Atomic Theory

Democritus

  • Lifespan: 460 B.C. – 370 B.C.

  • Proposed the existence of atoms (from the Greek word "atomos" meaning "not to be cut").

  • Atomic model: the “solid sphere” model.

  • Core ideas:

    • Atoms are indivisible and indestructible pieces of matter.
    • Atoms are the smallest unit of matter.
    • Atoms are infinite in number, always moving, and capable of joining together to form larger structures.
    • Atoms are small, hard particles that are all made of the same material but differ in shape and size.

  • Key terminology:

    • atomos = "not to be cut".

  • Major limitations:

    • Did not explain chemical behavior.
    • Was not based on the scientific method; relied on philosophy rather than empirical testing.

  • Connections to broader themes:

    • Transition from philosophical speculation to empirical inquiry; foreshadowed the need for experimental validation in atomic theory.

  • Ethical/philosophical implications:

    • Demonstrates early tension between metaphysical ideas and testable science; highlights the importance of method and evidence in scientific progress.

John Dalton

  • Lifespan: 1766 – 1844.

  • Dalton’s postulates (early atomic theory):

    • All matter is made of atoms.
    • Atoms of an element are identical.
    • Each element has different atoms.
    • Atoms of different elements combine in constant ratios to form compounds.
    • Atoms are rearranged in reactions.

  • Dalton’s model: the “billiard ball” or “solid sphere” model.

  • Improvements over Democritus:

    • Theory grounded in scientific experimentation, not mere philosophy.
    • Accounts for the Law of Conservation of Mass: atoms are neither created nor destroyed in reactions.
    • Agrees with the Law of Constant Composition: elements combine in fixed ratios to form compounds.

  • Mathematical/empirical relations linked to Dalton’s view:

    • Law of Conservation of Mass: in any closed system, total mass remains constant during a chemical reaction.
    • Law of Constant Composition (Definite Proportions): compounds contain elements in fixed, definite ratios.

  • Problems with Dalton’s theory:

    • Later discoveries showed atoms are divisible into subatomic particles (electrons, protons, neutrons) and that there are many other particles.

  • Connections to previous/next topics:

    • Sets the stage for identifying subatomic particles and revising the model of the atom as science progresses.

  • Practical implications:

    • Provided a robust framework for understanding chemical reactions, stoichiometry, and the conservation of mass in chemistry.

  • Notable formulas/equations:

    • Law of Conservation of Mass can be written as m_ ext{reactants} = m_ ext{products}.
    • Law of Constant Composition (definite proportions) can be summarized as the constant ratio of element masses in a given compound, e.g. for a general compound rac{mA}{mB} = ext{constant}.

J. J. Thomson

  • Thomson’s role:

    • Proposed the existence of subatomic particles; provided the first hint that atoms are divisible.

  • Model: the "plum pudding" model.

    • Atoms composed of a positively charged substance with negatively charged electrons scattered within, like raisins in a pudding.

  • Experimental hint:

    • Studied the passage of an electric current through a gas; the current produced rays of negatively charged particles.

  • Key outcomes:

    • Discovery of the electron.
    • Confirmation that atoms contain smaller, charged particles and thus are divisible.

  • Limitations:

    • Could not discover the proton (at this stage).

  • Metaphor/metaphor usage:

    • The plum pudding metaphor illustrated how negative charges could be embedded in a positively charged matrix.

  • Connections to broader themes:

    • Transition from indivisible-atom concept to a composite structure with subatomic particles.

  • Significance for science:

    • Demonstrated that atoms have internal structure, prompting subsequent experiments to locate and characterize subatomic particles.

  • Notable formulas/equations:

    • While Thomson’s work mainly identified electrons and charge/mass relationships, no specific standard equation attributed to his model is provided in the transcript. The qualitative insight is that electrons exist and reside within a positively charged matrix.

