WordNet

an abundant tasteless odorless multivalent nonmetallic element; best known in yellow crystals; occurs in many sulphide and sulphate minerals and even in native form (especially in volcanic regions) (同)S, sulphur, atomic number 16
treat with sulphur in order to preserve; "These dried fruits are sulphured" (同)sulfur

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出典(authority):フリー百科事典『ウィキペディア（Wikipedia）』「2016/04/06 01:12:19」(JST)

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「いおう」はこの項目へ転送されています。たかみ型掃海艇の2番艇については「いおう (掃海艇)」をご覧ください。

硫黄（いおう、英: sulfur, sulphur, 羅: sulphur）は原子番号 16、原子量 32.1 の元素である。元素記号は S。酸素族元素の一つ。多くの同素体や結晶多形が存在し、融点、密度はそれぞれ異なる。沸点 444.674 ℃。大昔から存在が知られており、発見者は不明になっている。硫黄の英名 sulphur は、ラテン語で「燃える石」を意味する言葉に語源を持っている^[2]。

用途

硫黄から製造される硫酸は化学工業上、最も重要な酸である。一般的に酸として用いられるのは希硫酸、脱水剤や乾燥剤に用いられるのは濃硫酸である。また、種々の硫黄を含んだ化合物が合成されている。

硫黄は黒色火薬の原料であり、合成繊維、医薬品や農薬、また抜染剤などの重要な原料であり、さまざまな分野で硫化物や各種の化合物が構成されている。農家における干し柿、干しイチジクなどの漂白剤には、硫黄を燃やして得る二酸化硫黄が用いられる（燻蒸して行われる）。

ゴムに数%の硫黄を加えて加熱すると（架橋により）弾性が増し、さらに添加量を増やすと硬さを増して行き、最終的にはエボナイトとなる。

また、金属の硫化鉱物は半導体の性質を示すものが多く、シリコン鉱石検波機やゲルマニウムダイオードが実用化される以前は、鉱石検波機の主要部品として重用された。

同素体

S₈硫黄

硫黄は30以上の同素体を形成するが、これは他の元素に比べてもかなり多い^[3]。通常、天然に見られる同素体は環状の S₈ 硫黄である^[4]。

常温、常圧で固体である S₈ 硫黄は3つの結晶形を持つ。

α硫黄（斜方硫黄） - 融点 112.8 ℃、比重 2.07、淡黄色斜方晶
β硫黄（単斜硫黄） - 融点 119.6 ℃、比重 1.96、淡黄色単斜晶
γ硫黄（単斜硫黄） - 融点 106.8 ℃、比重 1.955、淡黄針状晶

いずれも、S₈ 硫黄を単位構造とする結晶であるが、95.6 ℃以下では斜方硫黄が安定であり、それ以上の温度では単斜硫黄系が安定である。また、250 ℃まで加熱すると50万個以上の硫黄原子が繋がった直鎖状硫黄 (S_n) となる。これはゴム状硫黄またはプラスチック硫黄とも呼ばれる。

性質

硫黄は融解すると血赤色の液体となり、燃やすと青い炎を上げる

S₈ 硫黄は融点直上の温度では黄色をしており、粘性も低いが、温度が上昇するにつれて直鎖状硫黄へと変化が進み、159.4 ℃以上では暗赤色となり粘性が増大し殆ど流動性を失う。この温度以上では S₈ 硫黄の環が解裂し、直鎖状のビラジカルが発生し、直鎖状 S₁₆, S₂₄ などのオリゴマー化が進行し直鎖状硫黄 (S_n) が形成され粘性が急速に増大する。さらに加温すると、直鎖状の分子が切れて再び流動性を取り戻し、沸点の444.674 ℃にいたる。暗赤色の150 - 195 ℃の硫黄を冷水に投入すると、褐色を帯びたゴム状硫黄が得られる。鉄分など不純物を含む場合は黒褐色、不純物が微量である（純度が99 %を超える）場合は黄色のゴム状硫黄となるという報告も成されているが^[5]^{[リンク切れ]}、実際の硫黄の研究においては純度99.9999%以上の原料などが用いられており、褐色を帯びるのを単に不純物に帰するのは不正確であると言える。そもそもゴム状硫黄（amorphous sulfur）と呼ばれる物質は高温でS₈の環状構造が開裂、様々な長さの鎖状構造や、濃い色を示すS₃などの小さな分子の混合体となったものであり、その生成時の加熱温度や冷却速度などにより異なる組成を示す。このため色や粘度、ヤング率などの物理特性は合成条件に大きく依存する。例えば急速圧縮法^[6]を利用すると黄色透明なゴム状硫黄が得られるが^[7]、これは通常の手法で得られる褐色のゴム状硫黄とは熱力学的な特性が大きく異なるアモルファス相である。つまり、不純物により着色すると言うよりは、「ゴム状硫黄」としてまとめられている不定形化合物には様々なものが存在し、作り方によっては黄色透明な種類のゴム状硫黄も作成可能であったり、不純物の存在によりS₈環の開裂や鎖状構造の伸張・再開裂速度が異なり同じ加熱時間でも異なる組成のものが生成する、と捉えた方がよい。なお、準安定状態であるゴム状硫黄は放置すると斜方硫黄に徐々に変化していく。

他の同素体として、硫黄蒸気の分子量測定から S₂, S₄, S₆, S₇ などが存在することが判明している。また、ハッブル宇宙望遠鏡での木星の衛星「イオ」のスペクトル観測では S₂, S₃, S₄ の存在が観測されている。2200 ℃以上、低圧下では原子状硫黄が主となる^[8]。

また、硫黄の同素体は環状硫黄分子として人為的に合成されてきており、シクロ-S₆ を筆頭に、シクロ-S₇、シクロ-S₉、シクロ-S₁₀、シクロ-S₁₁、シクロ-S₁₂、シクロ-S₁₈、シクロ-S₂₀ 等が合成され、X線結晶構造解析でその構造が確認されている。

水には溶けにくいが、二硫化炭素に溶解しやすく、ベンゼンおよびトルエンにも少量溶解する。アルカリ水溶液と加熱すると多硫化物およびチオ硫酸塩を生じて溶解する。金、白金以外の多くの金属と反応して硫化物を形成する。銀や銅とは接触により室温でも反応して黒色の硫化銀や硫化銅を生成する。

シクロ-S₆ はアルケンの硫化に用いる際の反応性が S₈ 硫黄より高いことが知られている。

硫黄は臭気を発しないが、噴火口や硫黄泉の周囲など天然の硫黄が存在する場所で多く発生する硫黄化合物である硫化水素や二酸化硫黄は刺激臭があり、日本語ではこれらの臭気を俗に「硫黄の臭い」と表現することがある^[9]。また、硫黄は空気中で表面がわずかに酸化されて不安定な酸化物を作る。この酸化物がさらに酸化段階が高くなる過程で水分の存在で硫黄を還元し、硫化水素を生成する。こうして発生した硫黄酸化物と硫化水素の臭いが「硫黄の臭い」と誤認されている。

硫黄の所在・製法

イジェン火山(Ijen)での硫黄採取。赤い流れは液体の硫黄。

天然には数多くの硫黄鉱物（硫化鉱物、硫酸塩鉱物）として産出する。単体でも産出する（自然硫黄）。深海では熱水噴出口付近で鉄などの金属と結合した硫化物や温泉（硫黄泉）では硫黄が昇華した硫黄華や、湯の花としてコロイド状硫黄が見られ、白く濁って見える。そして人体では硫黄を含むシステインや必須アミノ酸のメチオニンとして存在する。

火山性ガスには硫化水素、二酸化硫黄が含まれ、それが冷えると硫黄が析出する。これを応用したのが昇華硫黄（火口硫黄ともいう）であり、噴気孔から石で煙道を造り、内部に適宜石を入れて、この石に昇華した硫黄を付着させる採取法であった。ガスから分離し、煙道内に溜まった硫黄は最初の内は液状であるが、温度の低下に伴い次第に粘度を増してゆき、採取口に近づく頃にはほぼ固化した状態で純度の高い硫黄が得られた。那須岳、十勝岳、九重山などの活火山ではこのような方法で硫黄採掘に従事する鉱山が点在していた。これとは別に、鉱床から得られる硫黄も存在しており、こちらは採掘・選鉱した後、製錬所において焼き釜に鉱石を入れて硫黄分を溶出させていた。

2 H₂S + SO₂ → 3 S + 2 H₂O

単体硫黄を産出することで、古来からイタリアのシチリア島が有名である。また現代ではハーマン・フラッシュが1891年に開発した、165 °Cの過熱水蒸気を鉱床に吹き込み硫黄をガスとして回収するフラッシュ法で、アメリカのテキサス州やルイジアナ州、メキシコ、チリ、南アフリカの鉱山で大量に採掘される^[10]。取り出されたガスを冷やすと硫黄が析出する。この方法は、上記の火山性ガスからの硫黄の析出の逆反応である。

3 S + 2 H₂O → 2 H₂S + SO₂（高温で進行）

2 H₂S + SO₂ → 3 S + 2H₂O（低温で進行）

この他に、火口湖の湖底から硫黄を採取する方法も取られた。この場合は、湖上に浚渫船を浮かべ、湖底に沈殿している硫黄分を多く含む泥を採取していた。

また石油精製の脱硫による副産物として大量の硫黄が供給されている。石油精製における製法については硫黄回収装置の項に説明されている。

日本での硫黄の生産

知床硫黄山の噴煙

日本には火山が多く、火口付近に露出する硫黄を露天掘りにより容易に採掘することが可能であることから、古くから硫黄の生産が行われ、8世紀の「続日本紀」には信濃国（長野県米子鉱山）から朝廷へ硫黄の献上があったことが記されている。鉄砲の伝来により火薬の材料として、中世以降は日本各地の硫黄鉱山開発が活発になった。江戸時代には硫黄付け木として火を起こすのに用いられた。明治期の産業革命に至り鉱山開発は本格化する。明治時代においては安田財閥は釧路の硫黄（アトサヌプリ#硫黄鉱山を参照）で築かれたと揶揄されるほどであった。純度の高い国産硫黄は、マッチ（当時の主要輸出品目）の材料に大量に用いられ、各地の鉱山開発に拍車が掛かった。1889年には知床硫黄山が噴火と共にほぼ純度100 %の溶解硫黄を沢伝いに海まで流出させるほど大量に産出したため、当時未踏の地だった同地に鉱業関係者が殺到したという。

