Discovery of Electrons
Thomson's cathode ray experiments and Millikan's oil drop experiment revealed the electron, the first subatomic particle ever identified. These discoveries shattered the long-held idea that atoms were indivisible and set the stage for uncovering the internal structure of the atom.
Thomson's Cathode Ray Experiments
Thomson studied cathode rays, which are produced by applying high voltage between electrodes inside a vacuum tube. He noticed that these rays were deflected by both electric and magnetic fields, which meant they had to be negatively charged particles with mass (neutral rays wouldn't deflect).
By measuring how much the rays bent in known fields, Thomson determined the charge-to-mass ratio () of these particles. The value he found was about 1,000 times larger than that of hydrogen ions, the lightest known charged particles at the time. That huge ratio meant one of two things: either the charge was enormous, or the mass was tiny. Thomson concluded it was the latter.
He called these particles "corpuscles" (later renamed electrons) and argued they were components of all atoms. This was a direct challenge to the prevailing view that atoms were the smallest, indivisible units of matter.
Millikan's Oil Drop Experiment
While Thomson found the charge-to-mass ratio, Millikan's experiment pinpointed the actual charge of a single electron. Here's how it worked:
- Tiny oil droplets were sprayed into a chamber between two parallel metal plates.
- An X-ray source ionized air molecules, which attached to the droplets and gave them a negative charge.
- An electric field was applied between the plates, pushing the negatively charged droplets upward against the downward pull of gravity.
- Millikan adjusted the electric field strength until individual droplets hovered motionless, meaning the upward electric force exactly balanced the downward gravitational force.
- Using the known electric field strength, the mass of the droplet (calculated from its size and the density of oil), and the droplet's velocity, he calculated the charge on each droplet.
The key result: the charges on different droplets were always whole-number multiples of a single fundamental value. That value turned out to be the charge of one electron:
This was the first accurate measurement of the electron's charge and confirmed that electric charge is quantized, meaning it only comes in discrete packets.
Discovery of the Atomic Nucleus
Rutherford's Gold Foil Experiment
At the time of this experiment (1911), the accepted picture of the atom was Thomson's "plum pudding" model: a sphere of uniformly spread positive charge with electrons embedded throughout, like raisins in a pudding. If this model were correct, alpha particles fired at atoms should pass through with only slight deflections, since the positive charge would be too spread out to exert a strong force.
Rutherford tested this by firing alpha particles (positively charged helium nuclei) at a very thin sheet of gold foil and observing where they ended up. The results were striking:
- Most alpha particles passed straight through with little or no deflection, as expected.
- A small fraction deflected at large angles (greater than 90°).
- An even smaller fraction bounced almost straight back toward the source.
The large deflections were impossible to explain with the plum pudding model. Rutherford realized that the only way an alpha particle could bounce backward was if it encountered a concentration of positive charge and mass dense enough to repel it strongly. His conclusion:
The positive charge and nearly all the mass of an atom are packed into a tiny, dense core at the center: the nucleus. The rest of the atom is mostly empty space, which is why most alpha particles sail right through.
This became the Rutherford (nuclear) model of the atom.
Rutherford's Model vs. Previous Atomic Models
| Model | Key Ideas |
|---|---|
| Dalton's model (early 1800s) | Atoms are the smallest, indivisible units of matter. Different elements have atoms with different masses and properties. |
| Thomson's plum pudding model (early 1900s) | Atoms are spheres of positive charge with negative electrons embedded throughout. The charges balance, making the atom neutral. |
| Rutherford's planetary model (1911) | A small, dense, positively charged nucleus sits at the center. Electrons orbit the nucleus at relatively large distances, like planets orbiting the sun. Most of the atom's volume is empty space. |
Key features of Rutherford's model:
- It explained the gold foil results, including the rare but dramatic backscattering of alpha particles.
- It introduced the concept of a compact, positively charged nucleus that contains most of the atom's mass but occupies a tiny fraction of its volume.
- It established that atoms are mostly empty space.
Limitations of Rutherford's model:
- Classical physics predicts that an orbiting electron should continuously radiate energy and spiral into the nucleus, so the model couldn't explain why atoms are stable.
- It didn't account for the discrete energy levels of electrons, which were later explained by quantum mechanics (Bohr's model and beyond).
Subatomic Particles and Nuclear Structure
Once the nucleus was discovered, further work revealed its composition:
- Proton: A positively charged particle in the nucleus. The number of protons defines the atomic number and determines which element an atom is.
- Neutron: An electrically neutral particle in the nucleus. Neutrons contribute to the atom's mass but don't change its atomic number.
- Isotopes: Atoms of the same element (same number of protons) that have different numbers of neutrons. For example, carbon-12 and carbon-14 are both carbon, but carbon-14 has two extra neutrons.
- Radioactivity: The spontaneous emission of particles or energy from unstable nuclei. Studying radioactive decay was one of the key tools that led to the discovery of nuclear structure in the first place. Alpha particles used in Rutherford's experiment, for instance, came from a radioactive source.