CIE A-Level Chemistry Molecular Shapes and Geometry Guide — A-Level 化学：分子形状与几何构型解析

一、价层电子对互斥理论：分子形状的核心原理 | VSEPR Theory: The Core Principle Behind Molecular Shapes

在 CIE A-Level 化学中，预测分子形状最常用、也是考试必考的工具就是价层电子对互斥理论（Valence Shell Electron Pair Repulsion，简称 VSEPR）。这个理论的核心思想非常朴素：分子中心原子周围的电子对彼此带负电荷，负电荷之间相互排斥，因此电子对会尽可能彼此远离，使排斥力降到最低。分子的实际形状，就是电子对在三维空间”尽量分开”之后所呈现的排布方式。

In CIE A-Level Chemistry, the most frequently used and exam-required tool for predicting molecular shapes is the Valence Shell Electron Pair Repulsion theory, abbreviated as VSEPR. The core idea of this theory is simple: electron pairs around the central atom of a molecule all carry negative charge, and negative charges repel each other. Therefore, electron pairs arrange themselves as far apart as possible to minimise repulsion. The actual shape of a molecule is the three-dimensional arrangement that results when electron pairs “spread out as much as they can”.

理解 VSEPR 理论时，最关键的一步是分清”电子对排布”和”分子形状”这两个概念。电子对排布描述的是中心原子周围所有电子对（包括成键电子对和孤对电子）的空间位置；而分子形状只描述原子的相对位置，即只考虑成键电子对连接出来的原子骨架。例如水分子的电子对排布是四面体形，但由于只有两对是成键电子对，水的分子形状是弯曲形（V 形）。这个区别是考试中最常设置的陷阱之一。

When understanding VSEPR theory, the most critical step is to distinguish between “electron pair arrangement” and “molecular shape”. Electron pair arrangement describes the positions of all electron pairs around the central atom, including both bonding pairs and lone pairs; molecular shape describes only the relative positions of atoms, that is, the atomic skeleton formed by bonding pairs alone. For example, the electron pair arrangement of a water molecule is tetrahedral, but because only two of the pairs are bonding pairs, the molecular shape of water is bent (V-shaped). This distinction is one of the most common traps set in exams.

VSEPR 理论还给出了一个实用的预测流程：先画出中心原子的路易斯结构（Lewis structure），数出中心原子周围的电子对总数；再判断其中有几对是成键电子对、几对是孤对电子；最后根据电子对总数确定空间排布，再根据孤对电子数目确定实际分子形状。CIE 考卷中凡是涉及”预测形状并解释原因”的题目，几乎都可以用这个四步流程完成。

VSEPR theory also provides a practical prediction procedure: first draw the Lewis structure of the central atom and count the total number of electron pairs around it; then determine how many are bonding pairs and how many are lone pairs; next use the total number of electron pairs to determine the arrangement, and finally use the number of lone pairs to determine the actual molecular shape. Almost every CIE exam question that asks you to “predict the shape and explain your reasoning” can be completed using this four-step procedure.

二、成键电子对与孤对电子：两种电子对如何决定空间排布 | Bonding Pairs vs Lone Pairs: How Two Types of Electron Pairs Determine Geometry

中心原子周围的电子对分为两大类：成键电子对（bonding pairs）和孤对电子（lone pairs）。成键电子对是两个原子共享的电子对，它们同时受到两个原子核的吸引，因此”活动空间”比较集中，占据的空间体积相对较小。孤对电子只属于中心原子本身，只受到一个原子核的吸引，因此电子云更加弥散，占据的空间更大。

There are two types of electron pairs around a central atom: bonding pairs and lone pairs. A bonding pair is a pair of electrons shared between two atoms; it is attracted by two nuclei simultaneously, so its “activity space” is concentrated and it occupies a relatively small volume. A lone pair belongs only to the central atom and is attracted by a single nucleus, so its electron cloud is more diffuse and occupies a larger space.

正因为孤对电子占据的空间更大，孤对电子对邻近电子对的排斥力也更强。排斥力的大小排序是：孤对电子-孤对电子（lp-lp）大于孤对电子-成键电子对（lp-bp），大于成键电子对-成键电子对（bp-bp）。这一排斥力排序是整个 VSEPR 理论预测键角的基础，也是解释氨、水键角为什么小于甲烷键角的关键。

Because lone pairs occupy more space, they exert stronger repulsion on neighbouring electron pairs. The order of repulsion strength is: lone pair-lone pair (lp-lp) greater than lone pair-bonding pair (lp-bp), which is greater than bonding pair-bonding pair (bp-bp). This repulsion order is the foundation of all VSEPR bond angle predictions and is the key to explaining why ammonia and water have smaller bond angles than methane.

