📚 IB Chemistry: Chemical Bonding Key Points Review | IB 化学：化学键 考点精讲
Chemical bonding is the heart of IB Chemistry, linking structure to properties and reactivity. A clear grasp of ionic, covalent, and metallic bonding, along with intermolecular forces, allows you to predict shapes, polarity, boiling points, and conductivity. This article distils the essential exam-focussed knowledge, from electronegativity trends to VSEPR and hybridisation, helping you avoid common pitfalls and answer with confidence.
化学键是 IB 化学的核心，它将结构与性质、反应活性紧密联系起来。清晰掌握离子键、共价键和金属键，以及分子间作用力，能让你预测分子的形状、极性、沸点和导电性。本文提炼了考试必考的核心知识，从电负性规律到 VSEPR 理论和杂化，帮助大家避开常见陷阱，自信作答。
1. Types of Bonding: An Overview | 化学键类型概述
Chemical bonding arises from electrostatic attractions between nuclei and electrons. The three primary types are ionic (transfer of electrons), covalent (sharing of electrons), and metallic (delocalised electrons in a lattice). The type of bonding is primarily determined by the electronegativity difference (ΔEN) between the combining atoms.
化学键源自原子核与电子之间的静电吸引。三种主要类型包括离子键（电子转移）、共价键（电子共享）和金属键（晶格中的离域电子）。成键类型主要由结合原子间的电负性差值（ΔEN）决定。
A ΔEN > 1.8 typically leads to ionic bonding, while ΔEN < 0.5 gives a pure covalent bond. Between 0.5 and 1.8, polar covalent bonds form, with unequal sharing of electrons. Metallic bonding occurs between metal atoms with low electronegativity, forming a 'sea of electrons'.
通常 ΔEN > 1.8 形成离子键，ΔEN < 0.5 为纯共价键。介于 0.5 到 1.8 之间则形成极性共价键，电子共享不均匀。金属键发生于低电负性的金属原子间，形成“电子海洋”。
2. Ionic Bonding and Lattice Energy | 离子键与晶格能
Ionic bonding involves the complete transfer of valence electrons, generating cations and anions held together in a giant ionic lattice. The strength of an ionic bond is measured by lattice enthalpy (ΔHlatt⦵), the energy released when gaseous ions form one mole of solid crystal. More highly charged and smaller ions give more exothermic lattice enthalpies, resulting in higher melting points.
离子键涉及价电子的完全转移，生成阳离子和阴离子，它们通过巨型离子晶格结合在一起。离子键的强度用晶格焓（ΔHlatt⦵）衡量，即气态离子形成一摩尔固态晶体时释放的能量。电荷越高、半径越小，晶格焓越负，熔点也越高。
The Born–Haber cycle is an application of Hess's Law linking formation enthalpy to ionisation energy, electron affinity, atomisation, and lattice enthalpy. Exam questions often ask you to calculate an unknown lattice enthalpy or to explain trends in thermal stability of Group 1/2 carbonates and nitrates using ionic polarisation.
玻恩-哈伯循环是赫斯定律的应用，将生成焓与电离能、电子亲和能、原子化焓及晶格焓联系起来。考试常要求计算未知的晶格焓，或利用离子极化作用解释第1、2族碳酸盐和硝酸盐的热稳定性趋势。
3. Covalent Bonding: Sharing Electrons | 共价键：共享电子
Covalent bonds form when atomic orbitals overlap and electron pairs are shared between two nuclei. A single bond is one shared pair (σ bond), a double bond contains one σ and one π bond, and a triple bond contains one σ and two π bonds. Bond length decreases and bond strength increases with more shared pairs.
当原子轨道重叠、两个原子核共用电子对时，便形成共价键。单键对应一个共享电子对（σ 键），双键含一个 σ 键和一个 π 键，三键含一个 σ 键和两个 π 键。共享电子对越多，键长越短，键能越高。
Resonance structures arise when a molecule or ion cannot be represented by a single Lewis structure; the actual structure is a hybrid, with delocalised π electrons. Classic examples are the carbonate ion CO₃²⁻ and benzene C₆H₆, where all C–O or C–C bonds are identical in length and strength. For IB, you must be able to draw resonance hybrids with delocalised electron clouds.
当一个分子或离子无法用单一路易斯结构表示时，就会出现共振结构；实际结构为共振杂化体，具有离域 π 电子。典型例子是碳酸根离子 CO₃²⁻ 和苯 C₆H₆，其中所有 C–O 或 C–C 键的键长和强度都相同。IB 考试要求能画出带离域电子云的共振杂化体。
4. Dative Covalent (Coordinate) Bonds | 配位共价键
A dative covalent bond occurs when both shared electrons originate from the same atom. The donor must have a lone pair, and the acceptor must be electron deficient. A classic example is the ammonium ion NH₄⁺, where the nitrogen lone pair is donated to a H⁺ ion. Another is the formation of Al₂Cl₆, where chlorine lone pairs bridge the electron-deficient aluminium centres.
