関: Cl、Cl2

WordNet

a common nonmetallic element belonging to the halogens; best known as a heavy yellow irritating toxic gas; used to purify water and as a bleaching agent and disinfectant; occurs naturally only as a salt (as in sea water) (同)Cl, atomic number 17
an explosive gas (ClO2) used chiefly in bleaching paper or starch or soap or flour and in water purification
an aqueous solution of chlorine used as a bleaching agent
a radioactive isotope of chlorine

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塩素(化学記号は『CL』)
chlorineの化学記号

Wikipedia preview

出典(authority):フリー百科事典『ウィキペディア（Wikipedia）』「2014/05/08 22:15:44」(JST)

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[Wiki ja表示]

塩素（えんそ、英: chlorine）は原子番号17の元素。元素記号は Cl。ハロゲン元素の一つ。

一般に「塩素」という場合は、塩素の単体である塩素分子（Cl₂、二塩素、塩素ガス）を示すことが多い。ここでも合わせて述べる。塩素分子は常温常圧では特有の臭いを持つ黄緑色の気体で、毒性と腐食性を持つ。

性質

塩素原子の電子親和力は非常に大きく、通常イオン化する際は1価の陰イオンとなる。EA = 3.617 eV^[2]

単体（塩素ガス）は、常温常圧では特有の臭いを有する黄緑色の気体。融点-101 °C、沸点-34.1 °C、比重2.49。非常に反応性が高く、多くの金属や有機物と反応し塩化物を形成する。

強い漂白・殺菌作用をもつため、パルプや衣類の漂白剤や、水道水やプールの殺菌剤として使用される。ただし、気体を扱うのは困難であり、また保存性の点から水酸化ナトリウム水溶液と反応させた次亜塩素酸ナトリウムの形で利用されることが多い^[3]。

地球上の塩素の存在

地球上において、92ある天然元素のうち18番目に多く存在し、鉱物やイオン、気体などとしてマントルに99.6 %、地殻に0.3 %、海水に0.1 %が保有されている^[4]。

マントル - アルステア・キャメロンによる隕石の分析で、ケイ素10,000原子に対し塩素190原子が含まれていると考えられていることから、地球の質量約6,000 Ygに対し22 Yg (22×10²⁴ g) の塩素が存在すると推測される。
地殻 - 塩素が地殻の総重量の0.19 %を占めると考えられていることから、約60 Zg (60×10²¹ g) の塩素が地殻に存在すると推測される。火山噴火により毎年0.4 - 11 Tg (0.4 - 11×10¹² g) の塩素が主に塩化水素の形で対流圏に放出され、その大部分が地表や海洋に降下する。
海水 - 約1.36×10⁸ km³の海水の総量のうち平均塩素濃度19.354 g/kgであることから、26 Zgが主に塩化ナトリウムとして存在すると推測される。海面上で生じる波により、年間6 - 18 Pg (6 - 18×10¹⁵ g) が大気中に放出される。大部分が海洋に戻るが、一部は揮発性塩素となる。
河川水・湖水 - 河川・湖の水の総量は約1×10⁵ km³であり、河川水には約5.8 mg/L含まれていることから、河川・湖水総量に対し580 Tgが含まれていると推測される。
地下水 - 帯水層および土壌中の地下水は地球の水量のおよそ8 %であり、標準的塩素濃度が40 mg/Lであることから、地下水中の塩素含有総量は320 Pgと推測される。
雪氷圏 - 極地や大陸の氷原には、0.5 Gg (0.5×10⁹ g) の塩素が存在すると推測される。
対流圏 - 大気中では主に塩化水素やクロロメタンの状態で存在し、塩化水素は地表近くでは100 - 300 pptv（1 pptvは1兆分の1）、都市部の高濃度域では3,000 pptvが測定される。クロロメタンや、より高層の塩化水素、海水からのエアロゾルなどを含めると、5.3 Tgが存在していると推測される。対流圏から成層圏へは年間約0.03 Tgが放出され、成層圏から対流圏へもほぼ同量が移動する。
成層圏 - 成層圏には約3 pptvの塩素が含まれ、約0.4 Tgの塩素が存在すると推測される。

生産

現在では一般的に塩化ナトリウム水溶液からイオン交換と電気分解とを併用するイオン交換膜法によって水酸化ナトリウムと共に生産される^[5]。塩素ガスの2008年度日本国内生産量は3,911,492トン、消費量は3,497,936トン、液体塩素の2008年度日本国内生産量は519,817トン、消費量は243,098トンである^[6]。高圧ガス保安法に基づく容器保安規則により、黄色いボンベに保管するように決められている^[7]。また液化塩素専用タンク車のタキ5450形も塗装は黄色である。

塩酸やクロロホルムなど各種塩化物の原料、ポリ塩化ビニルやポリ塩化ビニリデンなどの合成樹脂原料として多方面で使用されるほか、合成中間体としてシリコーンやポリウレタン、各種ポリマーなど塩素を含まない製品の製造にも用いられる^[8]。

人体・環境への影響

消毒

塩素は水道水の消毒に使用されており、水道法の規定で、各家庭の蛇口で1リットル当たり0.1 mg以上の濃度を保つように規定されている。一方、有機物と塩素が反応することにより、塩素臭（カルキ臭）が発生するほか^[9]発癌性が疑われるトリハロメタンを生成するといわれ、同様に塩素で汚水処理を行うと水路に塩素化有機物が流れ出てしまうのではないかという懸念の声もある。ただし、コレラなどの病気がほとんどの国で駆逐されたのは塩素を含んだ水道水のおかげでもある。近年は水道水の高度処理が進み、塩素臭は以前に比べて弱まっている^[9]。

単体の毒性

塩素は強い毒性を持つ為、人類初の本格的な化学兵器としても使われた。第一次世界大戦中の1915年4月22日、イープル戦線でのことである。この時にドイツ軍の化学兵器部隊の司令官を務めていたのは後年（1918年）ノーベル化学賞を受賞するフリッツ・ハーバーである。

塩素を吸引するとまず呼吸器に損傷を与える。空気中である程度以上の濃度では、皮膚の粘膜を強く刺激する。目や呼吸器の粘膜を刺激して咳や嘔吐を催し、重大な場合には呼吸不全で死に至る場合もある。液体塩素の場合には、塩素に直接触れた部分が炎症を起こす。

塩素を浴びてしまった場合、直ちにその場から離れ、着ていた衣服を脱ぎ、毛布に包まるなどして体を温めなければならない。直ちに医療機関での処置を要する。呼吸が停止している場合には一刻も早く人工呼吸による蘇生を行わなければならない。呼吸が苦しい場合には酸素マスクの着用を要する。

