📚 Electron Configuration for IB and WJEC Chemistry: Key Points | IB WJEC 化学：电子排布 考点精讲
Electron configuration describes the arrangement of electrons in an atom’s orbitals. Understanding these arrangements is crucial for predicting chemical properties, bonding behaviour, and the trends seen across the periodic table. For IB and WJEC Chemistry students, mastering electron configurations involves learning the rules for filling orbitals, writing notations correctly, and explaining exceptions observed in transition metals. This guide breaks down every essential concept you need to excel, from energy sublevels to shorthand noble gas notation, with a focus on exam-style clarity.
电子排布描述了原子中电子在轨道上的排列方式。理解这些排列对于预测化学性质、成键行为以及元素周期表中的递变规律至关重要。对于学习 IB 和 WJEC 化学的学生来说，掌握电子排布需要掌握轨道填充规则、正确书写表示法，并能解释过渡金属中出现的特例。本指南将分解每一个你需要掌握的核心概念，从能级亚层到稀有气体短写形式，以清晰的考试导向帮助你取得优异成绩。
1. What is Electron Configuration? | 什么是电子排布？
Electron configuration is the distribution of electrons of an atom or molecule in atomic or molecular orbitals. It tells us which energy levels and sublevels the electrons occupy in the ground state. The configuration determines the chemical reactivity, valency, and magnetic properties of the element. For example, the electron configuration of carbon (atomic number 6) is written as 1s² 2s² 2p², indicating two electrons in the 1s orbital, two in the 2s, and two in the 2p sublevel.
电子排布是指原子或分子中电子在原子轨道或分子轨道上的分布。它告诉我们电子在基态时占据哪些能级和亚层。这种排布决定了元素的化学反应活性、化合价及磁学性质。例如，碳（原子序数6）的电子排布写作 1s² 2s² 2p²，表示 1s 轨道有2个电子，2s 轨道有2个电子，2p 亚层有2个电子。
The arrangement follows specific rules based on quantum mechanics. Electrons are not randomly distributed; they fill atomic orbitals in order of increasing energy, often visualised using the Aufbau principle. We represent configurations using notation where the energy level is a number, the sublevel is a letter (s, p, d, f), and a superscript shows the number of electrons in that sublevel.
电子排布遵循基于量子力学的特定规则。电子并非随机分布；它们按照能量递增的顺序填充原子轨道，通常用构造原理（Aufbau principle）进行可视化。我们使用的表示法中，能级用数字表示，亚层用字母（s, p, d, f）表示，上标数字表示该亚层中的电子数。
2. Energy Levels and Sublevels | 能级与亚层
An atom contains principal energy levels (shells) labelled n = 1, 2, 3, and so on. As n increases, the energy of the level increases and the electrons are, on average, further from the nucleus. Each principal energy level is divided into sublevels: s, p, d, f. The number of sublevels in a given shell equals n. For instance, the first shell (n=1) has only the 1s sublevel, the second shell (n=2) has 2s and 2p, and the third (n=3) has 3s, 3p, and 3d.
原子包含主能级（电子层），用 n = 1, 2, 3 …… 表示。n 越大，能级的能量越高，电子平均离原子核越远。每个主能级又分为若干亚层：s, p, d, f。给定电子层中亚层的数量等于 n。例如，第一层（n=1）只有 1s 亚层，第二层（n=2）有 2s 和 2p，第三层（n=3）则有 3s, 3p, 3d。
The sublevels differ in shape and energy. An s sublevel has 1 orbital, a p sublevel has 3 orbitals, a d sublevel has 5, and an f sublevel has 7. Each orbital can hold a maximum of 2 electrons due to the Pauli exclusion principle. Therefore, an s sublevel can accommodate 2 electrons, p can hold 6, d can hold 10, and f can hold 14 electrons.
亚层在形状和能量上各不相同。s 亚层有 1 个轨道，p 亚层有 3 个轨道，d 亚层有 5 个，f 亚层有 7 个。根据泡利不相容原理，每个轨道最多可容纳 2 个电子。因此，s 亚层最多容纳 2 个电子，p 容纳 6 个，d 容纳 10 个，f 容纳 14 个。
Total electrons in shell = 2n²
各层电子总数 = 2n²
For n=1: 2 electrons; n=2: 8 electrons; n=3: 18 electrons (though the outermost shell rarely holds more than 8 for stability reasons).
n=1 时最多 2 个电子；n=2 时 8 个；n=3 时 18 个（但由于稳定性原因最外层通常不超过 8 个电子）。
3. Orbitals: s, p, d, f | 轨道：s, p, d, f
Orbitals are regions in space where there is a high probability of finding an electron. Each type has a characteristic shape: s orbitals are spherical, p orbitals are dumbbell-shaped and exist in three orientations (pₓ, pᵧ, p₂), d orbitals have cloverleaf shapes (except for d₂² which is different), and f orbitals have more complex shapes. These shapes influence how atoms bond with each other.
