📚 Ionic Bonding and Crystal Structures | 离子键的形成与晶体结构
Ionic bonding is one of the fundamental types of chemical bonding studied in IB Chemistry. It occurs when a metal transfers electrons to a non-metal, forming ions that are held together by electrostatic attraction. Understanding how ionic bonds form and how they produce ordered crystal structures is essential for explaining the physical properties of many inorganic compounds.
离子键是 IB 化学中学习的基本化学键类型之一。当金属将电子转移给非金属时形成离子，离子之间通过静电吸引而结合。理解离子键如何形成以及如何产生有序晶体结构，是解释许多无机化合物物理性质的关键。
1. The Nature of Ionic Bonding | 离子键的本质
Ionic bonding is the electrostatic attraction between oppositely charged ions. A metal atom typically loses one or more valence electrons to become a positively charged cation, while a non-metal atom gains those electrons to become a negatively charged anion. The bond is not localised between two specific atoms; instead, it acts in all directions, producing a three-dimensional lattice.
离子键是带相反电荷离子之间的静电吸引。金属原子通常失去一个或多个价电子成为带正电荷的阳离子，而非金属原子获得这些电子成为带负电荷的阴离子。离子键并不局限于两个特定原子之间；而是在所有方向共同作用，形成三维晶格。
For example, sodium reacts with chlorine to form sodium chloride. The sodium atom has the electron configuration 1s²2s²2p⁶3s¹, and the chlorine atom has 1s²2s²2p⁶3s²3p⁵. During the reaction, sodium loses its 3s electron and becomes Na⁺ with the configuration 1s²2s²2p⁶, which is isoelectronic with neon. Chlorine gains that electron and becomes Cl⁻ with 1s²2s²2p⁶3s²3p⁶, which is isoelectronic with argon.
例如，钠与氯反应生成氯化钠。钠原子的电子构型为 1s²2s²2p⁶3s¹，氯原子为 1s²2s²2p⁶3s²3p⁵。反应中，钠失去 3s 电子变成 Na⁺，构型为 1s²2s²2p⁶，与氖等电子；氯获得该电子成为 Cl⁻，构型为 1s²2s²2p⁶3s²3p⁶，与氩等电子。
Na → Na⁺ + e⁻ Cl + e⁻ → Cl⁻
The electron transfer produces a cation and an anion, and the subsequent electrostatic attraction between Na⁺ and Cl⁻ is the ionic bond. The overall process is highly exothermic because the lattice energy released is greater than the energy required to remove electrons and atomise the elements.
电子转移产生阳离子和阴离子，随后 Na⁺ 与 Cl⁻ 之间的静电吸引即为离子键。整个过程中释放的晶格能大于电离和原子化所需的能量，因此总体上高度放热。
2. Electron Configurations and Dot-and-Cross Diagrams | 电子构型与电子点叉图
In IB Chemistry, dot-and-cross diagrams are used to show the transfer of electrons during ionic bond formation. Only the valence electrons are shown, and the electrons from the metal are drawn as crosses while those from the non-metal are drawn as dots, or vice versa. The square brackets around each ion indicate a full charge, not a partial charge.
在 IB 化学中，电子点叉图用于表示离子键形成过程中的电子转移。图中只显示价电子，金属给出的电子用叉表示，非金属的电子用点表示，反之亦可。每个离子周围的方括号表示完整电荷，而非部分电荷。
For sodium chloride, the dot-and-cross diagram shows the sodium atom losing its single 3s electron and the chlorine atom gaining it, resulting in Na⁺ and Cl⁻. Both ions have complete outer shells with eight electrons, satisfying the octet rule.
对于氯化钠，电子点叉图显示钠原子失去其单个 3s 电子，氯原子获得该电子，形成 Na⁺ 和 Cl⁻。两种离子都具有完整的八电子外层，满足八隅体规则。
For magnesium oxide, magnesium loses two electrons and oxygen gains two electrons. The ionic bond in MgO is stronger than in NaCl because the ions carry higher charges, attracting each other more powerfully. The dot-and-cross diagram shows Mg²⁺ with an empty outer shell and O²⁻ with a full octet.
对于氧化镁，镁失去两个电子，氧获得两个电子。MgO 中的离子键比 NaCl 更强，因为离子所带电荷更高，相互吸引更强。电子点叉图显示 Mg²⁺ 外层已空，O²⁻ 具有完整的八电子。
It is important to remember that ionic compounds do not exist as individual ion pairs. In reality, each cation is surrounded by several anions and each anion by several cations. The formula of an ionic compound, such as NaCl or MgO, represents the simplest whole-number ratio of ions in the lattice.
