The Evolution of Atomic Models | 原子模型及其演变

📚 The Evolution of Atomic Models | 原子模型及其演变

The atomic model has undergone a remarkable transformation over the past two centuries, driven by experimental discoveries and evolving theories. From Dalton’s indivisible solid spheres to the modern quantum mechanical picture of electron clouds, each model has built upon its predecessor while correcting its limitations. This article traces this evolution, highlighting the key experiments, scientists, and concepts that are essential for A-Level chemistry.

原子模型在过去两个世纪中经历了非凡的演变，实验发现和理论发展不断推动着这一进程。从道尔顿不可分割的实心球体到现代量子力学中电子云的图像，每一个模型都在继承前人成果的同时修正了其局限性。本文梳理了这一演变过程，重点介绍A-Level化学中必需的实验、科学家和概念。

1. Early Ideas and Dalton’s Atomic Model | 早期思想与道尔顿原子模型

The concept of atoms dates back to ancient Greek philosophers like Democritus, who proposed that matter is composed of indivisible particles called “atomos”. However, these ideas were purely philosophical and lacked experimental support. It was John Dalton who, in 1803, formulated the first scientific atomic theory based on the law of conservation of mass and the law of constant composition.

原子的概念可以追溯到古希腊哲学家，如德谟克利特，他提出物质由不可分割的粒子（称为“原子”）组成。然而，这些想法纯粹是哲学性的，缺乏实验支持。直到1803年，约翰·道尔顿基于质量守恒定律和定比定律，提出了第一个科学的原子理论。

Dalton proposed that atoms are tiny, indivisible, and indestructible spheres. All atoms of a given element are identical in mass and properties, while atoms of different elements differ. Chemical reactions involve the rearrangement of atoms, but atoms themselves are not created or destroyed. This model explained the laws of chemical combination and provided a foundation for stoichiometry.

道尔顿提出原子是微小、不可分割且不可摧毁的球体。同一种元素的所有原子在质量和性质上完全相同，而不同元素的原子则不同。化学反应涉及原子的重新排列，但原子本身不会创生或毁灭。该模型解释了化学结合定律，并为化学计量学奠定了基础。

2. Discovery of the Electron and Thomson’s Plum Pudding Model | 电子的发现与汤姆逊葡萄干布丁模型

In 1897, J.J. Thomson discovered the electron through his experiments with cathode rays. He found that these rays consisted of negatively charged particles with a very small mass-to-charge ratio, about 1800 times lighter than a hydrogen atom. This discovery proved that atoms are not indivisible, contradicting Dalton’s postulate.

1897年，J.J. 汤姆逊通过阴极射线实验发现了电子。他发现这些射线由带负电的粒子组成，其质荷比非常小，质量约为氢原子的1/1800。这一发现证明原子并非不可分割，反驳了道尔顿的假设。

Thomson proposed the “plum pudding” model, in which the atom is a sphere of positive charge with electrons embedded throughout, like raisins in a pudding. The total positive charge balances the negative charge of the electrons, making the atom electrically neutral. This model explained that atoms could be ionized by gaining or losing electrons, but it could not account for the scattering of alpha particles observed later.

汤姆逊提出了“葡萄干布丁”模型，认为原子是一个带正电荷的球体，电子像布丁中的葡萄干一样镶嵌其中。总正电荷与电子的负电荷相平衡，使原子呈电中性。该模型解释了原子可以通过获得或失去电子而电离，但无法解释后来观察到的α粒子散射现象。

3. Rutherford’s Gold Foil Experiment and Nuclear Model | 卢瑟福金箔实验与核模型

In 1909, Ernest Rutherford and his colleagues Hans Geiger and Ernest Marsden performed the famous gold foil experiment. They fired a beam of alpha particles at a very thin sheet of gold foil and observed their deflections using a fluorescent screen. Most alpha particles passed straight through, but a small fraction were deflected at large angles, and about 1 in 8000 bounced back.

1909年，欧内斯特·卢瑟福与同事汉斯·盖革和欧内斯特·马斯登进行了著名的金箔实验。他们用α粒子束轰击极薄的金箔，并通过荧光屏观察粒子的偏转。大多数α粒子直接穿过，但一小部分以大角度偏转，约1/8000的粒子被反弹回来。

To explain these results, Rutherford proposed a new model in 1911. The atom consists of a very tiny, dense, positively charged nucleus at the centre, containing nearly all the mass. Electrons orbit the nucleus at relatively large distances, making the atom mostly empty space. This nuclear model accounted for the large-angle scattering: a rare alpha particle that came close to the concentrated positive charge would be strongly repelled. The radius of the nucleus is about 10⁻¹⁵ m, whereas the atom’s radius is about 10⁻¹⁰ m.

