📚 IB Chemistry: Periodic Table Key Points | IB 化学：元素周期表 考点精讲
The periodic table is the chemist’s most powerful organisational tool. By arranging elements in order of increasing atomic number, patterns in electron configuration give rise to periodic trends that allow us to predict and explain physical and chemical properties. For IB Chemistry, mastering these trends—from atomic radius and ionisation energy to the behaviour of period 3 oxides—is essential for Paper 1, Paper 2, and the Internal Assessment.
元素周期表是化学家最强大的组织工具。按原子序数递增排列元素，电子排布的模式产生了周期性规律，使我们能预测和解释物理和化学性质。对于 IB 化学，掌握这些趋势——从原子半径、电离能到第三周期氧化物的行为——是应对试卷一、试卷二和内部评估的关键。
1. Structure of the Periodic Table | 周期表的结构
The table is divided into blocks (s, p, d, f) based on the electron subshell being filled. Periods are horizontal rows; groups are vertical columns. Elements in the same group have the same number of valence electrons, leading to similar chemical properties.
周期表根据电子填充的亚层分为 s 区、p 区、d 区和 f 区。横行是周期，纵列是族。同族元素具有相同的价电子数，因此化学性质相似。
The s-block includes Groups 1 and 2 (alkali and alkaline earth metals), where s orbitals are being filled. The p-block spans Groups 13 to 18, filling p orbitals. The d-block contains transition metals (Groups 3–12), filling d orbitals; f-block elements (lanthanides and actinides) fill f orbitals.
s 区包括第 1 和 2 族（碱金属和碱土金属），填充 s 轨道。p 区覆盖第 13 至 18 族，填充 p 轨道。d 区含有过渡金属（第 3 至 12 族），填充 d 轨道；f 区元素（镧系和锕系）填充 f 轨道。
2. Atomic Radius Across a Period | 同周期原子半径的变化
Across a period from left to right, atomic radius generally decreases. This occurs because the nuclear charge (number of protons) increases while electrons are added to the same principal energy level. The increasing effective nuclear charge pulls the electron cloud closer to the nucleus.
从左到右穿过一个周期，原子半径通常减小。这是因为核电荷（质子数）增加，而电子添加到同一个主能层。逐渐增强的有效核电荷将电子云更紧地拉向原子核。
For example, in Period 3: Na (186 pm) > Mg (160 pm) > Al (143 pm) > Si (118 pm) > P (110 pm) > S (104 pm) > Cl (99 pm). Note that the noble gas at the end (Ar, 98 pm) is measured differently but still fits the general trend.
例如，第三周期：Na (186 pm) > Mg (160 pm) > Al (143 pm) > Si (118 pm) > P (110 pm) > S (104 pm) > Cl (99 pm)。注意末尾的稀有气体（Ar, 98 pm）测量方式不同，但仍符合总体趋势。
3. Atomic Radius Down a Group | 同族原子半径的变化
Going down a group, atomic radius increases. Each successive element has an extra electron shell, so the outer electrons are farther from the nucleus despite the increase in nuclear charge. The shielding effect of inner electrons reduces the effective nuclear attraction on valence electrons.
沿族向下，原子半径增大。每向下一周期就增加一个电子层，因此外层电子离核更远，尽管核电荷也在增加。内层电子的屏蔽效应减弱了核对价电子的有效吸引。
For Group 1: Li (152 pm) < Na (186 pm) < K (231 pm) < Rb (244 pm) < Cs (262 pm). This trend explains the increasing reactivity of alkali metals down the group.
第 1 族：Li (152 pm) < Na (186 pm) < K (231 pm) < Rb (244 pm) < Cs (262 pm)。这一趋势解释了碱金属由上到下反应性增强的原因。
4. Ionic Radius | 离子半径
Cations are smaller than their parent atoms because the loss of valence electrons reduces electron-electron repulsion and often leaves one fewer occupied shell. Anions are larger than their parent atoms because the addition of electrons increases electron-electron repulsion, causing the cloud to expand.