特に塩素を含む漂白剤（次亜塩素酸ナトリウム）と酸性の物質（主にトイレ用の洗剤）を混合すると、有毒な単体の塩素ガスが遊離し危険な状態となる。このため、漂白剤や酸性のトイレ用の洗剤には「混ぜるな危険」との大きく目立つ表示がある（しかしこれだけの表示では、具体的に何と混ぜると危険なのかが示されていない）。このような表示がされる前（当時も小さな注意書き自体は存在した）には1986年には徳島県で、1989年には長野県で、実際に塩素系漂白剤と酸性洗浄剤を混ぜたことにより、塩素ガスが発生し死亡した事故が起こっている。

オゾン層への影響

「一酸化塩素」も参照

塩素はオゾンホールの原因物質としても指摘されている。フロンなどの塩素原子を含む化合物が紫外線に当たると、結合が切断され塩素ラジカルが生じる。塩素ラジカルは周囲のオゾンと反応して触媒的にオゾンを酸素分子へと分解するため、オゾン層の破壊効果が大きい。

歴史

1774年にスウェーデンのカール・ヴィルヘルム・シェーレが、海塩酸（塩酸）と二酸化マンガンを加熱させることによって単体を分離し、｢脱フロギストン海塩酸気｣と命名。1810年にハンフリー・デービーが元素であると認め、気体が黄緑色である点から、ギリシャ語で「黄緑色」を意味する χλωρος (Chloros) を取って chlorine と命名した。日本語に直訳すれば緑気（りょっき）である。

日本語の「塩素」は、食塩の主成分である点による命名である。

塩素の化合物

塩化物イオンあるいは置換基として塩素を含む化合物は塩化物あるいは塩素化合物と呼ばれる。塩素はほとんどすべての元素と安定な化合物を形成し、また有機化合物にも塩素を含むものが多く知られている（記事塩化物に詳しい）。個々の化合物については、「塩化物のカテゴリ」及び「有機ハロゲン化合物のカテゴリ」を参照されたい。

有機塩素化合物は、安定で、かつ安価に合成できるために、クロロホルムやジクロロメタンのような代表的な有機溶媒として、あるいはポリ塩化ビニルなどのプラスチックとして、大量に生産・使用されている。

反面、多くは毒性を持ち、環境中に放出された際に化学分解され難い点、更に焼却時にはダイオキシンを発生する点から、法令等で規制されている物質も多い。

塩素のオキソ酸

塩素のオキソ酸は慣用名をもつ。次にそれらを挙げる。

オキソ酸の名称	化学式（酸化数）	オキソ酸塩の名称	備考
次亜塩素酸 (hypochlorous acid)	HClO (+I)	次亜塩素酸塩 ( - hypochlorite)	次亜塩素酸塩は塩基性を示し、遊離酸よりも安定で漂白剤、殺菌剤として使用される。
亜塩素酸 (chlorous acid)	HClO₂ (+III)	亜塩素酸塩 ( - chlorite)	亜塩素酸は中程度の酸 (pKa 2.31)。亜塩素酸塩は危険物第1類。
塩素酸 (chloric acid)	HClO₃ (+V)	塩素酸塩 ( - chlorate)	塩素酸は強酸。塩素酸塩は危険物第1類でマッチや火薬などの酸化剤として用いられる。
過塩素酸 (perchloric acid)	HClO₄ (+VII)	過塩素酸塩 ( - perchlorate)	過塩素酸は強酸で危険物第6類。過塩素酸塩は危険物第1類。

※オキソ酸塩名称の '-' にはカチオン種の名称が入る。

塩素のオキソ酸はいずれも酸化力が強い。代表的な化合物に次のようなものがある。

次亜塩素酸ナトリウム (NaClO)
さらし粉（CaCl(ClO)•H₂O) - 不純物として原料の Ca(OH)₂ を含む

同位体

詳細は「塩素の同位体」を参照

出典

^ Magnetic susceptibility of the elements and inorganic compounds, in Lide, D. R., ed. (2005), CRC Handbook of Chemistry and Physics (86th ed.), Boca Raton (FL): CRC Press, ISBN 0-8493-0486-5
^ 日本化学会編　『化学便覧基礎編　改訂4版』　丸善、1993年
^ 『化学大辞典』　共立出版、1993年
^ 『塩素白書』p10-21
^ F.A. コットン, G. ウィルキンソン著, 中原勝儼訳　『コットン・ウィルキンソン無機化学』　培風館、1987年
^ 化学工業統計月報 - 経済産業省
^ 容器保安規則（法令データ提供システム）
^ 『塩素白書』p45
^ ^a ^b 尼崎市水道局

参考文献

IUPAC編、宮本純之監訳　『塩素白書』　化学工業日報社、2000年。ISBN 4-87326-346-8。

外部リンク

塩素 - 「健康食品」の安全性・有効性情報（国立健康・栄養研究所）
『舎密開宗』からたどる，和名「塩酸」｢塩素｣の名称の起源について

ハロゲン間化合物
	フッ素	塩素	臭素	ヨウ素
フッ素	F₂
塩素	ClF　ClF₃　ClF₅	Cl₂
臭素	BrF　BrF₃　BrF₅	BrCl　BrCl₃	Br₂
ヨウ素	IF　IF₃　IF₅　IF₇	ICl　I₂Cl₆	IBr　IBr₃	I₂

wiki en

[Wiki en表示]

Chlorine is a chemical element with symbol Cl and atomic number 17. Chlorine is in the halogen group (17) and is the second lightest halogen after fluorine. The element is a yellow-green gas under standard conditions, where it forms diatomic molecules. It has the highest electron affinity and the fourth highest electronegativity of all the reactive elements; for this reason, chlorine is a strong oxidizing agent. Free chlorine is rare on Earth, and is usually a result of direct or indirect oxidation by oxygen.

The most common compound of chlorine, sodium chloride (common salt), has been known since ancient times. Around 1630 chlorine gas was first synthesized in a chemical reaction, but not recognized as a fundamentally important substance. Characterization of chlorine gas was made in 1774 by Carl Wilhelm Scheele, who supposed it to be an oxide of a new element. In 1809 chemists suggested that the gas might be a pure element, and this was confirmed by Sir Humphry Davy in 1810, who named it from Ancient Greek: χλωρóς khlôros “pale green”.

Nearly all chlorine in the Earth's crust occurs as chloride in various ionic compounds, including table salt. It is the second most abundant halogen and 21st most abundant chemical element in Earth's crust. Elemental chlorine is commercially produced from brine by electrolysis. The high oxidizing potential of elemental chlorine led commercially to free chlorine's bleaching and disinfectant uses, as well as its many uses of an essential reagent in the chemical industry. Chlorine is used in the manufacture of a wide range of consumer products, about two-thirds of them organic chemicals such as polyvinyl chloride, as well as many intermediates for production of plastics and other end products which do not contain the element. As a common disinfectant, elemental chlorine and chlorine-generating compounds are used more directly in swimming pools to keep them clean and sanitary.