为解释这些结果，卢瑟福于1911年提出了新模型。原子中心是一个极其微小、致密且带正电荷的原子核，几乎包含了全部质量。电子在较远距离绕核运动，因此原子大部分是空的。该核模型解释了大角度散射：少数α粒子靠近集中的正电荷时会被强烈排斥。原子核半径约为10⁻¹⁵米，而原子半径约为10⁻¹⁰米。

4. Bohr’s Model and Energy Levels | 玻尔模型与能级

Rutherford’s model could not explain why electrons, which are accelerating around the nucleus, do not spiral into the nucleus and emit continuous radiation. In 1913, Niels Bohr resolved this paradox by applying quantum ideas to the hydrogen atom. Bohr proposed that electrons move in fixed circular orbits, or energy levels, around the nucleus, each with a specific energy. As long as an electron stays in a given orbit, it does not emit or absorb energy.

卢瑟福模型无法解释为什么绕核加速运动的电子不会螺旋落入原子核并发出连续辐射。1913年，尼尔斯·玻尔将量子思想应用于氢原子，解决了这一矛盾。玻尔提出电子在核外固定的圆形轨道（即能级）上运动，每个轨道具有特定的能量。只要电子停留在某一轨道，它就不会发射或吸收能量。

Energy is emitted or absorbed only when an electron jumps from one orbit to another. The energy difference ΔE between two levels corresponds to a photon of frequency ν, given by ΔE = hν, where h is Planck’s constant. Bohr’s model successfully explained the hydrogen emission spectrum, predicting the wavelengths of the Balmer series and other spectral lines. However, it failed for atoms with more than one electron, because electron-electron repulsion complicates the simple picture.

只有当电子从一个轨道跃迁到另一个轨道时，才会发射或吸收能量。两个能级之间的能量差ΔE对应频率为ν的光子，关系为ΔE = hν，其中h是普朗克常数。玻尔模型成功解释了氢原子发射光谱，预测了巴尔末系等谱线的波长。然而，该模型对多电子原子失效，因为电子间的排斥作用使得简单图像不再适用。

ΔE = hν = hc / λ

Here, c is the speed of light and λ is the wavelength of the emitted or absorbed photon.

其中，c是光速，λ是发射或吸收光子的波长。

5. The Quantum Mechanical Model | 量子力学模型

In the 1920s, Erwin Schrödinger and Werner Heisenberg developed a more advanced model based on quantum mechanics. Schrödinger formulated the wave equation that describes the electron’s behaviour as a wave, replacing Bohr’s fixed orbits with three-dimensional regions of probability called orbitals. Heisenberg’s uncertainty principle states that it is impossible to know both the exact position and momentum of an electron simultaneously.

在20世纪20年代，埃尔温·薛定谔和维尔纳·海森堡发展了基于量子力学的更先进模型。薛定谔提出了描述电子波动行为的波方程，用三维概率区域（即轨道）取代了玻尔的固定轨道。海森堡不确定性原理表明，不可能同时精确知道电子的位置和动量。

In this model, an electron is described by a wave function ψ, whose square ψ² gives the probability density of finding the electron at a particular point. The electron cloud is denser in regions where the electron is most likely to be found. The quantum numbers n, l, mₗ, and mₛ define the energy, shape, orientation, and spin of the orbital. This model accurately predicts the behavior of all atoms, including multi-electron systems, and is the basis of chemical bonding theories.

在该模型中，电子由波函数ψ描述，ψ²给出在特定位置找到电子的概率密度。电子云在电子最可能出现的区域密度更大。量子数n、l、mₗ和mₛ分别定义轨道的能量、形状、取向和自旋。该模型精确预测包括多电子系统在内的所有原子的行为，是化学键理论的基础。

Unlike the Bohr model, quantum mechanics does not define a precise path for the electron. Instead, it provides a probabilistic map of where the electron might be. The s, p, d, and f orbitals correspond to different values of the azimuthal quantum number l = 0, 1, 2, and 3, respectively.