在 CIE 考试中，解释形状变化时一定要写清楚两层意思：第一，孤对电子占据更大空间、排斥力更强；第二，更强的排斥力把成键电子对”挤”得更近，导致键角减小。只写”因为有孤对电子所以键角变小”而不说明排斥力排序，通常只能得到一半分数。把 lp-lp 大于 lp-bp 大于 bp-bp 这个排序写出来，是拿满解释分的标准写法。

In CIE exams, when explaining shape changes you must write two layers of reasoning: first, lone pairs occupy more space and exert stronger repulsion; second, the stronger repulsion pushes the bonding pairs closer together, reducing the bond angle. Writing only “there is a lone pair so the bond angle is smaller” without mentioning the repulsion order usually earns only half marks. Stating the order lp-lp greater than lp-bp greater than bp-bp is the standard way to score full marks on explanations.

三、直线形与平面三角形：两对和三对电子对的空间构型 | Linear and Trigonal Planar: Two and Three Electron Domains

当中心原子周围只有两对电子对时，两对电子对会尽量远离，彼此夹角为 180 度，分子呈直线形（linear）。典型例子是二氧化碳 CO2 和氯化铍 BeCl2。二氧化碳分子中碳原子与两个氧原子各形成双键，双键仍按一对电子对处理，因此 CO2 是直线形分子，键角 180 度。这里要注意：无论成键是单键、双键还是三键，在 VSEPR 计数时都只算作一对电子对。

When there are only two electron pairs around the central atom, the two pairs spread as far apart as possible with a 180-degree angle between them, giving a linear shape. Typical examples are carbon dioxide CO2 and beryllium chloride BeCl2. In carbon dioxide, the carbon atom forms a double bond with each oxygen atom; each double bond still counts as one electron pair, so CO2 is linear with a 180-degree bond angle. Note that single, double and triple bonds all count as one electron pair in VSEPR counting.

当中心原子周围有三对电子对时，三对电子对在同一平面内彼此相隔 120 度排布，形成平面三角形（trigonal planar）。典型例子是三氟化硼 BF3 和三氯化硼 BCl3。BF3 中硼原子只有三对成键电子对，没有孤对电子，所以硼的电子对排布和分子形状都是平面三角形，键角 120 度。

When there are three electron pairs around the central atom, the three pairs lie in the same plane, 120 degrees apart, forming a trigonal planar arrangement. Typical examples are boron trifluoride BF3 and boron trichloride BCl3. In BF3, the boron atom has only three bonding pairs and no lone pairs, so both its electron pair arrangement and molecular shape are trigonal planar with 120-degree bond angles.

如果三对电子对中有一对是孤对电子，分子形状就变成弯曲形（bent 或 V 形）。典型例子是二氧化硫 SO2 和二氧化氮 NO2。SO2 中硫原子周围有三对电子对，其中两对是成键电子对，一对是孤对电子。孤对电子的排斥使 O-S-O 键角从 120 度略微压缩到约 119 度，分子呈弯曲形。这类”电子对排布与分子形状不同”的例子，是 CIE 选择题的常客。

If one of the three electron pairs is a lone pair, the molecular shape becomes bent (V-shaped). Typical examples are sulfur dioxide SO2 and nitrogen dioxide NO2. In SO2, the sulfur atom has three electron pairs, two bonding pairs and one lone pair. The lone pair repulsion compresses the O-S-O bond angle slightly from 120 degrees to about 119 degrees, giving a bent shape. Examples where “electron pair arrangement differs from molecular shape” are frequent guests in CIE multiple-choice questions.

