Introduction to Atomic Theory and Subatomic Structure

Fundamental Laws of Chemical Composition

Law of Conservation of Mass:
- States that matter/mass can neither be created nor destroyed in a chemical reaction.
- The total mass remaining after a reaction is identical to the total mass present before the reaction occurred.
Law of Definite Proportion (Law of Definite Composition):
- Formulated by Joseph Clue in 1799.
- States that all samples of a given chemical compound have the exact same elemental composition by mass, regardless of sample size or source.
- If a small sample, a larger sample, and a much larger sample of the same compound are analyzed, each sample will yield identical percentage mass compositions for every constituent element.
- Elemental Composition Example (Water):
  - Water is composed of $88.8\%$ oxygen and $11.2\%$ hydrogen by mass.
  - State Independence: The elemental mass composition of water remains fixed at $88.8\%$ oxygen and $11.2\%$ hydrogen regardless of whether it exists as solid ice, liquid water, or gaseous steam.
Law of Multiple Proportions:
- Applies when two or more distinct chemical compounds are formed from the same constituent elements.
- States that if a fixed mass of one element combines with varying masses of a second element across different compounds, the masses of the second element will relate to one another in ratios of small whole numbers (e.g., $2:1$ , $3:2$ , $3:1$ , $4:3$ ).
- Detailed Comparison Example (Carbon Monoxide vs. Carbon Dioxide):
  - Consider samples of carbon monoxide ( $\text{CO}$ ) and carbon dioxide ( $\text{CO}_2$ ), fixing the mass of carbon at $12\,g$ in both samples.
  - In carbon monoxide, $12\,g$ of carbon combines with $16\,g$ of oxygen.
  - In carbon dioxide, $12\,g$ of carbon combines with $32\,g$ of oxygen.
  - Taking the ratio of the oxygen masses yields $\frac{32\,g}{16\,g} = \frac{2}{1}$ , establishing a simple whole-number ratio of $2:1$

John Dalton and the First Atomic Theory

Context and Theoretical Basis:
- John Dalton published the first formal atomic theory, synthesizing previously presented and accepted natural laws.
- A scientific theory provides an overarching explanation of why natural phenomena occur based on observed physical laws, whereas scientific laws merely describe what occurs.
Dalton's Initial Physical Picture of the Atom:
- Visualized the atom as a solid sphere of mass.
- Posited that these mass spheres were indivisible and could not be broken apart.
- Hypothesized that atomic spheres interact by colliding; upon collision, they either bounce off one another or adhere together to form a new substance (unlike macroscopic billiard balls, which bounce without adhering or transforming).
The Four Main Postulates of Dalton's Atomic Theory and Modern Modifications:
1. Postulate 1: Each chemical element is composed of tiny particles called atoms.
  - Dalton's Original Phrasing: Explicitly defined these particles as indivisible fundamental units.
  - Modern Revision: While the atom remains the fundamental unit in chemistry, modern physics shows that atoms are divisible and composed of subatomic particles.
2. Postulate 2: Atoms of a given element are completely identical, whereas atoms of different elements differ in fundamental ways.
  - Historical Context: Atomic mass was widely considered the fundamental differentiating factor between elements.
  - Modern Revision: Modern chemistry recognizes that atoms of the same element can differ in mass (e.g., isotopes).
3. Postulate 3: Chemical compounds are formed when atoms of different elements combine with one another. A given compound always contains the same relative number and types of atoms.
  - Relationship to Observed Laws: Provides the theoretical explanation for the Law of Definite Proportion by translating fixed mass ratios into fixed atomic numerical ratios.
4. Postulate 4: Chemical reactions involve the reorganization of atoms—specifically changing the ways in which they are bound together. The individual atoms themselves are neither created, destroyed, nor transformed in a chemical reaction.
  - Relationship to Observed Laws: Direct microscopic explanation for the macroscopic Law of Conservation of Mass.

Gas Volume Relationships and Particle Theory

Avogadro's Hypothesis:
- Proposed by Alba Lautau to explain observed volumetric ratios in gas-phase reactions.
- Hypothesized that equal volumes of different gases measured at the same temperature and pressure contain identical numbers of particles.
Volumetric Gas Examples:
- Water Synthesis: $2$ volumes of hydrogen gas combine with $1$ volume of oxygen gas to produce $2$ volumes of water vapor ( $2 + 1 = 2$ in volumetric terms).
- Hydrogen Chloride Synthesis: $1$ volume of hydrogen gas combines with $1$ volume of chlorine gas to yield $2$ volumes of hydrogen chloride ( $\text{HCl}$ ) gas.

