Thomson's plum-pudding model
J. J. Thomson, who had identified the electron in 1897, proposed in 1898 the first concrete model of how charge is arranged within an atom. He pictured the atom as a sphere of radius approximately $10^{-10}\ \text{m}$ in which the positive charge is distributed uniformly, with the electrons embedded into it in the arrangement that gives the most stable electrostatic configuration. Because the negative electrons are evenly seeded through a continuous positive cloud, the atom as a whole is electrically neutral.
The model is known by several food-based names — the plum-pudding model, the raisin-pudding model, or the watermelon model. The image is of a pudding or watermelon of positive charge with the electrons sitting in it like plums or seeds. A point that NEET frequently tests is that, in this model, the mass of the atom is assumed to be spread uniformly over the whole atom rather than concentrated anywhere.
Figure 1. Thomson's atom: electrons (the "plums") set into a continuous, uniformly distributed sphere of positive charge. Mass and positive charge are spread evenly, so the model could explain neutrality but not later scattering data.
Thomson's model successfully explained the overall electrical neutrality of the atom and was an important conceptual advance, earning him the 1906 Nobel Prize in Physics for his work on the conduction of electricity by gases. However, it could not survive the experimental tests that followed. The decisive evidence came from a study of how positively charged α-particles behave when fired at thin metal foil.
It is worth being precise about what the model did and did not commit to. By insisting that both the positive charge and the mass were smeared out over the whole sphere, Thomson left the atom with no dense interior anywhere — there was no region capable of exerting a large, sudden force on a fast-moving projectile. This is the feature that the scattering experiment was about to test, because a diffuse charge distribution and a concentrated one make sharply different predictions for how an incoming particle is deflected.
Rutherford's α-scattering experiment
By the early twentieth century, radioactivity had been discovered by Henri Becquerel and developed by Marie Curie, Pierre Curie, Rutherford, and Frederick Soddy. Three kinds of emission — $\alpha$, $\beta$, and $\gamma$ rays — were known. Rutherford found that α-rays consist of high-energy particles carrying two units of positive charge and four units of atomic mass, and concluded they are helium nuclei: combining $\alpha$-particles with two electrons yields helium gas, $\ce{He^2+ + 2e^- -> He}$.
Rutherford and his students Hans Geiger and Ernest Marsden bombarded a very thin gold foil (thickness about $100\ \text{nm}$) with a stream of high-energy α-particles from a radioactive source. A circular fluorescent zinc sulphide screen surrounded the foil, so that wherever an α-particle struck the screen, a tiny flash of light was produced and could be counted. The geometry let the experimenters record both how many particles passed straight through and how many were deflected to large angles.
Figure 2. Schematic of the gold-foil experiment. The overwhelming majority of α-particles pass through undeflected (teal); a few are deflected through small angles (purple); about one in twenty thousand rebound through nearly $180^\circ$ (coral).
Observations and conclusions
On Thomson's model the gold atom's mass and positive charge would be spread evenly across each atom, so the energetic α-particles should pass through this diffuse distribution with at most small, gentle changes of direction. The results, however, were quite unexpected and are summarised below alongside the inference Rutherford drew from each.
| Observation | Conclusion drawn |
|---|---|
| Most α-particles passed through the foil undeflected. | Most of the space inside the atom is empty. |
| A small fraction of α-particles were deflected through small angles. | The positive charge is not spread throughout the atom (as Thomson presumed) but is concentrated in a very small volume that repels the positive α-particles. |
| A very few α-particles (about 1 in 20,000) bounced back, deflected by nearly $180^\circ$. | An enormous repulsive force acted on these particles, so the dense positive core (the nucleus) occupies a negligibly small volume compared with the whole atom. |
The size contrast that follows from the third conclusion is dramatic. The radius of the atom is about $10^{-10}\ \text{m}$, while the radius of the nucleus is about $10^{-15}\ \text{m}$. NCERT offers a memorable scaling: if a cricket ball represented the nucleus, the atom would have a radius of about 5 km. Almost all of the atom is empty space, with its mass and positive charge packed into a core five orders of magnitude smaller in radius.