轨道是空间中发现电子概率较高的区域。每种轨道都有特征形状：s 轨道为球形，p 轨道呈哑铃形且存在于三种取向（pₓ, pᵧ, p₂），d 轨道多为花瓣形（d₂² 例外），f 轨道形状更为复杂。这些形状影响了原子间的成键方式。
When writing electron configurations, we treat the three p orbitals, five d orbitals, and seven f orbitals as sets. Electrons will occupy empty degenerate (equal-energy) orbitals singly before pairing up, which is Hund’s rule. Understanding orbital capacity is essential for deducing bonding patterns and magnetic behaviour.
书写电子排布时，我们视三个 p 轨道、五个 d 轨道和七个 f 轨道为简并（能量相等）的轨道组。电子在配对之前会优先单独占据所有空的简并轨道，这就是洪特规则。理解轨道容量对于推断成键模式和磁学行为至关重要。
4. Electron Filling Rules (Aufbau, Pauli, Hund) | 电子填充规则（构造原理、泡利不相容原理、洪特规则）
Three fundamental rules govern the filling of electrons into atomic orbitals for ground-state atoms:
基态原子中电子填入原子轨道遵循三条基本规则：
- Aufbau principle: Electrons occupy the lowest energy orbitals available first. The order of filling is not simply 1, 2, 3… because the 4s orbital is lower in energy than 3d for atoms up to calcium. A common mnemonic is the diagonal rule or using the energy sequence: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, etc. However, after scandium, 3d becomes lower in energy than 4s, which explains ionic configurations.
- 构造原理（Aufbau principle）：电子优先填入能量最低的轨道。填充顺序并非简单的 1、2、3…… 因为对于钙之前的原子，4s 轨道能量低于 3d。常用的记忆方法是斜线规或使用能量顺序：1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p 等。然而，在钪之后 3d 的能量变得低于 4s，这可以解释离子的电子排布。
- Pauli exclusion principle: No two electrons in an atom can have the same set of four quantum numbers. Practically this means an orbital can hold a maximum of two electrons, and they must have opposite spins. We represent this with up (↑) and down (↓) arrows.
- 泡利不相容原理（Pauli exclusion principle）：原子中不可能有两个电子具有完全相同的四个量子数。实际上这意味着一个轨道最多容纳两个电子，且它们必须自旋相反。我们用向上的箭头 (↑) 和向下的箭头 (↓) 来表示。
- Hund’s rule: Electrons will occupy all degenerate orbitals singly with parallel spins before any orbital is doubly occupied. This minimises electron-electron repulsion and stabilises the atom. For example, for nitrogen (1s² 2s² 2p³), all three p orbitals each contain one unpaired electron with the same spin.
- 洪特规则（Hund’s rule）：在任何一个轨道被双双占据之前，电子会优先以平行自旋单独占据所有简并轨道。这使电子间排斥最小，原子更稳定。例如，氮原子 (1s² 2s² 2p³) 的三个 p 轨道各有一个未成对电子且自旋相同。
5. Writing Electron Configurations | 书写电子排布
To write the electron configuration of an element, start with its atomic number (Z), which equals the number of electrons in a neutral atom. Follow the energy order, adding electrons to sublevels up to the total. For example, oxygen (Z=8): 1s² 2s² 2p⁴. Check your work: the sum of the superscripts must equal Z.
要书写某种元素的电子排布，先从它的原子序数 (Z) 开始，它等于中性原子的电子数。遵循能量顺序，向各亚层添加电子直到总数吻合。例如，氧 (Z=8)：1s² 2s² 2p⁴。检查你的写法：所有上标数字之和必须等于 Z。
For elements in the d-block, remember that after filling 4s with two electrons, the next electrons go into 3d. Vanadium (Z=23): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d³. Notice how 3d is written after 4s, though the energy order changes across the series. For exam purposes, IB and WJEC both accept the configuration with 4s written before 3d for neutral atoms, but the filling order should be understood.
