IB Chemistry: Chemical Bonding and Structure Core Concepts | IB化学：化学键与物质结构核心概念

📚 IB Chemistry: Chemical Bonding and Structure Core Concepts | IB化学：化学键与物质结构核心概念

Chemical bonding and structure form the foundation of IB Chemistry, explaining why substances behave the way they do. From ionic lattices to molecular orbitals, this revision guide distills the essential concepts you need for exam success.

化学键与物质结构是IB化学的核心基础，解释了物质为何表现出特定的性质。从离子晶格到分子轨道，本复习指南浓缩了你考试成功所需的核心概念。

1. Ionic Bonding | 离子键

Ionic bonding occurs when electrons are transferred from a metal to a non-metal, resulting in oppositely charged ions held together by electrostatic attraction. This transfer typically happens between atoms with large electronegativity differences, usually greater than 1.7 on the Pauling scale.

离子键发生在电子从金属转移至非金属时，产生带相反电荷的离子，通过静电吸引力结合在一起。这种转移通常发生在电负性差异较大的原子之间，在鲍林标度上通常大于1.7。

Key features of ionic compounds include:

离子化合物的关键特征包括：

High melting and boiling points due to strong electrostatic forces in the lattice | 高熔沸点，因为晶格中存在强烈的静电作用力
Conduct electricity when molten or dissolved in water, but not in solid state | 熔融或溶于水时导电，固态时不导电
Often soluble in polar solvents like water | 通常溶于极性溶剂如水
Hard but brittle: layers shift under stress, causing repulsion between like charges | 硬而脆：在应力下发生层间滑移，导致同种电荷相互排斥

Lattice enthalpy is a measure of the strength of ionic bonding. The larger the lattice enthalpy, the stronger the ionic bond. Lattice enthalpy depends on ionic charge and ionic radius: higher charge and smaller radius lead to stronger attractions.

晶格焓是衡量离子键强度的指标。晶格焓越大，离子键越强。晶格焓取决于离子电荷和离子半径：电荷越高、半径越小，吸引力越强。

2. Covalent Bonding | 共价键

Covalent bonding involves the sharing of electron pairs between non-metal atoms. A single covalent bond consists of one shared pair, a double bond consists of two shared pairs, and a triple bond consists of three shared pairs. Bond strength increases with bond order, while bond length decreases.

共价键是非金属原子之间共享电子对形成的键。单键由一对共享电子组成，双键由两对共享电子组成，三键由三对共享电子组成。键强度随着键级增加而增大，键长则随之缩短。

For example, in nitrogen gas (N₂), the triple bond has a bond enthalpy of 945 kJ/mol, making it one of the strongest bonds in nature. Oxygen (O₂) has a double bond (498 kJ/mol), and fluorine (F₂) has a single bond (158 kJ/mol).

例如，氮气（N₂）中的三键键焓为945 kJ/mol，是自然界中最强的键之一。氧气（O₂）含有双键（498 kJ/mol），氟气（F₂）是单键（158 kJ/mol）。

N ≡ N (945 kJ/mol) > O = O (498 kJ/mol) > F − F (158 kJ/mol)

Coordinate (dative) covalent bonds occur when both shared electrons come from one atom. This is seen in the ammonium ion (NH₄⁺) and the hydronium ion (H₃O⁺).

配位共价键发生在一个原子提供两个共享电子的情况。这出现在铵离子（NH₄⁺）和水合氢离子（H₃O⁺）中。

3. Metallic Bonding | 金属键

Metallic bonding arises from the electrostatic attraction between a lattice of positive metal ions and delocalized valence electrons. These ‘sea of electrons’ are free to move throughout the structure, giving metals their characteristic properties.