昭和20年代の朝鮮戦争時には「黄色いダイヤ」と呼ばれるほど硫黄価格が高騰し、鉱工業の花形に成長したが、昭和30年代に入ると資源の枯渇に加え、石油の脱硫装置からの硫黄生産が可能となったことで生産方法は一変する。エネルギー転換に加え大気汚染の規制が強化されたことから、石油精製の過程で発生する硫黄の生産も急増し、硫黄の生産者価格の下落が続いた結果、昭和40年代半ばには国内の硫黄鉱山は全て閉山に追い込まれた（岩手県の松尾鉱山など）。現在、国内に流通している硫黄は、全量が脱硫装置起源のものである。

硫黄の化合物

硫黄のオキソ酸

詳細は「硫黄のオキソ酸」を参照

硫黄は数種のオキソ酸を作る。最も有名なのものに硫酸 H₂SO₄ がある。

その他の硫黄化合物

硫化水素 (H₂S)
二酸化硫黄 (SO₂)
三酸化硫黄 (SO₃)
六フッ化硫黄 (SF₆)
二塩化硫黄 (SCl₂)
二硫化炭素 (CS₂)
有機硫黄化合物

塩

チオ硫酸ナトリウム（ハイポ） (Na₂S₂O₃)
硫酸バリウム (BaSO₄)

生物における硫黄化合物

硫黄化合物は生物でも不可欠な役割を果たしている。ビタミンB1とB7 (ビオチン、ビタミンHとも) に含まれる。

植物の根では、硫黄は硫酸イオンの形で吸収され、還元されて最終的に硫化水素となってから、システインやその他の有機化合物に取り込まれる。

アミノ酸ではシステインとメチオニンが硫黄を含み、それらがさらにペプチド・蛋白質に取り込まれる。そのほか含硫アミノ酸としてはホモシステインとタウリンがあり、これらはペプチド・蛋白質には取り込まれないが代謝上は重要である。

蛋白質のシステイン残基にあるチオール基は、システインプロテアーゼなどの活性中心として機能する。また1対のシステイン残基の間にジスルフィド結合（S-S結合）が形成され、蛋白質の高次構造（三次構造・四次構造）を形成・維持する上で重要である。顕著な例として、羽毛や毛髪が力学的・化学的に頑丈なのは、主要蛋白質ケラチンに多数の S-S 結合が含まれていることが大きな要因である。これらを燃やしたときの特異なにおい、またゆで卵のにおいも、主に硫黄化合物による。

硫黄を含む低分子ペプチドとして特に重要なのはグルタチオンで、細胞内でそのチオール基により還元剤として、あるいは解毒代謝に働いている。またアシル基に関係した多くの反応は、例えば補酵素A、α-リポ酸などの、チオール基を含む補欠分子を必要とする。

一部の光合成・化学合成細菌では、硫化水素が水の代わりに電子供与体として使われる。多くの生物の電子伝達系で、硫黄と鉄からなる鉄-硫黄クラスターが働いている（フェレドキシンなど）。また呼吸鎖のシトクロムc酸化酵素の銅中心 Cu_A にも含まれる。

脚注

[ヘルプ]

^ Magnetic susceptibility of the elements and inorganic compounds, in Handbook of Chemistry and Physics. CRC press. (2000). ISBN 0849304814.
^ ロナルド・ルイス・ボネウィッツ著、青木正博訳『ROCK and GEM 岩石と宝石の大図鑑』誠文堂新光社 2007年 120ページ
^ Ralf Steudel, Bodo Eckert (2003). “Solid Sulfur Allotropes Sulfur Allotropes”. Topics in Current Chemistry 230: 1–80. doi:10.1007/b12110.
^ Steudel, R. (1982). “Homocyclic Sulfur Molecules”. Topics Curr. Chem. 102: 149.
^ ゴム状硫黄「黄色」です―17歳が実験、教科書変えた 2009年1月5日閲覧
^ S. M. Hong, L. Y. Chen, X. R. Liu, X. H. Wu and L. Su, Rev. Sci. Instrum., 76, 053905 (2005).
^ P. Yu, W. H. Wang, R. J. Wang, S. X. Lin, X. R. Liu, S. M. Hong and H. Y. Bai, App. Phys. Lett., 94, 011910 (2009).
^ 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
^ 温泉などの「硫黄の臭い」は，本当は何のにおいなのか？ - Q&Aで学ぶ/第1章物質の構成、浜島書店
^ 硫黄の主な生産国は、アメリカ、カナダ、ポーランド、フランス、ロシア、メキシコ、日本である

参考文献

エリック・シャリーン　『図説　世界史を変えた50の鉱物』　上原ゆうこ訳　原書房、2013年、ISBN 978-4-562-04871-7。

外部リンク

硫黄化合物の合成技術により様々な製品が製造されている参考：旭化学工業㈱
硫黄 - 「健康食品」の安全性・有効性情報（国立健康・栄養研究所）
硫黄　理科ねっとわーく（一般公開版） - 科学技術振興機構
sulfur (Chemical element) - ブリタニカ百科事典
日本大百科全書(ニッポニカ)『硫黄』 - コトバンク

wiki en

[Wiki en表示]

Sulfur or sulphur (see spelling differences) is a chemical element with symbol S and atomic number 16. It is an abundant, multivalent non-metal. Under normal conditions, sulfur atoms form cyclic octatomic molecules with chemical formula S₈. Elemental sulfur is a bright yellow crystalline solid at room temperature. Chemically, sulfur reacts with all elements except for nitrogen and the noble gases.

Elemental sulfur occurs naturally as the element (native sulfur), but most commonly occurs in combined forms as sulfide and sulfate minerals. Being abundant in native form, sulfur was known in ancient times, being mentioned for its uses in ancient India, ancient Greece, China, and Egypt. Sulfur is referred to in the Bible as brimstone.^[5] Today, almost all elemental sulfur is produced as a byproduct of removing sulfur-containing contaminants from natural gas and petroleum. The element's largest commercial use (after mostly being converted to sulfuric acid) is to produce sulfate and phosphate fertilizers, because of the relatively high requirement of plants for sulfur and phosphorus. Sulfuric acid is also a primary industrial chemical outside fertilizer manufacture. Other uses for the element are in matches, insecticides and fungicides. Many sulfur compounds are odoriferous, and the smells of odorized natural gas, skunk scent, grapefruit, and garlic are due to organosulfur compounds. Hydrogen sulfide imparts the characteristic odor to rotting eggs and other biological processes.

Sulfur is an essential element for all life, but almost always in the form of organosulfur compounds or metal sulfides. Three amino acids (cysteine, cystine, and methionine) and two vitamins (biotin and thiamine) are organosulfur compounds. Many cofactors also contain sulfur including glutathione and thioredoxin and iron–sulfur proteins. Disulfides, S–S bonds, confer mechanical strength and insolubility of the protein keratin, found in outer skin, hair, and feathers. Elemental sulfur is not common in higher forms of life, but is both a product and an oxidant for various bacteria.

1 Spelling and etymology
2 Characteristics
- 2.1 Physical properties
- 2.2 Chemical properties
- 2.3 Allotropes
- 2.4 Isotopes
- 2.5 Natural occurrence
3 Production
4 Compounds
- 4.1 Sulfur polycations
- 4.2 Sulfides
- 4.3 Oxides, oxoacids and oxoanions
- 4.4 Halides and oxyhalides
- 4.5 Pnictides
- 4.6 Metal sulfides
- 4.7 Organic compounds
5 History
- 5.1 Antiquity
- 5.2 Modern times
6 Applications
- 6.1 Sulfuric acid
- 6.2 Other large-scale sulfur chemicals
- 6.3 Fertilizer
- 6.4 Fine chemicals
- 6.5 Fungicide and pesticide
- 6.6 Bactericide in winemaking and food preservation
- 6.7 Pharmaceutical use
  - 6.7.1 Mechanism of action
7 Biological role
- 7.1 Protein and organic cofactors
- 7.2 Metalloproteins and inorganic cofactors
- 7.3 Sulfur metabolism and the sulfur cycle
8 Precautions
9 See also
10 References
11 External links

Spelling and etymology

Sulfur is derived from the Latin word sulpur, which was Hellenized to sulphur. The spelling sulfur appears toward the end of the Classical period. (The true Greek word for sulfur, θεῖον, is the source of the international chemical prefix thio-.) In 12th-century Anglo-French, it was sulfre; in the 14th century the Latin ph was restored, for sulphre; and by the 15th century the full Latin spelling was restored, for sulfur, sulphur. The parallel f~ph spellings continued in Britain until the 19th century, when the word was standardized as sulphur.^[6] Sulfur was the form chosen in the United States, whereas Canada uses both. The IUPAC adopted the spelling sulfur in 1990, as did the Nomenclature Committee of the Royal Society of Chemistry in 1992, restoring the spelling sulfur to Britain.^[7] The Oxford Dictionaries note that "in chemistry ... the -f- spelling is now the standard form in all related words in the field in both British and US contexts."^[8]

The late Latin form also continues in the Romance languages: French soufre, Italian zolfo (from solfo), Spanish azufre (from açufre, from earlier çufre), Portuguese enxofre (from xofre). The Spanish and Portuguese forms are prefixed with the Arabic article, despite not being Arabic words.^[6] The root has been traced back to reconstructed proto-Indo-European *swépl̥ (genitive *sulplós), a nominal derivative of *swelp 'to burn', a lineage also preserved in the Germanic languages, where it is found for example as modern German Schwefel, Dutch zwavel, and Swedish svavel, and as Old English swefl.^[9]

Characteristics

When burned, sulfur melts to a blood-red liquid and emits a blue flame that is best observed in the dark.