与玻尔模型不同，量子力学不定义电子的精确路径，而是提供电子可能出现的位置的概率图。s、p、d、f轨道分别对应角量子数l = 0、1、2和3。

6. Subatomic Particles: Protons, Neutrons, Electrons | 亚原子粒子：质子、中子、电子

Modern atomic theory recognizes three main subatomic particles within an atom. Protons and neutrons are located in the nucleus, while electrons occupy the surrounding space. The properties of these particles are summarised in the table below.

现代原子理论确认原子内存在三种主要亚原子粒子。质子和中子位于原子核中，而电子占据核外空间。这些粒子的性质总结如下表。

Particle	Relative Mass	Relative Charge	Location
Proton	1	+1	Nucleus
Neutron	1	0	Nucleus
Electron	1/1836	–1	Electron cloud

The proton was discovered by Rutherford in 1919, and the neutron by James Chadwick in 1932. The number of protons determines the element’s identity, while the number of neutrons contributes to the isotope. Electrons are responsible for chemical bonding and the periodic properties of elements.

质子由卢瑟福于1919年发现，中子由詹姆斯·查德威克于1932年发现。质子数决定元素的种类，中子数决定同位素。电子负责化学键合和元素的周期性性质。

7. Atomic Number, Mass Number, and Isotopes | 原子序数、质量数与同位素

The atomic number (Z) is the number of protons in the nucleus of an atom. It is unique for each element and defines its position in the periodic table. The mass number (A) is the total number of protons and neutrons. In nuclear notation, the mass number is written as a superscript and the atomic number as a subscript before the element symbol.

原子序数（Z）是原子核中的质子数。每种元素具有独特的原子序数，决定了它在周期表中的位置。质量数（A）是质子数和中子数的总和。核素符号中，质量数写为左上标，原子序数写为左下标，位于元素符号之前。

²³⁵₉₂U

Isotopes are atoms of the same element that have the same number of protons but different numbers of neutrons. For example, carbon-12 (¹²₆C) and carbon-14 (¹⁴₆C) are both isotopes of carbon. They have identical chemical properties because the electron configuration depends on the proton number, but their physical properties, such as mass and stability, differ. Some isotopes are radioactive, which has applications in radiocarbon dating and medical imaging.

同位素是质子数相同而中子数不同的同一种元素的原子。例如，碳-12（¹²₆C）和碳-14（¹⁴₆C）是碳的同位素。它们的化学性质相同，因为电子排布取决于质子数，但物理性质（如质量和稳定性）不同。一些同位素具有放射性，可用于放射性碳定年和医学成像。

The relative atomic mass of an element is a weighted average of the masses of its isotopes, taking into account their natural abundances. This is calculated using the formula: Aᵣ = Σ (isotopic mass × relative abundance) / 100. This concept is crucial for understanding mass spectra and calculating molar masses.

元素的相对原子质量是其同位素质量的加权平均值，考虑天然丰度。计算公式为：Aᵣ = Σ（同位素质量 × 相对丰度）/ 100。这一概念对于理解质谱和计算摩尔质量至关重要。

8. Electron Configuration and Orbitals | 电子排布与轨道

The arrangement of electrons in an atom is described by its electron configuration. Electrons fill orbitals according to the Aufbau principle, which states that orbitals are occupied in order of increasing energy. The order is 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, and so on. Each orbital can hold a maximum of two electrons with opposite spins, as stated by the Pauli exclusion principle.

原子中电子的排列由其电子排布描述。电子按照构造原理填充轨道，即轨道按能量升高的顺序被占据。顺序为1s、2s、2p、3s、3p、4s、3d、4p等。根据泡利不相容原理，每个轨道最多容纳两个自旋相反的电子。

Hund’s rule states that electrons occupy degenerate orbitals (orbitals of the same energy) singly before pairing up. For example, the electronic configuration of oxygen (Z = 8) is 1s² 2s² 2p⁴. The 2p subshell has three orbitals, and the four electrons fill them with one electron in each of the three orbitals and one extra electron pairing with another.

洪德规则规定，电子在简并轨道（能量相同的轨道）中先单独占据，再配对。例如，氧（Z = 8）的电子排布为1s² 2s² 2p⁴。2p亚层有三个轨道，四个电子先分别占据三个轨道，剩余一个电子与其中一个配对。

Shorthand notation using inert gas cores, such as [Ar] 3d⁶ 4s² for iron (Z = 26), simplifies long configurations. Exceptions exist for certain elements, such as chromium and copper, where a half-filled or fully-filled d subshell provides extra stability, leading to configurations [Ar] 3d⁵ 4s¹ and [Ar] 3d¹⁰ 4s¹, respectively.

