Atomic Theory History and Periodic Table Study Notes

Democritus (400 B.C.)

  • First suggested the idea of the atom.
  • Atoms were described as:
    • indivisible
    • indestructible
    • always in motion
    • separated by empty space

John Dalton (1808)

  • Developed the first scientific atomic theory.
  • Key propositions:
    • All matter is made of atoms.
    • Atoms of the same element are identical (this was later disproved by isotopes).
    • Atoms cannot be subdivided, created, or destroyed (disproved by nuclear reactions).
    • Atoms combine in whole number ratios to form compounds.
    • In chemical reactions, atoms are rearranged, combined, or separated.
  • Model: Billiard Ball sphere.

J.J. Thomson (1897)

  • Discovered the electron using the Cathode Ray Tube Experiment.
  • Electrons were originally called corpuscles.
  • Proposed the Plum Pudding Model: atom = positive sphere with electrons scattered inside.
  • Analogy: like a blueberry muffin (positive interior with embedded electrons).

Ernest Rutherford (1908)

  • Performed the Gold Foil Experiment.
  • Discovered the nucleus with protons inside.
  • Showed that atoms are mostly empty space.
  • Model: Nuclear model of the atom.

Niels Bohr (1922)

  • Electrons move in fixed orbits or energy levels around the nucleus.
  • Electrons can jump by absorbing or releasing energy.
  • When electrons drop to lower energy levels, they release energy as light.

Evolution of Models (in order)

  • 1) Democritus – indivisible particles
  • 2) Dalton – solid sphere
  • 3) Thomson – plum pudding model
  • 4) Rutherford – nucleus discovered
  • 5) Bohr – electrons in fixed orbits

Guided Notes: The Periodic Table

  • Link: https://youtu.be/diT5qYHriJQ?si=MqmWShvVDphr3uo5

Dmitri Mendeleev and the Periodic Table (1869)

  • Organized elements to show repeating properties.
  • The modern table shows properties and trends of elements.

Structure of the Periodic Table

  • Periods (Rows) → show number of electron shells/energy levels.
  • Groups (Columns) → show valence electrons and bonding patterns.

Element Box

  • Name and symbol of element.
  • Atomic Number = number of protons.
  • Atomic Mass = average of protons + neutrons (decimal value is the average of isotopes).
  • Number of electrons = number of protons (in a neutral atom).

Energy Levels

  • Electrons arranged in levels around the nucleus.
  • 1st level holds 22 max.
  • 2nd level holds 88 max.
  • Higher levels may hold more, but the rule of thumb for stability is an octet rule: 88 electrons in the outer shell.

Valence Electrons & Bonding

  • Atoms want a full outer shell to achieve stability.
  • Will form ions by gaining or losing electrons to complete the outer shell.
  • Common charges by group (loss/gain corresponds to achieving a full outer shell):
    • Group 1 → +1+1 charge (lose 1 e⁻)
    • Group 2 → +2+2 charge (lose 2 e⁻)
    • Group 13 → +3+3 charge (lose 3 e⁻)
    • Group 14 → ±4±4 charge (can lose or gain 4 e⁻)
    • Group 15 → −3-3 charge (gain 3 e⁻)
    • Group 16 → −2-2 charge (gain 2 e⁻)
    • Group 17 → −1-1 charge (gain 1 e⁻)
    • Group 18 → 00 charge (inert/stable, no gain/loss)

Metals, Nonmetals, Metalloids

  • Metals: located on the left side of the zigzag line; typically shiny, malleable, ductile, good conductors.
  • Nonmetals: located on the right side; can be gases or brittle solids; generally poor conductors.
  • Metalloids: located along the zigzag line; properties of both metals and nonmetals.

Compounds

  • Ionic compounds = metal + nonmetal (electrons are transferred).
  • Covalent compounds = nonmetals (electrons are shared).
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