WordNet

a nonmetallic bivalent element that is normally a colorless odorless tasteless nonflammable diatomic gas; constitutes 21 percent of the atmosphere by volume; the most abundant element in the earths crust (同)O, atomic number 8

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『酸素』(化学記号はO)

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出典(authority):フリー百科事典『ウィキペディア（Wikipedia）』「2016/07/25 09:26:51」(JST)

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酸素（さんそ、英: oxygen）は原子番号8、原子量16.00の非金属元素である。元素記号は O。周期表では第16族元素（カルコゲン）および第2周期元素に属し、電気陰性度が大きいため反応性に富み、他の殆どの元素と化合物（特に酸化物）を作る。標準状態では2個の酸素原子が二重結合した無味無臭無色透明の二原子分子である酸素分子 O₂ として存在する。宇宙では水素、ヘリウムに次いで3番目に多くの質量を占め^[4]、ケイ素量を10⁶としたときの比率は 2.38 × 10⁷ である^[5]。地球地殻の元素では質量が最も多く^[6]47%が酸素である^[1]。気体の酸素分子は大気の体積の20.95%^[7]、質量で23%を占める^[1]。

1 名称
2 性質
- 2.1 物理的性質
- 2.2 化学的性質
3 酸素分子
- 3.1 物理的性質
- 3.2 構造
- 3.3 その他の特徴
4 生物学的役割
- 4.1 光合成と呼吸
- 4.2 大気成分中の酸素形成
5 歴史
- 5.1 初期の実験
- 5.2 フロギストン説
- 5.3 発見
- 5.4 ラヴォアジエの功罪
- 5.5 量産・工業化
6 製造
7 用途
8 化合物
9 同素体
10 同位体
11 安全と注意
- 11.1 酸素中毒
- 11.2 燃焼と他の危険
12 大気中酸素濃度の減少
13 脚注
- 13.1 注釈
- 13.2 出典
14 参考文献
15 関連項目

名称

スウェーデンの薬剤師、カール・ヴィルヘルム・シェーレが1771年に初めて見つけた^[1]。しかし、これはすぐに公にされず、その後1774年にジョゼフ・プリーストリーがそれとは独立して見付けた後に広く知られるようになった^[8]。そのため、化学史上の発見者はプリーストリーとされている^[9]。

酸素は発見当初、「酸を生む物」と誤解され、ギリシャ語の oxys（酸）と genen（生む）を合わせた名称で呼ばれていた。これは、アントワーヌ・ラヴォアジエが、酸素が「酸を生む物」であると誤解して、oxygène（仏語）と名付けた^[1]ことに由来する。英語でも「oxygen（オキシジン）」といい、独語でも「Sauerstoff（ザウアーシュトッフ）」といい、日本語でもこれらを宇田川榕菴が直訳して「酸素」と呼んだ。

一方、中国語圏では「酸」という字を用いず、「氧」（中国語読み：ヤン。日本語読み：よう）という字を充て、氧や氧氣（ようき）という。

性質

物理的性質

同位体については、3種類の安定同位体と10種の放射性同位体（いずれも半減期3分未満）が知られている。

酸素は、地球の地殻（質量比で約46.7%）およびマントルに最も多く含まれている元素であり、多くは岩石中に酸化物・ケイ酸塩・炭酸塩などの形で存在する。

地球外でも酸素は多く存在している。主な存在形態である氷は地球の他、惑星や、彗星、小惑星などにも見られる。火星の極には二酸化炭素が固体のドライアイスとして存在している。星が生まれる元となる分子雲では一酸化炭素が分子の中で2番目に存在量の多い分子である。酸素の起源は恒星核におけるヘリウムの核融合であり、酸素のスペクトルが検出される恒星も存在している。

約 90 K で液体、約 54 K で青みがかった固体となる。ダイアモンドアンビルなどで100万気圧を超えた高圧下では金属光沢を持ち、125万気圧、0.6 K では超伝導金属となる。

化学的性質

酸素は、フッ素に次いで2番目に電気陰性度が大きい^[10]ため酸化力が強く、ほとんどの元素と発熱反応を起こして化合物をつくる^[11]。1962年以降には希ガスであるキセノンも、酸素と化合して三酸化キセノン (XeO₃) などの化合物を作ることがわかった^[12]。

酸素分子

物理的性質

酸素分子（英: dioxygen、化学式：O₂）は、常温常圧では無色無臭で助燃性をもつ気体として存在する。分子量 32.00、沸点 −183 ℃ (90 K)、融点 −218.9 ℃ (54.3 K)。水100グラムに溶解する量は0℃で6.945ミリグラム、25℃で3.931ミリグラム、50℃で2.657ミリグラム^[3]。液体酸素は淡青色を示し、比重は1.14である^[3]。基底状態の三重項状態では不対電子を持つため常磁性体である。また活性酸素の一種で反磁性である励起状態の一重項酸素も存在する。

構造

標準状態において一般の^[13]酸素は、2つの酸素原子が縮退した三重項の電子配置で化学結合した分子構造（三重項酸素分子）を持つ無色無臭の気体である。この結合次数は2であり、一般に二重結合^[14]、または1個の2電子結合と2個の3電子結合と表記される^[15]。三重項酸素分子とは電子の全スピン量子数が1となる状態で、具体的には2つの不対電子が酸素分子に2つあるπ^*反結合性軌道^[16]をひとつずつ占め、しかも同じ向きのスピンを取っている^[17]。このとき、酸素分子のエネルギーは基底状態にある^[18]。また、酸素分子の二重結合は反結合軌道にも電子が存在するため、結合軌道のみで電子を充足させる三重結合の窒素よりも安定さは下がり、また、2つの電子が対を作らずビラジカルとして存在するため、結果として酸素分子は窒素分子よりも少ないエネルギーで他の物質と反応しやすくなる^[18]^[19]。

通常の三重項酸素分子は常磁性を持つ。これは、不対電子のスピン磁気モーメント（スピンの向きが同じ電子がπ*反結合性軌道に入る^[20]）とふたつの酸素分子間に働く交換相互作用による^[21]。液体酸素は磁石に吸い付けられ、実験では磁極間で自重を支えるに充分強い橋をつくる程である^[22]^[23]。

これに対し、外部から高エネルギーが加わり不対電子の一つがスピンを逆方向へ変え^[24]、全スピン量子数が0となった酸素を一重項酸素といい、有機化合物との反応性が高い。自然界で一重項酸素は、光合成の過程で水から作られたり^[25]、対流圏で短波長の光によってオゾンの分解から発生したり^[26]、または免疫システムの中で活性酸素の原料として用いられたりする^[27]。

その他の特徴

熱力学的に反応性が高く不安定な分子ではあるが、地球上では初期には光合成を行なう嫌気性菌により、後の時代には植物の光合成によって年間約10¹¹トン^[9]供給され続けているため多量に存在する。酸素呼吸を行なう生物によって消費される。実際、生命が発生する以前の原始大気では酸素分子はほとんど存在せず、二酸化炭素など他の原子と結合した状態であった。現在の大気中の酸素分子はそのほぼ全てが光合成由来だと考えられている^[28]^{[注 1]}。逆に、他の天体の大気中に遊離酸素の存在が確認されれば、生命の存在する間接的証拠となると考えられている。

酸素は、呼吸をする生物によっては必須であるが、同時に有害でもある^[30]。呼吸の過程や光反応などで生じる活性酸素は、DNAなどの生体構成分子を酸化して変性させる^[31]。純酸素の長時間吸引は生体にとって有害である。未熟児網膜症の原因になったり、60%以上の高濃度酸素を12時間以上吸引すると、肺の充血がみられたりし、最悪の場合、失明や死亡する危険性がある。

25℃で標準気圧下では、淡水は1リットル中に酸素を6.04ミリリットル含んでいるが、海水では1リットルあたり4.95ミリリットルしか含んでいない^[32]。5℃での溶解度は、淡水では 9.0 mL L⁻¹、海水では 7.2 mL L⁻¹ まで増加している。

液体酸素は液体空気を分留して得られ、強い酸化剤である^[3]。液体空気を放置すると、沸点の低い窒素が先に蒸発するため、酸素分子が濃縮される^[3]。1リットルの液化酸素が気化すると約800リットルの酸素ガスになる。

酸素は紫外線や無声放電などによってオゾン (O₃) へと変換される。また、酸素分子のイオンとしてスーパーオキシドアニオン O₂^- とジオキシゲニル O₂⁺ が知られている。

生物学的役割

光合成と呼吸

光合成は水を分解し、酸素を放出して水素を二酸化炭素と反応させて糖類を得る過程である

自然界において遊離酸素は、光合成によって水が光分解されることで生じ、海洋中の緑藻類やシアノバクテリアが地球大気中の酸素70%を、残りは陸上の植物が作り出している^[33]。

簡易な光合成の反応式は以下の通りである^[34]。

6 CO₂ + 6H₂O + 光子 → C₆H₁₂O₆ + 6 O₂ （二酸化炭素 + 水 + 日光 → グルコース + 酸素）

光分解による酸素発生は葉緑体のチラコイド膜中で起こる。光をエネルギーとするこの作用は多くの段階を経て、ATP を光リン酸化 (photophosphorylation) させるプロトンの濃度勾配を起こす^[35]。この際、水を酸化することで酸素ガスが発生し、大気中に放出される^[36]。

酸素ガスは好気性生物が呼吸を行い、ミトコンドリアで酸化的リン酸化反応を経てATPを発生させるために使われる。酸素呼吸の反応は本質的に光合成の逆である。

C₆H₁₂O₆ + 6 O₂ → 6 CO₂ + 6 H₂O + 2880 kJ mol^-1

脊椎動物では酸素ガスは肺の膜を通して血液中に拡散し赤血球中のヘモグロビンと結びつき、その色を紫がかった赤から明るい赤へ変える^[37]^[38]。他の動物ではヘモシアニン（軟体動物や節足動物の一種など）やヘムエリスリン（クモやロブスターなど）が使われる例もある^[39]。1リットルの血液が溶かせる酸素ガスは 200 cm³ である^[39]。

超酸化物イオンや過酸化水素などの活性酸素は酸素呼吸を行う生体にとって非常に危険な副産物であり^[37]^[39]、ミトコンドリアを取り込んだ真核生物は進化の過程で DNA を酸素から保護するために核膜を獲得した^[31]。その一方で高等生物は免疫系で細菌を破壊するために過酸化物を用いている^[37]^[40]。また、植物が病原体に抵抗して起こす過敏感反応 (hypersensitive response) でも活性酸素は重要な役割を果たす^[41]^[42]。

成人が消費する酸素は1分あたり約250 mLであり^[43]、これは約0.36 gに相当する。ここから計算すると、人類全体が1年間に消費する量は13億トンに相当する^{[注 2]}。

なお、酸素を利用しない呼吸の形態を嫌気呼吸という。最初の地球に酸素が存在しなかったことから、これが最初の呼吸のあり方と考えられる。これは好気呼吸の経路にも、解糖系という形態で残っている。現在も、酸素を全く使わずに生活する微生物も存在し、そのような微生物は、酸素の存在下では死滅する（嫌気性生物）。おそらく、初期の微生物にとっても、酸素は有毒物質であったと考えられる。

大気成分中の酸素形成

地球大気における酸素含有量の変遷。1. 酸素が作られない期間 2. 酸素生成が始まるが海水や海底岩石に吸収される 3. 海洋から酸素ガスが放出されるが地表への吸収やオゾン層形成のため消費される期間 4, 5. 酸素吸収が飽和し大気中に溜まる

