Unit 5
General and Physical Chemistry
Class 11 Chemistry
Chemical Bonding and Shapes of Molecules
On mobile, swipe inside the PDF to read all pages and pinch to zoom.
NEB/CDC syllabus scope: Unit 5 – Chemical Bonding and Shapes of Molecules is a 9-teaching-hour unit covering valence shell/electrons and octet theory; ionic, covalent and coordinate covalent bonding; properties of ionic/covalent compounds; Lewis structures; resonance; VSEPR shapes of BeF₂, BF₃, CH₄, CH₃Cl, PCl₅, SF₆, H₂O, NH₃, CO₂, H₂S and PH₃; elementary valence-bond theory; hybridization involving s and p orbitals; bond length, ionic character and dipole moment; van der Waals forces and molecular solids; hydrogen bonding; and metallic bonding.
1. Why Do Atoms Form Chemical Bonds?
Chemical bond
A chemical bond is the attractive interaction that holds atoms or ions together in a stable chemical species.
A chemical bond is the attractive interaction that holds atoms or ions together in a stable chemical species.
Bond formation is favored when the bonded arrangement has lower energy than the separated particles. Valence electrons are the electrons mainly involved in ordinary chemical bonding.
Stability idea
The octet rule is a useful introductory pattern, not a universal law. Many stable species do not obey a simple octet.
The octet rule is a useful introductory pattern, not a universal law. Many stable species do not obey a simple octet.
Diagram 1: Stable bonding corresponds to lower energy
2. Valence Shell, Valence Electrons and Octet Theory
Valence shell
The outermost occupied shell of an atom.
The outermost occupied shell of an atom.
Valence electrons
Electrons present in the valence shell and most directly involved in bond formation.
Electrons present in the valence shell and most directly involved in bond formation.
Octet rule
Many main-group atoms tend to gain, lose or share electrons so that their valence shell resembles a noble-gas configuration with eight electrons.
Many main-group atoms tend to gain, lose or share electrons so that their valence shell resembles a noble-gas configuration with eight electrons.
Common Exceptions
- Duet: H and He are stable with two electrons in the first shell.
- Incomplete octet: BeF₂ and BF₃ are common electron-deficient examples.
- Odd-electron species: some radicals cannot give every atom an octet.
- Expanded valence shell: species such as PCl₅ and SF₆ are represented with more than eight electrons around the central atom in standard Lewis structures.
3. Lewis Symbols
Lewis symbols show the valence electrons as dots around an element symbol.
| Element | Valence electrons | Lewis-symbol idea |
|---|---|---|
| Na | 1 | Na· |
| Mg | 2 | ·Mg· |
| C | 4 | four single dots around C |
| O | 6 | two lone pairs + two single electrons |
| Cl | 7 | three lone pairs + one single electron |
4. Ionic / Electrovalent Bond
Ionic bond
An electrostatic attraction between oppositely charged ions produced after electron transfer.
An electrostatic attraction between oppositely charged ions produced after electron transfer.
Formation of NaCl
Na → Na⁺ + e⁻ Cl + e⁻ → Cl⁻ Na⁺ + Cl⁻ → NaCl(s)The solid is not made of separate NaCl molecules; it forms an extended ionic lattice.
Diagram 2: Electron transfer and ion attraction in NaCl
5. Properties of Ionic Compounds
- Usually crystalline solids.
- Generally high melting and boiling points due to strong electrostatic attractions.
- Hard but often brittle.
- Do not conduct electricity as intact solids because ions are not free to move.
- Conduct when molten or dissolved in a suitable solvent because ions become mobile.
- Often soluble in polar solvents such as water; solubility varies with lattice and solvation energies.
6. Covalent Bond
Covalent bond
A bond formed by sharing one or more pairs of electrons between atoms.
A bond formed by sharing one or more pairs of electrons between atoms.
| Bond | Shared pairs | Example |
|---|---|---|
| Single | 1 | H–H |
| Double | 2 | O=O |
| Triple | 3 | N≡N |
Shared pair
Each covalent bond contains a shared electron pair contributed by the bonded atoms in an ordinary Lewis description.
Each covalent bond contains a shared electron pair contributed by the bonded atoms in an ordinary Lewis description.
