📚 Mastering Covalent Bonding for WJEC A-Level Chemistry | A-Level WJEC 化学：共价键考点精讲
Covalent bonding is central to understanding molecular structure and reactivity in A-Level Chemistry. For WJEC students, this topic requires a firm grasp of electron sharing, molecular shape, bond polarity, and sigma/pi bonding. This article distils the essential points, providing clear explanations, diagrams, and examples aligned with the specification. You will learn to draw dot-and-cross diagrams, predict molecular geometry using VSEPR, and interpret bond properties confidently.
共价键是理解A-Level化学中分子结构和反应性的核心。对于WJEC学生而言，掌握电子共享、分子形状、键的极性以及σ/π键至关重要。本文提炼了必考知识要点，提供清晰的解释、图示和契合考纲的实例。你将学会绘制点交叉图、运用VSEPR预测分子几何结构，并自信地解读键的性质。
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 pair is attracted to the nuclei of both atoms, creating a strong electrostatic force that holds the atoms together. Covalent bonding typically occurs between non-metal atoms with similar electronegativities.
当两个原子共享一对或多对电子时，便形成共价键。共享电子对同时受到两个原子核的吸引，产生强大的静电引力将原子结合在一起。共价键通常发生在电负性相近的非金属原子之间。
The simplest example is the hydrogen molecule H₂. Each hydrogen atom contributes one electron, forming a single covalent bond (H−H). The pair of electrons occupies a molecular orbital that concentrates electron density between the two nuclei.
最简单的例子是氢分子 H₂。每个氢原子提供一个电子，形成一个单共价键 (H−H)。这对电子占据的分子轨道使电子密度集中在两个原子核之间。
In a structural formula, a covalent bond is represented by a single line (−) for a single bond, two lines (=) for a double bond, and three lines (≡) for a triple bond. For example, O=O in oxygen and N≡N in nitrogen.
在结构式中，单键用一条线 (−) 表示，双键用两条线 (=) 表示，三键用三条线 (≡) 表示。例如氧气中 O=O，氮气中 N≡N。
2. Dot-and-Cross Diagrams | 点交叉图
Dot-and-cross diagrams are a visual way to show the origin of electrons in a covalent bond. Dots represent electrons from one atom, crosses from another. Only outer-shell electrons are drawn.
点交叉图是一种直观展示共价键中电子来源的方式。点表示来自一个原子的电子，叉表示来自另一个原子的电子。只绘制最外层电子。
To draw the diagram for methane CH₄: carbon has four outer electrons (shown as dots), each hydrogen brings one electron (shown as crosses). Four bonding pairs form, giving carbon a full outer shell of eight electrons and hydrogen a duplet.
绘制甲烷 CH₄ 的点交叉图：碳有四个最外层电子（用点表示），每个氢提供一个电子（用叉表示）。形成四对共用电子，使得碳最外层满足八电子结构，氢满足二电子结构。
For molecules with multiple bonds, such as carbon dioxide CO₂, the diagram must show two double bonds. Oxygen atoms each contribute two unpaired electrons, and carbon shares two of its electrons with each oxygen, resulting in O=C=O.
对于含有重键的分子，如二氧化碳 CO₂，图必须显示两个双键。每个氧原子提供两个未成对电子，碳与每个氧共用两个电子，结果形成 O=C=O。
Lone pairs (non-bonding pairs) are also shown. In water H₂O, oxygen has two bonding pairs (O−H) and two lone pairs, which influence molecular shape.
孤对电子（非键合电子对）也要标出。在水 H₂O 中，氧有两个键合电子对 (O−H) 和两个孤对电子，这将影响分子形状。
3. The Octet Rule and Its Exceptions | 八隅体规则及其例外
Atoms tend to form covalent bonds until they have eight electrons in their outer shell (an octet). This gives them a stable noble-gas electron configuration. Hydrogen is an exception because it achieves a duplet (2 electrons).
原子倾向于形成共价键，直至最外层达到八个电子（八隅体）。这使其具有稳定的稀有气体电子构型。氢是个例外，因为它只需达到 2 电子（二电子结构）。
Elements in Period 3 and beyond can expand their octet by using low-lying d orbitals. For instance, phosphorus in PCl₅ has ten electrons around it, and sulfur in SF₆ has twelve. These are called expanded octets.
