Development of the Atom - Comprehensive Study Notes

Early Atomic Theory

  • Two major historical periods laid the groundwork for atomic theory: the development of Ancient Greek philosophy and the rise of alchemy. Discoveries from these eras later formed the basis of atoms and matter.

Ancient Greeks

  • THALES OF MILETUS (ca. 624-546 BCE)
    • Asked: What is matter composed of?
    • Proposed water as the ultimate substance because it could change into any kind of matter.
  • EMPEDOCLES OF AGRIGENTUM (ca. 490-430 BCE)
    • Suggested matter could be composed of fire or air.
    • Proposed four essential elements: air, water, earth, fire as the sources of all matter.
  • LEUCIPPUS (5th century BCE)
    • Proposed that matter consists of fragments that are indivisible yet retain the properties of the original material.
    • Considered the first atomist (less well-known than Democritus).
  • DEMOCRITUS (ca. 460-370 BCE)
    • Expanded Leucippus’ hypothesis; called the “father of atomism.”
    • Believed atoms were solid and indestructible; the universe consisted of atoms and empty space (the void).

Ancient Greeks (continued)

  • ARISTOTLE (384-322 BCE)
    • Stated that only observable things could be proven true.
    • Argued atoms were not observable, hence their existence could not be verified by observation.
    • Accepted the four fundamental elements (water, air, fire, earth) and their associated properties (hot, cold, dry, wet).
  • LUCRETIUS (ca. 100-51 BCE)
    • Wrote On the Nature of Things, reaffirming atomism.
    • Writings largely forgotten until rediscovered after the invention of the printing press in 1400 CE.

The Alchemists

  • Alchemy persisted through the 1700s and sought to change one substance into another (transmutation).
  • Practitioners were called alchemists; they studied properties of many substances and how they reacted.
  • Alchemy contributed to chemistry by advancing experimentation and observation.
  • Goals included the elixir of life and the philosopher’s stone (secretly guarded via symbols to protect notes).

Paracelsus and iatrochemistry

  • PARACELSUS (1493-1541)
    • Founded iatrochemistry (medical chemistry).
    • Produced metallic medicines to cure illnesses, promoting laboratory drug preparation rather than transmutation of metals.
    • Coined the term “alcohol” to describe spirits of wine.
  • By the late 18th century, traditional alchemy separated from chemistry.

Observations that led to the Modern Atomic Theory

  • ROBERT BOYLE (1627-1691)
    • Argued that elements water, fire, earth, and air are not true elements because they can be decomposed chemically into simpler substances (iron, oxygen, lead, hydrogen, sulfur).
    • Believed atoms exist, though not directly visible.
    • Key experiment: a J-shaped glass tube closed at one end and open at the other. Trapping a small amount of air at the closed end with mercury showed that increasing pressure decreased volume, suggesting air is composed of minute particles (atoms) separated by space and that they are compressed under pressure.

French Experiments

  • ANTOINE-LAURENT LAVOISIER (1743-1794)
    • Investigated chemical reactions by measuring masses of reactants and products.
    • Found that burning sulfur and phosphorus altered mass due to the involvement of oxygen from air.
    • Law of Conservation of Mass: mass is neither created nor destroyed in a chemical reaction.
  • JOSEPH-LOUIS PROUST (1754-1826)
    • Decomposed copper carbonate to copper, carbon, and oxygen; found constant and identical mass ratios among the elements across samples.
    • Established the 5:4:1 mass ratio for copper carbonate (CuCO₃) decomposition, demonstrating constant mass ratios in compounds (Definite Proportions).
    • Realized that ratios are whole numbers, suggesting atoms are discrete units (supporting Democritus’ idea).

Dalton’s Atomic Theory

  • JOHN DALTON (1766-1844)
    • All matter is composed of atoms — extremely small, indestructible, indivisible particles.
    • All atoms of a given element are identical in size, weight, and chemical properties.
    • Atoms differ between elements.
    • Atoms cannot be created or destroyed; they cannot be divided into smaller particles.
    • In chemical reactions, atoms are separated, combined, or rearranged, but not destroyed to form new atoms or molecules.
    • This theory laid the groundwork for modern chemistry and atomic structure.

