Ionic Bonding — Concepts, Formulas & Examples

Ionic bonding is the complete transfer of electrons from one atom to another, forming cations and anions held together by electrostatic attraction. CBSE Class 11 and NEET test formation, lattice energy and properties.

Core Concepts

Formation of ionic bonds

Between metal (low ionisation energy) and non-metal (high electron affinity). Metal loses electrons to form cation; non-metal gains to form anion. Both reach noble gas configuration.

Conditions favouring ionic bond formation:

Low ionisation energy of the metal (easy to lose electrons) — alkali metals are ideal
High electron affinity of the non-metal (eager to gain electrons) — halogens are ideal
Large electronegativity difference (typically > 1.7 on the Pauling scale)

Example — Formation of NaCl:

\text{Na} \rightarrow \text{Na}^+ + e^- \quad (\text{IE}_1 = 496\;\text{kJ/mol})

\text{Cl} + e^- \rightarrow \text{Cl}^- \quad (\text{EA} = -349\;\text{kJ/mol})

Na loses one electron (2,8,1 → 2,8) and Cl gains it (2,8,7 → 2,8,8). Both achieve noble gas configuration. The electrostatic attraction between Na $^+$ and Cl $^-$ forms the ionic bond.

Lattice energy

Energy released when gaseous ions combine to form one mole of solid crystal. Higher for smaller ions and higher charges. $\text{LE} \propto z_+ z_- / r$ .

U = k \frac{z_+ z_- e^2}{r_+ + r_-}

Where $z_+$ and $z_-$ are ion charges, $e$ is electronic charge, and $r_+$ , $r_-$ are ionic radii. The constant $k$ includes the Madelung constant (accounts for the geometry of the crystal lattice).

Factors affecting lattice energy:

Factor	Effect	Example
Smaller ions	Higher LE	LiF > NaF > KF
Higher charge	Much higher LE	MgO >> NaF (both similar size, but 2+ and 2- vs 1+ and 1-)
Crystal structure	Affects Madelung constant	NaCl vs CsCl structures

Lattice energy values to remember:

NaCl: 786 kJ/mol
MgO: 3850 kJ/mol (about 5x NaCl — because charges are doubled)
LiF: 1037 kJ/mol (small ions, high LE)

The huge lattice energy of MgO explains why it melts at 2852°C while NaCl melts at just 801°C.

Born-Haber cycle

A thermochemical cycle used to calculate lattice energy indirectly, using enthalpies of sublimation, ionisation, electron affinity and bond dissociation.

The cycle for NaCl:

\text{Na(s)} \rightarrow \text{Na(g)} \quad \Delta H_{sub} = +108\;\text{kJ/mol}

\text{Na(g)} \rightarrow \text{Na}^+\text{(g)} + e^- \quad \text{IE}_1 = +496\;\text{kJ/mol}

\frac{1}{2}\text{Cl}_2\text{(g)} \rightarrow \text{Cl(g)} \quad \frac{1}{2}\Delta H_{diss} = +121\;\text{kJ/mol}

\text{Cl(g)} + e^- \rightarrow \text{Cl}^-\text{(g)} \quad \text{EA} = -349\;\text{kJ/mol}

\text{Na}^+\text{(g)} + \text{Cl}^-\text{(g)} \rightarrow \text{NaCl(s)} \quad \Delta H_{lattice} = ?

By Hess’s law: $\Delta H_f = \Delta H_{sub} + \text{IE} + \frac{1}{2}\Delta H_{diss} + \text{EA} + \Delta H_{lattice}$

Given $\Delta H_f(\text{NaCl}) = -411$ kJ/mol:

-411 = 108 + 496 + 121 + (-349) + \Delta H_{lattice}

\Delta H_{lattice} = -411 - 376 = -787\;\text{kJ/mol}

Properties of ionic compounds

High melting and boiling points. Conduct electricity only when molten or dissolved (ions must be mobile). Soluble in polar solvents (water). Hard but brittle.

Why ionic compounds have high melting points: Breaking an ionic lattice requires overcoming the strong electrostatic attraction between millions of ions arranged in a regular pattern. NaCl requires 786 kJ/mol of lattice energy to separate — that translates to needing 801°C to melt.

Why they conduct only when molten or dissolved: In the solid state, ions are locked in fixed positions — they vibrate but cannot move through the lattice. When melted or dissolved in water, ions become mobile and can carry charge. This is why electrolysis works only with molten or aqueous ionic compounds.