第三周期及以上的元素可以利用低能级 d 轨道扩展八隅体。例如 PCl₅ 中的磷周围有十个电子，SF₆ 中的硫周围有十二个电子，这称为扩展八隅体。
Some molecules have an incomplete octet. Boron in BF₃ has only six electrons in its outer shell. BF₃ acts as an electron-pair acceptor (Lewis acid) and can form a dative bond with a species like NH₃.
有些分子的八隅体不完整。BF₃ 中的硼最外层只有六个电子，它可以作为电子对受体（路易斯酸），与 NH₃ 等物种形成配位键。
Free radicals such as NO and NO₂ contain an odd number of electrons. The nitrogen atom in NO has only seven valence electrons, which makes the molecule paramagnetic and reactive.
自由基如 NO 和 NO₂ 含有奇数个电子。NO 中的氮原子只有七个价电子，这使得分子具有顺磁性且反应活性高。
4. Coordinate (Dative) Covalent Bonds | 配位共价键
A dative covalent bond occurs when one atom provides both electrons for the shared pair. Once formed, it is indistinguishable from an ordinary covalent bond. It is represented by an arrow → from the donor to the acceptor atom.
当一个原子提供共享电子对的两个电子时，便形成配位共价键。一旦形成，它与普通共价键没有区别。通常用箭头 → 从供体指向受体原子来表示。
In the ammonium ion NH₄⁺, the nitrogen atom of ammonia has a lone pair, which it donates to an H⁺ ion (which has no electrons). All four N−H bonds become equivalent.
在铵根离子 NH₄⁺ 中，氨分子的氮原子有一对孤对电子，它将其提供给 H⁺ 离子（没有电子）。结果四个 N−H 键都变得等价。
Another example is the formation of Al₂Cl₆. Each aluminium atom in AlCl₃ is electron-deficient and accepts a lone pair from a chlorine on the other AlCl₃ unit, forming two dative bonds.
另一个例子是 Al₂Cl₆ 的形成。AlCl₃ 中的每个铝原子缺电子，从另一个 AlCl₃ 单元的氯原子接受一对孤对电子，形成两个配位键。
Hydronium ion H₃O⁺ is also formed when a water molecule donates a lone pair to H⁺. Recognising dative bonding is vital for understanding the structures of complex ions and acids.
水分子将孤对电子给予 H⁺ 也可形成水合氢离子 H₃O⁺。识别配位键对于理解复杂离子和酸的结构至关重要。
5. Electronegativity and Bond Polarity | 电负性和键的极性
Electronegativity is the ability of an atom to attract the bonding electrons in a covalent bond. The Pauling scale is commonly used; fluorine (3.98) is the most electronegative element.
电负性是原子在共价键中吸引键合电子的能力。常用鲍林标度，氟 (3.98) 是电负性最强的元素。
In a bond between identical atoms (e.g., Cl−Cl), electrons are shared equally – a pure covalent (non-polar) bond. When the electronegativity difference is large, the bond becomes polar, with a partial negative charge (δ⁻) on the more electronegative atom and a partial positive charge (δ⁺) on the other.
相同原子（如 Cl−Cl）之间的键，电子均等共享，是纯共价（非极性）键。当电负性差值较大时，键具有极性，电负性较大的原子带部分负电荷 (δ⁻)，另一个原子带部分正电荷 (δ⁺)。
A general guideline: a difference of less than 0.5 gives a non-polar bond; between 0.5 and 1.7 gives a polar covalent bond; above 1.7 the bond is considered ionic. However, there is no sharp boundary between ionic and covalent.
通常指南：差值小于 0.5 为非极性键；0.5 至 1.7 为极性共价键；大于 1.7 可视为离子键。但离子键和共价键之间没有绝对界限。
The polarity of bonds determines many physical properties, such as solubility and boiling points. For example, HCl contains a polar bond because chlorine is more electronegative than hydrogen.