地球誕生当初、大気中には遊離酸素ガスはほとんど存在しなかったが、やがて古細菌やバクテリアが生じ、少なくとも30億年前には酸素が作られ始めた^[28]。当初は海水中の溶解鉄と化合し縞状鉄鉱床を形成したが、酸素ガスが海洋から溢れ始めたのは大規模な大陸変動によって浅瀬が作られた27億年前頃からであった^[30]。17億年前には大気中の酸素含有比率は10%に達し^[44]、二酸化炭素と酸素の比率が逆転したのは7 - 8億年前と考えられる^[28]。

24億年前の酸素の大量発生（英語版）が起こった期間、他の元素と結合していない多くの遊離酸素が海中や大気中に溢れ、当時の嫌気性生物の大量絶滅を引き起こしたと考えられる。しかしながら、酸素を用いる細胞呼吸を獲得した好気性生物はより多くのATPを作り出せるようになり、地球に生物圏を形成した^[45]。この光合成と酸素呼吸は真核生物への進化をもたらし、これが植物や動物などの複雑な多細胞生物が生まれるに至る第一歩となった。

5億4千万年前のカンブリア紀が始まったころからは、大気中の酸素比率は15%–30%の間で推移した^[28]。それは石炭紀の終わりに当たる3億年前頃には最大35 %まで達し^[28]、昆虫や両生類の大型化に作用した可能性がある^[28]。人類は年間70億トンの化石燃料を使用するにあたり酸素を消費し続けているが、これによって大気中の酸素比率に与える影響は微々たるものである^[21]。

歴史

初期の実験

フィロンの実験は後の研究者たちに影響を与えた

燃焼と空気の間には何らかの関係があるのでは、と行われた最も古い実験のひとつは、紀元前2世紀の古代ギリシアのビザンチウムのフィロンが著した『プネウマティカ (Pneumatica)』に記録されている。器に据えた蝋燭を灯してガラスの壷を上から被せ、壷の口が漬かるまで器に水を満たす。すると、壷の中へ水が吸い上がる様子を観察できた^[46]。フィロンは、壷の中の空気が「四大元素の火」に変換され、これが壷のガラス壁を透過して逃げたと考えた。それから遥か時代が下った中世のルネサンス期に、レオナルド・ダ・ヴィンチはフィロンの実験に考察を加え、燃焼や呼吸を通じて空気が一部消費されると考えた^[47]。

17世紀後半にロバート・ボイルは、燃焼には空気が必要不可欠であることを立証した。これをジョン・メーヨーは、必要なものは彼が「硝気精^[48] (spiritus nitroaereus、nitroaereus)」と名づけた空気の構成要素だという説を提唱した^[49]。メイヨーの実験はフィロンと同じように水で封じた逆さの容器にそれぞれ蝋燭とマウスを入れ、どちらも水位が1/14程度上昇したことを確認した^[50]。これから、メイヨーは燃焼と呼吸のいずれでも硝気精が消費されるとの確証を得た。またメイヨーは、アンチモンを加熱すると質量が増えることも確認し、これは金属に硝気精が結合したためと考えた^[49]。呼吸については、硝気精は肺の中で空気から取り出されて血液に受け渡され、動物の体温や筋肉の動きを生み出す反応に使われると考察し^[49]、1668年に発表した^[50]。

フロギストン説

ゲオルク・シュタール。彼はフロギストン説の構築と普及に寄与した

詳細は「フロギストン説」を参照

17世紀から18世紀にかけて、酸素はロバート・フック、オーレ・ボッシュ（英語版）、ミハイル・ロモノーソフ、ピエール・バイエンらが実験で作り出していたが、いずれもがそれを元素とは認識しなかった^[51]。そこには、フロギストン説と呼ばれる燃焼と腐食に関する広く知られた学説が影響を及ぼしていた。

1667年にドイツの錬金術師ヨハン・ベッヒャーが発案し1731年までにゲオルク・シュタールが理論構築したフロギストン説^[52]は、可燃物とは燃素（フロギストン）と他の物質の2つが結合した状態にあり、燃焼が起こると燃素が遊離し、残りの物質もしくは石灰が残るというものだった^[47]。この説では、木材や石炭などは燃素の含有率が高く、鉄など不燃性のものはほとんど含まないと考えられた。空気の効果は無視され、わずかに行われた実証試験でも可燃物を燃やすと軽くなるという点から確かに何かが失われているという考察がされたに過ぎず^[47]、発生ガスへ意識が向けられることは無かった。このフロギストン説が否定される契機は、金属を空気中で燃やすと重量が増すという報告だった。

発見

カール・ヴィルヘルム・シェーレ。惜しくも酸素発見者の栄誉を逃した

酸素は1771年^[2]、スウェーデンのカール・ヴィルヘルム・シェーレが酸化水銀(II)と様々な硝酸塩混合物を加熱する過程で発見した^[7]^[47]。シェーレはこの気体を「火素 (fire air)」と名づけ1775年に論文を作成したが、出版社の都合で^[2]発表されたのは1777年となった^[53]。

ジョゼフ・プリーストリー。一般的には彼が酸素の発見者とされる

シェーレが発見を知らしめるのに手間取っていた1774年8月1日、イギリスのジョゼフ・プリーストリーはガラス管に入れた酸化水銀(II)に日光を照射して得たガスに「脱フロギストン空気(dephlogisticated air)」と命名した^[7]。彼はこのガスの中では蝋燭がより明るく燃え、マウスが活発かつ長寿になることを確かめた。さらに自分でこのガスを吸い、「吸い込んだ時には普通の空気と大差ないと思ったが、少し後になると呼吸が軽く楽になった」と書き残した^[51]。1775年、プリーストリーは新聞紙上にこの発見を発表し、2冊目の著作 Experiments and Observations on Different Kinds of Air でも論述した^[47]。このように、彼の発表がシェーレよりも先に行われたため、酸素発見者はプリーストリーということになった。

フランスの高名な化学者アントワーヌ・ラヴォアジエは後に自分が新元素を発見していたと主張したが、1774年10月にラヴォアジエはプリーストリーの訪問を受け、ガス発生手段など実験の概要を耳にしている。また、それに先立つ9月30日にプリーストリーは前もって新発見したガスの説明を記した書簡をラヴォアジエに送っているが、ラヴォアジエはこれを受け取っていないと主張した。なおプリーストリーの死後、彼の私物の中から書簡の写しが見つかっている^[53]。

ラヴォアジエの功罪

アントワーヌ・ラヴォアジエ。旧来のフロギストン説を葬り去った

ラヴォアジエは、厳密な物質量確認を伴う酸化の実験を通じて、燃焼の実態を正しく説明することに貢献した^[7]。彼はフロギストン説を否定し、プリーストリーらが発見したガスが元素のひとつであると立証するため、1774年以来行われた実験の追試に乗り出した。

ラヴォアジエは、スズと空気を密閉した容器を加熱しても全体の重さに変化が無いことを観測し^[7]、開封すると外気が流れ込む事から空気の一部が減少していると確認し、またスズが重くなっていることも計測した。そして、この流入空気質量とスズの質量増分が同じであることを確認した。1777年、彼はこの実験結果などをまとめた書籍 Sur la combustion en général を発表した^[7]。この中でラヴォアジエは、空気は燃焼と呼吸に深く関わる vital air と、これらに関与しない azote（古希: ζωτον、「生気のない」の意）」の2種類のガスが混合したものと証明した。azote は後に窒素とされた^[7]。

1777年ラヴォアジエは「vital air」に、古代ギリシア語 ὀξύς（oxys、味覚の酸味を由来とする「鋭い」の意）と -γενής（-genēs、生み出す者を由来とする「製作者」の意）を合成したフランス語「oxygène」という命名を施した^[9]。これは、彼が酸素こそすべての酸性の源泉だという誤解を持っていたためこれらの単語が選択されたものだった^[54]。後に、酸性の根本となる元素は水素であることが判明したが、その頃には単語が既に定着していたため変更はできなかった。

イギリス科学界は同国人のプリーストリーが分離に成功したガスにこの名称を用いることに反対だったが、1791年に詩人でもあるエラズマス・ダーウィン（チャールズ・ダーウィンの祖父）が出版した有名な書籍『植物の園』 (The Botanic Garden) の中で、このガスを称賛する詩 oxygen を載せたため既に一般に広まっていたこともあり、「oxygen」の単語は英語に組み込まれてしまった^[53]。

量産・工業化

ジョン・ドルトンの原子論では、当初すべての元素は「単元素」であり、原子比も単純なものであるという仮定があり、水は水素と酸素が1対1のHO というみなしの元で酸素の原子量を8と判断していた^[55]。これは1805年にジョセフ・ルイ・ゲイ＝リュサックとアレクサンダー・フォン・フンボルトによって原子比が1対2に改められ、1811年にアメデオ・アヴォガドロがアボガドロの法則に則って水の正しい構成を解釈した^[56]。

19世紀には空気の構成も判明してきた。1877年にスイスのラウル・ピクテ（英語版）^[57]とフランスのルイ・ポール・カイユテ^[57]が相次いで酸素の液体化に成功したと発表し、安定状態での液体酸素はヤギェウォ大学のジグムント・ヴルブレフスキとカオル・オルシェフスキ（英語版）が初めて得た^[58]。

1891年にはイギリスのジェイムズ・デュワーが研究で用いるに充分な液体酸素の製法を見つけ^[21]、1895年にはドイツのカール・フォン・リンデとイギリスのウィリアム・ハンプソンがそれぞれ液化分留による商業ベースに乗る量産法を確立した^[59]。この酸素を工業的に用いる例として、1901年にはアセチレンと圧縮酸素を用いた溶接法のデモンスチレーションが行われた^[59]。

製造

実験室的には過酸化水素を触媒で分解することで得られる^[9]。触媒としては二酸化マンガンまたは、カタラーゼおよびそれらを含むレバーやジャガイモなどが利用できる。

そのほか、水の電気分解でも得られる。純粋な水は電気を通さないため少量の水酸化ナトリウムを加える。酸素は陽極で発生し、陰極では水素が発生する。

黒色に塗られた酸素ボンベ（MRI用）

工業的には空気の分留で得られる。空気を圧縮冷却し、沸点の差を利用して窒素やアルゴンなど他の成分と分けられる^[3]。酸素が圧縮充填されるボンベは内部圧力が14.7メガパスカルで、容器の色は黒と定められている^[3]（特に高純度品は表面積の半分を超えない範囲で水色も加えられる）。液体充填されている容器は断熱構造をしており圧力は1メガパスカル以下（およそ700キロパスカル）程度であり色は地金（ステンレスやアルミ合金の場合）か灰色に黒の帯を配したものである。ただし工業的にはほとんど液体酸素をタンクローリーで1回あたり9–10トンが輸送され、低温液化ガス貯槽（コールドエバポレーター）で受け入れされる^[3]。

用途

酸化剤: 化学工業などでは最も安価な酸化剤として多用される。
吸入用: 呼吸に不可欠な元素であるため、医療分野での酸素吸入に使われている^[60]。また傷病人に限らず、空気中の酸素濃度が低い場所での呼吸を助けるために、飛行機や青海チベット鉄道などの酸素放出装置や、高山に登る時などのボンベの中身にも使われている。他にテクニカルダイビングにおいて、減圧用ガスとして用いられる。
助燃剤: ガス溶接や鉄鋼の製造工程で助燃剤として使用されている^[60]。アセチレンを酸素とともに吹き出してえられる酸素アセチレン炎は 3000–4000 °C もの高温が得られ、鉄材の溶接や切断に利用されている。特に液体酸素はロケットエンジンの推進剤の酸化剤として用いられている。

