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Covalent Bonding and Electron Domains: Molecular Geometries (HL) | 共价键与电子域：分子几何构型 (HL)

2026-08-03 00:00 · 本文部分内容由AI辅助创作

📚 Covalent Bonding and Electron Domains: Molecular Geometries (HL) | 共价键与电子域：分子几何构型 (HL)

Covalent bonding is a fundamental concept in chemistry, describing how atoms share electrons to achieve a stable electron configuration. In the IB Chemistry Higher Level (HL) curriculum, students explore advanced topics such as sigma and pi bonds, formal charge, resonance, and the prediction of molecular shapes using VSEPR theory. Understanding electron domains—both bonding and non-bonding pairs—is key to determining three-dimensional molecular geometries, which in turn influence physical and chemical properties.

共价键是化学中的基本概念，描述了原子通过共享电子以达到稳定电子构型的方式。在IB化学高级课程（HL）中，学生将深入学习σ键与π键、形式电荷、共振以及运用VSEPR理论预测分子形状等高级主题。理解电子域——成键电子对和非成键电子对——是确定三维分子几何构型的关键，而分子几何构型又影响着物理和化学性质。

1. Covalent Bond Formation and Nature | 共价键的形成与本质

A covalent bond forms when two atoms, typically nonmetals, share one or more pairs of electrons. The shared electrons are simultaneously attracted to both nuclei, creating a strong, directional bond that holds the atoms together. The bond arises from the overlap of atomic orbitals, which concentrates electron density in the internuclear region. The strength of a covalent bond is measured by its bond enthalpy, and the distance between the two nuclei is the bond length. Multiple bonds (double or triple) involve sharing more electron pairs, leading to shorter and stronger bonds compared to single bonds.

共价键在通常是非金属的两个原子共享一对或多对电子时形成。共享的电子同时受到两个原子核的吸引，产生一个强大的方向性键，将原子结合在一起。该键源于原子轨道的重叠，使电子密度集中在核间区域。共价键的强度由键焓衡量，两个原子核之间的距离为键长。多重键（双键或三键）涉及共享更多电子对，因此比单键更短、更强。

2. Sigma and Pi Bonds | σ键与π键 (HL)

A sigma (σ) bond results from the end-to-end overlap of atomic orbitals along the internuclear axis. This cylindrical symmetry allows free rotation of the bonded atoms around the bond. A pi (π) bond is formed by the sideways overlap of two parallel p orbitals, creating electron density above and below the bond axis; this restricts rotation. A single bond is always a σ bond. A double bond contains one σ bond and one π bond. A triple bond consists of one σ bond and two π bonds. The presence of π bonds explains the rigidity and planar nature of molecules like ethene (C₂H₄) and the linear shape of ethyne (C₂H₂).

σ键由原子轨道沿核间轴头对头重叠而成。这种圆柱对称性允许键合原子绕键自由旋转。π键由两个平行p轨道肩并肩重叠形成，在键轴上下产生电子密度，从而限制了旋转。单键总是σ键。双键包含一个σ键和一个π键。三键包含一个σ键和两个π键。π键的存在解释了诸如乙烯（C₂H₄）的刚性和平面性，以及乙炔（C₂H₂）的线性形状。

3. Formal Charge | 形式电荷 (HL)

Formal charge (FC) is a bookkeeping tool used to assess the most plausible Lewis structure among several possibilities. It is calculated for each atom as:

FC = V − L − ½B

where V is the number of valence electrons on the free atom, L is the number of lone pair electrons belonging to the atom in the structure, and B is the total number of bonding electrons shared with other atoms. The most stable Lewis structure minimizes the sum of absolute formal charges, and any negative formal charge should reside on the most electronegative atom. For example, in the sulfate ion (SO₄²⁻), assigning formal charges helps select the dominant resonance contributor where sulfur carries a +2 charge and each oxygen −1, but resonance delocalisation reduces real charges.

