as dense as air. The gas has a pale yellow-green colour and a characteristic sharp, suffocating smell detectable at concentrations as low as 0.3 ppm. Chlorine liquefies at −34.04 °C (−29.27 °F) at atmospheric pressure, forming a clear amber liquid. It solidifies at −101.5 °C (−150.7 °F) into a pale yellow crystalline solid. Liquid and gaseous chlorine are both poor conductors of electricity, but the element exhibits a relatively high dielectric constant in the liquid state. Chlorine is slightly soluble in water (about 3.15 g/L at 20 °C), reacting partially to form hypochlorous and hydrochloric acids; it is more soluble in organic solvents such as carbon tetrachloride and chloroform.

Atomic and chemical properties

A chlorine atom has 17 electrons, arranged in the electron configuration [Ne] 3s² 3p⁵, leaving it one electron short of the stable argon configuration. This makes chlorine highly electronegative (Pauling scale: 3.16) and gives it a very high electron affinity (349 kJ/mol). In compounds, chlorine most commonly exhibits a −1 oxidation state (as chloride, Cl⁻), but it also shows positive oxidation states up to +7 in compounds with oxygen and fluorine, such as chlorates (ClO₃⁻) and perchlorates (ClO₄⁻).

Chlorine reacts directly with almost all elements. With metals it forms ionic chlorides, often with vigorous or even explosive reactions (e.g., sodium burns brilliantly in chlorine gas). With non‑metals it forms covalent compounds, such as hydrogen chloride (HCl) gas, carbon tetrachloride (CCl₄), and phosphorus trichloride (PCl₃). Chlorine also forms interhalogen compounds with other halogens, such as chlorine trifluoride (ClF₃) and iodine monochloride (ICl). In the presence of water, chlorine disproportionates to form hydrochloric acid (HCl) and hypochlorous acid (HOCl), the latter being the active bleaching and disinfecting species.

Isotopes

Naturally occurring chlorine consists of two stable isotopes: chlorine‑35 (²⁵Cl) and chlorine‑37 (³⁷Cl), with relative abundances of approximately 75.76% and 24.24%, respectively, giving the element a standard atomic weight of 35.45. Both isotopes have nuclear spin 3/2 and are used in nuclear magnetic resonance (NMR) spectroscopy. A number of radioactive isotopes have been produced artificially, ranging from ²⁸Cl to ⁵¹Cl; the longest‑lived, ³⁶Cl, has a half‑life of 301,000 years and is produced in the atmosphere by spallation of argon and used in dating groundwater and ice cores.

Occurrence

Chlorine is too reactive to exist as the free element in nature, but chloride compounds are widespread. The principal source is sodium chloride (halite, rock salt), which occurs in vast evaporite deposits—remnants of ancient seas—and is also dissolved in seawater. The world’s oceans contain about 2.3 × 10¹⁶ tonnes of chloride ions, corresponding to an average salinity of about 3.5%. Other significant chloride minerals include sylvite (potassium chloride, KCl) and carnallite (KMgCl₃·6H₂O). Chlorine is the 20th most abundant element in the Earth’s crust, at about 130 ppm by weight.

Production

Industrial production

Most chlorine is manufactured by the electrolysis of brine (a concentrated solution of sodium chloride). Three main cell types are used: the mercury cell, the diaphragm cell, and the membrane cell. The membrane cell has become dominant because it avoids the use of mercury and asbestos and is more energy‑efficient. In the membrane process, an ion‑exchange membrane separates the anode and cathode compartments. Chlorine gas is generated at the anode:

2 Cl⁻ → Cl₂ + 2 e⁻

Hydrogen gas and sodium hydroxide are produced at the cathode:

2 H₂O + 2 e⁻ → H₂ + 2 OH⁻

The overall reaction is:

2 NaCl + 2 H₂O → Cl₂ + H₂ + 2 NaOH

Global chlorine production capacity exceeds 90 million tonnes per year. China is the largest producer, followed by the United States and Western Europe. Chlorine is typically produced on‑site or transported as a pressurised liquid in steel cylinders or tank cars, because of the hazards and cost of long‑distance transport.

Laboratory preparation

In the laboratory, small quantities of chlorine can be prepared by the oxidation of hydrochloric acid with strong oxidising agents such as manganese dioxide (MnO₂) or potassium permanganate (KMnO₄). The classic method using manganese dioxide:

4 HCl + MnO₂ → MnCl₂ + 2 H₂O + Cl₂

The gas is typically dried by passing it through concentrated sulfuric acid and collected by downward displacement of air.

Applications

Chlorine and its compounds are integral to modern society, with uses spanning water treatment, plastics, solvents, pharmaceuticals, and crop protection.

Water and waste treatment

Chlorine is widely used to disinfect drinking water and wastewater, eliminating bacteria, viruses, and protozoa. It is typically applied as chlorine gas, sodium hypochlorite (liquid bleach), or calcium hypochlorite (powder). Chlorination has been instrumental in reducing waterborne diseases such as typhoid and cholera globally. In swimming pools, chlorine‑based compounds maintain hygienic conditions. Chlorine is also used to control biofouling in industrial cooling water systems.