酸素ガスの2004年度日本国内生産量は10,422,238,000立方メートル、工業消費量は4,093,787,000立方メートル、液化酸素の2004年度日本国内生産量は855,476,000立方メートル、工業消費量は68,215,000立方メートルである^[61]。

化合物

「Category:酸素の化合物」も参照

酸素は電気陰性度が高く、ほとんどあらゆる元素と化学結合する。多くの有機化合物は構成元素として酸素を含み、無機化合物の酸素化合物は酸化物として多方面で利用されている。

同素体

オゾンは主に大気中に含まれる希有な気体である

地球上での主な同素体は酸素分子 O₂ であり、その結合長は121 pm、結合エネルギーは498 kJ/molである^[62]。酸素分子は生物の複雑な細胞呼吸に使われている。

三酸素 (O₃) はオゾンとしてよく知られる非常に反応性の大きい単体の気体で、吸入すると肺組織を破壊する^[63]^[38]。オゾンは高層大気において、酸素分子が紫外線によって分裂した酸素原子と別の酸素分子が結合することによって生成している^[54]。オゾンは紫外領域を強く吸収するため、高層大気にあるオゾン層は地球を放射線から保護するシールドとして機能している^[38]^[64]。地表近くでもオゾンは生成しているが、これは自動車の排気ガスなどとして生成されている大気汚染物質である^[65]。

準安定状態分子である四酸素 (O₄) が2001年に発見されたが^[66]^[67]、これは固体酸素の6種の相のうちの1種として存在が仮定されていた。2006年にこの相が証明され、O₂を20 GPaに加圧することで合成されたが、実際には菱面体晶の O₈ クラスターであった^[68]。このクラスターは O₂ や O₃ よりも強力な酸化剤であるためロケットの推進剤としての用途が考えられている^[66]^[67]。1990年には固体酸素に96 GPa以上の圧力を与えると金属状態となる事が分かり^[69]、1998年にはこの相を超低温条件におくことにより超伝導となることが発見された^[70]。

同位体

詳細は「酸素の同位体」を参照

酸素には安定同位体として¹⁶O, ¹⁷O, ¹⁸Oの3種類が知られるが、天然存在比は¹⁶Oが99.7 %以上を占めている。また、放射性同位体も作られている。

かつては酸素を16として原子量を定義していたが、物理学では¹⁶Oの原子量を16としたのに対して、化学においては安定核種の平均原子量を16と置く定義の差があったことから、酸素の同位体の存在が判明して以降混乱が起こり、1961年に炭素12を基準とするように置き換えられた。

安全と注意

酸素中毒は通常よりも高い気圧の酸素を肺が吸い込んだ時に起こりやすい。深度へのスクーバダイビング（ディープダイビング）などで起こる可能性がある

酸素中毒

詳細は「酸素中毒」を参照

酸素ガスは高い分圧状態で痙攣症状などの酸素中毒を引き起こす場合がある^[71]^[72]。これは通常、大気の2.5倍の酸素分圧に相当する50キロパスカル以上であるときに起こる。そこで、標準気圧30キロパスカルの医療用酸素マスクは、酸素ガス比率を30%に定めている^[51]。かつて未熟児用保育器の中は高い比率の酸素を含んだガスが使われていたが、視神経に悪影響を与える可能性が指摘されてからは用いられなくなった^[51]^[73]。

宇宙飛行などにおいて、アポロ計画では火災事故以前の初期段階で^[74]、また最新の宇宙服などにて比較的低圧で封じるため純酸素ガスが使用された^[75]^[76]。最新の宇宙服では、服内を0.3気圧程度まで減圧した純酸素で満たし、血液中の酸素分圧が上昇しない方法が取られている^[77]^[78]。

肺や中枢神経系に及ぼす酸素中毒は、深い水深へのスクーバダイビング（ディープダイビング）や送気式潜水でも起こる可能性がある^[51]^[71]。酸素分圧60キロパスカル以上の空気を長い時間呼吸していることは、恒久的な肺線維症に至ることがある^[79]。これがさらに高い160キロパスカル以上となると、ダイバーにとって致命的になる痙攣に繋がることもありうる。深刻な酸素中毒は、酸素比率21%の空気を用いながら66メートル以上潜水することで起こるが、同様のことは比率100%の空気ならばわずか6メートルの潜水で起こる^[79]^[80]^[81]^[82]。

アポロ1号実験中に発生した火災現場の写真。当時は純酸素が使われていた

燃焼と他の危険

高濃度酸素と可燃物が混在している状況で、そこに何らかの火種があれば火災や爆発など激しい燃焼が引き起こされる^[83]。酸素そのものは燃えないが、酸化剤として作用する。燃焼発生の危険は、酸素が酸化電位の高い物質、例えば過酸化物や塩素酸塩、硝酸塩や過塩素酸塩、クロム酸塩などと混在している場合も高い。

大気中酸素濃度の減少

現在、地球の大気中における酸素濃度は約21% (209,490 ppm) であるが、年平均 4 ppm ずつ減少している（1999年から2005年の平均値）という調査結果がある^[84]。一方で、大気中の二酸化炭素濃度は年平均2 ppmずつ増加しており、酸素濃度の減少もこれに関連して化石燃料の燃焼などが主な原因になっていると思われる。また、二酸化炭素濃度の増加量と酸素濃度の減少量の差は、二酸化炭素が海面で多く吸収されている（陸上の約2倍）ことや化石燃料燃焼時に二酸化炭素排出量より酸素消費量の方が1.4倍多いことなどに起因する。大気中酸素濃度の1年間を通した変動では、陸上における光合成量が呼吸量を上回る北半球の夏季には増加しており、冬季には減少している。

もっとも、大気中の二酸化炭素濃度は2006年時点で約0.038% (381 ppm) 程であり、約21%の酸素とは元々の大気中濃度がまったく異なっている。年平均 4 ppm の減少は1000年間で0.4%程度の酸素濃度の減少であり、地球上における生態圏への影響は微々たるものである。

脚注

注釈

^ 原初の地球大気にも、水蒸気が光分解されて発生するメカニズムが指摘されており、ごく微量ながら酸素ガスが存在した可能性はあるが、ほとんどはすぐ酸化反応で消費されるか、オゾンへ変化したものと思われ、いずれにしろ考慮に足る量ではなかった^[29]。
^ (0.36 g/分/人) × (60秒/時) × (24時/日) × (365日/年) × (70億人)/1,000,000 = 13.2億トン

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^ “2005年代表選抜試験で出題された問題とその解説”. 国際生物学オリンピック日本委員会. 2010年5月28日閲覧。
^ Morgenthaler, G. W.; Fester, D. A.; Cooley, C. G. (1994). “As assessment of habitat pressure, oxygen fraction, and EVA suit design for space operations”. Acta Astronaut 32 (1): 39–49. doi:10.1016/0094-5765(94)90146-5. PMID 11541018.
^ Wade, Mark (2007年). “Space Suits”. Encyclopedia Astronautica. 2007年12月16日閲覧。
^ 鳥居啓之（川崎重工業）. “船外活動技術における生命維持技術の検討”. 宇宙航空研究開発機構. 2010年5月28日閲覧。
^ “宇宙服と船外活動”. 宇宙航空研究開発機構. 2010年5月28日閲覧。
^ ^a ^b Wilmshurst, P. (1998). “Diving and oxygen”. BMJ 317 (7164): 996–999. PMC 1114047. PMID 9765173.
^ Donald, Kenneth W. (1992). Oxygen and the Diver. England: SPA in conjunction with K. Donald. ISBN 1854211765.
^ Donald, Kenneth W. (1947). “Oxygen Poisoning in Man: Part I”. British Medical Journal 1 (4506): 667–672. doi:10.1136/bmj.1.4506.667. PMC 2053251.
^ Donald, K. W. (1947). “Oxygen Poisoning in Man: Part II”. British Medical Journal 1 (4507): 712–717. doi:10.1136/bmj.1.4507.712. PMC 2053400.
^ Werley, Barry L., ed (1991). “Fire Hazards in Oxygen Systems”. ASTM Technical Professional training. Philadelphia: ASTM International Subcommittee G-4.05
^ 大気中酸素濃度の減少量から二酸化炭素の陸域生物圏吸収量の推定に成功（国立環境研究所）

参考文献

桜井弘『元素111の新知識』講談社、1998年（初版1997年）、第6刷。ISBN 4-06-257192-7。
玉尾皓平、桜井弘、福山秀敏（監修）『完全図解周期表』ニュートンプレス〈ニュートン別冊〉、2010年、第2版。ISBN 978-4-315-51876-4。
J・D・リー『リー無機化学』浜口博、菅野等（訳）、東京化学同人、2005年（初版1982年）、第1版第22刷。ISBN 4-8079-0185-0。
瀬名秀明、太田成男『ミトコンドリアと生きる』角川書店、2000年、第一刷。ISBN 4-04-704006-1。
Cook, Gerhard A.; Lauer, Carol M. (1968). “Oxygen”. In Hampel, Clifford A. The Encyclopedia of the Chemical Elements. New York: Reinhold Book Corporation. pp. 499–512. LCCN 68-29938. .
Emsley, John (2001). “Oxygen”. Nature's Building Blocks: An A-Z Guide to the Elements. Oxford, England, UK: Oxford University Press. pp. 297–304. ISBN 0198503407.
Raven, Peter H.; Ray F. Evert, Susan E. Eichhorn (2005). Biology of Plants, 7th Edition. New York: W.H. Freeman and Company. pp. 115–127. ISBN 0-7167-1007-2.

1 History
- 1.1 Early experiments
- 1.2 Phlogiston theory
- 1.3 Discovery
- 1.4 Lavoisier's contribution
- 1.5 Later history
2 Characteristics
- 2.1 Properties and molecular structure
- 2.2 Allotropes
- 2.3 Physical properties
- 2.4 Isotopes and stellar origin
- 2.5 Occurrence
- 2.6 Analysis
3 Biological role of O₂
- 3.1 Photosynthesis and respiration
- 3.2 Living organisms
- 3.3 Build-up in the atmosphere
4 Industrial production
5 Storage
6 Applications
- 6.1 Medical
- 6.2 Life support and recreational use
- 6.3 Industrial
7 Compounds
- 7.1 Oxides and other inorganic compounds
- 7.2 Organic compounds and biomolecules
8 Safety and precautions
- 8.1 Toxicity
- 8.2 Combustion and other hazards
9 See also
10 Notes
11 Citations
12 References
13 External links

History

Early experiments

Philo's experiment inspired later investigators.

One of the first known experiments on the relationship between combustion and air was conducted by the 2nd century BCE Greek writer on mechanics, Philo of Byzantium. In his work Pneumatica, Philo observed that inverting a vessel over a burning candle and surrounding the vessel's neck with water resulted in some water rising into the neck.^[9] Philo incorrectly surmised that parts of the air in the vessel were converted into the classical element fire and thus were able to escape through pores in the glass. Many centuries later Leonardo da Vinci built on Philo's work by observing that a portion of air is consumed during combustion and respiration.^[10]

In the late 17th century, Robert Boyle proved that air is necessary for combustion. English chemist John Mayow (1641–1679) refined this work by showing that fire requires only a part of air that he called spiritus nitroaereus or just nitroaereus.^[11] In one experiment, he found that placing either a mouse or a lit candle in a closed container over water caused the water to rise and replace one-fourteenth of the air's volume before extinguishing the subjects.^[12] From this he surmised that nitroaereus is consumed in both respiration and combustion.

