Atoms and Elements: Foundations of Atomic Theory

Atoms and Elements

Introduction to Atoms

Atoms are the fundamental building blocks of matter. The term atom is derived from the Greek word atomos, meaning "indivisible." Atoms retain the properties of an element and cannot be divided further without losing those properties.

• Atoms compose all ordinary matter.

• Understanding matter requires understanding atoms.

Modern Atomic Theory

Development of Atomic Theory

The concept that all matter is composed of atoms is supported by several fundamental laws:

• Law of Conservation of Mass

• Law of Definite Proportions

• Law of Multiple Proportions

Law of Conservation of Mass

Formulated by Antoine Lavoisier, this law states:

• In a chemical reaction, matter is neither created nor destroyed.

• The total mass of substances involved in a reaction remains constant.

Equation:

$\text{Total Mass of Reactants} = \text{Total Mass of Products}$

This law supports the idea that matter is composed of small, indestructible particles.

Law of Definite Proportions (Law of Constant Composition)

All samples of a given compound, regardless of their source or preparation, have the same proportions of their constituent elements.

• For example, water always decomposes into hydrogen and oxygen in a mass ratio of 2.0 g H to 16.0 g O, or 1:8.

Equation:

$\text{Mass ratio} = \frac{16.0\ \text{g O}}{2.0\ \text{g H}} = 8.0$

Law of Multiple Proportions (Dalton’s Law)

When two elements (A and B) form more than one compound, the masses of B that combine with a fixed mass of A are in ratios of small whole numbers.

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

• CO2: 2.67 g O per 1 g C; CO: 1.33 g O per 1 g C.

• The ratio of these masses is $\frac{2.67}{1.33} = 2$ (a small whole number).

Dalton’s Atomic Theory

John Dalton explained the above laws with his atomic theory:

1. Each element is composed of tiny, indestructible particles called atoms.

2. All atoms of a given element have the same mass and properties, distinguishing them from atoms of other elements.

3. Atoms combine in simple, whole-number ratios to form compounds.

4. Atoms of one element cannot change into atoms of another element; chemical reactions only rearrange how atoms are bonded.

Discovery of Subatomic Particles

Discovery of the Electron

• J.J. Thomson’s cathode ray experiments showed that atoms contain negatively charged particles (electrons).

• Electrons travel in straight lines, are independent of the cathode material, and carry a negative charge.

• Thomson measured the charge-to-mass ratio of the electron: $-1.76 \times 10^8$ C/g.

Millikan’s Oil Drop Experiment

• Robert Millikan measured the charge of a single electron: $-1.6 \times 10^{-19}$ C.

• Using Thomson’s ratio, the mass of the electron was calculated as $9.10 \times 10^{-28}$ g.

Structure of the Atom: Plum-Pudding Model

• J.J. Thomson proposed that electrons are embedded in a positively charged sphere (plum-pudding model).

Rutherford’s Gold Foil Experiment

• Ernest Rutherford’s experiment showed that most of the atom is empty space, with a small, dense, positively charged nucleus.

• Most alpha particles passed through gold foil, but some were deflected, indicating a concentrated positive charge.

Nuclear Theory of the Atom:

1. Most of the atom’s mass and all positive charge are in the nucleus.

2. Most of the atom’s volume is empty space with electrons dispersed throughout.

3. The number of electrons equals the number of protons, making the atom electrically neutral.

Discovery of the Neutron

• James Chadwick discovered the neutron, a neutral particle in the nucleus with a mass similar to the proton.

• Neutrons account for the difference between atomic number and atomic mass.

Subatomic Particles

Particle	Symbol	Charge	Mass (kg)
Proton	p+	+1	1.67262 × 10−27
Neutron	n	0	1.67493 × 10−27
Electron	e−	−1	9.1 × 10−31

Note: 1 atomic mass unit (amu) = 1.66054 × 10−27 kg.