阳离子比其母原子小，因为失去价电子减少了电子间排斥力，并往往减少一个占据的电子层。阴离子比其母原子大，因为电子增加使电子间排斥增大，导致电子云膨胀。
For an isoelectronic series (same number of electrons), ionic radius decreases with increasing nuclear charge. Example: N³⁻ (146 pm) > O²⁻ (140 pm) > F⁻ (133 pm) > Na⁺ (102 pm) > Mg²⁺ (72 pm) > Al³⁺ (53.5 pm).
在等电子系列（电子数相同）中，离子半径随核电荷增加而减小。例如：N³⁻ (146 pm) > O²⁻ (140 pm) > F⁻ (133 pm) > Na⁺ (102 pm) > Mg²⁺ (72 pm) > Al³⁺ (53.5 pm)。
5. First Ionisation Energy | 第一电离能
Ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms: X(g) → X⁺(g) + e⁻. It is a measure of how strongly an atom holds onto its outermost electron.
电离能是指从 1 mol 气态原子中移走 1 mol 电子所需的最低能量：X(g) → X⁺(g) + e⁻。它衡量原子对其最外层电子的束缚强度。
Across a period, first ionisation energy generally increases because of increasing effective nuclear charge and decreasing atomic radius. However, there are small dips between Groups 2 and 13 (e.g., Be to B) and between Groups 15 and 16 (e.g., N to O) due to electron pairing and subshell stability.
穿过一个周期，第一电离能总体上升，因为有效核电荷增加且原子半径减小。然而，在第 2 族与第 13 族之间（如 Be 到 B）以及第 15 族与第 16 族之间（如 N 到 O）会出现小幅下降，原因是电子配对和亚层稳定性。
Down a group, first ionisation energy decreases because the outer electron is farther from the nucleus and experiences more shielding, making it easier to remove.
沿族向下，第一电离能降低，因为外层电子离核更远，屏蔽效应更强，使之更易失去。
6. Successive Ionisation Energies | 逐级电离能
Successive ionisation energies provide evidence for electron shell structure. A large jump in ionisation energy indicates that an electron is being removed from a filled inner shell, closer to the nucleus. For example, for magnesium (1s² 2s² 2p⁶ 3s²), there is a big jump between the second and third ionisation energies because the third electron is removed from the 2p subshell.
逐级电离能为电子层结构提供了证据。电离能的大幅跳跃表明电子正从已填满的内层、离核更近的轨道上移走。例如，对于镁（1s² 2s² 2p⁶ 3s²），第二与第三电离能之间存在大幅跃升，因为第三个电子取自 2p 亚层。
7. Electron Affinity | 电子亲和能
Electron affinity is the energy change when one mole of electrons is added to one mole of gaseous atoms: X(g) + e⁻ → X⁻(g). Exothermic values (negative ΔH) indicate favourable electron capture. Across a period, electron affinity generally becomes more exothermic due to increased effective nuclear charge. Down a group, it becomes less exothermic because the added electron is farther from the nucleus.
电子亲和能是 1 mol 气态原子获得 1 mol 电子时的能量变化：X(g) + e⁻ → X⁻(g)。放热值（负 ΔH）表示更容易捕获电子。穿过一个周期，电子亲和能通常变得更放热，因为有效核电荷增大。沿族向下，放热减少，因为新增电子离核更远。
Notable exceptions: Group 2 elements have endothermic first electron affinities because the added electron would enter a new p subshell, opposed by the stability of a full s subshell. Group 15 elements also have low or endothermic values because a half-filled p subshell provides extra stability.
值得注意的例外：第 2 族元素的第一电子亲和能为吸热，因为加入的电子要进入新的 p 亚层，而全满的 s 亚层的稳定性对此不利。第 15 族元素的值也较低或吸热，因为半满的 p 亚层提供了额外稳定性。
8. Electronegativity | 电负性
Electronegativity is the ability of an atom to attract a bonding pair of electrons. The Pauling scale is commonly used. Fluorine (3.98) is the most electronegative element. Across a period, electronegativity increases with effective nuclear charge. Down a group, electronegativity decreases because the bonding pair is farther from the nucleus and more shielded.