需注意，离子化合物并不以单个离子对存在。实际上，每个阳离子周围有多个阴离子，每个阴离子周围有多个阳离子。离子化合物的化学式（如 NaCl 或 MgO）表示晶格中离子的最简整数比。
3. Lattice Enthalpy and Born–Haber Cycles | 晶格焓与玻恩–哈伯循环
Lattice enthalpy, ΔHlattice, is defined as the enthalpy change when one mole of an ionic compound is formed from its gaseous ions under standard conditions. For sodium chloride, this corresponds to the reaction:
晶格焓 ΔHlattice 定义为标准条件下由气态离子形成一摩尔离子化合物时的焓变。对于氯化钠，对应反应为：
Na⁺(g) + Cl⁻(g) → NaCl(s) ΔH = −787 kJ mol⁻¹
The lattice enthalpy is a measure of the strength of the ionic bonding in the crystal. A more negative ΔHlattice means a stronger attraction between ions and therefore a more stable lattice. Lattice enthalpy cannot be measured directly, so Born–Haber cycles are used to calculate it indirectly.
晶格焓衡量晶体中离子键的强度。ΔHlattice 越负，说明离子间吸引力越强，晶格越稳定。晶格焓无法直接测量，因此使用玻恩–哈伯循环间接计算。
A Born–Haber cycle uses Hess’s law to relate lattice enthalpy to other measurable enthalpy changes. For sodium chloride, the key steps are:
玻恩–哈伯循环利用赫斯定律，将晶格焓与其他可测量的焓变联系起来。对于氯化钠，关键步骤如下：
-
Standard formation of NaCl(s) from elements: ΔHf = −411 kJ mol⁻¹.
由单质生成 NaCl(s) 的标准生成焓：ΔHf = −411 kJ mol⁻¹。
-
Atomisation of sodium: Na(s) → Na(g), ΔH = +107 kJ mol⁻¹.
钠的原子化：Na(s) → Na(g)，ΔH = +107 kJ mol⁻¹。
-
Atomisation of chlorine: ½Cl₂(g) → Cl(g), ΔH = +122 kJ mol⁻¹.
氯的原子化：½Cl₂(g) → Cl(g)，ΔH = +122 kJ mol⁻¹。
-
First ionisation energy of sodium: Na(g) → Na⁺(g) + e⁻, ΔH = +496 kJ mol⁻¹.
钠的第一电离能：Na(g) → Na⁺(g) + e⁻，ΔH = +496 kJ mol⁻¹。
-
Electron affinity of chlorine: Cl(g) + e⁻ → Cl⁻(g), ΔH = −349 kJ mol⁻¹.
氯的电子亲和能：Cl(g) + e⁻ → Cl⁻(g)，ΔH = −349 kJ mol⁻¹。
Using Hess’s law, the lattice enthalpy can be calculated as:
根据赫斯定律，晶格焓可计算为：
ΔHlattice = ΔHf − ΔHatom(Na) − ΔHatom(Cl) − IE₁(Na) − EA(Cl)
ΔHlattice = (−411) − (+107) − (+122) − (+496) − (−349) = −787 kJ mol⁻¹
In the Born–Haber cycle, it is essential to use the correct signs for each step. Electron affinity values are usually negative for the first electron gain by non-metals, while ionisation energies and atomisation enthalpies are positive.
在玻恩–哈伯循环中，务必注意每一步焓变的符号。非金属获得第一个电子的电子亲和能通常为负，而电离能和原子化焓为正。
4. Factors Affecting Lattice Enthalpy | 影响晶格焓的因素
Two main factors determine the magnitude of lattice enthalpy: ionic charge and ionic radius. Both affect the strength of electrostatic attraction between ions in the lattice.
决定晶格焓大小的两个主要因素是离子电荷和离子半径。两者都会影响晶格中离子间静电吸引的强弱。
Ionic charge: The electrostatic force between ions is proportional to the product of their charges, Q₁Q₂. Doubling both charges from +1 to +2 and from −1 to −2 increases the force by a factor of four. Therefore, MgO has a much more exothermic lattice enthalpy than NaCl.
离子电荷：离子间的静电力与电荷乘积 Q₁Q₂ 成正比。将电荷从 +1 增至 +2、从 −1 增至 −2，会使作用力增大到原来的四倍。因此，MgO 的晶格焓远比 NaCl 更放热。
Ionic radius: The electrostatic force also depends on the distance between the ion centres, which is approximately the sum of the ionic radii. Smaller ions can pack more closely together, leading to stronger attraction. For example, LiF has a more negative lattice enthalpy than NaI, even though both contain +1 and −1 ions, because Li⁺ and F⁻ are much smaller than Na⁺ and I⁻.
离子半径：静电力还取决于离子中心之间的距离，约等于离子半径之和。较小的离子能排列得更紧密，产生更强的吸引。例如，LiF 的晶格焓比 NaI 更负，即使两者都含 +1 和 −1 离子，因为 Li⁺ 和 F⁻ 比 Na⁺ 和 I⁻ 小得多。
| Compound | Ionic charge | Sum of ionic radii / pm | ΔHlattice / kJ mol⁻¹ |
| NaCl | +1, −1 | 102 + 181 = 283 | −787 |
| LiF | +1, −1 | 76 + 133 = 209 | −1049 |
更多咨询请联系16621398022（同微信）
CommentsMore posts |
屏轩国际教育cambridge primary/secondary checkpoint, cat4, ukiset,ukcat,igcse,alevel,PAT,STEP,MAT, ibdp,ap,ssat,sat,sat2课程辅导，国外大学本科硕士研究生博士课程论文辅导