📚 Intermolecular Forces | 分子间作用力
Intermolecular forces are the attractive interactions that exist between discrete molecules, separate atoms of the noble gases, or between molecular ions. Although they are significantly weaker than the covalent and ionic bonds that hold atoms together within a molecule, these forces dictate crucial physical properties such as boiling point, melting point, viscosity, surface tension and solubility. A thorough grasp of intermolecular forces is essential for predicting and explaining the macroscopic behaviour of substances, from the anomalously high boiling point of water to the solubility of biomolecules.
分子间作用力是存在于独立分子、惰性气体原子或分子离子之间的吸引力。尽管它们远弱于将原子结合在分子中的共价键和离子键，但这些力决定了沸点、熔点、粘度、表面张力和溶解度等关键物理性质。透彻理解分子间作用力对于预测和解释物质的宏观行为至关重要，从水反常的高沸点到生物分子的溶解度，都离不开它们。
1. What Are Intermolecular Forces? | 什么是分子间作用力？
Intermolecular forces are the attractive or repulsive interactions that occur between neighbouring particles. The term encompasses a range of forces including van der Waals forces (London dispersion forces, dipole–dipole interactions and dipole–induced dipole forces) and hydrogen bonds. All condensed phases – liquids and solids – exist because these attractive forces overcome the kinetic energy of particles. The strength of intermolecular forces relative to thermal energy determines whether a substance is a gas, liquid or solid at a given temperature.
分子间作用力是相邻粒子之间发生的吸引或排斥相互作用。该术语涵盖了一系列力，包括范德华力（伦敦色散力、偶极-偶极相互作用和偶极-诱导偶极力）以及氢键。所有凝聚相——液体和固体——之所以存在，是因为这些吸引力克服了粒子的动能。分子间作用力相对于热能的强度决定了物质在给定温度下是气体、液体还是固体。
2. Comparison with Intramolecular Bonds | 与分子内化学键的比较
Intramolecular forces are the strong covalent, ionic or metallic bonds that hold the constituent atoms of a molecule or giant structure together. Typical covalent bond energies range from 150 to 800 kJ mol⁻¹. In sharp contrast, intermolecular forces are far weaker: London forces can be as little as 0.5 kJ mol⁻¹, permanent dipole–dipole interactions usually lie between 1 and 10 kJ mol⁻¹, and even the strongest hydrogen bonds, such as those in water, reach only about 20–40 kJ mol⁻¹. This large disparity explains why molecular substances melt or boil without breaking their covalent bonds – only the intermolecular forces are overcome during these phase changes.
分子内力是将分子或巨型结构的组成原子结合在一起的强共价键、离子键或金属键。典型的共价键能在150至800 kJ mol⁻¹之间。与此形成鲜明对比的是，分子间作用力要弱得多：伦敦力可低至0.5 kJ mol⁻¹，永久偶极-偶极相互作用通常在1至10 kJ mol⁻¹之间，即使是最强的氢键（例如水中的氢键）也仅达到约20–40 kJ mol⁻¹。这种巨大差异解释了为什么分子物质在熔化或沸腾时不会破坏其共价键——在这些相变过程中，只有分子间作用力被克服。
Covalent bond >> Hydrogen bond >> Dipole–dipole >> London forces
共价键 ＞＞ 氢键 ＞ 偶极-偶极作用 ＞ 伦敦力
3. Types of Intermolecular Forces | 分子间作用力的类型
Intermolecular forces are broadly classified into van der Waals forces and hydrogen bonding. Van der Waals forces include three subtypes: London dispersion forces (present in all particles), permanent dipole–dipole forces (only in polar molecules) and dipole–induced dipole forces. Hydrogen bonding is a particularly strong dipole–dipole interaction that occurs when hydrogen is bonded to highly electronegative nitrogen, oxygen or fluorine atoms. The table below summarises the key features of each type.
