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CBSE • Class 11 • Chemistry

Chemical Bonding and Molecular Structure

Ionic bonding, covalent bonding, VSEPR, hybridisation, and molecular orbital ideas.

Chapter 4

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What is Chemical Bonding and Molecular Structure?

Ionic bonding, covalent bonding, VSEPR, hybridisation, and molecular orbital ideas.

Chemical Bonding and Molecular Structure matters because it links chemical ideas, reactions, and reasoning patterns that recur throughout the syllabus. At Class 11 level, students are often expected to define terms accurately, explain processes clearly, and connect theory to reactions, observations, or applications.

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Summary

The One Thing

Chemical bonding reflects the tendency of atoms to adopt lower-energy, more stable arrangements through electron transfer or sharing. Lewis structures, VSEPR theory, hybridisation, and molecular orbital theory explain how these arrangements determine composition, geometry, bond strength, polarity, and magnetic behaviour.

Reactions, Processes and Experiments

What happensEquation or processWhat you observeType
Atoms form bonds through electron transfer or electron sharing, producing a more stable, lower-energy arrangement.Atoms arrange and share or transfer electrons in ways that lower energy.Greater stability than that of the separated atoms.General bonding process
A Lewis structure is constructed by counting total valence electrons, selecting a skeleton, completing the outer atoms, placing remaining electrons on the central atom, and checking formal charges.Counting total valence electrons → choosing a skeleton → completing outer atoms → placing remaining electrons on the central atom → checking formal charges.Shared electron pairs and lone pairs are shown; the sum of formal charges equals the overall charge.Lewis-structure construction
An ionic bond forms after electron transfer, commonly from a metal to a non-metal, followed by electrostatic attraction between the ions.Electron transfer → oppositely charged ions → electrostatic attraction.Formation of positive and negative ions; ionic character increases with greater electronegativity difference.Ionic bonding
Ionic compound formation can be analysed using the Born-Haber concept.Sublimation → ionisation → bond dissociation → electron gain → lattice formation.The process involves gaseous ions and formation of an ionic lattice.Born-Haber process
An ionic solid forms from gaseous ions, or the reverse process occurs, depending on the convention used.Lattice enthalpy: formation of one mole of an ionic solid from its gaseous ions, or the reverse process.Lattice enthalpy generally increases in magnitude when ionic charges increase and ionic sizes decrease.Lattice formation
A covalent bond forms when atoms mutually share one or more pairs of electrons.Mutual sharing of one or more pairs of electrons.Shared electron pairs connect the atoms; covalent bonds may be single, double, or triple.Covalent bonding
A coordinate bond forms when both shared electrons are donated by the same atom.Donation of both shared electrons by one atom.The shared pair originates from one atom.Coordinate covalent bonding
Covalent bonds may contain one, two, or three shared electron pairs.Single bond = one sigma bond; double bond = one sigma and one pi bond; triple bond = one sigma and two pi bonds.Additional bonds in a multiple bond are pi bonds; the first bond between two atoms is a sigma bond.Bond formation
Bond order affects bond length and strength.Higher bond order → shorter bond length and greater bond strength.Higher-order bonds are generally shorter and stronger.Bond-property relationship
Bond dipoles combine according to molecular structure.μ = q × rPolar bonds may produce a non-polar molecule when bond dipoles cancel through symmetry, as in carbon dioxide.Dipole formation
Electron pairs arrange around a central atom to minimise repulsions.Lone pair-lone pair > lone pair-bond pair > bond pair-bond pair.Lone pairs usually compress bond angles more strongly than bonding pairs.VSEPR arrangement
Two electron domains adopt a linear arrangement.2 electron domains, linear, 180°.Linear geometry with a 180° angle.Electron-pair geometry
Three electron domains adopt a trigonal planar arrangement.3, trigonal planar, 120°.Trigonal planar geometry with approximately 120° angles.Electron-pair geometry
Four electron domains adopt a tetrahedral arrangement.4, tetrahedral, 109.5°.Tetrahedral electron-pair geometry with 109.5° angles.Electron-pair geometry
Five electron domains adopt a trigonal bipyramidal arrangement.5, trigonal bipyramidal, 90° and 120°.Trigonal bipyramidal geometry; lone pairs prefer equatorial positions.Electron-pair geometry
