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CBSE Class 11 · Chemistry

Chemical Bonding and Molecular Structure

Official NCERT chapter from Chemistry Part I–II (book code kech1). ExamMaster notes are original teaching at CBSE Class 11 depth.

This lesson follows the official chapter “Chemical Bonding and Molecular Structure” in Chemistry Part I–II. The words below are ExamMaster’s teaching, not a paste from the book. Use the NCERT chapter for the classroom sequence; use these notes to hold the idea without copying exercises or figures.

  • CBSE Class 11
  • Medium level
  • 9 concepts

1Kössel-Lewis Approach to Chemical Bonding

Kössel–Lewis: atoms seek a taught noble-gas count by losing, gaining, or sharing. A Lewis-dot is the valence-pack drawn. The approach is a count-story, not a glow. Na· plus ·Cl → Na⁺ and Cl⁻ as the ionic end of that count.

Drawing eight dots on hydrogen “to be safe” fails the duet.

Figure. Kössel–Lewis: sodium starts 2,8,1 and chlorine 2,8,7. One electron leaves Na and joins Cl, so both ions hold a noble-gas count — Na⁺ is 2,8 and Cl⁻ is 2,8,8. The figure is a count transfer, not a drawn Lewis-dot cloud.

How it works

  1. Count valence electrons as taughtThe start.
  2. Lose, gain, or share toward the noble countThe approach.
  3. Keep H at two if the lesson used the duetHonest.

2Ionic or Electrovalent Bond

Ionic (electrovalent): a metal-to-nonmetal electron transfer as taught, then opposite-charge stick. NaCl is the school case. An ionic bond is a transfer-plus-stick, not “any two different elements”. A high melting look can support, not replace, the transfer-story.

Calling HCl ionic because two names appear is a steal.

Figure. An ionic bond is the electrostatic hold between ions after the electron has already moved. Na⁺ and Cl⁻ attract; the small four-box grid is a 2-by-2 slice of the NaCl pairing, not a full crystal drawing.

How it works

  1. Name the transfer (who loses, who gains)The electrovalent.
  2. Write the ions and the stickThe bond.
  3. Refuse a shared-pair compound as this headingIonic ≠ covalent.

3Bond Parameters

Bond parameters: length, angle, enthalpy, order as taught — numbers that describe one bond. A shorter bond is often a stronger one in the school trend, not a law for every pair. Parameters are measures, not a second Lewis-dot.

A pretty length with no unit is a mute parameter.

Figure. Bond length is a measured internuclear distance. Single, double and triple C–C bonds in ethane, ethene and ethyne sit near 154 pm, 134 pm and 120 pm. The bars use those lengths as values, so a taller bar is a longer bond, not a stronger one.

How it works

  1. Name length, angle, enthalpy, or orderThe parameter.
  2. Keep the unit (pm, kJ/mol, degrees)The measure.
  3. Read a trend only as taughtHonest.

4The Valence Shell Electron Pair Repulsion (VSEPR) Theory

VSEPR: valence pairs (bond + lone) repel and take a taught shape. Four pairs, no lone → tetrahedral (CH4); three bond + one lone → pyramidal (NH3 as taught). VSEPR is a pair-count, not a colour-model.

Ignoring lone pairs and calling NH3 tetrahedral is a miss.

Figure. VSEPR places electron pairs as far apart as a plane can show. BeCl₂ has two bonding pairs and stays in a line. H₂O has four pairs (two bonding, two lone); the lone-pair box sits opposite the hydrogens so the H–O–H path is bent. This is a pair-count sketch, not a 3-D tetrahedron.

How it works

  1. Count bond pairs and lone pairsThe input.
  2. Name the taught shapeVSEPR.
  3. Keep lone pairs in the countThey occupy space.

5Valence Bond Theory

Valence bond theory (school): a bond is an overlap of half-filled orbitals as framed. End-on overlap is sigma; side-on is pi if taught. VBT is an overlap-picture, not a replacement of Lewis for every exam line.

A no-overlap “stick” with no orbital names is not this theory.

Figure. Valence bond theory treats a covalent bond as two atomic regions that occupy the same space. The middle strip is the shared pair; the two outer boxes are the hydrogen 1s regions. The figure names overlap, not orbital lobes.

How it works

  1. Name the overlapping orbitalsThe VBT start.
  2. Say sigma or pi as taughtThe sort.
  3. Keep it an overlap, not a transferCovalent-side.

