Covalent Bonding Essentials for IB & CIE Chemistry | IB CIE 化学：共价键考点精讲

📚 Covalent Bonding Essentials for IB & CIE Chemistry | IB CIE 化学：共价键考点精讲

Covalent bonding lies at the heart of molecular chemistry. Whether you are preparing for IB or CIE A-level examinations, a thorough understanding of how atoms share electrons, the resulting molecular shapes, and the nuances of bond polarity, hybridisation, and resonance is essential. This article systematically covers every key concept you need to master, from the octet rule to VSEPR theory and hybridisation, with clear explanations and examination-focused insights.

共价键是分子化学的核心。无论你在准备IB还是CIE A-level考试，透彻理解原子如何共享电子、由此形成的分子形状，以及键的极性、杂化和共振等细节都至关重要。本文系统讲解你需要掌握的每一个关键概念，从八隅律到VSEPR理论和杂化，配以清晰的解释和紧扣考点的分析。

1. Formation of Covalent Bonds | 共价键的形成

A covalent bond forms when two non-metal atoms share one or more pairs of electrons. The shared electron pair is attracted to both nuclei, creating a stable balance between attractive and repulsive forces. This sharing allows each atom to attain a more stable electronic configuration, often a noble gas arrangement. For example, in a chlorine molecule Cl₂, each chlorine atom contributes one electron to the shared pair, achieving a full outer shell.

当两个非金属原子共享一对或多对电子时，便形成共价键。共享的电子对同时被两个原子核吸引，在吸引与排斥力之间达成稳定平衡。这种共享使每个原子获得更稳定的电子构型，通常为稀有气体构型。例如，在氯分子Cl₂中，每个氯原子各提供一个电子形成共享电子对，从而实现满壳层结构。

In Lewis structures, a covalent bond is represented by a line (—) between the two atoms. Lone pairs of electrons not involved in bonding are shown as dots. The total number of electrons around each atom should match the octet rule for Period 2 elements, though there are exceptions.

在路易斯结构中，共价键用两个原子之间的一条短线（—）表示。未参与成键的孤对电子以点表示。每个原子周围的总电子数应符合第二周期元素的八隅律，但存在例外。

2. Octet Rule and Exceptions | 八隅律及其例外

The octet rule states that atoms tend to form bonds until they are surrounded by eight valence electrons, achieving a stable noble gas configuration. This rule successfully explains bonding in most organic molecules and simple inorganic compounds like H₂O, NH₃, and CH₄. However, several important exceptions appear frequently in IB and CIE exam questions.

八隅律指出，原子倾向于通过成键使周围价电子数达到8个，从而获得稳定的稀有气体构型。这一规则成功地解释了大多数有机分子和简单无机化合物（如H₂O、NH₃、CH₄）中的成键。但考试中经常出现几个重要例外。

Incomplete octet: Boron in BF₃ has only six electrons around it. Beryllium in BeCl₂ has four. These electron-deficient compounds can act as Lewis acids.
不完整八隅体： BF₃中的硼周围仅有6个电子。BeCl₂中的铍只有4个电子。这些缺电子化合物可作为路易斯酸。
Expanded octet: Elements from Period 3 onwards, such as phosphorus in PCl₅ and sulfur in SF₆, can have 10 or 12 electrons around the central atom because they have accessible d-orbitals.
扩展八隅体： 第三周期及之后的元素，如PCl₅中的磷和SF₆中的硫，中心原子周围可容纳10或12个电子，因为它们有可用的d轨道。
Odd-electron species: Molecules like NO and NO₂ have an odd number of valence electrons and cannot satisfy the octet rule for every atom. These are called free radicals.
奇电子物种： NO和NO₂等分子具有奇数价电子，无法让每个原子都满足八隅律，被称为自由基。

3. Coordinate (Dative) Covalent Bonds | 配位键（配位共价键）

A coordinate bond is a covalent bond in which both electrons of the shared pair come from the same atom. Once formed, a coordinate bond is indistinguishable from an ordinary covalent bond. A common example is the ammonium ion NH₄⁺, where the nitrogen atom of ammonia donates its lone pair to a hydrogen ion, H⁺.

配位键是一种共价键，其中共享电子对的两个电子均来自同一个原子。一旦形成，配位键与普通共价键无法区分。常见的例子是铵根离子NH₄⁺，氨分子中的氮原子将其孤对电子提供给氢离子H⁺。

Another key example is the formation of the Al₂Cl₆ dimer. In this molecule, two AlCl₃ units join via coordinate bonds from chlorine atoms on one Al to the other. Carbon monoxide, CO, also contains a coordinate bond in addition to two normal covalent bonds, giving a triple bond overall. In Lewis structures, an arrow (→) is sometimes used to indicate the donor atom.

