📚 Electronegativity and Bond Polarity | 电负性与键的极性
Electronegativity is a fundamental concept in chemistry that explains how atoms attract bonding electrons. It determines whether a bond between two atoms will be covalent, polar covalent or ionic, and it underpins the physical and chemical properties of molecules. Understanding electronegativity and bond polarity is essential for predicting molecular behaviour, solubility, and reaction mechanisms at A-Level and beyond.
电负性是化学中一个基本概念，它解释原子如何吸引成键电子。它决定了两个原子之间的键是共价键、极性共价键还是离子键，并支撑着分子的物理和化学性质。理解电负性与键的极性对于预测分子行为、溶解性以及反应机理至关重要，无论是在A-Level阶段还是更深入的学习中。
1. What is Electronegativity? | 什么是电负性？
Electronegativity is the ability of an atom to attract a pair of electrons in a covalent bond towards itself. It is a relative property that applies only when an atom is part of a bond, not to isolated atoms. The concept was developed by Linus Pauling, who created the widely used Pauling scale. Atoms with high electronegativity, such as fluorine, oxygen and chlorine, have a strong pull on shared electrons, while those with low electronegativity, like caesium and francium, have a weak pull.
电负性是原子在共价键中吸引一对电子向自身偏移的能力。这是一种相对性质，仅适用于原子成键时，而非孤立原子。该概念由莱纳斯·鲍林提出，他建立了广泛使用的鲍林标度。电负性高的原子，如氟、氧和氯，对共用电子有强大的吸引力；而电负性低的原子，如铯和钫，吸引力很弱。
2. Trends in Electronegativity in the Periodic Table | 周期表中电负性的趋势
Electronegativity shows clear periodic trends. Across a period from left to right, electronegativity increases because the nuclear charge increases while the atomic radius decreases, making the nucleus more effective at attracting bonding electrons. For example, in Period 2, lithium has a low electronegativity (1.0), while fluorine has the highest (4.0). Down a group, electronegativity decreases because the atomic radius increases and shielding by inner electrons reduces the effective nuclear pull on outer electrons. Thus, iodine has a lower electronegativity than chlorine.
电负性呈现清晰的周期变化趋势。沿周期从左到右，电负性增加，因为核电荷增加而原子半径减小，使核更有效地吸引成键电子。例如，第二周期中锂的电负性较低（1.0），而氟最高（4.0）。沿族向下，电负性减小，因为原子半径增大，内层电子的屏蔽作用降低了核对最外层电子的有效吸引。因此，碘的电负性低于氯。
3. The Pauling Scale of Electronegativity | 鲍林电负性标度
The Pauling scale assigns fluorine, the most electronegative element, a value of 4.0. All other elements are assigned values relative to fluorine. The scale is dimensionless and based on bond dissociation energies. Key values include hydrogen (2.2), carbon (2.5), nitrogen (3.0), oxygen (3.5), chlorine (3.2) and sodium (0.9). Although other scales such as the Mulliken and Allred-Rochow scales exist, the Pauling scale is the most common in A-Level chemistry.
鲍林标度将电负性最强的氟元素赋值为4.0，其他所有元素都相对氟赋值。该标度无量纲，基于键离解能。关键数值包括氢（2.2）、碳（2.5）、氮（3.0）、氧（3.5）、氯（3.2）和钠（0.9）。尽管存在马利肯标度和阿莱-罗周标度等其他标度，但在A-Level化学中最常用的是鲍林标度。
| Element | Electronegativity (Pauling) |
|---|---|
| F | 4.0 |
| O | 3.5 |
| N, Cl | 3.0 |
| C, S | 2.5 |
| H, P | 2.2 |
| Na | 0.9 |
4. Bond Polarity: Unequal Sharing of Electrons | 键的极性：电子不均匀共享
When two atoms with different electronegativities form a covalent bond, the bonding electron pair is not shared equally. The more electronegative atom attracts the electron cloud more strongly, causing a partial negative charge (δ⁻) on that atom and a partial positive charge (δ⁺) on the other. This separation of charge creates a bond dipole, and the bond is described as polar. The greater the difference in electronegativity, the more polar the bond.
