Molecules and Molecular Structure: Chemical Bonding and Properties

Introduction to Chemical Bonding

Overview of Chemical Bonding

Chemical bonding refers to the forces that hold atoms together in molecules or compounds. The orientation and type of bonds determine the shape and properties of molecules. Atoms combine by losing, gaining, or sharing electrons to achieve a stable electron configuration, often an octet (eight electrons in the valence shell).

• Chemical bonds are essential for the formation of all substances.

• The octet rule states that atoms tend to react to achieve eight electrons in their valence shell.

Types of Chemical Bonding

• Ionic Bonding

• Covalent Bonding

• Metallic Bonding

Ionic Bonding

Formation of Ionic Bonds

Ionic (or electrovalent) bonding is the electrostatic attraction between oppositely charged ions. It involves the transfer of electrons from one atom (typically a metal) to another (typically a non-metal) to achieve an octet configuration.

• Metals lose electrons to form cations (positively charged ions).

• Non-metals gain electrons to form anions (negatively charged ions).

• Example: Formation of MgCl2 involves Mg losing two electrons and each Cl gaining one electron.

Bonding in Giant Ionic Lattices

A giant ionic lattice is a three-dimensional structure consisting of a repeating pattern of ions. The arrangement maximizes the attractive forces between oppositely charged ions, resulting in a stable structure.

• Lattice energy is the energy released when gaseous ions combine to form an ionic solid.

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

• Example: NaCl forms a giant lattice of alternating Na+ and Cl- ions, creating ionic crystals.

Strength of Ionic Bonds and Lattice Energy

• Ionic lattices require high energy to separate ions due to strong attractive forces.

• This results in high melting and boiling points.

• Lattice energy depends on:

  • Ionic charge: Higher charge increases attraction and lattice energy.

  • Ionic size: Smaller ions increase attraction and lattice energy.

Example Table: Effect of Ionic Charge and Size on Lattice Energy

Ion Pair	Ionic Radius (pm)	Lattice Energy (kJ/mol)
Na+ & F-	231	-910
Ca2+ & O2-	239	-3414

As ionic charge increases, lattice energy increases. As ionic size increases, lattice energy decreases.

Properties of Ionic Compounds

• Require high energy to break strong ionic bonds.

• Exist as crystalline solids at room temperature.

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

• Have high melting and boiling points.

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

Metallic Bonding

Electron-Sea Model

Metallic bonding is the force of attraction between positive metal ions and delocalized 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'.

• These delocalized electrons allow metals to conduct electricity and heat.

Properties of Metals

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

• Malleable and ductile (can be shaped without breaking).

• Good conductors of heat and electricity in both solid and liquid states.

Band Theory

When atomic orbitals overlap in metals, they form molecular orbitals that are very closely spaced, creating energy bands. The band theory explains the electrical properties of metals, insulators, and semiconductors.

• Valence band: Lower energy, contains valence electrons.

• Conduction band: Higher energy, contains empty molecular orbitals.

Classification by Band Theory:

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 Bonding and Coordinate Bonding

Formation of Covalent Bonds

Covalent bonds are formed when two atoms share electrons. Covalent compounds are typically formed between non-metal atoms. The shared electrons allow each atom to achieve a stable electron configuration.

• Bonding involves attraction between electrons and nuclei, and repulsion between like charges.

• The bond length is the distance at which these forces balance.

Bonding and Lone Pairs

• Bonding pairs are shared between atoms.

• Lone pairs are not shared and belong to a single atom.

• Example: In HF, there is 1 bonding pair and 3 lone pairs on F.

Lewis Electron-Dot Structures

Lewis structures represent valence electrons as dots around atomic symbols. They help visualize bonding and lone pairs.

• Example: CH3Cl and NF3 structures show how electrons are distributed.

• Follow the octet rule when drawing structures.

Exceptions to the Octet Rule

Type	Description	Examples
Expanded	More than 8 electrons (3rd period and beyond)	SF6, PCl5, SO42-
Incomplete	Less than 8 electrons (electron-deficient)	AlCl3, BF3, BeCl2
Odd	Odd number of electrons (free radicals)	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 Bonds

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

• Example: Formation of NH4+ (ammonium ion) and Al2Cl6 (dimer of AlCl3).

• Hydrated metal ions (e.g., [Al(H2O)6]3+) also involve coordinate bonds.

Resonance Structures

Some molecules can be represented by more than one valid Lewis structure, differing only in the placement of π electrons. These are called resonance structures.

• Rules: Atoms' positions remain the same; only electrons move.

• Examples: Ozone (O3), nitrate ion (NO3-).

Formal Charge

Formal charge helps determine the most stable Lewis structure. It is calculated as:

$\text{Formal charge} = \text{No. of valence electrons in free atom} - \frac{1}{2}(\text{No. of bonding electrons}) - \text{No. of lone pair electrons}$

• The most stable structure has zero or the smallest possible formal charges.

• Negative formal charges should be on the most electronegative atoms.

• Example: For nitrosyl chloride, the structure with all formal charges zero is most stable.

Covalent to Ionic - A Continuum

Bonding is often not purely ionic or covalent but exists on a continuum. The difference in electronegativity between atoms determines the degree of ionic or covalent character.

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

As the difference in electronegativity increases, the bond becomes more ionic in character.