Electron Pair Sharing Reactions in Detail | IB化学：电子对共享反应详解

📚 Electron Pair Sharing Reactions in Detail | IB化学：电子对共享反应详解

In IB Chemistry, the concept of electron pair sharing forms the foundation of covalent bonding and governs the behaviour of countless molecular systems. This article systematically unpacks the mechanisms, notation, and reactivity patterns surrounding electron pair sharing, with a focus on exam-relevant details for both SL and HL students.

在IB化学课程中，电子对共享的概念构成了共价键的基础，并支配着无数分子体系的行为。本文将系统剖析电子对共享相关的机理、表示方法和反应规律，重点关注SL和HL学生考试中涉及的核心细节。

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1. The Octet Rule and Electron Pair Formation | 八隅体规则与电子对的形成

The octet rule states that atoms tend to gain, lose, or share electrons in order to achieve a stable configuration of eight valence electrons, resembling the electron arrangement of a noble gas. Exceptions include hydrogen (which seeks two electrons), boron (often stable with six), and elements in Period 3 and beyond that can expand their octet.

八隅体规则指出，原子倾向于获得、失去或共享电子，以达到类似稀有气体的八电子稳定价电子构型。例外情况包括氢（只追求两个电子）、硼（常以六电子稳定存在），以及第三周期及以后可扩展八隅体的元素。

When two atoms have similar electronegativities, neither is willing to completely transfer electrons. Instead, they share one or more pairs of electrons. Each shared pair constitutes a single covalent bond, and this mutual sharing allows both atoms to complete their valence shells.

当两个原子的电负性相近时，任何一方都不愿完全转移电子。此时它们选择共享一对或多对电子。每对共享电子构成一个共价单键，这种共享使双方都能填满价电子层。

H · + · H → H : H (equivalent to H−H)

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2. Lewis Structures: Visualising Shared Pairs | Lewis结构：可视化共享电子对

A Lewis structure (or Lewis dot diagram) represents valence electrons as dots placed around atomic symbols. Shared electron pairs are shown as lines (or pairs of dots) between atoms, while lone pairs are depicted as unshared dot pairs on individual atoms.

Lewis结构（又称Lewis点式）以原子符号周围的点来表示价电子。共享电子对以原子之间的短线（或成对的点）表示，而孤对电子则以单个原子上的未共享点对表示。

To draw a Lewis structure correctly, follow these steps: count the total number of valence electrons, arrange the atoms (usually the least electronegative in the centre), connect them with single bonds, distribute remaining electrons as lone pairs, and finally form multiple bonds if any atom lacks an octet.

正确绘制Lewis结构需遵循以下步骤：计算价电子总数；排列原子（通常电负性最小的居中）；以单键连接；将剩余电子分配为孤对电子；若仍有原子未满足八隅体，则形成多重键。

• Water (H₂O): 8 valence electrons; two O−H single bonds and two lone pairs on oxygen.

  水（H₂O）：共8个价电子；两个O−H单键和氧上的两对孤对电子。

• Carbon dioxide (CO₂): 16 valence electrons; two C=O double bonds, no lone pairs on carbon.

  二氧化碳（CO₂）：共16个价电子；两个C=O双键，碳上没有孤对电子。

• Ammonium ion (NH₄⁺): 8 valence electrons (5 + 4×1 − 1); four N−H single bonds, positive formal charge on nitrogen.

  铵根离子（NH₄⁺）：共8个价电子（5 + 4×1 − 1）；四个N−H单键，氮上带正形式电荷。

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3. Coordinate (Dative) Bonding | 配位（授予）键

A coordinate bond, also called a dative covalent bond, is a special type of covalent bond in which both shared electrons are donated by the same atom. This occurs when one species has a lone pair and another species has an empty orbital capable of accepting electrons.

配位键，又称授予共价键，是一种特殊的共价键类型，其中两个共享电子均由同一个原子提供。这发生在一个物种拥有孤对电子、而另一个物种具有可接受电子的空轨道时。

The ammonium ion is a classic example: ammonia (NH₃) donates its lone pair to a hydrogen ion (H⁺), forming an N−H coordinate bond that is indistinguishable from the other three N−H bonds in the ion. Similarly, the hydronium ion (H₃O⁺) forms when water donates a lone pair to H⁺.

铵根离子是经典例子：氨（NH₃）将其孤对电子提供给氢离子（H⁺），形成一根N−H配位键，该键与离子中其他三根N−H键完全无法区分。类似地，水合氢离子（H₃O⁺）在水向H⁺提供孤对电子时形成。

Another important example is the reaction between boron trifluoride (BF₃) and ammonia, where ammonia’s lone pair fills boron’s vacant p-orbital, forming an adduct. In IB exams, students are often asked to identify the donor atom and the acceptor species in such reactions.