Discovery of Subatomic Particles and Radiation

Cathode Rays and the Electron (J.J. Thomson):
- Studied electrical discharges in cathode ray tubes.
- Demonstrated that cathode rays are streams of negatively charged particles (electrons) emitted from metallic surfaces (such as silver or gold).
- Proved that atoms are not indivisible solid spheres but contain internal negatively charged subatomic particles.
- Determined the fundamental mass-to-charge ratio of the electron.
Charge and Mass Quantification of the Electron (Robert Millikan):
- Conducted the oil drop experiment, imparting negative electrical charges onto fine oil droplets using electrons.
- Suspended the charged oil droplets within an adjustable electric and magnetic field to balance gravitational forces.
- Calculated the precise magnitude of the electric charge on an electron, enabling the determination of its exact physical mass.
X-Rays and Energy Emissions:
- Discovered by Wilhelm Röntgen (referenced historically as Billy Helms' Linskin / Bill Milton Rincken).
- Produced the first human X-ray image, capturing the skeletal structure and wedding ring of his wife's hand.
- Proved that atomic processes can emit high-energy radiation capable of penetrating soft human tissue.
Radioactivity (Murray Curie):
- Coined the term radioactivity to describe spontaneous nuclear emissions.
- Categorized radioactive emissions into high-energy rays, negatively charged particle streams, and positively charged particle streams.
The Gold Foil Experiment and Nuclear Structure (Ernest Rutherford, Beiger, and Martin):
- Experimental Setup: Directed positively charged radioactive alpha particles at an extremely thin sheet of gold foil.
- Research Team: Conducted alongside postdoctoral researcher Beiger and undergraduate researcher Martin.
- Initial Hypothesis: Based on Thomson's model (which assumed positive charge was evenly dispersed throughout the atom), alpha particles were expected to pass directly through the foil unimpeded.
- Experimental Results:
  - The vast majority of alpha particles passed straight through the foil without deflection.
  - A small fraction of particles were deflected at very wide angles.
  - A tiny fraction of particles bounced directly back toward the emission source.
- Rutherford's Theoretical Deductions:
  - Disproved Thomson's dispersed model, concluding that all positive charge and nearly all atomic mass are concentrated in a dense central core called the nucleus.
  - Undeviated particles pass through the vast empty space surrounding the nucleus.
  - Deflected particles experience strong electrostatic repulsion when passing close to the concentrated positive core.
  - Back-scattered particles collide directly with the dense positive nucleus.
Prediction and Discovery of the Neutron:
- Rutherford determined that the measured positive charge and mass of heavy elements did not balance using protons alone.
- Postulated the existence of an uncharged, massive neutral subatomic particle within the nucleus.
- James Chad formally discovered the neutron $20$ to $30$ years after Rutherford's nuclear hypothesis.

Evolution of Atomic Models

Billiard Ball Model (John Dalton):
- Pictures the atom as an indivisible, uniform solid sphere of mass with no internal subatomic components.
Plum Pudding Model (J.J. Thomson):
- Visualizes the atom as a diffuse sphere of uniform positive charge with negatively charged electrons distributed throughout like berries or raisins embedded in plum pudding.
Nuclear Model (Ernest Rutherford):
- Features a dense, positively charged nucleus at the exact center containing almost all atomic mass, surrounded by negatively charged electrons occupying the external space.
Quantum / Orbital Models:
- Subsequent models refined Rutherford's structure by defining specific discrete electron energy levels and spatial locations around the nucleus.

Properties of Subatomic Particles

Locations within the Atom:
- Electrons: Located outside the central nucleus.
- Protons: Located inside the central nucleus.
- Neutrons: Located inside the central nucleus.
Mass Relationships:
- Proton and neutron masses are nearly identical, differing by only $2$ parts in approximately $1700$ ( $1673$ mass units vs. $1675$ mass units).
- An electron is roughly $2000$ times lighter than a proton or neutron.
- Relative Atomic Masses:
  - Proton = $1$
  - Neutron = $1$
  - Electron = $0$ (negligible relative mass)
Electrical Charges:
- Absolute Charge Magnitude: $\pm 1.602 \times 10^{-19}\,\text{C}$ (coulombs) for protons and electrons.
- Relative Charges:
  - Proton = $+1$
  - Electron = $-1$
  - Neutron = $0$ (electrically neutral)

In-Class Assignment and Student Discussion

Assignment Details:
- Consists of $4$ conceptual questions on early atomic history and subatomic structure.
- Submission formats: Digital upload via Canvas or physical paper submission.
- Collaboration options: Individual completion or small group work (groups turn in a single document with all names listed at the top for equal grading).
Student Discussion Points:
- Students discuss creating a shared document for online submission without retyping prompt questions.
- Clarifying the requirements for compare-and-contrast analysis questions.