📚 Covalent Bonding in IB Chemistry | IB 化学：共价键 考点精讲
Covalent bonding lies at the heart of molecular chemistry. In IB Chemistry, understanding how atoms share electrons to form molecules, predict shapes, and explain properties is essential for both Standard Level (SL) and Higher Level (HL) students. This article breaks down every key concept you need to master, from Lewis structures and VSEPR theory to hybridization and delocalized bonding.
共价键是分子化学的核心。在 IB 化学中，无论是标准级别（SL）还是高级别（HL），理解原子如何通过共用电子形成分子、预测形状并解释性质都是必须掌握的关键。本文将深入剖析你需要掌握的每一个重要概念，从路易斯结构和 VSEPR 理论到杂化与离域键。
1. What Is a Covalent Bond? | 什么是共价键？
A covalent bond forms when two atomic nuclei simultaneously attract a shared pair of electrons. This electrostatic attraction between the positively charged nuclei and the negatively charged shared electrons holds the atoms together in a molecule or polyatomic ion. Unlike ionic bonding, which involves electron transfer, covalent bonding arises from electron sharing, typically between non‑metals with similar electronegativity values.
共价键形成于两个原子核同时吸引一对共用电子之时。带正电的原子核与带负电的共用电子对之间的静电引力将原子结合在分子或多原子离子中。与涉及电子转移的离子键不同，共价键源于电子的共用，通常发生在电负性值相近的非金属原子之间。
2. The Electrostatic Nature and Bond Strength | 静电本质与键的强度
Although often described as electron sharing, a covalent bond is fundamentally an electrostatic interaction. The shared pair is more likely to be found in the region between the two nuclei, creating a region of high electron density that both nuclei are drawn toward. This electrostatic attraction determines bond length and bond strength: shorter bonds are generally stronger, as the shared electrons are held more tightly. In IB, you may need to compare bond enthalpies and bond lengths for single, double, and triple bonds.
虽然常被描述为电子共用，但共价键本质上是静电相互作用。共用电子对更可能出现在两个原子核之间的区域，形成一个高电子密度的区域，两个原子核都受到该区域的吸引。这种静电引力决定了键长和键的强度：较短通常更强，因为共用电子被拉得更紧。在 IB 中，你需要会比较单键、双键和三键的键焓与键长。
3. Lewis (Electron Dot) Structures | 路易斯（电子点式）结构
Lewis structures represent the valence electrons of atoms within a molecule. Dots represent electrons, and lines represent shared pairs (single, double, or triple bonds). To draw a Lewis structure:
- Count total valence electrons from all atoms; adjust for charge on ions.
- Draw a skeleton with the least electronegative atom as the centre (except hydrogen).
- Place single bonds first, then distribute remaining electrons as lone pairs to satisfy the octet rule.
- Form multiple bonds if needed to complete octets for central atoms (especially C, N, O).
路易斯结构表示分子中原子的价电子。点代表电子，线代表共用电子对（单、双或三键）。绘制路易斯结构的步骤：
- 计算所有原子的总价电子数；对于离子要调整电荷。
- 画出骨架，将电负性最低的原子作为中心原子（氢除外）。
- 先放置单键，然后分配剩余电子作为孤对电子以满足八隅体规则。
- 如有需要，形成多重键来完成中心原子（特别是 C、N、O）的八隅体。
Example: CO₂ → O=C=O with two double bonds and no lone pairs on carbon.
例如：CO₂ → O=C=O，碳上无双键，无孤对电子。
4. Formal Charge and the Most Stable Lewis Structure | 形式电荷与最稳定路易斯结构
Formal charge helps you decide which of several possible Lewis structures is most plausible. The formula is:
FC = V − N − ½B
where V = number of valence electrons of the free atom, N = number of non‑bonding (lone pair) electrons, and B = total number of bonding electrons shared. The best structure minimizes formal charges on all atoms, placing any negative formal charge on the more electronegative element. HL students are expected to use formal charge routinely.
形式电荷帮助你判断几种可能的路易斯结构中哪一种最合理。公式为：
FC = V − N − ½B
其中 V = 自由原子的价电子数，N = 非键（孤对）电子数，B = 共享的键电子总数。最佳结构使所有原子的形式电荷最小化，并将负的形式电荷放在电负性更高的元素上。HL 学生需要熟练运用形式电荷。
5. Resonance and Delocalisation | 共振与离域
When a molecule or ion can be represented by two or more valid Lewis structures that differ only in the distribution of electrons (not atom positions), we say it has resonance. The true structure is a resonance hybrid — a blend of the contributing forms. In the carbonate ion, CO₃²⁻, the three C–O bonds are identical and intermediate in length between single and double bonds because the pi electrons are delocalised over the entire ion. IB students should draw resonance structures using double‑headed arrows (↔) and describe delocalisation as sharing of electrons across more than two nuclei.
