Nitrogen and Sulfur | 氮和硫

📚 Nitrogen and Sulfur | 氮和硫

Nitrogen and sulfur are two essential non-metallic elements in the p‑block of the Periodic Table. Nitrogen, a colourless diatomic gas, makes up about 78% of the atmosphere and is vital for the synthesis of ammonia and nitric acid, while sulfur, a yellow solid found in volcanic deposits and fossil fuels, serves as the foundation for sulfuric acid, one of the most important industrial chemicals. A-level Chemistry explores their structures, bonding, manufacturing processes, characteristic reactions, and the environmental consequences of their oxides, particularly acid rain.

氮和硫是周期表中p区两种重要的非金属元素。氮是无色双原子气体，约占大气的78%，是合成氨和硝酸的基础；硫是一种黄色固体，存在于火山沉积物和化石燃料中，是硫酸这一最重要的工业化学品之一的原料。A-level化学深入探究它们的结构、键合、工业制备方法、特征反应及其氧化物（尤其是酸雨）的环境影响。

1. Position in the Periodic Table | 元素周期表中的位置

Nitrogen (atomic number 7) belongs to Group 15, with a valence electronic configuration of 2s²2p³. It forms three covalent bonds and can also expand its octet through coordinate bonding. Sulfur (atomic number 16) sits in Group 16, with a valence configuration of 3s²3p⁴. Its larger size allows for expansion beyond the octet, leading to compounds such as SF₆ and SO₄²⁻.

氮（原子序数7）位于第15族，价电子构型为2s²2p³。它通常形成三个共价键，也可通过配位键拓展八隅体。硫（原子序数16）位于第16族，价电子构型为3s²3p⁴，较大的原子尺寸使其能容纳超过八隅体的电子，形成如SF₆和SO₄²⁻等化合物。

Both elements exhibit a range of oxidation states: from −3 in NH₃ to +5 in NO₃⁻ for nitrogen, and from −2 in H₂S to +6 in SO₄²⁻ for sulfur. This versatile redox behaviour underpins much of their descriptive chemistry.

这两种元素都展现出多种氧化态：氮从NH₃中的−3到NO₃⁻中的+5，硫从H₂S中的−2到SO₄²⁻中的+6。这种多变的氧化还原行为是它们描述性化学的基石。

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2. The Unreactive Nature of Dinitrogen | 氮气的惰性本质

Dinitrogen, N₂, contains a triple bond (N≡N) with a bond enthalpy of +944 kJ mol⁻¹, one of the strongest known bonds. This high stability renders N₂ remarkably inert at room temperature, which is why it is used as a protective atmosphere in food packaging and chemical reactions.

氮气N₂含有一个三键（N≡N），键焓高达+944 kJ mol⁻¹，是已知最强的化学键之一。这一高稳定性使N₂在室温下极为惰性，因此常被用作食品包装和化学反应中的保护气氛。

To break the N≡N bond and make nitrogen reactive, extreme conditions or highly specialised catalysts are required. In nature, nitrogen fixation is carried out by lightning strikes and nitrogen‑fixing bacteria; industrially, the Haber process uses an iron catalyst at high pressure and temperature.

要断裂N≡N键并使氮气参与反应，需要极端条件或高度专一的催化剂。自然界中，固氮作用通过闪电和固氮菌实现；工业上，哈伯法则利用铁催化剂在高温高压下进行。

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3. The Haber Process for Ammonia Synthesis | 哈伯法制氨

Ammonia (NH₃) is manufactured by the Haber process, which combines dinitrogen from fractional distillation of liquid air and hydrogen derived mainly from methane (natural gas) by steam reforming. The reaction is reversible and exothermic:

氨（NH₃）通过哈伯法生产，氮气来源于液态空气的分馏，氢气主要来自甲烷（天然气）的蒸汽重整。该反应为可逆放热反应：

N₂(g) + 3H₂(g) ⇌ 2NH₃(g)    ΔH = −92 kJ mol⁻¹

Optimum industrial conditions are a pressure of 150–200 atm, a temperature of 400–450 °C, and a finely divided iron catalyst with promoters such as Al₂O₃ and K₂O. Although lower temperatures favour the forward equilibrium, too low a temperature slows the rate unacceptably; thus a compromise temperature is chosen to balance yield and kinetics.