电负性是原子吸引成键电子对的能力。常用鲍林标度。氟（3.98）是电负性最高的元素。穿过一个周期，电负性随有效核电荷增加而升高。沿族向下，电负性降低，因为成键电子对离核更远且屏蔽更强。
Electronegativity differences help predict bond type: ΔEN < 0.5 → non-polar covalent; 0.5 ≤ ΔEN < 1.7 → polar covalent; ΔEN ≥ 1.7 → largely ionic. Beware of borderline cases like AlCl₃ (ΔEN ~1.5) which is covalent in gaseous state but ionic lattice in solid.
电负性差值有助于预测键型：ΔEN < 0.5 → 非极性共价键；0.5 ≤ ΔEN < 1.7 → 极性共价键；ΔEN ≥ 1.7 → 以离子键为主。需注意边界情况，如 AlCl₃（ΔEN ≈ 1.5），气态时为共价分子，固态时形成离子晶格。
9. Melting Points of Period 3 Elements | 第三周期元素的熔点
The melting points of Period 3 elements vary in a characteristic pattern linked to structure and bonding. Na (98 °C), Mg (650 °C) and Al (660 °C) are metallic; their melting points rise from Na to Al due to increasing charge of the positive ions and increasing number of delocalised electrons, strengthening the metallic bond.
第三周期单质的熔点因结构和键合方式的不同而呈特征性变化。Na (98 °C)、Mg (650 °C) 和 Al (660 °C) 是金属晶体；从 Na 到 Al 熔点升高，是因为阳离子电荷增加、离域电子数增多，增强了金属键。
Silicon has a giant covalent structure (diamond-like) with a very high melting point (~1410 °C). Phosphorus (P₄), sulfur (S₈) and chlorine (Cl₂) exist as simple molecular substances with weak van der Waals’ forces, so their melting points are low. Argon has the lowest melting point as monatomic with only weak London dispersion forces.
硅具有巨型共价结构（类金刚石），熔点极高 (~1410 °C)。白磷 (P₄)、硫 (S₈) 和氯 (Cl₂) 均为简单分子，分子间存在微弱的范德华力，因此熔点很低。氩是单原子分子，只有微弱的伦敦色散力，熔点最低。
Trend: Na < Mg > Al << Si >> P₄ < S₈ < Cl₂ < Ar (lowest)
10. Period 3 Oxides – Acid-Base Behaviour | 第三周期氧化物——酸碱行为
Moving across Period 3, oxides change from basic to amphoteric to acidic. Sodium oxide (Na₂O) and magnesium oxide (MgO) are basic, reacting with water to form alkaline hydroxides and neutralising acids.
沿第三周期从左到右，氧化物从碱性经两性变为酸性。氧化钠 (Na₂O) 和氧化镁 (MgO) 是碱性氧化物，与水反应生成碱性氢氧化物，并能中和酸。
Aluminium oxide (Al₂O₃) is amphoteric, reacting with both acids and bases: Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O and Al₂O₃ + 2NaOH + 3H₂O → 2NaAl(OH)₄. Silicon dioxide (SiO₂) is weakly acidic (reacts slowly with hot concentrated alkalis). Phosphorus(V) oxide (P₄O₁₀), sulfur trioxide (SO₃) and dichlorine heptoxide (Cl₂O₇) are acidic, forming oxyacids with water.
