Molecules and Molecular Structure: Chemical Bonding 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: The interaction that leads to the formation of molecules and compounds.

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

Types of Chemical Bonding

There are three main 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 of electrons from one atom to another, resulting in the formation of cations (positive ions) and anions (negative ions). The process follows the octet rule, where elements react to achieve eight electrons in their valence shell.

• Metals: Lose electrons to form cations.

• Non-metals: Gain electrons to form anions.

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

Bonding in Giant Ionic Lattices

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

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

• Ionic Crystals: Example is the arrangement of NaCl in a giant lattice.

Strength of Ionic Bonds

• High energy is required to separate ions in a lattice due to strong attractive forces.

• Results in high melting and boiling temperatures.

• Lattice energy is affected by:

  • 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	Lattice Energy (kJ/mol)
Na+ & F-	~910
Ca2+ & O2-	~3414
Na+ & Cl-	~787
Na+ & Br-	~751
Na+ & I-	~705

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, but insoluble in non-polar solvents.

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 in metals, molecular orbitals are formed, resulting in energy bands. The band theory explains the electrical properties of metals, insulators, and semiconductors.

• 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	Silicon, Germanium

Covalent Bonding and Coordinate Bonding

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.

• Attraction between electrons and nuclei balances repulsive forces, resulting in a specific bond length.

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

Lewis Electron-Dot Structures

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

• Example: CH3Cl (methyl chloride) and NF3 (nitrogen trifluoride).

• Steps to draw Lewis structures:

  1. Count total valence electrons.

  2. Place the least electronegative atom at the center.

  3. Distribute electrons to form bonds and satisfy the octet rule.

  4. Assign remaining electrons as lone pairs.

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+), Al2Cl6 dimer, hydrated metal ions.

Resonance Structures

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

• Rules for resonance:

  • Same placement of atoms, different locations of electrons.

  • 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 smallest 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 electronegativity between atoms.

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

*Additional info: These notes cover foundational concepts in 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 formulas are provided for clarity and comparison.*