Mayow observed that antimony increased in weight when heated, and inferred that the nitroaereus must have combined with it.^[11] He also thought that the lungs separate nitroaereus from air and pass it into the blood and that animal heat and muscle movement result from the reaction of nitroaereus with certain substances in the body.^[11] Accounts of these and other experiments and ideas were published in 1668 in his work Tractatus duo in the tract "De respiratione".^[12]

Phlogiston theory

Stahl helped develop and popularize the phlogiston theory.

Robert Hooke, Ole Borch, Mikhail Lomonosov, and Pierre Bayen (fr) all produced oxygen in experiments in the 17th and the 18th century but none of them recognized it as a chemical element.^[13] This may have been in part due to the prevalence of the philosophy of combustion and corrosion called the phlogiston theory, which was then the favored explanation of those processes.

Established in 1667 by the German alchemist J. J. Becher, and modified by the chemist Georg Ernst Stahl by 1731,^[14] phlogiston theory stated that all combustible materials were made of two parts. One part, called phlogiston, was given off when the substance containing it was burned, while the dephlogisticated part was thought to be its true form, or calx.^[10]

Highly combustible materials that leave little residue, such as wood or coal, were thought to be made mostly of phlogiston; non-combustible substances that corrode, such as iron, contained very little. Air did not play a role in phlogiston theory, nor were any initial quantitative experiments conducted to test the idea; instead, it was based on observations of what happens when something burns, that most common objects appear to become lighter and seem to lose something in the process.^[10] The fact that a substance like wood gains overall weight in burning was hidden by the buoyancy of the gaseous combustion products. Indeed, one of the first clues that the phlogiston theory was incorrect was that metals gain weight in rusting (when they were supposedly losing phlogiston).

Discovery

Carl Wilhelm Scheele beat Priestley to the discovery but published afterwards.

Oxygen was first discovered by Swedish pharmacist Carl Wilhelm Scheele. He had produced oxygen gas by heating mercuric oxide and various nitrates by about 1772.^[5]^[10] Scheele called the gas "fire air" because it was the only known supporter of combustion, and wrote an account of this discovery in a manuscript he titled Treatise on Air and Fire, which he sent to his publisher in 1775. That document was published in 1777.^[15]

Joseph Priestley is usually given priority in the discovery.

In the meantime, on August 1, 1774, an experiment conducted by the British clergyman Joseph Priestley focused sunlight on mercuric oxide (HgO) inside a glass tube, which liberated a gas he named "dephlogisticated air".^[5] He noted that candles burned brighter in the gas and that a mouse was more active and lived longer while breathing it. After breathing the gas himself, he wrote: "The feeling of it to my lungs was not sensibly different from that of common air, but I fancied that my breast felt peculiarly light and easy for some time afterwards."^[13] Priestley published his findings in 1775 in a paper titled "An Account of Further Discoveries in Air" which was included in the second volume of his book titled Experiments and Observations on Different Kinds of Air.^[10]^[16] Because he published his findings first, Priestley is usually given priority in the discovery.

The French chemist Antoine Laurent Lavoisier later claimed to have discovered the new substance independently. Priestley visited Lavoisier in October 1774 and told him about his experiment and how he liberated the new gas. Scheele also posted a letter to Lavoisier on September 30, 1774 that described his discovery of the previously unknown substance, but Lavoisier never acknowledged receiving it (a copy of the letter was found in Scheele's belongings after his death).^[15]

Lavoisier's contribution

What Lavoisier did (although this was disputed at the time) was to conduct the first adequate quantitative experiments on oxidation and give the first correct explanation of how combustion works.^[5] He used these and similar experiments, all started in 1774, to discredit the phlogiston theory and to prove that the substance discovered by Priestley and Scheele was a chemical element.

Antoine Lavoisier discredited the phlogiston theory.

In one experiment, Lavoisier observed that there was no overall increase in weight when tin and air were heated in a closed container.^[5] He noted that air rushed in when he opened the container, which indicated that part of the trapped air had been consumed. He also noted that the tin had increased in weight and that increase was the same as the weight of the air that rushed back in. This and other experiments on combustion were documented in his book Sur la combustion en général, which was published in 1777.^[5] In that work, he proved that air is a mixture of two gases; 'vital air', which is essential to combustion and respiration, and azote (Gk. ἄζωτον "lifeless"), which did not support either. Azote later became nitrogen in English, although it has kept the name in French and several other European languages.^[5]

Lavoisier renamed 'vital air' to oxygène in 1777 from the Greek roots ὀξύς (oxys) (acid, literally "sharp", from the taste of acids) and -γενής (-genēs) (producer, literally begetter), because he mistakenly believed that oxygen was a constituent of all acids.^[8] Chemists (such as Sir Humphry Davy in 1812) eventually determined that Lavoisier was wrong in this regard (hydrogen forms the basis for acid chemistry), but by then the name was too well established.

Oxygen entered the English language despite opposition by English scientists and the fact that the Englishman Priestley had first isolated the gas and written about it. This is partly due to a poem praising the gas titled "Oxygen" in the popular book The Botanic Garden (1791) by Erasmus Darwin, grandfather of Charles Darwin.^[15]

Later history

Robert H. Goddard and a liquid oxygen-gasoline rocket

John Dalton's original atomic hypothesis presumed that all elements were monatomic and that the atoms in compounds would normally have the simplest atomic ratios with respect to one another. For example, Dalton assumed that water's formula was HO, giving the atomic mass of oxygen was 8 times that of hydrogen, instead of the modern value of about 16.^[17] In 1805, Joseph Louis Gay-Lussac and Alexander von Humboldt showed that water is formed of two volumes of hydrogen and one volume of oxygen; and by 1811 Amedeo Avogadro had arrived at the correct interpretation of water's composition, based on what is now called Avogadro's law and the diatomic elemental molecules in those gases.^[18]^[a]

By the late 19th century scientists realized that air could be liquefied and its components isolated by compressing and cooling it. Using a cascade method, Swiss chemist and physicist Raoul Pierre Pictet evaporated liquid sulfur dioxide in order to liquefy carbon dioxide, which in turn was evaporated to cool oxygen gas enough to liquefy it. He sent a telegram on December 22, 1877 to the French Academy of Sciences in Paris announcing his discovery of liquid oxygen.^[19] Just two days later, French physicist Louis Paul Cailletet announced his own method of liquefying molecular oxygen.^[19] Only a few drops of the liquid were produced in each case and no meaningful analysis could be conducted. Oxygen was liquified in a stable state for the first time on March 29, 1883 by Polish scientists from Jagiellonian University, Zygmunt Wróblewski and Karol Olszewski.^[20]

In 1891 Scottish chemist James Dewar was able to produce enough liquid oxygen for study.^[21] The first commercially viable process for producing liquid oxygen was independently developed in 1895 by German engineer Carl von Linde and British engineer William Hampson. Both men lowered the temperature of air until it liquefied and then distilled the component gases by boiling them off one at a time and capturing them.^[22] Later, in 1901, oxyacetylene welding was demonstrated for the first time by burning a mixture of acetylene and compressed O
2. This method of welding and cutting metal later became common.^[22]

In 1923, the American scientist Robert H. Goddard became the first person to develop a rocket engine that burned liquid fuel; the engine used gasoline for fuel and liquid oxygen as the oxidizer. Goddard successfully flew a small liquid-fueled rocket 56 m at 97 km/h on March 16, 1926 in Auburn, Massachusetts, US.^[22]^[23]

Oxygen levels in the atmosphere are trending slightly downward globally, possibly because of fossil-fuel burning.^[24]

Characteristics

Properties and molecular structure

Orbital diagram, after Barrett (2002),^[25] showing the participating atomic orbitals from each oxygen atom, the molecular orbitals that result from their overlap, and the Aufbau filling of the orbitals with the 12 electrons, 6 from each O atom, beginning from the lowest energy orbitals, and resulting in covalent double bond character from filled orbitals (and cancelation of the contributions of the pairs of σ and σ^* and

π

and

π

^* orbital pairs).

At standard temperature and pressure, oxygen is a colorless, odorless, and tasteless gas with the molecular formula O
2, referred to as dioxygen.^[26]

As dioxygen, two oxygen atoms are chemically bound to each other. The bond can be variously described based on level of theory, but is reasonably and simply described as a covalent double bond that results from the filling of molecular orbitals formed from the atomic orbitals of the individual oxygen atoms, the filling of which results in a bond order of two. More specifically, the double bond is the result of sequential, low-to-high energy, or Aufbau, filling of orbitals, and the resulting cancellation of contributions from the 2s electrons, after sequential filling of the low σ and σ^* orbitals; σ overlap of the two atomic 2p orbitals that lie along the O-O molecular axis and $π$ overlap of two pairs of atomic 2p orbitals perpendicular to the O-O molecular axis, and then cancellation of contributions from the remaining two of the six 2p electrons after their partial filling of the lowest $π$ and $π$ ^* orbitals.^[25]

This combination of cancellations and σ and $π$ overlaps results in dioxygen's double bond character and reactivity, and a triplet electronic ground state. An electron configuration with two unpaired electrons, as is found in dioxygen (see the filled $π$ * orbitals in the diagram) orbitals that are of equal energy—i.e., degenerate—is a configuration termed a spin triplet state. Hence, the ground state of the O
2 molecule is referred to as triplet oxygen.^[27]^[b] The highest energy, partially filled orbitals are antibonding, and so their filling weakens the bond order from three to two. Because of its unpaired electrons, triplet oxygen reacts only slowly with most organic molecules, which have paired electron spins; this prevents spontaneous combustion.^[28]

A trickle of liquid oxygen is deflected by a magnetic field, illustrating its paramagnetic property

In the triplet form, O
2 molecules are paramagnetic. That is, they impart magnetic character to oxygen when it is in the presence of a magnetic field, because of the spin magnetic moments of the unpaired electrons in the molecule, and the negative exchange energy between neighboring O
2 molecules.^[21] Liquid oxygen is so magnetic that, in laboratory demonstrations, a bridge of liquid oxygen may be supported against its own weight between the poles of a powerful magnet.^[29]^[c]

Singlet oxygen is a name given to several higher-energy species of molecular O
2 in which all the electron spins are paired. It is much more reactive with common organic molecules than is molecular oxygen per se. In nature, singlet oxygen is commonly formed from water during photosynthesis, using the energy of sunlight.^[30] It is also produced in the troposphere by the photolysis of ozone by light of short wavelength,^[31] and by the immune system as a source of active oxygen.^[32] Carotenoids in photosynthetic organisms (and possibly animals) play a major role in absorbing energy from singlet oxygen and converting it to the unexcited ground state before it can cause harm to tissues.^[33]

Allotropes

Ozone is a rare gas on Earth found mostly in the stratosphere.

Space-filling model representation of dioxygen (O₂) molecule

The common allotrope of elemental oxygen on Earth is called dioxygen, O
2, the major part of the Earth's atmospheric oxygen (see Occurrence). O₂ has a bond length of 121 pm and a bond energy of 498 kJ·mol⁻¹,^[34] which is smaller than the energy of other double bonds or pairs of single bonds in the biosphere and responsible for the exothermic reaction of O₂ with any organic molecule.^[28]^[35] Due to its energy content, O₂ is used by complex forms of life, such as animals, in cellular respiration (see Biological role). Other aspects of O
2 are covered in the remainder of this article.