当一个分子或离子可以用两个或多个仅在电子分布上不同（原子位置不变）的有效路易斯结构表示时，我们称其存在共振。真实结构是共振杂化体 —— 所有参与结构的混合。在碳酸根离子 CO₃²⁻ 中，三条 C–O 键完全相同，键长介于单键和双键之间，因为 π 电子离域扩展到整个离子。IB 学生应使用双向箭头（↔）画出共振结构，并将离域描述为电子在超过两个原子核之间的共享。
6. VSEPR Theory: Electron Domain Geometry | VSEPR 理论：电子域几何构型
Valence Shell Electron Pair Repulsion (VSEPR) theory states that electron domains around a central atom arrange themselves to minimise repulsion. An electron domain can be a single bond, double bond, triple bond, or a lone pair — each counts as one domain. The repulsion strength follows: lone pair–lone pair > lone pair–bonding pair > bonding pair–bonding pair. This determines the basic electron‑domain geometry, upon which molecular shape is based.
价层电子对互斥（VSEPR）理论指出，中心原子周围的电子域会自行排列以最小化排斥。电子域可以是单键、双键、三键或孤对电子 —— 每个都计为一个域。排斥强度顺序为：孤对–孤对 > 孤对–键对 > 键对–键对。这决定了基本的电子域几何构型，分子形状在此基础上得出。
| Electron Domains | Geometry | Bond Angle |
|---|---|---|
| 2 | Linear | 180° |
| 3 | Trigonal planar | 120° |
| 4 | Tetrahedral | 109.5° |
| 5 | Trigonal bipyramidal | 90°, 120° |
| 6 | Octahedral | 90°, 180° |
电子域数量、基本几何构型和理想键角如上表所示。
7. Molecular Shapes and the Effect of Lone Pairs | 分子形状与孤对电子的影响
When lone pairs are present, the actual molecular shape differs from the electron‑domain geometry because lone pairs are not ‘seen’ in the final molecular skeleton. Examples include:
- 3 bonding domains + 1 lone pair → trigonal pyramidal (e.g., NH₃, bond angle ~107°)
- 2 bonding domains + 2 lone pairs → bent (e.g., H₂O, bond angle ~104.5°)
- 4 bonding domains + 2 lone pairs → square planar (e.g., XeF₄, with two lone pairs opposite each other in an octahedral arrangement)
当存在孤对电子时，实际的分子形状与电子域几何构型不同，因为孤对电子在最终的分子骨架中不被“看见”。例子包括：
- 3 个键域 + 1 个孤对 → 三角锥形（如 NH₃，键角约 107°）
- 2 个键域 + 2 个孤对 → V 形（如 H₂O，键角约 104.5°）
- 4 个键域 + 2 个孤对 → 平面正方形（如 XeF₄，在八面体排布中对位两个孤对）
IB requires you to predict and draw shapes for both SL (up to 4 electron domains) and HL (5 and 6 domains). Always indicate bond angles and name the shape.
IB 要求你预测并画出 SL（最多 4 个电子域）和 HL（5 和 6 个域）的形状。务必要标明键角并说出形状名称。
8. Bond Polarity and Molecular Polarity | 键极性与分子极性
A bond is polar if the two atoms have different electronegativity values; the more electronegative atom pulls electron density toward itself, creating a dipole with partial charges (δ⁺ and δ⁻). However, a molecule with polar bonds can be non‑polar if the dipoles cancel due to symmetry. Carbon dioxide (CO₂) is linear and non‑polar, while water (H₂O) is bent and polar. Molecular polarity influences physical properties like solubility and boiling point.