分子间作用力大致分为范德华力和氢键。范德华力包括三种亚型：伦敦色散力（存在于所有粒子中）、永久偶极-偶极作用力（仅存在于极性分子中）以及偶极-诱导偶极力。氢键是一种特别强的偶极-偶极相互作用，当氢与电负性很强的氮、氧或氟原子键合时发生。下表总结了每种类型的主要特征。
| Type 类型 | Origin 来源 | Present in 存在于 | Relative strength 相对强度 |
|---|---|---|---|
| London (dispersion) forces 伦敦色散力 | Instantaneous dipole – induced dipole 瞬时偶极-诱导偶极 | All atoms and molecules 所有原子和分子 | Weak; increases with Mᵣ 弱；随分子量增加 |
| Permanent dipole–dipole 永久偶极-偶极 | Electrostatic attraction between permanent dipoles 永久偶极间的静电吸引 | Polar molecules only 仅极性分子 | Moderate 中等 |
| Dipole–induced dipole 偶极-诱导偶极 | Polar molecule induces a dipole in a non‑polar one 极性分子诱导非极性分子产生偶极 | Mixtures of polar and non‑polar species 极性与非极性物种混合物 | Weak 弱 |
| Hydrogen bond 氢键 | Strong dipole–dipole involving H–N, H–O or H–F 涉及H–N、H–O或H–F的强偶极-偶极 | Molecules with N–H, O–H or H–F 含N–H、O–H或H–F的分子 | Strongest intermolecular force 最强的分子间力 |
The relative contributions of these forces determine the physical properties of a substance. In non‑polar molecules such as halogens or alkanes, only London forces operate. In polar molecules like hydrogen chloride, both London forces and permanent dipole–dipole attractions are present. When hydrogen is bonded to N, O or F, hydrogen bonding often dominates the intermolecular interactions.
这些力的相对贡献决定了物质的物理性质。在卤素或烷烃等非极性分子中，只有伦敦力起作用。在像氯化氢这样的极性分子中，伦敦力和永久偶极-偶极吸引力同时存在。当氢与N、O或F键合时，氢键通常在分子间相互作用中占主导地位。
4. London Dispersion Forces | 伦敦色散力
London dispersion forces, named after Fritz London, arise from temporary fluctuations in electron distribution. At any instant, an atom or molecule may develop a transient dipole moment, which distorts the electron cloud of a neighbouring particle, inducing an opposing dipole. The resulting electrostatic attraction is known as the instantaneous dipole–induced dipole force. All particles – even monatomic noble gases – exhibit these forces, making them universal.
伦敦色散力以弗里茨·伦敦命名，源自电子分布的瞬时波动。在任何瞬间，一个原子或分子都可能产生瞬时偶极矩，该偶极矩会扭曲邻近粒子的电子云，诱导出一个相反的偶极。由此产生的静电吸引称为瞬时偶极-诱导偶极力。所有粒子——即使是单原子惰性气体——都表现出这些力，因此它们是普遍存在的。
The potential energy (E) of the London interaction is proportional to the product of the polarisabilities (α) of the two particles and inversely proportional to the sixth power of their separation distance (r):
伦敦相互作用的势能（E）与两个粒子的极化率（α）的乘积成正比，与它们之间距离（r）的六次方成反比：
E ∝ α₁α₂ / r⁶
Polarisability describes the ease with which the electron cloud can be distorted; large, many‑electron atoms or molecules are more polarisable and therefore experience stronger London forces.
极化率描述了电子云变形的容易程度；大的、多电子的原子或分子更容易极化，因此经历更强的伦敦力。
5. Factors Affecting London Forces | 影响伦敦力的因素
The strength of London dispersion forces depends primarily on the number of electrons (and hence the molecular size or relative molecular mass, Mᵣ) and the shape of the molecule. More electrons result in a larger and more easily polarisable electron cloud, so London forces become stronger as Mᵣ increases. For instance, the boiling points of the noble gases rise from helium (−269 °C) to radon (−62 °C) as the atomic number increases, reflecting increasingly stronger dispersion forces.