Trioxygen (O
3) is usually known as ozone and is a very reactive allotrope of oxygen that is damaging to lung tissue.^[36] Ozone is produced in the upper atmosphere when O
2 combines with atomic oxygen made by the splitting of O
2 by ultraviolet (UV) radiation.^[8] Since ozone absorbs strongly in the UV region of the spectrum, the ozone layer of the upper atmosphere functions as a protective radiation shield for the planet.^[8] Near the Earth's surface, it is a pollutant formed as a by-product of automobile exhaust.^[36] The metastable molecule tetraoxygen (O
4) was discovered in 2001,^[37]^[38] and was assumed to exist in one of the six phases of solid oxygen. It was proven in 2006 that this phase, created by pressurizing O
2 to 20 GPa, is in fact a rhombohedral O
8 cluster.^[39] This cluster has the potential to be a much more powerful oxidizer than either O
2 or O
3 and may therefore be used in rocket fuel.^[37]^[38] A metallic phase was discovered in 1990 when solid oxygen is subjected to a pressure of above 96 GPa^[40] and it was shown in 1998 that at very low temperatures, this phase becomes superconducting.^[41]

Physical properties

Oxygen discharge (spectrum) tube. The green color is similar to the color of an "aurora borealis"

Oxygen dissolves more readily in water than nitrogen, and in freshwater more readily than seawater. Water in equilibrium with air contains approximately 1 molecule of dissolved O
2 for every 2 molecules of N
2 (1:2), compared with an atmospheric ratio of approximately 1:4. The solubility of oxygen in water is temperature-dependent, and about twice as much (14.6 mg·L⁻¹) dissolves at 0 °C than at 20 °C (7.6 mg·L⁻¹).^[13]^[42] At 25 °C and 1 standard atmosphere (101.3 kPa) of air, freshwater contains about 6.04 milliliters (mL) of oxygen per liter, and seawater contains about 4.95 mL per liter.^[43] At 5 °C the solubility increases to 9.0 mL (50% more than at 25 °C) per liter for water and 7.2 mL (45% more) per liter for sea water.

Oxygen gas dissolved in water at sealevel
	5 °C	25 °C
Freshwater	9.0 mL	6.04 mL
Seawater	7.2 mL	4.95 mL

Oxygen condenses at 90.20 K (−182.95 °C, −297.31 °F), and freezes at 54.36 K (−218.79 °C, −361.82 °F).^[44] Both liquid and solid O
2 are clear substances with a light sky-blue color caused by absorption in the red (in contrast with the blue color of the sky, which is due to Rayleigh scattering of blue light). High-purity liquid O
2 is usually obtained by the fractional distillation of liquefied air.^[45] Liquid oxygen may also be condensed from air using liquid nitrogen as a coolant.^[46]

Oxygen is a highly reactive substance and must be segregated from combustible materials.^[46]

The spectroscopy of molecular oxygen is associated with the atmospheric processes of aurora, airglow and nightglow.^[47] The absorption in the Herzberg continuum and Schumann–Runge bands in the ultraviolet produces atomic oxygen that is important in the chemistry of the middle atmosphere.^[48] Excited state singlet molecular oxygen is responsible for red chemiluminescence in solution.^[49]

Isotopes and stellar origin

Late in a massive star's life, ¹⁶O concentrates in the O-shell, ¹⁷O in the H-shell and ¹⁸O in the He-shell.

Naturally occurring oxygen is composed of three stable isotopes, ¹⁶O, ¹⁷O, and ¹⁸O, with ¹⁶O being the most abundant (99.762% natural abundance).^[50]

Most ¹⁶O is synthesized at the end of the helium fusion process in massive stars but some is made in the neon burning process.^[51] ¹⁷O is primarily made by the burning of hydrogen into helium during the CNO cycle, making it a common isotope in the hydrogen burning zones of stars.^[51] Most ¹⁸O is produced when ¹⁴N (made abundant from CNO burning) captures a ⁴He nucleus, making ¹⁸O common in the helium-rich zones of evolved, massive stars.^[51]

Fourteen radioisotopes have been characterized. The most stable are ¹⁵O with a half-life of 122.24 seconds and ¹⁴O with a half-life of 70.606 seconds.^[50] All of the remaining radioactive isotopes have half-lives that are less than 27 s and the majority of these have half-lives that are less than 83 milliseconds.^[50] The most common decay mode of the isotopes lighter than ¹⁶O is β⁺ decay^[52]^[53]^[54] to yield nitrogen, and the most common mode for the isotopes heavier than ¹⁸O is beta decay to yield fluorine.^[50]

Occurrence

Ten most common elements in the Milky Way Galaxy estimated spectroscopically^[55]
Z	Element	Mass fraction in parts per million
1	Hydrogen	739,000	71 × mass of oxygen (red bar)
2	Helium	240,000	23 × mass of oxygen (red bar)
8	Oxygen	10,400	10400
6	Carbon	4,600	4600
10	Neon	1,340	1340
26	Iron	1,090	1090
7	Nitrogen	960	960
14	Silicon	650	650
12	Magnesium	580	580
16	Sulfur	440	440

Oxygen is the most abundant chemical element by mass in the Earth's biosphere, air, sea and land. Oxygen is the third most abundant chemical element in the universe, after hydrogen and helium.^[4] About 0.9% of the Sun's mass is oxygen.^[5] Oxygen constitutes 49.2% of the Earth's crust by mass^[6] as part of oxide compounds such as silicon dioxide and is most abundant element by mass in the Earth's crust. It is also the major component of the world's oceans (88.8% by mass).^[5] Oxygen gas is the second most common component of the Earth's atmosphere, taking up 20.8% of its volume and 23.1% of its mass (some 10¹⁵ tonnes).^[5]^[56]^[d] Earth is unusual among the planets of the Solar System in having such a high concentration of oxygen gas in its atmosphere: Mars (with 0.1% O
2 by volume) and Venus have much less. The O
2 surrounding those planets is produced solely by ultraviolet radiation on oxygen-containing molecules such as carbon dioxide.

The unusually high concentration of oxygen gas on Earth is the result of the oxygen cycle. This biogeochemical cycle describes the movement of oxygen within and between its three main reservoirs on Earth: the atmosphere, the biosphere, and the lithosphere. The main driving factor of the oxygen cycle is photosynthesis, which is responsible for modern Earth's atmosphere. Photosynthesis releases oxygen into the atmosphere, while respiration, decay, and combustion remove it from the atmosphere. In the present equilibrium, production and consumption occur at the same rate of roughly 1/2000th of the entire atmospheric oxygen per year.

Cold water holds more dissolved O

2.

Free oxygen also occurs in solution in the world's water bodies. The increased solubility of O
2 at lower temperatures (see Physical properties) has important implications for ocean life, as polar oceans support a much higher density of life due to their higher oxygen content.^[57] Water polluted with plant nutrients such as nitrates or phosphates may stimulate growth of algae by a process called eutrophication and the decay of these organisms and other biomaterials may reduce the O
2 content in eutrophic water bodies. Scientists assess this aspect of water quality by measuring the water's biochemical oxygen demand, or the amount of O
2 needed to restore it to a normal concentration.^[58]

Analysis

500 million years of climate change vs ¹⁸O

Paleoclimatologists measure the ratio of oxygen-18 and oxygen-16 in the shells and skeletons of marine organisms to determine the climate millions of years ago (see oxygen isotope ratio cycle). Seawater molecules that contain the lighter isotope, oxygen-16, evaporate at a slightly faster rate than water molecules containing the 12% heavier oxygen-18, and this disparity increases at lower temperatures.^[59] During periods of lower global temperatures, snow and rain from that evaporated water tends to be higher in oxygen-16, and the seawater left behind tends to be higher in oxygen-18. Marine organisms then incorporate more oxygen-18 into their skeletons and shells than they would in a warmer climate.^[59] Paleoclimatologists also directly measure this ratio in the water molecules of ice core samples as old as hundreds of thousands of years.

Planetary geologists have measured the relative quantities of oxygen isotopes in samples from the Earth, the Moon, Mars, and meteorites, but were long unable to obtain reference values for the isotope ratios in the Sun, believed to be the same as those of the primordial solar nebula. Analysis of a silicon wafer exposed to the solar wind in space and returned by the crashed Genesis spacecraft has shown that the Sun has a higher proportion of oxygen-16 than does the Earth. The measurement implies that an unknown process depleted oxygen-16 from the Sun's disk of protoplanetary material prior to the coalescence of dust grains that formed the Earth.^[60]

Oxygen presents two spectrophotometric absorption bands peaking at the wavelengths 687 and 760 nm. Some remote sensing scientists have proposed using the measurement of the radiance coming from vegetation canopies in those bands to characterize plant health status from a satellite platform.^[61] This approach exploits the fact that in those bands it is possible to discriminate the vegetation's reflectance from its fluorescence, which is much weaker. The measurement is technically difficult owing to the low signal-to-noise ratio and the physical structure of vegetation; but it has been proposed as a possible method of monitoring the carbon cycle from satellites on a global scale.

Biological role of O₂

Photosynthesis and respiration

Photosynthesis splits water to liberate O
2 and fixes CO

2 into sugar in what is called a Calvin cycle.

In nature, free oxygen is produced by the light-driven splitting of water during oxygenic photosynthesis. According to some estimates, green algae and cyanobacteria in marine environments provide about 70% of the free oxygen produced on Earth, and the rest is produced by terrestrial plants.^[62] Other estimates of the oceanic contribution to atmospheric oxygen are higher, while some estimates are lower, suggesting oceans produce ~45% of Earth's atmospheric oxygen each year.^[63]

A simplified overall formula for photosynthesis is:^[64]

6 CO₂ + 6 H
2O + photons → C
6H
12O
6 + 6 O
2

or simply

carbon dioxide + water + sunlight → glucose + dioxygen

Photolytic oxygen evolution occurs in the thylakoid membranes of photosynthetic organisms and requires the energy of four photons.^[e] Many steps are involved, but the result is the formation of a proton gradient across the thylakoid membrane, which is used to synthesize adenosine triphosphate (ATP) via photophosphorylation.^[65] The O
2 remaining (after production of the water molecule) is released into the atmosphere.^[f]

Molecular dioxygen, O
2, is essential for cellular respiration in all aerobic organisms. Oxygen is used in mitochondria to generate ATP during oxidative phosphorylation. The reaction for aerobic respiration is essentially the reverse of photosynthesis and is simplified as:

C
6H
12O
6 + 6 O
2 → 6 CO₂ + 6 H
2O + 2880 kJ·mol⁻¹

In vertebrates, O
2 diffuses through membranes in the lungs and into red blood cells. Hemoglobin binds O
2, changing color from bluish red to bright red^[36] (CO
2 is released from another part of hemoglobin through the Bohr effect). Other animals use hemocyanin (molluscs and some arthropods) or hemerythrin (spiders and lobsters).^[56] A liter of blood can dissolve 200 cm³ of O
2.^[56]

Until the discovery of anaerobic metazoa,^[66] oxygen was thought to be a requirement for all complex life.^[67]

Reactive oxygen species, such as superoxide ion (O−
2) and hydrogen peroxide (H
2O
2), are dangerous by-products of oxygen use in organisms.^[56] Parts of the immune system of higher organisms create peroxide, superoxide, and singlet oxygen to destroy invading microbes. Reactive oxygen species also play an important role in the hypersensitive response of plants against pathogen attack.^[65] Oxygen is toxic to obligately anaerobic organisms, which were the dominant form of early life on Earth until O
2 began to accumulate in the atmosphere about 2.5 billion years ago during the Great Oxygenation Event, about a billion years after the first appearance of these organisms.^[68]^[69]

An adult human at rest inhales 1.8 to 2.4 grams of oxygen per minute.^[70] This amounts to more than 6 billion tonnes of oxygen inhaled by humanity per year.^[g]

Living organisms

Partial pressures of oxygen in the human body (PO₂)
Unit	Alveolar pulmonary gas pressures	Arterial blood oxygen	Venous blood gas
kPa	14.2	11^[71]-13^[71]	4.0^[71]-5.3^[71]
mmHg	107	75^[72]-100^[72]	30^[73]-40^[73]

The free oxygen partial pressure in the body of a living vertebrate organism is highest in the respiratory system, and decreases along any arterial system, peripheral tissues, and venous system, respectively. Partial pressure is the pressure that oxygen would have if it alone occupied the volume.^[74]

Build-up in the atmosphere

O
2 build-up in Earth's atmosphere: 1) no O
2 produced; 2) O
2 produced, but absorbed in oceans & seabed rock; 3) O
2 starts to gas out of the oceans, but is absorbed by land surfaces and formation of ozone layer; 4–5) O

2 sinks filled and the gas accumulates

Free oxygen gas was almost nonexistent in Earth's atmosphere before photosynthetic archaea and bacteria evolved, probably about 3.5 billion years ago. Free oxygen first appeared in significant quantities during the Paleoproterozoic eon (between 3.0 and 2.3 billion years ago).^[75] For the first billion years, any free oxygen produced by these organisms combined with dissolved iron in the oceans to form banded iron formations. When such oxygen sinks became saturated, free oxygen began to outgas from the oceans 3–2.7 billion years ago, reaching 10% of its present level around 1.7 billion years ago.^[75]^[76]

The presence of large amounts of dissolved and free oxygen in the oceans and atmosphere may have driven most of the extant anaerobic organisms to extinction during the Great Oxygenation Event (oxygen catastrophe) about 2.4 billion years ago. Cellular respiration using O
2 enables aerobic organisms to produce much more ATP than anaerobic organisms.^[77] Cellular respiration of O
2 occurs in all eukaryotes, including all complex multicellular organisms such as plants and animals.