Production of chemicals and plastics

About 30% of chlorine production is used to manufacture polyvinyl chloride (PVC) via the intermediate vinyl chloride monomer. PVC is employed in pipes, window frames, flooring, medical devices, and countless other products. Chlorine is also a raw material for the production of chlorinated solvents (e.g., methylene chloride, trichloroethylene), epichlorohydrin (for epoxy resins), and phosgene (COCl₂), which is a key intermediate for polycarbonates and polyurethanes.

Pharmaceuticals and agrochemicals

Chlorine is present in around 85% of pharmaceutical compounds either as a structural component or as a reagent in their synthesis. It is also essential in the production of many crop protection agents, including herbicides (e.g., atrazine, 2,4‑D), insecticides (e.g., DDT, chlorpyrifos), and fungicides. Chlorinated organic compounds often show enhanced biological activity and stability.

Disinfection and bleaches

Household and industrial bleach products rely on chlorine chemistry. Sodium hypochlorite solution (typically 3–6% available chlorine) is a common laundry bleach and surface disinfectant. Chlorine dioxide (ClO₂) is used for bleaching wood pulp in the paper industry and as a disinfectant with reduced formation of chlorinated organic by‑products compared to elemental chlorine.

Other uses

Chlorine is employed in the extraction and purification of metals such as titanium and magnesium from their ores. It is used in the manufacture of synthetic rubber, dyes, and explosives. In the electronics industry, high‑purity hydrogen chloride is used for etching semiconductor materials. Additionally, chlorine chemistry is central to the production of chlorine‑containing flame retardants and refrigerants, although many chlorofluorocarbons (CFCs) have been phased out due to ozone‑depletion concerns.

Biological role and toxicity

Essentiality

The chloride ion is an essential nutrient for all animal life. It is the most abundant anion in the extracellular fluid, playing critical roles in maintaining osmotic pressure, acid–base balance, and the generation of nerve impulses. The stomach uses chloride to produce hydrochloric acid (gastric acid), which is essential for digestion and controlling microbial growth. The recommended dietary intake for humans is around 2.3 g per day; deficiency is rare outside cases of severe vomiting or certain kidney disorders.

Toxicity of elemental chlorine

Chlorine gas is highly toxic and was used as a chemical weapon during the First World War. Inhalation causes irritation of the respiratory tract, coughing, choking, and pulmonary oedema at higher concentrations. The characteristic sharp odour provides some warning, but prolonged exposure can deaden the sense of smell. Concentrations above 30 ppm can cause immediate chest pain and vomiting; above 430 ppm, death occurs within 30 minutes. Modern occupational exposure limits (e.g., OSHA PEL) are set at 1 ppm as a ceiling value. Treatment for chlorine exposure is supportive, often involving humidified oxygen and bronchodilators.

Environmental considerations

While chlorine itself is not persistent in the environment, many chlorinated organic compounds are resistant to degradation and can bioaccumulate. Accidental releases of chlorine gas can cause localised harm to vegetation and wildlife. Strict regulations govern handling, storage, and transport. The chlorine industry follows a “Responsible Care” programme and promotes the safe management of chlorinated chemicals.

Compounds

Chlorine forms an enormous range of compounds, reflecting its versatility. The most significant are:

• Hydrogen chloride (HCl): A colourless gas that dissolves in water to form hydrochloric acid, a fundamental laboratory and industrial chemical.
• Sodium hypochlorite (NaOCl): The active constituent of household bleach.
• Chlorine dioxide (ClO₂): A selective oxidant used in pulp bleaching and disinfection.
• Chlorates and perchlorates: Oxidising agents used in fireworks, matches, and explosives (e.g., potassium chlorate, KClO₃; ammonium perchlorate, NH₄ClO₄).
• Organochlorines: A vast class including PVC, solvents like dichloromethane (CH₂Cl₂), chloroform (CHCl₃), carbon tetrachloride (CCl₄), and numerous pesticides and pharmaceuticals.

Safety and handling

Chlorine is classified as a toxic, corrosive, and environmentally hazardous gas. Specialised materials such as steel, Monel, or PTFE are used for piping and containers because dry chlorine does not attack these materials. However, in the presence of moisture, the corrosive combination of hydrochloric and hypochlorous acids requires strict moisture exclusion. Chlorine storage areas must be well ventilated, and emergency scrubbers are often installed to neutralise accidental releases. Because chlorine is an oxidiser, it can support combustion and react dangerously with fuels, ammonia, and organic materials.

See also

• Halogen
• Chlorination
• Polyvinyl chloride
• Chlorine cycle

References

• Emsley, J. (2011). Nature’s Building Blocks: An A–Z Guide to the Elements. Oxford University Press.
• Greenwood, N. N., & Earnshaw, A. (1997). Chemistry of the Elements (2nd ed.). Butterworth‑Heinemann.
• U.S. EPA. (2023). Chlorine and Drinking Water.
• Chlorine Institute. (2024). Pamphlet 1: Chlorine Basics.
• World Health Organization. (2017). Guidelines for Drinking‑water Quality.