Sharing, Not Transfer
The Kössel-Lewis approach offered two routes to the stable noble-gas octet. The first, electron transfer, produces ions and the ionic bond. The second, set out by Lewis in 1916 and refined by Langmuir in 1919, achieves the octet by the mutual sharing of a pair of electrons between two atoms. The attractive force arising from this shared pair is the covalent bond, and the species it holds together are called covalent compounds.
The classic illustration is the chlorine molecule. Each chlorine atom has the configuration $\ce{[Ne]}3s^2 3p^5$ and is one electron short of the argon octet. Rather than one atom stripping an electron from the other, the two atoms pool one electron each into a shared pair. Both atoms then "count" all eight electrons of the shared region, and each attains the argon configuration.
Two chlorine atoms, each with seven valence electrons, contribute one electron each to a single shared pair. The pair is owned jointly, completing the octet of both atoms.
Such electron-dot diagrams are called Lewis dot structures. In them, the electrons shown in the bond are the bond pair, while the pairs not involved in bonding — three on each chlorine here — are the lone pairs. The nature of these pairs governs molecular shape, a thread picked up in VSEPR theory.
Conditions for a Covalent Bond
NCERT distils the Lewis-Langmuir picture into three conditions that every covalent bond must satisfy. These conditions apply equally whether the combining atoms are identical, as in $\ce{Cl2}$, or different, as in $\ce{H2O}$ and $\ce{CCl4}$.
| Condition | Meaning | Example |
|---|---|---|
| Sharing of an electron pair | Every bond is one shared pair of electrons between the two atoms | $\ce{H-H}$ in $\ce{H2}$ |
| Each atom contributes | Each combining atom donates at least one electron to the shared pair | 1 e⁻ from each Cl in $\ce{Cl2}$ |
| Octet (or duplet) attained | Sharing leaves each atom with the noble-gas configuration | Ar octet for Cl; He duplet for H |
Hydrogen is the standing exception to the octet: it completes the duplet of helium with just two shared electrons. This is why in water, carbon tetrachloride and methane the hydrogen atoms each carry one bond pair and no lone pairs, while the central atom satisfies a full octet.
Single, Double and Triple Bonds
The number of electron pairs shared between the same two atoms defines the bond multiplicity. When one pair is shared the atoms are joined by a single bond; sharing two pairs gives a double bond; sharing three pairs gives a triple bond. Multiple bonding is invoked whenever a single bond cannot satisfy the octet of every atom.
| Bond type | Pairs shared | Electrons | Representative molecules |
|---|---|---|---|
| Single bond | 1 | 2 | $\ce{H2}$, $\ce{Cl2}$, $\ce{F2}$, C–C in ethane |
| Double bond | 2 | 4 | $\ce{O2}$, each C=O of $\ce{CO2}$, C=C in ethene ($\ce{C2H4}$) |
| Triple bond | 3 | 6 | $\ce{N2}$, C≡C in ethyne ($\ce{C2H2}$) |
In carbon dioxide each oxygen forms a double bond to the central carbon, giving $\ce{O=C=O}$; in the nitrogen molecule the two atoms share three pairs, $\ce{N#N}$, the triple bond that makes $\ce{N2}$ so inert. The carbon atoms of ethene are joined by a double bond and those of ethyne by a triple bond, a pattern NEET tests directly by asking how many σ and π bonds lie between a given pair of carbons.
Bond order rises as bond length falls
Students often assume "more bonds means a longer bond." The opposite is true. As more electron pairs are shared, the two nuclei are drawn closer, so a triple bond is shorter and stronger than a double bond, which is shorter and stronger than a single bond. Bond order and bond length move in opposite directions.
Single < double < triple in strength; single > double > triple in length.
Lewis Representation of Molecules
A Lewis structure makes the bonding in a molecule or ion visible in terms of shared pairs and lone pairs. NCERT prescribes a systematic procedure for drawing them, which converts a molecular formula into a correct electron-dot picture.
| Step | What to do |
|---|---|
| 1 | Add the valence electrons of all combining atoms. For $\ce{CH4}$ this is 4 (C) + 4 (H) = 8 electrons. |
| 2 | For an anion add one electron per negative charge; for a cation subtract one per positive charge. $\ce{CO3^2-}$ gains 2; $\ce{NH4+}$ loses 1. |
| 3 | Place the least electronegative atom at the centre (e.g. N in $\ce{NF3}$, C in $\ce{CO3^2-}$); distribute pairs as single bonds. |
| 4 | Use leftover pairs for lone pairs or, where the octet is incomplete, for multiple bonds. |
Write the Lewis dot structure of the carbon monoxide molecule, $\ce{CO}$.
Step 1. Valence electrons: carbon $2s^2 2p^2$ gives 4, oxygen $2s^2 2p^4$ gives 6, total $4 + 6 = 10$.
Step 2. Skeleton C–O. A single bond plus completing oxygen's octet leaves carbon short of an octet.
Step 3. Resort to multiple bonding — a triple bond between C and O, $\ce{C#O}$, satisfies the octet on both atoms.
The same logic applied to the nitrite ion $\ce{NO2-}$ — total $5 + (2\times 6) + 1 = 18$ electrons — forces a double bond between nitrogen and one oxygen so that nitrogen attains its octet. Lewis structures do not show the true shapes of molecules, but they reliably account for which bonds form and how many lone pairs remain.