Physical properties

Sulfur forms polyatomic molecules with different chemical formulas, with the best-known allotrope being octasulfur, cyclo-S₈. The point group of cyclo-S₈ is D_4d and its dipole moment is 0 D.^[10] Octasulfur is a soft, bright-yellow solid with only a faint odor, similar to that of matches.^[11] It melts at 115.21 °C (239.38 °F), boils at 444.6 °C (832.3 °F) and sublimes easily.^[5] At 95.2 °C (203.4 °F), below its melting temperature, cyclo-octasulfur changes from α-octasulfur to the β-polymorph.^[12] The structure of the S₈ ring is virtually unchanged by this phase change, which affects the intermolecular interactions. Between its melting and boiling temperatures, octasulfur changes its allotrope again, turning from β-octasulfur to γ-sulfur, again accompanied by a lower density but increased viscosity due to the formation of polymers.^[12] At even higher temperatures, the viscosity decreases as depolymerization occurs. Molten sulfur assumes a dark red color above 200 °C (392 °F). The density of sulfur is about 2 g·cm⁻³, depending on the allotrope; all of its stable allotropes are excellent electrical insulators.

Chemical properties

Sulfur burns with a blue flame with formation of sulfur dioxide, which has a suffocating odor. Sulfur is insoluble in water but soluble in carbon disulfide and, to a lesser extent, in other nonpolar organic solvents, such as benzene and toluene. The first and second ionization energies of sulfur are 999.6 and 2252 kJ·mol⁻¹, respectively. Despite such figures, the +2 oxidation state is rare, with +4 and +6 being more common. The fourth and sixth ionization energies are 4556 and 8495.8 kJ·mol⁻¹, the magnitude of the figures caused by electron transfer between orbitals; these states are only stable with strong oxidants such as fluorine, oxygen, and chlorine.^{[citation needed]}

Allotropes

The structure of the cyclooctasulfur molecule, S₈.

Sulfur forms over 30 solid allotropes, more than any other element.^[13] Besides S₈, several other rings are known.^[14] Removing one atom from the crown gives S₇, which is more deeply yellow than S₈. HPLC analysis of "elemental sulfur" reveals an equilibrium mixture of mainly S₈, but with S₇ and small amounts of S₆.^[15] Larger rings have been prepared, including S₁₂ and S₁₈.^[16]^[17]

Amorphous or "plastic" sulfur is produced by rapid cooling of molten sulfur—for example, by pouring it into cold water. X-ray crystallography studies show that the amorphous form may have a helical structure with eight atoms per turn. The long coiled polymeric molecules make the brownish substance elastic, and in bulk this form has the feel of crude rubber. This form is metastable at room temperature and gradually reverts to crystalline molecular allotrope, which is no longer elastic. This process happens within a matter of hours to days, but can be rapidly catalyzed.

Isotopes

Sulfur has 25 known isotopes, four of which are stable: ³²S (95.02%), ³³S (0.75%), ³⁴S (4.21%), and ³⁶S (0.02%). Other than ³⁵S, with a half-life of 87 days and formed in cosmic ray spallation of ⁴⁰Ar, the radioactive isotopes of sulfur have half-lives less than 3 hours.

When sulfide minerals are precipitated, isotopic equilibration among solids and liquid may cause small differences in the δS-34 values of co-genetic minerals. The differences between minerals can be used to estimate the temperature of equilibration. The δC-13 and δS-34 of coexisting carbonate minerals and sulfides can be used to determine the pH and oxygen fugacity of the ore-bearing fluid during ore formation.

In most forest ecosystems, sulfate is derived mostly from the atmosphere; weathering of ore minerals and evaporites contribute some sulfur. Sulfur with a distinctive isotopic composition has been used to identify pollution sources, and enriched sulfur has been added as a tracer in hydrologic studies. Differences in the natural abundances can be used in systems where there is sufficient variation in the ³⁴S of ecosystem components. Rocky Mountain lakes thought to be dominated by atmospheric sources of sulfate have been found to have different δ³⁴S values from lakes believed to be dominated by watershed sources of sulfate.

Natural occurrence

Nearly exhausted sulfur vat from which railroad cars are loaded, Freeport Sulphur Co., Hoskins Mound, Texas (1943).

Most of the yellow and orange hues of Io are due to elemental sulfur and sulfur compounds, produced by active volcanoes.

A man carrying sulfur blocks from Kawah Ijen, a volcano in East Java, Indonesia, 2009

³²S is created inside massive stars, at a depth where the temperature exceeds 2.5×10⁹ K, by the fusion of one nucleus of silicon plus one nucleus of helium.^[18] As this is part of the alpha process that produces elements in abundance, sulfur is the 10th most common element in the universe.

Sulfur, usually as sulfide, is present in many types of meteorites. Ordinary chondrites contain on average 2.1% sulfur, and carbonaceous chondrites may contain as much as 6.6%. It is normally present as troilite (FeS), but there are exceptions, with carbonaceous chondrites containing free sulfur, sulfates and other sulfur compounds.^[19] The distinctive colors of Jupiter's volcanic moon Io are attributed to various forms of molten, solid and gaseous sulfur.^[20]

Sulfur occurs in fumaroles such as this one in Vulcano, Italy

On Earth, elemental sulfur can be found near hot springs and volcanic regions in many parts of the world, especially along the Pacific Ring of Fire; such volcanic deposits are currently mined in Indonesia, Chile, and Japan. Such deposits are polycrystalline, with the largest documented single crystal measuring 22×16×11 cm.^[21] Historically, Sicily was a large source of sulfur in the Industrial Revolution.^[22]

Native sulfur is synthesised by anaerobic bacteria acting on sulfate minerals such as gypsum in salt domes.^[23]^[24] Significant deposits in salt domes occur along the coast of the Gulf of Mexico, and in evaporites in eastern Europe and western Asia. Native sulfur may be produced by geological processes alone. Fossil-based sulfur deposits from salt domes were until recently the basis for commercial production in the United States, Russia, Turkmenistan, and Ukraine.^[25] Currently, commercial production is still carried out in the Osiek mine in Poland. Such sources are now of secondary commercial importance, and most are no longer worked.

Common naturally occurring sulfur compounds include the sulfide minerals, such as pyrite (iron sulfide), cinnabar (mercury sulfide), galena (lead sulfide), sphalerite (zinc sulfide) and stibnite (antimony sulfide); and the sulfates, such as gypsum (calcium sulfate), alunite (potassium aluminium sulfate), and barite (barium sulfate). On Earth, just as upon Jupiter's moon Io, elemental sulfur occurs naturally in volcanic emissions, including emissions from hydrothermal vents.

Production

Sulfur may be found by itself and historically was usually obtained in this form; pyrite has also been a source of sulfur.^[26] In volcanic regions in Sicily, in ancient times, it was found on the surface of the Earth, and the "Sicilian process" was used: sulfur deposits were piled and stacked in brick kilns built on sloping hillsides, with airspaces between them. Then, some sulfur was pulverized, spread over the stacked ore and ignited, causing the free sulfur to melt down the hills. Eventually the surface-borne deposits played out, and miners excavated veins that ultimately dotted the Sicilian landscape with labyrinthine mines. Mining was unmechanized and labor-intensive, with pickmen freeing the ore from the rock, and mine-boys or carusi carrying baskets of ore to the surface, often through a mile or more of tunnels. Once the ore was at the surface, it was reduced and extracted in smelting ovens. The conditions in Sicilian sulfur mines were horrific, prompting Booker T. Washington to write "I am not prepared just now to say to what extent I believe in a physical hell in the next world, but a sulphur mine in Sicily is about the nearest thing to hell that I expect to see in this life."^[27]

Elemental sulfur was once extracted from salt domes where it sometimes occurs in nearly pure form, but this method has been obsolete since the late 20th century. Today's sulfur production is as a side product of other industrial processes such as oil refining; in these processes, sulfur often occurs as undesired or detrimental compounds that are extracted and converted to elemental sulfur. As a mineral, native sulfur under salt domes is thought to be a fossil mineral resource, produced by the action of ancient bacteria on sulfate deposits. It was removed from such salt-dome mines mainly by the Frasch process.^[25] In this method, superheated water was pumped into a native sulfur deposit to melt the sulfur, and then compressed air returned the 99.5% pure melted product to the surface. Throughout the 20th century this procedure produced elemental sulfur that required no further purification. Due to a limited number of such sulfur deposits and the high cost of working them, this process for mining sulfur has not been employed in a major way anywhere in the world since 2002.^[28]^[29]

Sulfur recovered from hydrocarbons in Alberta, stockpiled for shipment in North Vancouver, British Columbia

Today, sulfur is produced from petroleum, natural gas, and related fossil resources, from which it is obtained mainly as hydrogen sulfide. Organosulfur compounds, undesirable impurities in petroleum, may be upgraded by subjecting them to hydrodesulfurization, which cleaves the C–S bonds:^[28]^[29]

R-S-R + 2 H₂ → 2 RH + H₂S

The resulting hydrogen sulfide from this process, and also as it occurs in natural gas, is converted into elemental sulfur by the Claus process. This process entails oxidation of some hydrogen sulfide to sulfur dioxide and then the comproportionation of the two:^[28]^[29]

3 O₂ + 2 H₂S → 2 SO₂ + 2 H₂O

SO₂ + 2 H₂S → 3 S + 2 H₂O

Production and price (US market) of elemental sulfur

Owing to the high sulfur content of the Athabasca Oil Sands, stockpiles of elemental sulfur from this process now exist throughout Alberta, Canada.^[30] Another way of storing sulfur is as a binder for concrete, the resulting product having many desirable properties (see sulfur concrete).^[31] Sulfur is still mined from surface deposits in poorer nations with volcanoes, such as Indonesia, and worker conditions have not improved much since Booker T. Washington's days.^[32]

The world production of sulfur in 2011 amounted to 69 million tonnes (Mt), with more than 15 countries contributing more than 1 Mt each. Countries producing more than 5 Mt are China (9.6), US (8.8), Canada (7.1) and Russia (7.1).^[33] Production has been slowly increasing from 1900 to 2010; the price was unstable in the 1980s and around 2010.^[34]

Compounds

Common oxidation states of sulfur range from −2 to +6. Sulfur forms stable compounds with all elements except the noble gases.