Thomson’s Plum Pudding Model

  • JOSEPH JOHN THOMSON (1856-1940)
    • In 1897, discovered the electron via the cathode ray tube experiment; showed electrons exist.
    • Proposed that atoms could be divided further, contradicting Dalton’s indivisibility.
    • plum pudding model: atom is a sphere of positive charge with embedded electrons (plums) inside (pudding).

Rutherford’s Planetary Model

  • ERNEST RUTHERFORD (1871-1937)
    • Critically evaluated Thomson’s model through experiments (gold foil experiment not detailed here) and concluded Thomson’s model was incorrect.
    • Proposed a planetary model: a small, dense, positively charged nucleus at the center with electrons surrounding in mostly empty space.
    • Electrons orbit the positively charged nucleus like planets around the sun.

Bohr’s Model of the Atom

  • NIELS BOHR (1885-1962)
    • Worked in Rutherford’s lab (1913) to address flaws in the planetary model using hydrogen’s emission spectrum.
    • Core ideas:
    • Electrons occupy definite orbits around the nucleus (stationary orbits).
    • Each orbit has a specific energy level.
    • Electron transitions between levels involve emission (downward) or absorption (upward) of light.
    • The emitted/absorbed light has energy and a frequency determined by the energy difference between levels:
      ΔE=hν\Delta E = h\nu
    • The color (wavelength) of light is tied to this energy difference.

Emission Spectrum

  • EMISSION SPECTRUM
    • A series of narrow spectral lines at specific colors when emitted light passes through a prism.
    • According to Bohr’s model, electrons jumping between high- and low-energy orbits emit photons with characteristic wavelengths.
    • Bohr’s model is primarily applicable to the hydrogen atom.
    • Bohr received the Nobel Prize in Physics in 1922 for his work on atomic structure and spectra.

Nuclear Model of the Atom (Goldstein)

  • EUGEN GOLDSTEIN (1850-1930)
    • Performed a gas-discharge tube experiment with a perforated cathode.
    • Observed canal rays traveling from the anode to the cathode; proposed the existence of a positively charged particle to balance the negatively charged electrons.
    • Identified the positive particle as a proton; mass-to-charge ratio varied with different gases, indicating protons are hydrogenic ions of the gases.

James Chadwick

  • JAMES CHADWICK
    • Discovered the neutron (massive neutral particle).
    • Demonstrated that the particle emitted by bombarded beryllium with alpha particles was the neutron.
    • Concluded that the nucleus contains protons and neutrons, while electrons reside outside the nucleus.

Atomic Models Summary

  • The historical sequence of models:
    • Dalton’s atomic model
    • Thomson’s Plum Pudding model
    • Rutherford’s Planetary model
    • Bohr’s quantized orbits

Neutrality and Ions

  • An object is neutral when it has no net charge; in atoms, the number of protons equals the number of electrons.
  • Atoms with extra positive or negative charge are called cations (positive) and anions (negative), respectively.
  • The number of protons determines the identity of the element and is represented by the atomic number $Z$.

Subatomic Particles

  • PROTONS
    • Symbol: $p$; Relative charge: $+1$; Location: in the nucleus; Relative mass: $1$ amu; Discoverer: Eugen Goldstein (proton concept) and Rutherford associated with the term in broader context.
  • NEUTRONS
    • Symbol: $n$; Relative charge: $0$; Location: in the nucleus; Relative mass: $1$ amu; Discoverer: James Chadwick.
  • ELECTRONS
    • Symbol: $e$; Relative charge: $-1$; Location: outside the nucleus; Relative mass: 11840\dfrac{1}{1840} amu (≈ 9.11×10−319.11\times 10^{-31} kg); Discoverer: J. J. Thomson.

