Ionic or Electrovalent Bond: Formation, Lattice Enthalpy & Properties
Kössel’s road: a metal hands its valence electrons to a non-metal, two octets complete, and electrostatics takes over. The strength of what forms — the lattice enthalpy — is a simple product of charge and closeness, and it decides everything from melting point to solubility.
The Transfer Bond — Complete Theory
An ionic (electrovalent) bond is the electrostatic attraction between oppositely charged ions, formed by the complete transfer of one or more electrons from one atom to another. The donor is typically a metal with low ionization enthalpy — it parts with electrons cheaply; the acceptor is a non-metal with high negative electron gain enthalpy — it profits from taking them. NaCl is the archetype: Na (3s¹) hands its electron to Cl (3p⁵); both reach noble-gas octets, and Na⁺Cl⁻ crystallises.
Why the transfer pays. The ledger of § 4.1 decides the direction: few dots on one side (cheap to lose), nearly-full shell on the other (profitable to fill). The energy accounting runs: strip Na (+496 kJ/mol) and fill Cl (−349 kJ/mol) — a net cost so far. The transfer only becomes profitable when the lattice enthalpy — the energy released packing Na⁺ and Cl⁻ into the crystal — pays it back with interest (−788 kJ/mol). The lattice is not decoration; it is the business model of ionic bonding.
low IE (metal)
Acceptorhigh −EA (non-metal)
Payofflarge lattice enthalpy
Lattice enthalpy — the strength dial. Defined as the energy required to completely separate one mole of solid ionic compound into gaseous ions (NaCl: 788 kJ/mol). Two factors, and both are Coulomb’s law: it rises with ionic charge — MgO (2+/2−) is roughly four times NaCl’s league — and falls with ionic size — CsCl is weaker than NaCl because caesium is fatter. Every melting point in the ionic world is this dial being read aloud.
| Compound | Ions | Lattice enthalpy (kJ/mol) | Charge · size story |
|---|---|---|---|
| MgO | Mg²⁺ / O²⁻ | ≈ 3795 | Double charge both sides — Coulomb ×4 |
| CaO | Ca²⁺ / O²⁻ | ≈ 3460 | Double charge, slightly larger ions than MgO |
| NaF | Na⁺ / F⁻ | ≈ 894 | Single charge, small ions — strong for 1+/1− |
| NaCl | Na⁺ / Cl⁻ | 788 | The reference point |
| NaBr | Na⁺ / Br⁻ | ≈ 732 | Bigger anion → weaker glue |
| NaI | Na⁺ / I⁻ | ≈ 682 | Biggest anion of the set → weakest |
| CsCl | Cs⁺ / Cl⁻ | ≈ 657 | Big cation → distance kills the attraction |
Properties of ionic compounds — every one traces to the lattice. They are hard, brittle crystalline solids (the lattice is strong, but sliding a layer by one ion puts like charges face-to-face and it shatters). They have high melting and boiling points (breaking the lattice costs the numbers in Table 1 — MgO melts near 2852 °C). They conduct in the molten state and in aqueous solution but not as solids — solid ions are locked in place; melted or dissolved, they migrate and carry current. And they are soluble in polar solvents like water — water’s dipoles claw at the lattice ions — but insoluble in non-polar solvents like benzene, which offer nothing to the ions.
| Property | Observation | Lattice explanation |
|---|---|---|
| Nature | Hard, brittle crystals | Strong lattice; layer-slip brings like charges together → shatter |
| Melting/boiling points | High (NaCl 801 °C · MgO 2852 °C) | Lattice enthalpy must be paid in full |
| Conduction (solid) | No | Ions locked at lattice sites — nothing mobile |
| Conduction (molten/aqueous) | Yes | Ions freed to migrate — they are the charge carriers |
| Solubility | Water: yes · Benzene: no | Polar solvent claws at the lattice; non-polar cannot |
| Vapour phase species | Ion pairs, not molecules | No discrete “NaCl molecule” exists in the crystal |
Visualising the Transfer & Ranking the Lattices
Ek transfer, ek comparator — NaCl’s electron handover first, then the lattice league table ranked live.
Lattice Enthalpy Comparator
Charge badha do ya size ghatao — bars ko apne aap rank hote dekho. Har compound ke neeche uska Coulomb story likha hai.
Read the pattern: MgO and CaO tower over the single-charge set (charge ×4), and within any charge family the bars sink as the ions grow (NaF → NaI, and CsCl last). Two dials — charge and distance — and the whole ionic world obeys.
Bars are scaled to MgO’s ≈ 3795 kJ/mol. Values follow NCERT/standard convention; treat exact figures as order-of-magnitude anchors.
Solved Examples (Step-by-Step)
Transfer → lattice → property. Jo chain yahan chalti hai, wahi comparator me live chalti hai.
MgO beats MgCl in the formation stakes
Explain, using electron transfer, why magnesium oxide is MgO and why its bond is far stronger than that of NaCl.
- TransferMg (2,8,2) loses 2 e⁻ → Mg²⁺ (Ne config); O (2,6) gains 2 → O²⁻ (Ne config). 1 : 1 balanced.