Sulfur polycations

Sulfur polycations, S₈²⁺, S₄²⁺ and S₁₆²⁺ are produced when sulfur is reacted with mild oxidising agents in a strongly acidic solution.^[35] The colored solutions produced by dissolving sulfur in oleum were first reported as early as 1804 by C.F Bucholz, but the cause of the color and the structure of the polycations involved was only determined in the late 1960s. S₈²⁺ is deep blue, S₄²⁺ is yellow and S₁₆²⁺ is red.^[12]

Sulfides

Treatment of sulfur with hydrogen gives hydrogen sulfide. When dissolved in water, hydrogen sulfide is mildly acidic:^[5]

H₂S HS⁻ + H⁺

Hydrogen sulfide gas and the hydrosulfide anion are extremely toxic to mammals, due to their inhibition of the oxygen-carrying capacity of hemoglobin and certain cytochromes in a manner analogous to cyanide and azide (see below, under precautions).

Reduction of elemental sulfur gives polysulfides, which consist of chains of sulfur atoms terminated with S⁻ centers:

2 Na + S₈ → Na₂S₈

This reaction highlights arguably the single most distinctive property of sulfur: its ability to catenate (bind to itself by formation of chains). Protonation of these polysulfide anions gives the polysulfanes, H₂S_x where x = 2, 3, and 4.^[36] Ultimately reduction of sulfur gives sulfide salts:

16 Na + S₈ → 8 Na₂S

The interconversion of these species is exploited in the sodium-sulfur battery. The radical anion S₃⁻ gives the blue color of the mineral lapis lazuli.

Lapis lazuli owes its blue color to a trisulfur radical anion (S−

3)

Two parallel sulfur chains grown inside a single-wall carbon nanotube (CNT, a). Zig-zag (b) and straight (c) S chains inside double-wall CNTs.^[37]

Oxides, oxoacids and oxoanions

The principal sulfur oxides are obtained by burning sulfur:

S + O₂ → SO₂

2 SO₂ + O₂ → 2 SO₃

Other oxides are known, sulfur-rich oxides e.g. sulfur monoxide and disulfur mono- and dioxides and higher oxides containing peroxo groups.

Sulfur forms sulfur oxoacids, some of which cannot be isolated and are only known through their salts. Sulfur dioxide and sulfites (SO2−
3) are related to the unstable sulfurous acid (H₂SO₃). Sulfur trioxide and sulfates (SO2−
4) are related to sulfuric acid. Sulfuric acid and SO₃ combine to give oleum, a solution of pyrosulfuric acid (H₂S₂O₇) in sulfuric acid.

Thiosulfate salts (S
2O2−
3), sometimes referred as "hyposulfites", used in photographic fixing (hypo) and as reducing agents, feature sulfur in two oxidation states. Sodium dithionite (Na
2S
2O
4), contains the more highly reducing dithionite anion (S
2O2−
4).

Halides and oxyhalides

The two main sulfur fluorides are sulfur hexafluoride, a dense gas used as an insulator gas in high voltage transformers and a nonreactive and nontoxic propellant, and sulfur tetrafluoride, a rarely used organic reagent that is highly toxic.^[38] Sulfur dichloride and disulfur dichloride are important industrial chemicals. Sulfuryl chloride and chlorosulfuric acid are derivatives of sulfuric acid; thionyl chloride (SOCl₂) is a common reagent in organic synthesis.^[39]

Pnictides

An important S–N compound is the cage tetrasulfur tetranitride (S₄N₄). Heating this compound gives polymeric sulfur nitride ((SN)_x), which has metallic properties even though it does not contain any metal atoms. Thiocyanates contain the SCN⁻ group. Oxidation of thiocyanate gives thiocyanogen, (SCN)₂ with the connectivity NCS-SCN. Phosphorus sulfides are numerous, the most important commercially being the cages P₄S₁₀ and P₄S₃.^[40]^[41]

Metal sulfides

The principal ores of copper, zinc, nickel, cobalt, molybdenum, and other metals are sulfides. These materials tend to be dark-colored semiconductors that are not readily attacked by water or even many acids. They are formed, both geochemically and in the laboratory, by the reaction of hydrogen sulfide with metal salts. The mineral galena (PbS) was the first demonstrated semiconductor and found a use as a signal rectifier in the cat's whiskers of early crystal radios. The iron sulfide called pyrite, the so-called "fool's gold," has the formula FeS₂.^[42] The upgrading of these ores, usually by roasting, is costly and environmentally hazardous. Sulfur corrodes many metals via the process called tarnishing.

Organic compounds

Illustrative organosulfur compounds
Allicin, the active ingredient in garlic
(R)-cysteine, an amino acid containing a thiol group
Methionine, an amino acid containing a thioether
Diphenyl disulfide, a representative disulfide
Perfluorooctanesulfonic acid, a controversial surfactant
Dibenzothiophene, a component of crude oil
Penicillin a antibiotic where "R" is the variable group

Some of the main classes of sulfur-containing organic compounds include the following:^[43]

Thiols or mercaptans (as they are mercury capturers as chelators) are the sulfur analogs of alcohols; treatment of thiols with base gives thiolate ions.
Thioethers are the sulfur analogs of ethers.
Sulfonium ions have three groups attached to a cationic sulfur center. Dimethylsulfoniopropionate (DMSP) is one such compound, important in the marine organic sulfur cycle.
Sulfoxides and sulfones are thioethers with one and two oxygen atoms attached to the sulfur atom, respectively. The simplest sulfoxide, dimethyl sulfoxide, is a common solvent; a common sulfone is sulfolane.
Sulfonic acids are used in many detergents.

Compounds with carbon–sulfur multiple bonds are uncommon with the exception of carbon disulfide, a volatile colorless liquid that is structurally similar to carbon dioxide. It is used as a reagent to make the polymer rayon and many organosulfur compounds. Unlike carbon monoxide, carbon monosulfide is only stable as a dilute gas, as in the interstellar medium.^[44]

Organosulfur compounds are responsible for some of the unpleasant odors of decaying organic matter. They are used in the odoration of natural gas and cause the odor of garlic and skunk spray. Not all organic sulfur compounds smell unpleasant at all concentrations: the sulfur-containing monoterpenoid grapefruit mercaptan in small concentrations is responsible for the characteristic scent of grapefruit, but has a generic thiol odor at larger concentrations. Sulfur mustard, a potent vesicant, was used in World War I as a disabling agent.^[45]

Sulfur-sulfur bonds are a structural component to stiffen rubber, in a way similar to the biological role of disulfide bridges to rigidify proteins (see biological below). In the most common type of industrial "curing" or hardening and strengthening of natural rubber, elemental sulfur is heated with the rubber to the point that chemical reactions form disulfide bridges between isoprene units of the polymer. This process, patented in 1843, allowed rubber to become a major industrial product, especially in automobile tires. Because of the heat and sulfur, the process was named vulcanization, after the Roman god of the forge and volcanism.

History

Antiquity

Pharmaceutical container for sulfur from the first half of the 20th century. From the Museo del Objeto del Objeto collection

Being abundantly available in native form, sulfur was known in ancient times and is referred to in the Torah (Genesis). English translations of the Bible commonly referred to burning sulfur as "brimstone", giving rise to the term "fire-and-brimstone" sermons, in which listeners are reminded of the fate of eternal damnation that await the unbelieving and unrepentant. It is from this part of the Bible that Hell is implied to "smell of sulfur" (likely due to its association with volcanic activity). According to the Ebers Papyrus, a sulfur ointment was used in ancient Egypt to treat granular eyelids. Sulfur was used for fumigation in preclassical Greece;^[46] this is mentioned in the Odyssey.^[47] Pliny the Elder discusses sulfur in book 35 of his Natural History, saying that its best-known source is the island of Melos. He mentions its use for fumigation, medicine, and bleaching cloth.^[48]

A natural form of sulfur known as shiliuhuang (石硫黄) was known in China since the 6th century BC and found in Hanzhong.^[49] By the 3rd century, the Chinese discovered that sulfur could be extracted from pyrite.^[49] Chinese Daoists were interested in sulfur's flammability and its reactivity with certain metals, yet its earliest practical uses were found in traditional Chinese medicine.^[49] A Song dynasty military treatise of 1044 AD described different formulas for Chinese black powder, which is a mixture of potassium nitrate (KNO
3), charcoal, and sulfur. It remains an ingredient of black gunpowder.

Indian alchemists, practitioners of "the science of mercury" (sanskrit rasaśāstra, रसशास्त्र), wrote extensively about the use of sulfur in alchemical operations with mercury, from the eighth century AD onwards.^[50] In the rasaśāstra tradition, sulfur is called "the smelly" (sanskrit gandhaka, गन्धक).