Do not say "most particles bounced back"
The headline result is that most α-particles passed straight through — that is what proved the atom is mostly empty. Only a very few (about 1 in 20,000) bounced back, and it was these rare large-angle deflections that revealed the tiny dense nucleus. Confusing the frequency of each outcome inverts the entire argument.
Frequent + undeflected → empty space. Rare + back-scattered → concentrated positive nucleus.
Rutherford's nuclear model
From these observations and conclusions Rutherford proposed the nuclear model of the atom. Three features define it. First, the positive charge and most of the mass of the atom are densely concentrated in an extremely small region that Rutherford named the nucleus. Second, the nucleus is surrounded by electrons that revolve around it at very high speed in circular paths called orbits. Third, the electrons and the nucleus are held together by electrostatic forces of attraction.
The picture deliberately resembles the solar system: the nucleus plays the role of the massive Sun while the electrons behave like the lighter, orbiting planets. This planetary analogy is intuitive, but as the closing section shows, it is precisely the orbital motion of the charged electron that the classical model could not reconcile with atomic stability.
Figure 3. Rutherford's nuclear atom. Almost all the mass and the entire positive charge sit in a minute central nucleus (radius $\sim 10^{-15}\,\text{m}$); electrons revolve in circular orbits within a region $\sim 10^{-10}\,\text{m}$ across. Drawn far larger than scale.
Thomson vs Rutherford compared
The two models differ on every structural question that matters: where the positive charge sits, where the mass sits, and whether the electrons are static or moving. The table contrasts them directly.
| Feature | Thomson (plum-pudding) | Rutherford (nuclear) |
|---|---|---|
| Distribution of positive charge | Spread uniformly through the whole sphere | Concentrated in a tiny central nucleus |
| Distribution of mass | Spread uniformly over the atom | Almost entirely in the nucleus |
| Electrons | Embedded (static) in the positive cloud | Revolving in circular orbits around the nucleus |
| Empty space | Effectively none; atom is filled charge | Most of the atomic volume is empty |
| Explains scattering data? | No — predicts only small deflections | Yes — explains large-angle back-scattering |
| Analogy | Plum pudding / watermelon | Solar system (Sun and planets) |
The instability flaw of Rutherford's model and the spectral evidence that follows feed straight into developments leading to the Bohr model.
Atomic number and mass number
With a nucleus established, the identity of an atom can be defined by counting its nucleons. The positive charge of the nucleus is due to the protons, whose charge is equal and opposite to that of the electron. The atomic number $Z$ is the number of protons in the nucleus, and for electrical neutrality it also equals the number of electrons in a neutral atom.
$$Z = \text{number of protons} = \text{number of electrons (in a neutral atom)}$$For example, hydrogen has $Z = 1$ and sodium has $Z = 11$. The mass of the nucleus is due to both protons and neutrons, which are collectively called nucleons. The total number of nucleons is the mass number $A$.
$$A = \text{number of protons }(Z) + \text{number of neutrons }(n)$$The composition of any atom is written using the element symbol with the mass number as a left superscript and the atomic number as a left subscript, $\ce{^{A}_{Z}X}$. The number of neutrons is always $A - Z$.
Calculate the number of protons, neutrons, and electrons in $\ce{^{80}_{35}Br}$, and assign the symbol to a species with 18 electrons, 16 protons, and 16 neutrons.
For $\ce{^{80}_{35}Br}$: $Z = 35$, $A = 80$, neutral species, so protons $=$ electrons $= 35$ and neutrons $= 80 - 35 = 45$.
For the second species: protons $= 16$, so the element is sulphur (S). Mass number $A = 16 + 16 = 32$. Electrons (18) exceed protons (16) by 2, so it is an anion of charge $2-$: $\ce{^{32}_{16}S^{2-}}$.
Charge changes electrons, not protons or neutrons
When a species is an ion, decide whether protons outnumber electrons (cation) or electrons outnumber protons (anion) by the magnitude of the charge. The number of neutrons is always $A - Z$, and the number of protons is always $Z$ — neither depends on the charge.
Neutrons $= A - Z$ always · electrons $= Z -$ (charge) for a cation, $Z +$ (charge) for an anion.