伦敦色散力的强度主要取决于电子数（从而分子大小或相对分子质量Mᵣ）以及分子的形状。电子数越多，电子云越大、越容易极化，因此随着Mᵣ增加，伦敦力变强。例如，随着原子序数增加，惰性气体的沸点从氦的−269 °C上升到氡的−62 °C，反映了越来越强的色散力。
Molecular shape is equally important. Straight‑chain alkanes have greater surface area of contact between adjacent molecules than their branched isomers, allowing more extensive dispersion interactions. For example, pentane (CH₃CH₂CH₂CH₂CH₃) boils at 36 °C, whereas its branched isomer, 2,2‑dimethylpropane (neopentane), boils at only 10 °C, even though both have the same Mᵣ. The spherical neopentane molecules pack less efficiently and have smaller surface contact, resulting in weaker London forces.
分子形状同样重要。直链烷烃相邻分子间的接触表面积大于其支链异构体，从而允许更广泛的色散相互作用。例如，戊烷（CH₃CH₂CH₂CH₂CH₃）沸点为36 °C，而其支链异构体2,2-二甲基丙烷（新戊烷）沸点仅为10 °C，尽管两者具有相同的Mᵣ。球形的新戊烷分子堆积效率较低，表面接触较小，导致伦敦力较弱。
6. Permanent Dipole–Dipole Interactions | 永久偶极-偶极相互作用
Polar molecules possess a permanent electric dipole moment due to an uneven distribution of electron density. This arises when two atoms with different electronegativities form a polar covalent bond, giving one end a partial negative charge (δ⁻) and the other a partial positive charge (δ⁺). Adjacent polar molecules orient themselves so that the positive end of one molecule is attracted to the negative end of its neighbour, giving a net electrostatic attraction.
极性分子由于电子密度分布不均匀而具有永久电偶极矩。当两个电负性不同的原子形成极性共价键时，就会产生这种不均匀分布，使一端带部分负电荷（δ⁻），另一端带部分正电荷（δ⁺）。相邻的极性分子会定向排列，使一个分子的正端被吸引到相邻分子的负端，从而产生净静电吸引。
Permanent dipole– ndon forces, so polar substances generally have higher melting and boiling points than non‑polar substances of comparable molar mass. For example, butane (C₄H₁₀, non‑ 6 °C) have similar molecular masses, but propanone boils much higher because it experiences both London forces and strong permanent dipole–dipole attractions.
永久偶极-偶极作用力是在伦敦力的基础上额外存在的，因此极性物质的熔点和沸点通常高于摩尔质量相近的非极性物质。例如，丁烷（C₄H₁₀，非极性，Mᵣ = 58，沸点−0.5 °C）和丙酮（CH₃COCH₃，极性，Mᵣ = 58，沸点56 °C）具有相似的分子质量，但丙酮的沸点高得多，因为它同时经历伦敦力和强烈的永久偶极-偶极吸引力。
7. Hydrogen Bonding | 氢键
Hydrogen bonding is a special, stronger type of dipole–dipole interaction. It occurs when a hydrogen atom is covalently bonded to a highly electronegative atom with at least one lone pair of electrons – specifically nitrogen (N), oxygen (O) or fluorine (F). The large electronegativity difference polarises the H–X bond strongly, making the hydrogen highly δ⁺. This partially positive hydrogen is then attracted to a lone pair on an electronegative atom (N, O or F) in a neighbouring molecule.
氢键是一种特殊、更强的偶极-偶极相互作用。当氢原子与至少有一对孤对电子的强电负性原子——具体是氮（N）、氧（O）或氟（F）——共价键合时，就会发生氢键。巨大的电负性差使H–X键强烈极化，使氢带有显著的δ⁺。这个部分带正电的氢随后被邻近分子中电负性原子（N、O或F）上的孤对电子所吸引。
Hydrogen bonds are roughly ten times stronger than ordinary dipole–dipole interactions, with bond strengths between 10 and 40 kJ mol⁻¹. The most famous example is water: each H₂O molecule can form up to four hydrogen bonds, giving water its exceptionally high boiling point (100 °C), high specific heat capacity and maximum density at 4 °C. Hydrogen bonding is also responsible for the double‑helix structure of DNA, the folding of proteins and the solubility of alcohols and sugars in water.