金属键源于金属正离子晶格与离域价电子之间的静电吸引。这些“电子海”遍布整个结构，赋予金属特有的性质。

Properties explained by metallic bonding:

金属键可解释的性质：

Electrical conductivity – free electrons carry charge when a potential difference is applied | 导电性 – 在外加电势差时自由电子传导电荷
Thermal conductivity – delocalized electrons transfer kinetic energy quickly | 导热性 – 离域电子快速传递动能
Malleability – layers of cations can slide without breaking the metallic bond | 延展性 – 阳离子层可以滑动而不断裂金属键
High melting points – strong electrostatic forces between cations and electrons (except for mercury and the Group 1 metals) | 高熔点 – 阳离子与电子之间存在强静电作用（汞和第1族金属例外）

The strength of metallic bonding increases with the number of valence electrons and decreases with atomic radius. For example, magnesium (12 protons, 2 valence electrons per atom) has a higher melting point (650°C) than sodium (98°C), because Mg²⁺ has a higher charge and a smaller radius than Na⁺.

金属键强度随价电子数目增加而增强，随原子半径增大而减弱。例如，镁（12个质子，每个原子2个价电子）的熔点（650°C）高于钠（98°C），因为Mg²⁺比Na⁺电荷更高、半径更小。

4. Electronegativity and Bond Polarity | 电负性与键的极性

Electronegativity is the ability of an atom to attract bonding electrons toward itself. On the Pauling scale, fluorine (4.0) is the most electronegative element, followed by oxygen (3.5) and nitrogen (3.0).

电负性是原子吸引成键电子对自身的能力。在鲍林标度上，氟（4.0）电负性最强，其次是氧（3.5）和氮（3.0）。

When two atoms with different electronegativities form a covalent bond, the electron pair is unevenly shared, creating a polar bond. The atom with higher electronegativity acquires a partial negative charge (δ⁻), while the other acquires a partial positive charge (δ⁺).

当电负性不同的两个原子形成共价键时，电子对不均匀共享，产生极性键。电负性较高的原子获得部分负电荷（δ⁻），而另一个原子获得部分正电荷（δ⁺）。

ΔEN = 0 (non-polar) | 0 < ΔEN < 2.0 (polar covalent) | ΔEN > 2.0 (often ionic)

For example, HCl: ΔEN = 3.2 − 2.2 = 0.8, so the bond is polar covalent. The hydrogen atom is δ⁺ and the chlorine atom is δ⁻. This polarity is crucial for the reactivity of HCl in solution and for its solubility in water.

例如，HCl：ΔEN = 3.2 − 2.2 = 0.8，因此该键是极性共价键。氢原子为δ⁺，氯原子为δ⁻。这种极性对HCl在水溶液中的反应性和水溶性至关重要。

5. Molecular Polarity and Dipole Moments | 分子极性与偶极矩

A molecule is polar if it has an overall dipole moment, which depends on both bond polarity and molecular geometry. Symmetrical molecules can have polar bonds but zero net dipole moment because individual bond dipoles cancel out.

分子是否极性取决于总偶极矩，这受键的极性和分子几何构型共同影响。对称分子可能含有极性键，但因为各键偶极相互抵消，总偶极矩为零。

Key examples:

关键示例：

CO₂ – linear geometry, bond dipoles cancel, non-polar | CO₂ – 直线形，键偶极抵消，非极性
H₂O – bent geometry, bond dipoles do not cancel, polar | H₂O – 弯折形，键偶极不抵消，极性
CCl₄ – tetrahedral, bond dipoles cancel, non-polar | CCl₄ – 正四面体，键偶极抵消，非极性
NH₃ – trigonal pyramidal, lone pair creates asymmetry, polar | NH₃ – 三角锥形，孤对电子造成不对称，极性

To determine molecular polarity, first draw the Lewis structure, then apply VSEPR theory to determine geometry, and finally sum the individual bond dipoles considering symmetry.

要判断分子极性，首先画出路易斯结构，然后应用VSEPR理论确定几何构型，最后根据对称性将各键偶极矢量求和。

6. VSEPR Theory and Molecular Geometry | VSEPR理论与分子几何构型

The Valence Shell Electron Pair Repulsion (VSEPR) theory states that electron pairs around a central atom arrange themselves as far apart as possible to minimize repulsion. Both bonding pairs and lone pairs are considered, but lone pairs exert greater repulsion than bonding pairs.

