❓People Also Ask — Chemical Bonding
Questions students most frequently search about this chapter. Answers are direct and exam-focused.
QWhat is the difference between electron geometry and molecular geometry?
Electron geometry considers all electron pairs (bonding + lone) around the central atom. Molecular geometry describes only the positions of the atoms (bonding pairs), ignoring lone pairs. E.g., H₂O has tetrahedral electron geometry but bent molecular geometry.
QWhy does the bond angle in H₂S (92°) differ so much from H₂O (104.5°)?
Sulfur is larger and less electronegative, so it uses nearly pure p-orbitals for bonding with hydrogen (Drago's Rule). Pure p-orbitals are at 90° to each other — hence the ~92° angle. Oxygen, being smaller and more electronegative, uses sp³ hybridization, giving ~104.5°.
QWhat is the hybridization of the central atom in PCl₅?
sp³d. Phosphorus has 5 bond pairs and 0 lone pairs → steric number 5 → trigonal bipyramidal geometry with 90° and 120° angles.
QWhy is SF₆ non-polar despite having polar S-F bonds?
SF₆ has perfect octahedral symmetry. The six identical S-F bond dipoles cancel each other exactly, giving a net dipole of zero → non-polar molecule.
QWhat is the difference between polar covalent and ionic bonds?
Polar covalent: electrons are shared unequally between atoms (0.1 < ΔEN < 1.7). Ionic: electrons are transferred from one atom to another (ΔEN > 1.7). Both are extreme ends of a continuum, not discrete categories.
QWhat is the significance of the ΔEN = 1.7 threshold?
According to Pauling, when electronegativity difference reaches 1.7, the bond is about 50% ionic. Beyond 1.7, ionic character predominates; below 1.7, covalent character predominates.
QHow is percentage ionic character calculated from dipole moment?
% ionic character = (μ_observed / μ_100% ionic) × 100, where μ_100% ionic = q × d (full electron charge × bond length). For example, HCl has ~17% ionic character.
QWhy does NF₃ have a much smaller dipole moment than NH₃?
In NH₃, N-H bond dipoles point toward N (up) and the lone pair points up — vectors add. In NF₃, N-F bond dipoles point toward F (down), partially cancelling the upward lone pair. Net result: NF₃ has only 0.23 D vs NH₃'s 1.47 D.
QWhy is the F₂ bond energy anomalously low?
F atoms are extremely small, causing strong lone pair–lone pair repulsion between adjacent F atoms in F₂. This repulsion weakens the F-F sigma bond, making its bond energy lower than expected (158 kJ/mol).
QWhat is the difference between sigma and pi bonds in terms of overlap?
Sigma (σ): end-on (axial) overlap, giving maximum electron density along internuclear axis. Pi (π): sideways (lateral) overlap of parallel p-orbitals, with electron density above and below the axis.
QWhat is the hybridization of carbon in diamond?
Diamond is a giant covalent solid where each carbon is sp³ hybridized and bonded tetrahedrally to 4 other carbons. This gives diamond its extreme hardness.
QHow many sigma and pi bonds are in a C=C double bond?
One sigma (σ) + one pi (π). The sigma bond is the direct end-on overlap; the pi bond is the sideways overlap of unhybridized p-orbitals. Double bonds restrict rotation.
QWhat is the hybridization of nitrogen in the nitrate ion (NO₃⁻)?
sp². N forms 3 sigma bonds (one to each oxygen) and has 0 lone pairs. The double bond counts as one sigma domain only. Trigonal planar geometry with 120° angles.
QHow does bond order relate to bond length?
Higher bond order → shorter bond length. More shared electrons pull nuclei closer. C≡C (triple, ~1.20 Å) < C=C (double, ~1.34 Å) < C–C (single, ~1.54 Å).
QWhat is Drago's rule?
Drago's Rule states that for central atoms in Period 3 or below (S, P, Se, etc.) bonded to low-electronegativity atoms (H, F), hybridization is minimal — the atoms use nearly pure p-orbitals. Hence H₂S angle (~92°), not tetrahedral 109.5°.
