What Is an Ionic Bond
The attractive force that holds the constituent atoms or ions together in any chemical species is called a chemical bond; bonding is nature's way of lowering the energy of a system to attain stability. The ionic bond is the simplest realisation of this idea. Following the Kössel–Lewis treatment, a highly electropositive metal atom transfers one or more electrons to a highly electronegative non-metal atom, so that each attains the stable octet (or duplet) configuration of the nearest noble gas.
The resulting cation (positive) and anion (negative) are then held together by electrostatic attraction. NCERT defines this bond, formed as a result of the electrostatic attraction between the positive and negative ions, as the electrovalent bond. NIOS states it equivalently: the ionic bond is the electrostatic force of attraction that holds the cation and anion together, and the compounds so formed are termed ionic or electrovalent compounds.
The magnitude of charge carried by an ion is its electrovalence, equal to the number of unit charges on that ion. Calcium, which forms $\ce{Ca^2+}$, is assigned a positive electrovalence of two; chlorine, which forms $\ce{Cl-}$, a negative electrovalence of one.
"Ionic bond" is transfer, not sharing
An ionic (electrovalent) bond arises from the complete transfer of electrons and an electrostatic attraction between discrete ions. A covalent bond arises from the sharing of an electron pair. Examiners frequently pair these as a single match-the-following or assertion item.
Transfer → ions → coulombic attraction = ionic. Sharing → common pair = covalent.
Electron Transfer: How the Bond Forms
Consider sodium chloride. Sodium ($\ce{[Ne]3s^1}$) is one electron beyond a noble-gas core, so it loses that electron readily to become $\ce{Na+}$; chlorine ($\ce{[Ne]3s^2 3p^5}$) is one electron short of argon, so it accepts the electron to become $\ce{Cl-}$. Both ions now possess a complete octet, and the bond is the coulombic attraction between them.
The stepwise electron transfer can be written in mhchem as:
$$\ce{Na ->[\,-e^-\,] Na+}\qquad \ce{Cl + e^- -> Cl^-}\qquad \ce{Na+ + Cl^- -> NaCl}$$
and the overall ion-pair formation as $\ce{Na+ + Cl- -> NaCl}$.
The same logic, applied to a divalent metal, gives compounds in which the formula reflects charge balance rather than a discrete molecule. For calcium fluoride, calcium loses two electrons and each of two fluorine atoms gains one:
$$\ce{Ca -> Ca^2+ + 2e^-}\qquad \ce{2F + 2e^- -> 2F^-}\qquad \ce{Ca^2+ + 2F^- -> CaF2}$$
The compound is written $\ce{Ca^2+(F^-)2}$, i.e. $\ce{CaF2}$. NEET commonly tests this charge-balancing step: knowing the valencies of the ions, you write the simplest whole-number formula.
Factors That Favour an Ionic Bond
From the Kössel–Lewis treatment, the formation of an ionic compound depends on two things: the ease of forming the positive and negative ions from neutral atoms, and the arrangement of those ions into the crystal lattice. Three energy quantities therefore decide whether an ionic bond is favourable.
| Factor | Requirement for ionic bonding | Why it helps |
|---|---|---|
| Ionisation enthalpy of the metal | Should be low | Less energy is needed to remove the electron and form the cation $\ce{M(g) -> M+(g) + e^-}$. Ionisation is always endothermic. |
| Electron gain enthalpy ($\Delta_{eg}H$) of the non-metal | Should be high negative | More energy is released as the anion forms $\ce{X(g) + e^- -> X^-(g)}$, making anion formation favourable. |
| Lattice enthalpy of the crystal | Should be high | A large amount of energy is released when gaseous ions assemble into the lattice; this is the decisive stabilising term. |
NCERT puts the first two together: ionic bonds form more easily between elements with comparatively low ionisation enthalpies and elements with comparatively high negative values of electron gain enthalpy. This is exactly why ionic compounds typically pair a Group 1 or 2 metal (low ionisation enthalpy) with a halogen or oxygen (high negative electron gain enthalpy). NIOS lists the same three favourable conditions: low ionisation energy of the metal, high electron affinity of the non-metal, and high lattice energy.
Sign convention for electron gain enthalpy
"High electron affinity" and "high negative electron gain enthalpy" describe the same favourable situation. Electron affinity is the negative of the energy change accompanying electron gain, so a large positive electron affinity corresponds to a large negative $\Delta_{eg}H$. Do not read "high $\Delta_{eg}H$" as a large positive number.
Favourable anion formation ⇒ $\Delta_{eg}H$ strongly negative ⇒ electron affinity large and positive.
The third factor deserves its own treatment. See Lattice Enthalpy & the Born–Haber Cycle for the full thermochemical accounting and the factors that govern lattice enthalpy.
Energetics: Why NaCl Forms
A natural objection arises: forming $\ce{Na+}$ costs more energy than is released when $\ce{Cl-}$ forms, so how can the compound be stable? The resolution is that octet attainment in the gas phase is not the measure of stability — lattice formation is. NCERT states the qualitative measure of the stability of an ionic compound is provided by its enthalpy of lattice formation, and not simply by achieving an octet around the ionic species in the gaseous state.