6Hybridisation

Hybridisation: s and p (and d if taught) mix to new equal-energy hybrids that match the shape. sp³ for four tetrahedral, sp² for trigonal, sp for linear — as named. Hybrid is a mix-word, not a breed of atom.

Calling every carbon sp³ because “organic” is a steal.

Figure. Hybridisation remixes one s and three p atomic boxes into four equal sp³ boxes. The count is 1 + 3 = 4 on both sides. The drawing is a count-and-equivalence statement, not a set of lobes.

How it works

  1. Count the regions around the atomThe cue.
  2. Name sp, sp², or sp³ as taughtThe hybrid.
  3. Match the shape you already namedThe use.

7Molecular Orbital Theory

MOT (school): electrons sit in molecule-wide orbitals, bonding and antibonding as taught. Bond order is (bonding − antibonding)/2. MOT is a second picture for the same pair, not a poster to memorise without a fill.

A bond order of 3 with no electron-count is a guess.

Figure. A molecular-orbital energy ladder for H₂. The lower rung is the bonding σ 1s level and holds both electrons (up and down). The upper rung is the empty antibonding σ* 1s. Atomic H 1s boxes sit at the foot. Bond order is (2 − 0)/2 = 1. Levels are rungs, not orbital shapes.

How it works

  1. Fill the taught MO ladderThe picture.
  2. Compute (Nb − Na*)/2Bond order.
  3. Keep a zero order as “no stable bond” if taughtHonest.

8Bonding in Some Homonuclear Diatomic Molecules

Homonuclear diatomic (school): H2, N2, O2, F2 as named — same-element pairs. N2’s triple and O2’s paramagnetism (if taught) are MOT-uses, not extra biographies. One fill for N2 or O2 is enough at this class.

Treating CO as homonuclear because both are “small” is a miss.

Figure. Bond order from the MO count (bonding minus antibonding electrons)/2. H₂ and F₂ are 1, O₂ is 2, N₂ is 3. The bar height is that order, so N₂ is the tallest. These are the usual homonuclear exam set, drawn as orders rather than as orbital cartoons.

How it works

  1. Stay on X2, same elementHomonuclear.
  2. Fill or quote the taught bond-order / magnetismThe case.
  3. Refuse a mixed-element pair hereThat is hetero.

9Hydrogen Bonding

Hydrogen bonding: H on N, O, or F (as taught) feels a second stick to another N, O, or F. Water’s high boiling is the school use. H-bond is a special extra, not “any H nearby”. Intra- versus inter- if the lesson split them.

Calling every hydrogen compound H-bonded is a steal.

Figure. A hydrogen bond is a dashed extra hold from a hydrogen already bonded to O (or N, F) toward a second O. Two water-like H–O–H units share that dashed H-bond. The solid edges are covalent O–H; the dashed edge is the weaker intermolecular link. Not a second covalent bond.

How it works

  1. Need H on N/O/F and a nearby N/O/FThe test.
  2. Name the extra stickHydrogen bond.
  3. Use it for a boiling or a DNA-mention only as taughtThis class.
VSEPR shape starts from
  1. Counting bond pairs plus lone pairs
  2. The atomic mass
  3. A glow

Pairs occupy space.

Notes

  • Mapped to the official NCERT chapter “Chemical Bonding and Molecular Structure”. Original teaching only — no textbook sentences.
  • Science here is Physics, Chemistry and Biology ideas at this class, never a language or social-science chapter.

Formulas

  • bond order = (Nb − Na*)/2
  • VSEPR: regions = bond pairs + lone pairs

Recap

Hold these pegs from the official chapter “Chemical Bonding and Molecular Structure”. The wording is ExamMaster’s teaching, not a textbook recap.

Kössel-Lewis Approach to Chemical Bonding
Kössel–Lewis: atoms seek a taught noble-gas count by losing, gaining, or sharing.
Ionic or Electrovalent Bond
Ionic (electrovalent): a metal-to-nonmetal electron transfer as taught, then opposite-charge stick.
Bond Parameters
Bond parameters: length, angle, enthalpy, order as taught — numbers that describe one bond.
The Valence Shell Electron Pair Repulsion (VSEPR) Theory
VSEPR: valence pairs (bond + lone) repel and take a taught shape.
Valence Bond Theory
Valence bond theory (school): a bond is an overlap of half-filled orbitals as framed.
Hybridisation
Hybridisation: s and p (and d if taught) mix to new equal-energy hybrids that match the shape.

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