Why they are hard but brittle: The strong electrostatic forces resist deformation (hardness). But when a force shifts one layer of ions, like charges suddenly face each other — the resulting repulsion shatters the crystal.

Solubility rule: An ionic compound dissolves in water if the hydration energy (energy released when water molecules surround ions) exceeds the lattice energy. Small, highly charged ions have high hydration energy. This is why NaCl dissolves easily but BaSO $_4$ does not — BaSO $_4$ has very high lattice energy that water cannot overcome.

Fajan’s rules

Predict covalent character in ionic bonds. More covalent character when cation is small and highly charged, anion is large and highly charged. Applies to compounds like AgCl (mostly ionic) vs AgI (more covalent).

Summary of Fajan’s rules:

Small cation with high charge → strong polarisation → more covalent
Large anion with high charge → easily polarised → more covalent
Cation with pseudo-noble gas configuration (d $^{10}$ ) like Cu $^+$ , Ag $^+$ , Zn $^{2+}$ → extra polarising power → more covalent

Application: Among silver halides, AgF is most ionic and AgI is most covalent. Why? F $^-$ is small (hard to polarise), while I $^-$ is large (easy to polarise). Ag $^+$ has a d $^{10}$ configuration, making it a strong polariser.

This explains colour too — AgCl is white, AgBr is pale yellow, AgI is yellow. More covalent character = more colour (because covalent compounds absorb visible light more easily).

Coordination number

In an ionic crystal, each ion is surrounded by a fixed number of oppositely charged ions — this is the coordination number.

Structure	Cation CN	Anion CN	Example
Rock salt (NaCl)	6	6	NaCl, KCl, MgO
Caesium chloride (CsCl)	8	8	CsCl, CsBr
Zinc blende (ZnS)	4	4	ZnS, CuCl
Fluorite (CaF $_2$ )	8	4	CaF $_2$ , BaF $_2$

The radius ratio ( $r_+/r_-$ ) determines which structure an ionic compound adopts. This is a favourite JEE question.

Worked Examples

Na and Cl have a large electronegativity difference; electron transfer is complete. H and Cl have a smaller difference; electrons are shared but unequally, giving a polar covalent bond.

Slipping layers of ions brings like charges close together — strong repulsion shatters the crystal. Metals are not brittle because their delocalised electrons rearrange without this effect.

Both have similar internuclear distances. But NaF has charges +1 and -1, while MgO has +2 and -2.

\frac{U_{MgO}}{U_{NaF}} \approx \frac{2 \times 2}{1 \times 1} = 4

So MgO should have roughly 4 times the lattice energy of NaF. Experimental values: NaF = 923 kJ/mol, MgO = 3850 kJ/mol — the ratio is about 4.2. The charge effect dominates lattice energy.

Solved Problems (Exam Style)

Problem 1 (JEE Main pattern): Using the Born-Haber cycle, calculate the electron affinity of chlorine given: $\Delta H_f(\text{NaCl}) = -411$ kJ/mol, $\Delta H_{sub}(\text{Na}) = 108$ kJ/mol, IE $_1$ (Na) = 496 kJ/mol, $\frac{1}{2}$ D(Cl $_2$ ) = 121 kJ/mol, $U(\text{NaCl}) = -787$ kJ/mol.

\Delta H_f = \Delta H_{sub} + \text{IE} + \frac{1}{2}D + \text{EA} + U

-411 = 108 + 496 + 121 + \text{EA} + (-787)

-411 = -62 + \text{EA}

\text{EA} = -411 + 62 = -349\;\text{kJ/mol}

Problem 2 (NEET pattern): Arrange in order of increasing lattice energy: NaCl, KCl, MgO, CaO.

For same charge: smaller ions → higher LE. KCl < NaCl (K $^+$ is larger than Na $^+$ ).

For different charges: higher charge → much higher LE. MgO and CaO (2+ and 2-) >> NaCl, KCl (1+ and 1-).

Between MgO and CaO: Mg $^{2+}$ is smaller than Ca $^{2+}$ , so MgO > CaO.

Order: KCl < NaCl < CaO < MgO

Common Mistakes

Saying ionic bonds are non-directional. That is correct for the bond itself, but the crystal lattice has a specific geometry.

Confusing lattice energy and ionisation energy.

Writing that ionic compounds conduct in solid state. They do not — ions are fixed.