四、四面体构型：甲烷与氨和水的关键对比 | Tetrahedral Geometry: Methane vs Ammonia vs Water

四面体（tetrahedral）是 A-Level 化学中出现频率最高的空间构型。当中心原子周围有四对电子对且全部是成键电子对时，四对电子对在三维空间中以 109.5 度的夹角彼此分开，形成正四面体。最典型的例子是甲烷 CH4。碳原子周围有四对成键电子对，没有孤对电子，所以 CH4 的键角是精确的 109.5 度，分子呈正四面体形。

The tetrahedron is the most frequently appearing geometry in A-Level Chemistry. When there are four electron pairs around the central atom and all of them are bonding pairs, the four pairs separate at 109.5 degrees in three-dimensional space, forming a regular tetrahedron. The most typical example is methane CH4. The carbon atom has four bonding pairs and no lone pairs, so the bond angle of CH4 is exactly 109.5 degrees and the molecule is tetrahedral.

氨气 NH3 是四面体电子对排布下最经典的”变形”案例。氮原子周围有四对电子对，其中三对是成键电子对，一对是孤对电子。孤对电子的排斥力强于成键电子对，把三对 N-H 键”压”得更近，键角从 109.5 度减小到约 107 度，分子形状称为三角锥形（trigonal pyramidal）。考试中必须同时写出”电子对排布为四面体、分子形状为三角锥形”这一对概念，缺一不可。

Ammonia NH3 is the classic “deformed” case under a tetrahedral electron pair arrangement. The nitrogen atom has four electron pairs, three bonding pairs and one lone pair. The lone pair repels more strongly than bonding pairs, pushing the three N-H bonds closer together, so the bond angle decreases from 109.5 degrees to about 107 degrees; the molecular shape is called trigonal pyramidal. In exams you must write both concepts together: “electron pair arrangement is tetrahedral, molecular shape is trigonal pyramidal”.

水 H2O 则更进一步。氧原子周围有四对电子对，其中两对是成键电子对，两对是孤对电子。两对孤对电子的双重排斥把 O-H 键压得更紧，键角进一步减小到约 104.5 度，分子形状为弯曲形（bent）。把 CH4、NH3、H2O 三个分子放在一起对比，是理解孤对电子数目如何逐步压缩键角的最佳素材：孤对电子从 0 到 1 再到 2，键角从 109.5 度到 107 度再到 104.5 度。

Water H2O goes one step further. The oxygen atom has four electron pairs, two bonding pairs and two lone pairs. The double repulsion of two lone pairs squeezes the O-H bonds even closer, reducing the bond angle to about 104.5 degrees, giving a bent molecular shape. Comparing CH4, NH3 and H2O side by side is the best material for understanding how the number of lone pairs progressively compresses bond angles: as lone pairs go from 0 to 1 to 2, bond angles go from 109.5 degrees to 107 degrees to 104.5 degrees.

五、三角双锥与八面体：五对和六对电子对的空间构型 | Trigonal Bipyramidal and Octahedral: Five and Six Electron Domains

当中心原子周围有五对电子对时，电子对排布为三角双锥形（trigonal bipyramidal）。三角双锥由两个”轴向”位置和三个”赤道”位置组成，轴向位置与赤道位置的夹角为 90 度，赤道位置之间的夹角为 120 度。典型例子是五氯化磷 PCl5。磷原子周围有五对成键电子对，没有孤对电子，因此 PCl5 是三角双锥形，分子中同时存在 90 度和 120 度两类键角。

When there are five electron pairs around the central atom, the electron pair arrangement is trigonal bipyramidal. A trigonal bipyramid consists of two “axial” positions and three “equatorial” positions; axial-equatorial angles are 90 degrees while equatorial-equatorial angles are 120 degrees. A typical example is phosphorus pentachloride PCl5. The phosphorus atom has five bonding pairs and no lone pairs, so PCl5 is trigonal bipyramidal with both 90-degree and 120-degree bond angles present.

当中心原子周围有六对电子对时，电子对排布为八面体形（octahedral）。八面体可以理解为六个方向均匀指向三维空间，所有相邻键角都是 90 度。典型例子是六氟化硫 SF6。硫原子周围有六对成键电子对，没有孤对电子，因此 SF6 是八面体形，六个 S-F 键完全等价，键角均为 90 度。SF6 是 CIE 考纲中”扩展八电子”（expanded octet）的经典例子，第三周期元素可以容纳超过四对电子对。

When there are six electron pairs around the central atom, the electron pair arrangement is octahedral. An octahedron can be understood as six directions pointing evenly into three-dimensional space, with all adjacent bond angles equal to 90 degrees. A typical example is sulfur hexafluoride SF6. The sulfur atom has six bonding pairs and no lone pairs, so SF6 is octahedral with six completely equivalent S-F bonds, all at 90 degrees. SF6 is the classic example of the “expanded octet” in the CIE syllabus: elements of period 3 can accommodate more than four electron pairs.

