Atomic Theory: From Democritus to Modern Understanding

Posted on Feb 15, 2025 in Chemistry

Atomic Theory: A Historical Perspective

Democritus: Postulated the existence of atoms, suggesting that matter is discrete.

Aristotle: Denied the existence of atoms, proposing that matter is continuous.

Lavoisier (1777): Using a scale, he measured mass by weighing and established the Law of Conservation of Mass.

Mass and energy are related: E = mc² (Energy can be converted to mass and vice versa).

Dalton’s Atomic Theory (1880)

John Dalton postulated:

  1. Matter consists of atoms, which are structural units and indivisible.
  2. Atoms of different elements are different.
  3. Chemical compounds are composed of different elements in whole number ratios.
  4. Chemical reactions involve mass conservation.

Structure of the Atom

Discovery of the electron and its characterization (charge, mass).

Cathode Rays (1800s)

Air: A mixture of O₂ and N₂ + CO₂, etc., including Argon (Ar).

E + (molecules of air and glass)

E⁻ + x > x*

X + > e⁻ + x⁺ + light (visible)

Thomson’s Cathode Ray Experiment (1897)

  1. Application of an electric field.
  2. No electric or external magnetic fields.
  3. Application of a magnetic field (fluorescent display under condition c).

Thomson’s Discovery (1897)

Thomson (1897): Discovered negatively charged particles called electrons and determined the charge-to-mass ratio (qₑ⁻/mₑ⁻). He received the Nobel Prize in 1906.

Millikan’s Oil Drop Experiment (1908-1917)

Robert Millikan (1908-1917): Characterized the electron, determining its mass and charge using the oil drop experiment. He calculated the charge (q) of the oil droplets.

The Mole

Mole: The amount of substance that contains 6.02 x 10²³ particles.

Chemical Formulas

Chemical Formula: An abbreviated expression of the composition of a substance, both qualitatively and quantitatively. Example: NaCl, H₂O.

Empirical and Molecular Formulas

Empirical Formula: The simplest whole-number ratio of atoms in a compound.

Molecular Formula: The actual number of atoms of each element in a molecule of a compound.

Quantization

Quantization is always equal to or a multiple of: 1.6 x 10⁻¹⁹ Coulombs (charge of 1 electron).

Radioactivity

Becquerel (1896): Uranium ore (pitchblende) spontaneously emitted high-energy radiation, termed radioactivity.

Rutherford, Geiger, and Marsden

Plum Pudding Model (1900): A compact, positively charged sphere with electrons embedded within.

Rutherford’s Gold Foil Experiment (1910)

Particles deflected > repulsion due to collisions with positive electrostatic charges.

  • Most of the matter is empty space.
  • There is a positive charge concentrated in a nucleus.

Rutherford’s Atomic Model

Existence of a positively charged nucleus. Electrons move around the nucleus at large distances.

Discovery of the Proton

Rutherford (1919): Discovered the existence of protons, determining their mass (m) and charge (q).

Discovery of the Neutron

Chadwick (1932): Discovered the neutron.

Atomic Number and Mass Number

Atomic Number: Equal to the number of protons and electrons, determining the element’s position in the periodic table.

Mass Number: The total number of protons and neutrons in the nucleus.

Atomic Weight and Isotopes

Atomic Weight: The weighted average of the masses of all isotopes of an element.

Isotopes: Different forms of the same element that have different mass numbers.

Criticisms of Dalton’s Postulates

  1. Atoms are divisible and have subatomic particles.
  2. Atoms of the same element can have different atomic weights (isotopes).
  3. Existence of non-stoichiometric compounds.
  4. Relativistic conditions are not considered (but do not apply to normal laboratory and industrial conditions).

Limiting Reagent

Limiting Reagent: The reactant that is completely consumed in a chemical reaction, determining the amount of product formed.