Since the beginning of the Cambrian period 540 million years ago, atmospheric O
2 levels have fluctuated between 15% and 30% by volume.^[78] Towards the end of the Carboniferous period (about 300 million years ago) atmospheric O
2 levels reached a maximum of 35% by volume,^[78] which may have contributed to the large size of insects and amphibians at this time.^[79]

Variations of oxygen shaped the climates of the past. When oxygen declined, atmospheric density dropped and this in turn increased surface evaporation, and led to precipitation increases and warmer temperatures.^[80]

At the current rate of photosynthesis it would take about 2,000 years to regenerate the entire O
2 in the present atmosphere.^[81]

Industrial production

Hofmann electrolysis apparatus used in electrolysis of water.

One hundred million tonnes of O
2 are extracted from air for industrial uses annually by two primary methods.^[15] The most common method is fractional distillation of liquefied air, with N
2 distilling as a vapor while O
2 is left as a liquid.^[15]

The other primary method of producing O
2 is passing a stream of clean, dry air through one bed of a pair of identical zeolite molecular sieves, which absorbs the nitrogen and delivers a gas stream that is 90% to 93% O
2.^[15] Simultaneously, nitrogen gas is released from the other nitrogen-saturated zeolite bed, by reducing the chamber operating pressure and diverting part of the oxygen gas from the producer bed through it, in the reverse direction of flow. After a set cycle time the operation of the two beds is interchanged, thereby allowing for a continuous supply of gaseous oxygen to be pumped through a pipeline. This is known as pressure swing adsorption. Oxygen gas is increasingly obtained by these non-cryogenic technologies (see also the related vacuum swing adsorption).^[82]

Oxygen gas can also be produced through electrolysis of water into molecular oxygen and hydrogen. DC electricity must be used: if AC is used, the gases in each limb consist of hydrogen and oxygen in the explosive ratio 2:1. Contrary to popular belief, the 2:1 ratio observed in the DC electrolysis of acidified water does not prove that the empirical formula of water is H₂O unless certain assumptions are made about the molecular formulae of hydrogen and oxygen themselves. A similar method is the electrocatalytic O
2 evolution from oxides and oxoacids. Chemical catalysts can be used as well, such as in chemical oxygen generators or oxygen candles that are used as part of the life-support equipment on submarines, and are still part of standard equipment on commercial airliners in case of depressurization emergencies. Another air separation method is forcing air to dissolve through ceramic membranes based on zirconium dioxide by either high pressure or an electric current, to produce nearly pure O
2 gas.^[58]

In large quantities, the price of liquid oxygen in 2001 was approximately $0.21/kg.^[83] Since the primary cost of production is the energy cost of liquefying the air, the production cost will change as energy cost varies.

Storage

Oxygen storage methods include high pressure oxygen tanks, cryogenics and chemical compounds. For reasons of economy, oxygen is often transported in bulk as a liquid in specially insulated tankers, since one liter of liquefied oxygen is equivalent to 840 liters of gaseous oxygen at atmospheric pressure and 20 °C (68 °F).^[15] Such tankers are used to refill bulk liquid oxygen storage containers, which stand outside hospitals and other institutions that need large volumes of pure oxygen gas. Liquid oxygen is passed through heat exchangers, which convert the cryogenic liquid into gas before it enters the building. Oxygen is also stored and shipped in smaller cylinders containing the compressed gas; a form that is useful in certain portable medical applications and oxy-fuel welding and cutting.^[15]

Applications

Medical

An oxygen concentrator in an emphysema patient's house

Uptake of O
2 from the air is the essential purpose of respiration, so oxygen supplementation is used in medicine. Treatment not only increases oxygen levels in the patient's blood, but has the secondary effect of decreasing resistance to blood flow in many types of diseased lungs, easing work load on the heart. Oxygen therapy is used to treat emphysema, pneumonia, some heart disorders (congestive heart failure), some disorders that cause increased pulmonary artery pressure, and any disease that impairs the body's ability to take up and use gaseous oxygen.^[84]

Treatments are flexible enough to be used in hospitals, the patient's home, or increasingly by portable devices. Oxygen tents were once commonly used in oxygen supplementation, but have since been replaced mostly by the use of oxygen masks or nasal cannulas.^[85]

Hyperbaric (high-pressure) medicine uses special oxygen chambers to increase the partial pressure of O
2 around the patient and, when needed, the medical staff.^[86] Carbon monoxide poisoning, gas gangrene, and decompression sickness (the 'bends') are sometimes addressed with this therapy.^[87] Increased O
2 concentration in the lungs helps to displace carbon monoxide from the heme group of hemoglobin.^[88]^[89] Oxygen gas is poisonous to the anaerobic bacteria that cause gas gangrene, so increasing its partial pressure helps kill them.^[90]^[91] Decompression sickness occurs in divers who decompress too quickly after a dive, resulting in bubbles of inert gas, mostly nitrogen and helium, forming in the blood. Increasing the pressure of O
2 as soon as possible helps to redissolve the bubbles back into the blood so that these excess gasses can be exhaled naturally through the lungs.^[84]^[92]^[93]

Oxygen is also used medically for patients who require mechanical ventilation, often at concentrations above the 21% found in ambient air.

Life support and recreational use

Low pressure pure O

2 is used in space suits.

An application of O
2 as a low-pressure breathing gas is in modern space suits, which surround their occupant's body with pressurized air. These devices use nearly pure oxygen at about one third normal pressure, resulting in a normal blood partial pressure of O
2.^[94]^[95] This trade-off of higher oxygen concentration for lower pressure is needed to maintain suit flexibility.

Scuba divers and submariners also rely on artificially delivered O
2, but most often use normal pressure, and/or mixtures of oxygen and air. Pure or nearly pure O
2 use in diving at higher-than-sea-level pressures is usually limited to rebreather, decompression, or emergency treatment use at relatively shallow depths (~6 meters depth, or less).^[96]^[97] Deeper diving requires significant dilution of O
2 with other gases, such as nitrogen or helium, to prevent oxygen toxicity.^[96]

People who climb mountains or fly in non-pressurized fixed-wing aircraft sometimes have supplemental O
2 supplies.^[h] Pressurized commercial airplanes have an emergency supply of O
2 automatically supplied to the passengers in case of cabin depressurization. Sudden cabin pressure loss activates chemical oxygen generators above each seat, causing oxygen masks to drop. Pulling on the masks "to start the flow of oxygen" as cabin safety instructions dictate, forces iron filings into the sodium chlorate inside the canister.^[58] A steady stream of oxygen gas is then produced by the exothermic reaction.

Oxygen, as a supposed mild euphoric, has a history of recreational use in oxygen bars and in sports. Oxygen bars are establishments found in Japan, California, and Las Vegas, Nevada since the late 1990s that offer higher than normal O
2 exposure for a fee.^[98] Professional athletes, especially in American football, sometimes go off-field between plays to don oxygen masks to boost performance. The pharmacological effect is doubted; a placebo effect is a more likely explanation.^[98] Available studies support a performance boost from enriched O
2 mixtures only if it is breathed during aerobic exercise.^[99]

Other recreational uses that do not involve breathing include pyrotechnic applications, such as George Goble's five-second ignition of barbecue grills.^[100]

Industrial

Most commercially produced O

2 is used to smelt iron into steel.

Smelting of iron ore into steel consumes 55% of commercially produced oxygen.^[58] In this process, O
2 is injected through a high-pressure lance into molten iron, which removes sulfur impurities and excess carbon as the respective oxides, SO
2 and CO
2. The reactions are exothermic, so the temperature increases to 1,700 °C.^[58]

Another 25% of commercially produced oxygen is used by the chemical industry.^[58] Ethylene is reacted with O
2 to create ethylene oxide, which, in turn, is converted into ethylene glycol; the primary feeder material used to manufacture a host of products, including antifreeze and polyester polymers (the precursors of many plastics and fabrics).^[58]

Most of the remaining 20% of commercially produced oxygen is used in medical applications, metal cutting and welding, as an oxidizer in rocket fuel, and in water treatment.^[58] Oxygen is used in oxyacetylene welding burning acetylene with O
2 to produce a very hot flame. In this process, metal up to 60 cm (24 in) thick is first heated with a small oxy-acetylene flame and then quickly cut by a large stream of O
2.^[101]

Compounds

Water (H

2O) is the most familiar oxygen compound.

The oxidation state of oxygen is −2 in almost all known compounds of oxygen. The oxidation state −1 is found in a few compounds such as peroxides.^[102] Compounds containing oxygen in other oxidation states are very uncommon: −1/2 (superoxides), −1/3 (ozonides), 0 (elemental, hypofluorous acid), +1/2 (dioxygenyl), +1 (dioxygen difluoride), and +2 (oxygen difluoride).

Oxides and other inorganic compounds

Water (H
2O) is an oxide of hydrogen and the most familiar oxygen compound. Hydrogen atoms are covalently bonded to oxygen in a water molecule but also have an additional attraction (about 23.3 kJ·mol⁻¹ per hydrogen atom) to an adjacent oxygen atom in a separate molecule.^[103] These hydrogen bonds between water molecules hold them approximately 15% closer than what would be expected in a simple liquid with just van der Waals forces.^[104]^[i]

Oxides, such as iron oxide or rust, form when oxygen combines with other elements.

Due to its electronegativity, oxygen forms chemical bonds with almost all other elements to give corresponding oxides. The surface of most metals, such as aluminium and titanium, are oxidized in the presence of air and become coated with a thin film of oxide that passivates the metal and slows further corrosion. Many oxides of the transition metals are non-stoichiometric compounds, with slightly less metal than the chemical formula would show. For example, the mineral FeO (wüstite) is written as Fe
1 − xO, where x is usually around 0.05.^[105]

Oxygen is present in the atmosphere in trace quantities in the form of carbon dioxide (CO
2). The Earth's crustal rock is composed in large part of oxides of silicon (silica SiO
2, as found in granite and quartz), aluminium (aluminium oxide Al
2O
3, in bauxite and corundum), iron (iron(III) oxide Fe
2O
3, in hematite and rust), and calcium carbonate (in limestone). The rest of the Earth's crust is also made of oxygen compounds, in particular various complex silicates (in silicate minerals). The Earth's mantle, of much larger mass than the crust, is largely composed of silicates of magnesium and iron.

