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Chemical Bonding and Molecular Structure

Build molecular structure from valence electrons, then connect shape, bonding, polarity, bond order and intermolecular forces to observable properties.

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Build the idea step by step

Move through the core concepts in order, then compare exceptions and common traps.

Core concept

Lewis Symbols, Octet Rule and Bond Formation

Valence electrons provide the accounting language for ionic and covalent bonding.

Lewis symbols

A Lewis symbol shows only valence electrons around an element symbol. Main-group valence-electron counts follow the group pattern, with helium treated as a filled duet.

Place one electron on each side before pairing; the symbol represents an atom, not a molecule.

Ionic and covalent ideas

Ionic bonding is dominated by electrostatic attraction after electron transfer; covalent bonding involves shared electron pairs between atoms.

Bond type is a model of electron distribution, not a claim that real bonds are perfectly ionic or perfectly covalent.

Octet rule and limits

Many main-group atoms form bonds to approach noble-gas configurations. Incomplete octets, odd-electron species and expanded valence shells are recognised exceptions.

Apply the octet rule as a useful guide, then check authorised exceptions.

Worked connection

How many valence electrons must be counted for CO₂?

  1. Carbon contributes 4.
  2. Each oxygen contributes 6, giving 12 from oxygen.
  3. Total=4+12=16 valence electrons.

Answer: CO₂ has 16 valence electrons.

Common Trap: Do not count all electrons in the atoms; Lewis structures use valence electrons.

Core concept

Ionic Bonding and Lattice Enthalpy

Ion formation and lattice stabilisation together determine whether an ionic arrangement is favourable.

Formation factors

Low ionisation enthalpy favours cation formation, while favourable electron gain and strong lattice stabilisation support an ionic solid.

Do not judge ionic-compound formation from a single isolated energy step.

Lattice enthalpy

Lattice enthalpy reflects electrostatic interactions throughout an ionic crystal. Its magnitude increases with greater ionic charge and generally with smaller ion size.

Electrostatic attraction scales broadly with the product of charges and inversely with interionic distance.

Polarisation

A small highly charged cation can distort a large anion, increasing covalent character. Charge density and anion polarisability matter.

Greater polarisation means greater covalent character in the bond.

Worked connection

Which should have the larger lattice-enthalpy magnitude, NaCl or MgO?

  1. Compare charge products: 1×1 for NaCl and 2×2 for MgO.
  2. Mg²⁺ and O²⁻ are also relatively compact.
  3. Stronger electrostatic attraction gives MgO the larger magnitude.

Answer: MgO has the larger lattice-enthalpy magnitude.

Common Trap: Compare both ionic charge and size; charge alone is not the complete electrostatic picture.

Core concept

Bond Parameters, Polarity and Resonance

Bond length, strength, order, polarity and resonance describe different aspects of bonding.

Bond length and enthalpy

Bond length is the equilibrium internuclear distance. Stronger bonds are often shorter, but comparisons should involve related atoms and bond types.

Within a comparable pair, increasing bond order generally shortens and strengthens the bond.

Bond order

Lewis bond order counts shared pairs between two atoms; molecular-orbital bond order uses half the difference between bonding and antibonding electrons.

State which model supplies the bond order.

Polarity and dipole moment

Bond polarity comes from electronegativity difference; molecular dipole moment is the vector sum of all bond moments and depends on shape.

Polar bonds can cancel in a symmetrical molecule.

Worked connection

Why is CO₂ non-polar although each C=O bond is polar?

  1. CO₂ is linear.
  2. The two equal C=O bond moments point in opposite directions.
  3. Their vector sum is zero.

Answer: Symmetrical cancellation makes CO₂ non-polar.

Common Trap: Do not decide molecular polarity from one bond without considering geometry.

Core concept

Lewis Structures and Formal Charge

A disciplined electron-accounting method prevents incomplete or impossible structures.

Construction method

Count valence electrons, choose a sensible skeleton, add single bonds, complete terminal octets, place remaining electrons centrally and form multiple bonds when required.

Every electron placed must be included in the original total.

Formal charge

Formal charge=valence electrons−non-bonding electrons−half the bonding electrons. Preferred contributors usually minimise charge separation and place negative charge on more electronegative atoms.

Formal charge is bookkeeping, not the measured partial charge.

Resonance

When more than one valid Lewis contributor has the same atom arrangement, the real species is a resonance hybrid with delocalised electrons.

Atoms do not move between resonance contributors; only electron placement changes.

Worked connection

Find the formal charge on N in NH₄⁺.

  1. N has 5 valence electrons.
  2. It has no lone-pair electrons and eight bonding electrons.
  3. FC=5−0−8/2=+1.

Answer: The nitrogen formal charge is +1.

Common Trap: Count half of all bonding electrons assigned to the atom, not the number of bonds twice.

Core concept

VSEPR Shapes

Electron domains arrange to minimise repulsion; molecular shape names only atom positions.

Electron domains

Each single, double or triple bond counts as one electron domain; each lone pair is another domain around the central atom.

