Modern chemistry began with Eighteenth century discoveries leading to the formulation of two basic laws of chemical combination:

The law of conservation of mass and the law of constant composition (definite proportions). Dalton proposed another law of chemical combination, the law of multiple proportions.

• The first clues to the structure of atoms came through the discovery and characterization of cathode rays (electrons). Key experiments were those that established the mass-to-charge ratio and then the charge on an electron.
• The principal types of radiation emitted by radioactive substances are alpha particles, beta particles, and gamma rays.
• Studies on the scattering of particles by thin metal foils (Rutherford’s atomic model) led to the concept of the nuclear atom − a tiny, but massive, positively charged nucleus surrounded by lightweight, negatively charged electrons.

Dalton’s Atomic Theory

Proposed by John Dalton in the early 19th century. States that:

• All matter is composed of tiny, indivisible particles called atoms.
• Atoms of the same element are identical in size, mass, and properties.
• Atoms of different elements combine in simple whole-number ratios to form compounds.
• Atoms cannot be created, divided into smaller particles, or destroyed in chemical reactions.

Modern Atomic Theory

Developed in the early 20th century based on experimental evidence, including discoveries of subatomic particles.States that:

• Atoms are composed of subatomic particles: protons, neutrons, and electrons.
• Protons and neutrons reside in the nucleus, while electrons orbit the nucleus in specific energy levels.
• Atoms of the same element can have different numbers of neutrons (isotopes).
• Chemical reactions involve the rearrangement of atoms; atoms are not destroyed or created.

Subatomic particles

Expands on the modern atomic theory by describing the characteristics and properties of subatomic particles:

• Protons: Positively charged particles found in the nucleus.
• Neutrons: Neutral particles found in the nucleus.
• Electrons: Negatively charged particles that orbit the nucleus in specific energy levels.

• Subatomic particles have distinct masses and charges that contribute to the overall structure and behavior of atoms.

Isotopes

• Isotopes are different forms of the same element that have the same number of protons (same atomic number) but different numbers of neutrons (different atomic mass).
• Isotopes have similar chemical properties due to identical numbers of protons, but their atomic masses differ.
• Examples include carbon-12 (12C) and carbon-14 (14C), both isotopes of carbon with different numbers of neutrons.

Average Atomic Mass

• The average atomic mass of an element is the weighted average mass of all naturally occurring isotopes of that element.
• It takes into account the relative abundance of each isotope
• The average atomic mass is found on the periodic table for each element and is useful in chemical calculations and understanding atomic structure.