📚 Covalent Bonding: IB & OCR Chemistry Key Points | 共价键：IB 与 OCR 化学考点精讲
Understanding covalent bonding is fundamental to IB and OCR A-level Chemistry. This topic explains how atoms share electrons to form molecules with distinct shapes, polarities, and reactivities. This article systematically covers the key concepts, from Lewis structures and VSEPR theory to hybridisation and bond enthalpy, ensuring you master the essential ideas required for examinations.
理解共价键是 IB 与 OCR A-level 化学的基础。这一主题解释原子如何通过共享电子形成具有特定形状、极性与反应活性的分子。本文系统梳理了从路易斯结构、价层电子对互斥理论到杂化与键焓的核心概念，帮助你掌握考试所要求的重点内容。
1. What is a Covalent Bond? | 什么是共价键？
A covalent bond forms when two atoms share one or more pairs of electrons. The shared electrons are attracted to the nuclei of both atoms, creating a stable balance of electrostatic forces. This typically occurs between non-metal atoms with similar electronegativities.
当两个原子共用一对或多对电子时，便形成共价键。共用电子同时受到两个原子核的吸引，形成稳定的静电平衡。这通常发生在电负性相近的非金属原子之间。
The shared pair of electrons is often represented as a line between the atomic symbols, e.g., H–H. A single covalent bond contains one shared pair, a double bond two shared pairs (O=O), and a triple bond three shared pairs (N≡N).
共用电子对通常用原子符号之间的短线表示，例如 H–H。单键含有一对共用电子，双键含有两对共用电子（O=O），三键则含有三对共用电子（N≡N）。
The driving force for bond formation is the tendency of atoms to achieve the stable electronic configuration of a noble gas. For most main-group elements, this is an octet of electrons in the outermost shell, although there are notable exceptions such as hydrogen (duet) and elements like boron or sulfur that can have expanded octets.
键形成的驱动力在于原子倾向于达到稀有气体的稳定电子构型。对多数主族元素而言，这意味着最外层满足八电子（八隅律），但氢（双电子稳定）以及硼、硫等可拥有扩展八隅体的元素为例外。
2. Lewis Structures and the Octet Rule | 路易斯结构与八隅律
Lewis structures (electron-dot diagrams) display the valence electrons of atoms within a molecule. Dots represent unpaired electrons, and lines or pairs of dots symbolise bonding pairs. Lone pairs (non-bonding electrons) are also shown around each atom.
路易斯结构（电子点式图）展示分子中各个原子的价电子。点表示未成对电子，线或点对代表成键电子对。孤对电子（非键电子）也围绕每个原子画出。
To construct a Lewis structure, first count the total number of valence electrons from all atoms. Then arrange atoms so that the most electronegative element is usually in the centre (if it’s not hydrogen). Distribute electrons to satisfy octets, starting with outer atoms, and form multiple bonds if the central atom still lacks an octet.
构建路易斯结构时，首先计算所有原子的价电子总数。然后将原子排列，通常电负性最高的原子（若非氢）位于中心。分配电子以满足八隅律，从外围原子开始，若中心原子仍未满八电子则形成多重键。
Common pitfalls include forgetting to count charges for ions (add an electron for each negative charge, subtract for positive charges) and attempting to give elements of period 3 or below more than eight electrons only when necessary. The octet rule is a useful guide, but radical species like NO₂ or electron-deficient compounds like BF₃ are important exceptions.
常见错误包括忘记计算离子的电荷（每带一个负电荷增加一个电子，正电荷则减少），以及只有在必要时才让第三周期及以下的元素扩展八隅体。八隅律是有用的向导，但自由基如 NO₂ 或缺电子化合物如 BF₃ 是重要的例外。
3. Bond Polarity and Electronegativity | 键的极性与电负性
Electronegativity is the ability of an atom to attract the bonding electrons towards itself. In a covalent bond between two different atoms, the electron cloud is not shared evenly: it is displaced towards the more electronegative atom, creating a polar covalent bond. The greater the difference in electronegativity (ΔEN), the more polar the bond.
电负性指原子吸引成键电子的能力。在不同原子间的共价键中，电子云并非均匀共享：它会偏向电负性更强的原子，形成极性共价键。电负性差值（ΔEN）越大，键的极性越强。
A bond with ΔEN less than about 0.4 is generally considered non-polar covalent, between 0.4 and 1.7 polar covalent, and above 1.7 primarily ionic. In IB and OCR chemistry, this boundary is often discussed with the Pauling scale. For example, the O–H bond has a significant dipole moment due to the high electronegativity of oxygen (3.44) compared to hydrogen (2.20).
