Chlorine

Chlorine
Atomic number	17
Symbol	Cl
Group	17 (halogens)
Boiling point	−34.04 °C
Electronegativity	3.16 (Pauling)
Electron configuration	[Ne] 3s2 3p5
Period	3
Main isotopes	35Cl, 37Cl
Oxidation states	−1, +1, +7
Phase STP	Gas
Wikidata	Q688

Chlorine (symbol Cl, Z=17) is a chemical element in group 17 (the halogens) of the periodic table. At standard conditions (25 °C, 1 atm) it is a pungent yellow-green diatomic gas (Cl₂). As a halogen, chlorine is one electron short of a full octet (electronic configuration [Ne] 3s²3p⁵), making it very reactive and a strong oxidizer. This reactivity has made chlorine invaluable as a disinfectant (for drinking water and swimming pools) and bleach (e.g. household bleach, sodium hypochlorite), and as a chemical feedstock. Industrially, tens of millions of tonnes of chlorine are produced each year by electrolyzing salt (sodium chloride brine) in the chlor-alkali process. Chlorine is a key ingredient in thousands of products, especially chlorinated organic chemicals such as polyvinyl chloride (PVC) plastic. In nature, chlorine is never found as the free element; it occurs mainly as chloride ions (Cl⁻) in saltwater and minerals. In the body, chloride ions are essential electrolytes (e.g. in nerves, muscles, stomach acid), but elemental chlorine gas itself is toxic and was even used as a chemical weapon in World War I.

Overview and Key Facts

Chlorine is a nonmetallic halogen (Group 17, period 3, p-block) with atomic number 17. Its standard atomic weight is about 35.45 u (determined by its stable isotopes). In its elemental form at room temperature and pressure, chlorine is a yellow-green gas composed of Cl₂ molecules. The element is known for its high reactivity – among the strongest of the nonmetallic oxidizers – due to its electron configuration [Ne] 3s²3p⁵. With seven electrons in its outer shell, chlorine readily accepts one more electron to reach the noble-gas configuration of argon. As a result, chlorine commonly occurs in oxidation state –1 (as in metal chlorides like NaCl or acids like HCl), but it also exhibits positive states +1, +3, +5, and +7 in various oxyanions and oxyacids (such as ClO⁻, ClO₂⁻, ClO₃⁻, and ClO₄⁻).

At standard conditions chlorine is a gas (Cl₂) with a sharp, irritating odor. Under higher pressure or lower temperature it condenses to a clear amber liquid (at –34.0 °C) and then solidifies (melting at –101.5 °C) into pale yellow orthorhombic crystals. The gas is about 2.5 times as dense as air (≈3.2 kg/m³ at 0 °C, 1 atm). Chlorine’s electronegativity is 3.16 (Pauling scale), which is the third-highest of all elements (behind only oxygen and fluorine), and its electron affinity (349 kJ/mol) is the highest of any element. These values reflect chlorine’s strong pull on electrons and its tendency to form Cl⁻. In practice, chlorine chemistry is dominated by chloride salts (e.g. NaCl, MgCl₂) and by its oxyanions and oxyacids (hypochlorite ClO⁻, chlorite ClO₂⁻, chlorate ClO₃⁻, perchlorate ClO₄⁻). Aside from Cl₂ gas, elemental chlorine is not found in nature; it occurs only as chloride compounds. It is the second-most abundant halogen in the Earth’s crust (after fluorine) – about 0.02% by mass – and exists mainly in dissolved form in seawater (about 19 g of chlorine per liter as Cl⁻) or in evaporite minerals (halite NaCl, sylvite KCl, carnallite KCl·MgCl₂·6H₂O, etc.).

Atomic Structure and Electron Configuration

Chlorine has 17 protons in its nucleus and (when neutral) 17 electrons. Its ground-state electron configuration is [Ne] 3s²3p⁵, which means its valence shell (n=3) contains two 3s electrons and five 3p electrons, for a total of seven valence electrons. This one-electron shortfall from the argon configuration (3s²3p⁶) gives chlorine a strong tendency to gain an electron and become Cl⁻. The first ionization energy of chlorine is about 1251 kJ/mol (12.97 eV), which means a fair amount of energy is required to remove an electron. This ionization energy is lower than fluorine’s but higher than that of its heavier congeners (bromine, iodine), reflecting periodic trends. Chlorine’s electronegativity (3.16) and electron affinity (349 kJ/mol) are very high, underscoring its powerful ability to attract electrons in chemical bonds.

