Electron Configuration Writing Rules | 电子构型的书写规律

📚 Electron Configuration Writing Rules | 电子构型的书写规律

Understanding how to write electron configurations is one of the most fundamental skills in A-Level Chemistry. It not only helps you predict the chemical behaviour of elements but also forms the basis for understanding periodic trends, bonding, and transition metal chemistry. In this article, we will explore the rules, exceptions, and common pitfalls you need to master for your CIE examinations.

掌握电子构型的书写是A-Level化学中最基本的技能之一。它不仅帮助你预测元素的化学行为，更是理解周期性规律、化学键以及过渡金属化学的基础。在本文中，我们将深入探讨CIE考试中你需要掌握的书写规则、特例和常见易错点。

1. The Basics: Orbitals and Energy Levels | 基础：轨道与能级

Before writing electron configurations, you must understand the hierarchy of atomic structure. Electrons occupy regions of space called orbitals, which are grouped into subshells (s, p, d, f), and these subshells are organised within principal energy levels (n = 1, 2, 3, 4…). Each orbital can hold a maximum of two electrons with opposite spins, and each subshell contains a fixed number of orbitals: s has 1, p has 3, d has 5, and f has 7.

在书写电子构型之前，你必须理解原子结构的层级关系。电子占据称为轨道的空间区域，轨道被分为亚层（s、p、d、f），而这些亚层又排列在主能级（n = 1, 2, 3, 4…）之内。每个轨道最多可容纳两个自旋相反的电子，每个亚层含有固定数量的轨道：s亚层有1个，p亚层有3个，d亚层有5个，f亚层有7个。

The maximum number of electrons each subshell can hold is therefore: s = 2, p = 6, d = 10, f = 14. For example, the third principal energy level (n = 3) contains 3s, 3p and 3d subshells, and can hold a total of 2 + 6 + 10 = 18 electrons.

因此，每个亚层能容纳的最大电子数为：s = 2，p = 6，d = 10，f = 14。例如，第三主能级（n = 3）包含3s、3p和3d亚层，总共可以容纳2 + 6 + 10 = 18个电子。

2. The Aufbau Principle | 构造原理

The Aufbau principle (from German “aufbauen”, meaning “to build up”) states that electrons fill orbitals in order of increasing energy. The lowest-energy orbitals are filled first before moving to higher-energy orbitals. The standard filling order can be remembered using the diagonal rule (also called the Madelung rule), which gives the sequence: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p.

构造原理（源自德语”aufbauen”，意为”构建”）指出：电子按照能量递增的顺序填充轨道。能量最低的轨道先被填充，然后才填充更高能量的轨道。标准填充顺序可以用对角线规则（又称马德隆规则）记忆，其顺序为：1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p。

A common source of confusion is the 4s/3d ordering. Although 3d is written before 4s in the periodic table, 4s has a slightly lower energy than 3d and is therefore filled first. For example, potassium (Z = 19) has the configuration [Ar] 4s¹, and calcium (Z = 20) is [Ar] 4s² — not [Ar] 3d¹ or [Ar] 3d².

一个常见的困惑点是4s/3d的排列顺序。尽管在元素周期表中3d写在4s之前，但4s的能量略低于3d，因此在填充时4s优先。例如，钾（Z = 19）的构型是[Ar] 4s¹，钙（Z = 20）是[Ar] 4s²——而不是[Ar] 3d¹或[Ar] 3d²。

Filling order: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p → 7s

填充顺序：1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p → 7s

3. Pauli Exclusion Principle and Hund’s Rule | 泡利不相容原理与洪特规则

The Pauli exclusion principle states that no two electrons in the same atom can have the same set of all four quantum numbers. In practical terms, this means that each orbital can hold at most two electrons, and these two electrons must have opposite spins (one spin-up ↑ and one spin-down ↓). A filled orbital is written with the notation ↑↓.

