Sulfur and Its Oxides | 硫及其氧化物

📚 Sulfur and Its Oxides | 硫及其氧化物

Sulfur is a non-metallic element found in Group 16 of the periodic table. It displays a wide range of oxidation states and forms several important oxides, most notably sulfur dioxide (SO₂) and sulfur trioxide (SO₃). These compounds play vital roles in industrial chemistry, biological systems, and environmental science. Understanding the structure, bonding, and reactivity of sulfur and its oxides is essential for A-Level Chemistry, particularly within the Cambridge curriculum's treatment of Period 3 elements, redox chemistry, and atmospheric pollution.

硫是元素周期表第16族的非金属元素。它表现出广泛的氧化态，并形成多种重要的氧化物，最典型的是二氧化硫（SO₂）和三氧化硫（SO₃）。这些化合物在工业化学、生物体系和环境科学中发挥着关键作用。理解硫及其氧化物的结构、键合与反应活性，对于A-Level化学至关重要，尤其是在剑桥课程对第三周期元素、氧化还原化学和大气污染的处理中。

1. Occurrence and Extraction of Sulfur | 硫的存在与提取

Sulfur occurs naturally in elemental form near volcanic regions and in underground deposits. It is also widely found in metal sulfide ores such as pyrite (FeS₂), sphalerite (ZnS), and galena (PbS). In addition, sulfur is present in fossil fuels and in biological molecules like amino acids. The Frasch process, which uses superheated water to melt underground sulfur and force it to the surface with compressed air, was historically important for extracting elemental sulfur. Today, much of the sulfur produced industrially comes from the desulfurisation of petroleum and natural gas.

硫以单质形式天然存在于火山地区和地下矿床中。它也广泛存在于金属硫化物矿石中，如黄铁矿（FeS₂）、闪锌矿（ZnS）和方铅矿（PbS）。此外，硫还存在于化石燃料以及氨基酸等生物分子中。弗拉希法曾是利用过热水熔化地下硫，并用压缩空气将其压出地表的历史性提取方法。如今，工业生产的大部分硫来自石油和天然气的脱硫过程。

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2. Allotropes of Sulfur | 硫的同素异形体

The most stable allotrope at room temperature is rhombic α-sulfur, which consists of puckered S₈ rings. When heated above about 96 °C, rhombic sulfur slowly converts to monoclinic β-sulfur, another crystalline form also built from S₈ rings but with a different packing arrangement. If molten sulfur is heated further and then poured into cold water, plastic (amorphous) sulfur is produced. Plastic sulfur contains long chains of sulfur atoms and is rubbery and elastic, but it gradually reverts to the rhombic form over time. Understanding these allotropes highlights the ability of sulfur to catenate – to form chains and rings of atoms.

室温下最稳定的同素异形体是斜方α-硫，它由皱折状的S₈环组成。加热到约96°C以上时，斜方硫缓慢转变为单斜β-硫，这是另一种晶形，同样由S₈环构成，但堆积方式不同。若将熔融硫进一步加热后倒入冷水中，可制得塑性（无定形）硫。塑性硫含有长链硫原子，具有橡胶弹性和弹性，但会逐渐恢复为斜方硫。认识这些同素异形体体现了硫的成链能力——即原子形成链和环的能力。

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3. Physical Properties of Sulfur | 硫的物理性质

Sulfur is a yellow, brittle solid at room temperature. It is insoluble in water but dissolves in non-polar solvents such as carbon disulfide (CS₂) and to some extent in toluene. It is a poor conductor of heat and electricity. The melting point of rhombic sulfur is about 113 °C, and the boiling point of liquid sulfur is around 445 °C. When heated, sulfur undergoes a series of interesting physical changes: it melts to a pale yellow mobile liquid, then darkens and becomes highly viscous near 200 °C as the S₈ rings break open and form long chains. Above this temperature, the viscosity decreases again as the chains break into shorter fragments.