五对电子对含有孤对电子的情况需要特别注意。例如四氟化硫 SF4（一对孤对电子）中，孤对电子会优先占据排斥最小的赤道位置，形成变形四面体（seesaw 形）；三氟化氯 ClF3（两对孤对电子）和三碘化氙 XeF2（三对孤对电子）也遵循”孤对电子优先占赤道位”的规则。这部分内容在 CIE A-Level 中属于较高要求，但理解”孤对电子抢占赤道位置”这一规律后，推导并不困难。

Cases with lone pairs among five electron pairs deserve special attention. In sulfur tetrafluoride SF4 (one lone pair), the lone pair preferentially occupies the equatorial position where repulsion is smallest, forming a seesaw shape; chlorine trifluoride ClF3 (two lone pairs) and xenon difluoride XeF2 (three lone pairs) also follow the rule that “lone pairs occupy equatorial positions first”. This content is at a higher level in CIE A-Level, but once you understand the “lone pairs take equatorial positions” rule, the derivation is not difficult.

六、孤对电子的压缩效应：键角为什么变小 | Lone Pair Compression: Why Bond Angles Shrink

键角变化的根本原因是孤对电子与成键电子对排斥力的差异。孤对电子只受一个原子核吸引，电子云更扩散，占据更大空间，因此它对邻近电子对的排斥比成键电子对更强。更强的排斥会把成键电子对之间的夹角压缩，使键角小于理想值。这就是为什么 NH3 的键角（107 度）和 H2O 的键角（104.5 度）都小于 CH4 的 109.5 度。

The fundamental reason for bond angle changes is the difference in repulsion between lone pairs and bonding pairs. A lone pair is attracted by only one nucleus, its electron cloud is more diffuse and occupies more space, so it repels neighbouring electron pairs more strongly than a bonding pair does. The stronger repulsion compresses the angle between bonding pairs, making the bond angle smaller than the ideal value. This is why the bond angles of NH3 (107 degrees) and H2O (104.5 degrees) are both smaller than the 109.5 degrees of CH4.

用排斥力排序可以系统解释所有键角偏差：孤对电子-孤对电子之间的排斥最大，孤对电子-成键电子对次之，成键电子对-成键电子对最小。H2O 中有两对孤对电子，存在 lp-lp 排斥；NH3 中只有一对孤对电子，主要是 lp-bp 排斥；CH4 没有孤对电子，只有 bp-bp 排斥。排斥力越大，键角被压缩得越多，所以 H2O 的键角比 NH3 更小。

The repulsion order systematically explains all bond angle deviations: lone pair-lone pair repulsion is greatest, lone pair-bonding pair is intermediate, and bonding pair-bonding pair is smallest. H2O has two lone pairs and therefore lp-lp repulsion; NH3 has one lone pair and mainly lp-bp repulsion; CH4 has no lone pairs and only bp-bp repulsion. The stronger the repulsion, the more the bond angle is compressed, so H2O has a smaller bond angle than NH3.

CIE 的简答题经常要求”比较 NH3 和 NF3 的键角大小”。这是一个进阶考点：虽然 NH3 和 NF3 都有孤对电子，但氟原子电负性更强，把 N-F 成键电子对拉向自身，使成键电子对离氮原子更远、排斥变小，因此 NF3 的键角（约 102 度）反而小于 NH3（107 度）。这类题目考查的是电负性对成键电子对位置的影响，答题时要同时考虑孤对电子排斥和成键电子对被拉远两个因素。

CIE structured questions often ask you to “compare the bond angles of NH3 and NF3”. This is an advanced point: although both NH3 and NF3 have a lone pair, fluorine is more electronegative and pulls the N-F bonding pairs towards itself, so the bonding pairs lie farther from nitrogen and repel less; therefore the bond angle of NF3 (about 102 degrees) is actually smaller than that of NH3 (107 degrees). Such questions test the effect of electronegativity on the position of bonding pairs, and your answer should consider both lone pair repulsion and the pulling away of bonding pairs.