In the form of chloride ions, chlorine is necessary to all known species of life. Other types of chlorine compounds are rare in living organisms, and artificially produced chlorinated organics range from inert to toxic. In the upper atmosphere, chlorine-containing organic molecules such as chlorofluorocarbons have been implicated in ozone depletion. Small quantities of elemental chlorine are generated by oxidation of chloride to hypochlorite in neutrophils, as part of the immune response against bacteria. Elemental chlorine at high concentrations is extremely dangerous and poisonous for all living organisms, and was used in World War I as the first gaseous chemical warfare agent.

1 Characteristics
- 1.1 Physical characteristics of chlorine and its compounds
- 1.2 Chemical characteristics
  - 1.2.1 Hydrolysis of free chlorine or disproportionation in water
  - 1.2.2 Chemistry and compounds
    - 1.2.2.1 Chlorides
    - 1.2.2.2 Chlorine oxides
    - 1.2.2.3 Interhalogen compounds
    - 1.2.2.4 Organochlorine compounds
- 1.3 Occurrence
- 1.4 Isotopes
2 History
3 Production
- 3.1 Laboratory methods
4 Applications
- 4.1 Production of industrial and consumer products
- 4.2 Public sanitation, disinfection, and antisepsis
  - 4.2.1 Combating putrefaction
  - 4.2.2 Against infection and contagion
  - 4.2.3 Semmelweis and experiments with antisepsis
  - 4.2.4 Public sanitation
- 4.3 Use as a weapon
  - 4.3.1 World War I
  - 4.3.2 Iraq War
5 Health effects of the free element and hazards
- 5.1 Chlorine induced cracking in structural materials
- 5.2 Chlorine-iron fire
6 Organochlorine compounds as pollutants
7 See also
8 References
9 Bibliography
10 External links

Characteristics

Physical characteristics of chlorine and its compounds

Chlorine, liquefied under a pressure of 7.4 bar at room temperature, displayed in a quartz ampule embedded in acrylic glass.

At standard temperature and pressure, two chlorine atoms form the diatomic molecule Cl₂.^[3] This is a yellow-green gas that has a distinctive strong odor, familiar to most from common household bleach.^[4] The bonding between the two atoms is relatively weak (only 242.580 ± 0.004 kJ/mol), which makes the Cl₂ molecule highly reactive. The boiling point at standard pressure is around −34 ˚C, but it can be liquefied at room temperature with pressures above 740 kPa (107 psi).^[5]

Although elemental chlorine is yellow-green, chloride ion, in common with other halide ions, has no color in either minerals or solutions (example, table salt). Similarly, (again as with other halogens) chlorine atoms impart no color to organic chlorides when they replace hydrogen atoms in colorless organic compounds, such as tetrachloromethane. The melting point and density of these compounds is increased by substitution of hydrogen in place of chlorine. Compounds of chlorine with other halogens, however, as well as many chlorine oxides, are visibly colored.

Chemical characteristics

Along with fluorine, bromine, iodine, and astatine, chlorine is a member of the halogen series that forms the group 17 (formerly VII, VIIA, or VIIB) of the periodic table. Chlorine forms compounds with almost all of the elements to give compounds that are usually called chlorides. Chlorine gas reacts with most organic compounds, and will even sluggishly support the combustion of hydrocarbons.^[6]

Hydrolysis of free chlorine or disproportionation in water

At 25 °C and atmospheric pressure, one liter of water dissolves 3.26 g or 1.125 L of gaseous chlorine.^[7] Solutions of chlorine in water contain chlorine (Cl₂), hydrochloric acid, and hypochlorous acid:

Cl₂ + H₂O HCl + HClO

This conversion to the right is called disproportionation, because the ingredient chlorine both increases and decreases in formal oxidation state. The solubility of chlorine in water is increased if the water contains dissolved alkali hydroxide, and in this way, chlorine bleach is produced.^[8]

Cl₂ + 2 OH^– → ClO^– + Cl^– + H₂O

Chlorine gas only exists in a neutral or acidic solution.

Chemistry and compounds

Chlorides

Chlorine combines with almost all elements to give chlorides. Compounds with oxygen, nitrogen, xenon, and krypton are known, but do not form by direct reaction of the elements.^[11] Chloride is one of the most common anions in nature. Hydrogen chloride and its aqueous solution, hydrochloric acid, are produced on megaton scale annually both as valued intermediates but sometimes as undesirable pollutants.

Chlorine oxides

Chlorine forms a variety of oxides, as seen above: chlorine dioxide (ClO₂), dichlorine monoxide (Cl₂O), dichlorine hexoxide (Cl₂O₆), dichlorine heptoxide (Cl₂O₇). The anionic derivatives of these same oxides are also well known including chlorate (ClO−
3), chlorite (ClO−
2), hypochlorite (ClO⁻), and perchlorate (ClO−
4). The acid derivatives of these anions are hypochlorous acid (HOCl), chloric acid (HClO₃) and perchloric acid (HClO₄). The chloroxy cation chloryl (ClO₂⁺) is known and has the same structure as chlorite but with a positive charge and chlorine in the +5 oxidation state.^[12] The compound "chlorine trioxide" does not occur, but rather in gas form is found as the dimeric dichlorine hexoxide (Cl₂O₆) with a +6 oxidation state. This compound in liquid or solid form disproportionates to a mixture of +5 and +7 oxidation states, occurring as the ionic compound chloryl perchlorate, [ClO
2]+
[ClO
4]−
.^[13]

In hot concentrated alkali solution hypochlorite disproportionates:

2 ClO⁻ → Cl⁻ + ClO−
2

ClO⁻ + ClO−
2 → Cl⁻ + ClO−
3

Sodium chlorate and potassium chlorate can be crystallized from solutions formed by the above reactions. If their crystals are heated to a high temperature, they undergo a further, final disproportionation:

4 ClO−
3 → Cl⁻ + 3 ClO−
4

This same progression from chloride to perchlorate can be accomplished by electrolysis. The anode reaction progression is:^[14]

Reaction	Electrode potential
Cl⁻ + 2 OH⁻ → ClO⁻ + H₂O + 2 e⁻	+0.89 volts
ClO⁻ + 2 OH⁻ → ClO− 2 + H₂O + 2 e⁻	+0.67 volts
ClO− 2 + 2 OH⁻ → ClO− 3 + H₂O + 2 e⁻	+0.33 volts
ClO− 3 + 2 OH⁻ → ClO− 4 + H₂O + 2 e⁻	+0.35 volts

Each step is accompanied at the cathode by

2 H₂O + 2 e⁻ → 2 OH⁻ + H₂ (−0.83 volts)

Interhalogen compounds

Chlorine oxidizes bromide and iodide salts to bromine and iodine, respectively. However, it cannot oxidize fluoride salts to fluorine. It makes a variety of interhalogen compounds, such as the chlorine fluorides, chlorine monofluoride (ClF), chlorine trifluoride (ClF
3), chlorine pentafluoride (ClF
5). Chlorides of bromine and iodine are also known.^[15]

Organochlorine compounds

Chlorine is used extensively in organic chemistry in substitution and addition reactions. Chlorine often imparts many desired properties to an organic compound, in part owing to its electronegativity.