Early European alchemists gave sulfur its own alchemical symbol, a triangle at the top of a cross (, there also was in use.). In traditional skin treatment before the modern era of scientific medicine, elemental sulfur was used, mainly in creams, to alleviate conditions such as scabies, ringworm, psoriasis, eczema, and acne. The mechanism of action is unknown—though elemental sulfur does oxidize slowly to sulfurous acid, which in turn (through the action of sulfite) acts as a mild reducing and antibacterial agent.^[51]^[52]^[53]

Modern times

Sicilian kiln used to obtain sulfur from volcanic rock.

In 1777, Antoine Lavoisier helped convince the scientific community that sulfur was an element, not a compound.

Sulfur deposits in Sicily were the dominant supply source for over half a century. Approximately 2000 tons per year of sulfur were imported into Marseilles, France for the production of sulfuric acid via the Leblanc process by the late 18th century. In industrializing Britain, with the repeal of tariffs on salt in 1824, demand for sulfur from Sicily surged upward. The increasing British control and exploitation of the mining, refining and transportation of the sulfur, coupled with the failure of this lucrative export to transform Sicily's backward and impoverished economy led to the 'Sulfur Crisis' of 1840, when King Ferdinand II gave a monopoly of the sulfur industry to a French firm, violating an earlier 1816 trade agreement with Britain. A peaceful negotiated solution was eventually mediated by France.^[54]^[55]

In 1867, sulfur was discovered in underground deposits in Louisiana and Texas. The highly successful Frasch process was developed to extract this resource.^[56]

In the late 18th century, furniture makers used molten sulfur to produce decorative inlays in their craft. Because of the sulfur dioxide produced during the process of melting sulfur, the craft of sulfur inlays was soon abandoned. Molten sulfur is sometimes still used for setting steel bolts into drilled concrete holes where high shock resistance is desired for floor-mounted equipment attachment points. Pure powdered sulfur was used as a medicinal tonic and laxative.^[25] With the advent of the contact process, the majority of sulfur today is used to make sulfuric acid for a wide range of uses, particularly fertilizer.^[57]

Applications

Sulfuric acid

Elemental sulfur is mainly used as a precursor to other chemicals. Approximately 85% (1989) is converted to sulfuric acid (H₂SO₄):

2 S + 3 O₂ + 2 H₂O → 2 H₂SO₄

Sulfuric acid production in 2000

Because of its importance, sulfuric acid was considered an excellent indicator of a country's industrial well-being.^[58] For example, with 32.5 million tonnes in 2010, the United States produces more sulfuric acid every year than any other inorganic industrial chemical.^[34] The principal use for the acid is the extraction of phosphate ores for the production of fertilizer manufacturing. Other applications of sulfuric acid include oil refining, wastewater processing, and mineral extraction.^[25]

Other large-scale sulfur chemicals

Sulfur reacts directly with methane to give carbon disulfide, which is used to manufacture cellophane and rayon.^[25] One of the direct uses of sulfur is in vulcanization of rubber, where polysulfide chains crosslink organic polymers. Sulfites are heavily used to bleach paper and as preservatives in dried fruit. Many surfactants and detergents, e.g. sodium lauryl sulfate, are produced are sulfate derivatives. Calcium sulfate, gypsum, (CaSO₄^·2H₂O) is mined on the scale of 100 million tons each year for use in Portland cement and fertilizers.

When silver-based photography was widespread, sodium and ammonium thiosulfate were widely used as "fixing agents." Sulfur is a component of gunpowder.

Fertilizer

Sulfur is increasingly used as a component of fertilizers. The most important form of sulfur for fertilizer is the mineral calcium sulfate. Elemental sulfur is hydrophobic (that is, it is not soluble in water) and, therefore, cannot be directly utilized by plants. Over time, soil bacteria can convert it to soluble derivatives, which can then be utilized by plants. Sulfur improves the use efficiency of other essential plant nutrients, particularly nitrogen and phosphorus.^[59] Biologically produced sulfur particles are naturally hydrophilic due to a biopolymer coating. This sulfur is, therefore, easier to disperse over the land (via spraying as a diluted slurry), and results in a faster release.

Plant requirements for sulfur are equal to or exceed those for phosphorus. It is one of the major nutrients essential for plant growth, root nodule formation of legumes and plants' protection mechanisms. Sulfur deficiency has become widespread in many countries in Europe.^[60]^[61]^[62] Because atmospheric inputs of sulfur continue to decrease, the deficit in the sulfur input/output is likely to increase, unless sulfur fertilizers are used.

Fine chemicals

A molecular model of the pesticide malathion.

Organosulfur compounds are used in pharmaceuticals, dyestuffs, and agrochemicals. Many drugs contain sulfur, early examples being antibacterial sulfonamides, known as sulfa drugs. Sulfur is a part of many bacterial defense molecules. Most β-lactam antibiotics, including the penicillins, cephalosporins and monolactams contain sulfur.^[43]

Magnesium sulfate, known as Epsom salts when in hydrated crystal form, can be used as a laxative, a bath additive, an exfoliant, magnesium supplement for plants, or (when in dehydrated form) as a desiccant.

Fungicide and pesticide

Sulfur candle originally sold for home fumigation

Elemental sulfur is one of the oldest fungicides and pesticides. "Dusting sulfur," elemental sulfur in powdered form, is a common fungicide for grapes, strawberry, many vegetables and several other crops. It has a good efficacy against a wide range of powdery mildew diseases as well as black spot. In organic production, sulfur is the most important fungicide. It is the only fungicide used in organically farmed apple production against the main disease apple scab under colder conditions. Biosulfur (biologically produced elemental sulfur with hydrophilic characteristics) can also be used for these applications.

Standard-formulation dusting sulfur is applied to crops with a sulfur duster or from a dusting plane. Wettable sulfur is the commercial name for dusting sulfur formulated with additional ingredients to make it water miscible.^[63]^[64] It has similar applications and is used as a fungicide against mildew and other mold-related problems with plants and soil.

Elemental sulfur powder is used as an "organic" (i.e. "green") insecticide (actually an acaricide) against ticks and mites. A common method of use is to dust clothing or limbs with sulfur powder.

Diluted solutions of lime sulfur (made by combining calcium hydroxide with elemental sulfur in water), are used as a dip for pets to destroy ringworm (fungus), mange and other dermatoses and parasites. Sulfur candles consist of almost pure sulfur in blocks or pellets that are burned to fumigate structures. It is no longer used in the home due to the toxicity of the products of combustion.

Bactericide in winemaking and food preservation

Small amounts of sulfur dioxide gas addition (or equivalent potassium metabisulfite addition) to fermented wine to produce traces of sulfurous acid (produced when SO₂ reacts with water) and its sulfite salts in the mixture, has been called "the most powerful tool in winemaking."^[65] After the yeast-fermentation stage in winemaking, sulfites absorb oxygen and inhibit aerobic bacterial growth that otherwise would turn ethanol into acetic acid, souring the wine. Without this preservative step, indefinite refrigeration of the product before consumption is usually required. Similar methods go back into antiquity but modern historical mentions of the practice go to the fifteenth century. The practice is used by large industrial wine producers and small organic wine producers alike.

Sulfur dioxide and various sulfites have been used for their antioxidant antibacterial preservative properties in many other parts of the food industry also. The practice has declined since reports of an allergy-like reaction of some persons to sulfites in foods.

Pharmaceutical use

Sulfur (specifically octasulfur, S₈) is used in pharmaceutical skin preparations for the treatment of acne and other conditions. It acts as a keratolytic agent and also kills bacteria, fungi, scabies mites and other parasites.^[66] Precipitated sulfur and colloidal sulfur are used, in form of lotions, creams, powders, soaps, and bath additives, for the treatment of acne vulgaris, acne rosacea, and seborrhoeic dermatitis.^[67]

Common adverse effects include irritation of the skin at the application site, such as dryness, stinging, itching and peeling.^[68]

Mechanism of action

Sulfur is converted to hydrogen sulfide (H₂S) through reduction, partly by bacteria. H₂S kills bacteria (possibly including Propionibacterium acnes which plays a role in acne^[69]), fungi, and parasites such as scabies mites.^[66]

Biological role

Protein and organic cofactors

Sulfur is an essential component of all living cells. It is the seventh or eighth most abundant element in the human body by weight, being about as common as potassium, and a little more common than sodium or chlorine. A 70 kg (150 lb) human body contains about 140 grams of sulfur.

In plants and animals, the amino acids cysteine and methionine contain most of the sulfur. The element is thus present in all polypeptides, proteins, and enzymes that contain these amino acids. In humans, methionine is an essential amino acid that must be ingested. However, save for the vitamins biotin and thiamine, cysteine and all sulfur-containing compounds in the human body can be synthesized from methionine. The enzyme sulfite oxidase is needed for the metabolism of methionine and cysteine in humans and animals.

Disulfide bonds (S-S bonds) formed between cysteine residues in peptide chains are very important in protein assembly and structure. These covalent bonds between peptide chains confer extra toughness and rigidity.^[70] For example, the high strength of feathers and hair is in part due to their high content of S-S bonds and their high content of cysteine and sulfur. Eggs are high in sulfur because large amounts of the element are necessary for feather formation, and the characteristic odor of rotting eggs is due to hydrogen sulfide. The high disulfide bond content of hair and feathers contributes to their indigestibility and to their characteristic disagreeable odor when burned.

Homocysteine and taurine are other sulfur-containing acids that are similar in structure, but not coded by DNA, and are not part of the primary structure of proteins. Many important cellular enzymes use prosthetic groups ending with -SH moieties to handle reactions involving acyl-containing biochemicals: two common examples from basic metabolism are coenzyme A and alpha-lipoic acid.^[70] Two of the 13 classical vitamins, biotin and thiamine contain sulfur, with the latter being named for its sulfur content. Sulfur plays an important part, as a carrier of reducing hydrogen and its electrons, for cellular repair of oxidation. Reduced glutathione, a sulfur-containing tripeptide, is a reducing agent through its sulfhydryl (-SH) moiety derived from cysteine. The thioredoxins, a class of small protein essential to all known life, using neighboring pairs of reduced cysteines to act as general protein reducing agents, to similar effect.