七、配位键与复杂离子形状：铵根离子、水合氢离子与碳酸根 | Coordinate Bonds and Complex Ion Shapes: NH4+, H3O+ and CO3 2-

配位键（dative bond 或 coordinate bond）是指一对电子完全由一个原子提供的共价键。在 VSEPR 计数时，配位键与普通共价键完全一样，只算一对成键电子对。铵根离子 NH4+ 是配位键的经典例子：氮原子用三对电子与三个氢原子成键后还剩一对孤对电子，这对孤对电子与 H+ 形成配位键，生成 NH4+。氮周围有四对成键电子对、零孤对电子，所以 NH4+ 是正四面体形，键角 109.5 度。

A dative bond (or coordinate bond) is a covalent bond in which both electrons come from one atom. In VSEPR counting, a dative bond is treated exactly like an ordinary covalent bond and counts as one bonding pair. The ammonium ion NH4+ is the classic example: after nitrogen uses three pairs to bond with three hydrogen atoms, one lone pair remains, and this lone pair forms a dative bond with H+, producing NH4+. Nitrogen has four bonding pairs and zero lone pairs, so NH4+ is tetrahedral with 109.5-degree bond angles.

水合氢离子 H3O+ 则是配位键与孤对电子共同作用的例子。水分子中的氧有一对孤对电子，与 H+ 形成配位键后，氧周围变为四对电子对，其中三对是成键电子对、一对是孤对电子。因此 H3O+ 的电子对排布是四面体，分子形状是三角锥形，键角约 107 度，与 NH3 类似。这类”离子也能用 VSEPR 分析”的题目，需要先正确画出路易斯结构并确定总电子数。

The hydronium ion H3O+ is an example where a dative bond and lone pairs act together. After the oxygen of water, which has a lone pair, forms a dative bond with H+, oxygen has four electron pairs: three bonding pairs and one lone pair. Therefore the electron pair arrangement of H3O+ is tetrahedral, its molecular shape is trigonal pyramidal with a bond angle of about 107 degrees, similar to NH3. For such questions, where “ions can also be analysed with VSEPR”, you must first draw the correct Lewis structure and determine the total electron count.

碳酸根离子 CO3 2- 是平面三角形的典型离子例子。碳原子是三配位，与三个氧原子成键，其中两个 C-O 键是单键、一个 C-O 键是双键，通过共振结构（resonance）三个 C-O 键完全等价。碳周围有三对电子对、零孤对电子，所以 CO3 2- 是平面三角形，键角 120 度。类似地，硝酸根 NO3- 和硫酸根 SO4 2- 也都可以用同样的方法分析，SO4 2- 中硫周围有四对成键电子对，呈正四面体形。

The carbonate ion CO3 2- is a typical ionic example of trigonal planar geometry. Carbon is three-coordinate, bonded to three oxygen atoms with two single C-O bonds and one double C-O bond; through resonance, the three C-O bonds are completely equivalent. Carbon has three electron pairs and zero lone pairs, so CO3 2- is trigonal planar with 120-degree bond angles. Similarly, the nitrate ion NO3- and the sulfate ion SO4 2- can be analysed the same way; in SO4 2-, sulfur has four bonding pairs and the ion is tetrahedral.

八、分子极性：形状如何决定分子是否极性 | Molecular Polarity: How Shape Determines Whether a Molecule Is Polar

分子的极性取决于两个条件：分子中含有极性键，且这些极性键的偶极不能相互抵消。判断偶极是否抵消的关键就是分子形状。以二氧化碳 CO2 为例，C=O 键是极性键，但 CO2 是直线形分子，两个 C=O 偶极方向相反、大小相等，完全抵消，因此 CO2 是非极性分子，尽管它含有极性键。

The polarity of a molecule depends on two conditions: the molecule contains polar bonds, and the bond dipoles do not cancel each other out. The key to judging whether dipoles cancel is molecular shape. Taking carbon dioxide CO2 as an example, the C=O bonds are polar, but CO2 is linear: the two C=O dipoles point in opposite directions with equal magnitude and cancel completely, so CO2 is a non-polar molecule even though it contains polar bonds.