Like the other halides, chlorine undergoes electrophilic addition reactions, the most notable one being the chlorination of alkenes and aromatic compounds with a Lewis acid catalyst. Organic chlorine compounds tend to be less reactive in nucleophilic substitution reactions than the corresponding bromine or iodine derivatives, but they tend to be cheaper. They may be activated for reaction by substituting with a tosylate group, or by the use of a catalytic amount of sodium iodide.^{[citation needed]}

Occurrence

Isotopes

Chlorine has a wide range of isotopes. The two stable isotopes are ³⁵Cl (75.77%) and ³⁷Cl (24.23%).^[19] Together they give chlorine an atomic weight of 35.4527 g/mol. The half-integer value for chlorine's weight caused some confusion in the early days of chemistry, when it had been postulated that atoms were composed of even units of hydrogen (see Proust's law), and the existence of chemical isotopes was unsuspected.^[20]

Trace amounts of radioactive ³⁶Cl exist in the environment, in a ratio of about 7x10⁻¹³ to 1 with stable isotopes. ³⁶Cl is produced in the atmosphere by spallation of ³⁶Ar by interactions with cosmic ray protons. In the subsurface environment, ³⁶Cl is generated primarily as a result of neutron capture by ³⁵Cl or muon capture by ⁴⁰Ca. ³⁶Cl decays to ³⁶S and to ³⁶Ar, with a combined half-life of 308,000 years. The half-life of this hydrophilic nonreactive isotope makes it suitable for geologic dating in the range of 60,000 to 1 million years. Additionally, large amounts of ³⁶Cl were produced by irradiation of seawater during atmospheric detonations of nuclear weapons between 1952 and 1958. The residence time of ³⁶Cl in the atmosphere is about 1 week. Thus, as an event marker of 1950s water in soil and groundwater, ³⁶Cl is also useful for dating waters less than 50 years before the present. ³⁶Cl has seen use in other areas of the geological sciences, including dating ice and sediments.^[19]

History

The most common compound of chlorine, sodium chloride, has been known since ancient times; archaeologists have found evidence that rock salt was used as early as 3000 BC and brine as early as 6000 BC.^[21] Around 1630, chlorine was recognized as a gas by the Belgian chemist and physician Jan Baptist van Helmont.^[22]

Carl Wilhelm Scheele

Elemental chlorine was first prepared and studied in 1774 by Swedish chemist Carl Wilhelm Scheele, and, therefore, he is credited for its discovery.^[23] He called it "dephlogisticated muriatic acid air" since it is a gas (then called "airs") and it came from hydrochloric acid (then known as "muriatic acid").^[23] However, he failed to establish chlorine as an element, mistakenly thinking that it was the oxide obtained from the hydrochloric acid (see phlogiston theory).^[23] He named the new element within this oxide as muriaticum.^[23] Regardless of what he thought, Scheele did isolate chlorine by reacting MnO₂ (as the mineral pyrolusite) with HCl:^[22]

4 HCl + MnO₂ → MnCl₂ + 2 H₂O + Cl₂

Scheele observed several of the properties of chlorine: the bleaching effect on litmus, the deadly effect on insects, the yellow green color, and the smell similar to aqua regia.^[24]

At the time, common chemical theory was: any acid is a compound that contains oxygen (still sounding in the German and Dutch names of oxygen: sauerstoff or zuurstof, both translating into English as acid stuff), so a number of chemists, including Claude Berthollet, suggested that Scheele's dephlogisticated muriatic acid air must be a combination of oxygen and the yet undiscovered element, muriaticum.^[25]^[26]^[27]

In 1809, Joseph Louis Gay-Lussac and Louis-Jacques Thénard tried to decompose dephlogisticated muriatic acid air by reacting it with charcoal to release the free element muriaticum (and carbon dioxide).^[23] They did not succeed and published a report in which they considered the possibility that dephlogisticated muriatic acid air is an element, but were not convinced.^[28]

In 1810, Sir Humphry Davy tried the same experiment again, and concluded that it is an element, and not a compound.^[23] He named this new element as chlorine, from the Greek word χλωρος (chlōros), meaning green-yellow.^[29] The name halogen, meaning "salt producer," was originally used for chlorine in 1811 by Johann Salomo Christoph Schweigger. However, this term was later used as a generic term to describe all the elements in the chlorine family (fluorine, bromine, iodine), after a suggestion by Jöns Jakob Berzelius in 1842.^[30]^[31] In 1823, Michael Faraday liquefied chlorine for the first time,^[32]^[33] and demonstrated that what was then known as "solid chlorine" had a structure of chlorine hydrate (Cl₂·H₂O).^[22]

Chlorine gas was first used by French chemist Claude Berthollet to bleach textiles in 1785.^[34]^[35] Modern bleaches resulted from further work by Berthollet, who first produced sodium hypochlorite in 1789 in his laboratory in the town of Javel (now part of Paris, France), by passing chlorine gas through a solution of sodium carbonate. The resulting liquid, known as "Eau de Javel" ("Javel water"), was a weak solution of sodium hypochlorite. However, this process was not very efficient, and alternative production methods were sought. Scottish chemist and industrialist Charles Tennant first produced a solution of calcium hypochlorite ("chlorinated lime"), then solid calcium hypochlorite (bleaching powder).^[34] These compounds produced low levels of elemental chlorine, and could be more efficiently transported than sodium hypochlorite, which remained as dilute solutions because when purified to eliminate water, it became a dangerously powerful and unstable oxidizer. Near the end of the nineteenth century, E. S. Smith patented a method of sodium hypochlorite production involving electrolysis of brine to produce sodium hydroxide and chlorine gas, which then mixed to form sodium hypochlorite.^[36] This is known as the chloralkali process, first introduced on an industrial scale in 1892, and now the source of essentially all modern elemental chlorine and sodium hydroxide production (a related low-temperature electrolysis reaction, the Hooker process, is now responsible for bleach and sodium hypochlorite production).

Elemental chlorine solutions dissolved in chemically basic water (sodium and calcium hypochlorite) were first used as anti-putrification agents and disinfectants in the 1820s, in France, long before the establishment of the germ theory of disease. This work is mainly due to Antoine-Germain Labarraque, who adapted Berthollet's "Javel water" bleach and other chlorine preparations for the purpose (for a more complete history, see below). Elemental chlorine has since served a continuous function in topical antisepsis (wound irrigation solutions and the like) as well as public sanitation (especially of swimming and drinking water).