价层电子对互斥理论（VSEPR）指出，中心原子周围的电子对会尽可能远离排列，以最小化排斥力。成键电子对和孤对电子都要考虑，但孤对电子产生的排斥力大于成键电子对。

Repulsion order:

排斥力顺序：

lone pair–lone pair > lone pair–bonding pair > bonding pair–bonding pair

Common geometries and bond angles you must memorize:

必须牢记的常见几何构型和键角：

Electron domains \| 电子域	Lone pairs \| 孤对电子	Geometry \| 几何构型	Bond angle \| 键角	Example \| 示例
2	0	Linear \| 直线形	180°	CO₂, BeCl₂
3	0	Trigonal planar \| 平面三角形	120°	BF₃
4	0	Tetrahedral \| 四面体	109.5°	CH₄
4	1	Trigonal pyramidal \| 三角锥形	107°	NH₃
4	2	Bent (V-shaped) \| 角形	104.5°	H₂O
5	0	Trigonal bipyramidal \| 三角双锥	90°, 120°	PCl₅
6	0	Octahedral \| 八面体	90°	SF₆

Remember that lone pairs compress bond angles because they occupy more space. For example, water’s bond angle is 104.5° instead of the ideal tetrahedral 109.5° due to two lone pairs on oxygen.

记住孤对电子会压缩键角，因为它们占据更大空间。例如，水的键角为104.5°，而不是理想四面体的109.5°，因为氧上有两对孤对电子。

7. Hybridization of Atomic Orbitals | 原子轨道杂化

Hybridization explains the shapes of molecules by mixing atomic orbitals to form new hybrid orbitals. Carbon in methane (CH₄) undergoes sp³ hybridization, forming four equivalent orbitals directed toward the corners of a tetrahedron.

杂化理论通过混合原子轨道形成新的杂化轨道，解释分子的形状。甲烷（CH₄）中的碳发生sp³杂化，形成四个等价的轨道，指向四面体的顶点。

Three main types of hybridization you need to know:

需要掌握的三种主要杂化方式：

sp – two linear hybrid orbitals, found in alkynes and compounds like BeCl₂ | sp杂化 – 两个直线形杂化轨道，见于炔烃和BeCl₂等化合物
sp² – three trigonal planar hybrid orbitals, found in alkenes and aromatic compounds | sp²杂化 – 三个平面三角形杂化轨道，见于烯烃和芳香化合物
sp³ – four tetrahedral hybrid orbitals, found in alkanes and compounds like NH₃ and H₂O | sp³杂化 – 四个四面体杂化轨道，见于烷烃以及NH₃和H₂O等化合物

For ethene (C₂H₄), each carbon is sp² hybridized. The unhybridized p orbitals overlap side-by-side to form the π bond. Similarly, in ethyne (C₂H₂), each carbon is sp hybridized, leaving two unhybridized p orbitals that form two perpendicular π bonds.

在乙烯（C₂H₄）中，每个碳为sp²杂化。未杂化的p轨道侧面重叠形成π键。类似地，在乙炔（C₂H₂）中，每个碳为sp杂化，剩余两个未杂化的p轨道形成两个相互垂直的π键。

8. Sigma and Pi Bonds | σ键与π键

Sigma (σ) bonds form when orbitals overlap head-on along the internuclear axis. Pi (π) bonds form when p orbitals overlap side-by-side, above and below the plane of the nuclei. Every covalent bond contains one σ bond; multiple bonds add π bonds on top.