硫在室温下是黄色脆性固体。它不溶于水，但溶于二硫化碳（CS₂）等非极性溶剂，也可部分溶于甲苯。硫是热和电的不良导体。斜方硫的熔点约为113°C，液态硫的沸点约为445°C。加热时，硫会发生一系列有趣的物理变化：先熔为浅黄色可流动的液体，然后在接近200°C时颜色变深，粘度急剧增大，这是因为S₈环断裂形成了长链。高于此温度，长链断裂成短链碎片，粘度再次降低。

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4. Chemical Properties of Elemental Sulfur | 单质硫的化学性质

Sulfur reacts with many metals and non-metals on heating. It burns in air or oxygen with a blue flame to form sulfur dioxide:

S(s) + O₂(g) → SO₂(g)

. With many metals, such as iron and zinc, sulfur reacts when heated to form sulfides, e.g. Fe(s) + S(s) → FeS(s). Sulfur also reacts directly with hydrogen at elevated temperatures to give hydrogen sulfide, H₂S. As an oxidising agent, sulfur accepts electrons from metals; as a reducing agent, it donates electrons to oxygen. This dual redox behaviour arises because sulfur can adopt oxidation states ranging from −2 to +6.

硫在加热时能与许多金属和非金属反应。它在空气或氧气中燃烧，产生蓝色火焰，生成二氧化硫：

S(s) + O₂(g) → SO₂(g)

。与许多金属（如铁和锌）在加热时反应生成硫化物，例如 Fe(s) + S(s) → FeS(s)。硫在高温下还可与氢气直接反应生成硫化氢（H₂S）。作为氧化剂，硫从金属接受电子；作为还原剂，它向氧提供电子。这种双重的氧化还原行为源于硫的氧化态可在−2到+6之间变化。

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5. Preparation of Sulfur Dioxide | 二氧化硫的制备

In the laboratory, sulfur dioxide gas is conveniently prepared by the action of a dilute acid on a sulfite salt, typically sodium sulfite or sodium hydrogensulfite. For example:

Na₂SO₃(s) + 2HCl(aq) → 2NaCl(aq) + H₂O(l) + SO₂(g)

. The gas is collected by downward delivery because it is denser than air. Industrially, SO₂ is produced by burning sulfur in air or by roasting sulfide ores such as pyrite: 4FeS₂(s) + 11O₂(g) → 2Fe₂O₃(s) + 8SO₂(g). This large-scale production is the first step in the Contact Process for sulfuric acid manufacture.

在实验室中，二氧化硫气体通常通过稀酸与亚硫酸盐（如亚硫酸钠或亚硫酸氢钠）反应来制备。例如：

Na₂SO₃(s) + 2HCl(aq) → 2NaCl(aq) + H₂O(l) + SO₂(g)

。由于SO₂密度大于空气，采用向下排空气法收集。工业上，二氧化硫通过硫在空气中燃烧或焙烧黄铁矿等硫化物矿石来生产：4FeS₂(s) + 11O₂(g) → 2Fe₂O₃(s) + 8SO₂(g)。这一大规模生产是接触法制硫酸的第一步。

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6. Structure and Bonding of Sulfur Dioxide | 二氧化硫的结构与键合

is a lone pair on the sulfur, which repels the bonding pairs and reduces the bond angle from the ideal 120°. The S– e structure can be represented by two resonance hybrids with one S=O double bond and one S– le and explains its ability to act as both an electrophile and a nucleophile in reactions.

SO₂分子呈角形，键角约为119°。硫原子为sp²杂化，与两个氧原子形成两个σ键。硫上有一个孤电子对，它会排斥成键电子对，使键角从理想的120°减小。S–O键长介于单键和双键之间，表明存在共振：该结构可用两个共振杂化体表示，其中一个S=O双键和一个S→O配位键，两个键均具有部分双键特征。这种共振稳定了分子，也解释了其在反应中既可作亲电试剂又可作亲核试剂的能力。

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7. Sulfur Dioxide as an Acidic Oxide | 二氧化硫作为酸性氧化物

Sulfur dioxide dissolves readily in water to form a weakly acidic solution containing sulfurous acid, H₂SO₃:

SO₂(g) + H₂O(l) ⇌ H₂SO₃(aq)

. The acid is diprotic, dissociating in two steps, but it cannot be isolated in pure form because it readily decomposes back into SO₂ and water. The solution reacts with alkalis to produce sulfite salts. For example, with sodium hydroxide, the reaction depends on the molar ratio: with an excess of NaOH, the normal sulfite Na₂SO₃ is formed; with a limited amount of base, the hydrogensulfite NaHSO₃ is obtained. This is typical behaviour of an acidic oxide derived from a non-metal in oxidation state +4.