In 1826, silver chloride was used to produce photographic images for the first time.^[37] Chloroform was first used as an anesthetic in 1847.^[37]

Polyvinyl chloride (PVC) was invented in 1912, initially without a purpose.^[37]

Chlorine gas was first introduced as a weapon on April 22, 1915, at Ypres by the German Army,^[38]^[39] and the results of this weapon were disastrous because gas masks had not been mass distributed and were tricky to get on quickly.

Production

Liquid chlorine analysis

In industry, elemental chlorine is usually produced by the electrolysis of sodium chloride dissolved in water. This method, the chloralkali process industrialized in 1892, now provides essentially all industrial chlorine gas.^[40] Along with chlorine, the method yields hydrogen gas and sodium hydroxide (with sodium hydroxide actually being the most crucial of the three industrial products produced by the process). The process proceeds according to the following chemical equation:^[10]

2 NaCl + 2 H₂O → Cl₂ + H₂ + 2 NaOH

The electrolysis of chloride solutions all proceed according to the following equations:

Cathode: 2 H⁺_(aq) + 2 e⁻ → H_2(g)

Anode: 2 Cl⁻_(aq) → Cl_2(g) + 2 e⁻

Overall process: 2 NaCl (or KCl) + 2 H₂O → Cl₂ + H₂ + 2 NaOH (or KOH)

In diaphragm cell electrolysis, an asbestos (or polymer-fiber) diaphragm separates a cathode and an anode, preventing the chlorine forming at the anode from re-mixing with the sodium hydroxide and the hydrogen formed at the cathode.^[41] The salt solution (brine) is continuously fed to the anode compartment and flows through the diaphragm to the cathode compartment, where the caustic alkali is produced and the brine is partially depleted. Diaphragm methods produce dilute and slightly impure alkali but they are not burdened with the problem of preventing mercury discharge into the environment and they are more energy efficient. Membrane cell electrolysis employ permeable membrane as an ion exchanger. Saturated sodium (or potassium) chloride solution is passed through the anode compartment, leaving at a lower concentration.^[42] This method is more efficient than the diaphragm cell and produces very pure sodium (or potassium) hydroxide at about 32% concentration, but requires very pure brine.

Membrane cell process for chloralkali production

Laboratory methods

Small amounts of chlorine gas can be made in the laboratory by combining hydrochloric acid and manganese dioxide. Alternatively a strong acid such as sulfuric acid or hydrochloric acid reacts with sodium hypochlorite solution to release chlorine gas but reacts with sodium chlorate to produce chlorine gas and chlorine dioxide gas as well. In the home, accidents occur when hypochlorite bleach solutions are combined with certain acidic drain-cleaners.

Applications

Production of industrial and consumer products

Principal applications of chlorine are in the production of a wide range of industrial and consumer products.^[43]^[44] For example, it is used in making plastics, solvents for dry cleaning and metal degreasing, textiles, agrochemicals and pharmaceuticals, insecticides, dyestuffs, household cleaning products, etc.

Many important industrial products are produced via organochlorine intermediates. Examples include polycarbonates, polyurethanes, silicones, polytetrafluoroethylene, carboxymethyl cellulose, and propylene oxide. Like the other halogens, chlorine participates in free-radical substitution reactions with hydrogen-containing organic compounds. When applied to organic substrates, reaction is often—but not invariably—non-regioselective, and, hence, may result in a mixture of isomeric products. It is often difficult to control the degree of substitution as well, so multiple substitutions are common. If the different reaction products are easily separated, e.g., by distillation, substitutive free-radical chlorination (in some cases accompanied by concurrent thermal dehydrochlorination) may be a useful synthetic route. Industrial examples of this are the production of methyl chloride, methylene chloride, chloroform, and carbon tetrachloride from methane, allyl chloride from propylene, and trichloroethylene, and tetrachloroethylene from 1,2-dichloroethane.

Quantitatively, about 63% and 18% of all elemental chlorine produced is used in the manufacture of organic and inorganic chlorine compounds, respectively,.^[40] About 15,000 chlorine compounds are being used commercially.^[24] The remaining 19% is used for bleaches and disinfection products.^[40] The most significant of organic compounds in terms of production volume are 1,2-dichloroethane and vinyl chloride, intermediates in the production of PVC. Other particularly important organochlorines are methyl chloride, methylene chloride, chloroform, vinylidene chloride, trichloroethylene, perchloroethylene, allyl chloride, epichlorohydrin, chlorobenzene, dichlorobenzenes, and trichlorobenzenes. The major inorganic compounds include HCl, Cl₂O, HOCl, NaClO₃, chlorinated isocyanurates, AlCl₃, SiCl₄, SnCl₄, PCl₃, PCl₅, POCl₃, AsCl₃, SbCl₃, SbCl₅, BiCl₃, S₂Cl₂, SCl₂, SOCI₂, ClF₃, ICl, ICl₃, TiCl₃, TiCl₄, MoCl₅, FeCl₃, ZnCl₂, etc.^[40]^[45]

Public sanitation, disinfection, and antisepsis

Combating putrefaction

Antoine-Germain Labarraque

In France (as elsewhere) there was a need to process animal guts in order to make musical instrument strings, Goldbeater's skin and other products. This was carried out in "gut factories" (boyauderies) as an odiferous and unhealthy business. In or about 1820, the Société d'encouragement pour l'industrie nationale offered a prize for the discovery of a method, chemical or mechanical, that could be used to separate the peritoneal membrane of animal intestines without causing putrefaction.^[46]^[47] It was won by Antoine-Germain Labarraque, a 44 year-old French chemist and pharmacist who had discovered that Berthollet's chlorinated bleaching solutions ("Eau de Javel") not only destroyed the smell of putrefaction of animal tissue decomposition, but also retarded the decomposition process itself.^[47]^[48]

Labarraque's research resulted in chlorides and hypochlorites of lime (calcium hypochlorite) and of sodium (sodium hypochlorite) being employed not only in the boyauderies but also for the routine disinfection and deodorisation of latrines, sewers, markets, abattoirs, anatomical theatres and morgues.^[49] They were also used, with success, in hospitals, lazarets, prisons, infirmaries (both on land and at sea), magnaneries, stables, cattle-sheds, etc.; and for exhumations,^[50] embalming, during outbreaks of epidemic illness, fever, blackleg in cattle, etc.^[46]

Against infection and contagion

Labarraque's chlorinated lime and soda solutions have been advocated since 1828 to prevent infection (called "contagious infection", and presumed to be transmitted by "miasmas") and also to treat putrefaction of existing wounds, including septic wounds.^[51] In this 1828 work, Labarraque recommended for the doctor to breathe chlorine, wash his hands with chlorinated lime, and even sprinkle chlorinated lime about the patient's bed, in cases of "contagious infection." In 1828, it was well known that some infections were contagious, even though the agency of the microbe was not to be realized or discovered for more than half a century.