通常 ΔEN 小于 0.4 为非极性共价键，0.4 至 1.7 之间为极性共价键，大于 1.7 则主要认为是离子键。在 IB 与 OCR 化学中，常以鲍林标度讨论此界限。例如 O–H 键因氧的电负性（3.44）远大于氢（2.20）而产生显著的偶极矩。
In molecules with polar bonds, the overall molecular polarity depends on geometry. Symmetrical molecules like CO₂ or CCl₄ are non-polar despite having polar bonds because the bond dipoles cancel. Asymmetrical molecules such as H₂O or NH₃ are polar, leading to higher boiling points and solubility in polar solvents.
在含有极性键的分子中，整体分子极性取决于几何构型。对称分子如 CO₂ 或 CCl₄ 虽存在极性键，但由于键矩相互抵消而表现为非极性。水（H₂O）或氨（NH₃）等不对称分子为极性分子，因此具有更高的沸点，并易溶于极性溶剂。
4. Coordinate (Dative) Covalent Bonds | 配位（给予）共价键
A coordinate or dative covalent bond forms when both electrons in the shared pair originate from the same atom. The donating atom must have a lone pair of electrons, and the accepting atom must have an empty orbital or a deficiency of electrons.
当共用电子对的两个电子均来自同一原子时，即形成配位共价键（又称给予键）。提供电子的原子必须具有孤对电子，接受电子的原子则须有空轨道或缺电子。
Classic examples include the ammonium ion NH₄⁺, where a lone pair on nitrogen in NH₃ is donated to a H⁺ ion, and the hydronium ion H₃O⁺. In transition metal complexes, ligands such as H₂O: or :NH₃ donate lone pairs to the central metal ion, forming coordinate bonds that are otherwise identical in length and strength to classical covalent bonds.
经典例子包括铵根离子 NH₄⁺，氨分子中氮上的孤对电子提供给 H⁺；以及水合氢离子 H₃O⁺。在过渡金属配合物中，配体如 H₂O: 或 :NH₃ 将孤对电子提供给中心金属离子，形成的配位键在长度和强度上与传统共价键并无区别。
5. Shapes of Molecules: VSEPR Theory | 分子形状：VSEPR 理论
Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the three-dimensional shape of a molecule by assuming that electron pairs (bonding and non-bonding) around a central atom repel each other and arrange themselves as far apart as possible.
价层电子对互斥理论（VSEPR）通过假定中心原子周围的电子对（成键与孤对）相互排斥并尽可能远离，来预测分子的三维形状。
The shape depends on the number of electron domains (bonding regions and lone pairs) around the central atom. Common shapes with their ideal bond angles are shown in the table below.
分子形状取决于中心原子周围的电子域（成键区域与孤对电子）数目。常见形状及其理想键角见下表。
| Electron domains | Bonding pairs | Lone pairs | Shape | Example | Bond angle |
|---|---|---|---|---|---|
| 2 | 2 | 0 | Linear | CO₂ | 180° |
| 3 | 3 | 0 | Trigonal planar | BF₃ | 120° |
| 3 | 2 | 1 | Bent (V-shaped) | SO₂ | < 120° |
| 4 | 4 | 0 | Tetrahedral | CH₄ | 109.5° |
| 4 | 3 | 1 | Trigonal pyramidal | NH₃ | 107° |
| 4 | 2 | 2 | Bent (V-shaped) | H₂O | 104.5° |
| 5 | 5 | 0 | Trigonal bipyramidal | PCl₅ | 90°, 120° |
| 6 | 6 | 0 | Octahedral | SF₆ | 90° |
6. Bond Angle Variations | 键角的变化
Lone pairs occupy more space than bonding pairs because they are attracted to only one nucleus. Consequently, lone pair–bonding pair repulsion > bonding pair–bonding pair repulsion. This causes a reduction in bond angles from the ideal tetrahedral angle when lone pairs are present.
孤对电子由于只受单个核的吸引，占据的空间比成键电子对更大。因此，孤对-成键电子对排斥力 > 成键-成键电子对排斥力。当存在孤对电子时，键角会从理想四面体角收缩。
For instance, methane (CH₄) has no lone pairs and a bond angle of 109.5°. Ammonia (NH₃) has one lone pair, compressing the H–N–H angle to about 107°. Water (H₂O) with two lone pairs reduces the H–O–H angle further to 104.5°. Understanding these deviations is crucial for predicting molecular shapes correctly in exams.
