Atomic Models & Bohr Theory

1. Dalton (1803) — atoms are tiny, indivisible, solid spheres; different elements have different atoms.
2. Thomson (1897) — discovered the electron; proposed the "plum pudding" model, negative charges embedded in a positive sphere.
3. Rutherford (1911) — gold-foil experiment showed atoms are mostly empty space with a tiny, dense, positively charged nucleus.
4. Bohr (1913) — electrons move in fixed, quantized circular orbits (energy levels) around the nucleus; explained line spectra.
5. Quantum mechanical model (1926–present) — electrons exist in probability clouds (orbitals), not fixed paths; used for precise chemistry today.

No model was thrown out on a whim — each one was replaced because a new experiment produced a result it could not explain. Thomson's cathode-ray beam bent toward a positively charged plate, so it had to be made of negative particles, and it looked the same no matter which metal the electrodes were made of, so every atom must contain these particles (electrons). That broke Dalton's idea of an indivisible atom. Thomson then needed somewhere to put the positive charge that balances the electrons, so he spread it through the whole atom like pudding around plums. Rutherford's team tested that picture by firing positively charged alpha particles at gold foil less than a micrometre thick (on the order of a thousand atoms). If the positive charge were spread thinly, every alpha particle should punch straight through with barely a wobble.

The nucleus is roughly 100 000 times smaller across than the whole atom — if the atom were a hockey arena, the nucleus would be a pea at centre ice. Rutherford's model had one fatal flaw: an orbiting negative charge should steadily radiate energy and spiral into the nucleus in a fraction of a second, and it gave no reason why atoms emit only certain colours. Bohr fixed both problems. (The neutron, which accounts for the missing mass of the nucleus, was found by Chadwick in 1932.)

The Bohr model is the one you'll actually draw and calculate with in this course. Its big idea: an electron cannot have just any amount of energy. It can only exist in specific, allowed energy levels (shells), numbered $n=1,2,3,\dots$ outward from the nucleus. Energy is quantized — like being allowed to stand only on the rungs of a ladder, never in between.

Because every element has a unique arrangement of electrons and energy levels, every element produces a unique pattern of coloured lines when its light is passed through a prism — its line (emission) spectrum. This is effectively a fingerprint: astronomers identify the elements inside stars light-years away by reading their spectra. The size of the drop sets the colour: a big energy gap releases a high-energy photon toward the violet end of the spectrum, while a small gap releases lower-energy red light. For hydrogen, the drop from $n=3$ to $n=2$ gives red light (656 nm), while the larger drop from $n=4$ to $n=2$ gives blue-green light (486 nm). Flame tests use the same idea on a lab bench: sodium glows yellow-orange, lithium crimson, potassium lilac and copper blue-green. Bohr's model predicts hydrogen's lines exactly, but it fails for any atom with more than one electron — which is why the quantum mechanical model eventually replaced it.

To draw a Bohr–Rutherford diagram: put the nucleus (protons and neutrons, with their count) in the centre, then draw concentric rings for each occupied shell, filling from the inside out using the maximums above. For a neutral atom, the number of electrons equals the atomic number $Z$, because the atom's positive and negative charges must cancel. The neutron count comes from the mass number of the particular isotope, $A-Z$.