当两个共享电子都来自同一个原子时，形成配位共价键。供体必须含有孤对电子，受体必须缺电子。经典例子是铵根离子 NH₄⁺，氮原子的孤对电子提供给 H⁺；另一个例子是 Al₂Cl₆ 的形成，氯的孤对电子桥连了缺电子的铝中心。
Once formed, a dative bond is indistinguishable from an ordinary covalent bond; both are electron pairs shared between atoms. In Lewis structures, a dative bond is often indicated by an arrow pointing from donor to acceptor, but in exam answers, you may simply show the bond and state that it is dative.
一旦形成，配位键与普通共价键无法区分，都是原子间共享的电子对。路易斯结构中常用从供体指向受体的箭头表示配位键，但在考试作答中，只需画出该键并注明其为配位键即可。
5. Electronegativity and Bond Polarity | 电负性与键的极性
Electronegativity (EN) is the ability of an atom to attract the bonding pair of electrons in a covalent bond. Pauling scale values increase across a period and decrease down a group. Fluorine (4.0) is the most electronegative element. Non-polar bonds have equal sharing, while polar bonds have unequally shared electron pairs, producing partial charges δ+ and δ−.
电负性（EN）是原子在共价键中吸引成键电子对的能力。鲍林标度值在同一周期中从左到右递增，同一族从上到下递减。氟（4.0）是电负性最大的元素。非极性键中电子均等共享，极性键则共享不均，产生部分电荷 δ+ 和 δ−。
Bond polarity is a vector quantity; the overall molecular polarity depends on both bond polarities and molecular geometry. A molecule with polar bonds can be non-polar if the dipoles cancel due to symmetry, as in CO₂ (linear) and CCl₄ (tetrahedral). Exam tip: always draw the shape and show resultant dipole moment to justify polarity.
键的极性是矢量；分子的整体极性取决于键的极性和分子几何构型。如果因对称性使得键的偶极相互抵消，即便含有极性键，分子仍为非极性，例如 CO₂（直线形）和 CCl₄（四面体形）。考试技巧：一定要画出分子形状并标出净偶极矩，以此论证极性。
6. VSEPR Theory and Molecular Shapes | VSEPR 理论与分子形状
Valence Shell Electron Pair Repulsion (VSEPR) theory states that electron pairs around a central atom arrange themselves to minimise repulsion. The order of repulsion is: lone pair–lone pair > lone pair–bond pair > bond pair–bond pair. This determines the molecular geometry, not the electron domain geometry.
价层电子对互斥理论（VSEPR）认为，中心原子周围的电子对会通过排布使互斥力最小化。斥力大小顺序为：孤对-孤对 > 孤对-键对 > 键对-键对。这决定了分子的几何形状，而不仅仅是电子域几何。
mains → octahedral (90°). Lone pairs reduce bond angles: e.g., NH₃ is trigonal pyramidal (~107°), and H₂O is bent (~104.5°). You must be able to predict and draw shapes from Lewis structures.
IB 常见形状：2 个电子域 → 直线形（180°）；3 个域 → 平面三角形（120°）；4 个域 → 四面体形（109.5°）；5 个域 → 三角双锥形（120° 和 90°）；6 个域 → 八面体形（90°）。孤对电子会使键角减小，例如 NH₃ 为三角锥形（~107°），H₂O 为 V 形（~104.5°）。必须能够根据路易斯结构预测并画出分子形状。
7. Hybridisation: sp, sp², sp³ | 杂化：sp, sp², sp³
Hybridisation describes the mixing of atomic orbitals to form new hybrid orbitals suitable for bonding. In IB HL, you deal with sp, sp², and sp³ hybridisation. sp³ hybridisation occurs when one s and three p orbitals blend to give four equivalent orbitals, as in methane (CH₄) with a tetrahedral shape and 109.5° bond angles.
杂化描述的是原子轨道混合形成新的、适合成键的杂化轨道。在 IB 高阶水平（HL）中，你需要掌握 sp、sp² 和 sp³ 杂化。sp³ 杂化由一个 s 轨道和三个 p 轨道混合，产生四个等价的杂化轨道，如甲烷（CH₄）具有四面体形和 109.5° 键角。
sp² hybridisation results from one s and two p orbitals, yielding three hybrid orbitals and one unhybridised p orbital, forming trigonal planar geometry (120°), as in ethene (C₂H₄). sp hybridisation involves one s and one p orbital, producing two linear orbitals (180°) and two unhybridised p orbitals, as in ethyne (C₂H₂). The unhybridised p orbitals form π bonds.
sp² 杂化由一个 s 和两个 p 轨道产生，给出三个杂化轨道和一个未杂化的 p 轨道，形成平面三角形几何（120°），如乙烯（C₂H₄）。sp 杂化涉及一个 s 和一个 p 轨道，产生两个直线形杂化轨道（180°）和两个未杂化的 p 轨道，如乙炔（C₂H₂）。未杂化的 p 轨道用于形成 π 键。
8. Intermolecular Forces: London, Dipole-Dipole, Hydrogen Bonds | 分子间作用力：伦敦力、偶极-偶极、氢键
Intermolecular forces (IMFs) determine physical properties like boiling point, melting point, and solubility. The weakest are London (dispersion) forces, present in all molecules due to instantaneous dipoles. Their strength increases with the number of electrons (larger molar mass and greater polarisability).