During the Paris cholera outbreak of 1832, large quantities of so-called chloride of lime were used to disinfect the capital. This was not simply modern calcium chloride, but contained chlorine gas dissolved in lime-water (dilute calcium hydroxide) to form calcium hypochlorite (chlorinated lime). Labarraque's discovery helped to remove the terrible stench of decay from hospitals and dissecting rooms, and, by doing so, effectively deodorised the Latin Quarter of Paris.^[52] These "putrid miasmas" were thought by many to be responsible for the spread of "contagion" and "infection" – both words used before the germ theory of infection. The use of chloride of lime was based on destruction of odors and "putrid matter." One source has claimed that chloride of lime was used by Dr. John Snow to disinfect water from the cholera-contaminated well feeding the Broad Street pump in 1854 London.^[53] Three reputable sources that described the famous Broad Street pump cholera epidemic do not mention Snow performing any disinfection of water from that well.^[54]^[55]^[56] Instead, one reference makes it clear that chloride of lime was used to disinfect the offal and filth in the streets surrounding the Broad Street pump—a common practice in mid-nineteenth century England.^[54]^:296

Semmelweis and experiments with antisepsis

Ignaz Semmelweis

Perhaps the most famous application of Labarraque's chlorine and chemical base solutions was in 1847, when Ignaz Semmelweis used (first) chlorine-water (simply chlorine dissolved in pure water), then cheaper chlorinated lime solutions, to deodorize the hands of Austrian doctors, which Semmelweis noticed still carried the stench of decomposition from the dissection rooms to the patient examination rooms. Semmelweis, still long before the germ theory of disease, had theorized that "cadaveric particles" were somehow transmitting decay from fresh medical cadavers to living patients, and he used the well-known "Labarraque's solutions" as the only known method to remove the smell of decay and tissue decomposition (which he found that soap did not). The solutions proved to be far more effective germicide antiseptics than soap (Semmelweis was also aware of their greater efficacy, but not the reason), and this resulted in Semmelweis's (later) celebrated success in stopping the transmission of childbed fever ("puerperal fever") in the maternity wards of Vienna General Hospital in Austria in 1847.^[57]

Much later, during World War I in 1916, a standardized and diluted modification of Labarraque's solution, containing hypochlorite (0.5%) and boric acid as an acidic stabilizer, was developed by Henry Drysdale Dakin (who gave full credit to Labarraque's prior work in this area). Called Dakin's solution, the method of wound irrigation with chlorinated solutions allowed antiseptic treatment of a wide variety of open wounds, long before the modern antibiotic era. A modified version of this solution continues to be employed in wound irrigation in the modern era, where it remains effective against multiply antibiotic resistant bacteria (see Century Pharmaceuticals).

Public sanitation

By 1918, the US Department of Treasury called for all drinking water to be disinfected with chlorine. Chlorine is presently an important chemical for water purification (such as in water treatment plants), in disinfectants, and in bleach. Chlorine in water is more than three times as effective as a disinfectant against Escherichia coli than an equivalent concentration of bromine, and is more than six times more effective than an equivalent concentration of iodine.^[58]

Chlorine is usually used (in the form of hypochlorous acid) to kill bacteria and other microbes in drinking water supplies and public swimming pools. In most private swimming pools, chlorine itself is not used, but rather sodium hypochlorite, formed from chlorine and sodium hydroxide, or solid tablets of chlorinated isocyanurates. The drawback of using chlorine in swimming pools is that the chlorine reacts with the proteins in human hair and skin (see Hypochlorous acid). Once the chlorine reacts with the hair and skin, it becomes chemically bonded. Even small water supplies are now routinely chlorinated.^[6]

It is often impractical to store and use poisonous chlorine gas for water treatment, so alternative methods of adding chlorine are used. These include hypochlorite solutions, which gradually release chlorine into the water, and compounds like sodium dichloro-s-triazinetrione (dihydrate or anhydrous), sometimes referred to as "dichlor", and trichloro-s-triazinetrione, sometimes referred to as "trichlor". These compounds are stable while solid and may be used in powdered, granular, or tablet form. When added in small amounts to pool water or industrial water systems, the chlorine atoms hydrolyze from the rest of the molecule forming hypochlorous acid (HOCl), which acts as a general biocide, killing germs, micro-organisms, algae, and so on.^[59]^[60]

Use as a weapon

World War I

Chlorine gas, also known as bertholite, was first used as a weapon in World War I by Germany on April 22, 1915 in the Second Battle of Ypres.^[61] As described by the soldiers it had a distinctive smell of a mixture between pepper and pineapple. It also tasted metallic and stung the back of the throat and chest. Chlorine can react with water in the mucosa of the lungs to form hydrochloric acid, an irritant that can be lethal. The damage done by chlorine gas can be prevented by the activated charcoal commonly found in gas masks, or other filtration methods, which makes the overall chance of death by chlorine gas much lower than those of other chemical weapons. It was pioneered by a German scientist later to be a Nobel laureate, Fritz Haber of the Kaiser Wilhelm Institute in Berlin, in collaboration with the German chemical conglomerate IG Farben, who developed methods for discharging chlorine gas against an entrenched enemy. It is alleged that Haber's role in the use of chlorine as a deadly weapon drove his wife, Clara Immerwahr, to suicide.^[62] After its first use, chlorine was utilized by both sides as a chemical weapon, but it was soon replaced by the more deadly phosgene and mustard gas.^[63]

Iraq War

Chlorine gas has also been used by insurgents against the local population and coalition forces in the Iraq War in the form of chlorine bombs. On March 17, 2007, for example, three chlorine-filled trucks were detonated in the Anbar province killing two and sickening over 350.^[64] Other chlorine bomb attacks resulted in higher death tolls, with more than 30 deaths on two separate occasions.^[65] Most of the deaths were caused by the force of the explosions rather than the effects of chlorine, since the toxic gas is readily dispersed and diluted in the atmosphere by the blast. The Iraqi authorities have tightened security for elemental chlorine, which is essential for providing safe drinking water to the population.

Health effects of the free element and hazards

Chlorine is a toxic gas that irritates the respiratory system. Because it is heavier than air, it tends to accumulate at the bottom of poorly ventilated spaces. Chlorine gas is a strong oxidizer, which may react with flammable materials.^[66]

Chlorine is detectable with measuring devices in concentrations of as low as 0.2 parts per million (ppm), and by smell at 3 ppm. Coughing and vomiting may occur at 30 ppm and lung damage at 60 ppm. About 1000 ppm can be fatal after a few deep breaths of the gas.^[24] Breathing lower concentrations can aggravate the respiratory system, and exposure to the gas can irritate the eyes.^[67] The toxicity of chlorine comes from its oxidizing power. When chlorine is inhaled at concentrations above 30 ppm, it begins to react with water and cells, which change it into hydrochloric acid (HCl) and hypochlorous acid (HClO).