另一个重要例子是Al₂Cl₆二聚体的形成。两个AlCl₃单元通过一个氯原子向另一个铝原子提供配位键而结合。一氧化碳CO除了两个普通共价键外，还有一个配位键，总体上形成三键。在路易斯结构中，有时会用箭头（→）标明电子对给予体。

4. Bond Length, Bond Energy, and Bond Order | 键长、键能与键级

Bond length is the distance between the nuclei of two bonded atoms. Bond energy is the energy required to break one mole of covalent bonds in the gaseous state. Bond order is the number of electron pairs shared between two atoms. Single bonds have bond order 1, double bonds 2, triple bonds 3.

键长是两个成键原子核之间的距离。键能是断裂气态中一摩尔共价键所需的能量。键级是两个原子间共享的电子对数。单键键级为1，双键为2，三键为3。

As bond order increases, bond length decreases and bond energy increases. For carbon–carbon bonds: C–C (bond order 1) has length 154 pm and energy 348 kJ·mol⁻¹; C=C (2) length 134 pm, energy 614 kJ·mol⁻¹; C≡C (3) length 120 pm, energy 839 kJ·mol⁻¹. These trends are frequently tested in relation to reactivity and bond strength.

随着键级升高，键长缩短，键能增大。以碳-碳键为例：C–C（键级1）键长154 pm，键能348 kJ·mol⁻¹；C=C（键级2）键长134 pm，键能614 kJ·mol⁻¹；C≡C（键级3）键长120 pm，键能839 kJ·mol⁻¹。这些趋势常结合反应活性和键强度进行考查。

Bond	Bond Order	Length / pm	Energy / kJ·mol⁻¹
C–C	1	154	348
C=C	2	134	614
C≡C	3	120	839

5. Polar and Non-polar Covalent Bonds | 极性共价键与非极性共价键

A non-polar covalent bond occurs when the two atoms involved have identical or very similar electronegativities, resulting in an equal sharing of electrons. Examples include H–H, Cl–Cl, and C–H bonds. In a polar covalent bond, the atom with the higher electronegativity attracts the bonding electrons more strongly, creating partial charges. The symbol 𝛿⁺ is used for the electron-poor end and 𝛿⁻ for the electron-rich end.

当成键的两个原子电负性相同或极为相近时，电子对平均共享，形成非极性共价键。例子包括H–H、Cl–Cl和C–H键。在极性共价键中，电负性更高的原子更强烈地吸引成键电子，产生部分电荷。缺电子端标为𝛿⁺，富电子端标为𝛿⁻。

Hydrogen fluoride, HF, is a classic example: fluorine has a much higher electronegativity (4.0) than hydrogen (2.2), so the bond is highly polar. The polarity of bonds affects the overall dipole moment of the molecule, which in turn influences physical properties such as boiling point and solubility. A molecule can have polar bonds but be non-polar overall if the shape is symmetric, such as CO₂ or CCl₄.

氟化氢HF是一个典型例子：氟的电负性（4.0）远高于氢（2.2），因此键的极性很强。键的极性会影响分子的整体偶极矩，进而影响沸点、溶解度等物理性质。分子可以含有极性键，但如果形状对称，如CO₂或CCl₄，分子整体仍为非极性。

6. Shapes of Molecules: VSEPR Theory | 分子形状：VSEPR理论

Valence Shell Electron Pair Repulsion (VSEPR) theory states that electron pairs around a central atom arrange themselves to minimise repulsion. The order of repulsion strength is: lone pair–lone pair > lone pair–bonding pair > bonding pair–bonding pair. The shape of a molecule is determined by the number of bonding pairs and lone pairs on the central atom.

价层电子对互斥理论（VSEPR）指出，中心原子周围的电子对会尽量远离以减小排斥力。排斥力强弱顺序为：孤对–孤对 > 孤对–键对 > 键对–键对。分子形状取决于中心原子上的键对和孤对数目。

Common shapes tested in IB and CIE include: linear (2 bonding pairs, 0 lone pairs, e.g. BeCl₂, CO₂, angle 180°); trigonal planar (3, 0, e.g. BF₃, 120°); tetrahedral (4, 0, e.g. CH₄, 109.5°); trigonal pyramidal (3, 1, e.g. NH₃, ~107°); bent (2, 2, e.g. H₂O, ~104.5°); and octahedral (6, 0, e.g. SF₆, 90°). You must also be able to predict shapes with expanded octets, such as trigonal bipyramidal (5, 0, PCl₅) and square planar (4, 2, XeF₄).