另一个重要例子是三氟化硼（BF₃）与氨的反应：氨的孤对电子填入硼的空p轨道，形成加合物。在IB考试中，学生常被要求识别此类反应中的供电子原子和接受电子物种。

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4. Resonance Structures and Electron Pair Delocalisation | 共振结构与电子对离域

Some molecules cannot be accurately represented by a single Lewis structure. Instead, they exhibit resonance: the true structure is a hybrid of two or more contributing structures, with electron pairs delocalised across the molecule.

某些分子无法用单一的Lewis结构准确表示。此时会出现共振：真实结构是两个或多个贡献结构的杂化体，电子对在整个分子中离域。

The ozone molecule (O₃) is a typical example. Its two contributing resonance structures show one O=O double bond and one O−O single bond, but the actual molecule has two identical O−O bonds of intermediate character (bond order 1.5). The delocalisation of electron pairs lowers the overall energy and increases stability.

臭氧分子（O₃）是典型例子。其两个贡献共振结构分别显示一个O=O双键和一个O−O单键，但实际分子的两根O−O键完全等同且具有中间键级（1.5）。电子对离域降低了体系总能量并增强了稳定性。

O=O−O ↔ O−O=O

In the nitrate ion (NO₃⁻), the negative charge and the π electrons are spread equally over all three oxygen atoms. This delocalisation is often indicated by a dashed circle inside the triangle, and it explains the unusual stability of nitrate salts.

在硝酸根离子（NO₃⁻）中，负电荷和π电子均等地分布在三个氧原子之间。这种离域通常以三角形内部的虚线圆圈表示，它解释了硝酸盐的异常稳定性。

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5. VSEPR Theory and the Geometry of Shared Pairs | VSEPR理论与共享电子对的几何构型

The Valence Shell Electron Pair Repulsion (VSEPR) theory states that electron pairs around a central atom arrange themselves as far apart as possible to minimise repulsion. Both bonding pairs and lone pairs exert repulsion, but lone pairs repel more strongly because they occupy more space closer to the nucleus.

价层电子对互斥（VSEPR）理论指出，中心原子周围的电子对会尽可能远离排列以最小化斥力。成键电子对和孤对电子都会产生斥力，但孤对电子的斥力更强，因为它占据更靠近原子核的更大空间。

• Two electron pairs: linear geometry, 180° bond angle (e.g. BeCl₂, CO₂).

  两对电子对：直线形，180°键角（如BeCl₂、CO₂）。

• Three electron pairs: trigonal planar, 120° bond angle (e.g. BF₃, SO₃).

  三对电子对：平面三角形，120°键角（如BF₃、SO₃）。

• Four electron pairs: tetrahedral, 109.5° bond angle (e.g. CH₄, NH₄⁺).

  四对电子对：正四面体，109.5°键角（如CH₄、NH₄⁺）。

• Five electron pairs: trigonal bipyramidal, 90° and 120° bond angles (e.g. PCl₅).

  五对电子对：三角双锥，90°和120°键角（如PCl₅）。

• Six electron pairs: octahedral, 90° bond angles (e.g. SF₆).

  六对电子对：正八面体，90°键角（如SF₆）。

Lone pairs compress bond angles slightly. For example, ammonia (NH₃) has one lone pair, giving a trigonal pyramidal shape with a 107° bond angle, while water (H₂O) has two lone pairs, giving a bent shape with a 104.5° bond angle.

孤对电子会略微压缩键角。例如，氨（NH₃）有一对孤对电子，呈三角锥形，键角约107°；水（H₂O）有两对孤对电子，呈V形（弯曲形），键角约104.5°。

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6. Hybridisation of Atomic Orbitals | 原子轨道杂化

Hybridisation is the mixing of atomic orbitals of similar energy on the same atom to produce a set of degenerate hybrid orbitals. This concept reconciles the observed geometry of molecules with the valence bond theory.

杂化是指同一原子上能量相近的原子轨道相互混合，产生一组简并的杂化轨道。该概念将分子观测到的几何构型与价键理论统一起来。

• sp³ hybridisation: four equivalent orbitals, tetrahedral geometry (CH₄, NH₃, H₂O).

  sp³杂化：四个等价轨道，正四面体构型（CH₄、NH₃、H₂O）。

• sp² hybridisation: three equivalent orbitals, trigonal planar geometry, one unhybridised p-orbital for π bonding (ethene C₂H₄, BF₃).

  sp²杂化：三个等价轨道，平面三角形构型，剩一个未杂化p轨道用于π成键（乙烯C₂H₄、BF₃）。

• sp hybridisation: two equivalent orbitals, linear geometry, two unhybridised p-orbitals for π bonds (ethyne C₂H₂, CO₂).

  sp杂化：两个等价轨道，直线形构型，剩两个未杂化p轨道用于π键（乙炔C₂H₂、CO₂）。

In carbon, the ground state electron configuration is 1s²2s²2p², which has only two unpaired electrons. Yet carbon forms four equivalent bonds in methane. The promotion of a 2s electron to the empty 2p orbital, followed by sp³ hybridisation, explains this observed tetravalence.

