Chlorine (Cl)

What is Chlorine?

Chlorine is a chemical element with the atomic number 17 and symbol Cl. On the periodic table, it is positioned between fluorine and bromine, making it the second-lightest of the halogens. The majority of its characteristics lie in the midway of the two.

At room temperature, chlorine is a gas with a yellow-green hue. It has the highest electron affinity and the third-highest electronegativity among the elements on the revised Pauling scale, behind only oxygen and fluorine. It is a very reactive and potent oxidizing agent.

In the experiments of the medieval alchemists, chlorine was a key ingredient. Typically, these experiments involved heating chloride salts such as sodium chloride (common salt) and ammonium chloride (sal ammoniac) to.

Produce a variety of chemical substances containing chlorine, including aqua regia, hydrogen chloride, and mercury(II) chloride (corrosive sublimate).

However, it wasn’t until Jan Baptist van Helmont discovered free chlorine gas’s distinct nature in the early 1630s.

In 1774, Carl Wilhelm Scheele described chlorine gas, speculating that it was an oxide of a recently discovered element. Chemists proposed in 1809 that the gas might be a pure element.

Sir Humphry Davy confirmed this in 1810, naming it after the Ancient Greek word χλωρός (khlōrós, “pale green”) due to its color.

All of the chlorine in the crust of the Earth is found in ionic chloride compounds, which include table salt, due to its high reactivity.

In the crust of the Earth, it is the twenty-first most prevalent chemical element and the second most abundant halogen (after fluorine).

Despite this, seawater’s enormous stores of chloride dwarf these crustal deposits. Commercial electrolysis of brine, mostly in the chlor-alkali process, yields elemental chlorine.

Due to elemental chlorine’s strong oxidizing potential, commercial bleaches and disinfectants as well as reagents for numerous chemical industry processes were developed.

A vast array of consumer goods, about two-thirds of which are organic compounds like polyvinyl chloride (PVC), several plastics production intermediates, and other final products that don’t have the component.

In order to maintain swimming pools hygienic, elemental chlorine and chlorine-generating compounds are more frequently utilized as disinfectants.

High concentrations of elemental chlorine are exceedingly harmful and deadly to the majority of living things.

Chlorine was originally employed as a chemical warfare agent as a poison gas weapon in World War I.

All known kinds of life require chlorine in the form of chloride ions. Living things rarely contain other forms of chlorine compounds, and synthetically created chlorinated organics can be harmful or inert.

Ozone depletion in the upper atmosphere has been linked to organic compounds containing chlorine, such as chlorofluorocarbons.

As part of the immune system’s defense against germs, neutrophils oxidize chloride ions, producing trace amounts of elemental chlorine.

History

Since ancient times, scientists have known that sodium chloride, the most common form of chlorine, has been around. Evidence from archaeologists suggests that rock salt and brine were utilized as early as 3000 BC and 6000 BC, respectively.

Early Discoveries

The alchemist and physician Abu Bakr al-Razi (c. 865–925, Latin: Rhazes) and the authors of the Arabic writings attributed to Jabir ibn Hayyan (Latin: Geber) were experimenting with sal ammoniac (ammonium chloride) around 900.

When this mixture was distilled with vitriol (hydrated sulfates of various metals), hydrogen chloride was the result.

However, it seems that the gaseous byproducts of these early chloride salt studies were thrown away, and hydrogen chloride might have been created instead.

Countless times until its chemical application was identified. One of the earliest applications for this material was the creation of mercury(II) chloride (corrosive sublimate), which was first described in the.

Arabic text De aluminibus et salibus (“On Alums and Salts”), which was mistakenly attributed to Abu Bakr al-Razi in the eleventh or twelfth century and translated into Latin by Gerard of Cremona in the latter half of the same century.

Mercury is heated with alum and ammonium chloride or vitriol and sodium chloride to produce mercury(II) chloride.1144–1187).

The finding by pseudo-Geber (in the De inventione veritatis, “On the Discovery of Truth,” after approximately 1300).

That ammonium chloride could be added to nitric acid to create a powerful solvent that could dissolve gold (aqua regia) was another significant breakthrough.

Despite the fact that aqua regia is an unstable mixture that continuously releases fumes including free chlorine gas.

It appears that this gas was overlooked until the Brabantian physician and scientist Jan Baptist van Helmont recognized it as a distinct gaseous substance in about 1630.

Isolation

Swedish chemist Carl Wilhelm Scheele is credited with making the discovery of the element after conducting the first thorough investigation of it in 1774 Scheele reacted MnO2 (as the mineral pyrolusite) with HCl to make chlorine:
4 HCl + MnO₂ → MnCl₂ + 2 H₂O + Cl₂
Scheele noted that chlorine has various characteristics, including the ability to bleach litmus, the ability to kill insects, a yellow-green hue, and an odor akin to aqua regia.

Since it is a gas (referred to as “airs”) and is derived from hydrochloric acid (referred to as “muriatic acid” at the time), he named it “dephlogisticated muriatic acid air”.

Chlorine was not proved to be an element by him. Many chemists, including Claude Berthollet, proposed that Scheele’s dephlogisticated muriatic acid air must be a mixture of oxygen and the as-yet-undiscovered element muriaticum.

At the time, the common chemical theory held that an acid is a compound that contains oxygen (remains of this survive in the German and Dutch names of oxygen: sauerstoff or zuurstof).

In an attempt to release the free element muriaticum (together with carbon dioxide) from dephlogisticated muriatic acid air, Joseph Louis Gay-Lussac and Louis-Jacques Thénard reacted with the air in 1809.

They failed and released a paper in which they explored the idea that air that has been dephlogisticated of muriatic acid might be an element, but they were not persuaded.

Sir Humphry Davy repeated the experiment in 1810 and came to the conclusion that the substance was an element rather than a compound.

On November 15 of that year, he informed the Royal Society of his findings. Based on its hue, he called this new element “chlorine” at the time, derived from the Greek word χλωρoς (chlōros, “green-yellow”).

Originally, the term “halogen,” which means “salt producer,” was used by Johann Salomo Christoph Schweigger in 1811.

Following a suggestion by Jöns Jakob Berzelius in 1826, this name was later employed as a generic phrase to represent all the elements of the chlorine family (fluorine, bromine, and iodine).

Michael Faraday first liquefied chlorine in 1823, proving that the substance that was then referred to as “solid chlorine” actually had a structure made of chlorine hydrate (Cl₂·H₂O).

Later Uses

French chemist Claude Berthollet utilized chlorine gas for the first time in 1785 to bleach textiles. Further research by Berthollet, who created sodium hypochlorite in 1789.