使用惰性气体核心的简写符号，如铁（Z = 26）写作[Ar] 3d⁶ 4s²，简化了长排布。某些元素存在例外，如铬和铜，半满或全满的d亚层提供额外稳定性，因此分别采用[Ar] 3d⁵ 4s¹和[Ar] 3d¹⁰ 4s¹。

9. Limitations and Refinements | 局限性与修正

Each atomic model has advanced our understanding but also has limitations. Dalton’s model ignored the existence of subatomic particles. Thomson’s plum pudding model could not explain alpha particle scattering. Rutherford’s nuclear model was unable to account for the stability of atoms and discrete spectral lines. Bohr’s model worked for hydrogen but not for multi-electron atoms, as it neglected electron-electron interactions and assumed circular orbits.

每个原子模型都推进了我们的理解，但也存在局限性。道尔顿模型忽视了亚原子粒子的存在。汤姆逊葡萄干布丁模型无法解释α粒子散射。卢瑟福核模型无法解释原子的稳定性和离散谱线。玻尔模型适用于氢原子，但对多电子原子失效，因为它忽略了电子间的相互作用并假设了圆形轨道。

The quantum mechanical model addresses these issues by incorporating wave-particle duality and probabilistic behavior. However, even this model has constraints; it cannot predict the exact path of an electron, only probabilities. Additionally, relativity effects become significant for heavy elements, requiring corrections such as those in relativistic quantum chemistry.

量子力学模型通过引入波粒二象性和概率行为解决了这些问题。然而，即使该模型也有局限：它无法预测电子的精确路径，只能给出概率。此外，对于重元素，相对论效应变得显著，需要相对论量子化学等修正。

For A-Level chemistry, it is sufficient to understand the key points of each model, their experimental evidence, and the reasons for their eventual replacement. Recognizing these limitations demonstrates a deeper understanding of scientific progress.

对于A-Level化学，理解每个模型的关键点、其实验证据以及被取代的原因就足够了。认识到这些局限性体现了对科学进步的更深层理解。

10. Key Takeaways for Exams | 考试关键要点

To excel in examinations, focus on the following points:

为了在考试中取得好成绩，请关注以下要点：

Dalton’s model: atoms are indivisible spheres; later disproved by electron discovery.

道尔顿模型：原子是不可分割的球体；后被电子发现所否定。
Thomson’s plum pudding model: positive sphere with embedded electrons; based on cathode ray experiments.

汤姆逊葡萄干布丁模型：正电荷球体镶嵌电子；基于阴极射线实验。
Rutherford’s experiment: gold foil alpha scattering led to nuclear model; nucleus is small, dense, and positive.

卢瑟福实验：金箔α散射得出核模型；原子核小而致密，带正电。
Bohr’s model: fixed energy levels and quantized transitions; explains hydrogen spectrum but not multi-electron atoms.

玻尔模型：固定能级和量子化跃迁；解释氢光谱，但不适用于多电子原子。
Quantum model: orbitals and probability; use quantum numbers to describe electrons.

量子模型：轨道和概率；用量子数描述电子。
Know the relative masses and charges of protons, neutrons, and electrons.

记住质子、中子、电子的相对质量和电荷。
Understand atomic number, mass number, and isotopes; calculate relative atomic mass.

理解原子序数、质量数和同位素；计算相对原子质量。
Write electron configurations using Aufbau principle, Pauli exclusion principle, and Hund’s rule, including exceptions.

使用构造原理、泡利不相容原理和洪德规则书写电子排布，包括例外情况。

Additionally, practice interpreting experimental evidence and linking it to model predictions. For example, the number of lines in the hydrogen emission spectrum corresponds to the number of possible electron transitions between energy levels.

此外，练习解读实验证据并将其与模型预测联系起来。例如，氢原子发射光谱中谱线的数量对应于能级之间可能的电子跃迁次数。

Mastery of these concepts will allow you to answer both multiple-choice and long-answer questions confidently. Remember to highlight the evolution of models as a story of scientific progress driven by experimental evidence.

掌握这些概念将帮助您自信地解答选择题和简答题。请记住，将模型演变视为由实验证据驱动的科学进步故事。

Published by TutorHao | Chemistry Revision Series | aleveler.com

更多咨询请联系16621398022（同微信）

The Evolution of Atomic Models | 原子模型及其演变