Methanogenesis, the route to most of the world's methane, is a multistep biochemical transformation of carbon dioxide. This conversion requires several organosulfur cofactors. These include coenzyme M, CH₃SCH₂CH₂SO₃⁻, the immediate precursor to methane.^[71]

Metalloproteins and inorganic cofactors

Inorganic sulfur forms a part of iron-sulfur clusters as well as many copper, nickel, and iron proteins. Most pervasive are the ferrodoxins, which serve as electron shuttles in cells. In bacteria, the important nitrogenase enzymes contains an Fe-Mo-S cluster, is a catalyst that performs the important function of nitrogen fixation, converting atmospheric nitrogen to ammonia that can be used by microorganisms and plants to make proteins, DNA, RNA, alkaloids, and the other organic nitrogen compounds necessary for life.^[72]

Sulfur metabolism and the sulfur cycle

The sulfur cycle was the first of the biogeochemical cycles to be discovered. In the 1880s, while studying Beggiatoa (a bacterium living in a sulfur rich environment), Sergei Winogradsky found that it oxidized hydrogen sulfide (H₂S) as an energy source, forming intracellular sulfur droplets. Winogradsky referred to this form of metabolism as inorgoxidation (oxidation of inorganic compounds). He continued to study it together with Selman Waksman until the 1950s.

Sulfur oxidizers can use as energy sources reduced sulfur compounds, including hydrogen sulfide, elemental sulfur, sulfite, thiosulfate, and various polythionates (e.g., tetrathionate).^[73] They depend on enzymes such as sulfur oxygenase and sulfite oxidase to oxidize sulfur to sulfate. Some lithotrophs can even use the energy contained in sulfur compounds to produce sugars, a process known as chemosynthesis. Some bacteria and archaea use hydrogen sulfide in place of water as the electron donor in chemosynthesis, a process similar to photosynthesis that produces sugars and utilizes oxygen as the electron acceptor. The photosynthetic green sulfur bacteria and purple sulfur bacteria and some lithotrophs use elemental oxygen to carry out such oxidization of hydrogen sulfide to produce elemental sulfur (S⁰), oxidation state = 0. Primitive bacteria that live around deep ocean volcanic vents oxidize hydrogen sulfide in this way with oxygen; the giant tube worm is an example of a large organism that uses hydrogen sulfide (via bacteria) as food to be oxidized.

The so-called sulfate-reducing bacteria, by contrast, "breathe sulfate" instead of oxygen. They use organic compounds or molecular hydrogen as the energy source. They use sulfur as the electron acceptor, and reduce various oxidized sulfur compounds back into sulfide, often into hydrogen sulfide. They can grow on other partially oxidized sulfur compounds (e.g. thiosulfates, thionates, polysulfides, sulfites). The hydrogen sulfide produced by these bacteria is responsible for some of the smell of intestinal gases (flatus) and decomposition products.

Sulfur is absorbed by plants via the roots from soil as the sulfate and transported as a phosphate ester. Sulfate is reduced to sulfide via sulfite before it is incorporated into cysteine and other organosulfur compounds.^[74]

SO₄²⁻ → SO₃²⁻ → H₂S → cysteine → methionine

Precautions

Effect of acid rain on a forest, Jizera Mountains, Czech Republic

Elemental sulfur is non-toxic, as generally are the soluble sulfate salts, such as Epsom salts. Soluble sulfate salts are poorly absorbed and laxative. When injected parenterally, they are freely filtered by the kidneys and eliminated with very little toxicity in multi-gram amounts.

When sulfur burns in air, it produces sulfur dioxide. In water, this gas produces sulfurous acid and sulfites, which are antioxidants that inhibit growth of aerobic bacteria and allow its use as a food additive in small amounts. At high concentrations these acids harm the lungs, eyes or other tissues. In organisms without lungs such as insects or plants, it otherwise prevents respiration in high concentrations. Sulfur trioxide (made by catalysis from sulfur dioxide) and sulfuric acid are similarly highly corrosive, due to the strong acids that form on contact with water.

The burning of coal and/or petroleum by industry and power plants generates sulfur dioxide (SO₂), which reacts with atmospheric water and oxygen to produce sulfuric acid (H₂SO₄) and sulfurous acid (H₂SO₃). These acids are components of acid rain, which lower the pH of soil and freshwater bodies, sometimes resulting in substantial damage to the environment and chemical weathering of statues and structures. Fuel standards increasingly require that fuel producers extract sulfur from fossil fuels to prevent acid rain formation. This extracted and refined sulfur represents a large portion of sulfur production. In coal-fired power plants, flue gases are sometimes purified. More modern power plants that use synthesis gas extract the sulfur before they burn the gas.

Hydrogen sulfide is as toxic as hydrogen cyanide, and kills by the same mechanism (inhibition of the respiratory enzyme cytochrome oxidase),^[75] though hydrogen sulfide is less likely to cause surprise poisonings from small inhaled amounts, because of its disagreeable warning odor. Hydrogen sulfide quickly deadens the sense of smell-—so a victim may breathe increasing quantities and be unaware of its presence until severe symptoms occur, which can quickly lead to death. Dissolved sulfide and hydrosulfide salts are also toxic by the same mechanism.

