📚 IB Chemistry: Overview of Atomic Structure | IB化学：物质结构概览
In this revision guide, we explore the foundational concepts of atomic structure in IB Chemistry. From subatomic particles to periodic trends and bonding models, this overview is designed to help you grasp the essential ideas quickly and accurately.
本复习指南将带您系统梳理IB化学中物质结构的核心概念。从亚原子粒子到周期趋势和成键模型，旨在帮助您快速、准确地掌握考点必备知识。
1. Atomic Structure and Subatomic Particles | 原子结构与亚原子粒子
Atoms consist of three fundamental particles: protons, neutrons, and electrons. Protons carry a positive charge (+1), electrons carry a negative charge (−1), and neutrons are neutral. Protons and neutrons are located in the nucleus, while electrons occupy the surrounding electron cloud.
原子由三种基本粒子构成：质子、中子和电子。质子带正电荷（+1），电子带负电荷（−1），中子呈电中性。质子和中子位于原子核中，而电子分布在周围的电子云中。
- Proton: relative mass 1, charge +1
- Neutron: relative mass 1, charge 0
- Electron: relative mass 1/1836, charge −1
- 质子：相对质量1，电荷+1
- 中子：相对质量1，电荷0
- 电子：相对质量1/1836，电荷−1
The number of protons defines the element, while the number of electrons in a neutral atom equals the number of protons.
质子数决定元素种类，而中性原子中的电子数等于质子数。
2. 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. The mass number (A) is the total number of protons and neutrons. Atoms of the same element with different numbers of neutrons are called isotopes.
原子序数（Z）是原子核中的质子数。质量数（A）是质子数和中子数之和。同种元素具有不同中子数的原子称为同位素。
A = Z + N
For example, carbon-12 and carbon-14 both have Z = 6, but carbon-12 has 6 neutrons while carbon-14 has 8 neutrons. Isotopes have identical chemical properties but different physical properties due to mass differences.
例如，碳-12和碳-14的Z都是6，但碳-12有6个中子，碳-14有8个中子。同位素化学性质相同，但因质量不同而物理性质有差异。
3. Electron Configuration and Energy Levels | 电子排布与能级
Electrons are arranged in energy levels (shells) around the nucleus. The maximum number of electrons in a principal shell is given by 2n², where n is the principal quantum number.
电子围绕原子核分布在能量层（壳层）中。主壳层最多可容纳的电子数为2n²，其中n是主量子数。
IB Chemistry expects you to know the electron configurations for elements up to Z = 20. For example, calcium (Z = 20) has the configuration: 2, 8, 8, 2 or 1s² 2s² 2p⁶ 3s² 3p⁶ 4s².
IB化学要求掌握Z = 20以内元素的电子排布。例如，钙（Z = 20）的电子排布为：2, 8, 8, 2 或 1s² 2s² 2p⁶ 3s² 3p⁶ 4s²。
Remember the order of filling: 1s → 2s → 2p → 3s → 3p → 4s → 3d. The 4s orbital is filled before the 3d orbital because it has a lower energy.
记住填充顺序：1s → 2s → 2p → 3s → 3p → 4s → 3d。4s轨道先于3d轨道填充，因为其能量更低。
4. Atomic Orbitals and Quantum Numbers | 原子轨道与量子数
Atomic orbitals are regions of space where the probability of finding an electron is high. The main types are s, p, d, and f orbitals. An s orbital is spherical, while p orbitals are dumbbell-shaped and oriented along the x, y, and z axes.
原子轨道是电子出现概率高的空间区域。主要类型有s、p、d、f轨道。s轨道呈球形，p轨道呈哑铃形，并沿x、y、z轴方向取向。
- s orbital: 1 orientation, holds 2 electrons
- p orbitals: 3 orientations, hold 6 electrons
- d orbitals: 5 orientations, hold 10 electrons
- s轨道：1个取向，容纳2个电子
- p轨道：3个取向，容纳6个电子
- d轨道：5个取向，容纳10个电子
The quantum numbers describe the size, shape, and orientation of orbitals. In IB, you mainly need to understand principal quantum number (n) and orbital shapes (l and m).