Six electron domains adopt an octahedral arrangement.6, octahedral, 90°.Octahedral geometry with 90° angles.Electron-pair geometry
Hybrid orbitals form by mixing atomic orbitals of similar energy on the same atom.Mixing of atomic orbitals of similar energy → equivalent hybrid orbitals.Equivalent bond directions and characteristic geometries are produced.Hybridisation
The central atom in BeCl2 and the central carbon in CO2 use sp hybridisation.sp hybridisation gives linear geometry and 180° bond angles, as in BeCl2 and CO2 around the central atom.Linear arrangement and 180° bond angles.sp hybridisation
The boron atom in BF3 and the carbon atoms of ethene use sp2 hybridisation.sp2 hybridisation gives trigonal planar geometry and approximately 120° bond angles, as in BF3 and the carbon atoms of ethene.Trigonal planar arrangement with approximately 120° angles.sp2 hybridisation
sp3 hybridisation gives tetrahedral electron-pair geometry, modified by lone pairs where present.sp3 hybridisation gives tetrahedral electron-pair geometry; CH4 is tetrahedral, NH3 is trigonal pyramidal, and H2O is bent because of lone pairs.CH4 is tetrahedral; NH3 is trigonal pyramidal; H2O is bent.sp3 hybridisation
PCl5 has trigonal bipyramidal electron-pair geometry.sp3d hybridisation is associated with trigonal bipyramidal electron-pair geometry, as in PCl5; lone pairs prefer equatorial positions.Trigonal bipyramidal arrangement; equatorial positions are preferred by lone pairs.sp3d hybridisation
SF6 has octahedral electron-pair geometry.sp3d2 hybridisation is associated with octahedral electron-pair geometry, as in SF6.Octahedral arrangement with 90° angles.sp3d2 hybridisation
Sigma bonds form through head-on orbital overlap along the internuclear axis.Head-on overlap of orbitals along the internuclear axis.Electron density is concentrated along the internuclear axis.Sigma bonding
Pi bonds form through sideways overlap of parallel orbitals.Sideways overlap of parallel orbitals above and below the internuclear axis.Electron density occurs above and below the internuclear axis.Pi bonding
Atomic orbitals combine to form molecular orbitals when their energies, symmetry, and overlap are suitable.Comparable energies + suitable symmetry + appreciable overlap → molecular orbitals.The number of molecular orbitals formed equals the number of atomic orbitals combined.Molecular-orbital formation
Combining atomic orbitals produces bonding and antibonding molecular orbitals.Atomic orbitals → bonding molecular orbitals and antibonding molecular orbitals.Bonding molecular orbitals are lower in energy; antibonding molecular orbitals are higher in energy and marked with an asterisk.Molecular orbital formation
Electrons occupy molecular orbitals according to standard electron-filling rules.Aufbau principle, Pauli exclusion principle, and Hund's rule.Electrons fill lower-energy orbitals first, occupy orbitals consistently with spin restrictions, and remain unpaired where required by Hund's rule.Molecular-orbital electron arrangement
Molecular-orbital bond order measures net stabilisation.Bond order in molecular orbital theory = 1/2(number of electrons in bonding molecular orbitals - number of electrons in antibonding molecular orbitals).Positive bond order generally indicates stability; bond order zero indicates no net stabilisation.Molecular-orbital bond order
The hydrogen molecule forms a stable bond.The hydrogen molecule has bond order 1 and is diamagnetic.Bond order 1; all electrons are paired, so the molecule is diamagnetic.Molecular-orbital description of H2
The helium molecule is not stable as a normal molecule.The helium molecule would have bond order 0 and is not stable as a normal molecule.Bond order zero; no net stabilisation.Molecular-orbital description of He2
Unpaired molecular-orbital electrons produce paramagnetism.Unpaired electrons in molecular orbitals cause paramagnetism, while all paired electrons generally result in diamagnetism.Paramagnetic substances contain unpaired electrons; diamagnetic substances generally have all electrons paired.Magnetic behaviour
Hydrogen bonding occurs when hydrogen is bonded to a highly electronegative atom and is attracted to another electronegative atom.Hydrogen bonded to nitrogen, oxygen, or fluorine → strong intermolecular attraction.Hydrogen bonding raises boiling points, affects solubility, and helps explain the unusual properties of water and biological molecular structures.Hydrogen bonding
Covalent character in an ionic compound increases when the cation strongly polarises the anion.Fajan's rules: small, highly charged cation + large, easily polarised anion → increased covalent character.Ionic compounds may have partly ionic and partly covalent bonding rather than being completely one type.Polarisation and bonding character