References

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^ Standard Atomic Weights 2013. Commission on Isotopic Abundances and Atomic Weights
^ Magnetic susceptibility of the elements and inorganic compounds, in Handbook of Chemistry and Physics (PDF). CRC press. 2000. ISBN 0849304814.
^ "Sulfur History". Georgiagulfsulfur.com. Retrieved 2008-09-12.
^ ^a ^b ^c Greenwood, N. N.; & Earnshaw, A. (1997). Chemistry of the Elements (2nd ed.), Oxford:Butterworth-Heinemann. ISBN 0-7506-3365-4.
^ ^a ^b "sulphur". Oxford English Dictionary (3rd ed.). Oxford University Press. September 2005. (Subscription or UK public library membership required.)
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^ "Ask Oxford". Retrieved 12 November 2010.
^ Mallory & Adams (2006) The Oxford introduction to Proto-Indo-European and the Proto-Indo-European world, Oxford University Press
^ Rettig, S. J.; Trotter, J. (15 December 1987). "Refinement of the structure of orthorhombic sulfur, α-S8". Acta Crystallographica Section C Crystal Structure Communications 43 (12): 2260–2262. doi:10.1107/S0108270187088152.
^ A strong odor called "smell of sulfur" actually is given off by several sulfur compounds, such as hydrogen sulfide and organosulfur compounds.
^ ^a ^b ^c Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. pp. 645–665. ISBN 0-08-037941-9.
^ Steudel, Ralf; Eckert, Bodo (2003). "Solid Sulfur Allotropes Sulfur Allotropes". Topics in Current Chemistry. Topics in Current Chemistry 230: 1–80. doi:10.1007/b12110. ISBN 978-3-540-40191-9.
^ Steudel, R. (1982). "Homocyclic Sulfur Molecules". Topics in Current Chemistry. Topics in Current Chemistry 102: 149–176. doi:10.1007/3-540-11345-2_10. ISBN 978-3-540-11345-4.
^ Tebbe, Fred N.; Wasserman, E.; Peet, William G.; Vatvars, Arturs; Hayman, Alan C. (1982). "Composition of Elemental Sulfur in Solution: Equilibrium of S
6, S₇, and S₈ at Ambient Temperatures". Journal of the American Chemical Society 104 (18): 4971–4972. doi:10.1021/ja00382a050.
^ Meyer, Beat (1964). "Solid Allotropes of Sulfur". Chemical Reviews 64 (4): 429–451. doi:10.1021/cr60230a004.
^ Meyer, Beat (1976). "Elemental sulfur". Chemical Reviews 76 (3): 367–388. doi:10.1021/cr60301a003.
^ Cameron, A. G. W. (1957). "Stellar Evolution, Nuclear Astrophysics, and Nucleogenesis" (PDF). CRL-41.
^ Mason, B. (1962). Meteorites. New York: John Wiley & Sons. p. 160. ISBN 0-908678-84-3.
^ Lopes, Rosaly M. C.; Williams, David A. (2005). "Io after Galileo". Reports on Progress in Physics 68 (2): 303–340. Bibcode:2005RPPh...68..303L. doi:10.1088/0034-4885/68/2/R02.
^ Rickwood, P. C. (1981). "The largest crystals" (PDF). American Mineralogist 66: 885–907.
^ Kutney, Gerald (2007). Sulfur: history, technology, applications & industry. Toronto: ChemTec Publications. p. 43. ISBN 978-1-895198-37-9. OCLC 79256100.
^ Klein, Cornelis and Cornelius S. Hurlbut, Jr., Manual of Mineralogy, Wiley, 1985, 20th ed., p. 265-6 ISBN 0-471-80580-7
^ Sulphur on Mindat.org
^ ^a ^b ^c ^d ^e Nehb, Wolfgang; Vydra, Karel (2006). "Sulfur". Ullmann's Encyclopedia of Industrial Chemistry. Wiley-VCH Verlag. doi:10.1002/14356007.a25_507.pub2. ISBN 3-527-30673-0.
^ Riegel, Emil; Kent, James (2007). Kent and Riegel's handbook of industrial chemistry and biotechnology 1. New York: Springer. p. 1171. ISBN 978-0-387-27842-1. OCLC 74650396.
^ Washington, Booker T. (1912). The Man Farthest Down: A Record of Observation and Study in Europe. Doubleday, Page. p. 214.
^ ^a ^b ^c Eow, John S. (2002). "Recovery of sulfur from sour acid gas: A review of the technology". Environmental Progress 21 (3): 143–162. doi:10.1002/ep.670210312.
^ ^a ^b ^c Schreiner, Bernhard (2008). "Der Claus-Prozess. Reich an Jahren und bedeutender denn je". Chemie in unserer Zeit 42 (6): 378–392. doi:10.1002/ciuz.200800461.
^ Hyndman, A. W.; Liu, J. K.; Denney, D. W. (1982). "Sulfur: New Sources and Uses". ACS Symposium Series 183: 69–82. doi:10.1021/bk-1982-0183.ch005. ISBN 0-8412-0713-5. |chapter= ignored (help)
^ Mohamed, Abdel-Mohsen; El-Gamal, Maisa (2010). Sulfur concrete for the construction industry: a sustainable development approach. Fort Lauderdale: J. Ross Publishing. p. 109. ISBN 978-1-60427-005-1. OCLC 531718953.
^ McElvaney, Kevin. "The Men Who Mine Volcanos". theatlantic.com. The Atlantic. Retrieved 26 February 2015.
^ Apodaca, Lori E. (2012) Sulfur. Mineral Commodity Summaries. USGS
^ ^a ^b Apodaca, Lori E. "Mineral Yearbook 2010: Sulfur" (PDF). United States Geological Survey.
^ Shriver, Atkins. Inorganic Chemistry, Fifth Edition. W. H. Freeman and Company, New York, 2010; pp 416
^ Handbook of Preparative Inorganic Chemistry, 2nd ed. Edited by G. Brauer, Academic Press, 1963, NY. Vol. 1. p. 421.
^ Fujimori, Toshihiko; Morelos-Gómez, Aarón; Zhu, Zhen; Muramatsu, Hiroyuki; Futamura, Ryusuke; Urita, Koki; Terrones, Mauricio; Hayashi, Takuya; Endo, Morinobu; Young Hong, Sang; Chul Choi, Young; Tománek, David; Kaneko, Katsumi (2013). "Conducting linear chains of sulphur inside carbon nanotubes". Nature Communications 4. doi:10.1038/ncomms3162.
^ Hasek, W. R. (1961). "1,1,1-Trifluoroheptane". Org. Synth. 41: 104. doi:10.1002/0471264180.os041.28.
^ Rutenberg, M. W; Horning, E. C. (1950). "1-Methyl-3-ethyloxindole". Org. Synth. 30: 62. doi:10.15227/orgsyn.030.0062.
^ Heal, H. G. (1980). The Inorganic Heterocyclic Chemistry of Sulfur, Nitrogen, and Phosphorus. London: Academic Press. ISBN 0-12-335680-6.
^ Chivers, T. (2004). A Guide To Chalcogen-Nitrogen Chemistry. Singapore: World Scientific. ISBN 981-256-095-5.
^ Vaughan, D. J.; Craig, J. R. "Mineral Chemistry of Metal Sulfides" Cambridge University Press, Cambridge (1978) ISBN 0-521-21489-0
^ ^a ^b Cremlyn R. J.; "An Introduction to Organosulfur Chemistry" John Wiley and Sons: Chichester (1996). ISBN 0-471-95512-4.
^ Wilson, R. W.; Penzias, A. A.; Wannier, P. G.; Linke, R. A. (15 March 1976). "Isotopic abundances in interstellar carbon monosulfide". Astrophysical Journal 204: L135–L137. Bibcode:1976ApJ...204L.135W. doi:10.1086/182072.
^ Banoub, Joseph (2011). Detection of Biological Agents for the Prevention of Bioterrorism. Dordrecht: Springer. p. 183. ISBN 978-90-481-9815-3. OCLC 697506461.
^ Rapp, George Robert (4 February 2009). Archaeomineralogy. p. 242. ISBN 978-3-540-78593-4.
^ Odyssey, book 22, lines 480–495. www.perseus.tufts.edu. Retrieved on 16 August 2012.
^ Pliny the Elder on science and technology, John F. Healy, Oxford University Press, 1999, ISBN 0-19-814687-6, pp. 247–249.
^ ^a ^b ^c Zhang, Yunming (1986). "The History of Science Society: Ancient Chinese Sulfur Manufacturing Processes". Isis 77 (3): 487. doi:10.1086/354207.
^ White, David Gordon (1996). The Alchemical Body — Siddha Traditions in Medieval India. Chicago: University of Chicago Press. pp. passim. ISBN 978-0-226-89499-7.
^ Lin, A. N.; Reimer, R. J.; Carter, D. M. (1988). "Sulfur revisited". Journal of the American Academy of Dermatology 18 (3): 553–558. doi:10.1016/S0190-9622(88)70079-1. PMID 2450900.
^ Maibach, HI; Surber, C.; Orkin, M. (1990). "Sulfur revisited". Journal of the American Academy of Dermatology 23 (1): 154–156. doi:10.1016/S0190-9622(08)81225-X. PMID 2365870.
^ Gupta, A. K.; Nicol, K. (2004). "The use of sulfur in dermatology". Journal of drugs in dermatology : JDD 3 (4): 427–31. PMID 15303787.
^ Riall, Lucy (1998). Sicily and the Unification of Italy: Liberal Policy and Local Power, 1859–1866. Oxford University Press. ISBN 9780191542619. Retrieved 7 February 2013.
^ Thomson, D. W. (April 1995). "Prelude to the Sulphur War of 1840: The Neapolitan Perspective". European History Quarterly 25 (2): 163–180. doi:10.1177/026569149502500201.
^ Botsch, Walter (2001). "Chemiker, Techniker, Unternehmer: Zum 150. Geburtstag von Hermann Frasch". Chemie in unserer Zeit (in German) 35 (5): 324–331. doi:10.1002/1521-3781(200110)35:5<324::AID-CIUZ324>3.0.CO;2-9.
^ Kogel, Jessica (2006). Industrial minerals & rocks: commodities, markets, and uses (7th ed.). Colorado: Littleton. p. 935. ISBN 978-0-87335-233-8. OCLC 62805047.
^ Sulfuric Acid Growth. Pafko.com. Retrieved on 16 August 2012.
^ Sulfur as a fertilizer. Sulphurinstitute.org. Retrieved on 16 August 2012.
^ Zhao, F.; Hawkesford, MJ; McGrath, SP (1999). "Sulphur Assimilation and Effects on Yield and Quality of Wheat". Journal of Cereal Science 30 (1): 1–17. doi:10.1006/jcrs.1998.0241.
^ Blake-Kalff, M. M. A. (2000). "Diagnosing sulfur deficiency in field-grown oilseed rape (Brassica napus L.) and wheat ( Triticum aestivum L.)". Plant and Soil 225 (1/2): 95–107. doi:10.1023/A:1026503812267.
^ Ceccotti, S. P. (1996). "Plant nutrient sulphur-a review of nutrient balance, environmental impact and fertilizers". Fertilizer Research 43 (1–3): 117–125. doi:10.1007/BF00747690.
^ Mohamed, Abdel-Mohsen Onsy; El Gamal, M. M (13 July 2010). Sulfur Concrete for the Construction Industry: A Sustainable Development Approach. pp. 104–105. ISBN 978-1-60427-005-1.
^ Every, Richard L.; et al. (20 August 1968). "Method for Preparation of Wettable Sulfur" (PDF). Retrieved 20 May 2010.
^ Spencer, Benjamin Sulfur in wine demystified. intowine.com. Retrieved 26 October 2011.
^ ^a ^b Hagers Handbuch der Pharmazeutischen Praxis (in German) 6B (4th ed.). Berlin–Heidelberg–New York: Springer. 1978. pp. 672–9. ISBN 3-540-07738-3.
^ Arzneibuch-Kommentar. Wissenschaftliche Erläuterungen zum Europäischen Arzneibuch und zum Deutschen Arzneibuch [Pharmacopoeia Commentary. Scientific annotations to the European Pharmacopoeia and the German Pharmacopoeia] (in German) (23rd ed.). Stuttgart: Wissenschaftliche Verlagsgesellschaft. 2004. Monographie Schwefel zum äußerlichen Gebrauch [Monograph Sulfur for external use]. ISBN 978-3-8047-2575-1.
^ Multum Consumer Information: Sulfur topical.
^ Cite error: The named reference Drugs.com was invoked but never defined (see the help page).
^ ^a ^b Nelson, D. L.; Cox, M. M. (2000). Lehninger, Principles of Biochemistry (3rd ed.). New York: Worth Publishing. ISBN 1-57259-153-6.
^ Thauer, R. K. (1998). "Biochemistry of methanogenesis: a tribute to Marjory Stephenson:1998 Marjory Stephenson Prize Lecture". Microbiology 144 (9): 2377–2406. doi:10.1099/00221287-144-9-2377. PMID 9782487.
^ Lippard, S. J.; Berg, J. M. (1994). Principles of Bioinorganic Chemistry. University Science Books. ISBN 0-935702-73-3.
^ Pronk JT; Meulenberg R; Hazeu W; Bos P; Kuenen JG (1990). "Oxidation of reduced inorganic sulphur compounds by acidophilic thiobacilli" (PDF). FEMS Microbiology Letters 75 (2–3): 293–306. doi:10.1111/j.1574-6968.1990.tb04103.x.
^ Heldt, Hans-Walter (1996). Pflanzenbiochemie. Heidelberg: Spektrum Akademischer Verlag. pp. 321–333. ISBN 3-8274-0103-8.
^ Gresham et al, Medscape: Hydrogen sulfide toxicity

External links

Sulfur at The Periodic Table of Videos (University of Nottingham)
Atomic Data for Sulfur, NIST Physical Measurement Laboratory
Sulfur phase diagram, Introduction to Chemistry For Ages 13–17
Crystalline, liquid and polymerization of sulfur on Vulcano Island, Italy
Sulfur and its use as a pesticide
The Sulphur Institute
Nutrient Stewardship and The Sulphur Institute

UpToDate Contents

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English Journal

Reductive off-odors in wines: Formation and release of H2S and methanethiol during the accelerated anoxic storage of wines.

