Atoms, Molecules, and Ions – Foundations of Atomic Theory

Historical Background

Early Models of the Atom

The concept of the atom has evolved over centuries, beginning with ancient Greek philosophers who proposed early models of matter.

• Empedocles (5th century B.C.): Proposed that all matter is composed of four fundamental elements: Earth, Air, Water, and Fire.

• Democritus (ca. 370 B.C.): Suggested that matter is divisible into tiny, indivisible particles called atoms (from Greek "atomos" meaning indivisible).

• Aristotle (ca. 322 B.C.): Argued that matter is continuous, not atomistic, and opposed Democritus' theory.

Atomic Theory of Matter

Dalton’s Atomic Theory and Fundamental Laws

Experiments in the 18th and 19th centuries led to the development of an organized atomic theory by John Dalton in the early 1800s. Dalton’s theory was based on several fundamental laws:

• Law of Constant Composition (Definite Proportions): A compound always contains two or more elements combined in a definite proportion by mass.

• Law of Multiple Proportions: Elements may combine in different ratios to produce more than one compound.

• Law of Conservation of Mass: Matter can be neither created nor destroyed in chemical reactions.

The Law of Definite Proportions

Definition and Example

This law states that a chemical compound always contains its component elements in a fixed ratio by mass, regardless of the source or method of preparation.

• Example: Hydrogen and oxygen combine to form water (H2O) and hydrogen peroxide (H2O2), each with a specific mass ratio of hydrogen to oxygen.

Compound	By mass of H	By mass of O	Atomic composition
Water (H2O)	11.2%	88.8%	2H + 1O
Hydrogen peroxide (H2O2)	5.9%	94.1%	2H + 2O

Law of Multiple Proportions

Definition and Example

If two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in ratios of small whole numbers.

• Example: Carbon and oxygen form carbon monoxide (CO) and carbon dioxide (CO2).

• Ratio of oxygen in CO to oxygen in CO2 is 1:2.

Law of Conservation of Mass

Definition and Application

This law states that the total mass of substances present after a chemical reaction is the same as the total mass before the reaction.

• Example: When atoms of element X and element Y react to form a compound, the total number of atoms and their mass remain unchanged.

Equation:

$\text{Total mass before reaction} = \text{Total mass after reaction}$

Key Point: Matter can be neither created nor destroyed.

Dalton’s Atomic Theory (1808)

Postulates

Dalton’s atomic theory provided a foundation for modern chemistry:

• All matter is composed of extremely small particles called atoms.

• Atoms of a given element are identical in size, mass, and chemical properties.

• Atoms of different elements differ in size, mass, and properties.

• Compounds are composed of atoms of more than one element, combined in fixed ratios.

• Chemical reactions involve the rearrangement of atoms; atoms are not created or destroyed.

Additional info:

• These foundational laws and Dalton’s theory set the stage for the development of modern atomic models and the understanding of chemical reactions.