Chemical Bonding & Molecular Structure

The attractive force that holds two or more atoms together in a molecule or compound.

To reach a lower-energy, more stable state — like a ball rolling downhill — by gaining a full outer shell like the nearest noble gas.

Only the outermost valence electrons; inner electrons stay locked near the nucleus.

Noble gases are unreactive because they already have a full outer shell (8 electrons; He has 2).

Bonding using valence electrons. Kössel stressed the ionic bond (electron transfer); Lewis introduced electron sharing (covalent bond).

The element symbol written with one dot per valence electron — Na has 1 dot, O has 6.

Small atoms like hydrogen need only 2 electrons, copying helium, instead of an octet.

In 1916, Kössel (Germany) and Lewis (USA) explained bonding using valence electrons, working independently.

Lewis pictured the atom as a positive kernel (nucleus + inner electrons) with valence electrons at the corners of a cube.

Octet rule: atoms gain, lose or share electrons to get 8 electrons in the outer shell.

Electron-dot (Lewis) structure of methane (CH₄): carbon shares four electron pairs with four hydrogens

Electrons are fully transferred from a metal to a non-metal, forming + and − ions that attract. Example: NaCl.

Two non-metals share one or more electron pairs, both counting the shared pair in their octet. Example: H–H.

A covalent bond in which both shared electrons come from the same atom — e.g. the ammonium ion NH₄⁺ and ozone O₃.

Ionic: electron transfer, metal + non-metal, gives ions, e.g. NaCl. Covalent: electron sharing, non-metal + non-metal, gives molecules, e.g. H₂, CH₄.

Example: NaCl forms as sodium gives its outer electron to chlorine.

Example: the hydrogen molecule H–H is the simplest covalent bond.

Examples of dative bonds: ammonium ion NH4+ and ozone O3.

Ionic bond

Electron transfer

Covalent bond

Electron sharing

Ionic bond

Metal + non-metal

Covalent bond

Non-metal + non-metal

Ionic bond

Ions (cation + anion)

Covalent bond

Molecules

Ionic bond

NaCl

Covalent bond

H2, CH4

Single = one shared pair, H–H. Double = two pairs, O=O, C=C. Triple = three pairs, N≡N, C≡C.

A bond pair is a shared pair holding two atoms together; a lone pair is an unshared pair sitting on one atom.

A single bond shares one electron pair, drawn as one dash (H–H).

A Lewis (dot) structure shows both bond pairs and lone pairs; replacing each shared pair with a dash gives a dash structure.

Counting bond pairs and lone pairs is the first step to predicting a molecule's shape.

Incomplete: BeCl₂, BF₃, LiCl — stable with fewer than 8. Expanded: PCl₅, SF₆, H₂SO₄ — third-period elements hold more than 8.

Odd-electron molecules like NO and NO₂.

Formal charge = valence electrons − lone-pair electrons − ½(bonding electrons); it is used to pick the best Lewis structure.

Incomplete octet: some atoms are stable with fewer than 8 electrons, e.g. BeCl2, BF3, LiCl.

Expanded octet: third-period elements can hold more than 8 electrons, e.g. PCl5, SF6, H2SO4.

The octet rule says nothing about molecular shape or relative stability.

The energy needed to break one mole of a bond. A higher value means a stronger bond.

H₂ = 1, O₂ = 2, N₂ = 3, CO = 3. As bond order rises the bond gets shorter and stronger.

N₂ has bond order 3 and a bond enthalpy of 946 kJ/mol — very hard to break.

Bond length is the equilibrium distance between nuclei of two bonded atoms; each atom's share is its covalent radius.

Bond order is the number of bonds between two atoms: H2 = 1, O2 = 2, N2 = 3, CO = 3.