Isotopes and isobars
Two classification terms follow naturally from the $\ce{^{A}_{Z}X}$ notation. Isotopes are atoms with the same atomic number but different mass numbers; they differ only in the number of neutrons in the nucleus. Isobars are atoms with the same mass number but different atomic numbers, so they belong to different elements.
| Class | Same | Different | Examples |
|---|---|---|---|
| Isotopes | Atomic number $Z$ | Mass number $A$ (i.e. neutron count) | $\ce{^{1}_{1}H}$, $\ce{^{2}_{1}H}$, $\ce{^{3}_{1}H}$ · $\ce{^{12}_{6}C}$, $\ce{^{13}_{6}C}$, $\ce{^{14}_{6}C}$ · $\ce{^{35}_{17}Cl}$, $\ce{^{37}_{17}Cl}$ |
| Isobars | Mass number $A$ | Atomic number $Z$ (different elements) | $\ce{^{14}_{6}C}$ and $\ce{^{14}_{7}N}$ |
Hydrogen is the textbook illustration of isotopy: 99.985% of hydrogen atoms are protium $\ce{^{1}_{1}H}$ (one proton, no neutron); a small fraction are deuterium $\ce{^{2}_{1}H}$ or D (one proton, one neutron, 0.015%); and trace amounts are tritium $\ce{^{3}_{1}H}$ or T (one proton, two neutrons). Chlorine occurs as $\ce{^{35}_{17}Cl}$ and $\ce{^{37}_{17}Cl}$, and carbon as the isotopes with 6, 7, and 8 neutrons.
A central consequence is that all isotopes of a given element show the same chemical behaviour. Chemical properties are controlled by the number of electrons, which is fixed by the number of protons; the number of neutrons in the nucleus has very little effect. This is exactly why NEET statements asserting that "all isotopes of an element show the same chemical properties" are marked correct.
Drawbacks of Rutherford's model
Rutherford's model resembles a miniature solar system, with the nucleus as the massive Sun and the electrons as orbiting planets. The Coulomb force between an electron and the nucleus, $F = k\,q_1 q_2 / r^2$, is mathematically of the same inverse-square form as the gravitational force that governs the planets, so it is tempting to expect equally stable orbits. The analogy, however, breaks down on one decisive point.
A body moving in a circular orbit is constantly accelerated, even at constant speed, because its direction of motion keeps changing. According to Maxwell's electromagnetic theory, an accelerated charged particle must continuously emit electromagnetic radiation — a feature absent for the uncharged planets. An orbiting electron would therefore radiate away energy drawn from its motion, its orbit would steadily shrink, and calculations show it should spiral into the nucleus in only about $10^{-8}\ \text{s}$.
Since this collapse does not happen and atoms are in fact stable, the Rutherford model cannot explain the stability of an atom. This single failure is the launch point for the quantum picture that followed.
The model is also silent about how electrons are distributed around the nucleus and about the discrete line spectra emitted by atoms. A planetary atom radiating continuously as it collapsed would be expected to emit a continuous spread of frequencies, yet experiments showed that excited atoms emit light only at sharp, characteristic wavelengths. Rutherford's model offered no way to account for this discreteness, and no rule to fix the size of the electron's orbit. These shortcomings, especially the stability problem, motivated Niels Bohr to introduce quantised stationary orbits in which an electron does not radiate — the subject of the next subtopic.
Five lines to retain
- Thomson (1898): uniform sphere of positive charge with electrons embedded; mass spread evenly; explains neutrality only.
- Rutherford's α-scattering: most pass undeflected (empty atom), few deflected (concentrated +), ~1 in 20,000 rebound (tiny dense nucleus).
- Nuclear model: mass and positive charge in a tiny nucleus ($\sim10^{-15}$ m); electrons orbit; held by electrostatic attraction.
- Counting: $Z =$ protons $=$ electrons (neutral); $A = Z + n$; neutrons $= A - Z$ always; isotopes share $Z$, isobars share $A$.
- Fatal flaw: an orbiting electron is accelerated, must radiate (Maxwell), and would spiral in within $\sim10^{-8}$ s — cannot explain atomic stability.