泡利不相容原理指出：同一原子中，没有任何两个电子可以拥有完全相同的四个量子数。就实际操作而言，这意味着每个轨道最多只能容纳两个电子，且这两个电子必须具有相反的自旋（一个自旋向上↑，一个自旋向下↓）。已填满的轨道用符号↑↓表示。

Hund’s rule states that when electrons occupy degenerate orbitals (orbitals of the same energy, such as the three 2p orbitals), they first fill each orbital singly with parallel spins before any pairing occurs. For example, the electron configuration of nitrogen (Z = 7) is 1s² 2s² 2p³, where the three 2p electrons occupy the three 2p orbitals singly: ↑ ↑ ↑.

洪特规则指出：当电子填充简并轨道（能量相同的轨道，如三个2p轨道）时，电子首先以相同的自旋方向逐个占据每个轨道，之后才开始配对。例如，氮（Z = 7）的电子构型为1s² 2s² 2p³，其中三个2p电子分别单独占据三个2p轨道：↑ ↑ ↑。

Correct orbital diagram for oxygen (Z = 8):

氧（Z = 8）的正确轨道图：

2p: ↑↓ ↑ ↑ (not ↑↓ ↑↓ ↑, not ↑↓ ↑↓)

2p：↑↓ ↑ ↑ （而非 ↑↓ ↑↓ ↑，也非 ↑↓ ↑↓）

4. Writing Full Electron Configurations | 书写完整的电子构型

To write a full electron configuration, follow these steps: (1) determine the atomic number Z, which equals the number of electrons in a neutral atom; (2) fill orbitals according to the Aufbau order; (3) respect the Pauli principle and Hund’s rule when placing electrons into subshells; (4) write the configuration with the number of electrons in each subshell as a superscript.

书写完整电子构型的步骤如下：（1）确定原子序数Z，它等于中性原子中的电子数；（2）按照构造顺序填充轨道；（3）在将电子放入亚层时遵守泡利原理和洪特规则；（4）以每个亚层中的电子数作为上标写出构型。

Let us work through a full example. For iron, Z = 26. We count electrons: 1s² (2), 2s² (4), 2p⁶ (10), 3s² (12), 3p⁶ (18), 4s² (20), then remaining 6 electrons go into 3d, giving 3d⁶. Therefore the full configuration is:

让我们看一个完整示例。对于铁，Z = 26。我们逐层计数：1s²（2），2s²（4），2p⁶（10），3s²（12），3p⁶（18），4s²（20），剩余6个电子进入3d，即3d⁶。因此完整构型为：

Fe: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶

Useful table for quick reference:

常用快速查阅表：

Element	Z	Full Electron Configuration
Sodium 钠	11	1s² 2s² 2p⁶ 3s¹
Chlorine 氯	17	1s² 2s² 2p⁶ 3s² 3p⁵
Calcium 钙	20	1s² 2s² 2p⁶ 3s² 3p⁶ 4s²
Scandium 钪	21	1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹

5. Anomalous Configurations: Chromium and Copper | 异常构型：铬与铜

Some transition metals do not follow the expected filling pattern because a half-filled or fully-filled d subshell provides extra stability due to exchange energy and electron repulsion effects. The two most important exceptions you must memorise for CIE are chromium (Z = 24) and copper (Z = 29).

某些过渡金属并不遵循预期的填充模式，因为半充满或全充满的d亚层由于交换能和电子排斥效应而具有额外的稳定性。你在CIE考试中必须记住的两个最重要特例是铬（Z = 24）和铜（Z = 29）。

Chromium: the predicted configuration would be [Ar] 4s² 3d⁴, but the actual configuration is [Ar] 4s¹ 3d⁵. One electron is promoted from 4s to 3d to achieve a half-filled 3d⁵ configuration, which is exceptionally stable.

铬：预期的构型为[Ar] 4s² 3d⁴，但实际构型为[Ar] 4s¹ 3d⁵。一个电子从4s跃迁到3d，以实现半充满的3d⁵构型，这种构型格外稳定。

Copper: the predicted configuration would be [Ar] 4s² 3d⁹, but the actual configuration is [Ar] 4s¹ 3d¹⁰. Here the 3d subshell is completely filled, which is also exceptionally stable.