氧化铝 (Al₂O₃) 是两性氧化物，既能与酸也能与碱反应：Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O；Al₂O₃ + 2NaOH + 3H₂O → 2NaAl(OH)₄。二氧化硅 (SiO₂) 呈弱酸性（与热的浓碱缓慢反应）。五氧化二磷 (P₄O₁₀)、三氧化硫 (SO₃) 和七氧化二氯 (Cl₂O₇) 均为酸性氧化物，与水生成含氧酸。
| Oxide | Nature | Reaction with water |
|---|---|---|
| Na₂O | Basic | Na₂O + H₂O → 2NaOH |
| MgO | Basic (slightly soluble) | MgO + H₂O ⇌ Mg(OH)₂ |
| Al₂O₃ | Amphoteric | No reaction with water |
| SiO₂ | Weakly acidic | No reaction with water |
| P₄O₁₀ | Acidic | P₄O₁₀ + 6H₂O → 4H₃PO₄ |
| SO₃ | Acidic | SO₃ + H₂O → H₂SO₄ |
11. Period 3 Chlorides – Structure and Hydrolysis | 第三周期氯化物——结构与水解
Chlorides of Period 3 also show a pattern linked to bonding type. Sodium chloride (NaCl) and magnesium chloride (MgCl₂) are ionic solids. Aluminium chloride (AlCl₃) exists as a dimer Al₂Cl₆ in the gas phase but behaves covalently; it hydrolyses violently. Silicon tetrachloride (SiCl₄) and phosphorus trichloride (PCl₃)/pentachloride (PCl₅) are covalent liquids that hydrolyse readily.
第三周期的氯化物同样呈现出与键型相关的规律。氯化钠 (NaCl) 和氯化镁 (MgCl₂) 是离子型固体。氯化铝 (AlCl₃) 在气态中以二聚体 Al₂Cl₆ 形式存在，但表现出共价性，遇水剧烈水解。四氯化硅 (SiCl₄)、三氯化磷 (PCl₃) 和五氯化磷 (PCl₅) 是共价液体，容易水解。
Hydrolysis examples: SiCl₄ + 2H₂O → SiO₂ + 4HCl; PCl₃ + 3H₂O → H₃PO₃ + 3HCl; PCl₅ + H₂O → POCl₃ + 2HCl (partial); AlCl₃ dissolves with hissing sound and forms acid solution. NaCl and MgCl₂ simply dissolve without significant hydrolysis (MgCl₂ shows slight hydrolysis due to Mg²⁺ polarisation).
水解举例：SiCl₄ + 2H₂O → SiO₂ + 4HCl；PCl₃ + 3H₂O → H₃PO₃ + 3HCl；PCl₅ + H₂O → POCl₃ + 2HCl (部分水解)；AlCl₃ 溶于水发出嘶嘶声并形成酸性溶液。NaCl 和 MgCl₂ 只是溶解，无明显水解（MgCl₂ 因 Mg²⁺ 极化作用有微弱水解）。
12. Transition Metal Characteristics (d-block) | 过渡金属的特性（d 区）
Transition metals are defined as elements that have an incomplete d subshell in either their atoms or common ions. Zn is not a transition metal because it has a full d¹⁰ configuration in both atom and Zn²⁺ ion. Key IB properties: variable oxidation states, coloured compounds, catalytic activity, and formation of complex ions.
过渡金属是指其原子或常见离子具有未填满的 d 亚层的元素。Zn 不是过渡金属，因为它的原子和 Zn²⁺ 离子都具有全满 d¹⁰ 构型。IB 核心性质：可变氧化态、有色化合物、催化活性以及形成配离子。
Variable oxidation states arise because the energy difference between 4s and 3d is small, allowing successive removal of electrons. For example, Mn shows +2, +3, +4, +6, +7 oxidation states. Catalytic activity (e.g., Fe in Haber process, V₂O₅ in Contact process, Ni in hydrogenation) results from the ability to change oxidation state and adsorb reactants.
可变氧化态是由于 4s 与 3d 轨道能量相近，可以逐步失去不同数目的电子。例如 Mn 表现出 +2, +3, +4, +6, +7 氧化态。催化活性（如 Haber 法中的 Fe，接触法中的 V₂O₅，氢化反应中的 Ni）源于其能够改变氧化态并吸附反应物。
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