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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.
timeline
title Development of Atomic Theory
1803 : Dalton's Solid Sphere Model
: Atoms are indivisible
1897 : Thomson's Plum Pudding Model
: Discovery of the electron
1911 : Rutherford's Nuclear Model
: Gold foil experiment
1913 : Bohr's Model
: Quantised energy levels
1926 : Wave-Mechanical Model
: Schrodinger equation and orbitals
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:
How Thomson measured the electron: Thomson applied electric and magnetic fields to cathode rays. By balancing the deflection from each field, he calculated the charge-to-mass ratio (e/m) of the particles. The value was the same regardless of the cathode material or the gas in the tube, proving these were fundamental particles -- not fragments of atoms.
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:
| Observation | Approximate Fraction | Conclusion |
|---|---|---|
| Passed straight through | ~99.99% | Atom is mostly empty space |
| Deflected at small angles | ~0.01% | Nucleus has positive charge, repelling alpha particles |
| Bounced back (>90 degrees) | ~1 in 20,000 | Nucleus is very small, very dense, and positively charged |
Rutherford estimated the nucleus is roughly 10,000 times smaller than the atom -- if the atom were the size of a football stadium, the nucleus would be about the size of a marble at the centre.
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.
When an electron drops from a higher energy level to a lower one, a photon is emitted with energy exactly equal to the difference between the two levels:
Delta E = E(higher) - E(lower) = hf
where h is Planck's constant and f is the frequency of the emitted photon.
| Transition Series | Drops To | Region of Spectrum |
|---|---|---|
| Lyman | n = 1 | Ultraviolet |
| Balmer | n = 2 | Visible |
| Paschen | n = 3 | Infrared |
The lines converge at higher energies because the energy levels get closer together as n increases. The convergence limit corresponds to ionisation -- removing the electron entirely.
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 Schrodinger, 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:
You cannot simultaneously know both the exact position and exact momentum of an electron. The more precisely you know one, the less precisely you can know the other. This is why we must describe electrons in terms of probability distributions (orbitals) rather than definite paths (orbits).
All atoms are built from three fundamental particles:
| Particle | Relative Mass | Relative Charge | Location | Actual Mass (kg) |
|---|---|---|---|---|
| Proton | 1 | +1 | Nucleus | 1.673 x 10^-27 |
| Neutron | 1 | 0 | Nucleus | 1.675 x 10^-27 |
| Electron | 1/1836 (approx 0) | -1 | Orbitals around nucleus | 9.109 x 10^-31 |
An atom is represented as:
A over Z, then the element symbol (e.g. carbon-12: mass number 12, atomic number 6)
The number of neutrons = A - Z.
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.
| Isotope | Application |
|---|---|
| Carbon-14 | Radiocarbon dating of archaeological specimens |
| Iodine-131 | Medical diagnosis and treatment of thyroid conditions |
| Cobalt-60 | Radiotherapy for cancer treatment |
| Uranium-235 | Nuclear fission in power stations |
| Deuterium (H-2) | NMR spectroscopy, nuclear fusion research |
Misconception: "Isotopes of the same element have different chemical properties because they have different masses."
Correction: Chemical properties depend on electron configuration, which is determined by the number of protons (and hence electrons in a neutral atom). Since isotopes have the same number of protons, they have identical electron configurations and therefore identical chemical properties. The mass difference only affects physical properties such as rate of diffusion, density, and boiling point.
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.
Carbon-12 was chosen as the standard because:
Chlorine exists as 75.0% 35-Cl and 25.0% 37-Cl. Calculate the relative atomic mass.
Ar = (75.0/100 x 35) + (25.0/100 x 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.
Lithium has two stable isotopes: 6-Li (7.59%) and 7-Li (92.41%). Calculate the relative atomic mass.
Ar = (7.59/100 x 6) + (92.41/100 x 7) Ar = 0.4554 + 6.4687 Ar = 6.9 (to 1 d.p.)
Ar = Sum of (fractional abundance x 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.
When atoms gain or lose electrons, they form ions:
| Process | Result | Example |
|---|---|---|
| Metal atom loses electrons | Positive ion (cation) | Na -> Na+ + e- |
| Non-metal atom gains electrons | Negative ion (anion) | Cl + e- -> Cl- |
When identifying a species from its particle composition:
An atom has 20 protons, 20 neutrons, and 18 electrons. Identify the species.