Chemical Bonding and Molecular Structure

Lesson

Part 1 — why atoms join, how to draw their bonding electrons, and what the measurable properties of a bond tell us.

In this lesson you will

Explain the octet rule, ionic bonding and lattice enthalpy qualitatively
Draw Lewis structures and calculate formal charges, including resonance and octet exceptions
Use bond length, bond angle, bond enthalpy and bond order to compare bonds
Explain bond polarity and dipole moment, including NH₃ and NF₃, H₂O and CO₂, and BF₃

Table of three kinds of chemical bond with how each forms and examples
Type of bond	How it forms	Examples
Ionic	Transfer of electrons	NaCl, MgO
Covalent	Sharing of electron pairs	H₂, O₂, N₂, CH₄
Polar covalent	Unequal sharing of a pair	HCl, H₂O

Concept

The Kössel–Lewis approach and the octet rule

Atoms combine to reach the stable configuration of a noble gas: eight electrons in the outer shell (the octet), or two for hydrogen (the duplet). A Lewis symbol shows the valence electrons as dots around the atomic symbol.

Table of the number of valence electrons in each main group with an example element
Group	1	2	13	14	15	16	17	18
Valence electrons	1	2	3	4	5	6	7	8 (He: 2)
Example	Na	Mg	Al	C	N	O	Cl	Ne

Concept

Ionic bonds and lattice enthalpy

An ionic bond forms when a metal loses and a non-metal gains electrons, and the ions attract. It is favoured by a low ionisation enthalpy of the metal, a negative electron gain enthalpy of the non-metal and a high lattice enthalpy. Lattice enthalpy rises with higher ionic charge and smaller ionic size.

Table of approximate lattice enthalpies of four ionic compounds rising with smaller ions and higher charge
Compound	Ions	Lattice enthalpy / kJ mol⁻¹ (approx.)
KCl	K⁺, Cl⁻	717
NaCl	Na⁺, Cl⁻	788
NaF	Na⁺, F⁻	923
MgO	Mg²⁺, O²⁻	3850

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