When used at specified levels for water disinfection, the reaction of chlorine with water is not a major concern for human health. Other materials present in the water may generate disinfection by-products that are associated with negative effects on human health,^[68]^[69] however, the health risk is far lower than drinking undisinfected water.

Chlorine induced cracking in structural materials

Chlorine "attack" on an acetal resin plumbing joint.

The element is widely used for purifying water owing to its powerful oxidizing properties, especially potable water supplies and water used in swimming pools. Several catastrophic collapses of swimming pool ceilings have occurred owing to chlorine induced stress corrosion cracking of stainless steel rods used to suspend them.^[70] Some polymers are also sensitive to attack, including acetal resin and polybutene. Both materials were used in hot and cold water domestic supplies, and stress corrosion cracking caused widespread failures in the USA in the 1980s and 1990s. The picture on the right shows an acetal joint in a water supply system, which, when it fractured, caused substantial physical damage to computers in the labs below the supply. The cracks started at injection molding defects in the joint and slowly grew until finally triggered. The fracture surface shows iron and calcium salts that were deposited in the leaking joint from the water supply before failure.^[71]

Chlorine-iron fire

The element iron can combine with chlorine at high temperatures in a strong exothermic reaction, creating a chlorine-iron fire.^[72]^[73] Chlorine-iron fires are a risk in chemical process plants, where much of the pipework used to carry chlorine gas is made of steel.^[72]^[73]

Organochlorine compounds as pollutants

Some organochlorine compounds are serious pollutants. These are produced either as by-products or end products of industrial processes which are persistent in the environment, such as certain chlorinated pesticides and chlorofluorocarbons. Chlorine is added both to pesticides and pharmaceuticals to make the molecules more resistant to enzymatic degradation by bacteria, insects, and mammals, but this property also has the effect of prolonging the residence time of these compounds when they enter the environment. In this respect chlorinated organics have some resemblance to fluorinated organics.

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Bibliography

Greenwood, Norman N; Earnshaw, Alan (1997). Chemistry of the Elements (2 ed.). Oxford: Butterworth-Heinemann. ISBN 0-08-037941-9.
Wiberg, Egon; Wiberg, Nils; Holleman, Arnold Frederick (2001). Inorganic Chemistry. Academic Press. ISBN 0-12-352651-5.

External links

Chlorine at The Periodic Table of Videos (University of Nottingham)
Agency for Toxic Substances and Disease Registry: Chlorine
Electrolytic production
Production and liquefaction of chlorine
Chlorine Production Using Mercury, Environmental Considerations and Alternatives
National Pollutant Inventory – Chlorine^{[dead link]}
National Institute for Occupational Safety and Health – Chlorine Page
Chlorine Institute – Trade association representing the chlorine industry
Chlorine Online – the web portal of Eurochlor – the business association of the European chlor-alkali industry

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1. 反応性気道機能障害症候群および刺激に誘発された喘息 reactive airways dysfunction syndrome and irritant induced asthma
2. Chemical terrorism: Rapid recognition and initial medical management
3. 成人における腐食性の食道損傷 caustic esophageal injury in adults
4. 皮膚炎の概要 overview of dermatitis
5. レジオネラ感染の治療および予防 treatment and prevention of legionella infection

English Journal

The kinetics and mechanism of ultrasonic degradation of p-nitrophenol in aqueous solution with CCl(4) enhancement.

Wang X, Wei Y, Wang J, Guo W, Wang C.SourceSchool of Chemistry and Chemical Engineering, University of Jinan, Jinan 250022, China; School of Chemical Engineering, Shandong Institute of Light Industry, Jinan 250353, China.
Ultrasonics sonochemistry.Ultrason Sonochem.2012 Jan;19(1):32-7. Epub 2010 Dec 30.
The ultrasonic degradation of p-nitrophenol (p-NP) in aqueous solution with CCl(4) enhancement was studied. The effects of operating parameters such as CCl(4) dosage, ultrasonic power, media temperature, the initial concentration of p-NP and initial pH value of the aqueous solution on the degradatio
PMID 21741873

The controlled photoactivity of nanoparticles derived from ionic interactions between a water soluble polymeric photosensitizer and polysaccharide quencher.

Park W, Park SJ, Na K.SourceDepartment of Biotechnology, The Catholic University of Korea, 43-1 Yeokkok2-dong, Wonmi-gu, Bucheon-si, Gyeonggi-do 420 743, Republic of Korea.
Biomaterials.Biomaterials.2011 Nov;32(32):8261-70. Epub 2011 Jul 30.
In order to design a water soluble polymeric photosensitizer (WPS) with controllable photoactivity, a nano-photosensitizer (NPS) was prepared from a polyelectrolyte complex between polyethylene glycol-polyethylenimine-chlorine e6 conjugate (PEG-PEI-Ce6) and Black Hole Quencher-3 chondroitin sulfate
PMID 21803417

Rapid free chlorine decay in the presence of Cu(OH)(2): Chemistry and practical implications.

Nguyen CK, Powers KA, Raetz MA, Parks JL, Edwards MA.SourceVirginia Tech, 418 Durham Hall, Blacksburg, VA 24061, USA.
Water research.Water Res.2011 Oct 15;45(16):5302-12. Epub 2011 Aug 5.
A rapid reaction between free chlorine and the cupric hydroxide [Cu(OH)(2)] solids commonly found on pipe walls in premise plumbing can convert free chlorine to chloride and rapidly age Cu(OH)(2) to tenorite (CuO). This reaction has important practical implications for maintaining free chlorine resi
PMID 21868051

Japanese Journal

Molecular dynamics analysis of the formation of surface roughness during Si etching in chlorine-based plasmas (Special issue: Dry process)

Tsuda Hirotaka,Takao Yoshinori,Eriguchi Koji [他]
Japanese journal of applied physics 50(8), 08KB02-1-5, 2011-08
NAID 40018961085

Molecular dynamics of ion incident angle dependence of sputtering yield in chlorine-adsorbed GaN crystal (Special issue: Plasma processing)

Harafuji Kenji,Kawamura Katsuyuki
Japanese journal of applied physics 50(8), 08JG03-1-5, 2011-08
NAID 40018961076

Three-dimensional atomic-scale cellular model and feature profile evolution during Si etching in chlorine-based plasmas: analysis of profile anomalies and surface roughness (Special issue: Plasma processing)

Tsuda Hirotaka,Miyata Hiroki,Takao Yoshinori [他]
Japanese journal of applied physics 50(8), 08JE06-1-6, 2011-08
NAID 40018961058