工业上的最佳条件为：压力150–200 atm，温度400–450 °C，使用加入Al₂O₃和K₂O等助催化剂的细粉状铁催化剂。低温有利于正反应平衡，但温度过低会使反应速率过慢，因此选择折中温度以兼顾产率和动力学。

The ammonia is continuously removed by condensation, and unreacted gases are recycled. The Haber process remains one of the most important demonstrations of applying equilibrium and rate principles to maximise economic efficiency.

生成的氨通过冷凝不断分离，未反应的气体则循环使用。哈伯法至今仍是应用平衡与速率原理以实现经济效益最大化的最经典范例之一。

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4. Properties and Reactions of Ammonia | 氨的性质与反应

Ammonia acts as a weak Brønsted–Lowry base in water, establishing the dynamic equilibrium:

氨在水中表现为弱布朗斯特–劳里碱，建立以下动态平衡：

NH₃(aq) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq)

It neutralises acids to form ammonium salts, e.g. with hydrogen chloride gas:

它能与酸中和生成铵盐，例如与氯化氢气体反应：

NH₃(g) + HCl(g) → NH₄Cl(s)

This reaction produces thick white fumes of ammonium chloride, a classic test for ammonia and hydrogen chloride. Ammonia also reacts with concentrated sulfuric acid to form ammonium sulfate, a valuable fertiliser.

该反应产生浓厚的氯化铵白烟，是检验氨和氯化氢的经典方法。氨还与浓硫酸反应生成硫酸铵，这是一种重要的肥料。

Additionally, ammonia acts as a ligand, forming complex ions with transition metal cations. For instance, when aqueous ammonia is added to copper(II) sulfate solution, initially a pale‑blue precipitate of Cu(OH)₂ forms, which dissolves in excess ammonia to yield the deep‑blue tetraamminecopper(II) complex:

此外，氨可作为配体，与过渡金属阳离子形成配离子。例如，向硫酸铜溶液中加入氨水，先生成淡蓝色Cu(OH)₂沉淀，继续加入过量氨水则沉淀溶解，得到深蓝色的四氨合铜(II)配离子：

[Cu(H₂O)₆]²⁺ + 4NH₃ ⇌ [Cu(NH₃)₄(H₂O)₂]²⁺ + 4H₂O

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5. Thermal Decomposition of Ammonium Salts | 铵盐的热分解

Most ammonium salts decompose on heating, but the products depend on the anion. Salts with anions that are weak oxidising agents, such as ammonium chloride, sublime and dissociate:

大多数铵盐受热分解，但产物取决于阴离子的性质。当阴离子为弱氧化剂时，如氯化铵，会升华并解离：

NH₄Cl(s) ⇌ NH₃(g) + HCl(g)

For salts containing oxidising anions, a redox reaction occurs. Ammonium nitrate decomposes exothermically, initially to nitrous oxide, then at higher temperatures more vigorously:

若含有氧化性酸根，则发生氧化还原反应。硝酸铵受热分解，先产生一氧化二氮，在更高温度下分解更为剧烈：

NH₄NO₃(s) → N₂O(g) + 2H₂O(g)

Ammonium dichromate decomposes spectacularly ('volcano demonstration') to chromium(III) oxide, nitrogen and water. The explosive nature of ammonium nitrate when mixed with fuel or upon strong shock is also a crucial safety consideration.

重铬酸铵分解的'火山演示'反应生成氧化铬(III)、氮气和水，现象壮观。硝酸铵在与燃料混合或受到强烈撞击时的爆炸性亦是重要的安全考量。

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6. Oxides of Nitrogen: NO and NO₂ | 氮的氧化物：一氧化氮和二氧化氮

Nitric oxide (nitrogen monoxide), NO, is a colourless, neutral, odd‑electron molecule (free radical). It forms naturally during lightning strikes and in combustion engines. Direct synthesis in the laboratory involves the reduction of nitrite ions:

一氧化氮（NO）是一种无色、电中性、含奇电子的自由基分子。闪电和汽车引擎燃烧时自然产生。实验室中可通过亚硝酸根的还原直接制备：

3Cu(s) + 8H⁺(aq) + 2NO₃⁻(aq) → 3Cu²⁺(aq) + 2NO(g) + 4H₂O(l)

In air, NO is immediately oxidised to nitrogen dioxide, NO₂, a reddish‑brown, toxic gas that dimerises to colourless N₂O₄ at low temperatures:

空气中NO会迅速被氧化为二氧化氮NO₂，一种红棕色有毒气体，低温下二聚为无色的N₂O₄：

2NO(g) + O₂(g) → 2NO₂(g)

2NO₂(g) ⇌ N₂O₄(g)

Both NO and NO₂ are atmospheric pollutants (collectively NOₓ) that contribute to the formation of photochemical smog and acid rain.