σ键是轨道沿核间轴方向“头对头”重叠形成的。π键是p轨道“肩并肩”侧面重叠形成的，位于核平面上方和下方。每个共价键含有一个σ键；多重键则额外包含π键。

Properties of σ and π bonds:

σ键和π键的性质：

A single bond = 1 σ bond | 单键 = 1个σ键
A double bond = 1 σ + 1 π | 双键 = 1个σ键 + 1个π键
A triple bond = 1 σ + 2 π | 三键 = 1个σ键 + 2个π键

π bonds are weaker than σ bonds because the side-by-side overlap is less effective. This explains why alkenes undergo addition reactions: the π bond breaks more easily. For example, ethene reacts with bromine water, decolorizing from orange to colorless, while ethane does not react under normal conditions.

π键弱于σ键，因为侧面重叠效率较低。这解释了烯烃为何容易发生加成反应：π键更容易断裂。例如，乙烯使溴水由橙色褪为无色，而乙烷在相同条件下不反应。

9. Intermolecular Forces | 分子间作用力

Intermolecular forces are weaker than intramolecular bonds but determine physical properties such as boiling point, vapor pressure, and solubility. The three main types are London dispersion forces, dipole–dipole interactions, and hydrogen bonding.

分子间作用力弱于分子内化学键，但决定熔沸点、蒸气压和溶解性等物理性质。三种主要类型是伦敦色散力、偶极–偶极相互作用和氢键。

London dispersion forces (instantaneous induced dipoles) exist between all molecules, and increase with the number of electrons. Larger molecules and longer chains have greater surface area, leading to stronger dispersion forces. For example, the boiling point increases from CH₄ (−164°C) to C₄H₁₀ (−0.5°C).

伦敦色散力（瞬时感应偶极）存在于所有分子之间，随电子数目增加而增强。更大分子和更长链具有更大表面积，色散力更强。例如，沸点从CH₄（−164°C）升高到C₄H₁₀（−0.5°C）。

Dipole–dipole interactions occur between polar molecules. The partial positive end of one molecule attracts the partial negative end of another. These are stronger than dispersion forces but weaker than hydrogen bonds.

偶极–偶极相互作用发生在极性分子之间。一个分子的部分正端吸引另一个分子的部分负端。这类作用强于色散力，但弱于氢键。

Hydrogen bonds form when hydrogen is covalently bonded to a highly electronegative atom (N, O, or F) and is attracted to another electronegative atom. Hydrogen bonds are the strongest intermolecular force, explaining water’s anomalously high boiling point (100°C) compared to H₂S (−60°C).

氢键在氢与高电负性原子（N、O或F）共价结合，并受到另一个电负性原子吸引时形成。氢键是最强的分子间作用力，解释了水（100°C）与H₂S（−60°C）相比沸点异常升高。

10. Resonance and Delocalization | 共振与离域

Resonance occurs when a molecule or ion can be represented by two or more Lewis structures that differ only in the position of electrons. The actual structure is a hybrid, with bond lengths intermediate between the contributing structures.

共振发生在分子或离子可以用两个或更多仅电子位置不同的路易斯结构表示时。实际结构是这些结构之间的杂化体，键长介于各贡献结构之间。

Classic examples in IB Chemistry:

IB化学中的经典例子：

Ozone (O₃) — both O–O bonds are identical, with a bond order of 1.5 | 臭氧（O₃） — 两个O–O键完全相同，键级为1.5
Carbonate ion (CO₃²⁻) — each C–O bond has partial double bond character | 碳酸根离子（CO₃²⁻） — 每个C–O键具有部分双键特征
Benzene (C₆H₆) — six equal C–C bonds, all 140 pm, between single (154 pm) and double (134 pm) | 苯（C₆H₆） — 六个等长的C–C键均为140 pm，介于单键（154 pm）和双键（134 pm）之间
Nitrate ion (NO₃⁻) — all N–O bonds are equivalent | 硝酸根离子（NO₃⁻） — 所有N–O键等价

Resonance stabilizes molecules and ions through delocalization of electrons. Benzene’s delocalized π electron cloud accounts for its unusual stability and its preference for substitution reactions over addition reactions.

共振通过电子离域使分子和离子稳定。苯的离域π电子云解释了其异常稳定性和偏好取代反应而非加成反应的特征。

11. VSEPR vs. Hybridization: Applying Both | VSEPR与杂化：综合应用

In IB exams, you may be asked to determine the geometry, hybridization, and polarity of a molecule. The steps are straightforward once you practice.