Ernest Rutherford

  • Gold Foil Experiment (1911):

    • Positively charged alpha particles were directed at a thin sheet of gold foil.
    • Some alpha particles reflected back or were deflected—as if they had struck something solid.
    • Conclusion: positive charges repel positive charges; most of the atom’s mass and positive charge are concentrated in a small, dense region—the nucleus.

  • Rutherford’s findings:

    • The nucleus is small.
    • The nucleus is dense.
    • The nucleus is positively charged.
    • Electrons move around the nucleus much like planets orbiting the sun.
    • This led to the nuclear model (also called the planetary model) of the atom.

  • Problems and limitations:

    • Based on classical (pre-quantum) physics; later abandoned due to discoveries in quantum physics.

  • Implications:

    • Shifted the view from a uniform positive charge spread (plum pudding) to a concentrated nucleus with electrons circling around it.

  • Connections to previous topics:

    • Direct empirical evidence for subatomic structure, paving the way for quantum models.

  • Practical implications:

    • Laid groundwork for nuclear physics and later quantum theory refinement of atomic structure.

  • Notable equations:

    • The qualitative description doesn’t include a specific standard equation in the transcript, but the core idea is the nuclear model: a tiny, dense, positively charged nucleus with electrons surrounding it.

Niels Bohr

  • Bohr Model (1913):

    • Bohr proposed an improvement to Rutherford’s model; sometimes referred to as the Rutherford–Bohr model.

  • Key features:

    • Electrons orbit the nucleus in orbits that have a set size and energy.
    • The energy of the orbit is related to its size; the lowest energy is found in the smallest orbit.
    • Radiation is absorbed or emitted when an electron moves from one orbit to another (quantized transitions).

  • Problems/limitations:

    • We cannot predict the exact location and orbit of electrons (precisely); quantum behavior becomes significant.
    • Predictions fail for larger atoms (i.e., atoms more complex than hydrogen).

  • Significance:

    • Introduced quantization of energy levels and explained spectral lines of hydrogen, representing a major step toward quantum theory.

  • Connections to prior sections:

    • Builds on Rutherford’s nucleus with quantized electron orbits to account for discrete spectral lines.

  • Mathematical notes:

    • Energy levels are quantized; for a hydrogen-like system, energy levels are often described (in simplified terms) by a relation of the form E_n ext{ depends on } n ext{ (quantum number)}, with the lowest energy at the smallest orbit (n=1).
    • Transitions involve emission/absorption of photons with energy difference riangle E = Ef - Ei = h\nu = \frac{hc}{\lambda}.

James Chadwick

  • Discovery: 1932 – confirmed the existence of the neutron.

  • Neutron properties:

    • Charge: neutral (
    • Mass: nearly equal to that of a proton (
      -Notation: mass of neutron approximately equal to proton, i.e., mn \approx mp.)

  • Experimental approach:

    • Fire alpha particles into a beryllium target; allow the resulting radiation to interact with paraffin wax.
    • Interactions between radiation and the hydrogen in the wax led to the discovery of the neutron.

  • Improvements to the atomic model:

    • The neutron’s inclusion made the atomic model more complete.
    • Overall charges remained the same (i.e., the total positive charge from protons and the negative charge from electrons balance as before).
    • Resolves discrepancies between atomic mass and atomic number by accounting for a substantial portion of atomic mass being carried by neutrons (which contribute mass but not charge).

  • Connections to previous topics:

    • Completes the basic subatomic picture (protons + neutrons in the nucleus, electrons orbiting) needed to explain atomic mass and behavior more accurately.

  • Practical/ethical implications:

    • Understanding neutron properties has had profound implications for nuclear physics, energy, and radiation safety.

  • Notable equations:

    • Neutron mass relation: mn \approx mp (nearly equal masses).
    • Charge balance context is implicit in the statement that total charge remains the same; the nucleus contains protons (positive charge) while electrons provide negative charge to balance overall.