At his laboratory in the town of Javel (now a part of Paris, France) passing chlorine gas through a sodium carbonate solution, led to the development of modern bleaches.

The resultant beverage was a weak sodium hypochlorite solution and was dubbed “Eau de Javel” or “Javel water”.

Since this procedure was inefficient, other ways to produce goods were looked for. Charles Tennant, a Scottish chemist and manufacturer, created the first calcium hypochlorite solution.

“Chlorinated lime”), followed by solid hypochlorite (powder used for bleaching). These substances created small amounts of elemental chlorine and were more transportable than sodium hypochlorite.

Which stayed in diluted solutions because it turned into an extremely potent and unstable oxidant when it was dried out to remove water.

E. S. Smith received a patent around the close of the 1800s for a process that produced sodium hypochlorite by electrolyzing brine to produce sodium hydroxide and chlorine gas, which were then combined to create sodium hypochlorite.

This technique called the chloralkali process, was first used on an industrial scale in 1892 and is currently the main source of sodium hydroxide and elemental chlorine.

Another chloralkali technique was created in 1884 by the German company Chemischen Fabrik Griesheim, and it was put into commercial use in 1888.

Long before the germ theory of disease was established, in the 1820s, elemental chlorine solutions mixed in chemically basic water (calcium and sodium hypochlorite) were utilized as disinfectants and anti-putrefaction agents in France.

Antoine-Germain Labarraque invented this technique by adapting Berthollet’s “Javel water” bleach and other chlorine preparations.

Since then, the use of elemental chlorine has continued in public sanitation, especially in drinking and swimming water, and topical antisepsis (wound irrigation treatments and similar applications).

The German Army employed chlorine gas as a weapon for the first time on April 22, 1915, during the Second Battle of Ypres.

Due to their limited distribution and difficulty of usage, the existing gas masks had a catastrophic effect on the Allies.

Properties

Group 17 of the periodic table contains the nonmetal chlorine, which is the second halogen. Its characteristics are therefore intermediate between those of the first two and comparable to those of fluorine, bromine, and iodine.

The seven electrons in the third and outermost shell of chlorine, which have the electron configuration [Ne]3s²3p⁵, serve as its valence electrons.

Because it lacks an electron to form a complete octet, like all halogens, it is a potent oxidizer that must react with a variety of elements to complete its outer shell.

In line with recurring patterns, it has an electronegativity of 3.98, Cl: 3.16, Br: 2.96, I: 2.66, it lies in between fluorine and bromine, with less reactivity than fluorine and more reactivity than bromine.

Moreover, it is a stronger oxidizer than bromine but weaker than fluorine. On the other hand, compared to bromide, the chloride ion is a stronger reducing agent than fluoride.

Due to its intermediate atomic radius between fluorine and bromine, many of its atomic properties— such as electron affinity, enthalpy of dissociation of the X2 molecule (X = Cl, Br, I), ionic radius, and X–X bond length—continue the upward trend from iodine to bromine.

(Fluorine’s small size makes it abnormal.)Intermolecular van der Waals forces of attraction are present in all four stable halogens, and they become stronger the more electrons that each homonuclear diatomic halogen molecule has.

Chlorine melts at -101.0 °C and boils at -34.0 °C, making it an intermediate melting and boiling point between fluorine and bromine.

The density and heat of fusion and vaporization of chlorine are again intermediate between those of bromine and fluorine due to the halogens’ increasing molecular weights.

The group, however, all of their heats of vaporization are relatively low (resulting in high volatility) because of the structure of molecules in diatoms.

As the group descends, the halogens become darker in color; for example, chlorine is clearly yellow-green in color, but fluorine is a pale yellow gas.

This pattern emerges as a result of the halogens’ absorption of visible light at longer wavelengths along the group.

In particular, the electron transition between the lowest unoccupied antibonding σu molecular orbital and the highest occupied antibonding πg molecular orbital determines the color of a halogen, like chlorine.

Around low temperatures, the color of chlorine disappears, and around 195 °C, solid chlorine is nearly colorless.

Similar to solid iodine and bromine, solid chlorine forms a layered lattice of Cl₂ molecules in the orthorhombic crystal structure.

Compare the van der Waals radius of chlorine, 180 pm, with the Cl–Cl distance of 198 pm, which is close to the gaseous Cl–Cl distance of 199 pm.

The Cl···Cldistance between molecules is 332 pm within a layer and 382 pm between layers.

Because of its structure, chlorine is an extremely poor electrical conductor—in fact, its conductivity is so low as to be virtually undetectable.

Isotopes

There are two stable isotopes of chlorine: ³⁵Cl and ³⁷Cl. With ³⁵Cl making up 76% of natural chlorine and ³⁷Cl making up the remaining 24%, these are the.

Only two naturally occurring isotopes of the element. Both are produced in stars through the processes of burning silicon and oxygen.

Since both have nuclear spin 3/2+, they can be used for nuclear magnetic resonance; however, if the spin magnitude is larger than 1/2, the nuclear charge will not be spherical.

Dispersion due to a nonzero nuclear quadrupole moment and the ensuing quadrupolar relaxation, and hence resonance broadening.

The half-lives of the other chlorine isotopes are too short for them to exist in primordial nature, making them all radioactive.

Of them, ³⁶Cl (t1/2 = 3.0×105 y) and ³⁸Cl (t1/2 = 37.2 min) are most frequently utilized in the laboratory and can be created by neutron activation of natural chlorine.

³⁶Cl is the chlorine radioisotope that is most stable. Isotopes lighter than ³⁵Cl decay mostly by electron capture to sulfur isotopes; isotopes heavier than ³⁷Cl decay primarily by beta decay to argon isotopes.

And ³⁶Cl can decay by either mechanism to stable ³⁶S or ³⁶Ar. In nature, ³⁶Cl is found in trace amounts as a cosmogenic nuclide, with stable chlorine isotopes in a ratio of roughly 7–10 × 10−13 to 1.

It is created in the atmosphere through the spallation of ³⁶Ar by interactions with protons from cosmic rays.

The primary sources of ³⁶Cl in the upper meter of the lithosphere are the thermal neutron activation of ³⁵Cl and the spallation of ³⁹K and ⁴⁰Ca.

Muon uptake by ⁴⁰Ca becomes increasingly significant in the subsurface environment as a means of producing ³⁶Cl.

Chemistry and Compounds

One of the most reactive elements, chlorine is reactive in an intermediate way between fluorine and bromine.

Compared to fluorine, chlorine is a weaker oxidizing agent, but it is more potent than bromine or iodine.

The X2/X⁻ couples’ conventional electrode potentials (F, +2.866 V; Cl, +1.395 V; Br, +1.087 V; I, +0.615 V; At, around +0.3 V) show this.