NO和NO₂均为大气污染物（统称NOₓ），会促成光化学烟雾和酸雨的形成。

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7. Nitric Acid: Ostwald Process and Reactions | 硝酸：奥斯特瓦尔德法与反应

Nitric acid (HNO₃) is manufactured by the Ostwald process, which involves the catalytic oxidation of ammonia:

硝酸（HNO₃）通过奥斯特瓦尔德法生产，其核心是氨的催化氧化：

4NH₃(g) + 5O₂(g) → 4NO(g) + 6H₂O(g)   (Pt‑Rh catalyst, ~900 °C)

The NO is then further oxidised and absorbed in water to yield nitric acid. Concentrated nitric acid (~68%) is a strong acid and a powerful oxidising agent. Its reactions with metals vary with concentration:

生成的NO进一步氧化并被水吸收，制得硝酸。浓硝酸（约68%）既是强酸也是强氧化剂。它与金属的反应随浓度不同而有所差异：

Cu(s) + 4HNO₃(浓) → Cu(NO₃)₂(aq) + 2NO₂(g) + 2H₂O(l)

3Cu(s) + 8HNO₃(稀) → 3Cu(NO₃)₂(aq) + 2NO(g) + 4H₂O(l)

Nitric acid is essential in the production of fertilisers (ammonium nitrate), explosives (TNT, nitroglycerin) and polymers. Its oxidising power also means it must be handled with great care, as it can ignite organic matter.

硝酸广泛用于制造化肥（硝酸铵）、炸药（TNT、硝化甘油）和聚合物。其强氧化性也意味着它足以引燃有机物，必须谨慎操作。

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8. Allotropes of Sulfur: Rhombic and Monoclinic | 硫的同素异形体：斜方硫和单斜硫

Solid sulfur exists as several allotropes, the most important of which are rhombic (α‑sulfur) and monoclinic (β‑sulfur) sulfur. Both consist of crown‑shaped S₈ rings, but they differ in crystal packing. Rhombic sulfur is the thermodynamic stable form at room temperature, yellow and crystalline.

固体硫存在多种同素异形体，其中最重要的是斜方硫（α‑硫）和单斜硫（β‑硫）。两者均由皇冠状S₈环构成，但晶体堆积方式不同。室温下斜方硫是热力学稳定形貌，呈黄色结晶状。

Above 95.5 °C, rhombic slowly converts to monoclinic sulfur, which forms needle‑like crystals. This transition temperature is the point where the vapour pressures of the two allotropes are equal. Upon melting at about 119 °C, S₈ rings break open and polymerise, causing molten sulfur to darken and become viscous before solidifying to plastic sulfur on quenching.

温度升至95.5 °C以上时，斜方硫缓慢转化为针状的单斜硫，这一转变温度是两种同素异形体蒸气压相等的点。在约119 °C熔化时，S₈环断裂聚合，液态硫颜色变深、黏度剧增，骤然冷却后可得到弹性硫。

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9. Sulfur Dioxide: Preparation and Properties | 二氧化硫的制备与反应

Sulfur dioxide (SO₂) is a colourless, pungent gas. It is prepared in the laboratory by the action of dilute hydrochloric acid on sodium sulfite, or by the direct combustion of sulfur in air:

二氧化硫（SO₂）是一种无色、有刺激性气味的气体。实验室中可用稀盐酸与亚硫酸钠反应制备，也可通过硫在空气中燃烧得到：

S(s) + O₂(g) → SO₂(g)

SO₂ is highly soluble in water, forming an acidic solution of sulfurous acid (H₂SO₃), which accounts for its ability to act as a reducing agent and a bleach:

SO₂易溶于水，形成亚硫酸（H₂SO₃）的酸性溶液，这解释了其还原性和漂白作用：

SO₂(g) + H₂O(l) ⇌ H₂SO₃(aq)

In bleaching, SO₂ reduces coloured organic compounds to colourless species, but unlike chlorine, its action is reversible and does not destroy fibres. It also reduces Fe(III) to Fe(II), decolourises acidified potassium manganate(VII), and converts dichromate(VI) to green Cr(III).