7. Coordinate Covalent / Dative Bond
Coordinate covalent bond
A covalent bond in which the shared electron pair originates from one atom or species.
A covalent bond in which the shared electron pair originates from one atom or species.
Example: Ammonium Ion
NH₃ + H⁺ → NH₄⁺Nitrogen donates its lone pair to H⁺. Once formed, the N–H bonds in NH₄⁺ are equivalent in the ordinary structural description.
Donor and Acceptor
- Electron-pair donor: species with a lone pair, e.g. NH₃.
- Electron-pair acceptor: electron-deficient species, e.g. H⁺.
8. General Properties of Covalent Compounds
- Many molecular covalent substances have lower melting/boiling points than ionic solids because intermolecular forces can be much weaker than ionic lattice forces.
- Most molecular covalent substances are poor electrical conductors.
- Many are soluble in non-polar or weakly polar solvents, depending on molecular polarity.
- Covalent bonds are directional, strongly affecting molecular shape.
Do not overgeneralize
Network covalent solids such as diamond have extremely high melting points. “Covalent compounds have low melting points” applies mainly to molecular covalent substances.
Network covalent solids such as diamond have extremely high melting points. “Covalent compounds have low melting points” applies mainly to molecular covalent substances.
9. Lewis Dot Structures
General Procedure
- Count total valence electrons.
- Choose a central atom, usually the least electronegative atom other than hydrogen.
- Connect atoms by single bonds.
- Complete terminal-atom octets first.
- Place remaining electrons on the central atom.
- If needed, form multiple bonds to improve octets and formal charges.
- Check total electron count and charge.
| Species | Lewis-structure idea |
|---|---|
| H₂ | H:H or H–H |
| Cl₂ | Cl–Cl, each Cl with three lone pairs |
| H₂O | H–O–H, O with two lone pairs |
| NH₃ | three N–H bonds, one lone pair on N |
| CO₂ | O=C=O, each O with two lone pairs |
| BF₃ | three B–F bonds; B has an incomplete octet |
Diagram 3: Representative Lewis structures
10. Resonance
Resonance
When one Lewis structure cannot adequately represent electron distribution, two or more contributing structures may be written; the actual species is a resonance hybrid.
When one Lewis structure cannot adequately represent electron distribution, two or more contributing structures may be written; the actual species is a resonance hybrid.
Example: Carbonate Ion
CO₃²⁻ can be drawn with the C=O double bond in three equivalent positions. The actual ion has equivalent C–O bonds with delocalized π-electron density.
Important
The molecule or ion does not jump back and forth between resonance structures. The contributors are representations of one delocalized electronic structure.
The molecule or ion does not jump back and forth between resonance structures. The contributors are representations of one delocalized electronic structure.
Diagram 4: Equivalent resonance contributors of carbonate
11. VSEPR Theory
VSEPR
Valence Shell Electron Pair Repulsion theory predicts molecular geometry by arranging electron domains around a central atom so that repulsions are minimized.
Valence Shell Electron Pair Repulsion theory predicts molecular geometry by arranging electron domains around a central atom so that repulsions are minimized.
Approximate repulsion order:
lone pair–lone pair > lone pair–bond pair > bond pair–bond pair
A multiple bond counts as one electron domain for basic VSEPR geometry, though it can exert somewhat greater repulsion than a single bond.
| Electron domains | Ideal electron-domain geometry | Ideal angle |
|---|---|---|
| 2 | Linear | 180° |
| 3 | Trigonal planar | 120° |
| 4 | Tetrahedral | 109.5° |
| 5 | Trigonal bipyramidal | 90°, 120°, 180° |
| 6 | Octahedral | 90°, 180° |
12. Required Molecular Shapes
| Species | Central-atom domains | Lone pairs on central atom | Molecular shape | Approx. angle |
|---|---|---|---|---|
| BeF₂ | 2 | 0 | Linear | 180° |
| BF₃ | 3 | 0 | Trigonal planar | 120° |
| CH₄ | 4 | 0 | Tetrahedral | 109.5° |
| CH₃Cl | 4 | 0 | Tetrahedral around C | near 109.5° |
| PCl₅ | 5 | 0 | Trigonal bipyramidal | 90°,120°,180° |
| SF₆ | 6 | 0 | Octahedral | 90°,180° |
| H₂O | 4 | 2 | Bent / V-shaped | ~104.5° |
| NH₃ | 4 | 1 | Trigonal pyramidal | ~107° |
| CO₂ | 2 | 0 | Linear | 180° |
| H₂S | 4 electron domains | 2 | Bent | ~92° |
| PH₃ | 4 electron domains | 1 | Trigonal pyramidal | ~93.5° |
Why NH₃ and H₂O angles are smaller than 109.5°
Lone-pair repulsion compresses the bond angle: CH₄ ≈ 109.5° > NH₃ ≈ 107° > H₂O ≈ 104.5°.
