Chem Lecture 4

All matter is made of tiny particles called atoms. Atoms of a given element (e.g., every oxygen atom) are in fixed relationships when forming compounds.
In water, the formula is H₂O, implying a fixed ratio of hydrogen to oxygen. The correct atoms-to-atom ratio is 2 H : 1 O, which illustrates the Law of Definite Proportions, stating that a chemical compound always contains exactly the same proportion of elements by mass.
The definite composition idea: you cannot have fractional atoms in a real compound; entire atoms participate in bonding.
The law of conservation of atoms in chemical reactions: you cannot create or destroy atoms in a chemical change; atoms are just rearranged or transferred between substances.
“you cannot subdivide our atoms,” this reflects the early Daltonian idea that atoms are indivisible in chemistry (though subatomic structure was later discovered).
Dalton’s claims atoms of different elements have different masses and properties, which helps explain why compounds have fixed, whole-number ratios of elements.
Common misunderstandings about numerical ratios (e.g., 2.5 H per O). Correct interpretation: atoms combine in simple, whole-number ratios; the example water is 2:1 in atoms, not 2.5:1 or other fractional counts.
These concepts set the stage for modern chemistry by tying together atomic identity, bonding, and stoichiometry.

JJ Thomson conducted cathode ray experiments and identified a negatively charged particle inside atoms, later called the electron.
Key experimental observation: cathode rays are deflected by electric fields, indicating negatively charged constituents of atoms.
Oil drop-style reasoning (historical description): small negatively charged oil droplets could be suspended in an electric field when the electric force balances the weight of the droplet. The balance condition is:
$mg = qE$
where $m$ is the droplet’s mass, $g$ is gravity, $q$ is charge, and $E$ is the electric field.
By analyzing deflection and suspension, Thomson and contemporaries established the existence of a small, negatively charged constituent of atoms, attributed mass-to-charge properties that implied a very light particle (the electron).
Discovery pointed to positive charges inside atoms (to balance negative charges), setting the stage for the concept of a structured atom.
The historical model that Thomson proposed to explain this internal structure is the plum pudding model, where electrons were envisioned as negatively charged 'plums' embedded within a positively charged 'pudding' or soup, providing a rudimentary understanding of atomic structure before the discovery of the nucleus.

Plum pudding model (Thomson’s model): the atom is a positively charged substance with embedded electrons (like raisins in a pudding). In this picture, positive charge is distributed throughout the atom and there’s little to no empty space.
Rutherford’s alpha-particle experiment (early 1910s): he bombarded a thin metal foil (often described as gold foil) with helium nuclei (alpha particles). An alpha particle is a small, positively charged particle consisting of 2 protons and 2 neutrons (the He-4 nucleus), and a mass ~4 amu.
Most alpha particles passed straight through, some were deflected at small angles, and a tiny fraction were deflected at large angles or bounced back.
Interpretation: most of the atom is empty space; there exists a very small, dense, positively charged nucleus containing most of the atom’s mass. Electrons orbit around this nucleus, but are far outside the dense core.
The Rutherford/Nuclear model replaced the plum pudding model as the correct picture of atomic structure.
The transcript notes the expectation that alpha particles would pass through easily if the plum pudding model were correct, yet the observed scattering revealed a concentrated positive center.

James Chadwick discovered the neutron in 1932, completing the basic picture of the atomic nucleus alongside protons (positive charge) and neutrons (neutral).
Neutrons explained several puzzles, including why atomic nuclei could contain more mass than protons alone would suggest and how nuclei are held together without electrostatic repulsion tearing them apart.
The existence of neutrons helped explain isotopes and guided the development of the modern nuclear model of the atom.
Chadwick (1932) and neutron rather than a misattributed individual.

The atom is a composition of smaller parts, not a uniform blob: Electron (negative), Protons (positive), Neutrons (neutral) in the nucleus; electrons occupy surrounding space.
The mass and charge relationships in the early experiments led to quantifying electron properties and to the idea of a nucleus.
The ratio of elements in compounds is fixed (definite proportions); the ratios are typically whole-number ratios by atoms (e.g., H₂O has 2 H atoms per 1 O atom).
In chemical reactions, atoms are conserved (not created or destroyed) but electrons may be redistributed, which changes bonding and charge balance but not the total number of atoms.
The story illustrates the scientific method: hypothesis (plum pudding) → experimental test (Rutherford’s foil) → revision of the model (nuclear atom).
Distinguish accurate scientific claims from conversational or erroneous statements. The core scientific points above remain the reliable backbone.

Balancing forces in the oil-drop experiment (Millikan-Thomson era context):
- This led to determinations of the charge and, indirectly, the mass relationships of charge carriers.
Definite-proportions/stoichiometry: In a compound such as water, the atom count obeys a fixed ratio:
2 hydrogen atoms to 1 oxygen atom, illustrating that chemical compounds have a defined structure based on elemental composition.
Plum pudding vs. nuclear model: conceptual contrast rather than a numeric equation, but reflects the idea that positive charge could be distributed vs. concentrated in a nucleus.
Alpha particle description: an alpha particle has charge and consists of $2p, 2n$ (two protons and two neutrons).
Electron properties (historic context; values often taught):
- Electron charge: 0
  -1 e
Dalton (1803): Atoms are indivisible particles.
Thomson (1897): Electrons discovered → plum pudding model.
Rutherford (1909): Gold foil experiment → nucleus discovered.
Chadwick (1932): Neutron discovered → modern atomic model forms.