- Charge effectBoth ions carry double the NaCl charges → electrostatic attraction ≈ ×4 → lattice ≈ 3795 kJ/mol (vs NaCl 788).
- Bonus factorBoth ions are also smaller than Na⁺/Cl⁻ — distance helps too.
MgO: double charge + smaller ions — nearly 5× NaCl’s glue
Why CsCl is the weakling of its family
NaCl and CsCl carry identical 1+/1− charges and the same crystal structure. Yet CsCl’s lattice enthalpy is about 657 kJ/mol against NaCl’s 788. Explain.
- Same chargeBoth 1+/1− — the charge dial is identical.
- Distance dialCs⁺ is far larger than Na⁺ → the ion centres sit farther apart → Coulomb attraction falls with distance.
- Verdict~130 kJ/mol weaker — enough to show in melting points (NaCl 801 °C vs CsCl 645 °C). Within any charge family, size runs the ranking.
Same charge, bigger cation — weaker lattice
Solid NaCl is an insulator; molten NaCl is not
Solid sodium chloride does not conduct electricity, but molten sodium chloride does. Explain both observations with the same model.
- SolidIons are locked at lattice sites — charge exists but cannot migrate → no conduction.
- MoltenMelting collapses the fixed lattice; Na⁺ and Cl⁻ migrate freely under applied voltage → conduction.
- Model noteThe carriers are ions, never electrons — that is why solid ionic compounds differ from metals here (metals conduct via electrons even as solids).
Locked ions vs mobile ions — one model, both observations
Practice Questions (With Solutions)
Attempt first — options lock after one shot, exactly like the real exam. Then read the working, chahe galti ho ya na ho.
Attempted 0/4 · Correct 0
The formation of an ionic bond is favoured by:
Solution
- The three-factor answer: cheap electron out (low IE), profitable electron in (high −EA), and a lattice worth building (large lattice enthalpy).
- Option C inverts the size dial; A and D poison the energy ledger.
(B) The three-factor answer
The correct order of increasing lattice enthalpy is:
Solution
- CsCl (big ions, 1+/1−) < NaCl (smaller ions) < MgO (2+/2−, small) — ≈ 657 < 788 < 3795 kJ/mol.
- Charge dominates, then size — the comparator’s bars show the same ranking visually.
(D) CsCl < NaCl < MgO
The ionic bond is non-directional because:
Solution
- A point charge pulls equally in all directions — the field is a sphere, so no angle is preferred → non-directional.
- Option D belongs to covalent bonding, which IS directional (hybrids point somewhere, § 4.6).
(B) Spherical field
Ionic compounds are soluble in water but not in benzene because:
Solution
- Water’s dipoles surround and stabilise the ions — clawing them out of the lattice. Non-polar benzene offers ions nothing.
- The same polarity logic that gave CO₂ its zero dipole (§ 4.3) runs the solubility story here.
(B) Polar solvent claws the lattice
Key Rules & Takeaways
Eight lines that solve this topic
Lattice league: MgO 3795 · CaO 3460 · NaF 894 · NaCl 788 · NaBr 732 · NaI 682 · CsCl 657 kJ/mol · Melting anchors: NaCl 801 °C · MgO 2852 °C
- The lattice funds the transfer — ionisation costs money; only a fat lattice enthalpy repays it. No lattice, no ionic bond.
- Two dials run the ranking — charge first (×4 for 2+/2−), size second; the comparator’s bars are Coulomb’s law drawn as furniture.
- Non-directional explains the crystal — spherical fields pack ions like balls; no molecules, no angles, 6:6 coordination.
- Every property is the lattice speaking — hardness, melting, conduction qualifier, solubility: trace each answer back to the lattice enthalpy number.
FAQs
How is an ionic bond formed?
An ionic bond forms by the complete transfer of one or more electrons from an atom of low ionization enthalpy — typically a metal — to an atom of high negative electron gain enthalpy — typically a non-metal. The transfer produces oppositely charged ions that are held together by electrostatic attraction, as in NaCl: Na loses its single valence electron to chlorine, forming Na+ and Cl−.
What factors favour the formation of an ionic bond?
Three factors: low ionization enthalpy of the electropositive atom — easy to strip its electron; high negative electron gain enthalpy of the electronegative atom — eager to accept it; and large lattice enthalpy — the crystal's electrostatic glue, which rises with the charges of the ions and falls with their sizes. MgO out-bonds NaCl because Mg2+ and O2− carry double the charge.
What is lattice enthalpy?
Lattice enthalpy is the energy required to completely separate one mole of a solid ionic compound into gaseous ions — for NaCl it is 788 kJ/mol. It measures ionic bond strength: it increases with the magnitude of the ionic charges and decreases with increasing ionic size. Hence MgO (double charges, smaller ions) has a far larger lattice enthalpy than NaCl, and CsCl is weaker than NaCl because caesium is larger.
Why do ionic compounds conduct electricity only in the molten or dissolved state?
In the solid, the ions are locked in fixed lattice positions and cannot move, so no charge flows. Melting or dissolving in a polar solvent frees the ions to migrate — the molten compound or aqueous solution then conducts electricity. The charge carriers are the ions themselves, not free electrons as in metals.
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