QWhy are sigma bonds stronger than pi bonds?
Sigma bonds have head-on overlap, putting maximum electron density directly between nuclei. Pi bonds have sideways overlap, with less efficient overlap and less electron density between nuclei. Thus σ > π in bond strength.
QWhat is the shape of XeF₄ and its hybridization?
Square planar. Xenon uses sp³d² hybridization with 4 bond pairs and 2 lone pairs. The lone pairs occupy axial positions (to minimize 90° repulsions), leaving F atoms at corners of a square.
QWhich molecule has the highest dipole moment among CH₃F, CH₃Cl, CH₃Br, CH₃I?
CH₃Cl (1.87 D). Although F is most electronegative, the C-F bond is very short, reducing the μ = q × d product. C-Cl has both substantial polarity and longer bond, giving the maximum dipole.
QWhat is the significance of bond energy in thermochemistry?
Bond energies allow calculation of enthalpy changes: ΔH = Σ(BE bonds broken) – Σ(BE bonds formed). This is used in estimating reaction enthalpies when calorimetric data isn't available.
QWhich is stronger: C≡C triple or C=C double bond?
C≡C triple bond is stronger (839 kJ/mol vs 614 kJ/mol). More shared electrons (1σ + 2π vs 1σ + 1π) increase bond strength, shorten bond length, and reduce rotation even further.
📋Syllabus & Learning Outcomes
According to PMDC/PMC MDCAT syllabus, Chemical Bonding covers:
- 1 Use VSEPR Theory to describe the shapes of molecules.
- 2 Describe the features of sigma and pi-bonds.
- 3 Describe the shapes of simple molecules using orbital hybridization.
- 4 Determine the shapes of some molecules from the number of bonded pairs.
- 5 Predict the molecular polarity from the shapes of molecules.
- 6 Explain what is meant by the term ionic character of a covalent bond.
- 7 Describe how knowledge of molecular polarity can be used to explain some physical and chemical properties of molecules.
- 8 Define bond energies and explain how they can be used to compare bond strength of different chemical bonds.
📝Full Notes — Chemical Bonding
🔗 Chemical Bonding – Schematic Notes
📐 1. VSEPR Theory & Molecular Shapes
- VSEPR: Electron pairs (bonding + lone) around central atom repel each other; they arrange to minimise repulsion.
- Lone pair–lone pair > lone pair–bond pair > bond pair–bond pair (repulsion order).
| Total pairs | Bond pairs | Lone pairs | Shape | Examples |
| 2 | 2 | 0 | Linear (180°) | BeCl₂ |
| 3 | 3 | 0 | Trigonal planar (120°) | BF₃ |
| 3 | 2 | 1 | Bent (~117°) | SO₂ |
| 4 | 4 | 0 | Tetrahedral (109.5°) | CH₄ |
| 4 | 3 | 1 | Trigonal pyramidal (~107°) | NH₃ |
| 4 | 2 | 2 | Bent (~104.5°) | H₂O |
| 5 | 5 | 0 | Trigonal bipyramidal | PF₅ |
| 6 | 6 | 0 | Octahedral | SF₆ |
🔗 2. Sigma (σ) and Pi (π) Bonds
- σ bond: Head‑on overlap of orbitals (s–s, s–p, p–p); cylindrical symmetry around bond axis; free rotation possible. Stronger than π bond.
- π bond: Sideways overlap of parallel p‑orbitals; electron density above and below plane of nuclei; restricted rotation (causes cis/trans isomerism).
- Single bond: 1 σ bond; Double bond: 1 σ + 1 π; Triple bond: 1 σ + 2 π.
🌀 3. Orbital Hybridization & Shapes
- sp: linear (180°), 2 electron domains, e.g., BeCl₂, C₂H₂.
- sp²: trigonal planar (120°), 3 domains, e.g., BF₃, C₂H₄.
- sp³: tetrahedral (109.5°), 4 domains, e.g., CH₄, NH₃ (pyramidal), H₂O (bent).