铜：预期的构型为[Ar] 4s² 3d⁹，但实际构型为[Ar] 4s¹ 3d¹⁰。在这里3d亚层完全填满，同样格外稳定。

Cr: [Ar] 4s¹ 3d⁵ | Cu: [Ar] 4s¹ 3d¹⁰

铬：[Ar] 4s¹ 3d⁵ | 铜：[Ar] 4s¹ 3d¹⁰

Notice that the 4s orbital is occupied by only one electron (4s¹) in both exceptions. Also remember that in ion formation, 4s electrons are lost before 3d electrons — we will return to this in Section 7.

请注意在这两个特例中，4s轨道都只被一个电子占据（4s¹）。同时需要记住：在形成离子时，4s电子先于3d电子失去——我们将在第7节中讨论这一点。

6. Condensed Configurations and Valence Electrons | 简写构型与价电子

For elements with many electrons, writing the full configuration is tedious. The condensed (noble gas) notation uses the symbol of the preceding noble gas in square brackets to represent the core electrons, followed by the valence electrons. For example, titanium (Z = 22) is written as [Ar] 4s² 3d² instead of 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d².

对于具有很多电子的元素，书写完整构型非常繁琐。简写构型（稀有气体简写法）使用前一个稀有气体的元素符号加方括号来表示核心电子，后接价电子。例如，钛（Z = 22）书写为[Ar] 4s² 3d²，而不是1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d²。

Valence electrons are those in the outermost shell that participate in chemical bonding. For main-group elements, the valence electrons are simply the electrons in the highest principal quantum number s and p subshells. For transition metals, valence electrons include the ns electrons and typically some (n−1)d electrons; for example, iron has 4s² 3d⁶, and its valence electrons include both 4s and 3d electrons.

价电子是最外层壳层中参与化学成键的电子。对于主族元素，价电子就是最高主量子数s和p亚层中的电子。对于过渡金属，价电子包括ns电子以及通常部分(n−1)d电子；例如，铁的构型为4s² 3d⁶，其价电子包括4s和3d电子。

7. Electron Configurations of Ions | 离子的电子构型

When writing electron configurations for ions, the key rule is: for main-group elements, add electrons for negative ions (anions) or remove electrons for positive ions (cations) starting from the highest occupied orbital. However, for transition metal cations, the 4s electrons are removed FIRST, before the 3d electrons — even though 3d is at a higher principal quantum number? No: 3d is n = 3 and 4s is n = 4, so 4s is the outermost; but due to the lower energy of 3d once electrons are present, the 4s electrons are removed first when ionising.

在书写离子的电子构型时，关键规则是：对于主族元素，负离子（阴离子）加电子，正离子（阳离子）则从最高占据轨道开始移除电子。然而，对于过渡金属阳离子，4s电子先于3d电子被移除——尽管3d的主量子数更小？不：3d的n = 3，4s的n = 4，所以4s在最外层；但由于一旦有电子存在后3d的能量更低，因此在电离时4s电子首先被移除。

Let us examine two examples. Fe (Z = 26): [Ar] 4s² 3d⁶. To form Fe²⁺, remove the two 4s electrons: Fe²⁺ = [Ar] 3d⁶. To form Fe³⁺, remove one more 3d electron: Fe³⁺ = [Ar] 3d⁵. Note that 3d⁵ is half-filled, which explains the relative stability of Fe³⁺.

让我们来看两个例子。铁（Z = 26）：[Ar] 4s² 3d⁶。形成Fe²⁺时，移除两个4s电子：Fe²⁺ = [Ar] 3d⁶。形成Fe³⁺时，再移除一个3d电子：Fe³⁺ = [Ar] 3d⁵。注意3d⁵是半充满状态，这解释了Fe³⁺的相对稳定性。

For non-metals, anions gain electrons to achieve a noble gas configuration. For example, oxygen (Z = 8) gains two electrons to form O²⁻: 1s² 2s² 2p⁶, which is the same as neon.