水分子则相反。H-O 键是极性键，水的弯曲形结构使两个 O-H 偶极不能抵消，而是叠加出一个指向氧原子的净偶极，因此水是极性分子。同理，氨 NH3 是三角锥形，三个 N-H 偶极不能完全抵消，NH3 是极性分子；而 BF3 是平面三角形，三个 B-F 偶极在平面内对称分布，完全抵消，BF3 是非极性分子。

Water is the opposite. The H-O bonds are polar, and the bent structure of water prevents the two O-H dipoles from cancelling; instead they combine into a net dipole pointing towards the oxygen atom, making water a polar molecule. Similarly, ammonia NH3 is trigonal pyramidal and its three N-H dipoles do not cancel completely, so NH3 is polar; BF3 is trigonal planar and its three B-F dipoles are symmetrically arranged in the plane and cancel completely, so BF3 is non-polar.

CF4 与 CHCl3 的对比是 CIE 常考的极性判断题。CF4 是正四面体，四个 C-F 偶极完全对称、相互抵消，是非极性分子；CHCl3（氯仿）虽然也是四面体构型，但由于四个取代基不同，偶极不能抵消，是极性分子。答题时先写分子形状，再说明偶极是否对称抵消，最后下结论：形状对称则非极性，形状不对称则极性。

The comparison between CF4 and CHCl3 is a common polarity question in CIE. CF4 is tetrahedral with four completely symmetric C-F dipoles that cancel, making it non-polar; CHCl3 (chloroform), although also tetrahedral, has four different substituents so its dipoles do not cancel, making it polar. When answering, first state the molecular shape, then explain whether the dipoles cancel symmetrically, and finally conclude: symmetric shape means non-polar, asymmetric shape means polar.

九、CIE 考试题型与答题框架：电子对数计算四步法 | CIE Exam Questions and Answer Framework: The Four-Step Electron Counting Method

CIE A-Level 化学中关于分子形状的考题主要有三类：选择题（给出分子或离子，判断形状或键角）、简答题（预测形状并解释原因）、以及结合极性、电负性的综合题。无论哪类题目，掌握统一的分析框架都能稳定得分。下面给出针对”预测形状并解释”题型的四步答题框架。

CIE A-Level Chemistry questions on molecular shapes come in three main types: multiple choice (given a molecule or ion, determine the shape or bond angle), structured questions (predict the shape and explain the reason), and integrated questions combining polarity and electronegativity. Regardless of the question type, a unified analytical framework secures marks reliably. Here is the four-step framework for “predict the shape and explain” questions.

第一步，写出中心原子的价电子数，加上或减去电荷修正（阴离子加电子、阳离子减电子），再除以 2 得到电子对总数。第二步，画出路易斯结构，数出成键电子对和孤对电子的数目。第三步，根据电子对总数写出电子对排布名称（直线、平面三角、四面体、三角双锥、八面体）。第四步，根据孤对电子数目修正分子形状，并写出键角，若键角偏离理想值，用”孤对电子排斥更强”解释原因。

Step one: write down the valence electron count of the central atom, add or subtract electrons for charge (add for anions, subtract for cations), then divide by 2 to obtain the total number of electron pairs. Step two: draw the Lewis structure and count the numbers of bonding pairs and lone pairs. Step three: name the electron pair arrangement from the total pair count (linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral). Step four: correct the molecular shape using the number of lone pairs, state the bond angle, and if the angle deviates from the ideal value, explain using “lone pairs repel more strongly”.

以 CIE 2019 年的一道真题为例：预测 ClF3 的形状并解释。氯原子价电子数 7，三个氟原子各贡献 1 个电子，总电子数 10，电子对总数 5。其中三对是成键电子对，两对是孤对电子。五对电子对排布为三角双锥，孤对电子优先占据赤道位置，因此 ClF3 是 T 形（T-shaped），键角约 87.5 度。答题时把四步完整写出，即使结论略有偏差，过程分也能保住。

Take a real CIE question from 2019 as an example: predict the shape of ClF3 and explain. Chlorine has 7 valence electrons, each of the three fluorine atoms contributes 1 electron, giving 10 electrons in total and 5 electron pairs. Three are bonding pairs and two are lone pairs. Five electron pairs arrange as a trigonal bipyramid, the lone pairs preferentially occupy equatorial positions, so ClF3 is T-shaped with a bond angle of about 87.5 degrees. If you write out all four steps completely, you keep the method marks even when the final conclusion is slightly off.

