📚 IB Chemistry: Core Concepts of the Periodic Table | IB化学：元素周期表核心知识梳理
The periodic table is one of the most powerful tools in chemistry. It organises all known elements according to atomic number and electron configuration, allowing us to observe recurring patterns in properties known as periodicity. This article summarises the core concepts of the periodic table that are essential for IB Chemistry.
元素周期表是化学中最强大的工具之一。它按原子序数和电子构型组织所有已知元素，使我们能观察到性质上反复出现的规律，即周期性。本文梳理了IB化学中元素周期表的核心知识点。
1. Structure of the Periodic Table | 周期表的结构
The periodic table is arranged in order of increasing atomic number (Z). It consists of horizontal rows called periods (numbered 1–7) and vertical columns called groups (numbered 1–18). The position of an element is directly linked to its electron configuration.
周期表按原子序数（Z）递增排列。它由横向的周期（编号1–7）和纵向的族（编号1–18）组成。元素的位置直接与其电子构型相关。
Elements are also divided into four blocks based on the highest-energy subshell occupied: s-block, p-block, d-block and f-block. For example, sodium (Na) is an s-block element because its outermost electron is in a 3s orbital.
元素还根据最高能量亚层划分为四个区：s区、p区、d区和f区。例如，钠（Na）是s区元素，因为它的最外层电子位于3s轨道。
| Block | Subshell | Groups |
| s | s | 1–2 |
| p | p | 13–18 |
| d | d | 3–12 |
| f | f | lanthanides and actinides |
2. Electron Configuration and Position | 电子构型与位置
For main-group elements, the period number corresponds to the principal quantum number (n) of the valence shell, and the group number (using modern numbering) indicates the total number of valence electrons. For example, phosphorus (P) is in period 3 and group 15, so its valence configuration is 3s² 3p³.
对于主族元素，周期数对应价电子层的主量子数（n），族号（现代编号）表示价电子总数。例如，磷（P）位于第3周期、第15族，因此其价层构型为3s² 3p³。
Transition metals follow a different rule: the d subshell fills after the s subshell of the same period. Thus, the electron configuration of scandium (Sc) is [Ar] 3d¹ 4s², although the 4s orbital is removed first during ionisation.
过渡金属遵循不同规则：同一周期的d亚层在s亚层之后填充。因此，钪（Sc）的电子构型是[Ar] 3d¹ 4s²，但在电离时首先失去4s电子。
A common mistake is to write the electron configuration for cations incorrectly. For example, Fe²⁺ from Fe ([Ar] 3d⁶ 4s²) becomes [Ar] 3d⁶, not [Ar] 3d⁴ 4s².
常见的错误是写错阳离子的电子构型。例如，Fe（[Ar] 3d⁶ 4s²）形成Fe²⁺时，构型为[Ar] 3d⁶，而不是[Ar] 3d⁴ 4s²。
3. Atomic Radius | 原子半径
Atomic radius is the distance between the nucleus and the outermost electron shell. Across a period, atomic radius decreases because the effective nuclear charge increases as protons are added while the shielding by inner electrons remains roughly constant. The outer electrons are pulled more strongly toward the nucleus.
原子半径是指原子核到最外层电子壳层的距离。同周期从左到右，原子半径减小，因为质子数增加导致有效核电荷增大，而内层电子提供的屏蔽效应大致不变，外层电子被更强烈地拉向原子核。
Down a group, atomic radius increases because each additional period adds a new electron shell, making the atom larger despite the increased nuclear charge. For example, radius increases from Li to Cs.
同族从上到下，原子半径增大，因为每增加一个周期就增加一个新的电子壳层，使原子变大，尽管核电荷也增加。例如，从Li到Cs半径增大。
Effective nuclear charge (Zₑff) is the net positive charge felt by an outer electron. It is calculated as Zₑff = Z − S, where S is the shielding constant.
有效核电荷（Zₑff）是外层电子感受到的净正电荷，计算公式为Zₑff = Z − S，其中S为屏蔽常数。
Zₑff = Z − S
4. Ionization Energy | 电离能
First ionisation energy (IE₁) is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions. It is always positive because energy is absorbed.
第一电离能（IE₁）是指从一摩尔气态原子中各移去一摩尔电子，形成一摩尔气态一价正离子所需的能量。它始终为正值，因为需要吸收能量。
X(g) → X⁺(g) + e⁻
Across a period, ionisation energy generally increases due to increasing effective nuclear charge and decreasing atomic radius. However, there are exceptions. In period 2, boron has a lower IE₁ than beryllium because the 2p electron in boron is higher in energy and easier to remove than a 2s electron. Similarly, oxygen has a lower IE₁ than nitrogen because the paired 2p electrons in oxygen experience repulsion, making one easier to remove.