例如，甲烷（CH₄）没有孤对电子，键角为 109.5°。氨（NH₃）有一对孤对电子，使 H–N–H 角压缩至约 107°。水（H₂O）拥有两对孤对电子，进一步将 H–O–H 角减小至 104.5°。理解这些偏差对在考试中正确预测分子形状至关重要。
7. Hybridisation: sp, sp², sp³ | 杂化：sp、sp²、sp³
To account for the observed geometries of molecules, the model of atomic orbital hybridisation was developed. It involves the mixing of atomic orbitals to form new hybrid orbitals that have the same energy and are oriented in space to minimise repulsion. IB and OCR syllabi typically cover sp, sp², and sp³ hybridisation for carbon and other elements.
为了解释分子实际呈现的几何构型，人们提出了原子轨道杂化模型。该模型将原子轨道混合，形成能量相同、空间取向能最大程度减小排斥的新杂化轨道。IB 和 OCR 大纲通常涵盖碳和其他元素的 sp、sp²、sp³ 杂化。
- sp³ hybridisation: One s orbital and three p orbitals mix to form four equivalent sp³ hybrid orbitals pointing towards the corners of a tetrahedron (109.5° apart). Example: methane, ethane.
- sp³ 杂化： 一个 s 轨道与三个 p 轨道混合，形成四个等价的 sp³ 杂化轨道，指向四面体的四个顶点（相隔 109.5°）。示例：甲烷、乙烷。
- sp² hybridisation: One s and two p orbitals mix, leaving one unhybridised p orbital. The three sp² orbitals lie in a plane at 120°. Example: ethene (C₂H₄), where the leftover p orbital forms a π bond.
- sp² 杂化： 一个 s 与两个 p 轨道混合，剩余一个未杂化的 p 轨道。三个 sp² 轨道位于同一平面，夹角 120°。示例：乙烯（C₂H₄），剩余的 p 轨道形成 π 键。
- sp hybridisation: One s and one p orbital mix, producing two linear sp orbitals at 180°, with two unhybridised p orbitals. Example: ethyne (C₂H₂).
- sp 杂化： 一个 s 与一个 p 轨道混合，产生两个直线型 sp 轨道，夹角 180°，另有两个未杂化的 p 轨道。示例：乙炔（C₂H₂）。
Hybridisation provides a clear link between electron-domain geometry and the type of bonding. It helps explain why carbon can form four bonds in methane yet only three sigma bonds in ethene, despite having only two unpaired electrons in its ground state configuration.
杂化为电子域几何与成键类型之间提供了清晰的纽带，解释了为何碳在基态仅有两个未成对电子，却能在甲烷中形成四个键，在乙烯中形成三个 σ 键。
8. Sigma and Pi Bonds | σ 键与 π 键
A covalent bond can be described in terms of orbital overlap. A sigma (σ) bond results from the end-on overlap of orbitals, allowing free rotation. A pi (π) bond arises from the sideways overlap of adjacent p orbitals and restricts rotation. All single bonds are sigma bonds; a double bond consists of one sigma and one pi bond; a triple bond includes one sigma and two pi bonds.
共价键可用轨道重叠来描述。σ 键由轨道“头对头”重叠形成，可自由旋转。π 键则由相邻 p 轨道侧向重叠产生，限制旋转。所有单键都是 σ 键；双键含一个 σ 和一个 π 键；三键含一个 σ 和两个 π 键。
In ethene, the sp²-sp² overlap forms a C–C σ bond, while the unhybridised p orbitals overlap sideways to create a π bond. This π bond is weaker than the σ bond and is the site of electrophilic attack. Understanding the distinction is often tested in questions about bond strength, reactivity, and isomerism.
在乙烯中，sp²-sp² 重叠形成 C–C σ 键，而未杂化的 p 轨道侧向重叠产生 π 键。该 π 键比 σ 键弱，是亲电进攻的位点。区分二者常出现在关于键强度、反应活性与异构现象的考题中。
9. Bond Enthalpy and Bond Length | 键焓与键长
Bond enthalpy is the average energy required to break one mole of a specific bond in the gaseous state. It is a measure of bond strength. Triple bonds have the highest bond enthalpy and the shortest bond length; single bonds have the lowest enthalpy and longest length.
键焓是气态下断裂一摩尔特定键所需的平均能量，是衡量键强弱的指标。三键的键焓最高、键长最短；单键的键焓最低、键长最长。
For carbon–carbon bonds, approximate values are:
C–C: 347 kJ mol⁻¹, 154 pm; C=C: 614 kJ mol⁻¹, 134 pm; C≡C: 839 kJ mol⁻¹, 120 pm
碳–碳键的近似值为：
C–C: 347 kJ mol⁻¹, 154 pm; C=C: 614 kJ mol⁻¹, 134 pm; C≡C: 839 kJ mol⁻¹, 120 pm
Breaking a molecule’s bonds requires energy (endothermic), while making bonds releases energy (exothermic). The enthalpy change of a reaction can be estimated using average bond enthalpies: ΔH = Σ(bonds broken) − Σ(bonds formed). This is a common calculation on both IB and OCR papers.