十、常见易错点与对比表格：形状、键角与示例分子速查 | Common Mistakes and Comparison Table: Shapes, Bond Angles and Examples at a Glance

第一个易错点是把电子对排布和分子形状混为一谈。看到 NH3 就写”四面体”是典型错误：NH3 的电子对排布是四面体，但分子形状是三角锥形。第二个易错点是忘记考虑孤对电子对键角的压缩，例如把 H2O 的键角写成 109.5 度而不是 104.5 度。第三个易错点是忽略离子电荷对电子对数的修正，例如 NH4+ 是 4 对电子对而不是 3 对。

The first common mistake is confusing electron pair arrangement with molecular shape. Writing “tetrahedral” for NH3 is a typical error: the electron pair arrangement of NH3 is tetrahedral, but its molecular shape is trigonal pyramidal. The second mistake is forgetting lone pair compression of bond angles, for example writing 109.5 degrees for H2O instead of 104.5 degrees. The third mistake is ignoring the charge correction for ions, for example NH4+ has 4 electron pairs, not 3.

下表汇总了 CIE A-Level 最常考的形状、键角与代表分子或离子，建议考前反复默写。直线形 180 度：CO2、BeCl2；平面三角形 120 度：BF3、CO3 2-、NO3-；弯曲形约 119 度：SO2；四面体 109.5 度：CH4、NH4+、SO4 2-；三角锥形约 107 度：NH3、H3O+；弯曲形约 104.5 度：H2O；三角双锥 90 度和 120 度：PCl5；T 形：ClF3；八面体 90 度：SF6。把这张表记牢，选择题基本可以秒杀。

The table below summarises the most frequently examined shapes, bond angles and representative molecules or ions in CIE A-Level; it is recommended to recite it repeatedly before the exam. Linear 180 degrees: CO2, BeCl2; trigonal planar 120 degrees: BF3, CO3 2-, NO3-; bent about 119 degrees: SO2; tetrahedral 109.5 degrees: CH4, NH4+, SO4 2-; trigonal pyramidal about 107 degrees: NH3, H3O+; bent about 104.5 degrees: H2O; trigonal bipyramidal 90 and 120 degrees: PCl5; T-shaped: ClF3; octahedral 90 degrees: SF6. Memorise this table and the multiple-choice questions become almost instant.

第四个易错点是极性判断只数极性键而不看形状。CH4 有极性键却是非极性分子，H2O 有极性键也是极性分子，区别完全在于形状是否对称。第五个易错点是扩展八电子元素（P、S、Xe 等第三周期及以后元素）可以拥有 5 对或 6 对电子对，不要把 PCl5 或 SF6 强行写成不符合 VSEPR 的形状。考试前把这些易错点逐一对照检查，能有效减少低级失误。

The fourth mistake is judging polarity by counting polar bonds only, without considering shape. CH4 has polar bonds yet is non-polar, while H2O has polar bonds and is polar; the difference lies entirely in whether the shape is symmetric. The fifth mistake is forgetting that expanded-octet elements (P, S, Xe and other elements of period 3 and beyond) can hold 5 or 6 electron pairs, so PCl5 and SF6 should never be forced into shapes that violate VSEPR. Checking these pitfalls one by one before the exam effectively reduces careless errors.

Summary | 总结

分子形状与几何构型是 CIE A-Level 化学结构化学部分的核心内容，也是历年考试的高频考点。掌握 VSEPR 理论的关键在于三点：一是分清电子对排布与分子形状的区别，二是牢记孤对电子排斥强于成键电子对，三是熟练运用”数电子对、定排布、修正形状、写键角”的四步框架。只要把 CH4、NH3、H2O、BF3、PCl5、SF6 这些经典例子的形状与键角记牢，再配合对配位键、离子电荷和极性判断的理解，分子形状类题目可以稳定拿到高分。

Molecular shapes and geometry are the core of the structure and bonding section in CIE A-Level Chemistry, and a high-frequency exam topic year after year. The key to mastering VSEPR theory lies in three points: first, distinguish clearly between electron pair arrangement and molecular shape; second, remember that lone pairs repel more strongly than bonding pairs; third, practise the four-step framework of “count electron pairs, determine arrangement, correct shape, state bond angle”. As long as you memorise the shapes and bond angles of classic examples such as CH4, NH3, H2O, BF3, PCl5 and SF6, and combine this with an understanding of dative bonds, ionic charge and polarity judgement, you can consistently score high marks on molecular shape questions.

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