Lone-pair repulsion compresses the bond angle: CH₄ ≈ 109.5° > NH₃ ≈ 107° > H₂O ≈ 104.5°.
Diagram 5: Common VSEPR geometries and lone-pair effects
13. Elementary Valence Bond Theory
Valence Bond Theory (VBT)
A covalent bond forms when suitable half-filled atomic orbitals from two atoms overlap and their electrons pair with opposite spins.
A covalent bond forms when suitable half-filled atomic orbitals from two atoms overlap and their electrons pair with opposite spins.
Greater effective overlap generally produces a stronger bond.
Types of Overlap
- s–s overlap
- s–p overlap
- p–p head-on overlap
- p–p sidewise overlap
14. Sigma (σ) and Pi (π) Bonds
| Feature | σ bond | π bond |
|---|---|---|
| Overlap | Head-on / axial | Sidewise / lateral |
| Electron density | Along internuclear axis | Above/below or on sides of axis |
| Relative overlap | Generally greater | Generally smaller |
| Relative strength | Usually stronger | Usually weaker |
| Rotation | Single σ bond can usually rotate comparatively freely | π bonding restricts rotation |
Single bond = 1σ
Double bond = 1σ + 1π
Triple bond = 1σ + 2π
Diagram 6: Head-on σ overlap and sidewise π overlap
15. Hybridization Involving s and p Orbitals
Hybridization
The conceptual mixing of atomic orbitals on the same atom to form a new set of equivalent directional hybrid orbitals used in bonding models.
The conceptual mixing of atomic orbitals on the same atom to form a new set of equivalent directional hybrid orbitals used in bonding models.
| Hybridization | Orbitals mixed | Hybrid orbitals | Ideal geometry | Ideal angle | Example |
|---|---|---|---|---|---|
| sp | 1s + 1p | 2 | Linear | 180° | BeF₂ / carbon in CO₂ model |
| sp² | 1s + 2p | 3 | Trigonal planar | 120° | BF₃ |
| sp³ | 1s + 3p | 4 | Tetrahedral electron-domain geometry | 109.5° | CH₄; NH₃/H₂O electron-domain model |
Syllabus boundary
The curriculum explicitly says hybridization involving s and p orbitals only. Therefore sp³d and sp³d² are not required as hybridization topics here, even though PCl₅ and SF₆ shapes are required under VSEPR.
The curriculum explicitly says hybridization involving s and p orbitals only. Therefore sp³d and sp³d² are not required as hybridization topics here, even though PCl₅ and SF₆ shapes are required under VSEPR.
Diagram 7: Required s/p hybridization types
16. Bond Length
Bond length
The equilibrium internuclear distance between two bonded atoms.
The equilibrium internuclear distance between two bonded atoms.
General qualitative trends:
- Larger atoms generally form longer bonds.
- For the same pair of atoms, higher bond order usually means shorter bond length.
Typical trend: C–C > C=C > C≡C in bond length
17. Ionic Character of a Covalent Bond
When two bonded atoms have different electronegativities, electron density shifts toward the more electronegative atom. The bond becomes polar covalent and has partial ionic character.
Electronegativity difference
Larger electronegativity difference generally corresponds to greater bond polarity/ionic character, but bond type is a continuum rather than a perfectly sharp cutoff.
Larger electronegativity difference generally corresponds to greater bond polarity/ionic character, but bond type is a continuum rather than a perfectly sharp cutoff.