Key Terms

  • Chemical bond: An attractive force that holds atoms or ions together in a molecule or compound.
  • Octet rule: Atoms generally tend to gain, lose, or share electrons to obtain eight electrons in their valence shell, although there are exceptions.
  • Lewis structure: A diagram showing valence electrons, shared electron pairs, and lone pairs in a molecule or ion.
  • Ionic bond: The electrostatic attraction between oppositely charged ions formed after electron transfer, commonly from a metal to a non-metal.
  • Lattice enthalpy: The enthalpy change associated with the formation of one mole of an ionic solid from its gaseous ions, or the reverse process depending on the convention used.
  • Covalent bond: A bond formed by the mutual sharing of one or more pairs of electrons between atoms.
  • Coordinate bond: A covalent bond in which both shared electrons are donated by the same atom.
  • Bond length: The average distance between the nuclei of two bonded atoms.
  • Bond angle: The angle between two bonds around a central atom.
  • Bond enthalpy: The average energy required to break one mole of a particular type of bond in gaseous molecules.
  • Bond order: The number of bonding electron pairs shared between two atoms; higher bond order generally means shorter and stronger bonds.
  • Resonance: The representation of a molecule or ion by two or more valid Lewis structures when one structure cannot describe the electron distribution completely.
  • Formal charge: The hypothetical charge assigned to an atom in a Lewis structure by assuming equal sharing of bonding electrons.
  • Electronegativity: The tendency of an atom in a chemical bond to attract the shared electron pair toward itself.
  • Dipole moment: A measure of charge separation in a bond or molecule, represented by the product of charge and distance.
  • VSEPR theory: The theory that electron pairs around a central atom arrange themselves as far apart as possible to minimise repulsions.
  • Lone pair: A pair of valence electrons not involved in bonding; it occupies space and usually repels more strongly than a bonding pair.
  • Hybridisation: The mixing of atomic orbitals of similar energy on the same atom to form equivalent hybrid orbitals used in bonding.
  • Sigma bond: A covalent bond formed by head-on overlap of orbitals along the internuclear axis.
  • Pi bond: A covalent bond formed by sideways overlap of parallel orbitals above and below the internuclear axis.
  • Molecular orbital: An orbital spread over the entire molecule that can contain electrons contributed by more than one atom.
  • Bonding molecular orbital: A molecular orbital lower in energy than the combining atomic orbitals; electrons in it stabilise the molecule.
  • Antibonding molecular orbital: A molecular orbital higher in energy than the combining atomic orbitals; electrons in it oppose bonding and are marked with an asterisk.
  • Hydrogen bond: A strong intermolecular attraction involving hydrogen bonded to a highly electronegative atom such as nitrogen, oxygen, or fluorine.