However, because fluorine is solitary due to its tiny size, limited polarisability, and inability to demonstrate hypervalence, this trend is not visible in the bond energies.

Another distinction is that chlorine has a whereas fluorine does not exhibit significant chemistry in positive oxidation states.

Compared to bromination or iodination, chlorination frequently results in greater oxidation states; however, fluorination commonly produces lower oxidation states.

X	XX	HX	BX3	AlX3	CX4
F	159	574	645	582	456
Cl	243	428	444	427	327
Br	193	363	368	360	272
I	151	294	272	285	239

M–Cl bonds are typically formed when chlorine reacts with substances such as M–M, M–H, or M–C bonds. Since E°(1/₂ O₂/H₂O) = +1.229 V, which is lower than +1.395 V, chlorine should be able to oxidize water to produce hydrochloric acid and oxygen.

The electrolysis of aqueous chloride solutions, however, evolves chlorine gas rather than oxygen gas due to unfavorable reaction kinetics and the bubble overpotential effect.

which is highly helpful for the industrial synthesis of chlorine.

Hydrogen Chloride

Hydrogen chloride, or HCl, is the most basic form of chlorine and is a common chemical in both industry and laboratories.

It can be dissolved in water to form hydrochloric acid or as a gas. It is frequently created as a byproduct of chlorinating hydrocarbons or by burning hydrogen gas in chlorine gas.

Another method, referred to as the “salt-cake” technique, involves treating sodium chloride with strong sulfuric acid to make hydrochloric acid:
NaCl + H₂SO₄ 150 °C⟶  NaHSO₄ + HCl

NaCl + NaHSO₄ 540–600 °C⟶  Na₂SO₄ + HCl

By using concentrated sulfuric acid to dry the acid, hydrogen chloride gas can be produced in a lab. Benzoyl chloride and heavy water (D₂O) can react to form deuterium chloride or DCl.

Like all hydrogen halides except hydrogen fluoride, hydrogen chloride is a colorless gas at room temperature because hydrogen cannot form strong hydrogen bonds with the larger electronegative chlorine atom.

However, at low temperatures, solid crystalline hydrogen chloride exhibits weak hydrogen bonding, which is similar to the structure of hydrogen fluoride, before disorder sets in as the mercury rises.

Since there are insufficient hydrogen bonds between hydrogen and chlorine to prevent dissociation, hydrochloric acid has a severe acidity (pK_a = −7).

There are numerous hydrates in the HCl/H₂O system, HCl·nH₂O for n = 1, 2, 3, 4, and 6. After HCl and H₂O are mixed 1:1, the system fully splits into two distinct liquid phases.

At 20.22 g HCl per 100 g solution, hydrochloric acid forms an azeotrope with a boiling point of 108.58 °C; as a result, distillation is unable to concentrate hydrochloric acid beyond this point.

Hydrogen chloride is easier to work with as a solvent in anhydrous liquid form than in hydrogen fluoride because of its low boiling point, small liquid range, and low dielectric constant.

And lack of significant dissociation into H₂Cl⁺ and HCl⁻₂ ions. The latter is, however, much less stable than the fluoride ions (HF⁻₂) because of the extremely weak hydrogen bonding between hydrogen and chlorine.

Even though its salts have very large and weakly polarizing properties. It’s possible for cations like Cs⁺ and NR⁺ ₄ (R = Me, Et, Bun) to remain separated.

Anhydrous hydrogen chloride is a weak solvent; it can only dissolve very low lattice energy salts, like tetraalkylammonium halides, or tiny molecules, like phenol and nitrosyl chloride.

It easily protonates electrophiles with π bonds or lone pairs. In hydrogen chloride solution, solubility, ligand replacement processes, and oxidations are all well-characterized:

Ph₃SnCl + HCl ⟶ Ph₂SnCl₂ + PhH (solvolysis)
Ph₃COH + ₃ HCl ⟶ Ph₃C⁺HCl⁻₂ + H₃O⁺Cl⁻ (solvolysis)
Me₄N⁺HCl⁻₂ + BCl₃ ⟶ Me₄N⁺BCl⁻₄ + HCl (ligand replacement)
PCl₃ + Cl₂ + HCl ⟶ PCl⁺₄HCl⁻₂ (oxidation)

Other Binary Chlorides

Binary chlorides are formed by almost all elements in the periodic table. The rare exceptions are caused by one of three factors.

Extreme nuclear instability impeding chemical investigation prior to decay and transmutation (many of the heaviest elements); extreme inertness and reluctance to participate in chemical reactions (the noble gases.

With the exception of xenon in the highly unstable XeCl₂ and XeCl₄) above bismuth); and possessing electronegativity greater than that of chlorine (oxygen and fluorine).

So that the resulting binary compounds are technically oxides or fluorides of chlorine rather than chlorides.

Despite the negative charge on the nitrogen in NCl₃, the chemical is commonly referred to as nitrogen trichloride.

When various oxidation states are available, as in MoCl₅ and MoBr₃, chlorinating a metal with Cl₂ typically results in a higher oxidation state than brominating it with Br₂.

An element or its oxide, hydroxide, or carbonate can react with hydrochloric acid to produce chlorides, which are subsequently dehydrated at moderately high temperatures.

Mixed with anhydrous hydrogen chloride gas or low-pressure hydrogen chloride gas. If not, there are other options such as high-temperature oxidative.

Chlorination of the element with hydrogen or chlorine, high-temperature chlorination of a metal oxide or other halide by chlorine, volatile metal chloride, carbon tetrachloride, or organic chloride.

These methods are most effective when the chloride product is stable for hydrolysis. For example, under ordinary conditions, zirconium dioxide interacts with chlorine to form zirconium tetrachloride, while when heated, uranium trioxide reacts with hexachloropropene.

To produce uranium tetrachloride while refluxing. In the second example, a reduction in oxidation state is likewise possible by the use of hydrogen or metal as a reducing agent to lower a higher chloride.

As follows, thermal disproportionation or decomposition can also be used to accomplish this.

EuCl₃ + 1/₂ H₂ ⟶ EuCl₂ + HCl

ReCl₅ ^{at “bp”}_⟶ReCl₃ + Cl₂

AuCl₃ ^160 °C_⟶AuCl + Cl₂

The majority of metal chlorides are ionic when the metal is in low oxidation states (+1 to +3). Covalent molecular chlorides are typically formed by nonmetals and metals in high oxidation states (above +3).

For metals in oxidation state +3, both ionic and covalent chlorides are known (for example, aluminum chloride is not primarily ionic.

Whereas scandium chloride is). Since silver chloride is highly soluble in water, it is frequently employed as a qualitative indicator of chlorine.