对于非金属，阴离子获得电子以达到稀有气体构型。例如，氧（Z = 8）获得两个电子形成O²⁻：1s² 2s² 2p⁶，这与氖相同。

Species	Configuration
Fe²⁺	[Ar] 3d⁶
Fe³⁺	[Ar] 3d⁵
Cu⁺	[Ar] 3d¹⁰
Cu²⁺	[Ar] 3d⁹
Cl⁻	[Ne] 3s² 3p⁶

8. The Periodic Table Connection | 与周期表的关联

The electron configuration of an element determines its position in the periodic table. The group number for s-block and p-block elements equals the number of valence electrons (for groups 1, 2 and 13–18). The period number equals the highest principal quantum number that contains electrons in the ground state. The block (s, p, d or f) is determined by the subshell that receives the last electron.

元素的电子构型决定了它在周期表中的位置。对于s区和p区元素，族序数等于价电子数（对于第1、2族和第13–18族）。周期数等于基态中含有电子的最高主量子数。区（s、p、d或f）由最后填入电子的亚层决定。

For instance, bromine (Z = 35) has the configuration [Ar] 4s² 3d¹⁰ 4p⁵. It has 7 valence electrons (4s² 4p⁵), placing it in Group 17; its highest occupied principal level is n = 4, placing it in Period 4; and the last electron enters the 4p subshell, placing it in the p-block.

例如，溴（Z = 35）的构型为[Ar] 4s² 3d¹⁰ 4p⁵。它有7个价电子（4s² 4p⁵），位于第17族；其最高占据主能级为n = 4，位于第4周期；最后一个电子填入4p亚层，位于p区。

This relationship allows you to predict the configuration of an element directly from its position in the periodic table — a faster and more reliable strategy in timed examinations.

这种关系使你能够直接从元素在周期表中的位置推断其构型——在限时考试中，这是一种更快、更可靠的策略。

9. Common Mistakes and Exam Tips | 常见错误与备考建议

Students frequently lose marks in examinations for the following errors: (1) writing 3d before 4s in neutral atoms — always write 4s first when electrons are present; (2) forgetting the exceptions for Cr and Cu; (3) removing d electrons before 4s electrons when forming transition metal ions; (4) miscounting electrons, especially for ions; (5) using incorrect superscripts or omitting them entirely.

学生在考试中常因以下错误而失分：（1）在中性原子中把3d写在4s之前——有电子时4s总是先写；（2）忘记铬和铜的特例；（3）形成过渡金属离子时先移除d电子再移除4s电子；（4）数错电子数，特别是离子；（5）使用错误的上标或完全漏写上标。

Here is a checklist for exam success: always verify the atomic number; use the noble gas shortcut to save time; check whether the species is neutral, cation or anion; memorise the seven diatomic elements for electron counting; and always write 4s before 3d in neutral atoms but remove 4s first in ion formation.

以下是考试成功的检查清单：始终核验原子序数；使用稀有气体简写法以节省时间；检查物种是中性、阳离子还是阴离子；记住七种双原子元素以便计数电子；始终在中性原子中先写4s后写3d，但在形成离子时先移除4s。

[Ar] 4s² 3d⁶ → Fe²⁺: [Ar] 3d⁶ → Fe³⁺: [Ar] 3d⁵

[Ar] 4s² 3d⁶ → Fe²⁺：[Ar] 3d⁶ → Fe³⁺：[Ar] 3d⁵

Finally, remember the golden triad: Aufbau principle for the order of filling, Pauli exclusion principle for the maximum occupancy of each orbital, and Hund’s rule for the filling of degenerate orbitals. Master these three rules, and electron configurations become a straightforward and high-scoring topic in your A-Level examination.

最后，请记住黄金三原则：构造原理决定填充顺序，泡利不相容原理决定每个轨道的最大容量，洪特规则决定简并轨道的填充方式。掌握这三条规则，电子构型将成为你A-Level考试中一个简单且高得分的考点。

Published by TutorHao | Chemistry Revision Series | aleveler.com

更多咨询请联系16621398022（同微信）

Electron Configuration Writing Rules | 电子构型的书写规律