Atoms and Elements: Foundations of Atomic Theory and Structure

Atoms and Elements

Objectives

• Describe the laws that are the basis of modern atomic theory.

• Summarize experiments that led to the discovery of the electron and its properties.

• Explain the structure of an atom.

• Identify properties of subatomic particles and interpret isotope symbols.

• Apply the periodic law to the organization of the periodic table.

• Determine the charge of ions.

• Calculate atomic mass of elements.

• Use the mole concept.

Foundations of Atomic Theory

Early Ideas: Democritus and the Atom

The concept of the atom dates back to ancient Greece, where Democritus (c. 460–370 BC) proposed that all matter is composed of tiny, indivisible particles called atomos (meaning "uncuttable"). He suggested that everything is made of atoms and empty space. Although this idea was not scientifically tested at the time, it laid the groundwork for later atomic theory.

• Atomos: The smallest indivisible unit of matter.

• Historical context: Democritus' ideas were philosophical, not experimental.

Dalton's Atomic Theory

John Dalton (early 19th century) formalized atomic theory based on experimental evidence. His postulates form the basis of modern chemistry:

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

2. All atoms of a given element are identical in mass and properties.

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

4. Atoms of one element cannot change into atoms of another element in chemical reactions; they only change the way they are bound together.

Fundamental Laws Supporting Atomic Theory

Three key laws support Dalton's atomic theory and the concept of atoms:

Law	How Dalton's Theory Explains It
Law of Conservation of Mass	In a chemical reaction, matter is neither created nor destroyed.
Law of Definite Proportions	All samples of a given compound, regardless of source or preparation, have the same proportions of their constituent elements.
Law of Multiple Proportions	When two elements form different compounds, the masses of one element that combine with a fixed mass of the other are in a ratio of small whole numbers.

Law of Conservation of Mass

• In any chemical reaction, the total mass of reactants equals the total mass of products.

• Example: Decomposition of 10 kg of CaCO3 yields CaO, CO2, and H2O, with total mass conserved.

Law of Definite Proportions

• A given compound always contains the same elements in the same proportion by mass.

• Example: Water (H2O) always has a mass ratio of oxygen to hydrogen:

$\text{Mass ratio} = \frac{16\,\text{g O}}{2\,\text{g H}} = 8:1$

• If 36 g of water decomposes to 32 g O and 4 g H, the ratio remains 8:1.

Law of Multiple Proportions

• When 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 CO and CO2:

$\text{CO}_2: \frac{32\,\text{g O}}{12\,\text{g C}} = 8:3$

$\text{CO}: \frac{16\,\text{g O}}{12\,\text{g C}} = 4:3$

• The ratio of oxygen in CO2 to CO is $\frac{8}{4} = 2:1$ (a small whole number ratio).

Discovery of Subatomic Particles

Cathode Ray Tube Experiments

Experiments with cathode ray tubes demonstrated that atoms are divisible and contain subatomic particles:

• Electrons: Negatively charged particles discovered by J.J. Thomson.

• Electrons are much lighter than hydrogen atoms (about 2000 times lighter).

• Atoms are not indivisible; they contain smaller particles.

Properties of Subatomic Particles

Particle	Mass (kg)	Relative Mass (amu)	Charge
Proton	1.6726 × 10-27	1	+1
Neutron	1.6749 × 10-27	1	0
Electron	9.1094 × 10-31	~0	-1

• Protons and neutrons have similar mass (~1 amu), electrons are much lighter.

• Atoms are electrically neutral: number of protons equals number of electrons.

Atomic Number and Isotopes

• Atomic number (Z): Number of protons in the nucleus; defines the element.

• Mass number (A): Total number of protons and neutrons.

• Isotopes: Atoms of the same element with different numbers of neutrons.

Example: Hydrogen has three isotopes: protium (1 proton), deuterium (1 proton, 1 neutron), tritium (1 proton, 2 neutrons).

Additional info:

• Further slides likely cover periodic table organization, ions, atomic mass calculations, and the mole concept.