Polychlorine compounds

Despite having a high initial ionization energy and being a potent oxidizer, dichlorine can undergo severe oxidation to produce the [Cl2]⁺ cation.

The best way to describe this is as an extremely unstable substance that can only be manufactured in a low-pressure discharge tube based on its electronic band spectrum. More stable than the other cation, yellow [Cl3]⁺ can be created in the following ways:

Cl₂ + ClF + AsF₅ −^{78 °C}_⟶[Cl₃]⁺[AsF₆]⁻

Arsenic pentafluoride, an oxidizing solvent, is used in this synthesis. Similar to triiodide, the trichloride anion, [Cl₃]⁻, has also been characterized.

Chlorine Fluorides

A subset of the interhalogen compounds, all of which are diamagnetic, are composed of the three fluorides of chlorine.

There are known cationic and anionic derivatives, including ClF⁻₂, ClF⁻₄, ClF⁺₂, and Cl₂F⁺.

Chlorine cyanate (ClNCO), cyanogen chloride (ClCN, linear), chlorine thiocyanate (ClSCN, unlike its oxygen cousin), and chlorine azide (ClN₃) are some other known pseudohalides of chlorine.

The commercial form of chlorine monofluoride (ClF) comes in 500-gram steel lecture bottles and is incredibly thermally stable.

This colorless gas boils at -100.1 °C and melts at – 155.6 °C. It can be created by the elements reacting at 225 °C, but first, it needs to be isolated and purified from the reactants and chlorine trifluoride.

Its characteristics are primarily in the middle range of those of fluorine and chlorine. At room temperature and higher, it will react with a variety of metals and nonmetals to fluorinate them and release chlorine.

Additionally, it will add chlorine and fluorine as a chlorofluorinating agent. Fluorine through oxidation or a multiple bond; for instance.

It can attack carbon monoxide to create COFCl or carbonyl chlorofluoride. It will react similarly to hexafluoroacetone (CF₃)₂CO, nitriles RCN (RCF₂NCl₂), sulfur oxides SO₂ and SO₃ (ClSO₂F and ClOSO₂F, respectively.

And heptafluoroisopropyl hypochlorite (CF₃)₂CFOCl when combined with a potassium fluoride catalyst.

Additionally, it will react exothermically with substances like water that possess -OH and -NH groups:
H₂O + 2 ClF ⟶ 2 HF + Cl₂O
The volatile colorless liquid known as chlorine trifluoride (ClF₃) melts at -76.3 °C and boils at 11.8 °C.

Direct fluorination of gaseous chlorine or chlorine monofluoride at 200–300 °C can create it. One of the most reactive substances ever discovered, it ignites a wide range of elements.

Including powdered molybdenum, tungsten, rhodium, iridium, and iron, as well as hydrogen, potassium, phosphorus, arsenic, antimony, sulfur, selenium, tellurium, bromine, and iodine.

It will also set fire to water and a variety of materials that are normally thought to be chemically inert, like sand, concrete, asbestos, and glass.

Even noble metals like palladium, platinum, and gold will corrode when heated, and even the noble gases xenon and radon do not escape fluorination.

An impenetrable fluoride layer is generated by sodium, magnesium, aluminum, zinc, tin, and silver, which may be removed by heating.

Because an unreactive layer of metal fluoride forms on nickel, copper, and steel containers, they are typically utilized because of their exceptional resistance to attack by chlorine trifluoride.

It was utilized in experimental rocket engines to create hydrogen fluoride, nitrogen, and chlorine gasses through its reaction with hydrazine; nevertheless, the main cause of its troubles is that it ignites spontaneously due to its severe hypergolicity.

These days, it is primarily utilized in the manufacturing of nuclear fuel, where it is employed to clean.

Chemical vapor deposition chambers and oxidize uranium to uranium hexafluoride for enrichment and separation from plutonium.

Despite not significantly dissociating into ClF⁺₂ and ClF^–₄ ions, it can function as a donor or acceptor of fluoride ions (Lewis base or acid).

Large-scale production of chlorine pentafluoride (ClF₅) involves direct fluorination of chlorine with excess fluorine gas at 350 °C and 250 atm.

Small-scale production of ClF₅ involves the reaction of metal chlorides with fluorine gas at 100–300 °C.

It boils at -13.1 °C and melts at -103 °C. Even though it is still not as powerful as chlorine trifluoride, it is a fairly potent fluorinating agent.

Merely a few number of distinct stoichiometric reactions have been described. Water interacts violently with arsenic pentafluoride and antimony pentafluoride to generate ionic adducts of the form [ClF₄]⁺[MF₆]⁻ (M = As, Sb):

2 H₂O + ClF₅ ⟶ 4 HF + FClO₂

One of the five known forms of chlorine oxide fluoride is the product, chloryl fluoride. The other three are FClO₂, F₃ClO, and F₃ClO₂.

They vary from the chemically unreactive perchloric fluoride (FClO₃) to the thermally unstable FClO.

All five exhibit comparable structural and chemical properties to chlorine fluorides.

They can function as bases or Lewis acids by accumulating or shedding fluoride ions, respectively, or as potent oxidizing and fluorinating agents.

Chlorine Oxides

Despite being unstable, all of the chlorine oxides have been thoroughly researched since they are endothermic chemicals.

They are significant because they are created when the ozone layer is destroyed by the photolysis of chlorofluorocarbons in the upper atmosphere.

The elements cannot be directly reacted to create any of them.Cl₂O, or dichlorine monoxide, is a brownish-yellow gas that can be reddish-brown whether solid or liquid.

It is produced when yellow mercury(II) oxide reacts with chlorine gas. It is very soluble in water, where it forms anhydride with hypochlorous acid (HOCl) and is in equilibrium with the latter.

As a result, it works well as a bleach and is usually utilized to create hypochlorites. When heated, ignited, or exposed to ammonia gas, it bursts.

Humphry Davy made the first chlorine oxide discovery in 1811 with the discovery of chlorine dioxide (ClO₂). As may be expected from its odd number of electrons, it is a yellow paramagnetic gas (deep-red as a solid or liquid).

Because of the delocalization of the unpaired electron, it is stable towards dimerization. For wood-pulp bleaching and water treatment.

It must be manufactured at low concentrations because it bursts as a liquid above -40°C and as a gas under pressure. In order to prepare it, a chlorate is commonly reduced as follows:
ClO⁻₃ + Cl⁻ + 2 H⁺ ⟶ ClO₂ + ¹/₂Cl₂ + H₂O

As a result, its synthesis is closely related to the redox processes of the chlorine oxoacids. It reacts with sulfur, phosphorus, phosphorus halides, and potassium borohydride, making it a potent oxidizing agent.