Diagram 8: Polar covalent bond and partial charges
18. Dipole Moment
Dipole moment, μ
A measure of charge separation in a bond or molecule.
A measure of charge separation in a bond or molecule.
μ = q × r
where q is the magnitude of separated charge and r is the separation distance. A common molecular unit is the debye (D).
Bond Polarity vs Molecular Polarity
A molecule may contain polar bonds yet have zero overall dipole moment if the bond dipoles cancel by symmetry.
| Molecule | Bond polarity | Geometry | Net molecular dipole |
|---|---|---|---|
| CO₂ | C=O bonds polar | Linear | 0 by cancellation |
| BF₃ | B–F bonds polar | Trigonal planar | 0 by symmetry |
| H₂O | O–H bonds polar | Bent | Non-zero |
| NH₃ | N–H bonds polar | Trigonal pyramidal | Non-zero |
Diagram 9: Geometry controls net molecular polarity
19. van der Waals Forces and Molecular Solids
van der Waals forces are weak intermolecular attractions that operate between molecules or atoms. Important types include London dispersion forces and permanent dipole–dipole attractions.
London Dispersion Forces
- Present in all atoms and molecules.
- Arise from temporary fluctuations in electron distribution.
- Generally become stronger as electron cloud size/polarizability increases.
Molecular Solids
Molecular solids consist of molecules held together by intermolecular forces rather than an extended ionic, metallic or covalent network. They generally have lower melting points than strongly bonded ionic/network solids.
20. Hydrogen Bonding and Applications
Hydrogen bond
A strong, directional intermolecular attraction involving H covalently bonded to a highly electronegative atom such as N, O or F and a lone pair on another electronegative atom.
A strong, directional intermolecular attraction involving H covalently bonded to a highly electronegative atom such as N, O or F and a lone pair on another electronegative atom.
Intermolecular Hydrogen Bonding
Occurs between separate molecules, e.g. H₂O, HF and alcohol molecules.
Intramolecular Hydrogen Bonding
Occurs within one molecule when geometry permits a suitable donor and acceptor to approach each other.
Effects / Applications
- Raises boiling points of water and many N/O/F-containing substances.
- Contributes to water’s unusual physical properties.
- Important in biomolecular structures such as proteins and nucleic acids.
- Affects solubility and molecular association.
Diagram 10: Hydrogen-bond interaction between water molecules
21. Metallic Bonding and Properties of Metallic Solids
Metallic bond
The collective electrostatic attraction between a lattice of positive metal ion cores and delocalized valence electrons.
The collective electrostatic attraction between a lattice of positive metal ion cores and delocalized valence electrons.
Properties Explained by Metallic Bonding
- Electrical conductivity: delocalized electrons can respond to an electric field.
- Thermal conductivity: mobile electrons and lattice vibrations transfer energy.
- Malleability and ductility: metallic bonding is non-directional enough to allow layers to shift while attraction persists.
- Lustre: conduction electrons interact strongly with incident light.
Diagram 11: Electron-sea picture of metallic bonding
22. Worked Examples
Example 1: Identify the Bond Type
NaCl is formed from Na⁺ and Cl⁻ ions, so the dominant bonding is ionic.
Example 2: Number of σ and π Bonds
Ethene, H₂C=CH₂: the C=C double bond contains 1σ + 1π, and four C–H bonds are σ.
Total = 5σ and 1πExample 3: Acetylene / Ethyne
HC≡CH: C≡C contains 1σ + 2π; two C–H bonds are σ.
Total = 3σ and 2πExample 4: VSEPR Shape of NH₃
N has 3 bonding pairs and 1 lone pair → 4 electron domains → tetrahedral electron-domain geometry → trigonal pyramidal molecular shape.
Example 5: VSEPR Shape of H₂O
O has 2 bonding pairs and 2 lone pairs → 4 electron domains → bent/V-shaped molecule.
Example 6: Hybridization of BF₃
Three electron domains around B → trigonal planar arrangement → introductory hybridization model sp².
Example 7: Hybridization of CH₄
Four equivalent C–H σ bonds arranged tetrahedrally → sp³ carbon model.
