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

Key scientific laws: Lavoisier's law of conservation of mass states that mass is conserved in chemical reactions; Proust's law of definite proportions shows compounds have constant composition.

Dalton's atomic theory (1808) proposed that:

• All matter is made of solid, indivisible atoms.

• Atoms are indestructible and retain identity in reactions.

• Atoms of the same element are identical; different elements have different atoms with distinct masses.

• Compounds form from elements in small whole-number ratios.

Atoms are composed of three subatomic particles:

• Protons (positive charge, in nucleus)

• Neutrons (neutral, in nucleus)

• Electrons (negative charge, orbit 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: Charge=#protons-#electrons

Each atom is defined by its atomic number (Z, number of protons) and mass number (A, number of nucleons):

• A=#protons+#neutrons

• Z=#protons

Isotopes are atoms of the same element (same atomic number) with different mass numbers due to varying numbers of neutrons.

Isotope ratios are valuable 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 isotopes by separating atoms based on mass, producing a spectrum that shows the proportion of each isotope in a sample.

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=∑(fractional abundance)×(isotope mass)

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