Early atomic theory began with Ancient Greek philosophers, who believed all matter was made of four elements (air, earth, fire, water). Democritus proposed that matter could be divided until reaching indivisible particles called atoms.

Key scientific laws established:

• Law of Conservation of Mass (Lavoisier, 1785): Mass is neither created nor destroyed in chemical reactions.

• Law of Definite Proportions (Proust, 1794): Compounds have constant composition by mass.

Dalton's Atomic Theory (1808) included:

• All matter is made 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.

While foundational, some postulates have since been revised.

Atoms are composed of three subatomic particles:

• Protons (positive charge, in nucleus)

• Neutrons (no charge, in nucleus)

• Electrons (negative charge, orbit nucleus)

Most of the atom's mass is in the tiny, dense nucleus; most of its volume is empty space.

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

Each atom is defined by:

• Atomic number (Z): number of protons (determines the element)

• Mass number (A): total number of protons and neutrons (nucleons)

Changing the number of protons changes the element; atoms of the same element with different numbers of neutrons are called isotopes.

Isotopes have important applications:

• Used in tracing and dating samples in biology, geology, archaeology, and forensics.

• Example: 14C in tooth enamel can determine year of birth due to changes in atmospheric 14C from nuclear testing.

Isotopes are measured using mass spectrometry, which separates atoms based on mass and provides a spectrum showing 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, calculated as: A=∑iisotopes(fractional abundance × isotope mass)

Example calculations:

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

• Gallium's two isotopes and their abundances can be determined from the average atomic mass and individual isotope masses.