Ionic or electrovalent bond

What is an ionic bond?

In plain words: pulling an electron off sodium costs energy, and chlorine gives back less than that when it takes the electron, so making the two gaseous ions is uphill. What makes salt form anyway is the huge amount of energy released when vast numbers of Na⁺ and Cl⁻ ions pull together into a crystal. That final step pays for everything else.

Factors favouring an ionic bond

From the Kössel and Lewis treatment, forming an ionic compound depends on (NCERT Section 4.2, page 106):

• how easily the positive and negative ions form from the neutral atoms;
• how the ions arrange themselves in the solid, the lattice of the crystal.

Forming a positive ion means ionisation, removing electrons; forming a negative ion means adding electrons:

The ionisation enthalpy is always positive (endothermic). The electron gain enthalpy, Δ_egH, may be negative or positive; electron affinity uses the opposite sign for energy release, with thermal-convention differences usually neglected here (page 106).

Obviously ionic bonds will be formed more easily between elements with comparatively low ionization enthalpies and elements with comparatively high negative value of electron gain enthalpy.

NCERT Chemistry Part I (Class 11), Section 4.2, page 106

So most ionic compounds have cations from metals (low ionisation enthalpy, on the left of the table) and anions from non-metals (very negative electron gain enthalpy, on the right). The ammonium ion, NH₄⁺, made only of non-metals, is an exception: it is the cation of many ionic compounds, such as NH₄Cl (page 106).

The ionic crystal

In the solid, cations and anions form an orderly three-dimensional arrangement held together by coulombic attraction. The crystal structure depends on the sizes of the ions, how they pack and other factors (NCERT page 106). In sodium chloride (rock salt), each Na⁺ is surrounded by Cl⁻ ions and each Cl⁻ by Na⁺ ions.

The energy balance and lattice enthalpy

For sodium chloride (NCERT page 106):

• ionising Na(g) to Na⁺(g) costs 495.8 kJ mol⁻¹;
• adding an electron to Cl(g) releases 348.7 kJ mol⁻¹ (Δ_egH = −348.7 kJ mol⁻¹);
• so making the gaseous ions costs 495.8 − 348.7 = 147.1 kJ mol⁻¹ overall;
• but forming the crystal lattice releases 788 kJ mol⁻¹, which more than pays for it.

Thus a qualitative measure of the stability of an ionic compound is provided by its enthalpy of lattice formation and not simply by achieving octet of electrons around the ionic species in gaseous state.

NCERT Chemistry Part I (Class 11), Section 4.2, page 107

Lattice enthalpy is defined as the energy needed to separate one mole of a solid ionic compound completely into its gaseous ions (NCERT Section 4.2.1, page 107):

The Lattice Enthalpy of an ionic solid is defined as the energy required to completely separate one mole of a solid ionic compound into gaseous constituent ions.

NCERT Chemistry Part I (Class 11), Section 4.2.1, page 107

For NaCl it is 788 kJ mol⁻¹: 788 kJ is needed to separate one mole of solid NaCl into one mole of Na⁺(g) and one mole of Cl⁻(g) at infinite distance. Lattice enthalpy is positive (energy is absorbed); the enthalpy of lattice formation, the reverse change, is −788 kJ mol⁻¹ (pages 106–107).

The process involves both attraction between oppositely charged ions and repulsion between like ions. Because the crystal is three-dimensional, the lattice enthalpy cannot be calculated from these forces alone; factors that depend on the crystal geometry must be included (page 107). In practice it is found from measured enthalpy changes with a Born–Haber cycle, as in Example 5.

What makes a lattice enthalpy large (beyond the NCERT text, by Coulomb’s law): higher charges on the ions and smaller ions, which can come closer together. So MgO (2+ and 2− ions) has a much larger lattice enthalpy than NaCl (1+ and 1−), and among sodium halides it falls from NaF to NaI as the anion grows. A larger lattice enthalpy usually means a higher melting point: MgO melts near 2850 °C, NaCl at 801 °C.

