History of Atomic Theory and Electron Configurations

Atomic Theory History

• Early Theories:
  • Democritus (5th century BC): Proposed that matter is made of indivisible atoms.
  • John Dalton (1803): Introduced the Atomic Theory; suggested atoms are solid spheres and different elements have different atoms.
• J.J. Thomson (1897):
  • Discovered the electron using the cathode ray tube; proposed the "plum pudding model" where electrons are embedded in a positively charged sphere.
• Ernest Rutherford (1911):
  • Conducted the gold foil experiment; concluded that atoms have a small, dense nucleus surrounded by electrons.
• Niels Bohr (1913):
  • Developed the Bohr model, introducing quantized electron orbits around the nucleus.
• Quantum Mechanical Model (1920s):
  • Based on Schrödinger's wave equation; electrons are described by probability distributions (orbitals) rather than fixed orbits.

Electron Configurations

• Definition: Distribution of electrons in an atom’s orbitals.
• Energy Levels: Electrons occupy energy levels (shells), which are numbered (n=1, 2, 3...).
• Aufbau Principle: Electrons fill the lowest energy orbitals first.
• Pauli Exclusion Principle: No two electrons in an atom can have the same four quantum numbers.
• Hund’s Rule: Electrons will occupy degenerate orbitals singly before pairing up.
• Notation:
  • Example: Oxygen (O) - 1s² 2s² 2p⁴.

Quantum Mechanics And Atoms

• Wave-Particle Duality: Electrons exhibit both particle-like and wave-like behavior.
• Uncertainty Principle (Heisenberg): Cannot simultaneously know the exact position and momentum of an electron.
• Quantum Numbers:
  • Principal (n): Energy level.
  • Azimuthal (l): Shape of orbital (s, p, d, f).
  • Magnetic (m_l): Orientation of the orbital.
  • Spin (m_s): Direction of electron spin (+1/2 or -1/2).

Isotopes And Ions

• Isotopes:
  • Atoms of the same element with different numbers of neutrons.
  • Example: Carbon-12 (6 neutrons) and Carbon-14 (8 neutrons).
  • Identified by mass number (sum of protons and neutrons).
• Ions:
  • Atoms that have gained or lost electrons, resulting in a charge.
  • Cations: Positively charged ions (loss of electrons).
  • Anions: Negatively charged ions (gain of electrons).

Subatomic Particles

• Protons:
  • Positive charge (+1).
  • Mass: ~1 amu (atomic mass unit).
  • Located in the nucleus.
• Neutrons:
  • No charge (neutral).
  • Mass: ~1 amu.
  • Located in the nucleus.
• Electrons:
  • Negative charge (-1).
  • Mass: ~1/1836 amu (very small).
  • Orbit around the nucleus in electron clouds/orbitals.
• Charge Balance: In a neutral atom, the number of protons equals the number of electrons.

Atomic Theory History

• Democritus (5th century BC) proposed that matter consists of indivisible atoms.
• John Dalton (1803) introduced the Atomic Theory, describing atoms as solid spheres characteristic to each element.
• J.J. Thomson (1897) discovered the electron and developed the "plum pudding model," suggesting that electrons are embedded in a positively charged mass.
• Ernest Rutherford (1911) conducted the gold foil experiment, revealing that atoms contain a small, dense nucleus surrounded by electrons.
• Niels Bohr (1913) created the Bohr model, which introduced quantized electron orbits around the nucleus.
• The Quantum Mechanical Model (1920s) emerged, based on Schrödinger's wave equation, depicting electrons through probability distributions instead of fixed orbits.

Electron Configurations

• Electron configuration refers to the arrangement of electrons in an atom’s orbitals.
• Electrons occupy energy levels or shells, denoted by principal quantum number n (n=1, 2, 3...).
• The Aufbau Principle states that electrons fill the lowest energy orbitals first.
• According to the Pauli Exclusion Principle, no two electrons in an atom can possess identical quantum numbers.
• Hund’s Rule dictates that electrons must occupy degenerate orbitals singly before pairing occurs.
• Example of electron configuration: Oxygen (O) is represented as 1s² 2s² 2p⁴.

Quantum Mechanics And Atoms

• Electrons demonstrate wave-particle duality, exhibiting characteristics of both particles and waves.
• The Uncertainty Principle (Heisenberg) asserts that one cannot know both the exact position and momentum of an electron simultaneously.
• Quantum Numbers describe the properties of electrons:
  • Principal (n): identifies the energy level.
  • Azimuthal (l): defines the shape of the orbital (types: s, p, d, f).
  • Magnetic (m_l): indicates the orientation of the orbital.
  • Spin (m_s): denotes the direction of electron spin (+1/2 or -1/2).

Isotopes And Ions

• Isotopes are variants of the same element that have different neutron counts.
• Example: Carbon-12 has 6 neutrons; Carbon-14 has 8 neutrons. Identified by mass number (total protons and neutrons).
• Ions are charged atoms formed by the loss or gain of electrons.
• Cations are positively charged ions created by electron loss.
• Anions are negatively charged ions formed by electron gain.

Subatomic Particles

• Protons have a positive charge (+1), a mass of approximately 1 amu (atomic mass unit), and reside in the nucleus.
• Neutrons carry no charge (neutral), possess a mass of about 1 amu, and are also located in the nucleus.
• Electrons are negatively charged (-1), have a very small mass (~1/1836 amu), and orbit the nucleus in electron clouds or orbitals.
• Charge balance in neutral atoms is maintained by equal numbers of protons and electrons.