如果两个原子的电负性值不同，键具有极性；电负性更高的原子将电子密度拉向自身，产生具有部分电荷（δ⁺ 和 δ⁻）的偶极。然而，含有极性键的分子如果由于对称性导致偶极相互抵消，仍可以是非极性的。二氧化碳（CO₂）呈线性非极性，而水（H₂O）呈 V 形极性。分子极性影响着溶解度和沸点等物理性质。
9. Sigma and Pi Bonds (HL) | σ 键与 π 键（HL）
HL students must distinguish between sigma (σ) and pi (π) bonds. A sigma bond forms by head‑on overlap of orbitals (s–s, s–p, p–p end‑on, or hybrid orbitals). It is the first bond between two atoms and has electron density concentrated along the internuclear axis. A pi bond results from sideways overlap of p orbitals; its electron density lies above and below the bond axis. All single bonds are σ; a double bond consists of one σ and one π; a triple bond consists of one σ and two π bonds.
HL 学生必须区分 σ 键和 π 键。σ 键由轨道头对头重叠形成（s–s、s–p、p–p端接或杂化轨道）。它是两个原子间的第一条键，电子密度集中在核间轴上。π 键由 p 轨道肩并肩重叠产生，其电子密度位于键轴的上下方。所有单键均为 σ；双键包含一个 σ 和一个 π；三键包含一个 σ 和两个 π。
10. Hybridisation (HL) | 杂化（HL）
Hybridisation describes the mixing of atomic orbitals to form equivalent hybrid orbitals that point in specific geometries. For a central atom:
- sp hybridisation → linear, 180° (e.g., BeCl₂, CO₂ carbon)
- sp² hybridisation → trigonal planar, 120° (e.g., BF₃, C₂H₄ carbon)
- sp³ hybridisation → tetrahedral, 109.5° (e.g., CH₄, NH₃, H₂O)
- sp³d → trigonal bipyramidal; sp³d² → octahedral
IB HL questions may ask you to identify the hybridisation of a given atom based on the number of electron domains or to explain why hybrid orbitals form.
杂化描述的是原子轨道混合形成指向特定几何形状的等价杂化轨道。对于中心原子：
- sp 杂化 → 直线形，180°（如 BeCl₂、CO₂ 中的碳）
- sp² 杂化 → 三角平面形，120°（如 BF₃、C₂H₄ 中的碳）
- sp³ 杂化 → 四面体形，109.5°（如 CH₄、NH₃、H₂O）
- sp³d → 三角双锥形；sp³d² → 八面体形
IB HL 题目可能要求你根据电子域数识别给定原子的杂化类型，或解释为何形成杂化轨道。
11. Exceptions to the Octet Rule | 八隅体规则的例外
Not all molecules obey the octet rule. Common exceptions tested in IB include:
- Electron‑deficient molecules: BeCl₂ and BF₃ have fewer than 8 electrons around the central atom.
- Expanded octet: elements from Period 3 onwards (P, S, Cl, etc.) can accommodate more than 8 electrons due to available d orbitals. Examples are PCl₅, SF₆, and SO₄²⁻.
- Odd‑electron species (radicals): NO and NO₂ have an unpaired electron.
For HL, you must rationalise expanded octets using d‑orbital availability and formal charge arguments.
并非所有分子都遵守八隅体规则。IB 中常考的例外包括：
- 缺电子分子：BeCl₂ 和 BF₃ 中心原子周围少于 8 个电子。
- 扩展八隅体：第三周期及以后的元素（P、S、Cl 等）因有可用的 d 轨道，可容纳超过 8 个电子。例如 PCl₅、SF₆ 和 SO₄²⁻。
- 奇电子物种（自由基）：NO 和 NO₂ 具有未成对电子。
对于 HL，你必须用 d 轨道的可用性和形式电荷论证来解释扩展八隅体。
12. Linking Structure to Properties | 从结构到性质
Bonding theory directly explains observable properties. Giant covalent structures like diamond (sp³ network) and graphite (sp² sheets with π delocalisation) demonstrate very different hardness, electrical conductivity, and melting points. Molecular polarity dictates solubility: ‘like dissolves like’. Intermolecular forces, including London dispersion forces, dipole–dipole interactions, and hydrogen bonding, are understood only after you master covalent bond polarity and molecular shape. Mastery of these concepts gives you the ability to predict and compare boiling points, volatility, and electrical conductivity across a range of substances.
键合理论直接解释了可观察的性质。巨型共价结构如金刚石（sp³ 网络结构）和石墨（sp² 层状结构，π 电子离域）展示了截然不同的硬度、导电性和熔点。分子极性决定了溶解度：“相似相溶”。包括伦敦色散力、偶极–偶极相互作用和氢键在内的分子间作用力，只有在你掌握共价键极性和分子形状之后才能被真正理解。熟练掌握这些概念，你就能预测和比较多种物质的沸点、挥发性及导电性。
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