Chemistry - Atomic Structure and Bonding

1. Atomic Structure

Subatomic Particles

• Atomic number ($Z$): number of protons; defines the element
• Mass number ($A$): protons + neutrons
• Isotopes: same $Z$, different $A$ (same element, different neutron count)

Electron Configuration

Electrons occupy shells. The maximum number of electrons in each shell follows the $2n^2$ rule:

Electrons fill in order of increasing energy: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p…

Examples:

• Sodium ($Z=11$): $1s^2\,2s^2\,2p^6\,3s^1$ or $[\mathrm{Ne}]\,3s^1$
• Iron ($Z=26$): $[\mathrm{Ar}]\,3d^6\,4s^2$
• Chromium ($Z=24$): $[\mathrm{Ar}]\,3d^5\,4s^1$ (exception for half-filled stability)

2. Periodic Trends

Across a Period (left to right)

Down a Group

Group Properties

• Group I (Alkali metals): $+1$ charge; highly reactive; react with water to form hydroxides and hydrogen
• Group VII (Halogens): $-1$ charge; reactivity decreases down group; displacement reactions
• Group 0 (Noble gases): full outer shell; chemically inert; boiling point increases down group

3. Ionic Bonding

Formation

Transfer of electrons from a metal atom (to form a cation) to a non-metal atom (to form an anion). The resulting electrostatic attraction between oppositely charged ions is the ionic bond.

Example: $\mathrm{Na} \to \mathrm{Na}^+ + e^-$ and $\mathrm{Cl} + e^- \to \mathrm{Cl}^-$

Dot-Cross Diagrams

Show the transfer of electrons evidently:

• Metal: dot and cross in outer shell → becomes ion with empty outer shell
• Non-metal: gains electron(s) to fill outer shell

Properties of Ionic Compounds

4. Covalent Bonding

Formation

Sharing of electron pairs between non-metal atoms. Each shared pair constitutes one covalent bond ($\sigma$ bond).

Dot-Cross Diagrams for Covalent Molecules

Show shared pairs (one dot from each atom, overlapping) and lone pairs:

• $\mathrm{H_2}$: single bond (1 shared pair)
• $\mathrm{O_2}$: double bond (2 shared pairs)
• $\mathrm{N_2}$: triple bond (3 shared pairs)
• $\mathrm{H_2O}$: bent molecule, 2 bonds, 2 lone pairs on O

Coordinate (Dative) Bonds

A covalent bond where both electrons come from the same atom.

Example: $\mathrm{NH_4^+}$ — the fourth N–H bond is dative (N donates both electrons to $\mathrm{H}^+$).

Simple vs Giant Covalent Structures

Diamond: tetrahedral; each C bonded to 4 others; insulator Graphite: layered; each C bonded to 3 others; delocalised electrons → conducts electricity; lubricant (weak intermolecular forces between layers)

5. Metallic Bonding

Formation

Positive metal ions in a lattice surrounded by a sea of delocalised electrons. The attraction between ions and the delocalised electrons is the metallic bond.

Properties of Metals

Alloys (e.g. steel, brass) are mixtures of metals with other elements. Different-sized atoms disrupt the regular lattice, making alloys harder than pure metals.

6. Intermolecular Forces

London (Van der Waals) Forces

• Present between all molecules
• Caused by instantaneous (temporary) dipoles due to electron movement
• Strength increases with: number of electrons (larger $A_r$), molecular size, surface area
• Explains trends in boiling points within homologous series (e.g. alkanes: $C_nH_{2n+2}$)

Dipole-Dipole Forces

• Between polar molecules (those with permanent dipoles)
• Stronger than London forces but weaker than hydrogen bonds
• Example: $\mathrm{HCl}$, $\mathrm{SO_2}$

Hydrogen Bonding

• Strongest intermolecular force
• Occurs when H is bonded to N, O, or F (highly electronegative atoms)
• Requires a lone pair on N, O, or F of a neighbouring molecule
• Examples: $\mathrm{H_2O}$, $\mathrm{NH_3}$, $\mathrm{HF}$, alcohols, carboxylic acids

Effects of hydrogen bonding:

• Unusually high boiling points (e.g. $\mathrm{H_2O}$ boils at 100 °C vs $\mathrm{H_2S}$ at $-60$ °C)
• Ice is less dense than water (open H-bonded lattice)
• Solubility of polar molecules in water

7. Physical Properties Linked to Structure

Key principle: The type of bonding determines the physical properties. Metallic and ionic bonding → high m.p. Giant covalent → very high m.p. Simple covalent → low m.p. (stronger intermolecular forces = higher m.p. within this group).

Predicting Properties

To predict the physical properties of a substance:

1. Identify the type of bonding present
2. Consider the strength of bonding/forces
3. Predict: melting/boiling point, electrical conductivity, solubility
4. Explain in terms of the bonding model

Worked Examples

Example 1: Predicting Physical Properties

Problem: Predict the melting point, electrical conductivity (solid and molten), and solubility in water for silicon dioxide ($\mathrm{SiO_2}$). Solution: $\mathrm{SiO_2}$ has a giant covalent structure. Each Si atom is bonded to 4 O atoms in a continuous network. Therefore: very high melting point (strong covalent bonds throughout), does not conduct electricity (no free electrons or ions), and is insoluble in water (giant covalent structures are not attracted by water molecules).

Example 2: Identifying Bond Type from Properties

Problem: A substance melts at 801 °C, conducts electricity when molten but not when solid, and dissolves in water. Identify the type of bonding present. Solution: High melting point rules out simple covalent. Conducts when molten but not solid is characteristic of ionic compounds (ions free to move in liquid state but locked in lattice when solid). Solubility in water confirms ionic bonding (polar water molecules attract and separate the ions).

Common Pitfalls

• Confusing ionic and metallic conductivity: Ionic compounds conduct only when molten or aqueous (mobile ions). Metals conduct in all states (delocalised electrons).
• Stating graphite is an insulator: Graphite has delocalised electrons between its layers and does conduct electricity, unlike diamond.
• Ignoring exceptions in electron configuration: Chromium ($[\mathrm{Ar}]\,3d^5\,4s^1$) and copper ($[\mathrm{Ar}]\,3d^{10}\,4s^1$) are exceptions to the standard filling order due to the stability of half-filled and fully filled d-subshells.

Summary

Atomic structure and bonding covers subatomic particles, electron configuration, periodic trends, and the four main types of bonding: ionic (electron transfer, giant lattice), covalent (electron sharing, simple and giant structures), metallic (delocalised electron sea), and intermolecular forces (London, dipole-dipole, hydrogen bonding). The type of bonding determines physical properties including melting point, conductivity, and solubility.