Easily Confused

  • Ionic bond vs covalent bond: Ionic bonding involves electron transfer and electrostatic attraction between oppositely charged ions; covalent bonding involves mutual sharing of electron pairs.
  • Covalent bond vs coordinate bond: In an ordinary covalent bond, each atom contributes to the shared pair; in a coordinate bond, both electrons are donated by one atom.
  • Molecular shape vs electron-pair geometry: Molecular shape considers only atom positions; electron-pair geometry includes bonding pairs and lone pairs.
  • Bonding molecular orbital vs antibonding molecular orbital: A bonding molecular orbital is lower in energy and stabilising; an antibonding molecular orbital is higher in energy and destabilising.
  • Paramagnetism vs diamagnetism: Paramagnetism results from unpaired molecular-orbital electrons; diamagnetism generally results when all electrons are paired.
  • Bond order vs bond enthalpy: Bond order counts bonding electron pairs or expresses net molecular-orbital stabilisation; bond enthalpy is the average energy required to break a particular bond in gaseous molecules.
  • Polar bond vs polar molecule: A molecule may contain polar bonds but be non-polar overall if molecular symmetry causes the bond dipoles to cancel, as in carbon dioxide.
  • Lone pair vs bonding pair: A lone pair is not involved in bonding and generally repels more strongly than a bonding pair.
  • Single, double, and triple bonds: A single bond contains one sigma bond; a double bond contains one sigma and one pi bond; a triple bond contains one sigma and two pi bonds.
  • Duet rule vs octet rule: Hydrogen and helium follow a duet rule, whereas atoms generally tend toward eight valence-shell electrons.
  • Hybridisation vs molecular orbital theory: Hybridisation is a model based on mixing orbitals on one atom; molecular orbital theory describes orbitals spread across the entire molecule.

What Gets Asked

  • Constructing Lewis structures: Questions require valence-electron counting, skeleton selection, placement of lone pairs, and formal-charge checks. Marks are lost by omitting lone pairs, using the wrong total electron count, or failing to match the sum of formal charges to the overall charge.
  • Applying the octet rule: Questions may test hydrogen and helium's duet rule, incomplete octets in second-period elements, expanded valence shells from the third period onward, and odd-electron species. Marks are lost by treating the octet rule as universal.
  • Comparing ionic and covalent character: Questions may require lattice-enthalpy trends, electronegativity differences, Born-Haber steps, or Fajan's rules. Marks are lost by describing compounds as completely ionic or completely covalent.
  • Determining molecular shape and bond angles: Questions may provide electron domains or lone pairs and ask for VSEPR geometry. Marks are lost by confusing electron-pair geometry with molecular shape or by ignoring the stronger repulsion of lone pairs.
  • Assigning hybridisation and sigma/pi bonds: Questions may involve BeCl2, CO2, BF3, ethene, CH4, NH3, H2O, PCl5, or SF6. Marks are lost by counting every bond in a multiple bond as a sigma bond rather than identifying one sigma bond and the additional pi bonds.
  • Using molecular orbital theory: Questions may ask for molecular-orbital filling, bond order, stability, or magnetic behaviour for H2 and He2. Marks are lost by reversing bonding and antibonding energy relationships, miscalculating bond order, or overlooking unpaired electrons.

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Key ideas to master

  • Learn the precise terms, laws, and reaction patterns associated with Chemical Bonding and Molecular Structure.
  • Understand why each step or change happens instead of memorising the result only.
  • Practise writing balanced equations, comparisons, or structured explanations where relevant.
  • Revise common exceptions, observations, and applications that examiners often test.

Common exam prompts

  • Define the main idea in Chemical Bonding and Molecular Structure using correct chemical terminology.
  • Write or interpret the reactions, observations, or comparisons that belong to this topic.
  • Explain why a process happens, not just what happens.
  • Summarise the high-yield facts and exceptions examiners often choose from this chapter.

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What is Chemical Bonding and Molecular Structure in CBSE Class 11 Chemistry?

Ionic bonding, covalent bonding, VSEPR, hybridisation, and molecular orbital ideas.

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