The atomic radii reflect chlorine’s intermediate position in the halogens. Its covalent radius is about 100 pm, while its van der Waals radius is ~175 pm. When chlorine gains an electron to form Cl⁻, its ionic radius expands to about 181 pm. In periodic trends down group 17, atomic size increases (fluorine ~64 pm, chlorine ~100 pm, bromine ~115 pm, iodine ~133 pm covalent radius), so chlorine is larger than fluorine but smaller than bromine. Electronegativity decreases down the group, so chlorine (3.16) is less electronegative than fluorine (3.98) but more than bromine (2.96). These trends contribute to chlorine’s chemical behavior: it is extremely reactive (less so than fluorine) and is a strong oxidizing nonmetal.

Isotopes and Nuclear Properties

Natural chlorine consists of two stable isotopes: ^35Cl (about 75.8% abundance) and ^37Cl (24.2%). Their relative abundance gives chlorine an atomic weight of ≈35.45 u. Both ^35Cl and ^37Cl nuclei have spin 3/2⁺, making them NMR-active with large quadrupole moments (so their NMR signals are broadened). No other chlorine isotopes are stable in nature. The next isotope, ^36Cl, is radioactive (half-life ≈3.01×10^5 years) and is produced cosmogenically (by cosmic-ray spallation of argon in the atmosphere) or by neutron activation (for example in rocks or during nuclear tests). ^36Cl decays by β⁻ emission to ^36Ar and by electron capture to ^36S. Its long half-life makes it useful as an environmental tracer: ^36Cl/^35Cl ratios can date groundwater or ice up to ~10^6 years old. In fact, atmospheric nuclear bomb tests in the 1950s created a spike of ^36Cl worldwide, which now serves as a “time marker” for water and ice formed in that era. Tiny amounts of ^36Cl also occur naturally in rain and snow. Other chlorine radioisotopes (masses 34, 36–44) have very short half-lives (seconds to minutes) and no practical applications outside research.

Allotropes and Typical Compounds

Chlorine does not have multiple elemental allotropes beyond its molecular form. Its normal elemental form is diatomic chlorine (Cl₂) in the gas, liquid, or solid state. (Even in the solid phase, it exists as discrete Cl₂ molecules arranging in an orthorhombic lattice.) There are no polymeric or atomic allotropes of chlorine known under ordinary conditions.

Chlorine forms a vast array of compounds. As a halogen, it most commonly forms a –1 anion (Cl⁻) with metals, making ionic chlorides. Prime examples include sodium chloride (NaCl, table salt), potassium chloride (KCl), and calcium chloride (CaCl₂). Many metal chlorides (MgCl₂, AlCl₃, FeCl₃, etc.) are important chemical reagents or catalysts. Chlorine also bonds covalently with nonmetals and in organics: hydrogen chloride (HCl) is a covalent gas which dissolves in water to give hydrochloric acid (a strong mineral acid). Organic chlorides (alkyl halides) include solvents like chloroform (CHCl₃), dichloromethane (CH₂Cl₂), and carbon tetrachloride (CCl₄), as well as countless industrial chemicals (e.g. vinyl chloride CH₂CHCl for PVC, and halogenated pesticides and pharmaceuticals). In these compounds, the C–Cl bond is polar (Cl is more electronegative than C).

A key family of chlorine compounds is the oxyacids and oxyanions. Chlorine can have positive oxidation states up to +7 when bound to oxygen. Important oxyanions and their acids include:

• Hypochlorite: ClO⁻ (anion), HOCl (hypochlorous acid) – a weak, unstable acid (pKa ≈7.5) and the active ingredient in bleach (NaClO solution, corrosion protect).
• Chlorite: ClO₂⁻ (chlorite), from HClO₂ (chlorous acid; unstable).
• Chlorate: ClO₃⁻, from HClO₃ (chloric acid; strong, but tends to decompose).
• Perchlorate: ClO₄⁻, from HClO₄ (perchloric acid; a superacid, fully dissociated).