当两个电负性不同的原子形成共价键时，成键电子对并非均匀共享。电负性更强的原子更强烈地吸引电子云，使该原子带上部分负电荷（δ⁻），另一原子带上部分正电荷（δ⁺）。这种电荷分离产生键偶极矩，该键被描述为极性键。电负性差值越大，键的极性越强。
5. Non-Polar Covalent Bonds | 非极性共价键
If two atoms in a bond are identical or have very similar electronegativities, the electrons are shared equally, and the bond is non-polar. Examples include H₂, N₂, Cl₂ and C—H bonds (in hydrocarbons). The electronegativity difference is zero or very small (generally less than 0.4 on the Pauling scale). In such bonds, there is no dipole moment, and the electron density is symmetrical around the bond axis.
如果成键两个原子相同或电负性非常接近，则电子被均匀共享，键为非极性键。例子包括H₂、N₂、Cl₂以及烃类中的C—H键。电负性差值为零或极小（通常在鲍林标度上小于0.4）。在这类键中没有偶极矩，电子密度在键轴周围对称分布。
6. Polar Covalent Bonds and Dipole Moments | 极性共价键与偶极矩
A polar covalent bond forms when the electronegativity difference is between about 0.4 and 1.7–2.0. In water (H₂O), oxygen (3.5) is more electronegative than hydrogen (2.2), so the O—H bonds are polar, with oxygen bearing δ⁻ and hydrogen δ⁺. The separation of charge is measured by the dipole moment (μ), which depends on both the magnitude of the partial charges and the distance between them. The dipole is represented by an arrow pointing from the positive to the negative centre, often with a crossed tail.
当电负性差值约在0.4至1.7–2.0之间时，形成极性共价键。在水（H₂O）中，氧（3.5）的电负性比氢（2.2）高，因此O—H键具有极性，氧带δ⁻、氢带δ⁺。电荷分离程度用偶极矩（μ）衡量，它取决于部分电荷大小及它们之间的距离。偶极用箭头表示，从正电荷中心指向负电荷中心，常带有十字箭头尾。
7. Ionic Bonds as Extreme Polar Bonds | 离子键作为极端极性键
When the electronegativity difference is very large (commonly above 1.7–2.0), the bond is considered ionic. In sodium chloride, chlorine (3.2) and sodium (0.9) have a difference of 2.3; the bonding electrons are almost completely transferred from sodium to chlorine, forming Na⁺ and Cl⁻ ions. Thus, ionic bonding can be viewed as an extreme case of polar covalent bonding, where the electron pair is effectively localised on the more electronegative atom.
当电负性差值非常大（通常大于1.7–2.0）时，该键被认为是离子键。在氯化钠中，氯（3.2）和钠（0.9）的差值为2.3，成键电子几乎完全从钠转移至氯，形成Na⁺和Cl⁻离子。因此，离子键可视为极性共价键的极端情况，其中电子对实际上定域在电负性更强的原子上。
8. Using Electronegativity Differences to Predict Bond Type | 利用电负性差判断键的类型
The electronegativity difference (ΔEN) is a useful guide to predict bond character. A ΔEN of less than about 0.4 typically indicates a non-polar covalent bond. A ΔEN from 0.4 to about 1.7 suggests a polar covalent bond, and a difference greater than 1.7 often corresponds to an ionic bond. However, these boundaries are approximate and not absolute; beryllium chloride (BeCl₂) has a high ΔEN but shows covalent character due to polarisation of the chloride ion by the small Be²⁺ ion.