- sp³d: trigonal bipyramidal, 5 domains, e.g., PF₅.
- sp³d²: octahedral, 6 domains, e.g., SF₆.
- Hybridization explains electron‑pair geometry and approximate bond angles; lone pairs occupy hybrid orbitals.
📏 4. Shapes from Bonded Pairs (Quick Reference)
- 2 bonded pairs, 0 lone pairs → Linear.
- 3 bonded pairs, 0 lone pairs → Trigonal planar.
- 4 bonded pairs, 0 lone pairs → Tetrahedral.
- 5 bonded pairs, 0 lone pairs → Trigonal bipyramidal.
- 6 bonded pairs, 0 lone pairs → Octahedral.
- Lone pairs reduce symmetry and bond angles (e.g., NH₃, H₂O).
⚡ 5. Molecular Polarity from Shape
- Polar molecule: has net dipole moment (asymmetric charge distribution). Depends on both bond polarity (ΔEN) and molecular shape.
- Non‑polar molecule: bond dipoles cancel due to symmetric shape (e.g., CO₂ linear – opposite dipoles cancel; CCl₄ tetrahedral – identical bonds cancel).
- Examples: H₂O – bent, polar; NH₃ – pyramidal, polar; BF₃ – trigonal planar, non‑polar; CH₄ – tetrahedral, non‑polar if all substituents same.
🧲 6. Ionic Character of a Covalent Bond
- Due to electronegativity difference (ΔEN): the greater the ΔEN, the more ionic character.
- A bond is purely covalent if ΔEN = 0; partially ionic (polar covalent) if ΔEN > 0. Approx. > 1.7 ΔEN → predominantly ionic.
- % ionic character increases with ΔEN; estimated by dipole moment measurement.
- Even very polar molecules (e.g., HF) have some covalent character.
💧 7. Polarity & Physical/Chemical Properties
- Polar molecules have higher boiling/melting points (dipole–dipole, H‑bonding if applicable).
- Polar solvents dissolve polar/ionic solutes ("like dissolves like").
- Chemical reactivity: polar bonds are more reactive toward nucleophiles/electrophiles (e.g., C=O polar → nucleophilic attack).
- Non‑polar molecules are insoluble in water and have weaker intermolecular forces (London dispersion only).
🔥 8. Bond Energies
- Bond energy (bond dissociation energy): average energy required to break one mole of bonds in gaseous state (kJ/mol).
- Higher bond energy → stronger bond.
- Multiple bonds: C≡C > C=C > C–C in bond strength (triple bond strongest, shortest).
- Compare bond strength between different compounds using bond energies; weaker bonds break more easily.
- Used in thermochemistry: ΔH ≈ Σ(BE of bonds broken) – Σ(BE of bonds formed).
🔑 High‑Yield MCQs & Traps
- VSEPR: lone pairs compress bond angles; know common shapes (CH₄, NH₃, H₂O, BF₃, BeCl₂).
- Sigma bond: end‑on overlap, stronger, free rotation; Pi bond: sideways overlap, restricts rotation.
- Single = 1σ, double = 1σ+1π, triple = 1σ+2π.
- Hybridization and shape: sp → linear; sp² → trigonal planar; sp³ → tetrahedral; sp³d → TBP; sp³d² → octahedral.
- Polarity: symmetry cancels dipoles (CO₂, CCl₄, BF₃ non‑polar; H₂O, NH₃ polar).
- Ionic character increases with ΔEN; >1.7 mostly ionic.
- Bond energy: higher = stronger bond; C≡C (839 kJ/mol) > C=C (614) > C–C (347).
- ⚠️ Traps: "All tetrahedral molecules are polar" → False (CH₄ non‑polar). "Pi bond allows rotation" → False (restricts). "Lone pairs do not affect shape" → False (reduce bond angles). "Bond polarity alone determines molecular polarity" → False (shape matters).
💡 MDCAT focus: VSEPR shapes, hybridization, sigma vs pi bonds, molecular polarity, and bond energy trends.
❔Frequently Asked Questions
What is chemical bonding?