Bond order

1

Bond enthalpy (kJ/mol)

435.8

Length (pm)

74

Bond order

1

Bond enthalpy (kJ/mol)

155

Length (pm)

144

Bond order

2

Bond enthalpy (kJ/mol)

498

Length (pm)

121

Bond order

3

Bond enthalpy (kJ/mol)

946

Length (pm)

110

When one Lewis structure cannot describe a molecule, several canonical forms are drawn (O₃, CO₃²⁻). The real molecule is a resonance hybrid, more stable than any single form — it does not rapidly switch.

Non-polar joins identical atoms with equal sharing (H₂, Cl₂). Polar joins different atoms — the more electronegative one pulls the shared pair (HCl).

CO₂ is linear, so the two dipoles cancel. H₂O is bent, so the dipoles add up.

μ = charge × distance, measured in Debye — it measures bond or molecular polarity.

104.5°. Ammonia's H–N–H angle is 107° — lone pairs push the bonds closer.

Electron pairs around the central atom repel and arrange themselves as far apart as possible.

lone pair–lone pair > lone pair–bond pair > bond pair–bond pair. Lone pairs therefore reduce the bond angle.

2 = linear, 3 = trigonal planar, 4 = tetrahedral, 5 = trigonal bipyramidal, 6 = octahedral.

SF₆ octahedral (6 bond pairs, no lone pairs) · CH₄ tetrahedral · NH₃ pyramidal · H₂O bent.

sp³d = trigonal bipyramidal (PCl₅); sp³d² = octahedral (SF₆).

Lone pairs

0

Shape

Linear

Example

BeCl2, CO2

Lone pairs

0

Shape

Trigonal planar

Example

BF3

Lone pairs

0

Shape

Tetrahedral

Example

CH4

Lone pairs

1

Shape

Pyramidal

Example

NH3

Lone pairs

2

Shape

Bent

Example

H2O

Lone pairs

0

Shape

Octahedral

Example

SF6

By head-on, end-to-end overlap along the bond axis. It is strong and present in every covalent bond — always the first bond between two atoms.

By sideways overlap of p-orbitals above and below the axis. It is weaker than a sigma bond.

Single = 1σ. Double = 1σ + 1π. Triple = 1σ + 2π.

A sigma bond allows free rotation; a pi bond restricts rotation around the bond.

Valence Bond Theory (VBT): a covalent bond forms when half-filled orbitals of two atoms overlap.

The mixing of atomic orbitals of similar energy to form equal hybrid orbitals.

sp: linear, 180°, BeCl₂/C₂H₂. sp²: trigonal planar, 120°, BF₃/C₂H₄. sp³: tetrahedral, 109.5°, CH₄/NH₃/H₂O.

Gives a shorter, stronger bond and a larger bond angle.

sp = 2 hybrids, linear, 180° (BeCl2, C2H2).

sp2 = 3 hybrids, trigonal planar, 120° (BF3, C2H4).

sp3 = 4 hybrids, tetrahedral, 109.5° (CH4, NH3, H2O).

sp3d = trigonal bipyramidal (PCl5); sp3d2 = octahedral (SF6).

Shape

Linear

Angle

180°

Example

BeCl2, C2H2

Shape

Trigonal planar

Angle

120°

Example

BF3, C2H4

Shape

Tetrahedral

Angle

109.5°

Example

CH4

Shape

Trigonal bipyramidal

Angle

90°,120°

Example

PCl5

Shape

Octahedral

Angle

90°

Example

SF6

Bond order = ½ (bonding electrons − antibonding electrons). Positive means the molecule exists; zero means it does not (He₂ = 0). CO gives ½(10 − 4) = 3.

Why O₂ is paramagnetic — it has two unpaired electrons.

MOT: atomic orbitals combine to form molecular orbitals belonging to the whole molecule.

Combining gives a low-energy bonding orbital and a high-energy antibonding orbital (marked *).

When H attached to a highly electronegative atom (F, O or N) is attracted to another such atom. Weaker than a covalent bond but stronger than van der Waals forces.

Water's high boiling point, ice being less dense than water, and it holds together the DNA double helix and protein structures.

Intermolecular hydrogen bonds are between molecules (water, HF); intramolecular are within one molecule (o-nitrophenol).