Molecules and Molecular Structure: Chemical Bonding (Unit 6) Study Notes

Molecules and Molecular Structure

Introduction to Chemical Bonding

Chemical bonding refers to the forces that hold atoms together in molecules or compounds. The orientation of these bonds determines the shape of a molecule. Atoms combine by losing, gaining, or sharing electrons to attain an octet configuration, which is a stable arrangement of eight electrons in the valence shell.

• Chemical Bonding: Forces that hold atoms together.

• Octet Rule: Atoms react to achieve eight electrons in their valence shell.

Types of Chemical Bonding

• Ionic Bonding

• Covalent Bonding

• Metallic Bonding

Ionic Bonding

Formation and Properties of Ionic Bonds

Ionic (Electrovalent) Bonding is the electrostatic attraction between oppositely charged ions. It involves the transfer (loss/gain) of electrons from one atom to another to attain an octet configuration.

• Transfer of Electrons: Metals lose electrons to form cations; non-metals gain electrons to form anions.

• Octet Rule: Elements react to achieve eight electrons in their valence shell.

Bonding in Simple Ionic Compounds

Example: Formation of magnesium chloride, MgCl2

• Mg atom loses two electrons to form Mg2+

• Each Cl atom gains one electron to form Cl-

Bonding in Giant Ionic Lattices

A giant ionic lattice is a compound consisting of endlessly repeating ions. The energy released during the formation of the lattice solid from gaseous ions is called lattice energy.

• Lattice Energy: $\text{Na}^+(g) + \text{Cl}^-(g) \rightarrow \text{NaCl}(s)$

• Ionic Crystals: Example: NaCl forms a giant lattice of ions.

Strength and Properties of Ionic Compounds

• High energy required to break strong ionic bonds.

• Exist as crystalline solids at room temperature.

• Conduct electricity in molten and aqueous states, but not in solid state.

• High melting and boiling points.

• Soluble in water, insoluble in non-polar solvents.

Factors Affecting Lattice Energy

• Ionic Charge: Higher charge leads to stronger attraction and larger lattice energy.

• Ionic Size: Larger ions result in weaker attraction and smaller lattice energy.

Ion Pair	Ionic Radius (pm)	Lattice Energy (kJ/mol)
Na+ & F-	231	-910
Ca2+ & O2-	239	-3414
Na+ & Cl-	---	-787
Na+ & Br-	---	-751
Na+ & I-	---	-705

Metallic Bonding

Electron-Sea Model

Metallic bonding is the force of attraction between positive metal ions and delocalised electrons. Metal atoms are closely packed in a giant lattice, and their outer electrons are free to move throughout the structure, forming a 'sea of electrons'.

• Delocalised electrons allow metals to conduct electricity and heat.

• Metals are malleable and ductile due to the mobility of electrons.

• High melting and boiling points due to strong metallic bonds.

Band Theory

When atomic orbitals overlap, molecular orbitals are formed. In metals, the number of molecular orbitals is very large, and the energy separations between them are extremely small, forming an energy band.

• Valence Band: Contains valence electrons (lower energy).

• Conduction Band: Contains empty molecular orbitals (higher energy).

Type	Band Structure	Conductivity	Example
Metal	No energy gap between bands	Good conductor	Copper
Insulator	Large energy gap	Poor conductor	Diamond
Semiconductor	Small energy gap	Moderate conductor (increases with temperature)	Silicon, Germanium

Covalent Bond and Coordinate Bond

Covalent Bonding

A covalent bond is formed when two electrons are shared between two atoms, typically non-metals. Covalent compounds can be formed from atoms of the same element or different non-metals.

• Bonding pairs and lone pairs of electrons affect molecular structure.

• Bond length is the distance at which attractive and repulsive forces between atoms balance.

Lewis Electron-Dot Symbols

Lewis structures represent the arrangement of electrons in molecules, showing bonding pairs and lone pairs.

• Example: CH3Cl (methyl chloride)

• Example: NF3 (nitrogen trifluoride) – stepwise construction using valence electrons and octet rule.

Exceptions to the Octet Rule

Type	Description	Examples
Expanded	Central atom has more than 8 electrons	SF6, PCl5, SO42-
Incomplete	Central atom has less than 8 electrons	AlCl3, BF3, BeCl2
Odd	Central atom has odd number of electrons	ClO, NO, NO2

Multiple Covalent Bonds

• Single bond: One electron pair shared (e.g., H2)

• Double bond: Two electron pairs shared (e.g., O2)

• Triple bond: Three electron pairs shared (e.g., N2)

Coordinate (Dative) Covalent Bond

A coordinate bond is a covalent bond in which both electrons come from the same atom.

• Example: Ammonium ion (NH4+)

• Example: Dimer formation in AlCl3 (Al2Cl6)

• Example: Hydrated metal ions (e.g., [Al(H2O)6]3+)

Resonance Structures

Some molecules can be represented by more than one reasonable Lewis structure, differing only in the location of π electrons.

• Rules: Same atom placement, different electron locations; only move electrons in π bonds or lone pairs; overall charge and bond length remain unchanged.

• Examples: Ozone (O3), Nitrate ion (NO3-)

Formal Charge

Formal charge helps determine the most likely Lewis structure.

• Formula: $\text{Formal charge} = \text{No. of valence e}^- \text{ in free atom} - \frac{1}{2} \text{No. of bonding e}^- - \text{No. of lone pair e}^-$

• The most stable structure has zero or minimal formal charges, with negative charges on the most electronegative atoms.

Covalent to Ionic - A Continuum

Bonding in compounds can be partially ionic and partially covalent, depending on the difference in nuclear strength and electronegativity between elements.

Compound	Bond Character
Cl2	Pure covalent (equal sharing)
HCl	Polar covalent (unequal sharing)

*Additional info: These notes cover the foundational concepts of chemical bonding, including ionic, covalent, and metallic bonds, as well as advanced topics such as resonance, formal charge, and the continuum between covalent and ionic character. Tables and diagrams have been recreated in text and HTML format for clarity and completeness.*