Franco-Luesma E1, Ferreira V2.
Food chemistry.Food Chem.2016 May 15;199:42-50. doi: 10.1016/j.foodchem.2015.11.111. Epub 2015 Nov 27.
In order to better understand the processes involved in the development of H2S and methanethiol (MeSH) along anoxic storage of wines, 24 wines were stored in strict anoxia at 50°C for 3weeks. Free and total forms of H2S and MeSH were measured at different times. Results showed that: (1) all wines c
PMID 26775942

Iron-sulfur-based single molecular wires for enhancing charge transport in enzyme-based bioelectronic systems.

Mahadevan A1, Fernando T1, Fernando S2.
Biosensors & bioelectronics.Biosens Bioelectron.2016 Apr 15;78:477-82. doi: 10.1016/j.bios.2015.11.086. Epub 2015 Nov 30.
When redox enzymes are wired to electrodes outside a living cell (ex vivo), their ability to produce a sufficiently powerful electrical current diminishes significantly due to the thermodynamic and kinetic limitations associated with the wiring systems. Therefore, we are yet to harness the full pote
PMID 26657591

Grapevine-shoot stilbene extract as a preservative in red wine.

Raposo R1, Ruiz-Moreno MJ2, Garde-Cerdán T3, Puertas B1, Moreno-Rojas JM2, Gonzalo-Diago A3, Guerrero R1, Ortíz V2, Cantos-Villar E4.
Food chemistry.Food Chem.2016 Apr 15;197 Pt B:1102-11. doi: 10.1016/j.foodchem.2015.11.102. Epub 2015 Nov 26.
This paper reports the use of a grapevine-shoot stilbene extract (Vineatrol®) as a preservative in red wine. Its effectiveness to preserve red wine quality under two different winemaking systems (traditional and Ganimede) was studied at bottling and after twelve months of storage in bottle. Enologi
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Effects of different drying methods on the product quality and volatile compounds of whole shiitake mushrooms.

Tian Y1, Zhao Y2, Huang J2, Zeng H2, Zheng B3.
Food chemistry.Food Chem.2016 Apr 15;197(Pt A):714-22. doi: 10.1016/j.foodchem.2015.11.029. Epub 2015 Nov 10.
Various drying methods play important roles in the preservation of foods. However, how the different drying methods affect the quality of some foods is not clear. This paper evaluates the effects of hot air, vacuum, microwave, and microwave vacuum drying techniques on important qualities and volatil
PMID 26617008

Japanese Journal

平成24年度硫黄および硫酸需給見通し(見直し)

硫酸協会調査部
硫酸と工業 66(1), 1-8, 2013-01
NAID 40019563692

シンポジウム低硫黄燃料システムと機関艤装(その2)低硫黄燃料対応システムの提案 (特集みなおそう!環境にやさしい船造り(Part 2))

伊藤誠,黒川浩志,濱中伸太
Marine engineering : journal of the Japan Institute of Marine Engineering = マリンエンジニアリング :日本マリンエンジニアリング学会誌 48(1), 29-33, 2013-01
NAID 40019554940

シンポジウム低硫黄燃料システムと機関艤装(その1)現状の規制及び状況について (特集みなおそう!環境にやさしい船造り(Part 2))

伊藤誠,黒川浩志,濱中伸太
Marine engineering : journal of the Japan Institute of Marine Engineering = マリンエンジニアリング :日本マリンエンジニアリング学会誌 48(1), 24-28, 2013-01
NAID 40019554932

Reconstruction of Physicochemical Environment of Hydrothermal Mineralization at Malanjkhand Copper Deposit, Central India : Constraints from Sulfur Isotope Ratios in Pyrite, Molybdenite and Chalcopyrite

PANIGRAHI Mruganka K.,PANDIT Dinseh,NAIK Rajesh K.
Resource geology : an official journal of the Society of Resource Geology 63(1), 110-116, 2013
NAID 40019554255

Related Pictures

★リンクテーブル★

リンク元	「thionic」「sulphur」「硫黄」
拡張検索	「sulfur-reducing bacteria」「dilute sulfuric acid」「desulfuration」「technetium-99m-sulfur colloid」「sulfur isotope」

「thionic」

Octasulfur
Systematic (IUPAC) name
	Octathiocane
Clinical data
AHFS/Drugs.com	Multum Consumer Information
Routes of; administration	Topical, rarely oral
Legal status
Legal status	US: OTC; OTC;
Identifiers
CAS Number	10544-50-0
ATC code	D10AB02 (WHO)
PubChem	CID 66348
ChemSpider	59726
ChEBI	CHEBI:29385
Chemical data
Formula	S8
Molar mass	256.52 g/mol
	SMILES S1SSSSSSS1 ;
	InChI InChI=1S/S8/c1-2-4-6-8-7-5-3-1 ; Key:JLQNHALFVCURHW-UHFFFAOYSA-N ;

Sulfur, ₁₆S
	Spectral lines of sulfur
General properties
Name, symbol	sulfur, S
Appearance	lemon yellow sintered microcrystals
Pronunciation	/ˈsʌlfər/; SUL-fər
Alternative name	sulphur
Sulfur in the periodic table
	phosphorus ← sulfur → chlorine
Atomic number (Z)	16
Group, block	group 16 (chalcogens), p-block
Period	period 3
Element category	polyatomic nonmetal
Standard atomic weight (Ar)	32.06 (32.059–32.076)
Electron configuration	[Ne] 3s2 3p4
per shell	2, 8, 6
Physical properties
Phase	solid
Melting point	388.36 K (115.21 °C, 239.38 °F)
Boiling point	717.8 K (444.6 °C, 832.3 °F)
Density near r.t.	alpha: 2.07 g/cm3; beta: 1.96 g/cm3; gamma: 1.92 g/cm3
when liquid, at m.p.	1.819 g/cm3
Critical point	1314 K, 20.7 MPa
Heat of fusion	mono: 1.727 kJ/mol
Heat of vaporization	mono: 45 kJ/mol
Molar heat capacity	22.75 J/(mol·K)
Atomic properties
Oxidation states	6, 5, 4, 3, 2, 1, −1, −2 (a strongly acidic oxide)
Electronegativity	Pauling scale: 2.58
Ionization energies	1st: 999.6 kJ/mol; 2nd: 2252 kJ/mol; 3rd: 3357 kJ/mol; (more)
Covalent radius	105±3 pm
Van der Waals radius	180 pm
Miscellanea
Crystal structure	orthorhombic
Thermal conductivity	0.205 W/(m·K) (amorphous)
Electrical resistivity	2×1015 Ω·m (at 20 °C) (amorphous)
Magnetic ordering	diamagnetic
Bulk modulus	7.7 GPa
Mohs hardness	2.0
CAS Number	7704-34-9
History
Discovery	Chinese (before 2000 BCE)
Recognized as an element by	Antoine Lavoisier (1777)
Most stable isotopes of sulfur
	view; talk; edit; · references

外見
	黄色; ; ; 硫黄のスペクトル線
一般特性
名称, 記号, 番号	硫黄, S, 16
分類	非金属
族, 周期, ブロック	16, 3, p
原子量	32.065(5)
電子配置	[Ne] 3s2 3p4
電子殻	2, 8, 6（画像）
物理特性
相	固体
密度（室温付近）	(α) 2.07 g/cm3
密度（室温付近）	(β) 1.96 g/cm3
密度（室温付近）	(γ) 1.92 g/cm3
融点での液体密度	1.819 g/cm3
融点	388.36 K, 115.21 °C, 239.38 °F
沸点	717.8 K, 444.6 °C, 832.3 °F
臨界点	1314 K, 20.7 MPa
融解熱	(mono) 1.727 kJ/mol
蒸発熱	(mono) 45 kJ/mol
熱容量	(25 °C) 22.75 J/(mol·K)
蒸気圧
原子特性
酸化数	6, 5, 4, 3, 2, 1, -1, -2; (強酸性酸化物)
電気陰性度	2.58（ポーリングの値）
イオン化エネルギー; （詳細）	第1： 999.6 kJ/mol
	第2： 2252 kJ/mol
	第3： 3357 kJ/mol
共有結合半径	105±3 pm
ファンデルワールス半径	180 pm
その他
結晶構造	斜方晶系
磁性	反磁性
電気抵抗率	(20 °C) （無定形）2×1015Ω·m
熱伝導率	(300 K) （無定形）0.205 W/(m·K)
体積弾性率	7.7 GPa
モース硬度	2.0
CAS登録番号	7704-34-9
最安定同位体
	詳細は硫黄の同位体を参照
	表示

　　[★]

adj.

イオウの、硫黄の

関: S、sulfur、sulphur

「sulphur」

　　[★]

n.

イオウ、硫黄

関: S、sulfur、thionic

「硫黄」

　　[★]

英: sulfur、sulphur、thionic
関: イオウ

「sulfur-reducing bacteria」

　　[★]

硫酸還元菌、硫酸還元細菌

関: Desulfovibrio

「dilute sulfuric acid」

　　[★]

希硫酸

関: dil. H2SO4、diluted sulfuric acid

「desulfuration」

　　[★]

脱硫化

関: desulfurization、desulfurize

「technetium-99m-sulfur colloid」

　　[★]

テクネチウム99mイオウコロイド

「sulfur isotope」

　　[★]

イオウ同位体、硫黄同位体

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[10] 硫黄の主な生産国は、アメリカ、カナダ、ポーランド、フランス、ロシア、メキシコ、日本である

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