Electron Configuration | 电子排布 考点精讲

📚 Electron Configuration | 电子排布 考点精讲

Understanding electron configuration is fundamental to explaining the chemical properties and periodic trends of elements. In IB and AQA A-level Chemistry, you must be able to write electron configurations for atoms and ions, apply the Aufbau principle, Hund’s rule, and the Pauli exclusion principle, and recognise exceptions such as chromium and copper. This guide covers all key concepts, with examples, diagrams, and exam tips.

理解电子排布是解释元素化学性质和周期律的基础。在IB和AQA A-level化学中，你需要能够书写原子和离子的电子排布，运用构造原理、洪特规则和泡利不相容原理，并识别铬和铜等例外情况。本指南涵盖所有核心概念，并配有示例、图示和考试技巧。

1. Energy Levels and Sub-levels | 能层与亚层

Electrons in atoms are arranged in energy levels (principal quantum number, n = 1, 2, 3, …). Each energy level contains sub-levels: s, p, d, and f. The number of sub-levels in a given energy level equals n. For example, n=1 has only 1s; n=2 has 2s and 2p; n=3 has 3s, 3p, and 3d; n=4 has 4s, 4p, 4d, and 4f.

原子中的电子按能层（主量子数 n = 1, 2, 3 …）排列。每个能层包含若干亚层：s、p、d、f。某能层的亚层数等于 n。例如，n=1 只有 1s 亚层；n=2 有 2s 和 2p；n=3 有 3s、3p、3d；n=4 有 4s、4p、4d 和 4f。

Each type of sub-level contains a specific number of atomic orbitals, and each orbital can hold a maximum of two electrons. The s sub-level has 1 orbital (max 2 electrons), p has 3 orbitals (max 6 e^–), d has 5 orbitals (max 10 e^–), and f has 7 orbitals (max 14 e^–).

每种亚层包含特定数量的原子轨道，每个轨道最多容纳两个电子。s 亚层有 1 个轨道（最多 2 个电子），p 亚层有 3 个轨道（最多 6 e^–），d 亚层有 5 个轨道（最多 10 e^–），f 亚层有 7 个轨道（最多 14 e^–）。

2. Atomic Orbitals: s, p, d, f | 原子轨道：s, p, d, f

An atomic orbital is a region of space where there is a high probability of finding an electron. s orbitals are spherical in shape. p orbitals have a dumbbell shape and are oriented along the x, y, and z axes (p_x, p_y, p_z). d orbitals (important for transition metals) have more complex shapes, often described as cloverleaf, with one d_z² orbital having a different shape.

原子轨道是电子出现概率较高的空间区域。s 轨道呈球形。p 轨道呈哑铃形，并沿 x、y、z 轴取向（p_x、p_y、p_z）。d 轨道（对过渡金属重要）形状更复杂，常描述为四叶草形，其中一个 d_z² 轨道形状不同。

3. The Aufbau Principle | 构造原理

The Aufbau principle states that electrons occupy the lowest energy orbitals available. The order of filling orbitals is not simply by principal quantum number due to overlapping of energy levels. The sequence (increasing energy) is: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p < 7s < 5f < 6d < 7p. A useful mnemonic is the diagonal rule (Madelung's rule).

构造原理表明电子首先占据能量最低的可用轨道。轨道填充顺序并非严格按照主量子数递增，因为能级会发生交错。填充顺序（能量递增）为：1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p < 7s < 5f < 6d < 7p。一个有用的记忆方法是斜线规则（马德隆规则）。

4. Hund’s Rule and Pauli Exclusion Principle | 洪特规则与泡利不相容原理

Pauli Exclusion Principle: No two electrons in an atom can have the same set of four quantum numbers. This means an atomic orbital can hold at most two electrons, and they must have opposite spins (represented as ↑ and ↓).

泡利不相容原理：原子中不能有两个电子具有完全相同的四个量子数。这意味着一个原子轨道最多容纳两个电子，且它们的自旋方向必须相反（用 ↑ 和 ↓ 表示）。

Hund’s Rule: When filling degenerate orbitals (orbitals of the same energy, e.g., the three p orbitals), electrons fill each orbital singly with parallel spins before pairing up. This minimises electron-electron repulsion and gives the lowest energy arrangement.

洪特规则：在填充简并轨道（能量相同的轨道，如三个p轨道）时，电子会先以自旋平行的方式单独占据每个轨道，然后再配对。这样可以最小化电子间排斥力，使能量最低。

Carbon 2p orbitals: ↑ ↑ _ (not ↑↓ _ _)

5. Writing Full Electron Configurations | 书写完整电子排布

To write the electron configuration of an element, list the occupied sub-levels in order of increasing energy, with the number of electrons as a superscript. For example, oxygen (Z=8): 1s² 2s² 2p⁴. Calcium (Z=20): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s². Note that 4s is filled before 3d because it is lower in energy for neutral atoms.

要书写元素的电子排布，应按能量递增顺序列出被占据的亚层，并以右上标数字表示电子数。例如，氧（原子序数8）：1s² 2s² 2p⁴。钙（原子序数20）：1s² 2s² 2p⁶ 3s² 3p⁶ 4s²。注意，对于中性原子，4s 的能量低于 3d，因此优先填充 4s。

For transition metals, the 3d sub-level is filled after 4s. Examples: Sc (Z=21): 1s² 2s² 2p^{Published by TutorHao | IB Chemistry Revision Series | aleveler.com}

更多咨询请联系16621398022（同微信）