键的极性决定了许多物理性质，如溶解度和沸点。例如 HCl 含有极性键，因为氯的电负性大于氢。
6. Dipole Moments and Polar Molecules | 偶极矩和极性分子
A dipole moment arises when a bond has separation of charge. It is a vector quantity, with direction pointing from δ⁺ to δ⁻. The overall polarity of a molecule depends on both bond polarity and molecular geometry.
当键上电荷分离时产生偶极矩。它是一个矢量，方向从 δ⁺ 指向 δ⁻。分子的整体极性取决于键的极性和分子几何结构。
In CO₂, each C=O bond is polar, but the linear shape causes the two bond dipoles to cancel, so the molecule is non-polar. In H₂O, the bent shape prevents cancellation, giving a net dipole moment and a polar molecule.
在 CO₂ 中，每个 C=O 键是极性的，但直线形构型使两个键的偶极抵消，分子非极性。在 H₂O 中，弯曲构型不能抵消，产生净偶极矩，分子为极性。
To determine if a molecule is polar, draw its 3D shape and assign δ⁺/δ⁻ to each bond. If the resultant vector sum is nonzero, the molecule is polar. Symmetrical molecules like CCl₄ are non-polar.
判断分子是否极性，可画出 3D 构型并为每个键标上 δ⁺/δ⁻。如果矢量和不为零，分子是极性的。对称分子如 CCl₄ 是非极性的。
Examples of polar molecules: HF, NH₃, CHCl₃. Non-polar: CH₄, BF₃, PCl₅. This concept explains why polar molecules dissolve in polar solvents such as water.
极性分子例子：HF、NH₃、CHCl₃。非极性分子：CH₄、BF₃、PCl₅。这一概念解释了为什么极性分子能溶于水等极性溶剂。
7. Bond Length, Bond Energy and Bond Order | 键长、键能和键级
Bond length is the average distance between the nuclei of two bonded atoms. Bond energy (bond enthalpy) is the energy required to break one mole of a given bond in the gaseous state. Both depend on the bond order.
键长是两个成键原子核之间的平均距离。键能（键焓）是在气态下断裂一摩尔特定键所需的能量。两者都取决于键级。
As bond order increases, bond length decreases and bond energy increases. For carbon–carbon bonds: C−C bond length 154 pm, bond energy 347 kJ mol⁻¹; C=C length 134 pm, energy 612 kJ mol⁻¹; C≡C length 120 pm, energy 838 kJ mol⁻¹.
随着键级增加，键长减小，键能增大。就碳-碳键而言：C−C 键长 154 pm，键能 347 kJ mol⁻¹；C=C 键长 134 pm，能量 612 kJ mol⁻¹；C≡C 键长 120 pm，能量 838 kJ mol⁻¹。
Average bond energies are used to estimate enthalpy changes in reactions using the equation:
ΔH ≈ Σ(bond energies broken) − Σ(bond energies formed)
平均键能可用于估算反应焓变，公式为：
ΔH ≈ Σ(断裂键的键能) − Σ(形成键的键能)
Note that bond energies vary with molecular environment; data books provide mean values. In WJEC exams, you may be asked to calculate an unknown bond energy from given enthalpy changes.
注意键能随分子环境而变化，数据手册提供平均值。在 WJEC 考试中，可能要求根据给定焓变计算未知键能。
8. Sigma (σ) and Pi (π) Bonds | σ键和π键
A single covalent bond is always a sigma (σ) bond, formed by head-on overlap of orbitals along the internuclear axis. The electron density is concentrated between the nuclei, allowing free rotation around the bond.
单共价键始终是 σ 键，由轨道沿核间轴头对头重叠形成。电子密度集中在两核之间，允许绕键自由旋转。
A double bond consists of one σ bond and one π bond. The π bond results from sideways overlap of p orbitals above and below the plane of the σ bond. The π bond restricts rotation, leading to geometrical isomerism in alkenes.
双键由一个 σ 键和一个 π 键组成。π 键由 p 轨道在 σ 键平面上、下方侧向重叠形成。π 键限制旋转，导致烯烃出现几何异构。
In ethene C₂H₄, the carbon–carbon double bond is σ + π. The molecule is planar with bond angles about 120°. In ethyne C₂H₂, the triple bond is σ + 2π, giving a linear geometry.