IB和CIE常考的形状包括：直线形（2个键对，0个孤对，如BeCl₂、CO₂，键角180°）；平面三角形（3，0，如BF₃，120°）；四面体形（4，0，如CH₄，109.5°）；三角锥形（3，1，如NH₃，~107°）；V形（2，2，如H₂O，~104.5°）；以及八面体形（6，0，如SF₆，90°）。你还应能够预测扩展八隅体化合物的形状，如三角双锥形（5，0，PCl₅）和平面四边形（4，2，XeF₄）。

7. Hybridisation of Atomic Orbitals | 原子轨道的杂化

To explain molecular shapes and bond angles, VSEPR is combined with the concept of hybridisation. Hybrid orbitals are formed by mixing atomic orbitals on the same atom. The type of hybridisation depends on the number of electron domains (bonding + lone pairs) around the central atom.

为解释分子形状和键角，VSEPR常与杂化概念结合。杂化轨道由同一原子上不同原子轨道混合而成。杂化类型取决于中心原子周围的电子域（键对+孤对）数目。

2 electron domains → sp hybridisation, linear, 180°.
2个电子域 → sp杂化，直线形，180°。
3 electron domains → sp² hybridisation, trigonal planar, 120°.
3个电子域 → sp²杂化，平面三角形，120°。
4 electron domains → sp³ hybridisation, tetrahedral, 109.5°.
4个电子域 → sp³杂化，四面体形，109.5°。
5 electron domains → sp³d hybridisation, trigonal bipyramidal, 90° and 120°.
5个电子域 → sp³d杂化，三角双锥形，90°和120°。
6 electron domains → sp³d² hybridisation, octahedral, 90°.
6个电子域 → sp³d²杂化，八面体形，90°。

In IB HL and CIE A2, students must assign hybridisation to carbon atoms in organic molecules. For example, in ethene (C₂H₄), each carbon is sp² hybridised; in ethyne (C₂H₂), sp. Understanding hybridisation also helps explain the relative length and strength of sigma and pi bonds.

在IB HL和CIE A2中，学生需指出有机分子中碳原子的杂化类型。例如，乙烯(C₂H₄)中碳为sp²杂化；乙炔(C₂H₂)中碳为sp杂化。理解杂化也有助于解释σ键和π键的相对长度与强度。

8. Resonance Structures | 共振结构

When a molecule or ion cannot be adequately represented by a single Lewis structure, resonance structures are used. The actual structure is a resonance hybrid, a blend of the contributing forms. The classic example is the carbonate ion CO₃²⁻, which has three equivalent C–O bonds that are intermediate between single and double bonds. The bonding electrons are delocalised over the oxygen atoms.

当单一路易斯结构无法充分表示分子或离子时，便使用共振结构。实际结构是共振杂化体，是各贡献形式的混合。典型例子是碳酸根离子CO₃²⁻，它有三个等价的C–O键，键级介于单键和双键之间。成键电子离域于所有氧原子上。

Other important cases include the nitrate ion NO₃⁻, ozone O₃, and the benzene ring C₆H₆. In benzene, all C–C bonds have the same length and bond order 1.5. Resonance stabilises molecules, lowering their energy. When answering exam questions, always draw the double-headed arrow between resonance forms and state that the actual structure has bond lengths intermediate between the extremes.

其他重要例子包括硝酸根离子NO₃⁻、臭氧O₃和苯环C₆H₆。在苯中，所有C–C键长相等，键级为1.5。共振使分子更稳定，能量更低。在考试作答时，始终在共振式之间画双箭头，并说明实际结构的键长介于各极端形式之间。

9. Sigma and Pi Bonds | σ键与π键

Covalent bonds can be classified as sigma (σ) or pi (π) bonds according to the orbital overlap. A sigma bond results from head-on overlap of orbitals, with the electron density concentrated along the internuclear axis. All single bonds are sigma bonds. In multiple bonds, the first bond is sigma, and the additional bonds are pi bonds. A pi bond forms by sideways overlap of p orbitals, producing a region of electron density above and below the axis.

共价键可根据轨道重叠方式分为σ键和π键。σ键由轨道“头对头”重叠形成，电子密度集中在核间轴。所有单键均为σ键。在多重键中，第一个键是σ键，其余为π键。π键由p轨道的侧向重叠形成，在键轴上下方产生电子密度区域。

For example, the double bond in ethene consists of one sigma and one pi bond; the triple bond in ethyne is one sigma plus two pi bonds. Pi bonds are weaker than sigma bonds and restrict rotation, which is the reason for cis–trans isomerism in alkenes. This distinction is particularly important in IB HL and CIE organic chemistry.

