Early atomic theory began with Greek philosophers, who believed all matter was made of four elements; Democritus introduced the concept of the atom as an indivisible particle.

Key laws in atomic theory:

• Law of Conservation of Mass (Lavoisier): Matter is neither created nor destroyed in chemical reactions.

• Law of Definite Proportions (Proust): Compounds always contain the same elements in the same proportions by mass.

Dalton's Atomic Theory (1808) established that:

• All matter is made of indivisible atoms.

• Atoms of the same element are identical; atoms of different elements differ in mass and properties.

• Atoms combine in simple whole-number ratios to form compounds.

• Atoms retain their identity in chemical reactions (though later discoveries modified some of these points).

Atoms are composed of three subatomic particles:

• Protons (positive charge, in nucleus)

• Neutrons (no charge, in nucleus)

• Electrons (negative charge, outside nucleus)

Most of an atom's volume is empty space; the nucleus is extremely small but contains nearly all the atom's mass.

The overall charge of an atom is determined by the difference between the number of protons and electrons: Q=p−e where p = number of protons, e = number of electrons.

Each element is defined by its atomic number (Z, number of protons). The mass number (A) is the sum of protons and neutrons: A=p+n.

Isotopes are atoms of the same element (same Z) with different numbers of neutrons (different A).

Isotope ratios are valuable in fields like biology, geology, archaeology, and forensics (e.g., using 14C in tooth enamel to estimate year of birth).

Mass spectrometry is used to measure isotopic composition by separating isotopes based on mass and displaying their relative abundances as a spectrum.

Most elements exist as mixtures of isotopes; the atomic mass on the periodic table is a weighted average of all naturally occurring isotopes: Atomic mass=∑in(fractional abundance of isotope i)×(mass of isotope i)

Example: Silicon has three naturally occurring isotopes with different abundances and masses; the average atomic mass is calculated using their weighted contributions.