分子间作用力（IMFs）决定了沸点、熔点和溶解度等物理性质。最弱的是伦敦（色散）力，由瞬时偶极引起，存在于所有分子中。其强度随电子数增加而增强（更大的摩尔质量和更大的可极化性）。
Permanent dipole-dipole forces exist only between polar molecules and are stronger than London forces between molecules of comparable size. Hydrogen bonding – the strongest IMF considered at IB – occurs when H is bonded to N, O, or F and interacts with a lone pair on N, O, or F. It explains the anomalously high boiling points of H₂O, NH₃, and HF, as well as the base-pairing in DNA and the secondary structure of proteins.
永久偶极-偶极力仅存在于极性分子之间，在大小相近的分子中，这种力强于伦敦力。氢键——IB 化学中讨论的最强分子间作用力——出现在 H 与 N、O、F 成键，并与另一个 N、O、F 上的孤对电子相互作用的情况下。它解释了 H₂O、NH₃ 和 HF 异常高的沸点，以及 DNA 中的碱基配对和蛋白质的二级结构。
9. Metallic Bonding and Alloys | 金属键与合金
Metallic bonding consists of a regular lattice of positive metal ions surrounded by a 'sea' of delocalised valence electrons. This model explains high electrical and thermal conductivity (mobile electrons), malleability and ductility (layers can slide without breaking the metallic bond). Melting points generally increase with charge density of the metal ion.
金属键由规则排列的金属正离子和环绕周围的离域价电子“海洋”构成。这一模型可解释金属的高导电性和导热性（可移动的电子）、延展性（层间滑动而不断裂金属键）。金属的熔点通常随金属离子电荷密度的增大而升高。
Alloys are mixtures of a metal with other elements, often designed to alter properties. The introduction of differently sized atoms disrupts the regular lattice, making it harder for layers to slide – thus alloys are typically harder and less malleable than pure metals. Common examples include steel (iron with carbon) and brass (copper with zinc).
合金是金属与其他元素的混合物，旨在改变材料性质。引入不同尺寸的原子会打乱规则晶格，使层间更难滑动——因此合金通常比纯金属更硬、延展性更差。常见的例子包括钢（铁与碳）和黄铜（铜与锌）。
10. Physical Properties Related to Bonding | 与化学键相关的物理性质
Bonding type directly dictates physical behaviour. Giant ionic lattices are hard, brittle, and conduct electricity only when molten or dissolved, as ions become mobile. Giant covalent networks (diamond, SiO₂) have extremely high melting points and, apart from graphite, do not conduct electricity. Simple molecular substances have low melting/boiling points and are often insulators.
键型直接决定了物质的物理行为。巨型离子晶格坚硬、易碎，只有熔融或溶于水时才能导电，因为此时离子可以自由移动。巨型共价网络（金刚石、SiO₂）熔点极高，且除石墨外一般不导电。简单分子物质熔、沸点低，通常为绝缘体。
Graphite is a classic exception: each carbon is sp² hybridised, forming layers of hexagonal rings with delocalised π electrons between layers, allowing electrical conductivity along the planes. This contrast with diamond, where all four valence electrons form σ bonds in a tetrahedral, insulating structure, frequently appears in exam comparison questions.
石墨是一个经典例外：每个碳原子为 sp² 杂化，形成六边形环层，层间有离域 π 电子，故而能沿层平面导电。这与金刚石形成了鲜明对比——金刚石中四个价电子都用于形成四面体 σ 键，是绝缘体。这类比较题在考试中很常见。
11. Exam Tips and Common Pitfalls | 考试技巧与常见错误
Ensure you can distinguish between intermolecular and intramolecular forces – a frequent mistake is to cite 'breaking covalent bonds' when a substance boils, whereas boiling overcomes intermolecular forces, not bonds. Also remember that ionic compounds conduct when molten or in aqueous solution, but not as solids. Practice drawing VSEPR shapes with clear dash-wedge notation and identifying bond angles.
务必区分分子间力和分子内力——一个常见错误是认为物质沸腾时破坏了共价键，事实上沸腾克服的是分子间作用力，而非化学键。还要记住，离子化合物在熔融态或水溶液中导电，但固态时不导电。多加练习用楔形式清晰地画出 VSEPR 形状，并判定键角。
When explaining trends, always link to charge density, polarisability, or electronegativity. For solubility, ‘like dissolves like’ is a good starting point: polar solvents dissolve polar/ionic solutes; non-polar solvents dissolve non-polar solutes. Use precise terminology: ‘delocalised electrons’ in metals and graphite, ‘mobile ions’ in electrolytes, and ‘dipole-dipole’, not just ‘dipole’.
解释递变规律时，一定要联系电荷密度、可极化性或电负性。对于溶解性，可从“相似相溶”入手：极性溶剂溶解极性/离子型溶质，非极性溶剂溶解非极性溶质。使用准确的术语：金属和石墨中的“离域电子”，电解质中的“可移动离子”，以及“偶极-偶极力”而不仅仅是“偶极”。
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