断裂分子的键需吸收能量（吸热），形成键则释放能量（放热）。可用平均键焓估算反应焓变：ΔH = Σ(断裂键的键焓) − Σ(形成键的键焓)。这是 IB 与 OCR 试卷中常见的计算。
10. Resonance Structures | 共振结构
When a molecule cannot be accurately represented by a single Lewis structure, resonance structures are used. These are two or more valid Lewis structures that differ only in the position of electrons (not atoms). The true structure is a resonance hybrid, a weighted average of the contributing forms, which results in delocalised electrons and increased stability.
当一个分子无法用单一路易斯结构准确描述时，便采用共振结构。它们是两个或多个仅在电子位置上有别的有效路易斯结构。真实结构是共振杂化体，是各贡献形式的加权平均，导致电子离域并增加稳定性。
Key examples include the benzene ring (benzene), where all C–C bonds are equivalent (bond order 1.5); the carbonate ion CO₃²⁻, with three equivalent C–O bonds; and the nitrate ion NO₃⁻. In benzene, delocalised π electrons explain its exceptional stability and its resistance to addition reactions.
关键实例包括苯环（苯），所有 C–C 键等价（键级 1.5）；碳酸根离子 CO₃²⁻，三个等价的 C–O 键；以及硝酸根离子 NO₃⁻。在苯中，离域 π 电子解释了其特殊的稳定性及对加成反应的惰性。
11. Intermolecular Forces Arising from Covalent Bonds | 共价键产生的分子间作用力
While covalent bonds hold atoms together within a molecule, intermolecular forces determine many physical properties. London dispersion forces (instantaneous dipole–induced dipole) exist between all molecules and increase with molecular size and surface area. Polar molecules also experience permanent dipole–dipole interactions.
共价键将分子内的原子结合在一起，而分子间作用力则决定了许多物理性质。色散力（瞬时偶极-诱导偶极）存在于所有分子间，并随分子大小和表面积增大而增强。极性分子间还存在永久偶极-偶极作用。
Hydrogen bonding, a particularly strong dipole–dipole interaction, occurs when hydrogen is covalently bonded to a highly electronegative atom (N, O, or F) and is attracted to a lone pair on another N, O, or F atom. This explains the anomalously high boiling points of water, ammonia, and hydrogen fluoride, as well as the structure of DNA and proteins.
氢键是一种特别强的偶极-偶极作用，当氢与高电负性原子（N、O 或 F）共价键合并被另一 N、O 或 F 上的孤对电子吸引时形成。它解释了水、氨、氟化氢反常的高沸点，以及 DNA 和蛋白质的结构。
12. Comparison with Ionic and Metallic Bonding | 共价键与离子键、金属键的对比
In the broader bonding triangle, covalent bonding sits between ionic and metallic bonding. Ionic bonding involves electron transfer and electrostatic attraction between oppositely charged ions, typically between a metal and a non-metal. Metallic bonding consists of a lattice of positive ions immersed in a ‘sea’ of delocalised electrons.
在更广阔的化学键三角中，共价键介于离子键与金属键之间。离子键涉及电子转移及阴阳离子间的静电吸引，通常发生在金属与非金属之间。金属键则是由正离子晶格沉浸于离域电子“海洋”中构成。
Many compounds exhibit intermediate bonding character. For example, lithium iodide (LiI) is predominantly ionic but has a degree of covalent character due to polarisation. Understanding these subtleties allows students to explain trends in melting points, solubility, and electrical conductivity. A comparison table can help:
许多化合物表现出中间键型特征。例如碘化锂（LiI）主要为离子性，但因极化作用也具有一定共价性。理解这些细微差别有助于解释熔点、溶解度和电导率的变化趋势。对比表可助理解：
| Property | Covalent simple molecular | Covalent giant (network) | Ionic | Metallic |
|---|---|---|---|---|
| Melting/boiling point | Low | Very high | High | High (variable) |
| Electrical conductivity | Non-conductor | Non-conductor (except graphite) | Conducts when molten/aqueous | Good conductor (solid/liquid) |
| Particles | Molecules | Atoms | Ions | Cations and delocalised electrons |
复习时需注意，二氧化硅和金刚石属于巨型共价（网状）结构，而冰或碘晶体为简单分子型共价物质。这两类性质截然不同，考试常要求解释差异。
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