It dissipates exothermically in water to create dark green solutions that break down extremely slowly at night.

At low temperatures, crystalline clathrate hydrates ClO₂·nH₂O (n = 6–10) separate off.

Nevertheless, when there is light, these solutions rapidly photodecompose to generate a combination of chloric and hydrochloric acids.

While mostly chlorine, oxygen, and a small amount of ClO₃ and Cl₂O₆ are produced at room temperature.

Photolysis of individual ClO₂ molecules produces the radicals ClO and ClOO. Cl₂O₃, a dark brown solid that explodes below 0°C, is also produced when the solid is photolyzed at -78°C.

Because it causes the ozone layer in the atmosphere to thin, the ClO radical is significant for the environment in the following ways:
Cl• + O₃ ⟶ ClO• + O₂
ClO• + O• ⟶ Cl• + O₂

At ambient temperature, chlorine perchlorate (ClOClO₃), a pale yellow liquid that is less stable than ClO₂, breaks down into oxygen, chlorine, and dichlorine hexoxide (Cl₂O₆).

Similar to the thermally unstable chlorine derivatives of other oxoacids, such as chlorine nitrate (ClONO₂, highly reactive and explosive).

Chlorine fluorosulfate (ClOSO₂F, more stable but still highly reactive to moisture), and chlorine perchlorate can also be thought of as a chlorine derivative of perchloric acid (HOClO₃).

Dark crimson liquid dichlorine hexoxide freezes to produce a solid that then becomes yellow at -180°C: often produced by the interaction of oxygen with chlorine dioxide.

It reacts more like chloryl perchlorate, [ClO₂]⁺[ClO₄]⁻, which has been proven to be the right structure of the solid, despite attempts to rationalize it as the dimer of ClO₃.

In water, it hydrolyzes to produce a combination of perchloric and chloric acids; the corresponding reaction with anhydrous hydrogen fluoride is not completed.

The anhydride of perchloric acid (HClO4), dichlorine heptoxide (Cl2O₇), is easily extracted from it by dehydrating it with phosphoric acid at -10°C and then distilling the resultant mixture at -35°C and 1 mmHg.

It is an oily, colorless liquid that is susceptible to shock. Because it is the only chlorine oxide that does not ignite organic compounds at room temperature, it is the least reactive of the group.

It can be dissolved in aqueous alkalis to regenerate perchlorates or in water to regenerate perchloric acid.

On the other hand, it undergoes an explosive thermal breakdown when one of the core Cl–O bonds is broken, resulting in the radicals ClO₃ and ClO₄, which then instantly break down into the elements via intermediary oxides.

Chlorine Oxoacids and Oxyanions

Four oxoacids are formed by chlorine: perchloric acid (HOClO₃), hypochlorous acid (HOCl), chlorous acid (HOClO₂), and chlorous acid (HOClO).

The redox potentials shown in the following table demonstrate that chlorine is far more stable in acidic solutions than it is in alkaline solutions with regard to disproportionation:

Cl₂ + H₂O ⇌ HOCl + H⁺ + Cl⁻ K_ac = 4.2 × 10⁻⁴ mol² l−2
Cl₂ + 2 OH⁻ ⇌ OCl⁻ + H₂O + Cl− K_alk = 7.5 × 10¹⁵ mol⁻¹ l

Despite the relatively favorable equilibrium constant of 1027, the hypochlorite ions also disproportionately create chloride and chlorate (3 ClO⁻ ⇌ 2 Cl− + ClO⁻ ₃).

However, this reaction is quite sluggish at temperatures below 70 °C.

Even at 100 °C, the formation of chloride and perchlorate (4 ClO⁻₃ ⇌ Cl⁻ + 3 ClO⁻₄) is still extremely sluggish, even though the chlorate ions may themselves be disproportionately 1020 is the favorable equilibrium constant.

As chlorine’s oxidation state drops, so do the rates of reaction for the chlorine oxyanions.

Because of the rising delocalization of charge over an increasing number of oxygen atoms in their conjugate bases.

E°(couple)	a(H+) = 1 (acid)	E°(couple)	a(OH−) = 1 (base)
Cl2/Cl−	+1.358	Cl2/Cl−	+1.358
HOCl/Cl−	+1.484	ClO−/Cl−	+0.890
ClO−3/Cl−	+1.459
HOCl/Cl2	+1.630	ClO−/Cl2	+0.421
HClO2/Cl2	+1.659
ClO−3/Cl2	+1.468
ClO−4/Cl2	+1.277
HClO2/HOCl	+1.701	ClO−2/ClO−	+0.681
		ClO−3/ClO−	+0.488
ClO−3/HClO2	+1.181	ClO−3/ClO−2	+0.295
ClO−4/ClO−3	+1.201	ClO−4/ClO−3	+0.374

The strengths of the chlorine oxyacids increase very quickly as the oxidation state of chlorine increases. Utilizing these disproportionation processes can lead to the production of the majority of chlorine oxoacids.

Because of its strong reactivity and instability, hypochlorous acid (HOCl) and its salts are mostly utilized for their bleaching and sterilizing properties.

They transmit an oxygen atom to the majority of inorganic species, making them extremely potent oxidizers.

As demonstrated by the breakdown of aqueous chlorine dioxide, chlorous acid (HOClO) is considerably more unstable and cannot be isolated or concentrated without disintegration.

But sodium chlorite, a stable salt, works well as an oxidizing agent, a source of chlorine dioxide, and a bleaching and stripping agent for textiles.

The strong acid chloric acid (HOClO2) is quite stable up to a 30% concentration in cold water but releases chlorine and chlorine dioxide when heated.

It can be further concentrated to about 40% by evaporating under low pressure, but beyond that, it breaks down into perchloric acid, chlorine, oxygen, water, and chlorine dioxide.

Sodium chlorate is its most important salt; it is usually utilized to create chlorine dioxide, which is used to bleach paper pulp.

One common method for producing oxygen on a small scale in the laboratory is the breakdown of chlorate into chloride and oxygen. In the following ways, chloride and chlorate can combine to generate chlorine:

ClO⁻₃ + 5 Cl⁻ + 6 H⁺ ⟶ 3 Cl₂ + 3 H₂O

Given that chlorine compounds are most stable when the chlorine atom is in its lowest (−1) or highest (+7) conceivable oxidation states, perchlorates and perchloric acid (HOClO₃) are the most stable oxo-compounds of chlorine.

Due to the high activation energies for these processes for kinetic reasons, perchloric acid and aqueous perchlorates are mostly inert at ambient temperature, but when heated, they become strong and occasionally violent oxidizing agents.

Electrolytically oxidizing sodium chlorate yields perchlorates, while concentrated hydrochloric acid reacts with anhydrous sodium or barium perchlorate to produce perchloric acid.