Chemical bonding is the force of attraction that holds atoms together in a molecule or compound. The main types are ionic, covalent, coordinate covalent, hydrogen, and metallic bonds.
What is VSEPR theory?
VSEPR (Valence Shell Electron Pair Repulsion) theory states that electron pairs (bonding and lone) around a central atom arrange themselves to be as far apart as possible to minimize repulsion. This determines molecular shape.
What is the order of electron pair repulsion?
Lone pair–lone pair > lone pair–bond pair > bond pair–bond pair. Lone pairs occupy more space, causing stronger repulsion and smaller bond angles.
What is the difference between sigma and pi bonds?
Sigma (σ) bonds form by head-on (axial) overlap, are stronger, and allow free rotation. Pi (π) bonds form by sideways (lateral) overlap of p orbitals, are weaker, and restrict rotation (causing cis-trans isomerism).
What is orbital hybridization?
Hybridization is the mixing of atomic orbitals of similar energy to form new hybrid orbitals of equal energy and shape, explaining molecular geometry and bond angles.
What are the types of hybridization?
sp (linear, 180°), sp² (trigonal planar, 120°), sp³ (tetrahedral, 109.5°), sp³d (trigonal bipyramidal), and sp³d² (octahedral).
Why is a molecule polar or non-polar?
Molecular polarity depends on bond polarity (electronegativity difference) and shape. Symmetric shapes (CO₂, BF₃, CCl₄) cancel bond dipoles → non-polar. Asymmetric shapes (H₂O, NH₃) do not → polar.
What is dipole moment?
Dipole moment is a vector quantity measuring the net polarity of a molecule. Unit: Debye (D). Zero dipole moment → non-polar molecule; non-zero → polar.
What is ionic character of a covalent bond?
Ionic character is the partial charge separation in a covalent bond due to electronegativity difference (ΔEN). Higher ΔEN → more ionic character. ΔEN > 1.7 → predominantly ionic.
What is bond energy?
Bond energy (bond dissociation energy) is the average energy required to break one mole of bonds in the gaseous state, expressed in kJ/mol. Higher bond energy = stronger bond.
How does bond order affect bond strength?
Higher bond order means more bonds, shorter bond length, and higher bond energy. Order: C≡C > C=C > C–C.
Why does NH₃ have a smaller bond angle than CH₄?
NH₃ has one lone pair, which repels the bond pairs more than bond pairs repel each other. This compresses the H-N-H angle from 109.5° (CH₄) to 107°.
Why is CO₂ non-polar while SO₂ is polar?
CO₂ is linear (sp hybridized) → bond dipoles cancel → non-polar. SO₂ is bent (sp² with one lone pair) → bond dipoles do not cancel → polar.
Which is stronger: sigma or pi bond?
Sigma bond is stronger due to greater (head-on) overlap. Pi bond is weaker due to less effective lateral overlap.
What is the hybridization of carbon in ethyne (C₂H₂)?
sp hybridization. Each carbon forms two sigma bonds and two pi bonds. Geometry is linear with 180° bond angle.
What is the hybridization of carbon in ethene (C₂H₄)?
sp² hybridization. Each carbon forms three sigma bonds and one pi bond. Geometry is trigonal planar with 120° bond angle.
Which molecule has the highest dipole moment?
HF has the highest dipole moment among hydrogen halides due to the greatest electronegativity difference between H and F.
What happens to bond length as bond order increases?
Bond length decreases as bond order increases (C≡C < C=C < C–C). More shared electrons pull nuclei closer.
How many MCQs come from Chemical Bonding in MDCAT?
Chemical Bonding typically contributes 3–5 MCQs out of 81 Chemistry MCQs in MDCAT.
Which topics from Chemical Bonding are most frequently asked in MDCAT?
VSEPR shapes (CH₄, NH₃, H₂O, BF₃, PCl₅, SF₆), hybridization types, sigma vs pi bonds, polarity of molecules (CO₂, H₂O, NH₃, BF₃), dipole moment trends, and bond energy comparisons.