量子数描述轨道的大小、形状和取向。在IB中，主要需要理解主量子数（n）以及轨道形状（l和m）。
5. Ionization Energy | 电离能
First ionization energy is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous ions with a +1 charge.
第一电离能是指从一摩尔气态原子中移除一摩尔电子，形成一摩尔带+1电荷的气态离子所需的能量。
X(g) → X⁺(g) + e⁻
Ionization energy increases across a period due to increasing nuclear charge and decreasing atomic radius. It decreases down a group because the outer electrons are farther from the nucleus and are more shielded.
电离能在一个周期内从左到右增大，原因是核电荷增加和原子半径减小。在同一族内从上到下减小，因为外层电子离核更远且屏蔽更强。
6. Electron Affinity and Electronegativity | 电子亲和能与电负性
Electron affinity is the energy change when an electron is added to a gaseous atom. It generally becomes more negative across a period as atoms approach the noble gas configuration.
电子亲和能是气态原子获得一个电子时的能量变化。在一个周期内从左到右变得更负，因为原子趋于达到稀有气体构型。
Electronegativity is the ability of an atom to attract electrons in a chemical bond. Fluorine has the highest electronegativity (4.0 on the Pauling scale). This trend mirrors ionization energy: it increases across a period and decreases down a group.
电负性是原子在化学键中吸引电子的能力。氟的电负性最高（鲍林标度为4.0）。该趋势与电离能一致：同周期增大，同族减小。
7. Periodic Trends | 周期趋势
Several key properties show clear trends across a period and down a group. These include atomic radius, ionic radius, ionization energy, electron affinity, and electronegativity.
许多关键性质在周期和族内有明显趋势，包括原子半径、离子半径、电离能、电子亲和能和电负性。
| Property | Across a period → | Down a group ↓ |
| Atomic radius | Decreases | Increases |
| Ionization energy | Increases | Decreases |
| Electronegativity | Increases | Decreases |
These trends arise from the interplay between nuclear charge, shielding, and the distance of outer electrons from the nucleus.
这些趋势源于核电荷、屏蔽效应以及外层电子与核距离之间的相互作用。
8. Ionic and Covalent Bonding | 离子键与共价键
Ionic bonding occurs when electrons are transferred from a metal to a non-metal, forming cations and anions held together by electrostatic attraction. For example, sodium chloride: Na → Na⁺ + e⁻; Cl + e⁻ → Cl⁻.
离子键发生在电子从金属转移至非金属时，形成通过静电吸引联系在一起的阳离子和阴离子。例如氯化钠：Na → Na⁺ + e⁻；Cl + e⁻ → Cl⁻。
Covalent bonding occurs when two non-metals share electrons. A single bond shares one pair, a double bond shares two pairs, and a triple bond shares three pairs. Electronegativity differences determine bond polarity.
共价键发生在两个非金属共享电子时。单键共享一对电子，双键共享两对，三键共享三对。电负性差异决定键的极性。
A polar covalent bond forms when the electronegativity difference is between 0.5 and 1.7. Non-polar covalent bonds form when the difference is less than 0.5.
当电负性差在0.5至1.7之间时形成极性共价键。当差值小于0.5时形成非极性共价键。
9. Molecular Geometry and VSEPR Theory | 分子几何与VSEPR理论
The Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular shapes based on the idea that electron pairs around a central atom repel each other and arrange themselves as far apart as possible.