例如，乙烯的双键由一个σ键和一个π键组成；乙炔的三键为一个σ键加两个π键。π键弱于σ键且限制旋转，这就是烯烃存在顺反异构的原因。这一区别在IB HL和CIE有机化学中尤为重要。

10. Delocalised Electrons and the Metallic Bonding Contrast | 离域电子及与金属键的对比

In some covalent structures, pi electrons are spread over several adjacent atoms rather than being confined to a single bond. This delocalisation results in enhanced stability and is a key feature of conjugated systems and aromatic compounds. Graphite, an allotrope of carbon, contains layers of sp² carbons with delocalised electrons throughout the layer, which explains its electrical conductivity.

在某些共价结构中，π电子不再局限于单个键，而是分散在相邻的多个原子上。这种离域作用增强了稳定性，是共轭体系与芳香族化合物的关键特征。石墨作为碳的同素异形体，其层面内sp²碳原子具有遍布整个层面的离域电子，这解释了它的导电性。

It is useful to contrast covalent bonding with metallic bonding. In metals, valence electrons are delocalised into a ‘sea’ of electrons that binds positive metal ions. Covalent substances typically form discrete molecules or giant covalent networks (e.g. diamond, SiO₂) with strong directional bonds, whereas metallic bonds are non-directional. Both appear in exam questions comparing properties such as melting point, conductivity, and solubility.

将共价键与金属键进行对比很有帮助。金属中，价电子离域成为“电子海”，将金属正离子结合在一起。共价物质通常形成分立分子或巨型共价网络（如金刚石、SiO₂），具有强方向性键；而金属键是非方向性的。两者在比较熔点、导电性和溶解度的考题中均有涉及。

11. Intermolecular Forces vs Covalent Bonds | 分子间作用力与共价键的区别

Students often confuse intermolecular forces with covalent bonds. Covalent bonds are strong intramolecular forces holding atoms together within a molecule (bond energies typically 150–800 kJ·mol⁻¹). Intermolecular forces — London dispersion, dipole–dipole, and hydrogen bonding — are much weaker (2–40 kJ·mol⁻¹) and act between molecules.

学生常将分子间作用力与共价键混淆。共价键是分子内部将原子结合在一起的强作用力（键能通常为150–800 kJ·mol⁻¹）。分子间作用力——伦敦色散力、偶极–偶极力及氢键——则弱得多（2–40 kJ·mol⁻¹），作用于不同分子之间。

When asked to explain boiling points of covalent substances, always refer to the strength of intermolecular forces, not the strength of covalent bonds within the molecule. For example, H₂O has a relatively high boiling point due to extensive hydrogen bonding, while CH₄ is low boiling because only weak London forces exist. Breaking covalent bonds would correspond to decomposition, not phase change.

当被要求解释共价物质的沸点时，始终应提及分子间作用力的强弱，而非分子内部共价键的强度。例如，H₂O因广泛的氢键而具有较高沸点，而CH₄因只存在弱伦敦力而沸点很低。断裂共价键将意味着分解，而非相变。

12. Exam Tips for Covalent Bonding Questions | 共价键考题应试技巧

When tackling covalent bonding questions in IB or CIE exams, start by drawing a clear Lewis structure. Count total valence electrons, identify the central atom, and distribute electrons to satisfy the octet rule where possible. Then apply VSEPR to determine electron domain geometry and molecular shape. Remember to adjust bond angles for lone-pair repulsion.

在应对IB或CIE考试中的共价键题目时，先画出清晰的路易斯结构。计算总价电子数，找出中心原子，尽量按八隅律分配电子。然后运用VSEPR确定电子域几何构型和分子形状。记得根据孤对排斥调整键角。

For bond polarity questions, use electronegativity values to assign partial charges and deduce whether the molecule has an overall dipole.
遇到键极性问题时，用电负性数值标出部分电荷，并推断分子是否具有总偶极矩。
In hybridisation, link the number of electron domains directly to the hybrid type: sp, sp², sp³, etc.
杂化方面，直接将电子域数目与杂化类型相关联：sp、sp²、sp³等。
Resonance questions require you to draw all valid forms and explain delocalisation and bond order.
共振题要求画出所有合理形式，并解释离域现象和键级。
Remember that expanded octets are only possible for central atoms from Period 3 onwards.
记住，只有第三周期及以后的中心原子才可能具有扩展八隅体。
Always distinguish between breaking intermolecular forces (physical change) and breaking covalent bonds (chemical change).
始终区分断裂分子间作用力（物理变化）与断裂共价键（化学变化）。

Mastering these concepts will give you the confidence to answer both multiple-choice and structured questions accurately. Practice with past paper questions, paying special attention to drawing Lewis structures, predicting shapes and angles, and explaining properties through bonding models.

掌握这些概念将使你有信心准确回答选择题与结构题。使用历年真题练习，特别关注绘制路易斯结构、预测形状和键角，以及运用成键模型解释性质。

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Covalent Bonding Essentials for IB & CIE Chemistry | IB CIE 化学：共价键 考点精讲