The filtrate is then distilled after the precipitated chloride is removed.to focus it. Anhydrous perchloric acid is a colorless.

The shock-sensitive liquid that oxidizes silver and gold, ignites hydrogen iodide and thionyl chloride and bursts upon coming into contact with most organic molecules.

While it is a weak ligand, even weaker than water, there are known to be a few compounds incorporating coordinated ClO^–₄.

The typical oxidation states for the element chlorine as listed in secondary schools or universities.

In any case, it should be mentioned in university chemistry courses that there are more complex chemical compounds, such as cluster technetium chloride [(CH₃)₄N]₃[Tc₆Cl₁₄], in which 6 of the 14 chlorine atoms are formally divalent and the oxidation states are fractional.

These compounds’ structures can only be understood using contemporary quantum chemical techniques.

Furthermore, all of the aforementioned chemical regularities apply to “normal” or almost normal circumstances.

Chlorine can display an oxidation state of -3 and produce the non-traditional chemical compound Na₃Cl with sodium at extremely high pressures (such as those found in the cores of giant planets).

Chlorine oxidation state	−1	+1	+3	+5	+7
Name	chloride	hypochlorite	chlorite	chlorate	perchlorate
Formula	Cl−	ClO−	ClO−2	ClO−3	ClO−4
Structure

Organochlorine Compounds

The C–Cl bond is a typical functional group that is a component of fundamental organic chemistry, just like the other carbon–halogen bonds.

Formally, substances having this functional group could be regarded as organic chloride anion derivatives.

Owing to the electronegativity differential between carbon (2.55) and chlorine (3.16), Because it lacks electrons, carbon in a C–Cl bond is electrophilic.

The physical characteristics of hydrocarbons are altered by chlorination in a number of ways. For example.

Because of the increased atomic weight of chlorine relative to hydrogen, chlorocarbons are often denser than water, and since chloride is a leaving group, aliphatic organochlorides act as alkylating agents.

Alkanes and aryl alkanes can be chlorinated in the presence of UV radiation and free radicals.

Though this would be acceptable if the products are easily separated, the degree of chlorination is difficult to control because the reaction is not regioselective and frequently produces a mixture of different isomers with varying degrees of chlorination.

The Friedel-Crafts halogenation process, which uses chlorine and a Lewis acid catalyst, can be used to create amyl chlorides.

Methyl ketones and similar chemicals can also be converted into alkyl halides by the haloform process, which uses sodium hydroxide and chlorine.

Chlorine adds to the multiple bonds on alkenes and alkynes as well, producing molecules that are di- or tetrachloro.

However, because chlorine is expensive and reactive, it is more frequently used to make organochlorine compounds using hydrogen chloride or other chlorinating agents such as thionyl chloride (SOCl₂) or phosphorus pentachloride (PCl₅).

The last is particularly practical for use in laboratories since it requires no distillation because all of the byproducts are gaseous.

Numerous organochlorine chemicals have been identified in natural sources, including humans and microbes.

Nearly all classes of biomolecules, including fatty acids, alkaloids, terpenes, amino acids, flavonoids, and steroids, contain chlorinated organic compounds.

Dioxins are among the organochlorides that are created in the high temperatures of forest fires; dioxins have been discovered in the conserved ashes of fires that were started by lightning and occurred before synthesized dioxins.

Furthermore, a range of basic chlorinated hydrocarbons, such as carbon tetrachloride, dichloromethane, and chloroform, have been extracted from marine algae.

Volcanoes, forest fires, and biological breakdowns produce most of the chloromethane found in the environment.

Not all organochlorides are dangerous to humans or other animals, but some are extremely harmful to both.

When organic matter burns in the presence of chlorine, dioxins are created. Other persistent organic pollutants, like DDT, can also be dangerous if released into the environment.

For instance, DDT, which was extensively used to control insects in the middle of the 20th century, builds up in food chains and affects certain bird species’ ability to reproduce (e.g., by weakening their eggshells).

Because of the damage they cause to the ozone layer, chlorofluorocarbons have been phased out because of the easy homolytic fission of the C–Cl bond, which produces chlorine radicals in the high atmosphere.

Occurrence and Production

Although chlorine is too reactive to exist in nature as a free element, it is quite common in the form of its chloride salts.

Made composed of 126 parts per million, it is the twenty-first most prevalent element in the crust of Earth.

This is due to the vast concentrations of chloride minerals, particularly sodium chloride, that have evaporated from water bodies.

These are nothing compared to the stores of chloride ions in seawater: some inland seas and subterranean brine wells, such as.

The Great Salt Lake in Utah and the Dead Sea in Israel, contain smaller amounts of these ions at higher concentrations.

Since manganese dioxide and hydrochloric acid are readily available, small amounts of chlorine gas can be produced in the lab.

However, this is not often needed. In the industrial setting, electrolysis of dissolved sodium chloride in water is typically used to produce elemental chlorine.

Today, most industrial chlorine gas is produced using this process, the chloralkali process, which was industrialized in 1892.

The process also produces sodium hydroxide, which is the most valuable product, and hydrogen gas. The following chemical equation describes the process as it moves forward:
2 NaCl + 2 H₂O → Cl₂ + H₂ + 2 NaOH

The following equations govern how chloride solutions are electrolyzed in each case:

Cathode: 2 H₂O + 2 e⁻ → H₂ + 2 OH⁻
Anode: 2 Cl⁻ → Cl₂ + 2 e⁻

An asbestos (or polymer-fiber) diaphragm divides the cathode from the anode in diaphragm cell electrolysis, preventing the hydrogen and sodium hydroxide generated at the cathode from recombining with the chlorine generated at the anode.

Brine, or salt solution, is continuously supplied to the anode compartment and then travels to the cathode compartment via the diaphragm, where it partially depletes and produces caustic alkali.

The alkali produced using diaphragm methods are diluted and slightly impure, but they use less energy and do not have to deal with the disposal of mercury.

The permeable membrane is used as an ion exchanger in membrane cell electrolysis. After passing through the anode compartment, saturated sodium (or potassium) chloride solution exits at a reduced concentration.

Although this process has the drawback of requiring highly pure brine at high concentrations, it nevertheless yields very pure sodium (or potassium) hydroxide.

Chlorine is recovered from hydrogen chloride during the Deacon process, which is the process by which organochlorine chemicals are produced. The method depends on oxidation with oxygen:

4 HCl + O₂ → 2 Cl₂ + 2 H₂O

A catalyst is needed for the reaction. The first catalysts, as described by Deacon, were based on copper.

Commercial methods have shifted to catalysts based on ruthenium and chromium, such as the Mitsui MT-Chlorine Process.