Hypochlorite and chlorate salts (e.g. NaClO, NaClO₃) are strong oxidizers used in bleaching and water treatment. Perchlorate salts (e.g. NH₄ClO₄) are powerful oxidizers used in rocket propellants and fireworks.

Chlorine also forms several interhalogen compounds by combining with other halogens. Notable examples include chlorine monofluoride (ClF), chlorine trifluoride (ClF₃), and chlorine pentafluoride (ClF₅), as well as bromine chloride (BrCl) and iodine monochloride (ICl). These interhalogens are often extremely reactive and sometimes volatile (ClF₃ is a toxic fuming liquid). Chlorine oxides also exist (Cl₂O, ClO₂, Cl₂O₇, etc.), though they are typically unstable; for example, chlorine dioxide (ClO₂) is a yellow paramagnetic gas used as a bleach oxidant.

In summary, chlorine’s typical compounds include chlorides (with nearly all metals), hydrogen chloride, numerous oxygenated species (oxyanions and acids), interhalogens, and organic chlorides. This vast chemistry underlies chlorine’s many applications.

Physical Properties

Elemental chlorine is a diatomic molecular gas under ordinary conditions. It is visible as a dense yellow-green gas or cloud; its color is due to absorption of longer wavelengths and scattering of shorter (green) wavelengths of light. The gas has a pungent, suffocating odor detectable at very low concentrations (ppm). At 0 °C and 1 atm, chlorine gas has a density ≈3.214 kg/m³ (about 2.48 times the density of air).

When cooled or pressurized, chlorine condenses to a clear amber liquid at –34.04 °C (239.11 K). At 15 °C (under ≈597 kPa), liquid Cl₂ has a density around 1420–1460 kg/m³ (about 1.47 times the density of water). The volume expansion ratio (liquid to gas) is large: one volume of liquid Cl₂ equals about 456 volumes of gas at STP. Chlorine solidifies at –101.5 °C (171.6 K) into a pale yellow crystalline solid with an orthorhombic lattice (each Cl₂ molecule is held by van der Waals forces in a molecular solid). The melting point is about –100.98 °C. The triple point (gas-liquid-solid equilibrium) is around 1.2 kPa at –101.5 °C.

Chlorine’s critical temperature and pressure are 143.75 °C and 7977 kPa, respectively, above which Cl₂ cannot be liquefied by pressure alone. The specific heat of gaseous chlorine (at 0 °C, 1 atm) is about 0.521 kJ/(kg·K) at constant pressure.

Chlorine gas is only slightly soluble in water: about 8.3 g per liter at 15.6 °C, forming a mixture of hydrochloric and hypochlorous acids. Solubility decreases with increasing temperature. In organic solvents, chlorine is more soluble (it is nonpolar). Liquid chlorine is light amber and when illuminated (e.g. by sunlight), it decomposes slowly.

Chlorine is not flammable itself, but as a strong oxidizer it can support the combustion of other substances. The ignition temperature of hydrogen in chlorine is low. Chlorine gas does not conduct electricity, but when it dissolves in water it forms ions (H⁺, Cl⁻, HOCl⁻) that can conduct.

Spectroscopically, chlorine’s spectral lines include a prominent atomic green line near 553 nm (seen when Cl₂ is excited). The molecule has vibrational bands in the infrared around 600–750 cm⁻¹ (absorption peaks), but as a homonuclear diatomic it has no permanent dipole, so its IR absorption is via these vibrational transitions.

In summary, chlorine’s physical properties reflect its small molecular size and van der Waals bonds: low melting/boiling points, gas at STP, and a dense, nonpolar gas phase.

Chemical Reactivity and Trends

Chlorine is highly reactive, especially as an oxidizing agent. It has a standard reduction potential of +1.36 V for the half-reaction Cl₂ + 2e⁻ → 2Cl⁻ (in acidic aqueous solution), indicating that Cl₂ readily takes electrons to form chloride. For example, chlorine oxidizes hydrogen to form HCl (Cl₂ + H₂ → 2HCl) and metals to form metal chlorides (e.g. 2Fe + 3Cl₂ → 2FeCl₃). It will even react with some inert substances (e.g. carbon, white phosphorus) under suitable conditions. In many reactions chlorine is reduced by one electron per atom (forming Cl⁻) or in steps (Cl₂ → Cl· → Cl⁻).

