History of Atomic Theory and Electron Configurations

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.

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.

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:
- 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).

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.

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 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.

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 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.

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.