同周期内，电离能总体上增大，因为有效核电荷增大且原子半径减小。但存在例外。在第2周期中，硼的IE₁低于铍，因为硼的2p电子能量更高，比2s电子更容易失去。同样，氧的IE₁低于氮，因为氧的2p电子成对后存在排斥作用，使其中一个更容易失去。
Down a group, ionisation energy decreases because the outermost electron is farther from the nucleus and more shielded by inner shells.
同族中，电离能从下到上减小，因为最外层电子离原子核更远，受到内层电子的屏蔽更强。
5. Electron Affinity and Electronegativity | 电子亲和能与电负性
Electron affinity (EA) is the energy change when an electron is added to a gaseous atom. First electron affinity is usually exothermic because the added electron is attracted to the nucleus. For example, chlorine has a very negative EA, which explains why it readily forms Cl⁻.
电子亲和能（EA）是向气态原子加入一个电子时的能量变化。第一电子亲和能通常是放热的，因为加入的电子受到原子核的吸引。例如，氯的电子亲和能非常负，这解释了它为什么容易形成Cl⁻。
Cl(g) + e⁻ → Cl⁻(g) ΔH < 0
Electronegativity is a measure of the tendency of an atom to attract a shared pair of electrons in a covalent bond. The most commonly used scale is the Pauling scale. Electronegativity increases across a period and decreases down a group. Fluorine has the highest electronegativity (3.98) and francium the lowest.
电负性用于衡量原子在共价键中吸引成键电子对的能力。最常用的是鲍林标度。电负性同周期从左到右增大，同族从上到下减小。氟的电负性最高（3.98），钫最低。
The difference in electronegativity between two atoms helps predict bond type: a large difference (>1.7) usually indicates ionic character, while a small difference indicates non-polar covalent, and an intermediate difference indicates polar covalent character.
两个原子之间的电负性差有助于预测键型：差值大（>1.7）通常表现为离子性，差值小表现为非极性共价键，中间差值则表现为极性共价键。
6. Metallic and Non-metallic Character | 金属性与非金属性
Metallic character refers to the tendency of an element to lose electrons and form positive ions. It decreases across a period because ionisation energy increases, but increases down a group because ionisation energy decreases. Thus, metals are found on the left and bottom of the periodic table, while non-metals are on the right and top.
金属性是指元素失去电子形成正离子的趋势。同周期中金属性减弱，因为电离能增大；同族中金属性增强，因为电离能减小。因此，金属位于周期表的左侧和下方，非金属位于右侧和上方。
Non-metals tend to gain electrons and form negative ions or share electrons. The oxide of a metal is typically basic, whereas the oxide of a non-metal is typically acidic. Some elements are metalloids (e.g., silicon) and show intermediate behaviour.
非金属倾向获得电子形成负离子或共享电子。金属氧化物通常呈碱性，非金属氧化物通常呈酸性。有些元素是类金属（如硅），表现出中间性质。
There is also a diagonal relationship between certain elements in adjacent periods and groups, such as lithium and magnesium. This occurs because their ionic radii and charge densities are similar.
某些相邻周期和族的元素之间存在对角线关系，例如锂和镁。这是因为它们的离子半径和电荷密度相似。
7. Group 1 and Group 17: Contrasting Trends | 第1族和第17族的周期性对比
Group 1 (alkali metals) have a single valence electron. Their atomic radius increases down the group, ionisation energy decreases, and reactivity increases as it becomes easier to lose the outermost electron. For example, caesium reacts more violently with water than lithium.
第1族（碱金属）只有一个价电子。同族中原子半径增大，电离能减小，反应活性增强，因为失去最外层电子更容易。例如，铯与水的反应比锂剧烈得多。
Group 17 (halogens) have seven valence electrons. Their electronegativity decreases down the group, and their reactivity also decreases because the tendency to gain an electron is weaker. Fluorine is the most reactive halogen. In displacement reactions, a more reactive halogen can displace a less reactive one.
第17族（卤素）有七个价电子。同族中电负性减小，反应活性也减弱，因为获得电子的趋势变弱。氟是最活泼的卤素。在置换反应中，较活泼的卤素可以置换出较不活泼的卤素。
Cl₂(aq) + 2KBr(aq) → 2KCl(aq) + Br₂(aq)
8. Noble Gases: Stability and Exceptions | 稀有气体的稳定性与例外
Noble gases (group 18) have complete valence shells (ns² np⁶, except helium which has 1s²). Their high ionisation energies and very low electron affinities make them extremely unreactive. They exist as monatomic gases under standard conditions.