Chlorine’s trend in the halogen group (F > Cl > Br > I) means it is less reactive than fluorine but more reactive than bromine. It will displace bromine or iodine from their compounds. For example, Cl₂ + 2KI → 2KCl + I₂ (freezing bromine to precipitate as Bromide).

In water, chlorine undergoes complex redox hydrolysis. A key reaction is: $$\text{Cl}2 + \text{H}2\text{O} \rightleftharpoons \text{HCl} + \text{HOCl},$$ which yields hydrochloric acid and hypochlorous acid. Because HOCl is a weak acid (pKₐ ≈ 7.5), in basic or neutral water some hypochlorite ion (ClO⁻) also forms: $$\text{Cl}2 + 2\text{OH}^- \to \text{Cl}^- + \text{ClO}^- + \text{H}_2\text{O}.$$ These species (HOCl/OCl⁻) are strong oxidizers and are responsible for chlorine’s disinfectant properties. Chlorate and perchlorate ions are also oxidizing but are far less reactive in water.

As a base concept, chlorine is often encountered as the chloride ion (Cl⁻) in acids and salts. Cl⁻ is a very weak base (conjugate of HCl, a strong acid), so in aqueous solution it is generally inert (non-oxidizing). In contrast, chlorine in positive oxidation states (as above) yields acids/hypochlorites which act as moderate to strong oxidizers. Perchloric acid (HClO₄) is a strong acid and a strong oxidizer, whereas HCl (hydrochloric acid) is a strong acid but a poor oxidizer under normal conditions.

Chlorine also forms complexes with transition metals. For example, AlCl₃ (aluminum trichloride) is a covalent compound that dimerizes (Al₂Cl₆) and is a Lewis acid catalyst. Iron(III) chloride FeCl₃ is strongly ionic/covalent and is used as a catalyst and metal chloride pigment. Platinum and gold form stable complex chlorides (e.g. K₂PtCl₄, AuCl₃).

Corrosion wise, chloride ions promote pitting corrosion in steels by breaking down oxide layers. Stainless steels are designed to resist chloride attack, but high concentrations of chlorine or its compounds can still corrode many metals. Passivation can occur with metals that form protective oxide coats (e.g. aluminum), but chloride can penetrate many such layers.

Chlorine’s reactivity with organic compounds typically requires activation (heat, light, or catalysts). It will add across C=C double bonds (forming dichloroe.g. R–CHCl–CHCl–R) or substitute for hydrogen in many bonds (especially in the presence of UV light or radical initiators, forming R–Cl). However, because such reactions are often uncontrollable, chemists frequently use gaseous Cl₂ only in moderation (one example is the production of vinyl chloride by Cl₂ addition to acetylene, followed by elimination).

Acid/base trends: Hypochlorous acid (HOCl) is a weak acid and unstable (disproportionating to O₂ and Cl⁻). Chloric (HClO₃) is a strong but unstable acid (dissociates to ClO₃⁻). Perchloric acid (HClO₄) is a superacid (stronger than sulfuric acid) but strong oxidizer when pure.

Summarizing, chlorine is a strong one-electron acceptor. It forms Cl⁻ readily, and it forms oxidizing species like HOCl/OCl⁻. Its reactivity is governed by its high electronegativity and electron affinity, which decrease down the halogen group, making chlorine less fierce than fluorine but still highly aggressive compared to most substances.

Occurrence and Production

On Earth, chlorine is relatively abundant but is not found as the free element. About 0.02% of the continental crust is chlorine by weight (roughly 150–200 ppm), making it the 21st most abundant element in the crust. Its crustal abundance comes almost entirely from chloride salts. The most common chloride mineral is halite (rock salt, NaCl), often found in vast sedimentary deposits. Other important minerals are sylvite (potash, KCl) and carnallite (KMgCl₃·6H₂O) among potash ores. Seawater is the richest source of chlorine: it contains about 19 grams of chlorine per liter (as chloride), corresponding to roughly 1.9% by weight. There are also saline lakes, brine wells, and groundwater reservoirs that hold large quantities of chloride.