形式电荷（FC）是一种用于从多个可能的路易斯结构中评估最合理结构的记录工具。每个原子的计算式为：

FC = V − L − ½B

其中V是自由原子的价电子数，L是该原子在结构中的孤对电子数，B是与其它原子共享的成键电子总数。最稳定的路易斯结构使形式电荷的绝对值之和最小，并且任何负形式电荷应位于电负性最大的原子上。例如，在硫酸根离子（SO₄²⁻）中，分配形式电荷有助于选出硫带+2、每个氧带−1的主导共振贡献体，但共振离域降低了实际电荷。

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4. Resonance Structures | 共振结构 (HL)

Some molecules and polyatomic ions cannot be adequately described by a single Lewis structure. When two or more valid Lewis structures differ only in the arrangement of electrons (not atoms), the species exhibits resonance. The actual electronic structure is a resonance hybrid, an average of all contributing forms, with electrons delocalised over the molecule. Consequently, bond lengths and strengths become equalised. Classic examples include ozone (O₃), with one double and one single bond in each canonical form but two identical bonds in reality, and the carbonate ion (CO₃²⁻), where all C–O bonds are equivalent. Benzene (C₆H₆) is another prime example: two Kekulé structures with alternating double bonds combine to a hybrid with equivalent C–C bonds and a delocalised π electron cloud.

有些分子和多原子离子无法用单一路易斯结构充分描述。当两个或多个有效的路易斯结构仅在电子排布（而非原子位置）上不同时，该物种表现出共振。真实的电子结构是一个共振杂化体，是所有贡献形式的平均值，电子在整个分子上离域。因此，键长和键强变得均等。经典例子包括臭氧（O₃），在每个正则式中有一个双键和一个单键，但实际上有两个相同的键；以及碳酸根离子（CO₃²⁻），所有C–O键都是等价的。苯（C₆H₆）是另一个典型例子：两个交替双键的凯库勒结构结合成键长等同、π电子云离域的杂化体。

5. VSEPR Theory Basics | 价层电子对互斥理论基础

Valence Shell Electron Pair Repulsion (VSEPR) theory enables the prediction of three-dimensional molecular shapes from a Lewis structure. The central principle is that electron pairs around a central atom—whether in bonds or as lone pairs—repel each other and arrange themselves as far apart as possible to minimise repulsion. Both bonding pairs and non-bonding (lone) pairs count as electron domains. The arrangement of electron domains determines the electron geometry, whereas the molecular geometry (the actual shape of the molecule) is named based on the positions of atoms only. The repulsion strengths follow the order: lone pair–lone pair > lone pair–bonding pair > bonding pair–bonding pair, because lone pairs occupy a larger volume of space close to the nucleus.

价层电子对互斥理论（VSEPR）使我们能够由路易斯结构预测分子的三维形状。核心原理是，中心原子周围的电子对——无论是成键的还是孤对电子——彼此排斥，并尽可能远离以最小化排斥。成键电子对和非键（孤对）电子对都算作电子域。电子域的排列决定电子几何构型，而分子几何构型（分子的实际形状）仅根据原子的位置来命名。排斥力的大小顺序为：孤对-孤对 > 孤对-成键 > 成键-成键，因为孤对电子在靠近原子核处占据更大的空间体积。

6. Electron Domains: Types and Counting | 电子域：类型与计数

n. Therefore, to count electron domains, simply add the number of atoms bonded to the central atom (each bond counts as one) and the number of lone pairs on the central atom. For instance, in water (H₂O), oxygen is bonded to two hydrogen atoms and carries two lone pairs → four electron domains. In carbon dioxide (CO₂), carbon forms two double bonds (each counts as one domain) and has no lone pairs → two electron domains. This counting is the first step in applying VSEPR theory.

电子域是中心原子周围任何高电子密度区域。每个单键、双键或三键都视作一个电子域，因为多重键中的电子对大致占据相同的空间区域。一个孤对电子也构成一个电子域。因此，要计算电子域，只需将中心原子连接的原子数（每个键算作一个）与中心原子上的孤对电子数相加。例如，在水（H₂O）中，氧连接两个氢原子并带有两个孤对电子 → 四个电子域。在二氧化碳（CO₂）中，碳形成两个双键（每个算作一个域）且没有孤对电子 → 两个电子域。这种计数是应用VSEPR理论的第一步。

7. Geometries with 2–4 Electron Domains | 2–4电子域的几何构型

The following summarises the molecular geometries for two, three, and four electron domains. The examples illustrate both electron geometry and the resulting molecular shape.

以下总结了两个、三个和四个电子域的分子几何构型。示例同时说明了电子构型和最终的分子形状。