对于 d 区元素，记住在 4s 填入两个电子后，后续电子进入 3d。钒 (Z=23)：1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d³。注意这里 3d 写在 4s 之后，尽管能量顺序在该系列中会发生变化。考试中，IB 和 WJEC 均接受中性原子将 4s 写在 3d 之前的写法，但必须理解填充顺序。
Complete electron configurations can be long, so a compact notation using the previous noble gas is often used (see Section 7). A table helps summarise the maximum capacity of the first four shells:
完整的电子排布可能很长，因此常使用上一周期稀有气体的短写形式（参见第7节）。下表总结了前四个电子层的最大容量：
| Principal Level (n) | Sublevels | Orbitals | Max Electrons |
|---|---|---|---|
| 1 | 1s | 1 | 2 |
| 2 | 2s 2p | 1+3=4 | 8 |
| 3 | 3s 3p 3d | 1+3+5=9 | 18 |
| 4 | 4s 4p 4d 4f | 1+3+5+7=16 | 32 |
6. Orbital Diagrams (Box and Arrow) | 轨道表示法（框线箭头图）
Orbital diagrams use boxes or lines to represent orbitals and arrows for electrons. Each box corresponds to one orbital. A horizontal line can separate sublevels. For carbon (1s² 2s² 2p²), the diagram shows:
轨道表示法用方框或短线代表轨道，用箭头代表电子。每个方框对应一个轨道。可以用水平线区分亚层。碳 (1s² 2s² 2p²) 的轨道图如下所示：
1s: [↑↓] 2s: [↑↓] 2p: [↑ ][↑ ][ ]
Note that Hund’s rule is applied in the 2p sublevel: the two electrons occupy two separate orbitals with parallel spins. Orbital diagrams are excellent for visualising unpaired electrons, which determine paramagnetism. An atom with unpaired electrons is paramagnetic (attracted to a magnetic field); with all electrons paired, it is diamagnetic.
注意在 2p 亚层中应用了洪特规则：两个电子以平行自旋各自占据一个轨道。轨道图非常适合用来可视化未成对电子，它们决定了顺磁性。拥有未成对电子的原子是顺磁性的（被磁场吸引）；所有电子都成对的原子则是抗磁性的。
For heavier elements, drawing full orbital diagrams is tedious, so they are usually drawn only for the valence shell or the last-filled sublevels. However, IB questions may ask you to deduce the orbital diagram for specific d-block ions, for which you must remember to remove electrons from the 4s orbital first.
对于更重的元素，绘制完整的轨道图会非常繁琐，因此通常只画价层或最后填充的亚层。然而，IB 考题可能会要求你推断特定 d 区离子的轨道图，这种情况下需要记住优先从 4s 轨道移除电子。
7. Shorthand Noble Gas Notation | 稀有气体简化写法
To avoid writing long electron configurations, we use the previous noble gas in brackets to represent the core electrons. The noble gas core corresponds to the filled inner shells. For instance, sodium (Na, Z=11) can be written as [Ne] 3s¹ instead of 1s² 2s² 2p⁶ 3s¹. The neon core [He] 2s² 2p⁶ represents 10 electrons.
为了避免书写冗长的电子排布式，我们用方括号内的上一周期稀有气体来表示原子实电子。稀有气体原子实对应于已填满的内层电子。例如，钠 (Na, Z=11) 可以写作 [Ne] 3s¹，而不是 1s² 2s² 2p⁶ 3s¹。氖原子实 [He] 2s² 2p⁶ 代表 10 个电子。
Common noble gas cores and their numbers: [He]=2, [Ne]=10, [Ar]=18, [Kr]=36, [Xe]=54, [Rn]=86. To use this method, find the largest noble gas with fewer electrons than your atom, subtract its atomic number from the element’s atomic number, and then write the remaining configuration. For chlorine (Z=17): [Ne] 3s² 3p⁵.