Multiple bonds count as one domain in basic VSEPR, though their repulsion can be somewhat stronger.

Repulsion order

Lone pair-lone pair repulsion is greater than lone pair-bond pair, which is greater than bond pair-bond pair repulsion.

Lone pairs compress adjacent bond angles.

Geometry and shape

Electron-domain geometry includes lone pairs; molecular shape describes the arrangement of bonded atoms only.

For NH₃: tetrahedral electron geometry, trigonal-pyramidal molecular shape.

Worked connection

Predict the shape of H₂O.

  1. O has two O−H bonds and two lone pairs: four domains.
  2. Four domains give tetrahedral electron geometry.
  3. Ignoring lone-pair positions when naming atom arrangement gives a bent shape.

Answer: H₂O is bent, with an angle smaller than the tetrahedral value.

Common Trap: Do not call H₂O tetrahedral when the question asks for molecular shape.

Core concept

Valence Bond Theory and Hybridisation

Orbital overlap explains directional sigma and pi bonds; hybridisation organises central-atom geometry.

Sigma and pi bonds

Head-on overlap forms a sigma bond; sidewise overlap forms a pi bond. A double bond contains one sigma and one pi bond; a triple bond contains one sigma and two pi bonds.

Every bonded atom pair has at most one sigma bond.

Steric number

Steric number is the number of sigma-bond domains plus lone-pair domains on the central atom in the basic hybridisation model.

Steric numbers 2, 3, 4, 5 and 6 correspond to sp, sp², sp³, sp³d and sp³d².

Overlap and direction

Greater effective overlap generally produces a stronger bond. Hybrid orbitals form directional sigma frameworks, while unhybridised orbitals can form pi bonds.

Hybridisation is assigned to an atom in a structure, not to the whole molecule without qualification.

Worked connection

Determine carbon hybridisation in ethene, C₂H₄.

  1. Each carbon forms three sigma domains: two C−H and one C−C.
  2. Steric number 3 gives sp² hybridisation.
  3. One unhybridised p orbital on each carbon forms the pi bond.

Answer: Each carbon in ethene is sp² hybridised.

Common Trap: A double bond counts as one electron domain, not two, for steric number.

Core concept

Molecular Orbital Theory

Atomic orbitals combine into delocalised molecular orbitals whose occupancy predicts stability and magnetism.

Combination conditions

Atomic orbitals combine effectively when their energies are comparable, their symmetry permits overlap and their spatial overlap is appreciable.

The number of molecular orbitals formed equals the number of atomic orbitals combined.

Bond order

MO bond order=½(Nb−Na), where Nb and Na are bonding and antibonding electron counts.

Positive bond order supports a bonded species; larger bond order generally means a stronger, shorter bond.

Magnetism

A species is paramagnetic if any molecular-orbital electron remains unpaired and diamagnetic if all are paired.

Magnetic character comes from occupancy, not merely total electron count.

Worked connection

Why is O₂ paramagnetic?

  1. Fill the O₂ molecular orbitals in the accepted order.
  2. Two electrons occupy separate degenerate π*2p orbitals.
  3. These unpaired electrons produce paramagnetism.

Answer: O₂ is paramagnetic because it has two unpaired antibonding electrons.

Common Trap: A Lewis structure with paired dots does not replace the MO prediction of O₂ magnetism.

Core concept

Hydrogen Bonding and Molecular Properties

Strong directional attractions involving H−F, H−O or H−N groups influence structure and physical properties.

Formation

Hydrogen bonding occurs when H bonded to a small strongly electronegative atom interacts with a lone pair on F, O or N in the accepted scope.

A molecule containing hydrogen does not automatically form hydrogen bonds.

Intermolecular and intramolecular

Intermolecular hydrogen bonding joins different molecules; intramolecular hydrogen bonding occurs within a suitably arranged molecule.

Intramolecular bonding can reduce the ability to associate with neighbouring molecules.

Property effects

Intermolecular hydrogen bonding can raise boiling point, viscosity and association. The exact outcome also depends on molecular size and shape.

Compare structurally related substances before assigning one cause to a physical property.

Worked connection

Why does H₂O have a much higher boiling point than H₂S?

  1. O is smaller and more electronegative than S.
  2. Water molecules form an extensive intermolecular hydrogen-bond network.
  3. More energy is required to separate the molecules.

Answer: Strong intermolecular hydrogen bonding raises water’s boiling point.

Common Trap: Do not describe the intermolecular hydrogen bond itself as an O−H covalent bond.

Active learning

Test the trends yourself

Change an element and compare the cause, trend and important exception.

Chemical Bonding & MO Laboratory

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What you will master

  • Count valence electrons and construct complete Lewis structures.
  • Use formal charge, resonance and octet exceptions without treating resonance forms as separate molecules.
  • Predict molecular shape and bond angle from electron-domain geometry.
  • Connect hybridisation with steric number and sigma-pi bonding.
  • Use molecular-orbital occupancy to calculate bond order and magnetic character.
  • Relate polarity and hydrogen bonding to molecular properties.
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