Related Pictures

★リンクテーブル★

リンク元	「Cl」「塩素」
拡張検索	「organochlorine pesticide」

「Cl」

NFPA 704; fire diamond
	0 3 0 OX

Chlorine
	17Cl
	sulfur ← chlorine → argon
	Chlorine in the periodic table
Appearance
	pale yellow-green gas;
General properties
Name, symbol, number	chlorine, Cl, 17
Pronunciation	/ˈklɔəriːn/ KLOHR-een; or /ˈklɔərɨn/ KLOHR-ən
Element category	diatomic nonmetal
Group, period, block	17 (halogens), 3, p
Standard atomic weight	35.45(1)
Electron configuration	[Ne] 3s2 3p5; 2, 8, 7
Physical properties
Phase	gas
Density	(0 °C, 101.325 kPa); 3.2 g/L
Liquid density at b.p.	1.5625 g·cm−3
Melting point	171.6 K, -101.5 °C, -150.7 °F
Boiling point	239.11 K, -34.04 °C, -29.27 °F
Critical point	416.9 K, 7.991 MPa
Heat of fusion	(Cl2) 6.406 kJ·mol−1
Heat of vaporization	(Cl2) 20.41 kJ·mol−1
Molar heat capacity	(Cl2); 33.949 J·mol−1·K−1
Vapor pressure
Atomic properties
Oxidation states	7, 6, 5, 4, 3, 2, 1, -1; (strongly acidic oxide)
Electronegativity	3.16 (Pauling scale)
Ionization energies; (more)	1st: 1251.2 kJ·mol−1
	2nd: 2298 kJ·mol−1
	3rd: 3822 kJ·mol−1
Covalent radius	102±4 pm
Van der Waals radius	175 pm
Miscellanea
Crystal structure	orthorhombic
Magnetic ordering	diamagnetic
Electrical resistivity	(20 °C) > 10 Ω·m
Thermal conductivity	8.9×10−3 W·m−1·K−1
Speed of sound	(gas, 0 °C) 206 m·s−1
CAS registry number	7782-50-5
History
Discovery	Carl Wilhelm Scheele (1774)
First isolation	Carl Wilhelm Scheele (1774)
Recognized as an element by	Humphry Davy (1808)
Most stable isotopes
	Main article: Isotopes of chlorine
	v; t; e; · references

外見
	黄緑色気体;
一般特性
名称, 記号, 番号	塩素, Cl, 17
分類	ハロゲン
族, 周期, ブロック	17, 3, p
原子量	35.453(2) g·mol-1
電子配置	[Ne] 3s2 3p5
電子殻	2, 8, 7（画像）
物理特性
相	気体
密度	(0 °C, 101.325 kPa); 3.2 g/L
融点	171.6 K, -101.5 °C, -150.7 °F
沸点	239.11 K, -34.04 °C, -29.27 °F
臨界点	416.9 K, 7.991 MPa
融解熱	(Cl2) 6.406 kJ·mol-1
蒸発熱	(Cl2) 20.41 kJ·mol-1
熱容量	(25 °C) (Cl2) 33.949 J·mol-1·K-1
蒸気圧
原子特性
酸化数	7, 6, 5, 4, 3, 2, 1, -1（強酸性酸化物）
電気陰性度	3.16（ポーリングの値）
イオン化エネルギー; （詳細）	第1： 1251.2 kJ·mol-1
	第2： 2298 kJ·mol-1
	第3： 3822 kJ·mol-1
共有結合半径	102±4 pm
ファンデルワールス半径	175 pm
その他
結晶構造	斜方晶系
磁性	反磁性
電気抵抗率	(20 °C) > 10Ω·m
熱伝導率	(300 K) 8.9×10-3 W·m-1·K-1
音の伝わる速さ	（気体、0 °C）206 m/s
CAS登録番号	7782-50-5
最安定同位体
	詳細は塩素の同位体を参照
	表示

　　[★]

塩素イオン chloride ion

「塩素」

　　[★]

英: chlorine
同: クロル
関: 塩素イオン Cl^-

Cl₂

「organochlorine pesticide」

　　[★]

有機塩素系殺虫剤、有機塩素剤

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[daijiten-3] 『化学大辞典』　共立出版、1993年

[4] 『塩素白書』p10-21

[Cotton-5] F.A. コットン, G. ウィルキンソン著, 中原勝儼訳　『コットン・ウィルキンソン無機化学』　培風館、1987年

[6] 化学工業統計月報 - 経済産業省

[7] 容器保安規則（法令データ提供システム）

[8] 『塩素白書』p45

[amagasaki-9] 尼崎市水道局

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[29] Sir Humphry Davy (1811). "On a Combination of Oxymuriatic Gas and Oxygene Gas". Philosophical Transactions of the Royal Society 101: 155–162. doi:10.1098/rstl.1811.0008.

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[33] O'Connor J J and Robertson E F. "Michael Faraday". School of Mathematics and Statistics, University of St Andrews, Scotland. Retrieved 2010-05-08.

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[36] "How Products Are Made Volume 2". May 2011.

[history-37] Jacqueline Brazin. "Chlorine & its Consequences". Archived from the original on September 18, 2006. Retrieved 2008-07-10.

[38] "Chlorine – History". Retrieved 2008-07-10. ^{[dead link]}

[39] "Weaponry: Use of Chlorine Gas Cylinders in World War I". historynet.com. Retrieved 2008-07-10.

[FOOTNOTEGreenwood1997798-40] Greenwood 1997, p. 798.

[Euro_Chlor3-41] "The diaphragm cell process". Euro Chlor. Retrieved 2007-08-15.

[Euro_Chlor4-42] "The membrane cell process". Euro Chlor. Retrieved 2007-08-15.

[Euro_Chlor5-43] "What is chlorine used for?". Euro Chlor. October 2010. Archived from the original on 2011-07-26. Retrieved 2007-08-20.

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Oxidation state	Name	Formula	Characteristic compounds
−1	chlorides	Cl⁻	ionic chlorides, organic chlorides, hydrochloric acid
0	chlorine	Cl₂	elemental chlorine
+1	hypochlorites	ClO⁻	sodium hypochlorite, calcium hypochlorite
+3	chlorites	ClO− 2	sodium chlorite
+4	chlorine(IV)	ClO 2	chlorine dioxide
+5	chloryl, chlorates	ClO− 3 ClO+ 2	potassium chlorate, chloric acid, dichloryl trisulfate [ClO2]₂[S₃O₁₀].
+6	chlorine(VI)	Cl 2O 6	dichlorine hexoxide (gas). In liquid or solid disproportionates to mix of +5 and +7 oxidation states, as ionic chloryl perchlorate [ClO 2]+ [ClO 4]−
+7	perchlorates	ClO− 4	perchloric acid, perchlorate salts such as magnesium perchlorate, dichlorine heptoxide

匿名

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案内

chlorine