Lewis symbols, the octet rule and its limitations are set out in the Kössel–Lewis approach.
The Coordinate (Dative) Bond
In an ordinary covalent bond each atom contributes one electron to the shared pair. In a coordinate covalent or dative bond, both electrons of the shared pair are supplied by a single species — an electron-rich donor — while the acceptor contributes none. The textbook example is the reaction between ammonia and boron trifluoride: $\ce{NH3}$ carries a lone pair on nitrogen and is electron-rich, while $\ce{BF3}$ is electron-deficient and can accept a pair. The donated lone pair forms the new bond, written $\ce{H3N->BF3}$.
A coordinate bond is conventionally drawn as an arrow pointing from donor to acceptor. Critically, once formed it is identical to a normal covalent bond in polarity and strength; donor and acceptor origin cannot be told apart afterwards. The ammonium ion $\ce{NH4+}$ and nitric acid $\ce{HNO3}$ are further everyday examples.
Ammonia's lone pair on nitrogen is donated to a bare proton $\ce{H+}$ (which brings no electrons), creating a fourth N–H bond. All four bonds become equivalent once formed.
For NEET, the most common form of the question is "how many coordinate bonds are present" in a given species. In $\ce{NH4+}$ the answer is exactly one: three N–H bonds are ordinary covalent bonds inherited from ammonia, and the fourth arises from donation of nitrogen's lone pair to $\ce{H+}$. The same reasoning identifies one coordinate bond in $\ce{H3O+}$, where water donates a lone pair to a proton.
Ionic Versus Covalent Bonding
Both bonds chase the same goal — the noble-gas octet — but by opposite mechanisms. The ionic (electrovalent) bond arises from complete electron transfer and electrostatic attraction between oppositely charged ions; the covalent bond arises from shared electron pairs. The contrast is summarised below.
| Feature | Ionic (electrovalent) bond | Covalent bond |
|---|---|---|
| Mechanism | Transfer of electrons → ions | Sharing of electron pairs |
| Bound species | Cations and anions in a lattice | Discrete molecules |
| Typical partners | Electropositive metal + electronegative non-metal | Two non-metals of similar electronegativity |
| Example | $\ce{NaCl}$, $\ce{CaF2}$ | $\ce{H2}$, $\ce{Cl2}$, $\ce{H2O}$, $\ce{CCl4}$ |
| Directionality | Non-directional | Directional |
The boundary is not sharp. When the two bonded atoms differ in electronegativity, the shared pair lies nearer the more electronegative atom, giving a polar covalent bond with partial ionic character, as in $\ce{H^{δ+}-Cl^{δ-}}$. At an electronegativity difference of about 1.7 the bond is roughly 50% ionic; beyond that, polarisation is essentially complete and the bond is best described as ionic. Pure covalent bonds occur only between identical atoms — $\ce{H2}$, $\ce{O2}$, $\ce{N2}$ — where sharing is exactly equal and the bond is non-polar.
A covalent bond is not an intermolecular force
Intermolecular forces — dipole–dipole, dipole–induced dipole, hydrogen bonding and dispersion forces — act between molecules. A covalent bond holds the atoms within a molecule together and is far stronger. NEET 2023 explicitly excluded covalent bonding from the list of intermolecular forces, so never count it among them.
Covalent bond = intramolecular (atoms in a molecule); van der Waals/H-bonding = intermolecular.
Properties of Covalent Compounds
Because covalent compounds exist as discrete molecules held to one another only by weak intermolecular forces, their bulk properties differ sharply from those of ionic solids. NIOS §4.4.3 lists three characteristic properties that NEET likes to contrast against ionic behaviour.
| Property | Behaviour | Reason |
|---|---|---|
| Melting / boiling point | Low | Weak intermolecular forces between molecules are easily overcome |
| Electrical conductivity | Poor conductors | No free ions are present |
| Solubility | Generally insoluble in water; soluble in non-polar solvents | Dissolve readily in benzene, $\ce{CCl4}$ and similar non-polar media |
These properties flow directly from the molecular nature of the bond. The covalent bond inside each molecule is strong — a few hundred kilojoules per mole — but the forces holding one molecule to its neighbours are weak, so the substance melts and vaporises easily and stays electrically inert. Recognising this molecular-versus-ionic contrast resolves a large fraction of physical-property questions in the bonding chapter.
The Covalent Bond at a glance
- A covalent bond is the force from sharing one or more electron pairs; each atom contributes at least one electron and attains the octet (H, the duplet).
- Sharing one, two or three pairs gives single, double and triple bonds respectively; bond strength rises and bond length falls along this series.
- Examples: single — $\ce{H2}$, $\ce{Cl2}$; double — $\ce{O2}$, $\ce{CO2}$, ethene; triple — $\ce{N2}$, ethyne.
- A coordinate (dative) bond has both shared electrons from one donor; $\ce{NH4+}$ and $\ce{H3O+}$ each contain exactly one, identical to the others once formed.
- Unequal sharing gives polar covalent bonds with partial ionic character; identical atoms give pure non-polar bonds.
- Covalent compounds: low melting/boiling points, poor conductors, soluble in non-polar solvents — all because of weak intermolecular forces.