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The story of the atom is one of the most remarkable journeys in the history of science. Over barely a century, our understanding shifted from "atoms are indivisible spheres" to "atoms are mostly empty space with a dense nucleus surrounded by probability clouds of electrons." Each model built on the last, correcting its failures while introducing new questions.
John Dalton proposed that all matter is made of indivisible atoms — tiny, solid spheres that cannot be created or destroyed. Each element is composed of identical atoms, and compounds form when atoms of different elements combine in fixed ratios.
Dalton's model explained the law of conservation of mass and the law of definite proportions, but it assumed atoms had no internal structure. There were no subatomic particles in Dalton's universe.
Key limitation: Dalton could not explain electrical phenomena or the existence of isotopes.
J.J. Thomson discovered the electron using cathode ray tubes. He measured the charge-to-mass ratio of these particles and showed they were much lighter than atoms. Since atoms are electrically neutral, Thomson proposed that the atom was a sphere of positive charge with electrons embedded in it — like plums in a pudding.
Key features:
Key limitation: Thomson's model could not explain the results of Rutherford's gold foil experiment.
Ernest Rutherford directed alpha particles at a thin sheet of gold foil. Most passed straight through, but a small fraction were deflected at large angles, and a very few bounced straight back.
This was completely incompatible with Thomson's model. If the positive charge were spread out evenly, no alpha particles would bounce back. Rutherford concluded:
Key limitation: According to classical physics, orbiting electrons should continuously emit electromagnetic radiation, lose energy, and spiral into the nucleus. Rutherford's model could not explain why atoms are stable.
Niels Bohr modified Rutherford's model by proposing that electrons exist in fixed energy levels (shells) around the nucleus. Electrons can only occupy specific orbits and can move between them by absorbing or emitting a precise quantum of energy.
Key features:
This model successfully explained the line spectrum of hydrogen — the specific wavelengths of light emitted by hydrogen atoms.
Key limitation: Bohr's model worked well for hydrogen but failed for multi-electron atoms. It could not explain the fine structure of spectral lines or the shapes of orbitals.
The modern model, developed by Schrödinger, Heisenberg, and others, treats electrons not as particles in fixed orbits but as wave-like entities described by mathematical functions called wavefunctions. The square of the wavefunction gives the probability of finding an electron in a particular region of space — an orbital.
Key features:
All atoms are built from three fundamental particles:
| Particle | Relative Mass | Relative Charge | Location |
|---|---|---|---|
| Proton | 1 | +1 | Nucleus |
| Neutron | 1 | 0 | Nucleus |
| Electron | 1/1836 (≈ 0) | −1 | Orbitals around nucleus |
Isotopes are atoms of the same element (same number of protons) that have different numbers of neutrons. For example, chlorine has two stable isotopes:
Both have identical chemical properties because they have the same electron configuration. They differ in mass and in nuclear stability.
Because elements often exist as a mixture of isotopes, we define the relative atomic mass as the weighted mean mass of an atom of the element relative to 1/12 the mass of a carbon-12 atom.
Chlorine exists as 75.0% ³⁵Cl and 25.0% ³⁷Cl. Calculate the relative atomic mass.
Ar = (75.0/100 × 35) + (25.0/100 × 37) Ar = 26.25 + 9.25 Ar = 35.5
This is why chlorine's relative atomic mass is 35.5 rather than a whole number — it reflects the natural mixture of isotopes.
Ar = Σ (fractional abundance × isotopic mass)
If given percentages, divide each by 100 before multiplying. If given peak heights or ratios from a mass spectrum, use those as the relative abundances and divide by the total.