Elemental chlorine is produced industrially by the electrolysis of brine (sodium chloride solution) in the chlor-alkali process. In the most common membrane-cell process, brine is separated into two chambers by an ion-exchange membrane. At the anode, chloride ions are oxidized to chlorine gas: 2Cl⁻ → Cl₂ + 2e⁻. At the cathode, water is reduced to produce hydrogen gas and hydroxide: 2H₂O + 2e⁻ → H₂ + 2OH⁻. The net result is the simultaneous production of Cl₂ gas, H₂ gas, and aqueous NaOH. This process is energy-intensive but economical at large scale. Modern chlor-alkali plants produce tens of millions of tonnes of chlorine per year worldwide. Major chlorine producers include chemical companies in the United States, China, India, Europe, and elsewhere where large salt deposits or brine sources are available.

Historically, other production methods included the diaphragm cell (using asbestos or polymer diaphragms) and the mercury cell (where mercury forms an amalgam with sodium). These older methods are largely being phased out due to environmental concerns (mercury emissions) and lower efficiency.

In the laboratory, chlorine can be prepared on a small scale by certain chemical reactions, such as heating solid sodium chlorate (NaClO₃) or by reacting hydrochloric acid with an oxidizing agent (e.g. MnO₂ + 4HCl → 2Cl₂ + MnCl₂ + 2H₂O). Chlorine is commonly sold in metal cylinders as a compressed liquefied gas (for industrial or municipal use). Because chlorine is not gaseous in nature, there are no natural sources of pure Cl₂ gas; every use of elemental chlorine relies on industrial generation.

Cosmically, chlorine was formed by stellar nucleosynthesis (in supernovae) but is not one of the most common elements in the universe. Its cosmic abundance is small (on the order of 10⁻⁶ relative to hydrogen). It is found in some meteorites (often as sulfates or chlorides) and in interstellar gas, but its main reservoir on Earth is indeed ocean salt and evaporite minerals.

Applications and Technology

Chlorine’s utility spans many industries due to its chemical reactivity. The single largest application is in making polyvinyl chloride (PVC), a durable plastic used for pipes, cable insulation, clothing, and more. PVC production consumes about a third of industrial chlorine (through the monomer vinyl chloride). Chlorine is also a building block for many other organic chemicals. For example, chloromethanes (like CH₂Cl₂, CHCl₃) and trichloroethanes are solvents and refrigerants (some are now phased out). Chlorobenzene and dichlorobenzenes are used as intermediates for other chemicals. Chlorine enables the production of fluorocarbons (e.g. Freon) by providing Cl to react with fluorine precursors (though these have mostly been discontinued due to ozone concerns).

As a bleaching agent, chlorine is crucial. Sodium hypochlorite solution (“liquid bleach”) and solid calcium hypochlorite are widely used to whiten fabrics and paper. Chlorine dioxide (ClO₂) is used in paper mills and water treatment as an effective bleach that produces fewer chlorinated byproducts. Elemental chlorine itself is used in pulp bleaching and sanitization of water (added to drinking water to kill pathogens). In water treatment, chlorine kills bacteria, viruses, and other microbes, making water safe to drink. Swimming pools are routinely chlorinated (often via added hypochlorite) to maintain sanitation.

Chlorine and its derivatives play roles in metallurgy and material processing. Titanium and zirconium metals are obtained by reducing TiCl₄ and ZrCl₄ (made in chlorine-based reactions) respectively. Aluminum chloride (AlCl₃) and iron(III) chloride (FeCl₃), both derived from chlorine chemistry, catalyze polymerization and other organic processes. Hydrogen chloride (HCl), produced from chlorine, is a widely used strong acid in refining and pickling steel.

In the electronics industry, chlorine is used in processes such as silicon and compound semiconductor etching (for instance, Cl₂ or BCl₃ plasmas remove materials during chip fabrication). In analytical chemistry, liquid chlorine is used for disinfection of reagents and cleaning glassware.