在漂白过程中，SO₂将有色有机物还原为无色物质，但与氯不同，这种漂白是可逆的且不损伤纤维。它还能将Fe(III)还原为Fe(II)，使酸性高锰酸钾溶液褪色，并将重铬酸根还原为绿色的Cr(III)。

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10. Acid Rain: Formation and Effects | 酸雨：形成与影响

Acid rain is primarily caused by the emission of SO₂ and NOₓ. In the atmosphere, SO₂ is oxidised to SO₃, which reacts with water droplets to form sulfuric acid. Nitrogen dioxide reacts with water and oxygen to form nitric acid:

酸雨主要由SO₂和NOₓ的排放引起。大气中SO₂被氧化为SO₃，继而与水雾反应生成硫酸。二氧化氮则与水和氧气反应生成硝酸：

4NO₂(g) + O₂(g) + 2H₂O(l) → 4HNO₃(aq)

Rainwater with a pH below 5.6 is considered acid rain. The consequences are severe: leaching of toxic Al³⁺ ions from soil, damage to plant leaves, acidification of lakes and rivers, and corrosion of limestone buildings and marble statues.

pH 低于5.6的雨水即被视为酸雨。其后果十分严重：导致土壤中释放出有毒的Al³⁺离子、伤害植物叶片、酸化湖泊河流，以及对石灰岩建筑和大理石雕像的腐蚀。

Modern mitigation strategies include flue‑gas desulfurisation (using CaCO₃ or CaO to remove SO₂), catalytic converters in cars, and the use of low‑sulfur fuels. These measures are direct applications of the Chemistry of nitrogen and sulfur.

现代缓解措施包括烟道气脱硫（用CaCO₃或CaO去除SO₂）、汽车尾气催化转换器和使用低硫燃料，这些均为氮和硫化学的实际应用。

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11. Contact Process for Sulfuric Acid | 接触法制硫酸

Sulfuric acid is manufactured by the Contact process. The key step is the equilibrium‑limited oxidation of SO₂ to SO₃ using a vanadium(V) oxide (V₂O₅) catalyst:

硫酸通过接触法生产，核心步骤是SO₂在V₂O₅催化下氧化为SO₃的平衡反应：

2SO₂(g) + O₂(g) ⇌ 2SO₃(g)    ΔH = −197 kJ mol⁻¹

The reaction is carried out at about 450 °C and near atmospheric pressure. A lower temperature would increase the yield but make the rate unacceptably slow; the chosen temperature ensures both high activity of the catalyst and a reasonable equilibrium conversion. The V₂O₅ catalyst provides a low‑energy pathway via a redox cycle involving vanadium in +4 and +5 oxidation states.

反应通常在约450 °C和接近常压下进行。降低温度可提高产率，但速率过慢；所选温度既保证了催化剂的高活性，也维持了合理的平衡转化率。V₂O₅催化剂通过钒在+4和+5氧化态间的氧化还原循环提供了低能垒路径。

Sulfur trioxide is not directly absorbed in water because the reaction is violently exothermic and produces a fine mist of H₂SO₄ droplets. Instead, SO₃ is dissolved in concentrated sulfuric acid to form oleum (H₂S₂O₇), which is then carefully diluted with water:

SO₃不宜直接用水吸收，因为该反应剧烈放热并会形成大量酸雾。实际操作中先将SO₃溶于浓硫酸生成发烟硫酸（H₂S₂O₇），再用水稀释：

SO₃(g) + H₂SO₄(l) → H₂S₂O₇(l)

H₂S₂O₇(l) + H₂O(l) → 2H₂SO₄(l)

This two‑stage absorption ensures safe, high‑conversion production of 98% sulfuric acid.

这种两段吸收工艺确保了高转化率且安全地生产出98%硫酸。

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12. Properties and Uses of Concentrated Sulfuric Acid | 浓硫酸的性质与用途

Concentrated sulfuric acid exhibits three characteristic properties: it is a strong acid, a powerful oxidising agent,

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