价层电子对互斥（VSEPR）理论通过中心原子周围的电子对尽量远离彼此的原则来预测分子形状。
- 2 electron pairs: linear (180°), e.g. BeCl₂
- 3 electron pairs: trigonal planar (120°), e.g. BF₃
- 4 electron pairs: tetrahedral (109.5°), e.g. CH₄
- 5 electron pairs: trigonal bipyramidal (90° and 120°), e.g. PCl₅
- 6 electron pairs: octahedral (90°), e.g. SF₆
- 2个电子对：直线形（180°），如BeCl₂
- 3个电子对：平面三角形（120°），如BF₃
- 4个电子对：四面体形（109.5°），如CH₄
- 5个电子对：三角双锥形（90°和120°），如PCl₅
- 6个电子对：八面体形（90°），如SF₆
Lone pairs occupy more space than bonding pairs, causing bond angles to be slightly reduced. For example, NH₃ has a bond angle of 107° instead of 109.5°, and H₂O has 104.5°.
孤对电子占据的空间比成键电子对大，导致键角略微减小。例如，NH₃的键角为107°而非109.5°，H₂O的键角为104.5°。
10. Hybridization | 杂化轨道
Hybridization explains the observed geometry of covalent molecules by mixing atomic orbitals to form new hybrid orbitals. The number of hybrid orbitals formed equals the number of electron domains around the central atom.
杂化轨道理论通过混合原子轨道形成新的杂化轨道，来解释共价分子观测到的几何形状。杂化轨道数等于中心原子周围的电子区域数。
| Electron domains | Hybridization | Geometry |
| 2 | sp | Linear |
| 3 | sp² | Trigonal planar |
| 4 | sp³ | Tetrahedral |
For example, carbon in methane (CH₄) uses sp³ hybrid orbitals, while carbon in ethene (C₂H₄) uses sp² hybrid orbitals, leaving a p orbital for the π bond.
例如，甲烷（CH₄）中的碳使用sp³杂化轨道，而乙烯（C₂H₄）中的碳使用sp²杂化轨道，并留下一个p轨道用于π键。
11. Intermolecular Forces | 分子间作用力
Intermolecular forces (IMFs) are attractions between molecules. They include London dispersion forces, permanent dipole–dipole interactions, and hydrogen bonding. These forces determine physical properties such as boiling point and solubility.
分子间作用力（IMFs）是分子之间的吸引力，包括伦敦色散力、永久偶极-偶极相互作用和氢键。这些作用力决定沸点、溶解度等物理性质。
- London forces: present in all molecules; increase with surface area and electron count.
- Dipole–dipole: occur between polar molecules due to permanent charge separation.
- Hydrogen bonding: the strongest IMF, occurring between H and N, O, or F.
- 伦敦色散力：存在于所有分子中，随表面积和电子数增加而增强。
- 偶极-偶极相互作用：发生在极性分子之间，源于永久电荷分离。
- 氢键：最强的分子间作用力，出现在H与N、O或F之间。
For example, water has a higher boiling point than H₂S because oxygen is more electronegative than sulfur, and H₂O forms stronger hydrogen bonds.
例如，水的沸点高于H₂S，因为氧比硫电负性更强，H₂O能形成更强的氢键。
12. Metallic Bonding and Alloys | 金属键与合金
Metallic bonding involves the electrostatic attraction between the lattice of positive ions and the delocalized “sea of electrons.” This model explains the electrical conductivity, malleability, and ductility of metals.
金属键是正离子晶格和离域“电子海”之间的静电吸引。该模型解释了金属的导电性、延展性和可锻性。
In an alloy, atoms of different sizes disrupt the regular arrangement of the pure metal lattice, making it harder for the layers to slide over each other. Therefore, alloys are usually harder and stronger than the pure metal.
在合金中，不同大小的原子破坏了纯金属晶格的规则排列，使层间滑动更加困难。因此，合金通常比纯金属更硬更强。
Common examples include brass (copper and zinc) and steel (iron with carbon and other elements).
常见例子包括黄铜（铜和锌）和钢（铁中加入碳及其他元素）。
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