Household and consumer products often involve chlorine: bathroom cleaners often have bleach (NaOCl), drinking water (tap water) is chlorinated, and bleach tablets (NaOCl or Ca(OCl)₂) are used in pools and sanitation. Invented in the 19th century, bleaching powder (calcium hypochlorite) by Labarraque and later Dakin’s antiseptic solution (weak NaOCl) are early medical uses.

Chlorine-containing compounds are also important in medicine (many drugs contain chlorine atoms) and agriculture (some herbicides and pesticides, though environmental impact is carefully regulated). Chloroform (CHCl₃) was historically an anesthetic; many antibiotics and chemicals have chlorine substituents.

While not an energy source per se, some energetic materials involve chlorine: perchlorates are used in solid rocket propellants and fireworks. Chlorine trifluoride (ClF₃) is a hypergolic propellant/oxidizer in rocketry (highly dangerous to handle, used only in specialized situations).

In summary, chlorine’s major industries are water treatment/disinfection, plastics (PVC) and other chlorine-based organics, detergents and bleaches, and chemical intermediates. Its role as an intermediate in the manufacture of many everyday products (pharmaceuticals, solvents, polymers, disinfectants) makes it one of the most important industrial chemicals.

Biology, Environment, and Safety

Biology: Chlorine in the form of chloride ions is essential for life. In animals, Cl⁻ is a key electrolyte in bodily fluids: it helps regulate osmotic balance and blood pressure in conjunction with sodium and potassium ions. It is also vital in nerve and muscle function (neuronal Cl⁻ channels shape signals) and in digestion (stomach parietal cells secrete HCl using Cl⁻ ions, aiding protein breakdown). Chloride is also required by certain plants for osmotic regulation and photosynthesis (it acts as a cofactor in some enzyme processes). Some immune cells (neutrophils) generate hypochlorous acid (HOCl) from chloride ions and hydrogen peroxide to kill bacteria during respiratory burst; this is a natural defensive use of chlorine chemistry.

No organism uses elemental chlorine gas internally; chlorine’s biological roles are solely via Cl⁻ or the derived acids/oxidants. Conversely, many chlorine-containing organic compounds (pesticides, dioxins, PCBs) are toxic or persistent pollutants because biological systems cannot easily break the C–Cl bond. For example, polychlorinated biphenyls (PCBs) and certain chlorinated herbicides accumulate in fat and can cause long-term health issues.

Environment: When chlorine is released into the environment (e.g. from industrial effluent), it rapidly reacts. Chlorine gas dissolves in water and forms chloride and hypochlorite; hypochlorite rapidly forms stable chloride or oxidizes organic matter. Thus, free Cl₂ in the environment is usually short-lived, but its effects can be severe (e.g. a chlorine gas leak into a river can kill fish and aquatic organisms immediately). Long-lived products of chlorine use include chloride ions (which then join the global chlorine cycle: rivers carry them to oceans, where they remain as NaCl).

A major environmental concern has been chlorine’s role in ozone depletion. In the stratosphere, chlorine atoms derived from chlorofluorocarbons (CFCs) catalyze the breakdown of ozone molecules. This led to international regulations (the Montreal Protocol) to phase out CFCs and reduce atmospheric chlorine radicals.

Safety: Chlorine gas is highly toxic and corrosive. It causes immediate pain and injury at low ppm levels. Short-term inhalation symptoms include a burning sensation in the nose/throat, coughing, chest tightness, and eye irritation. High exposures (tens to hundreds of ppm) can cause respiratory distress, pulmonary edema, and death. Occupational exposure limits are set very low: for example, OSHA’s permissible exposure limit (PEL) is 1 ppm (parts per million) averaged over 8 hours, and the recommended short-term exposure limit (STEL) by NIOSH is 0.5 ppm (15-minute average) with a 0.5 ppm ceiling. The IDLH (Immediately Dangerous to Life or Health) value is about 10 ppm. Chlorine concentration above 1000 ppm in air is often fatal very quickly. For comparison, a sharp chlorinous odor is noticeable around 0.2–1.0 ppm.

Skin or eye contact with liquid chlorine or concentrated solutions (like bleach) causes chemical burns. Inhaled chlorine reacts with moisture in the respiratory tract to form HCl and HOCl, which are corrosive to tissues.