Water-soluble silicates in the form of Na
4SiO
4, Na
2SiO
3, and Na
2Si
2O
5 are used as detergents and adhesives.^[106]

Oxygen also acts as a ligand for transition metals, forming transition metal dioxygen complexes, which feature metal–O
2. This class of compounds includes the heme proteins hemoglobin and myoglobin.^[107] An exotic and unusual reaction occurs with PtF
6, which oxidizes oxygen to give O₂⁺PtF₆⁻.^[108]

Organic compounds and biomolecules

Acetone is an important feeder material in the chemical industry.

Oxygen

Carbon

Hydrogen

Oxygen represents more than 40% of the molecular mass of the ATP molecule.

Among the most important classes of organic compounds that contain oxygen are (where "R" is an organic group): alcohols (R-OH); ethers (R-O-R); ketones (R-CO-R); aldehydes (R-CO-H); carboxylic acids (R-COOH); esters (R-COO-R); acid anhydrides (R-CO-O-CO-R); and amides (R-C(O)-NR
2). There are many important organic solvents that contain oxygen, including: acetone, methanol, ethanol, isopropanol, furan, THF, diethyl ether, dioxane, ethyl acetate, DMF, DMSO, acetic acid, and formic acid. Acetone ((CH
3)
2CO) and phenol (C
6H
5OH) are used as feeder materials in the synthesis of many different substances. Other important organic compounds that contain oxygen are: glycerol, formaldehyde, glutaraldehyde, citric acid, acetic anhydride, and acetamide. Epoxides are ethers in which the oxygen atom is part of a ring of three atoms.

Oxygen reacts spontaneously with many organic compounds at or below room temperature in a process called autoxidation.^[109] Most of the organic compounds that contain oxygen are not made by direct action of O
2. Organic compounds important in industry and commerce that are made by direct oxidation of a precursor include ethylene oxide and peracetic acid.^[106]

The element is found in almost all biomolecules that are important to (or generated by) life. Only a few common complex biomolecules, such as squalene and the carotenes, contain no oxygen. Of the organic compounds with biological relevance, carbohydrates contain the largest proportion by mass of oxygen. All fats, fatty acids, amino acids, and proteins contain oxygen (due to the presence of carbonyl groups in these acids and their ester residues). Oxygen also occurs in phosphate (PO3−
4) groups in the biologically important energy-carrying molecules ATP and ADP, in the backbone and the purines (except adenine) and pyrimidines of RNA and DNA, and in bones as calcium phosphate and hydroxylapatite.

Safety and precautions

The NFPA 704 standard rates compressed oxygen gas as nonhazardous to health, nonflammable and nonreactive, but an oxidizer. Refrigerated liquid oxygen (LOX) is given a health hazard rating of 3 (for increased risk of hyperoxia from condensed vapors, and for hazards common to cryogenic liquids such as frostbite), and all other ratings are the same as the compressed gas form.

Toxicity

Main symptoms of oxygen toxicity^[110]

Oxygen toxicity occurs when the lungs take in 2 ¹⁄₂ higher than normal O

2 partial pressure, which can occur in deep scuba diving.

Oxygen gas (O
2) can be toxic at elevated partial pressures, leading to convulsions and other health problems.^[96]^[j]^[111] Oxygen toxicity usually begins to occur at partial pressures more than 50 kilopascals (kPa), equal to about 50% oxygen composition at standard pressure or 2.5 times the normal sea-level O
2 partial pressure of about 21 kPa. This is not a problem except for patients on mechanical ventilators, since gas supplied through oxygen masks in medical applications is typically composed of only 30%–50% O
2 by volume (about 30 kPa at standard pressure).^[13] (although this figure also is subject to wide variation, depending on type of mask).

At one time, premature babies were placed in incubators containing O
2-rich air, but this practice was discontinued after some babies were blinded by the oxygen content being too high.^[13]

Breathing pure O
2 in space applications, such as in some modern space suits, or in early spacecraft such as Apollo, causes no damage due to the low total pressures used.^[94]^[112] In the case of spacesuits, the O
2 partial pressure in the breathing gas is, in general, about 30 kPa (1.4 times normal), and the resulting O
2 partial pressure in the astronaut's arterial blood is only marginally more than normal sea-level O
2 partial pressure (for more information on this, see space suit and arterial blood gas).

Oxygen toxicity to the lungs and central nervous system can also occur in deep scuba diving and surface supplied diving.^[13]^[96] Prolonged breathing of an air mixture with an O
2 partial pressure more than 60 kPa can eventually lead to permanent pulmonary fibrosis.^[113] Exposure to a O
2 partial pressures greater than 160 kPa (about 1.6 atm) may lead to convulsions (normally fatal for divers). Acute oxygen toxicity (causing seizures, its most feared effect for divers) can occur by breathing an air mixture with 21% O
2 at 66 m (217 ft) or more of depth; the same thing can occur by breathing 100% O
2 at only 6 m (20 ft).^[113]^[114]^[115]^[116]

Combustion and other hazards

The interior of the Apollo 1 Command Module. Pure O

2 at higher than normal pressure and a spark led to a fire and the loss of the Apollo 1 crew.

Highly concentrated sources of oxygen promote rapid combustion. Fire and explosion hazards exist when concentrated oxidants and fuels are brought into close proximity; an ignition event, such as heat or a spark, is needed to trigger combustion.^[28]^[117] Oxygen is the oxidant, not the fuel, but nevertheless the source of most of the chemical energy released in combustion.^[28]^[35] Combustion hazards also apply to compounds of oxygen with a high oxidative potential, such as peroxides, chlorates, nitrates, perchlorates, and dichromates because they can donate oxygen to a fire.

Concentrated O
2 will allow combustion to proceed rapidly and energetically.^[117] Steel pipes and storage vessels used to store and transmit both gaseous and liquid oxygen will act as a fuel; and therefore the design and manufacture of O
2 systems requires special training to ensure that ignition sources are minimized.^[117] The fire that killed the Apollo 1 crew in a launch pad test spread so rapidly because the capsule was pressurized with pure O
2 but at slightly more than atmospheric pressure, instead of the ¹⁄₃ normal pressure that would be used in a mission.^[k]^[119]

Liquid oxygen spills, if allowed to soak into organic matter, such as wood, petrochemicals, and asphalt can cause these materials to detonate unpredictably on subsequent mechanical impact.^[117] As with other cryogenic liquids, on contact with the human body it can cause frostbites to the skin and the eyes.

Notes

^ These results were mostly ignored until 1860. Part of this rejection was due to the belief that atoms of one element would have no chemical affinity towards atoms of the same element, and part was due to apparent exceptions to Avogadro's law that were not explained until later in terms of dissociating molecules.
^ An orbital is a concept from quantum mechanics that models an electron as a wave-like particle that has a spatial distribution about an atom or molecule.
^ Oxygen's paramagnetism can be used analytically in paramagnetic oxygen gas analysers that determine the purity of gaseous oxygen. ("Company literature of Oxygen analyzers (triplet)". Servomex. Retrieved December 15, 2007. )
^ Figures given are for values up to 50 miles (80 km) above the surface
^ Thylakoid membranes are part of chloroplasts in algae and plants while they simply are one of many membrane structures in cyanobacteria. In fact, chloroplasts are thought to have evolved from cyanobacteria that were once symbiotic partners with the progenerators of plants and algae.
^ Water oxidation is catalyzed by a manganese-containing enzyme complex known as the oxygen evolving complex (OEC) or water-splitting complex found associated with the lumenal side of thylakoid membranes. Manganese is an important cofactor, and calcium and chloride are also required for the reaction to occur. (Raven 2005)
^ (1.8 grams/min/person)×(60 min/h)×(24 h/day)×(365 days/year)×(6.6 billion people)/1,000,000 g/t=6.24 billion tonnes
^ The reason is that increasing the proportion of oxygen in the breathing gas at low pressure acts to augment the inspired O
2 partial pressure nearer to that found at sea-level.
^ Also, since oxygen has a higher electronegativity than hydrogen, the charge difference makes it a polar molecule. The interactions between the different dipoles of each molecule cause a net attraction force.
^ Since O
2's partial pressure is the fraction of O
2 times the total pressure, elevated partial pressures can occur either from high O
2 fraction in breathing gas or from high breathing gas pressure, or a combination of both.
^ No single ignition source of the fire was conclusively identified, although some evidence points to arc from an electrical spark.^[118]

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References

Cook, Gerhard A.; Lauer, Carol M. (1968). "Oxygen". In Clifford A. Hampel. The Encyclopedia of the Chemical Elements. New York: Reinhold Book Corporation. pp. 499–512. LCCN 68-29938.
Emsley, John (2001). "Oxygen". Nature's Building Blocks: An A-Z Guide to the Elements. Oxford, England: Oxford University Press. pp. 297–304. ISBN 0-19-850340-7.
Raven, Peter H.; Evert, Ray F.; Eichhorn, Susan E. (2005). Biology of Plants (7th ed.). New York: W.H. Freeman and Company Publishers. pp. 115–27. ISBN 0-7167-1007-2.

External links

Oxygen at The Periodic Table of Videos (University of Nottingham)
Oxidizing Agents > Oxygen
Oxygen (O₂) Properties, Uses, Applications
Roald Hoffmann article on "The Story of O"
WebElements.com – Oxygen
Chemistry in its element podcast (MP3) from the Royal Society of Chemistry's Chemistry World: Oxygen
Oxygen on In Our Time at the BBC. (listen now)
Scripps Institute: Atmospheric Oxygen has been dropping for 20 years

UpToDate Contents

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1. 酸素供給と消費 oxygen delivery and consumption
2. 長期酸素補充療法 long term supplemental oxygen therapy
3. 低酸素血症の酸素供給およびメカニズム oxygenation and mechanisms of hypoxemia
4. 酸素毒性 oxygen toxicity
5. 携帯用の酸素投与機器および酸素節約機器 portable oxygen delivery and oxygen conserving devices

English Journal

An oxygen-independent and membrane-less glucose biobattery/supercapacitor hybrid device.

Xiao X1, Conghaile PÓ2, Leech D2, Ludwig R3, Magner E4.
Biosensors & bioelectronics.Biosens Bioelectron.2017 Dec 15;98:421-427. doi: 10.1016/j.bios.2017.07.023. Epub 2017 Jul 11.
PMID 28711029

Fluorescence-based CdTe nanosensor for sensitive detection of cytochrome C.

Amin RM1, Elfeky SA2, Verwanger T3, Krammer B3.
Biosensors & bioelectronics.Biosens Bioelectron.2017 Dec 15;98:415-420. doi: 10.1016/j.bios.2017.07.020. Epub 2017 Jul 8.
PMID 28711028

Biomass-derived heteroatoms-doped mesoporous carbon for efficient oxygen reduction in microbial fuel cells.

Lu Y1, Zhu N2, Yin F1, Yang T1, Wu P3, Dang Z4, Liu M5, Wei X1.
Biosensors & bioelectronics.Biosens Bioelectron.2017 Dec 15;98:350-356. doi: 10.1016/j.bios.2017.07.006. Epub 2017 Jul 5.
PMID 28704783

Zinc(II) phthalocyanine fused in peripheral positions octa-substituted with alkyl linked carbazole: Synthesis, electropolymerization and its electro-optic and biosensor applications.