The chlorine generated is available in cylinders from sizes ranging from 450 g to 70 kg, as well as drums (865 kg), tank wagons (15 tonnes on highways; 27–90 tonnes by rail), and barges (600–1200 tonnes).

Applications

Sodium chloride is the most prevalent chlorine compound and is the main source of chlorine for the demand of the chemical sector.

Approximately fifteen thousand compounds containing chlorine are traded commercially. These include a wide range of compounds such as ethanes.

Vinyl chloride, polyvinyl chloride (PVC), magnesium trichloride for catalysis, and the chlorides of hafnium, zirconium, titanium, and magnesium, are used as starting points to produce the element in its pure form.

Approximately 63% of all elemental chlorine produced is utilized in the production of organic compounds, while the remaining 18% is utilized in the production of inorganic chlorine compounds.

There are about 15,000 chlorine compounds in commercial use. Bleaches and disinfection products are made with the remaining 19% of chlorine generated.

Vinyl chloride and 1,2- dichloroethane, which an intermediates in the synthesis of PVC, are the most important organic compounds in terms of manufacturing volume.

Vinylidene, vinyl chloride, methylene chloride, and methyl chloride are further notably significant organochlorines. chloride, dichlorobenzenes, epichlorohydrin, allyl chloride, trichloroethylene, perchloroethylene, and chlorobenzene.

HCl, Cl₂O, HOCl, NaClO₃, chlorinated isocyanurates, AlCl₃, SiCl₄, SnCl₄, PCl₃, PCl₅, POCl₃, AsCl₃, SbCl₃, SbCl₅, BiCl₃, and ZnCl₂ are among the principal inorganic compounds.

Sanitation, Disinfection, and Antisepsis

Combating Putrefaction

Animal intestines were processed in France, as well as other countries, to create items such as Goldbeater’s skin and strings for musical instruments.

The procedure was foul-smelling and unhygienic, and it was carried out in “gut factories” (boyauderies).

Around 1820, the Société d’encouragement pour l’industrie nationale offered a prize for the development of a mechanical or chemical process that could separate animal intestine peritoneal membranes without causing putrefaction.

The winner of the prize was 44-year-old French chemist and pharmacist Antoine-Germain Labarraque, who had discovered Berthollet’s chlorinated bleaching solutions.

Known as “Eau de Javel,” not only eliminated the putrefaction-causing stench of animal tissue decomposition but also slowed down the process.

As a result of Labarraque’s study, the boyauderies now utilize chlorides and hypochlorites of sodium (sodium hypochlorite) and lime (calcium hypochlorite).

The regular cleaning and deodorization of restrooms, sewers, marketplaces, slaughterhouses, anatomical theaters, and morgues were shown to benefit from the same chemicals.

They were helpful during exhumations, embalming, epidemic illness outbreaks, fever, and blackleg in cattle.

They were successful in hospitals, lazarets, jails, infirmaries (both on land and at sea), magnaneries, stables, cow sheds, etc.

Disinfection

Since 1828, Labarraque’s chlorinated lime and soda solutions have been recommended for the treatment of putrefaction of existing wounds, particularly septic wounds, and for.

The prevention of infection (referred to as “contagious infection,” which is thought to be spread by “miasmas”).

In his 1828 publication, Labarraque advised physicians to treat “contagious infections” by breathing chlorine, cleaning their hands with chlorinated lime, and even sprinkling chlorinated lime around the patients’ beds.

Even though the source of the infection was not identified until almost fifty years later, the spread of diseases was well established by 1828.

Large amounts of so-called chloride of lime were used to sanitize the metropolis during the 1832 Paris cholera outbreak.

In order to create calcium hypochlorite or chlorinated lime, chlorine gas was dissolved in lime water, a diluted form of calcium hydroxide.

This was not just modern calcium chloride. Labarraque’s finding helped to remove the awful stink of decomposition from hospitals and dissecting rooms, and by doing so, effectively deodorised the Latin Quarter of Paris.

Many believed that these “putrid miasmas” were what caused “contagion” and “infection” to spread; these terms were used prior to the germ theory of infection.

Lime chloride was used to eliminate smells and “putrid matter”. One account claims chloride of lime was employed by.

Dr. John Snow to purify water from the cholera-contaminated well that was feeding the Broad Street pump in 1854 London, yet three other trustworthy sources that chronicle that famous cholera pandemic do not mention the episode.

It is evident from one source that the offal and dirt in the streets around the Broad Street pump were cleaned with lime chloride, a typical procedure in England in the middle of the 1800s.

Semmelweis and Experiments with Antisepsis

The most well-known use of Labarraque’s chemical base and chlorine solutions dates back to 1847 when Ignaz Semmelweis used chlorine water —a cheaper alternative to chlorinated lime solutions—to disinfect the hands of Austrian physicians.

Semmelweis observed that the doctors’ hands continued to carry the stench of decomposition from the dissection rooms to the patient examination rooms.

Semmelweis utilized the well-known “Labarraque’s” to argue, long before the germ theory of disease, that “cadaveric particles” were spreading rot from recently deceased medical cadavers to living patients.

solutions”—which he discovered soap did not—as the only known way to get rid of the stench of decay and tissue degradation.

The solutions turned out to be far more potent antiseptics than soap (Semmelweis knew this, too, but not why), and this led to his well-known achievement.

In 1847 of halting the spread of childbed fever, also known as “puerperal fever,” in the maternity wards of Vienna General Hospital in Austria.

Much later, in 1916, during World War I, Henry Drysdale Dakin (who gave full credit to Labarraque’s earlier work in this area) created a standardized and diluted variant of Labarraque’s solution incorporating hypochlorite (0.5%) and boric acid as an acidic stabilizer.

Long before the development of current antibiotics, a technique known as Dakin’s solution—wound irrigation with chlorinated solutions—permitted the antiseptic treatment of a wide range of open wounds.

Even today, a modified form of this solution is used for wound irrigation since it still works well against germs that are resistant to several antibiotics (see Century Pharmaceuticals).

Public Sanitation

Jersey City, New Jersey, established the nation’s first continuous chlorination system for drinking water in 1908.

The US Department of Treasury demanded that chlorine be used to sterilize all drinking water by 1918.

These days, chlorine is a crucial ingredient for bleach, disinfectants, and water purification (such as in water treatment plants).

These days, even tiny water sources get regular chlorination. Common uses for chlorine (as hypochlorous acid) include disinfecting public swimming pools and drinking water sources of bacteria and other microorganisms.

Most private swimming pools use sodium hypochlorite, which is made of sodium hydroxide and chlorine, or solid tablets of chlorinated isocyanurates instead of chlorine itself.