电负性差值（ΔEN）是预测键性的有用指导。ΔEN约小于0.4通常表明非极性共价键；0.4至约1.7之间提示极性共价键；大于1.7常对应离子键。但这些界限是近似的、并非绝对；氯化铍（BeCl₂）ΔEN较高但因小的Be²⁺离子极化氯离子而表现出共价特征。
ΔEN = |χ₁ − χ₂|
| ΔEN range | Bond type | Example |
|---|---|---|
| 0 – 0.4 | non-polar covalent | Cl₂, CH₄ (C—H) |
| 0.4 – 1.7 | polar covalent | HCl, H₂O |
| > 1.7 | mostly ionic | NaCl, MgO |
9. Partial Charges and the δ Symbol | 部分电荷与δ符号
The Greek letter delta (δ) is used to represent a partial charge because the charge is less than a full unit of electron charge. For example, in hydrogen chloride (HCl), chlorine is more electronegative (3.2) than hydrogen (2.2), so the bond is written as H(δ⁺)—Cl(δ⁻). These partial charges are not formal oxidation states but indicate unequal electron distribution. Dipole arrows are drawn from the δ⁺ atom towards the δ⁻ atom. In complex molecules, showing δ⁺ and δ⁻ helps visualise regions of electron richness and deficiency.
希腊字母δ用于表示部分电荷，因为该电荷小于一个完整的电子电荷单位。例如，在氯化氢（HCl）中，氯（3.2）的电负性比氢（2.2）高，因此键可表示为H(δ⁺)—Cl(δ⁻)。这些部分电荷不是形式氧化数，但指示了不均匀的电子分布。偶极箭头从δ⁺原子指向δ⁻原子。在复杂分子中，标出δ⁺和δ⁻有助于观察电子丰富与缺乏的区域。
10. Molecular Polarity vs Bond Polarity | 分子极性与键的极性
r C=O bonds, but the linear shape means the dipoles are equal and opposite, resulting in a net dipole of zero. In contrast, water (H₂O) is bent, so the O— is a key concept when explaining physical properties such as boiling point and solubility.
含有极性键的分子如果因对称性使单个键偶极相互抵消，则整体可能为非极性。二氧化碳（CO₂）有两个极性C=O键，但线形构型使偶极大小相等方向相反，净偶极为零。相比之下，水（H₂O）为弯曲形，O—H键偶极不能抵消，分子具有极性。因此，分子极性取决于键的极性和分子几何构型。这是解释沸点、溶解度等物理性质时的关键概念。
11. Hydrogen Bonding and Electronegativity | 氢键与电负性
Hydrogen bonding occurs when a hydrogen atom covalently bonded to a highly electronegative atom (N, O or F) experiences an attraction to a lone pair on another electronegative atom. The high electronegativity of N, O and F creates a strongly polar bond, leaving the hydrogen with a significant δ⁺ charge. This δ⁺ hydrogen can then form a strong intermolecular interaction with a δ⁻ centre on a neighbouring molecule. This explains the anomalously high boiling points of water, ammonia and hydrogen fluoride.
当氢原子与高电负性原子（N、O或F）共价键合时，会受到另一个电负性原子上孤对电子的吸引，形成氢键。N、O和F的高电负性产生了强极性键，使氢带明显的δ⁺电荷。这个δ⁺氢随后可与邻近分子上的δ⁻中心形成强烈的分子间作用。这解释了水、氨和氟化氢异常高的沸点。
12. Applications and Exam Tips | 应用与考试提示
In examinations, you may be asked to compare bond polarities, predict molecular polarity or explain trends in physical properties using electronegativity. Always refer to specific electronegativity values and differences where possible. When describing a bond as polar, indicate the direction of the dipole and assign δ⁺ and δ⁻ correctly. Remember that ionic character increases with ΔEN, but polarisation effects can blur the boundary. For molecules, draw a clear diagram showing bond dipoles and resultant molecular dipole if relevant. Use examples like CO₂ and H₂O to illustrate the role of symmetry.
在考试中，你可能会被要求比较键的极性、预测分子极性或用电负性解释物理性质的规律。请尽可能引用具体的电负性值和差值。将键描述为极性时，要标明偶极方向并正确分配δ⁺和δ⁻。记住离子性随ΔEN增加而增强，但极化效应可能模糊其界限。对于分子，画出清晰的示意图，标出键偶极及相关的分子偶极。使用CO₂和H₂O等例子来说明对称性的作用。
Published by TutorHao | Chemistry Revision Series | aleveler.com
更多咨询请联系16621398022（同微信）
或截图后用「扫一扫 → 相册」选择本图