Olgac R1, Soganci T1, Baygu Y1, Gök Y2, Ak M3.
Biosensors & bioelectronics.Biosens Bioelectron.2017 Dec 15;98:202-209. doi: 10.1016/j.bios.2017.06.028. Epub 2017 Jun 19.
PMID 28683412

Japanese Journal

Experimental study on the unique stability mechanism via miniaturization of jet diffusion flames (microflame) by utilizing preheated air system

Fujiwara Kakeru,Nakamura Yuji
Combustion And Flame 160(8), 1373-1380, 2013-08-00
… It is also suggested that the endothermic radical-chain reactions are promoted near the exit of the burner when the burner temperature is substantially heated, at which the back-diffused oxygen is penetrated. …
NAID 120005307931

α-Lipoic acid-induced inhibition of proliferation and met phosphorylation in human non-small cell lung cancer cells

Michikoshi Hiromitsu,Nakamura Takahiro,Sakai Katsuya,Suzuki Yoshinori,Adachi Eri,Matsugo Seiichi,Matsumoto Kunio
Cancer Letters 335(2), 472-478, 2013-07-28
… The enantiomers R-α-LA and S-α-LA suppressed cell proliferation and increased the level of reactive oxygen species in HCC-827 and PC-9 human non-small cell lung cancer cells in an indistinguishable dose-dependent fashion. …
NAID 120005254051

ポリ（１-トリメチルシリル-１-プロピン）の気体透過性の経時変化に関する研究

増田俊夫,Masuda Toshio,岡本皇馬,Okamoto Kouma,曽根健史,Sone Takehumi,阪口壽一,Sakaguchi Toshikazu,橋本保,Hashimoto Tamotsu,天谷直之,Amaya Naoyuki,実生史朗,Mibae Shiro
福井工業大学研究紀要 43, 293-302, 2013-07-18
… Its oxygen permeability coefficient (Po2) reaches up to about 10 000 barrers. …
NAID 120005301144

Atomic level observation of octahedral distortions at the perovskite oxide heterointerface.

Aso Ryotaro,Kan Daisuke,Shimakawa Yuichi,Kurata Hiroki
Scientific reports 3, 2013-07-16
… While perovskite oxide heterostructures offer a good platform for controlling functionalities, atomistic understanding of octahedral distortion at the interface has been a challenge as it requires precise measurements of the oxygen atomic positions. … Here we demonstrate an approach to clarify distortions at an atomic level using annular bright-field imaging in aberration-corrected scanning transmission electron microscopy, which provides precise mappings of cation and oxygen atomic positions from distortion-minimized images. …
NAID 120005312067

「酸素」

Oxygen, ₈O
	Spectral lines of oxygen
General properties
Name, symbol	oxygen, O
Pronunciation	/ˈɒksᵻdʒən/; OK-si-jən
Appearance	gas: colorless; liquid: pale blue
Oxygen in the periodic table
	nitrogen ← oxygen → fluorine
Atomic number (Z)	8
Group, block	group 16 (chalcogens), p-block
Period	period 2
Element category	diatomic nonmetal
Standard atomic weight (Ar)	15.999 (15.99903–15.99977)
Electron configuration	[He] 2s2 2p4
per shell	2, 6
Physical properties
Phase	gas
Melting point	54.36 K (−218.79 °C, −361.82 °F)
Boiling point	90.188 K (−182.962 °C, −297.332 °F)
Density at stp (0 °C and 101.325 kPa)	1.429 g/L
when liquid, at b.p.	1.141 g/cm3
Triple point	54.361 K, 0.1463 kPa
Critical point	154.581 K, 5.043 MPa
Heat of fusion	(O2) 0.444 kJ/mol
Heat of vaporization	(O2) 6.82 kJ/mol
Molar heat capacity	(O2) 29.378 J/(mol·K)
Atomic properties
Oxidation states	2, 1, −1, −2
Electronegativity	Pauling scale: 3.44
Ionization energies	1st: 1313.9 kJ/mol; 2nd: 3388.3 kJ/mol; 3rd: 5300.5 kJ/mol; (more)
Covalent radius	66±2 pm
Van der Waals radius	152 pm
Miscellanea
Crystal structure	cubic
Speed of sound	330 m/s (gas, at 27 °C)
Thermal conductivity	26.58×10−3 W/(m·K)
Magnetic ordering	paramagnetic
CAS Number	7782-44-7
History
Discovery	Carl Wilhelm Scheele (1772)
Named by	Antoine Lavoisier (1777)
Most stable isotopes of oxygen
	view; talk; edit; \| references

外見
	無色の気体（液体は淡青色）; 沸騰している液体酸素(酸素の沸点は1 atmで約−183 °C (−297 °F))。; ; 酸素のスペクトル線
一般特性
名称, 記号, 番号	酸素, O, 8
分類	非金属, カルコゲン
族, 周期, ブロック	16, 2, p
原子量	15.9994(3)
電子配置	1s2 2s2 2p4
電子殻	2, 6（画像）
物理特性
色	無色
相	気体
密度	(0 °C, 101.325 kPa); 1.429 g/L
融点	54.8 K, -218.4 °C, -361.82 °F
沸点	90.2 K, −182.96 °C, −297.31 °F
臨界点	154.59 K, 5.043 MPa
融解熱	(O2) 0.444 kJ/mol
蒸発熱	(O2) 6.82 kJ/mol
熱容量	(25 °C) (O2) 29.378 J/(mol·K)
蒸気圧
原子特性
酸化数	2, 1, −1, −2
電気陰性度	3.44（ポーリングの値）
イオン化エネルギー; （詳細）	第1： 1313.9 kJ/mol
	第2： 3388.3 kJ/mol
	第3： 5300.5 kJ/mol
共有結合半径	66 ± 2 pm
ファンデルワールス半径	152 pm
その他
結晶構造	立方晶系
磁性	反磁性
熱伝導率	(300 K) 26.58 × 10−3 W/(m·K)
音の伝わる速さ	（気体、27 °C）330 m/s
CAS登録番号	7782-44-7
最安定同位体
	詳細は酸素の同位体を参照
	表示

　　[★]

英: oxygen
関: 空気、酸素療法

大気中の乾燥空気に含まれる酸素 20.946%

http://ja.wikipedia.org/wiki/%E7%A9%BA%E6%B0%97

麻酔科

参考1

	酸素	笑気	空気	二酸化炭素
医療ガス配管	緑	青	黄	橙
ガスボンベの色(日本)				緑
ピンインデックス	・　　　・	・　　　・	・　・　　・	・　・
ピン	2	2	3	2
角度(時計回り)	180	135	120	-90

参考

1.

http://www.eonet.ne.jp/~hidarite/me2/anzenkanri05.html

「dioxygen」

　　[★]

二原子酸素、酸素分子、二酸素

関: molecular oxygen、oxygen

「oxygenated」

　　[★]

酸素を含ませた、酸素化した

関: oxygenate、oxygenation

「dioxygenase」

　　[★]

ジオキシゲナーゼ、二原子酸素添加酵素

「protocatechuate dioxygenase」

　　[★]

プロトカテキュ酸ジオキシゲナーゼ

「phosphorus-oxygen lyase」

　　[★]

リン・酸素リアーゼ

[30] 原初の地球大気にも、水蒸気が光分解されて発生するメカニズムが指摘されており、ごく微量ながら酸素ガスが存在した可能性はあるが、ほとんどはすぐ酸化反応で消費されるか、オゾンへ変化したものと思われ、いずれにしろ考慮に足る量ではなかった^[29]。

[45] (0.36 g/分/人) × (60秒/時) × (24時/日) × (365日/年) × (70億人)/1,000,000 = 13.2億トン

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[Chikyu-28] ^ ^a ^b ^c ^d ^e ^f A・オレイニコフ「I. 地球進化を探る 3.地球大気の形成」『地球時計』講談社、1978年、第一刷、21–35頁。

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[mito19-21-31] 瀬名、太田 (2000)、19–21頁。

[mito31-35-32] 瀬名、太田 (2000)、31–35頁。

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[Raven-36] Raven (2005), pp. 115–127.

[37] Raven 2005

[sakurai68-38] 桜井 (1997)、68頁。

[GuideElem48-39] Stwertka, Albert (1998). Guide to the Elements (revised ed.). Oxford University Press. pp. 48–49. ISBN 0-19-508083-1.

[NBB298-40] Emsley (2001), p. 298.

[mito71-74-41] 瀬名らp.71-74 2.危険なエネルギープラントミトコンドリアの姿と働き酸素と活性酸素

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[44] “東邦大学医学部「呼吸器系」ユニット講義録８”. 2012年7月1日閲覧。

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[86] 大気中酸素濃度の減少量から二酸化炭素の陸域生物圏吸収量の推定に成功（国立環境研究所）

[19] These results were mostly ignored until 1860. Part of this rejection was due to the belief that atoms of one element would have no chemical affinity towards atoms of the same element, and part was due to apparent exceptions to Avogadro's law that were not explained until later in terms of dissociating molecules.

[29] An orbital is a concept from quantum mechanics that models an electron as a wave-like particle that has a spatial distribution about an atom or molecule.

[32] Oxygen's paramagnetism can be used analytically in paramagnetic oxygen gas analysers that determine the purity of gaseous oxygen. ("Company literature of Oxygen analyzers (triplet)". Servomex. Retrieved December 15, 2007. )

[60] Figures given are for values up to 50 miles (80 km) above the surface

[69] Thylakoid membranes are part of chloroplasts in algae and plants while they simply are one of many membrane structures in cyanobacteria. In fact, chloroplasts are thought to have evolved from cyanobacteria that were once symbiotic partners with the progenerators of plants and algae.

[71] Water oxidation is catalyzed by a manganese-containing enzyme complex known as the oxygen evolving complex (OEC) or water-splitting complex found associated with the lumenal side of thylakoid membranes. Manganese is an important cofactor, and calcium and chloride are also required for the reaction to occur. (Raven 2005)

[77] (1.8 grams/min/person)×(60 min/h)×(24 h/day)×(365 days/year)×(6.6 billion people)/1,000,000 g/t=6.24 billion tonnes

[105] The reason is that increasing the proportion of oxygen in the breathing gas at low pressure acts to augment the inspired O
2 partial pressure nearer to that found at sea-level.

[113] Also, since oxygen has a higher electronegativity than hydrogen, the charge difference makes it a polar molecule. The interactions between the different dipoles of each molecule cause a net attraction force.

[120] Since O
2's partial pressure is the fraction of O
2 times the total pressure, elevated partial pressures can occur either from high O
2 fraction in breathing gas or from high breathing gas pressure, or a combination of both.

[129] No single ignition source of the fire was conclusively identified, although some evidence points to arc from an electrical spark.^[118]

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[mmHg-78] Derived from mmHg values using 0.133322 kPa/mmHg

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リンク元	「酸素」「dioxygen」
拡張検索	「oxygenated」「dioxygenase」「protocatechuate dioxygenase」「phosphorus-oxygen lyase」

v t e Diatomic chemical elements

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Contents

History

Early experiments

Phlogiston theory

Discovery

Lavoisier's contribution

Later history

Characteristics

Properties and molecular structure

Allotropes

Physical properties

Isotopes and stellar origin

Occurrence

Analysis

Biological role of O2

Photosynthesis and respiration

Living organisms

Build-up in the atmosphere

Industrial production

Storage

Applications

Medical

Life support and recreational use

Industrial

Compounds

Oxides and other inorganic compounds

Organic compounds and biomolecules

Safety and precautions

Toxicity

Combustion and other hazards

See also

Notes

Citations

References

External links

UpToDate Contents

English Journal

Japanese Journal

Related Links

Biological role of O₂