二氧化硫易溶于水，形成含有亚硫酸（H₂SO₃）的弱酸性溶液：

SO₂(g) + H₂O(l) ⇌ H₂SO₃(aq)

。该酸是二元酸，分两步解离，但无法以纯品形式分离，因为它容易分解回SO₂和水。该溶液与碱反应生成亚硫酸盐。例如，与氢氧化钠的反应取决于摩尔比：过量NaOH生成正盐亚硫酸钠Na₂SO₃，碱量有限时则得到亚硫酸氢钠NaHSO₃。这是非金属在+4氧化态下形成的酸性氧化物的典型行为。

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8. Redox Chemistry of Sulfur Dioxide: Reducing Agent | 二氧化硫的氧化还原化学：还原性

Sulfur dioxide is an important reducing agent because the sulfur atom in SO₂ is in the +4 oxidation state and can be oxidised to +6. A classic example is its reaction with acidified potassium dichromate(VI), where the orange dichromate ion Cr₂O₇²⁻ is reduced to green Cr³⁺ ions:

3SO₂(g) + Cr₂O₇²⁻(aq) + 2H⁺(aq) → 3SO₄²⁻(aq) + 2Cr³⁺(aq) + H₂O(l)

. It also reduces acidified potassium manganate(VII), decolourising the purple MnO₄⁻ solution, as Mn(VII) is reduced to Mn²⁺: 5SO₂ + 2MnO₄⁻ + 2H₂O → 5SO₄²⁻ + 2Mn²⁺ + 4H⁺. In aqueous solution, SO₂ or H₂SO₃ is used as a mild reducing agent in organic and inorganic syntheses.

二氧化硫是一种重要的还原剂，因为SO₂中的硫原子处于+4氧化态，能被氧化至+6。一个经典例子是其与酸化重铬酸钾(VI)的反应：橙色的重铬酸根离子Cr₂O₇²⁻被还原为绿色的Cr³⁺离子：

3SO₂(g) + Cr₂O₇²⁻(aq) + 2H⁺(aq) → 3SO₄²⁻(aq) + 2Cr³⁺(aq) + H₂O(l)

。它还能还原酸化高锰酸钾(VII)，使紫色的MnO₄⁻溶液褪色，Mn(VII)被还原为Mn²⁺：5SO₂ + 2MnO₄⁻ + 2H₂O → 5SO₄²⁻ + 2Mn²⁺ + 4H⁺。在水溶液中，SO₂或H₂SO₃被用作有机和无机合成中的温和还原剂。

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9. Sulfur Dioxide as an Oxidising Agent | 二氧化硫作为氧化剂

Although SO₂ is primarily known as a reducing agent, it can also act as an oxidising agent when reacting with stronger reductants. For instance, it oxidises hydrogen sulfide to sulfur in a reaction that produces water and a fine yellow precipitate of solid sulfur:

SO₂(g) + 2H₂S(g) → 3S(s) + 2H₂O(l)

. This reaction illustrates a redox process where sulfur in SO₂ (+4) is reduced to elemental sulfur (0), while sulfur in H₂S (−2) is oxidised to 0. This type of synproportionation reaction is used to remove toxic H₂S from industrial gas streams.

尽管SO₂主要被视作还原剂，但在与更强的还原剂反应时，它也能充当氧化剂。例如，它能将硫化氢氧化为硫，生成水和黄色细小的硫单质沉淀：

SO₂(g) + 2H₂S(g) → 3S(s) + 2H₂O(l)

。此反应展示了氧化还原过程：SO₂中的硫(+4)被还原为单质硫(0)，而H₂S中的硫(−2)被氧化至0。这种归中反应被用于脱除工业气流中有毒的H₂S。

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10. Bleaching Action and Uses of Sulfur Dioxide | 二氧化硫的漂白作用与用途

Sulfur dioxide bleaches by reduction, not by oxidation like chlorine. In the presence of moisture, SO₂ forms dithionite-type intermediates that reduce coloured organic molecules to colourless compounds. This bleaching is often reversible; exposure to air over time can re-oxidise the reduced forms and colour may return. For this reason, SO₂ is used to bleach wool, silk, and wood pulp for paper. It is also used as a food preservative (E220) because it inhibits microbial growth and enzymatic browning in dried fruits and wine. Additionally, SO₂ is a precursor in the manufacture of sulfuric acid and is utilised in refining petroleum and as a solvent in some organic reactions.