Example 8: Molecular Polarity of CO₂
Each C=O bond is polar, but the molecule is linear and symmetric. Equal bond dipoles cancel → non-polar molecule.
Example 9: Molecular Polarity of H₂O
O–H bonds are polar and the molecule is bent, so the dipoles do not cancel → polar molecule.
Example 10: Resonance
In CO₃²⁻, three equivalent Lewis contributors can be drawn. The real ion is one resonance hybrid with delocalized bonding.
23. High-Yield Comparison Tables
Ionic vs Covalent Bonding
| Feature | Ionic | Covalent |
|---|---|---|
| Basic model | Electron transfer + ion attraction | Electron-pair sharing |
| Typical participants | Metal + non-metal | Non-metals |
| Directionality | Mainly non-directional electrostatic lattice interaction | Directional orbital overlap |
| Conduction | Molten/aqueous ions conduct | Most molecular compounds poor conductors |
Intermolecular Forces
| Interaction | Origin | Typical occurrence |
|---|---|---|
| London dispersion | Instantaneous/induced dipoles | All atoms and molecules |
| Dipole–dipole | Permanent molecular dipoles | Polar molecules |
| Hydrogen bonding | Strong directional interaction involving H bonded to N/O/F | Water, HF, NH₃, alcohols, biomolecules |
24. Common Exam Mistakes
- Treating the octet rule as universal.
- Calling NaCl a discrete molecule in the solid state instead of an ionic lattice/formula unit.
- Saying ionic compounds conduct electricity as solids; ions are fixed in the solid lattice.
- Confusing coordinate covalent bonding with ionic bonding.
- Forgetting lone pairs in Lewis structures and VSEPR.
- Counting a double bond as two VSEPR electron domains; it counts as one domain.
- Calling electron-domain geometry and molecular shape identical when lone pairs are present.
- Writing H₂O as linear or NH₃ as tetrahedral molecular shape.
- Writing BF₃ as pyramidal; it is trigonal planar.
- Writing PCl₅ as octahedral; it is trigonal bipyramidal.
- Writing SF₆ as trigonal bipyramidal; it is octahedral.
- Using sp³d/sp³d² as required Unit 5 hybridization content despite the syllabus specifying s and p orbitals only.
- Confusing σ and π overlap; σ is head-on, π is sidewise.
- Writing a double bond as two π bonds. Ordinary double bond = 1σ + 1π.
- Writing a triple bond as three π bonds. Triple bond = 1σ + 2π.
- Assuming a molecule is polar just because it contains polar bonds.
- Forgetting dipole cancellation in CO₂ and BF₃.
- Calling resonance contributors different molecules that rapidly interconvert.
- Confusing hydrogen bonding with an ordinary covalent H–X bond.
- Calling all intermolecular attractions hydrogen bonds.
- Confusing metallic bonding with ionic bonding; metallic electrons are delocalized across the metal.
25. Important Exam Questions
Very Short / Short Questions
- Define valence shell and valence electrons.
- State the octet rule and two exceptions.
- Define ionic bond and explain formation of NaCl.
- State four properties of ionic compounds.
- Define covalent bond and distinguish single, double and triple bonds.
- Define coordinate covalent bond with an example.
- State properties of molecular covalent compounds.
- Draw Lewis structures of H₂O, NH₃, CO₂, BF₃ and other common s/p-block compounds.
- Define resonance and explain it using a suitable species.
- State VSEPR theory.
- Write the order of electron-pair repulsions.
- Predict shapes of BeF₂, BF₃, CH₄, CH₃Cl, PCl₅, SF₆, H₂O, NH₃, CO₂, H₂S and PH₃.
- Define valence bond theory.
- Differentiate σ and π bonds.
- Define hybridization.
- Explain sp, sp² and sp³ hybridization.
- Define bond length.
- Explain ionic character of a covalent bond.
- Define dipole moment.
- Why is CO₂ non-polar but H₂O polar?
- What are van der Waals forces?
- Define hydrogen bonding and distinguish inter- and intramolecular H bonding.
- Define metallic bonding and explain conductivity/malleability of metals.
Long Questions
- Explain ionic bonding and correlate it with properties of ionic solids.