Solved examples

Common mistakes to avoid

Quick revision summary

• Ionic bond: electron transfer, then electrostatic attraction between ions in a crystal lattice.
• Favoured by a low ionisation enthalpy (metal), a large negative electron gain enthalpy (non-metal) and a large lattice enthalpy.
• NaCl: 495.8 − 348.7 = +147.1 kJ mol⁻¹ to form the gaseous ions, more than repaid by the lattice (788 kJ mol⁻¹).
• Lattice enthalpy: energy to separate one mole of an ionic solid into gaseous ions (NaCl 788 kJ mol⁻¹); larger for higher ionic charges and smaller ions.
• Lattice enthalpy contributes to stability; assess the full energy balance, not just octets or lattice enthalpy alone.

Questions students ask

What is the difference between an ionic bond and a covalent bond?

An ionic bond forms when electrons are transferred from one atom to another and is the non-directional electrostatic attraction between the resulting ions, while a covalent bond forms when two atoms share a pair of electrons and points in a definite direction. Ionic bonds usually join a metal to a non-metal, as in KBr, giving hard, high-melting crystals such as NaCl that conduct well when molten or dissolved because ions can move; covalent bonds usually join non-metals, giving separate molecules such as CH₄ or giant networks such as diamond. Real bonds lie between the two: polar covalent bonds are partly ionic and ionic bonds partly covalent (Fajans’ rules).

Why is an ionic bond non-directional?

An ionic bond is non-directional because an ideal spherical ion has a radial electric field, so each ion attracts every oppositely charged ion around it rather than one particular partner. That is why ionic compounds form three-dimensional crystal lattices instead of separate molecules: in rock salt each Na⁺ has six Cl⁻ nearest neighbours and each Cl⁻ six Na⁺, and the formula NaCl gives only the ratio of the ions. A covalent bond is directional because it depends on the overlap of orbitals that point in particular directions.

What is the difference between lattice enthalpy and lattice energy?

Lattice enthalpy is the enthalpy change, at constant pressure, for separating one mole of an ionic solid into its gaseous ions, while lattice energy, strictly, is the internal energy change for the same process; for ideal gaseous ions and negligible solid volume, the two differ by Δn_gRT. For NaCl(s) → Na⁺(g) + Cl⁻(g), Δn_g = 2, so at 298 K the lattice energy is smaller by 2RT ≈ 5 kJ mol^⁻¹: about 783 kJ mol^⁻¹ against 788 kJ mol^⁻¹. The gap is under 1%, so many books use the two names interchangeably.

Is an electrovalent bond the same as an ionic bond?

Yes, an electrovalent bond is another name for an ionic bond, the electrostatic attraction between ions formed by the transfer of electrons. The name comes from electrovalence, the number of electrons an atom loses or gains, which equals the charge on its ion: in Al₂O₃ aluminium has an electrovalence of +3 and oxygen of −2. NCERT uses both names, while most modern books simply say ionic bond.

In which chapter of Class 11 is the Born–Haber cycle?

The Born–Haber cycle is in NCERT Class 11 Chemistry Part I, Unit 5 (Thermodynamics), Section 5.5, where it is used to find the lattice enthalpy of NaCl; Unit 4, Chemical Bonding and Molecular Structure, only defines lattice enthalpy (Section 4.2.1). Neither the JEE Main 2027 nor the NEET (UG) 2026 syllabus names the cycle, but both list the ‘calculation of lattice enthalpy’, which uses it, under Unit 3, Chemical Bonding and Molecular Structure. The cycle for NaCl is worked as Example 5 on this page.

Practice questions

Try each question first, then open the answer to check your working.

1. Q1. Define lattice enthalpy, and explain what “the lattice enthalpy of NaCl is 788 kJ mol⁻¹” means.

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   Lattice enthalpy is the energy needed to separate one mole of a solid ionic compound completely into its gaseous ions. For NaCl, 788 kJ must be supplied to turn one mole of solid NaCl into one mole of Na⁺(g) and one mole of Cl⁻(g) far apart.

2. Q2. The sum of the ionisation enthalpy of Na and the electron gain enthalpy of Cl is positive. Why does NaCl form at all?

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   Because forming the crystal lattice from the gaseous ions releases 788 kJ mol⁻¹, much more than the 147.1 kJ mol⁻¹ needed to form the ions. The overall change is strongly negative.

3. Q3. Arrange KCl, NaCl and MgO in order of increasing lattice enthalpy.

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   KCl < NaCl < MgO. K⁺ is larger than Na⁺, so KCl’s ions are farther apart; MgO has doubly charged ions.