For these reasons, industrial facilities handle chlorine with many precautions: gas leak detectors, proper ventilation, emergency neutralization (e.g. sodium bisulfite sprays), and protective gear. Cylinders of Cl₂ are stored in well-ventilated areas. Transport pipelines often have monitoring for leaks. In community water treatment, chlorine levels are kept low (e.g. 2–4 mg/L in drinking water) to avoid toxicity, and dechlorination is used if needed (e.g. treating waste water).

In environmental regulations, chlorinated disinfection byproducts (trihalomethanes, haloacetic acids from chlorinating organic-laden water) are controlled due to their health risks. Many chlorinated organic chemicals (dioxins, PCBs, certain solvents) are banned or limited because they persist and bioaccumulate.

In summary, chlorine must be handled with great care due to its acute toxicity and reactivity. Safety data sheets categorize Cl₂ as an immediate health hazard, and emergency protocols are in place in chemical plants to protect workers and the public. Nonetheless, when used responsibly, chlorine’s benefits in sanitation and chemistry are enormous.

History and Etymology

The name chlorine comes from the Greek word χλωρός (chloros) meaning “pale green,” reflecting the gas’s color. Chlorine gas was first recognized in 1774 by Swedish chemist Carl Wilhelm Scheele, who generated it by reacting hydrochloric acid with manganese dioxide. Scheele observed its distinctive properties but thought it contained oxygen (“dephlogisticated muriatic acid”). In 1809, it was suggested that this “acid air” might be a new element, and in 1810 Sir Humphry Davy confirmed it was indeed elemental chlorine. Davy proposed the name “chlorine” in 1810 based on its color. The chemical symbol “Cl” derives from this name.

Chloride salts (like NaCl) have been known and used since antiquity, but free chlorine gas was unknown until the late 18th century. In 1785, French chemist Claude Berthollet used chlorine in aqueous form for bleaching (“Eau de Javel”) and hypochlorite solutions; this began chlorine’s use in sanitation and textiles. Throughout the 19th century, chlorine became industrially important. The discovery of the Leblanc soda process (1790s) produced Cl₂ as a byproduct, spurring interest in chlorine’s uses. Large-scale electrolysis of brine (the chlor-alkali process) was developed in the late 19th century (notably by H. B. Smith and others).

During World War I, chlorine’s toxic properties were exploited militarily: German forces released chlorine gas against Allied troops at the Second Battle of Ypres in 1915, causing massive casualty and hysteria. This marked the first large-scale use of poison gas in warfare. The event brought public attention to chlorine’s lethal effects. On the other hand, chlorine’s medical uses were also noted around this time: André Labarraque and others had shown (in 1820s–1830s) that dilute NaOCl solutions were useful antiseptics. During WWI a stabilized sodium hypochlorite solution (Labarraque’s solution) was used for treating wounds (Dakin’s solution, 1916).

In the mid-20th century, the discovery of chlorofluorocarbons (CFCs) and their impact on ozone led to environmental laws such as the 1987 Montreal Protocol, greatly reducing chlorine use in refrigerants. Industrial practices evolved as well: older mercury-cell chlor-alkali plants were replaced by membrane-cell technology by the late 20th century to reduce mercury pollution.

Today, chlorine is one of the oldest known chemical elements (recognized in modern times in 1774) and one of the most utilized. Its name, coined by Davy over 200 years ago, persists in signifying that distinctive green gas which revolutionized chemistry, medicine, and technology.

Data Table

Property	Value
Symbol	Cl
Atomic number (Z)	17
Group/Period/Block	17 (halogen), Period 3, p-block
Standard atomic weight	35.45 u
Electron configuration	[Ne] 3s² 3p⁵
Valence electrons	7
Common oxidation states	–1; also +1, +3, +5, +7 (in oxyacids/oxyanions)
Electronegativity (Pauling)	3.16
Electron affinity	349 kJ/mol
First ionization energy	1251 kJ/mol (12.97 eV)
Atomic radius (covalent)	≈100 pm
Van der Waals radius	≈175 pm
Density (gas, STP)	3.2 kg/m³ (at 0 °C, 1 atm)
Melting point	–101.5 °C
Boiling point	–34.0 °C
Standard state (at 20 °C, 1 atm)	Gas (yellow-green)
CAS Registry Number	7782-50-5