氢键大约比普通的偶极-偶极作用强十倍，键强在10到40 kJ mol⁻¹之间。最著名的例子是水：每个H₂O分子最多可形成四个氢键，这赋予了水极高的沸点（100 °C）、高比热容以及在4 °C时密度最大的特性。氢键还负责DNA的双螺旋结构、蛋白质的折叠以及醇类和糖类在水中的溶解度。
8. Effects on Boiling Point | 对沸点的影响：趋势与异常
Boiling point trends provide clear evidence for different intermolecular forces. For Group 14 hydrides (CH₄, SiH₄, GeH₄, SnH₄), boiling points increase smoothly down the group. These molecules are non‑polar and experience only London forces, which strengthen as the number of electrons and polarisability increase with increasing period.
沸点趋势为不同的分子间作用力提供了明确证据。对于第14族氢化物（CH₄、SiH₄、GeH₄、SnH₄），沸点沿族平稳上升。这些分子是非极性的，仅经历伦敦力，随着周期数增加，电子数和极化率增加，伦敦力增强。
In contrast, the boiling points of hydrides of Groups 15, 16 and 17 show striking anomalies. NH₃, H₂O and HF each have much higher boiling points than expected from extrapolation of the heavier hydrides. This is because the second‑period elements N, O and F are so small and electronegative that hydrogen bonding becomes possible. Water (H₂O) exhibits the largest deviation: based purely on London forces, its boiling point would be about −80 °C, but hydrogen bonding raises it to +100 °C, allowing life to exist on Earth.
相比之下，第15、16和17族氢化物的沸点表现出显著异常。NH₃、H₂O和HF的沸点都比从较重氢化物外推得到的数值高得多。这是因为第二周期元素N、O和F非常小且电负性很强，使得氢键成为可能。水（H₂O）表现出最大的偏差：纯粹基于伦敦力，其沸点大约为−80 °C，但氢键将其提升至+100 °C，从而使地球上的生命得以存在。
A common exam question asks you to explain why H₂O boils higher than H₂S. Although H₂S has more electrons (hence stronger London forces), water molecules form intermolecular hydrogen bonds, which are far stronger than the dipole–dipole and London forces in H₂S. The extra energy required to break hydrogen bonds results in the much higher boiling point of water.
一道常见的考试题要求你解释为什么H₂O的沸点高于H₂S。虽然H₂S有更多的电子（因此伦敦力更强），但水分子之间形成分子间氢键，其强度远超过H₂S中的偶极-偶极和伦敦力。打破氢键所需的额外能量导致水的沸点高得多。
9. Effects on Solubility | 对溶解度的影响
The principle “like dissolves like” is a direct consequence of intermolecular forces. For a solute to dissolve, the intermolecular forces between solute particles and between solvent particles must be broken, which requires energy. New intermolecular forces between solute and solvent particles are then formed, releasing energy. Dissolution is favoured when the energy released is comparable to the energy required, which happens when the solute and solvent have similar types of intermolecular forces.
“相似相溶”原则是分子间作用力的直接结果。溶质溶解时，必须打破溶质粒子之间以及溶剂粒子之间的分子间作用力，这需要能量。然后，溶质与溶剂粒子之间形成新的分子间作用力，释放能量。当释放的能量与所需能量相当时，溶解易于发生，而这发生在溶质和溶剂具有相似类型的分子间作用力时。
Thus, polar and hydrogen‑bonding solutes dissolve well in polar solvents like water. Ethanol (C₂H₅OH) is completely miscible with water because the hydroxyl group (–OH) can both donate and accept hydrogen bonds with water molecules. In contrast, non‑polar substances such as hexane and iodine dissolve poorly in water because the only new interactions would be weak dipole–induced dipole forces, which cannot compensate for the loss of strong hydrogen bonds in water. However, hexane mixes readily with other non‑polar liquids such as cyclohexane, where
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