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): Mass is neither created nor destroyed in chemical reactions.

• Law of Definite Proportions (Proust): Compounds have constant composition by mass.

Dalton's Atomic Theory (1808):

• All matter is composed of solid, indivisible atoms.

• Atoms are indestructible and retain identity in chemical reactions.

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

• Compounds are formed from elements in small, whole-number ratios.

Atoms are made of three subatomic particles:

• Protons (positive charge, in nucleus)

• Neutrons (neutral, in nucleus)

• Electrons (negative charge, outside nucleus)

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

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

Atomic number (Z) is the number of protons and defines the element; mass number (A) is the total number of protons and neutrons (nucleons):

• A=p+n

Isotopes are atoms of the same element (same Z) with different mass numbers (A), due to varying numbers of neutrons.

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

Mass spectrometry is used to measure isotope abundances by separating isotopes based on mass and generating a spectrum showing their proportions.

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=∑i(fractional abundance of isotope i)×(mass of isotope i)

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