在IB考试中，你可能需要确定分子的几何构型、杂化方式和极性。一旦练习熟练，这些步骤就非常简单。

Step-by-step approach:

分步方法：

Draw the Lewis structure, counting all valence electrons | 画出路易斯结构，计算所有价电子
Count electron domains around the central atom | 数中心原子周围的电子域
Apply VSEPR to predict geometry | 应用VSEPR预测几何构型
Deduce hybridization from the number of electron domains | 根据电子域数目推断杂化方式
Check molecular polarity by evaluating symmetry | 通过对称性判断分子极性

Worked example: XeF₄. Xenon has 8 valence electrons, each fluorine contributes 1, giving a total of 12 valence electrons. The central Xe is surrounded by 4 bonding pairs and 2 lone pairs (6 electron domains). The geometry is square planar, and the hybridization is sp³d². Despite the polar Xe–F bonds, the molecule is non-polar due to symmetry.

实例：XeF₄。氙有8个价电子，每个氟贡献1个，总共12个价电子。中心Xe周围有4个成键电子对和2个孤对电子（6个电子域）。几何构型为平面正方形，杂化为sp³d²。尽管Xe–F键是极性的，但由于对称性，分子为非极性。

12. Common Exam Pitfalls and Tips | 常见考试陷阱与技巧

Many students lose marks on bonding questions due to avoidable errors. Here are the most frequent weaknesses and how to fix them.

许多学生在键合问题上因可避免的错误而失分。以下是最常见的弱点和解决方法。

Confusing intermolecular forces with intramolecular bonds — hydrogen bonding is an intermolecular force, NOT a covalent bond. Always describe the type of particle and the force BETWEEN particles. | 混淆分子间作用力与分子内化学键 — 氢键是分子间作用力，不是共价键。始终说明粒子的类型以及粒子之间的作用力。
Forgetting lone pairs — when determining molecular shape, you must count lone pairs. NH₃ is trigonal pyramidal, not trigonal planar. | 忘记孤对电子 — 判断分子形状时，必须计算孤对电子。NH₃是三角锥形，而不是平面三角形。
Assuming diatomic molecules are always polar — homonuclear diatomics like Cl₂ are non-polar. Heteronuclear diatomics like HCl are polar. | 假设双原子分子总是极性的 — 同核双原子分子如Cl₂是非极性的；异核双原子分子如HCl是极性的。
Incorrect bond angle for central atoms — if the central atom has no lone pairs (e.g., BF₃, CH₄), the angles equal the ideal values (120°, 109.5°). With lone pairs, angles are reduced. | 中心原子键角错误 — 若中心原子无孤对电子（如BF₃、CH₄），键角等于理想值（120°、109.5°）；有孤对电子时键角减小。
Mixing up ionic and covalent character — many bonds have partial ionic character. Use ΔEN to classify, but remember that all polar covalent bonds have some ionic character. | 混淆离子性和共价性 — 很多键具有部分离子性。用ΔEN分类，但记住所有极性共价键都有一定离子性。

Finally, practice drawing molecular geometries in 3D: use wedge-and-dash diagrams to show depth. IB examiners award marks for correct notation, so always show lone pairs and label bond angles when required.

最后，练习用3D视角画出分子几何构型：使用楔形线和虚线表示空间深度。IB考官根据规范记法给分，所以务必画出孤对电子并按题目要求标注键角。

This review covers the core concepts of chemical bonding and structure required for IB Chemistry. Master the fundamental rules, practice past-paper questions, and you will be well prepared to tackle bonding-related problems in your exams.

本复习涵盖IB化学所需的化学键与物质结构核心概念。掌握基本规则，练习往年真题，你将足以从容应对考试中的键合相关问题。

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IB Chemistry: Chemical Bonding and Structure Core Concepts | IB化学：化学键与物质结构核心概念