在乙烯 C₂H₄ 中，碳-碳双键为 σ + π。分子是平面形，键角约 120°。在乙炔 C₂H₂ 中，三键为 σ + 2π，呈直线形结构。
Nitrogen gas N₂ has a triple bond: one σ and two π bonds. The π bonds are weaker than the σ bond, explaining why the bond enthalpy is high but the π bonds break first in reactions.
氮气 N₂ 具有三键：一个 σ 键和两个 π 键。π 键比 σ 键弱，这解释了为什么键焓很高，但反应中 π 键首先断裂。
9. Shapes of Molecules: VSEPR Theory | 分子形状：VSEPR理论
The Valence Shell Electron Pair Repulsion (VSEPR) theory states that electron pairs around a central atom will arrange themselves to minimise repulsion. The shape is determined by the number of bonding pairs and lone pairs.
价层电子对互斥理论 (VSEPR) 指出，中心原子周围的电子对会排列成最小排斥的结构。形状由键合电子对和孤对电子的数量决定。
| Bonding pairs / Lone pairs | Shape | Bond angle (°) | Example |
|---|---|---|---|
| 2 / 0 | Linear | 180 | BeCl₂, CO₂ |
| 3 / 0 | Trigonal planar | 120 | BF₃, SO₃ |
| 4 / 0 | Tetrahedral | 109.5 | CH₄, NH₄⁺ |
| 3 / 1 | Pyramidal | 107 | NH₃ |
| 2 / 2 | Bent / V-shaped | 104.5 | H₂O |
| 5 / 0 | Trigonal bipyramidal | 90, 120 | PCl₅ |
| 6 / 0 | Octahedral | 90 | SF₆ |
Lone pairs repel more strongly than bonding pairs, reducing bond angles. For example, NH₃ has one lone pair, compressing the H−N−H angle from 109.5° to about 107°. In H₂O with two lone pairs, the angle is 104.5°.
孤对电子比键合电子对排斥更强，使键角减小。如 NH₃ 有一个孤对电子，将 H−N−H 角从 109.5° 压缩至约 107°；H₂O 有两个孤对电子，键角为 104.5°。
To predict a shape, count the total number of electron pairs (bonding + lone) around the central atom. This gives the electron pair geometry, then deduce the molecular shape by ignoring the lone pairs’ positions.
预测形状时，先计算中心原子周围的电子对总数（键合 + 孤对），确定电子对几何构型，再忽略孤对电子位置推导出分子形状。
10. Delocalised π Electrons and Resonance | 离域π电子和共振
Some molecules and ions cannot be represented by a single Lewis structure. The actual structure is a hybrid of several resonance forms, with delocalised electrons spread over several atoms.
有些分子和离子不能用单一路易斯结构表示。实际结构是几个共振形式的杂化体，电子离域分布在多个原子上。
In the nitrate ion NO₃⁻, three equivalent structures show the double bond in different positions. The true structure has equal N−O bond lengths, with the π electrons delocalised across all three oxygen atoms.
硝酸根离子 NO₃⁻ 有三个等价结构，双键在不同位置。真实结构中 N−O 键长均等，π 电子离域遍布三个氧原子。
Benzene C₆H₆ is a classic example. Two Kekulé structures with alternating single and double bonds contribute, but all C−C bonds are identical (139 pm, between C−C and C=C). The delocalised π cloud gives aromatic stability.
苯 C₆H₆ 是经典例子。两种凯库勒结构（单双键交替）都有贡献，但所有 C−C 键长相同（139 pm，介于 C−C 和 C=C 之间）。离域 π 电子云赋予芳香稳定性。
The concept of delocalisation helps explain why some species are more stable than expected, and why bond lengths are intermediate. In WJEC, you may be asked to draw resonance structures for NO₃⁻, CO₃²⁻, or benzene.
离域概念解释了为何某些物质比预期更稳定，以及键长为中间值的原因。在 WJEC 考试中，可能要求绘制 NO₃⁻、CO₃²⁻ 或苯的共振结构。
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