碳的基态电子构型为1s²2s²2p²，只有两个未配对电子。但碳在甲烷中形成四个等价键。2s电子被激发到空的2p轨道后，再经sp³杂化，这一过程解释了观测到的四价性。

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7. Sigma (σ) and Pi (π) Bonds | σ键与π键

A sigma (σ) bond forms by head-on overlap of atomic orbitals along the internuclear axis. It is the first bond formed between two atoms and allows free rotation around the bond axis. A pi (π) bond forms by side-on overlap of parallel p-orbitals above and below the internuclear axis, and it restricts rotation.

σ键由原子轨道沿核间轴方向“头碰头”重叠而成，是两个原子间形成的第一根键，允许绕键轴自由旋转。π键则是由平行的p轨道在核间轴上下方“肩并肩”重叠而成，它限制了旋转。

For a double bond, one σ bond and one π bond exist; a triple bond consists of one σ bond and two π bonds. The presence of π bonds makes molecules like ethene planar and restricts rotation, giving rise to geometric (cis-trans) isomerism.

双键由一根σ键和一根π键构成；三键由一根σ键和两根π键构成。π键的存在使乙烯等分子保持平面构型并限制旋转，从而产生几何（顺反）异构现象。

In IB, students should be able to identify σ and π bonds in given structures and explain the difference in bond strength and reactivity: π bonds are weaker than σ bonds due to poorer orbital overlap, making alkenes more reactive than alkanes in addition reactions.

在IB课程中，学生需要能够在给定结构中识别σ键和π键，并解释键强度和反应性的差异：π键因轨道重叠较差而弱于σ键，这使得烯烃在加成反应中比烷烃更具反应活性。

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8. Bond Enthalpy, Bond Length, and Bond Order | 键焓、键长与键级

Bond enthalpy (bond energy) is the energy required to break one mole of a specific covalent bond in the gaseous state. Bond length is the equilibrium distance between two bonded nuclei, and bond order is the number of shared electron pairs between two atoms.

键焓（键能）是指在气态下断裂一摩尔特定共价键所需的能量。键长是两个成键原子核之间的平衡距离，键级则是两个原子之间共享的电子对数目。

Bond Type	Average Bond Enthalpy (kJ/mol)	Bond Length (pm)
C−C	347	154
C=C	614	134
C≡C	839	120

As bond order increases, bond enthalpy increases and bond length decreases because more shared electron pairs pull the nuclei closer together with greater electrostatic attraction. This trend is consistent across C−C, C=C, and C≡C, and it is a frequent multiple-choice question in IB Paper 1.

随着键级增加，键焓增大而键长减小，因为更多共享电子对以更强的静电引力将原子核拉得更近。这一趋势在C−C、C=C和C≡C中保持一致，是IB Paper 1选择题中的高频考点。

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9. Polarity of Covalent Bonds | 共价键的极性

A bond is polar when two bonded atoms have different electronegativities, causing the shared electron pair to be attracted more strongly toward the more electronegative atom. This creates a partial positive charge (δ⁺) on the less electronegative atom and a partial negative charge (δ⁻) on the more electronegative atom.

当两个成键原子的电负性不同时，共价键呈现极性：共享电子对更强烈地偏向电负性较大的原子，使电负性较小的原子带部分正电荷（δ⁺），而电负性较大的原子带部分负电荷（δ⁻）。

Electronegativity increases across a period and decreases down a group in the periodic table. Fluorine is the most electronegative element (Pauling scale 4.0). In hydrogen chloride (HCl), the bond is polar with δ⁺ on H and δ⁻ on Cl, as chlorine’s electronegativity (3.0) exceeds hydrogen’s (2.2).

元素周期表中，电负性在同一周期从左到右递增，在同一族从上到下递减。氟是电负性最大的元素（Pauling标度4.0）。在氯化氢（HCl）中，键具有极性，H带δ⁺而Cl带δ⁻，因为氯的电负性（3.0）大于氢（2.2）。

In addition to bond polarity, molecular polarity depends on molecular geometry. Carbon dioxide has polar C=O bonds, but its linear geometry causes the bond dipoles to cancel, making the molecule non-polar overall. In contrast, water’s bent geometry results in a net dipole, giving it polarity and explaining its properties as a solvent.