常见的稀有气体原子实及其电子数：[He]=2, [Ne]=10, [Ar]=18, [Kr]=36, [Xe]=54, [Rn]=86。使用这种方法时，先找到原子序数小于且最接近该原子的稀有气体，用元素的原子序数减去稀有气体的原子序数，然后写出剩余的排布。对于氯 (Z=17)：[Ne] 3s² 3p⁵。
IB exam questions frequently require students to write shorthand configurations for d-block and f-block elements. For iron (Z=26), the shorthand configuration is [Ar] 4s² 3d⁶, not [Ar] 3d⁶ 4s²? Actually, the order may vary, but writing [Ar] 3d⁶ 4s² is also seen. It is safer to follow the convention that 4s is written before 3d for neutral atoms, unless the question specifies the energy-ordered form. WJEC often follows the same convention.
IB 考题常要求学生写出 d 区和 f 区元素的简化电子排布。对于铁 (Z=26)，简化排布是 [Ar] 4s² 3d⁶，而不是 [Ar] 3d⁶ 4s²？实际上顺序可以变化，但将 4s 写在 3d 之前也是常见的。更安全的做法是中性原子按 4s 在 3d 之前的惯例书写，除非题目指定了按能量顺序书写的形式。WJEC 通常也遵循相同的惯例。
8. Exceptions to the Rules | 规则例外
The Aufbau principle predicts many configurations correctly, but there are notable exceptions, particularly among transition metals. The most famous are chromium (Cr) and copper (Cu). Instead of the expected [Ar] 4s² 3d⁴, chromium adopts [Ar] 4s¹ 3d⁵. An atom is more stable when its d sublevel is half-filled (d⁵) or completely filled (d¹⁰), due to symmetry and exchange energy. Thus, one electron from the 4s orbital is promoted to 3d to achieve a half-filled 3d sublevel.
构造原理可以正确预测许多排布，但存在一些显著的例外，特别是在过渡金属中。最著名的是铬 (Cr) 和铜 (Cu)。铬的实际排布是 [Ar] 4s¹ 3d⁵，而不是预期的 [Ar] 4s² 3d⁴。当 d 亚层处于半充满 (d⁵) 或全充满 (d¹⁰) 状态时，由于对称性和交换能，原子会更稳定。因此，4s 轨道的一个电子被激发到 3d，以实现半充满的 3d 亚层。
Similarly, copper adopts [Ar] 4s¹ 3d¹⁰ instead of [Ar] 4s² 3d⁹. This gives a completely filled 3d set. Other exceptions exist (e.g., Mo, Ag, Au), but Cr and Cu are the ones most commonly tested in IB and WJEC. You should be able to explain these exceptions in terms of the extra stability of d⁵ and d¹⁰ configurations.
类似地，铜采用 [Ar] 4s¹ 3d¹⁰ 而不是 [Ar] 4s² 3d⁹，从而获得全充满的 3d 组态。还存在其他例外（如 Mo, Ag, Au），但铬和铜是 IB 和 WJEC 最常考查的。你应该能够用 d⁵ 和 d¹⁰ 排布具有额外稳定性来解释这些例外。
For ions, the 4s electrons are removed before 3d electrons. For example, Fe²⁺ is [Ar] 3d⁶, not [Ar] 4s² 3d⁴. This is because once the 3d orbitals are populated, they drop in energy below 4s. Always remove electrons from the highest principal quantum number n first, which corresponds to 4s.
对于离子，4s 电子会先于 3d 电子被移除。例如，Fe²⁺ 是 [Ar] 3d⁶，而不是 [Ar] 4s² 3d⁴。这是因为一旦 3d 轨道被填充，其能量就会降至 4s 以下。通常先移除主量子数 n 最大的电子，即 4s 电子。
9. Electron Configuration of Ions | 离子的电子排布
When an atom forms a positive ion (cation), electrons are lost from the outermost shell first. For main-group elements, this simply means removing electrons from the highest n level. For a sodium ion, Na⁺, the configuration is 1s² 2s² 2p⁶, the same as neon. For a magnesium ion, Mg²⁺, it also becomes 1s² 2s² 2p⁶.
当原子形成阳离子时，电子会从最外层首先失去。对于主族元素，这仅仅意味着从最高 n 能级移除电子。对于钠离子 Na⁺，电子排布为 1s² 2s² 2p⁶，与氖相同。对于镁离子 Mg²⁺，同样变为 1s² 2s² 2p⁶。
For transition metal ions, the 4s electrons are lost before the 3d electrons, even though 4s was filled first. Thus, Fe: [Ar] 4s² 3d⁶ loses two electrons to form Fe²⁺: [Ar] 3d⁶. Fe³⁺ would be [Ar] 3d⁵. This explains why many transition metal ions have partially filled d orbitals, which is responsible for their colour and catalytic properties.
