Atomic Theory and Chemical Bonding: Exploring Subatomic Particles

Posted on Nov 15, 2024 in Chemistry

Atomic Theory and Chemical Bonding

Early Atomic Experiments

In the late 19th and early 20th centuries, experiments involving electrical discharges through low-pressure gases and particle bombardment of radioactive gases and thin metal sheets revealed the divisibility of atoms and the existence of subatomic particles.

Key Discoveries:

  • 1897: J.J. Thomson discovered the electron.
  • 1911: Ernest Rutherford identified the atomic nucleus.
  • 1932: James Chadwick discovered the neutron.

Atomic Models

Thomson’s Model:

The atom is a sphere of positive charge with embedded negative electrons, neutralizing the overall charge.

Rutherford’s Model:

  • A small, dense, positive nucleus contains most of the atom’s mass.
  • Negative electrons orbit the nucleus.
  • Significant empty space exists within the atom.
  • The nucleus is about 10,000 times smaller than the atom’s volume.
  • The atom is electrically neutral.

Key Atomic Concepts:

  • Atomic Number: Number of protons.
  • Atomic Mass Number: Number of protons plus number of neutrons.
  • Isotopes: Atoms with the same atomic number but different mass numbers.
  • Isobars: Atoms with the same mass number but different atomic numbers.
  • Atomic Mass: Weighted average of the masses of an element’s isotopes.

Planck’s Quantum Theory:

Energy changes in atoms occur in discrete multiples of a quantum of energy: E = hν, where h = 6.626 x 10-34 Js.

Bohr’s Model:

  • Electrons orbit the nucleus in specific energy levels without emitting energy.
  • Electron orbits have quantized radii.
  • Atoms absorb or emit energy when electrons transition between orbits.

Vector Model of the Atom:

Four quantum numbers (n, l, m, s) determine an electron’s energy and orbital characteristics.

Electron Configuration

  • Energy Levels (n): 1, 2, 3, …
  • Sublevels (l): 0, 1, 2, …, n-1
  • Orbitals (m): -l, …, 0, …, +l
  • Spin (s): +1/2, -1/2

Hund’s Rule:

Electrons fill orbitals individually before pairing up.

Madelung’s Rule:

Orbitals fill in order of increasing n+l, and for equal n+l, in order of increasing n.

Aufbau Principle:

Electrons fill orbitals in order of increasing energy.

Periodic Table

Elements are arranged by atomic number and grouped by properties.

Groups:

Columns with similar properties and electron configurations.

Periods:

Rows with the same outermost energy level.

Periodic Properties:

  • Ionization Energy: Energy to remove an electron.
  • Electron Affinity: Energy released when an atom gains an electron.
  • Electronegativity: An atom’s attraction for electrons.
  • Atomic Radius: Size of an atom.

Chemical Bonding

Chemical Bond:

The union of atoms to achieve greater stability.

Bond Energy:

Energy released during bond formation.

Bond Length:

Distance between bonded atoms.

Lewis Structures:

Representation of valence electrons.

Types of Bonds:

  • Ionic Bond: Attraction between oppositely charged ions.
  • Covalent Bond: Sharing of electrons between atoms.
  • Polar Covalent Bond: Unequal sharing of electrons.
  • Metallic Bond: Delocalized electrons shared among metal atoms.
  • Hydrogen Bond: Attraction between a hydrogen atom and an electronegative atom.
  • Van der Waals Forces: Weak attractions between molecules.