Atomic Structure and Bonding

The study of atomic structure and bonding provides a fundamental understanding of how matter is composed and how substances interact. This guide covers the evolution of atomic theory, the arrangement of electrons, periodic trends, chemical bonding, and nuclear chemistry.

1. The Concept of Atoms, Molecules, and Ions

An atom is the smallest unit of an element that retains the chemical properties of that element. A molecule consists of two or more atoms chemically bonded together. An ion is an atom or group of atoms that has gained or lost electrons, resulting in a net electrical charge.

2. Historical Development of Atomic Theory

• John Dalton: Proposed the first modern atomic theory, suggesting atoms are indivisible spheres.
• J.J. Thompson: Discovered the electron using the cathode ray tube; proposed the 'Plum Pudding' model.
• Ernest Rutherford: Conducted the Gold Foil experiment; discovered the nucleus and proposed that atoms are mostly empty space.
• Robert Millikan: Determined the charge of an electron via the Oil Drop experiment.
• Henry Moseley: Determined the atomic number of elements using X-rays, leading to the modern periodic table arrangement.
• Niels Bohr: Proposed that electrons move in fixed orbits or shells around the nucleus.

3. Subatomic Particles and Isotopy

Atoms consist of Protons (positive), Neutrons (neutral), and Electrons (negative).
Atomic Number (Z): Number of protons in the nucleus.
Mass Number (A): Total number of protons and neutrons.
Isotopes: Atoms of the same element with the same atomic number but different mass numbers (e.g., Carbon-12 and Carbon-14).

4. Electron Configuration and Orbitals

Electrons are arranged in shells (K, L, M, N) and subshells (s, p, d, f).
s-orbital: Spherical shape, holds 2 electrons.
p-orbital: Dumbbell shape, consists of three orientations (px, py, pz), holds 6 electrons.
Aufbau Principle: Electrons fill lower energy orbitals first.
Pauli Exclusion Principle: An orbital can hold a maximum of 2 electrons with opposite spins.

5. The Periodic Table and Periodicity

Elements are arranged by increasing atomic number. Groups are vertical columns (families), and periods are horizontal rows.
Trends:

• Atomic Radius: Increases down a group, decreases across a period.
• Ionization Energy: Decreases down a group, increases across a period.
• Electronegativity: Ability of an atom to attract electrons; increases across a period, decreases down a group.

6. Chemical Bonding

• Electrovalent (Ionic): Transfer of electrons from a metal to a non-metal (e.g., NaCl).
• Covalent: Sharing of electrons between non-metals (e.g., Cl2, CH4).
• Coordinate (Dative): A type of covalent bond where both shared electrons come from one atom (e.g., [NH4]+, [Cu(NH3)4]2+).
• Metallic Bonding: Attraction between positive metal ions and a 'sea' of delocalized electrons.
• Hydrogen Bonding: A strong intermolecular force between Hydrogen and highly electronegative atoms (N, O, F).

7. Shapes of Simple Molecules

• Linear: H2, O2, Cl2, HCl, CO2 (Bond angle: 180°).
• Non-linear (Bent): H2O (Bond angle: 104.5°).
• Pyramidal: NH3 (Bond angle: 107°).
• Tetrahedral: CH4 (Bond angle: 109.5°).

8. Nuclear Chemistry

Radioactivity is the spontaneous decay of unstable nuclei. Natural radioactivity occurs in elements like Uranium. Artificial radioactivity is induced by bombarding stable nuclei with particles.

• Alpha (α): Helium nucleus, low penetrating power.
• Beta (β): High-speed electron, moderate penetrating power.
• Gamma (γ): Electromagnetic radiation, high penetrating power.

Half-life (t1/2): The time taken for half the atoms in a radioactive sample to decay.