稀有气体（第18族）具有完全填满的价壳层（ns² np⁶，氦为1s²）。它们的高电离能和极低的电子亲和能使它们极不活泼。在标准状况下以单原子气体存在。
However, noble gases are not completely inert. Xenon reacts with fluorine and oxygen to form compounds such as XeF₂ and XeO₃. Helium and neon do not form stable compounds under normal conditions.
然而，稀有气体并非完全惰性。氙可以与氟和氧反应形成如XeF₂和XeO₃的化合物。氦和氖在普通条件下不能形成稳定化合物。
In IB Chemistry, you should remember that “inert” is a misleading term for noble gases; “unreactive” or “stable” is more accurate.
在IB化学中，要记得“惰性”一词对稀有气体来说并不准确，用“不活泼”或“稳定”更恰当。
9. Transition Elements | 过渡元素
Transition elements are d-block elements that form at least one stable ion with a partially filled d subshell. They are characterised by variable oxidation states, coloured compounds, catalytic activity, and the formation of complex ions.
过渡元素是d区元素中至少能形成一种具有未填满d亚层的稳定离子的元素。它们的特征包括可变氧化态、化合物有色、催化活性以及形成配合离子。
For example, iron exists as Fe²⁺ and Fe³⁺; copper forms Cu⁺ and Cu²⁺. The presence of partially filled d orbitals allows electron transitions within the d subshell, which absorb visible light and produce colour.
例如，铁以Fe²⁺和Fe³⁺存在；铜可形成Cu⁺和Cu²⁺。部分填充的d轨道允许电子在d亚层内跃迁，吸收可见光从而产生颜色。
Transition metals are also good catalysts. For instance, iron is used in the Haber process for ammonia synthesis, and platinum–rhodium is used in the Ostwald process for nitric acid production.
过渡金属也是良好的催化剂。例如，铁用于哈伯法制氨，铂-铑用于奥斯特瓦尔德法制硝酸。
10. Predictive Power and Exceptions | 周期表的预测能力与例外
The periodic table can be used to predict unknown properties. Mendeleev famously predicted the existence and properties of gallium and germanium. For example, by examining elements above and below an unknown element, one can estimate atomic radius, ionisation energy, and formula of the oxide.
周期表可用于预测未知元素的性质。门捷列夫曾成功预测镓和锗的存在及性质。例如，通过观察未知元素上方和下方的元素，可以估计其原子半径、电离能以及氧化物的化学式。
However, there are exceptions and complications. Hydrogen is placed in group 1 but behaves as both a metal and a non-metal. The lanthanides and actinides are placed separately because their f subshells are filled. The d-block shows irregular electron configurations, such as chromium [Ar] 3d⁵ 4s¹ and copper [Ar] 3d¹⁰ 4s¹, which result from extra stability of half-filled and fully filled d subshells.
但周期表也存在例外和复杂情况。氢位于第1族，却表现为金属和非金属双重性质。镧系和锕系元素单独放置，因为它们的f亚层被填充。d区表现出不规则的电子构型，如铬[Ar] 3d⁵ 4s¹和铜[Ar] 3d¹⁰ 4s¹，这是因为半充满和全充满d亚层具有额外稳定性。
11. Exam Tips and Common Pitfalls | 考试要点与常见错误
When answering questions on periodicity, always give the reason based on effective nuclear charge, shielding and distance from the nucleus. Avoid simply stating “across a period radius decreases” without explaining the cause.
回答周期性相关问题时，必须从有效核电荷、屏蔽效应和核间距三个角度给出原因。不要只说“同周期半径减小”而不解释原因。
Remember the exceptions to ionisation energy: Be vs. B and N vs. O. Also note that second ionisation energy is always larger than the first because the electron is removed from a positively charged ion.
要记住电离能的两个例外：铍对硼、氮对氧。还要注意，第二电离能总是大于第一电离能，因为电子是从带正电的离子中移去的。
In questions about electron configuration, write the full configuration or use the noble gas shorthand as instructed. Do not forget that for transition metal ions, electrons are removed from the 4s orbital before the 3d orbital.
在电子构型题目中，应按要求写出完整构型或使用稀有气体简写。不要忘记过渡金属离子中，先失去4s电子，再失去3d电子。
Finally, use the correct terminology: “electronegativity” is a property of atoms in a bond, while “electron affinity” is an isolated atom’s energy change.
最后，使用正确术语：“电负性”是原子在成键时的性质，“电子亲和能”是孤立原子发生电子得失时的能量变化。
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