The disadvantage of utilizing chlorine in swimming pools is that it reacts with human skin and hair proteins’ amino acids.

The characteristic “chlorine aroma” connected to swimming pools is not what people think.

Is actually the product of chloramine, a chemical molecule created when free dissolved chlorine reacts with amines found in organic materials, such as perspiration and urine, rather than elemental chlorine alone.

Chlorine is more than three times and more than six times more efficient than bromine and iodine as a water disinfectant against Escherichia coli.

A growing number of drinking water systems are chlorinating their water by directly adding monochloramine. This technique is called chlorination.

Alternative methods are employed to add chlorine because it is frequently impractical to store and use toxic chlorine gas for water treatment.

These include substances such as sodium dichloro-s-triazinetrione (dihydrate or anhydrous), sometimes referred to as “dichlor,” and trichloro-s-triazinetrione, sometimes referred to as “trichlor.”

Hypochlorite solutions also gradually release chlorine into the water. These substances can be employed as tablets, powders, or granules since they remain stable when solid.

Small amounts of chlorine are added to industrial water systems or swimming pools to cause the chlorine atoms to hydrolyze and generate hypochlorous acid (HOCl), a general biocide that destroys bacteria, microbes, algae, and other microorganisms.

Use As a Weapon

World War I

Germany employed chlorine gas, sometimes referred to as batholith, as a weapon for the first time during World War I on April 22, 1915, during the Second Battle of Ypres.

The soldiers characterized its unique scent as a cross between pineapple and pepper.[Reference required] It stung the back of my throat and chest and tasted metallic as well.

In the lung mucosa, chlorine combines with water to create hydrochloric acid, which is poisonous and potentially fatal to living things.

Chlorine gas is substantially less deadly when it is kept from human respiratory systems by gas masks with activated charcoal or other filters.

Compared to alternative chemical weapons. It was invented by Fritz Haber, a German chemist who would go on to win the Nobel Prize, from the Kaiser Wilhelm Institute in Berlin.

They worked together with IG Farben, a German chemical company, to create techniques for firing chlorine gas against a well-established foe.

Chlorine was utilized as a chemical weapon by both sides of the fight when it was first deployed, but mustard gas and phosgene, which are more lethal, quickly supplanted it.

Middle East

Insurgents in Anbar Province, Iraq, in 2007 also employed chlorine gas when they packed truck bombs with mortar shells and chlorine tanks.

Over 350 people were made ill and two people were killed by explosives during the attacks.

Since the hazardous gas is easily spread and diluted in the environment by the blast, the majority of the deaths were caused by the force of the explosions rather than the effects of chlorine.

More than a hundred civilians were hospitalized in some bombings because they were having respiratory problems.

The security surrounding elemental chlorine, which is necessary to supply the populace with clean drinking water, was tightened by the Iraqi authorities.

The village of Duluiyah, Iraq, was the site of a suspected chlorine gas attack by the Islamic State of Iraq and the Levant on October 23, 2014.

A laboratory investigation of soil and clothing samples verified that Kurdish Peshmerga Forces had been attacked with a vehicle-borne IED on January 23, 2015, at the Highway 47 Kiske Junction close to Mosul.

Chlorine gas had been used against them. Syria, another Middle Eastern nation, has employed chlorine as a chemical weapon delivered by rockets and barrel bombs.

The OPCW-UN Joint Investigative Mechanism came to the conclusion in 2016 that the Syrian government had carried out three independent chemical weapons strikes using chlorine.

Subsequent analyses by the OPCW’s Identification and Investigation Team determined that the attacks in 2017 and 2018 that used chlorine were carried out by the Syrian Air Force.

Biological Role

One nutrient that is necessary for metabolism is the chloride anion. Both the stomach’s production of hydrochloric acid and the operation of the cellular pump depend on chlorine.

Sodium chloride, or table salt, is the primary dietary source. An excess of chloride in the blood or a deficiency in it are instances of electrolyte imbalances.

Seldom does hypochloremia, or having too little chloride, happen on its own when there are no other anomalies. It occasionally has a connection to underbreathing.

It may be connected to long-term respiratory acidosis. Having too much chloride, or hyperchloremia, typically doesn’t cause any symptoms.

When symptoms do appear, they usually mimic those of hypernatremia or an excess of salt in the blood.

Reduced blood chloride causes cerebral dehydration, which is typically followed by rapid rehydration that causes cerebral edema. Oxygen transport may be impacted by hyperchloremia.

Hazards

Toxic gases like chlorine harm the skin, eyes, and respiratory system. It tends to gather at the bottom of poorly ventilated areas because it is denser than air.

Strong oxidizers like chlorine gas have the potential to react with combustible compounds. Measurement tools can identify chlorine at as low as 0.2 parts per million (ppm), and at 3 ppm, it can be detected by scent.

At 30 ppm, there may be vomiting and coughing, and at 60 ppm, lung damage. After taking a few deep breaths of the gas, exposure to approximately 1000 parts per million can be lethal.

Ten parts per million is the IDLH (immediately dangerous to life and health) concentration. Lower quantities can irritate the respiratory system when breathed in.

And the eyes may become irritated by the gas exposure. Hydrochloric acid (HCl) and hypochlorous acid (HOCl) are produced in the lungs when chlorine is breathed at amounts higher than 30 ppm.

The reaction of chlorine with water is not a serious health risk when used at defined levels for water disinfection.

Additional substances in the water could produce disinfection byproducts that are linked to harmful consequences for human health.

The Occupational Safety and Health Administration (OSHA) in the US has established a 1 ppm, or 3 mg/m3, allowable exposure limit for elemental chlorine.

The recommended exposure limit, set by the National Institute for Occupational Safety and Health, is 0.5 parts per million during a period of 15 minutes.

Accidents happen in homes when acidic drain cleansers and hypochlorite bleach solutions react to release chlorine gas.

Another hazardous class of compounds known as chloramines is created when ammonia, another common laundry ingredient, and hypochlorite bleach are mixed together.

Chlorine-Induced Cracking in Structural Materials

Chlorine is frequently used to disinfect water, particularly pool and drinkable water sources. Stress corrosion cracking of stainless steel suspension rods caused by chlorine has resulted in some disastrous collapses of swimming pool ceilings.

Certain polymers, such as polybutene and acetal resin, are also vulnerable to assault. In the US in the 1980s and 1990s, stress corrosion cracking caused numerous failures in both materials, which were utilized in hot and cold water home plumbing.

Chlorine-Iron Fire

A powerful exothermic reaction between the element iron and chlorine at high temperatures can result in a chlorine-iron fire.

Because steel makes up a large portion of the piping used to transport chlorine gas, chemical processing facilities are susceptible to chlorine-iron fires.

FAQ