1. Structure of the atom

Definition of Atom: The smallest unit quantity of an element capable of existing either alone or in chemical combination.
Fundamental Particles:
- Protons: Positively charged particles found in the nucleus.
- Neutrons: Neutral particles also located in the nucleus.
- Electrons: Negatively charged particles found in electron shells surrounding the nucleus.

Nucleus: Contains protons and neutrons.
Electron Shells: Outermost regions containing electrons.
Mass Comparison: Protons and neutrons have approximately the same mass which is significantly greater than that of electrons.
Volume and Density: Over 99% of an atom's volume is empty space, yet solid objects cannot pass through due to electron repulsion.

Atomic Number (Z): Number of protons in an atom, determining the element.
Mass Number (A): Sum of protons and neutrons; contributions from electrons are negligible.

Neutrons can be determined by:
- Formula: Neutrons = Mass Number (A) - Atomic Number (Z)

Definition: Variants of an element with the same proton number but different neutron numbers.
Example:
- Carbon-12 (12C): 6 protons, 6 neutrons.
- Carbon-14 (14C): 6 protons, 8 neutrons, a radioisotope used in carbon dating (half-life = 5,730 years).

Democritus (5th Century BC): Proposed the concept of the atom as indivisible particles.
John Dalton (1766-1844): Showed that matter is composed of atoms with distinct weights.
Thomson's Model (1904): Suggested atoms are spheres of positive material with negatively charged electrons embedded (Plum Pudding Model).
Rutherford's Experiment (1900s): Established the nuclear model by his gold foil experiment, where α particles deflected, implying a dense nucleus.
Bohr Model (1913): Introduced quantized energy levels where electrons jump between orbits.

Heisenberg Uncertainty Principle (1927): Implies that one cannot know both the momentum and position of an electron simultaneously.
Schrödinger Equation: Introduced wave mechanics, describing electrons as waves distributed over space with probability density.

Principal Quantum Number (n): Energy level of electrons; can take integer values (1, 2, 3...).
Angular Momentum Quantum Number (l): Describes shape of the orbital; can take values from 0 to n-1.
Magnetic Quantum Number (ml): Indicates the orientation of the orbital in space, with allowed values ranging from -l to +l.
Spin Quantum Number (ms): Represents the two spin states of an electron (+1/2 or -1/2).

Definition: The distribution of electrons in atomic orbitals.
Filling Order: Electrons fill orbitals starting from lowest energy levels according to the Aufbau principle and the Pauli exclusion principle.
Hund’s Rule: Electrons will fill orbitals singly before pairing up to maximize number of unpaired electrons.

Arrangement: Elements are arranged by increasing atomic number (modern periodic law).
Periodic Properties: Include atomic radius, ionization energy, electron affinity, and electronegativity.

Definition: Distance from the nucleus to the outermost electron shell; generally increases down groups and decreases across periods.

Ionization Energy: Energy required to remove an electron; increases across a period and decreases down a group.
Electron Affinity: Energy change when an atom gains an electron; trends similarly with atomic radius.

Definition: An atom's ability to attract and bind with electrons; increases across periods and decreases down groups.
Pauling Scale: Common scale for quantifying electronegativity values for elements.