The numbers make this concrete. The ionisation enthalpy for $\ce{Na+(g)}$ formation from $\ce{Na(g)}$ is $+495.8\ \text{kJ mol}^{-1}$, while the electron gain enthalpy for $\ce{Cl(g) + e^- -> Cl^-(g)}$ is only $-348.7\ \text{kJ mol}^{-1}$. Their sum, $+147.1\ \text{kJ mol}^{-1}$, is endothermic. But the lattice enthalpy of $\ce{NaCl(s)}$ is $-788\ \text{kJ mol}^{-1}$, which more than compensates. Because the energy released exceeds the energy absorbed, the overall process is exothermic and the crystal is stable.
| Energy term | Process | Value (kJ mol⁻¹) |
|---|---|---|
| Ionisation enthalpy of Na | $\ce{Na(g) -> Na+(g) + e^-}$ | +495.8 |
| Electron gain enthalpy of Cl | $\ce{Cl(g) + e^- -> Cl^-(g)}$ | −348.7 |
| Sum (gas-phase ion pair) | net of the two above | +147.1 |
| Lattice enthalpy of NaCl | $\ce{Na+(g) + Cl^-(g) -> NaCl(s)}$ | −788 |
NIOS breaks the formation of NaCl from its elements into five steps and sums them by the law of conservation of energy:
a) Sublimation: $\ce{Na(s) -> Na(g)}$, $\Delta H = +108.7$
b) Ionisation: $\ce{Na(g) -> Na+(g) + e^-}$, $\Delta H = +493.8$
c) Dissociation: $\ce{1/2 Cl2(g) -> Cl(g)}$, $\Delta H = +120.9$
d) Electron gain: $\ce{Cl(g) + e^- -> Cl^-(g)}$, $\Delta H = -379.5$
e) Lattice formation: $\ce{Na+(g) + Cl^-(g) -> NaCl(s)}$, $\Delta H = -754.8$
Net: $\ce{Na(s) + 1/2 Cl2(g) -> NaCl(s)}$, $\Delta_f H = (108.7 + 493.8 + 120.9 - 379.5 - 754.8) = -410.9\ \text{kJ mol}^{-1}$.
Of the five terms, sublimation and dissociation are comparatively small, so ionisation energy, electron affinity and lattice energy are the deciding quantities. (Values are quoted from NIOS; the cycle illustrates the same principle as the NCERT NaCl figures above.)
The Crystal Lattice
Ionic compounds in the crystalline state consist of orderly three-dimensional arrangements of cations and anions held together by coulombic interaction energies. The lattice is not a collection of discrete molecules — there is a charge balance between the positive and negative ions across the whole crystal, and the structure is stabilised by the enthalpy of lattice formation. Each compound crystallises in a structure determined by the sizes of the ions, their packing arrangement and other factors; sodium chloride adopts the rock-salt structure.
Properties of Ionic Compounds
The physical behaviour of ionic compounds follows directly from the strong, non-directional coulombic forces that bind a three-dimensional lattice. NIOS §4.3.2 lists the characteristic properties, each of which is a standard NEET fact.
| Property | Observation | Reason |
|---|---|---|
| Physical state | Crystalline solids; generally hard and brittle | Ions occupy fixed positions in a regular three-dimensional lattice. |
| Melting / boiling point | High | Strong electrostatic interactions between ions must be overcome; large thermal energy is required. |
| Solubility | Generally soluble in water; less soluble in non-polar solvents (ether, alcohol) | Polar water molecules solvate and stabilise the separated ions; non-polar solvents cannot. |
| Electrical conductivity | Conduct in the molten state or in aqueous solution; do not conduct as solids | Conduction needs mobile ions. In the solid the ions are fixed; melting or dissolution frees them. |
Solid ionic compounds do NOT conduct
A common error is to mark "ionic compounds conduct electricity" without qualification. Conduction requires charge carriers free to move; in the solid the ions are locked in the lattice. Conduction occurs only in the molten state or in aqueous solution.
Solid = non-conductor; molten or aqueous = conductor (mobile ions).
Scope and Exceptions
Kössel's electron-transfer theory explains bonding well, but only for a small class of solids made of electropositive Group 1 and Group 2 elements combined with highly electronegative elements. It cannot account for molecules such as $\ce{O2}$ or $\ce{SO2}$ — in $\ce{O2}$ there is no reason for one oxygen atom to lose two electrons while the other accepts them. That gap is filled by the Lewis theory of the covalent bond.
Two exceptions are worth memorising. First, most ionic compounds have metal-derived cations and non-metal-derived anions, but the ammonium ion $\ce{NH4+}$ — built from two non-metallic elements — is the cation of a number of ionic compounds. Second, octet attainment alone never guarantees stability; the lattice enthalpy term decides, which is precisely why the gas-phase energetics of NaCl looked unfavourable yet the solid is stable.
Ionic / Electrovalent Bond at a glance
- Ionic bond = electrostatic attraction between a cation and an anion formed by complete electron transfer from metal to non-metal; e.g. $\ce{Na+ + Cl- -> NaCl}$.
- Electrovalence = number of unit charges on the ion ($\ce{Ca^2+}$ → +2, $\ce{Cl-}$ → −1).
- Favourable factors: low ionisation enthalpy of the metal, high negative $\Delta_{eg}H$ of the non-metal, high lattice enthalpy.
- NaCl energetics: IE of Na (+495.8) + $\Delta_{eg}H$ of Cl (−348.7) = +147.1, more than offset by lattice enthalpy (−788) ⇒ net exothermic, stable.
- Stability is set by lattice enthalpy, not by octet attainment in the gas phase.
- Properties: hard brittle crystalline solids; high m.p./b.p.; soluble in water; conduct only when molten or aqueous.
- Exception: $\ce{NH4+}$ is a non-metallic cation; Kössel's theory fails for $\ce{O2}$, $\ce{SO2}$.