Atomic Symbols, Atomic Number, and Mass Number

  • Each element is represented by a symbol and has associated values:
    • Atomic number: $Z$ (number of protons)
    • Mass number: $A$ (sum of protons and neutrons)
    • Mass number and atomic number balance in nuclear reactions to satisfy conservation laws.

Prediction and Discovery of New Elements

  • DMITRY MENDELEEV (1834-1907)
    • In 1870, arranged known elements by increasing atomic mass to reveal periodic recurrences of chemical properties.
    • Created early periodic table and left gaps for undiscovered elements.
  • HENRY GWYN JEFFREYS MOSELY (1887-1915)
    • In 1913, developed a version of the periodic table based on atomic number rather than atomic mass.
    • Found that the atomic structure relates to atomic number via X-ray wavelengths.
  • PERIODIC LAW
    • States that there is a periodic repetition of physical and chemical properties when elements are arranged in order of increasing atomic number $Z$.

Nuclear Reactions vs. Chemical Reactions

  • Nuclear reactions differ from chemical reactions because they involve protons, neutrons (and sometimes electrons) rather than just electron rearrangements.
  • Guidelines for writing nuclear reactions:
    • Fundamental particles involved must be written correctly.
    • Reactants are on the left; products are on the right; the arrow denotes transformation.
    • The sum of mass numbers $A$ and the sum of atomic numbers $Z$ must balance on both sides (conservation of matter).
  • Balancing tips:
    • Coefficients are placed before formulas/symbols to balance $A$ and $Z$.
    • Verify that the equation is balanced by checking net $Z$, $A$, and charge on both sides.

Sample Problem (Rutherford’s first nuclear reaction)

  • Problem: Nitrogen-14 bombarded by alpha particles to produce Oxygen-17. What particle balances the equation?
    • Reaction: $^{14}{7} ext{N} + ^{4}{2} ext{He}
      ightarrow ^{17}{8} ext{O} + ^{1}{1} ext{H}$
    • The balancing particle is a proton $^{1}_{1} ext{H}$ (since $A$ total is $14+4=18$ on both sides and $Z$ total is $7+2=9$ on left; $8+1=9$ on right).

Subatomic Particles: Quick Reference

  • Protons: $p$, charge $+1$, location in nucleus, mass ~ $1$ amu, discovered via Goldstein context (proton concept).
  • Neutrons: $n$, charge $0$, location in nucleus, mass ~ $1$ amu, discovered by James Chadwick.
  • Electrons: $e$, charge $-1$, location outside nucleus, mass ~ 11840\dfrac{1}{1840} amu, discovered by J. J. Thomson.

Extraneous Content in the Slides

  • Some slides discuss planetary topics (Mercury, Venus, Mars, Jupiter, Neptune) and planetary facts unrelated to atomic theory.
  • Examples mentioned include:
    • Mercury: closest to the Sun, smallest planet; misinterpretation of its name; rotation period ~ 58 days; extreme temperatures; non-habitable; naming origin.
    • Venus and Mars described with various planetary traits.
    • General statements on planets and solar system features appearing in slides 31–39, which are not part of atomic theory content.

Miscellaneous/Attributions (not central to the atomic theory content)

  • Credits: Slidesgo template, icons by Flaticon, images by Freepik (attribution shown in the presentation).

Notes on significance and connections:

  • The arc from early Greek thought and alchemy to modern atomic theory shows evolving ideas about matter, structure, and the methods of science (speculation, observation, experimentation).
  • The development of atomic theory connects experimental measurements (mass, definite proportions, conservation laws) with theoretical models (Dalton, Thomson, Rutherford, Bohr) and later refinements (Goldstein, Chadwick).
  • The periodic table emerged from systematic organization of elements by properties and later by atomic number, revealing underlying order in matter.
  • Nuclear reactions introduce a framework for understanding changes in the nucleus, beyond electron interactions, with conservation laws playing a central role.