对于过渡金属离子，尽管填充时 4s 先于 3d，但失去电子时 4s 电子先于 3d 电子失去。因此，Fe: [Ar] 4s² 3d⁶ 失去两个电子形成 Fe²⁺: [Ar] 3d⁶。Fe³⁺ 则为 [Ar] 3d⁵。这解释了为何许多过渡金属离子具有未填满的 d 轨道，从而表现出颜色和催化性质。
Negative ions (anions) gain electrons in the lowest available energy orbitals following the Aufbau principle. For example, O²⁻ gains two electrons: oxygen’s configuration is 1s² 2s² 2p⁴; adding two electrons fills the 2p sublevel, giving 1s² 2s² 2p⁶, which is isoelectronic with neon.
阴离子获得电子时，电子按构造原理填入能量最低的可用轨道。例如，O²⁻ 获得两个电子：氧的排布为 1s² 2s² 2p⁴；添加两个电子后填满 2p 亚层，得到 1s² 2s² 2p⁶，与氖等电子。
A table comparing some common ions:
| Ion | Electron Configuration | Remarks |
|---|---|---|
| Cl⁻ | 1s² 2s² 2p⁶ 3s² 3p⁶ (or [Ar]) | Isoelectronic with Ar |
| K⁺ | 1s² 2s² 2p⁶ 3s² 3p⁶ ([Ar]) | Loses 4s¹ electron |
| Cr³⁺ | [Ar] 3d³ | From [Ar] 4s¹ 3d⁵, lose 4s¹ and two 3d |
| Cu²⁺ | [Ar] 3d⁹ | From [Ar] 4s¹ 3d¹⁰, lose 4s¹ and one 3d |
10. Relationship with the Periodic Table | 与周期表的关系
The periodic table is structured around electron configurations. The period number corresponds to the highest principal quantum number (n) being filled. The group number (for main groups) often equals the number of valence electrons. The block (s, p, d, f) indicates which sublevel is being filled last.
元素周期表的结构基于电子排布。周期数对应于正在填充的最高主量子数 (n)。主族元素的族数通常等于价电子数。区（s, p, d, f）则指示最后填充的是哪个亚层。
- s-block: Groups 1 and 2, plus helium. Their valence configurations end in s¹ or s².
- s 区：第 1 族和第 2 族，外加氦。它们的价层排布以 s¹ 或 s² 结尾。
- p-block: Groups 13 to 18. Their configurations fill the p sublevel, from p¹ to p⁶ (except helium).
- p 区：第 13 族至第 18 族。它们的排布依次填充 p 亚层，从 p¹ 到 p⁶（氦除外）。
- d-block: Transition metals, Groups 3 to 12. Here the d sublevel is being filled across the series, with some exceptions.
- d 区：过渡金属，第 3 族至第 12 族。在这一系列中填充的是 d 亚层，但存在一些例外。
- f-block: Lanthanides and actinides. These involve the filling of the 4f and 5f sublevels.
- f 区：镧系和锕系。它们涉及 4f 和 5f 亚层的填充。
Knowing the block helps you quickly predict the outer electron configuration of an element. For example, any element in Group 16 (oxygen group) has an outer configuration of ns² np⁴. This predictive power is frequently tested in multiple-choice questions on both IB and WJEC papers.
了解元素所处区可以帮助你快速预测该元素的外层电子排布。例如，第 16 族（氧族）中的任何元素都具有 ns² np⁴ 的外层排布。这种预测能力在 IB 和 WJEC 试卷的选择题中经常受到考查。
11. IB and WJEC Exam Tips | IB 和 WJEC 考试技巧
Both IB Chemistry and WJEC Chemistry place strong emphasis on electron configuration in topics such as periodicity, bonding, and transition metal chemistry. Here are some tips to maximise your marks:
IB 化学和 WJEC 化学在周期性、化学键以及过渡金属化学等课题中都十分强调电子排布。以下是一些提升得分的技巧：
- Memorise exceptions: Cr and Cu are almost guaranteed to appear. Be ready to write their configurations and explain using half-filled and fully filled stability.
- 熟记特例：Cr 和 Cu 基本必考。做好书写它们的排布并用半充满和全充满的稳定性来解释的准备。
- Isoelectronic species: Questions often ask to identify ions or
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