二氧化硫通过还原作用漂白，而非像氯气那样通过氧化。在湿气存在下，SO₂生成连二亚硫酸盐类中间体，将有色有机物还原为无色化合物。这种漂白通常是可逆的；长时间暴露在空气中，还原形态可能被重新氧化，颜色会恢复。因此，SO₂被用于漂白羊毛、丝绸和造纸木浆。它也用作食品防腐剂（E220），能抑制微生物生长和果干、葡萄酒中的酶促褐变。此外，SO₂是制造硫酸的前体，用于石油精炼，并在某些有机反应中作为溶剂。

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11. Sulfur Trioxide: Preparation and Properties | 三氧化硫的制备与性质

Sulfur trioxide is produced industrially by the catalytic oxidation of sulfur dioxide in the Contact Process:

2SO₂(g) + O₂(g) ⇌ 2SO₃(g) ΔH = −196 kJ mol⁻¹

. A vanadium(V) oxide (V₂O₅) catalyst is used at a temperature of around 450 °C at a pressure slightly above atmospheric. SO₃ is a strong acidic oxide that reacts vigorously with water to form sulfuric acid. However, in practice, it is absorbed into concentrated sulfuric acid to produce oleum (H₂S₂O₇), which is then diluted with water to give sulfuric acid of any desired concentration: SO₃ + H₂SO₄ → H₂S₂O₇ then H₂S₂O₇ + H₂O → 2H₂SO₄. This avoids the formation of a corrosive and hard-to-condense mist of sulfuric acid droplets.

三氧化硫在工业上通过接触法中二氧化硫的催化氧化来生产：

2SO₂(g) + O₂(g) ⇌ 2SO₃(g) ΔH = −196 kJ mol⁻¹

。使用五氧化二钒（V₂O₅）催化剂，温度约450°C，压力略高于大气压。SO₃是一种强酸性氧化物，遇水剧烈反应生成硫酸。但实际上，它被吸收到浓硫酸中生成发烟硫酸（H₂S₂O₇），再用水稀释得到任意浓度的硫酸：SO₃ + H₂SO₄ → H₂S₂O₇，然后H₂S₂O₇ + H₂O → 2H₂SO₄。这避免了形成具有腐蚀性且难以冷凝的硫酸酸雾。

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12. Environmental Impact: Acid Rain | 环境影响：酸雨

Sulfur dioxide is a primary contributor to acid rain. Coal and oil-burning power stations release SO₂ into the atmosphere, where it undergoes oxidation to SO₃ and then reacts with water to form sulfuric acid. Alternatively, SO₂ dissolves in water droplets to give sulfurous acid, which is slowly oxidised to sulfuric acid. Acid rain lowers the pH of soils and freshwater bodies, damaging forests, aquatic life, and limestone buildings. Flue gas desulfurisation (FGD) technologies, such as using limestone slurry to absorb SO₂ and produce calcium sulfite, help mitigate these emissions. International regulations have significantly reduced SO₂ emissions from industrial sources in many regions.

二氧化硫是导致酸雨的主要物质。燃煤和燃油发电站向大气排放SO₂，在空气中被氧化为SO₃，随后与水反应生成硫酸。或者，SO₂溶解在水滴中形成亚硫酸，再缓慢氧化为硫酸。酸雨降低了土壤和淡水水体的pH，危害森林、水生生物和石灰岩建筑。烟气脱硫技术，如使用石灰石浆液吸收SO₂并生成亚硫酸钙，有助于减少这类排放。国际法规已大幅削减了许多地区工业源的SO₂排放。

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