除了键的极性，分子极性还取决于分子几何构型。二氧化碳具有极性的C=O键，但其直线形构型使键偶极矩相互抵消，因此分子总体呈非极性。相比之下，水的弯曲构型产生净偶极矩，使其具有极性，这也解释了水作为溶剂的特性。

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10. Homolytic and Heterolytic Bond Fission | 共价键的均裂与异裂

Covalent bonds can break in two distinct ways. Homolytic fission occurs when each atom retains one electron from the shared pair, producing two radicals (neutral species with an unpaired electron). Heterolytic fission occurs when one atom retains both electrons, producing a cation and an anion.

共价键可以两种不同方式断裂。均裂发生时，每个原子从共享对中各保留一个电子，产生两个自由基（具有未配对电子的中性物种）。异裂发生时，一个原子保留两个电子，产生一个阳离子和一个阴离子。

Cl−Cl → Cl· + ·Cl (homolytic)

H−Cl → H⁺ + Cl⁻ (heterolytic)

Homolytic fission typically occurs under ultraviolet light or high temperatures and is common in halogenation reactions of alkanes. Heterolytic fission occurs more readily in polar solvents and is characteristic of acid-base reactions, where the more electronegative atom takes both electrons.

均裂通常在紫外光或高温条件下发生，常见于烷烃的卤代反应。异裂在极性溶剂中更容易发生，是酸碱反应的典型特征：电负性较大的原子带走两个电子。

Radicals are highly reactive species that participate in chain reactions. In the chlorination of methane, initiation produces chlorine radicals, propagation steps generate methyl radicals and HCl, and termination steps combine radicals into stable products. This mechanism is a classic IB HL exam question.

自由基是高度活泼的物种，参与链式反应。在甲烷氯化反应中，引发步骤产生氯自由基，链增长步骤生成甲基自由基和HCl，终止步骤则将自由基结合为稳定产物。这一机理是IB HL考试的经典题型。

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11. Electron Pair Sharing vs. Electron Transfer | 电子对共享与电子转移的对比

Ionic bonding involves complete electron transfer from a metal to a non-metal, forming ions held together by electrostatic attraction. Covalent bonding involves electron pair sharing between non-metals with similar electronegativities. The difference is not absolute: many bonds exhibit both ionic and covalent character.

离子键涉及电子从金属到非金属的完全转移，形成由静电引力结合的离子。共价键则涉及电负性相近的非金属之间共享电子对。这种区分并非绝对：许多键同时表现出离子性和共价性特征。

Property	Covalent Bond	Ionic Bond
Electron behaviour	Shared pair	Transferred
Electronegativity difference	Small (usually < 1.7)	Large (usually > 1.7)
Typical participants	Non-metals	Metal + non-metal
Physical state at room temp	Often gases, liquids or soft solids	Crystalline solids

When the electronegativity difference is intermediate, a polar covalent bond results. In IB, the 1.7 threshold on the Pauling scale is often used to approximate the boundary between ionic and covalent character, though students should recognise that this is a guideline rather than a rigid boundary.

当电负性差值处于中间范围时，形成极性共价键。在IB课程中，Pauling标度上的1.7阈值常被用来近似划分离子键和共价键的边界，但学生应认识到这只是一个参考界限而非刚性分界。

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12. Exam-Focused Summary | 考点总结

To excel in IB Chemistry questions on electron pair sharing, master the following skills: draw Lewis structures accurately for neutral molecules and polyatomic ions; calculate formal charges; determine the number of σ and π bonds; apply VSEPR theory to predict shape and bond angles; identify hybridisation states (sp, sp², sp³); and compare bond properties across different bond orders.

要在IB化学中攻克电子对共享相关题目，需要掌握以下技能：准确绘制中性分子和多原子离子的Lewis结构；计算形式电荷；确定σ和π键的数目；应用VSEPR理论预测构型和键角；识别杂化类型（sp、sp²、sp³）；比较不同键级的键参数。

Practice consistently with past Paper 1 multiple-choice questions and Paper 2 structured questions on Lewis structures of ions like CN⁻, SO₄²⁻, and PO₄³⁻. Remember that resonance, formal charge, and delocalisation frequently appear in higher-mark Extended Response questions.

坚持用历年Paper 1选择题和Paper 2结构题练习CN⁻、SO₄²⁻、PO₄³⁻等离子的Lewis结构。记住：共振、形式电荷和离域常出现在高分值的详细解答题中。

Finally, avoid common errors: forgetting to include all valence electrons, placing the least electronegative atom in the wrong position, neglecting lone pair effects on bond angles, and confusing bond polarity with molecular polarity. Careful, systematic application of the rules outlined above will earn full marks.

最后，避免常见错误：遗漏价电子总数、将电负性最小的原子放错位置、忽略孤对电子对键角的影响、混淆键的极性与分子极性。系统而仔细地运用上述规则，即可获得满分。

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