Thinking that all ionic compounds dissolve in water. BaSO $_4$ , AgCl, PbSO $_4$ are ionic but insoluble — their lattice energy exceeds the hydration energy that water can provide.

Forgetting the factor-of-four effect of charge. Doubling both ion charges quadruples the lattice energy. Students often think it just doubles.

Exam Weightage and Revision

This topic is a repeat performer in board papers and entrance exams. NEET typically asks one to two questions on the core mechanisms, CBSE boards give three to six marks, and state PMT papers often include a diagram-based long answer. The PYQs cluster around a small set of facts — lock those and you clear the topic.

JEE Main 2023 had a Born-Haber cycle calculation. JEE Main 2024 tested lattice energy ordering. NEET 2022 asked about properties of ionic compounds (conductivity). CBSE boards love the five-mark question on ionic bond formation with electron dot structures.

When a question gives a scenario, identify the core mechanism first, then match it to the concepts above. Most wrong answers come from reading the scenario too quickly.

Memorise four properties of ionic compounds — high MP, conduct when molten, soluble in water, brittle. Covers most one-line questions.

Practice Questions

Q1. Why does MgO have a much higher melting point than NaCl?

MgO has charges of +2 and -2, while NaCl has +1 and -1. Lattice energy is proportional to $z_+ \times z_-$ . So MgO has roughly four times the lattice energy of NaCl (3850 vs 786 kJ/mol). Higher lattice energy means more energy needed to break the lattice, hence a higher melting point (2852°C vs 801°C).

Q2. Why is AgI coloured while NaCl is white?

Ag $^+$ has a d $^{10}$ (pseudo-noble gas) configuration, which gives it extra polarising power. I $^-$ is large and easily polarised. By Fajan’s rules, AgI has significant covalent character. Covalent character allows electronic transitions that absorb visible light, giving colour (yellow). NaCl is purely ionic with no such transitions — it is white.

Q3. What is the coordination number of Na $^+$ in NaCl crystal?

In the NaCl (rock salt) structure, each Na $^+$ is surrounded by 6 Cl $^-$ ions arranged at the corners of an octahedron. So the coordination number is 6. Similarly, each Cl $^-$ is surrounded by 6 Na $^+$ ions.

Q4. An ionic compound has very high melting point but is insoluble in water. Explain.

High melting point confirms strong ionic bonding (high lattice energy, likely due to high charges or small ions). Insolubility means the lattice energy is greater than the hydration energy. Water molecules cannot provide enough energy to pull the ions apart. Example: MgO has a very high melting point but is practically insoluble because the $2+ / 2-$ lattice is too strong for water to break.

Q5. Why do ionic compounds not conduct electricity in solid state but do so when molten?

Electrical conduction requires mobile charge carriers. In solid ionic compounds, ions are held in fixed positions in the crystal lattice — they can vibrate but not migrate. When melted, the lattice breaks down and ions become free to move, carrying charge. This is also why dissolved ionic compounds conduct — water separates the ions.

FAQs

Is there such a thing as a purely ionic bond? No. Even NaCl has about 51% ionic character (calculated from the Hannay-Smyth equation). All ionic bonds have some covalent character, and Fajan’s rules predict how much. The classification as ionic or covalent is a matter of degree, not a strict boundary.

Why are ionic compounds soluble in water but not in organic solvents? Water is highly polar (dielectric constant ~80) and can stabilise separated ions through strong ion-dipole interactions (hydration). Organic solvents like hexane have low dielectric constants and cannot stabilise separated ions. The energy gained from solvation is too small to overcome the lattice energy.

What determines whether an ionic compound takes the NaCl or CsCl structure? The radius ratio $r_+/r_-$ . When the ratio is between 0.414 and 0.732, the NaCl structure (coordination number 6) is favoured. When the ratio is above 0.732, the CsCl structure (coordination number 8) is preferred. Cs $^+$ is large enough relative to Cl $^-$ to accommodate 8 neighbours.

How does lattice energy relate to solubility? Higher lattice energy generally means lower solubility, because water must overcome a larger energy barrier to separate the ions. However, this is only part of the story — hydration energy also matters. If hydration energy exceeds lattice energy, the compound dissolves. That is why NaOH (LE ~900 kJ/mol but high hydration energy) dissolves easily.

Ionic bonding is the strongest simple bond type and gives materials some of their most useful properties. It is also the simplest to understand — just electron transfer.