Atomic Structure and Rutherford's Gold Foil Experiment

Foundations of Modern Chemistry and the Atomic Theory

• Historical Evolution of Matter Conceptions:

  • Philosophical Foundations vs. Empirical Science: Democritus conceived an early atomic idea based entirely on philosophical speculation, abstract logic, and reasoning without experimental verification. In contrast, John Dalton formulated a scientific theory grounded in systematic quantitative measurements and experimental data.
  • Pre-1790s Chemical Paradigms: Before the late 1700s, natural philosophers recognized that elements represent fundamental substances that cannot be broken down further, and that elements combine to form compounds possessing chemical and physical properties entirely distinct from their constituent elements (for example, carbon monoxide, $\text{CO}$, exhibits completely different properties from pure carbon and pure oxygen gas). However, early investigators lacked the theoretical foundation to understand the underlying mechanisms governing why and how these transformations occurred.
  • Late Eighteenth-Century Technological Advances: During the 1790s, the introduction of improved analytical technology—notably precision balances and refined microscopes—shifted chemistry from qualitative philosophical notions to rigorous, quantitative scientific measurement, giving rise to modern chemistry through the establishment of basic physical laws.

• Fundamental Laws of Chemical Combination:

  • Law of Conservation of Mass: Mass is neither created nor destroyed during ordinary chemical reactions or physical changes. The total mass of the reactants must equal the total mass of the products, representing the combination, separation, or rearrangement of indestructible atomic units.
  • Law of Definite Proportions: A specific chemical compound always contains exactly the same elements in the exact same proportions by mass, regardless of the size of the sample or the origin of the source.
  • Example: Pure baking soda ($\text{NaHCO}_3$) consistently exhibits a constant percentage composition by mass: $27.37\%$ sodium, $1.20\%$ hydrogen, $14.30\%$ carbon, and $57.14\%$ oxygen.
  • Law of Multiple Proportions: If two or more different compounds are composed of the same two elements, the ratio of the masses of the second element that combine with a constant, fixed mass of the first element is always expressed as a ratio of small whole numbers.
  • Example: Nitrogen oxides ($\text{NO}$ and $\text{NO}_2$). In $\text{NO}_2$, the mass ratio of oxygen to nitrogen is $32\,\text{g O} / 14.01\,\text{g N} = 2.28$, demonstrating that the ratio of oxygen atoms combining with a fixed quantity of nitrogen in differing compounds reduces to small whole numbers (such as $1:2$).
  • Example: Carbon oxides ($\text{CO}$ and $\text{CO}_2$). The mass of oxygen combining with a fixed mass of carbon in carbon dioxide is exactly twice that found in carbon monoxide, yielding a whole-number ratio of $1:2$.
  • Example: Synthesis Stoichiometry. When $2$ carbon atoms combine with $4$ oxygen atoms, exactly $2$ carbon dioxide ($\text{CO}_2$) molecules are produced, because each $\text{CO}_2$ molecule requires exactly $1$ carbon atom and $2$ oxygen atoms ($2 \times 1 = 2\,\text{C}$, $2 \times 2 = 4\,\text{O}$), leaving zero unreacted atoms.

• Dalton's Atomic Theory (1808):

  • Postulate 1: All matter is composed of extremely small particles called atoms.
  • Postulate 2: Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties.
  • Postulate 3: Atoms cannot be subdivided, created, or destroyed.
  • Postulate 4: Atoms of different elements combine in simple whole-number ratios to form chemical compounds.
  • Postulate 5: In chemical reactions, atoms are combined, separated, or rearranged.

• Modern Revisions and Modifications to Dalton's Theory:

  • Divisibility of the Atom: Dalton postulated that atoms are indivisible particles that cannot be subdivided. Modern atomic theory confirms that atoms are composed of smaller subatomic particles: electrons, protons, and neutrons.
  • Non-Identity of Same-Element Atoms: Dalton postulated that all atoms of a given element possess identical masses. The discovery of isotopes demonstrated that atoms of the same element can possess different masses due to variable quantities of neutrons.
  • Elemental Identity Criterion: Dalton attributed elemental identity to the atom as an indivisible whole. Modern chemistry recognizes that the identity of an element is determined exclusively by