- Explain covalent and coordinate covalent bonding with Lewis structures.
- Explain Lewis structures and resonance with suitable examples.
- State VSEPR theory and predict the syllabus-listed molecular shapes.
- Explain valence bond theory and distinguish σ and π bonds.
- Explain sp, sp² and sp³ hybridization with geometry and examples.
- Discuss bond length, ionic character and dipole moment.
- Explain intermolecular forces, molecular solids and hydrogen bonding.
- Explain metallic bonding and the properties of metallic solids.
Practice / Application Questions
- Count σ and π bonds in ethene, ethyne and simple multiple-bond molecules.
- Predict geometry from the number of electron domains and lone pairs.
- Determine molecular polarity from bond polarity and molecular shape.
- Choose an appropriate sp/sp²/sp³ hybridization model for simple molecules.
- Compare boiling-point trends using hydrogen bonding and van der Waals forces.
Exam Strategy
Study Unit 5 in six blocks: octet/Lewis → ionic/covalent/coordinate bonds → resonance → VSEPR shapes → VBT + σ/π + hybridization → bond polarity/intermolecular/metallic bonding. The VSEPR table and σ/π/hybridization concepts are especially high-yield.
Study Unit 5 in six blocks: octet/Lewis → ionic/covalent/coordinate bonds → resonance → VSEPR shapes → VBT + σ/π + hybridization → bond polarity/intermolecular/metallic bonding. The VSEPR table and σ/π/hybridization concepts are especially high-yield.
26. One-Minute Revision
- Unit 5: Chemical Bonding and Shapes of Molecules — 9 teaching hours.
- Valence electrons participate most directly in ordinary bonding.
- Octet rule is useful but has exceptions.
- Ionic bond = electrostatic attraction between oppositely charged ions.
- Covalent bond = shared electron pair(s).
- Coordinate bond = shared pair donated initially by one atom/species.
- Lewis structures show valence electrons and bonds.
- Resonance contributors represent one delocalized structure.
- VSEPR minimizes repulsions between electron domains.
- LP–LP > LP–BP > BP–BP repulsion.
- BeF₂ linear; BF₃ trigonal planar; CH₄ tetrahedral.
- PCl₅ trigonal bipyramidal; SF₆ octahedral.
- NH₃ trigonal pyramidal; H₂O bent.
- CO₂ linear.
- H₂S bent; PH₃ trigonal pyramidal.
- σ bond = head-on overlap.
- π bond = sidewise overlap.
- Single = 1σ; double = 1σ+1π; triple = 1σ+2π.
- sp = linear 180°.
- sp² = trigonal planar 120°.
- sp³ = tetrahedral electron-domain geometry 109.5°.
- Higher bond order generally means shorter bond length for the same atoms.
- Electronegativity difference produces bond polarity/ionic character.
- Dipole moment depends on both bond polarity and molecular geometry.
- CO₂ and BF₃ have dipole cancellation by symmetry.
- H₂O and NH₃ are polar.
- London dispersion forces occur in all atoms/molecules.
- Hydrogen bonding is important in water and biomolecules.
- Metallic bonding involves positive ion cores and delocalized electrons.
27. Diagram Practice
- Potential-energy view of bond formation.
- NaCl ionic-bond formation.
- Lewis structures of H₂O, NH₃ and CO₂.
- Resonance contributors of carbonate.
- VSEPR geometry chart.
- σ and π orbital overlap.
- sp, sp² and sp³ hybridization geometries.
- Polar H–Cl bond.
- Dipole cancellation in CO₂ vs addition in H₂O.
- Hydrogen bonding in water.
- Metallic bonding electron-sea model.
Source handling: The Nepal eNotes source page was identified as the Class 11 Chemistry Chemical Bonding & Shape of Molecules page, but it timed out during direct retrieval in this session. A matching chapter mirror exposed the embedded Google Drive file used above. Therefore the PDF is presented as a matching handwritten-note mirror and is not claimed to be independently verified as the exact Nepal eNotes Drive file. The typed section follows the current Grade 11 Unit 5 syllabus and is a searchable, responsive study companion rather than a word-for-word transcription.
Discussion
Share a helpful question, idea, or explanation with other students.