From Dalton to Rutherford: Evidence for the Nuclear Atom

How do you study something you cannot see?

Nobody has ever looked at an atom the way you look at a maple leaf. Everything we know about the inside of an atom comes from indirect evidence: firing things at atoms, heating them, shining light on them, and then asking which model best explains the results.

In this unit you will follow that story from simple balls to the quantum mechanical model that chemists use today, and then use it to explain the periodic table, the shapes of molecules and the properties of every kind of solid.

Remember from earlier grades: in Grade 9 you drew Bohr-Rutherford diagrams, and in Grade 11 Chemistry (SCH3U, Matter, Chemical Trends, and Chemical Bonding) you used protons, neutrons and electrons, atomic and mass numbers, isotopes, and the electron arrangements of the first 20 elements. This lesson asks why scientists believe atoms are built that way.

Dalton: atoms as solid spheres (1803 to 1808)

John Dalton, an English teacher and chemist, proposed that:

1. All matter is made of tiny, indivisible particles called atoms.
2. All atoms of one element are identical in mass and properties; atoms of different elements differ.
3. Compounds form when atoms combine in fixed, whole-number ratios.
4. In a chemical reaction atoms are rearranged, not created or destroyed.

Dalton's theory explained the laws of conservation of mass and definite proportions. Two of his ideas did not last: atoms can be divided (they contain smaller particles), and atoms of one element are not all identical in mass (isotopes, which you met in Grade 11).

Thomson: the electron (1897)

J. J. Thomson studied cathode rays, the glowing beams inside sealed glass tubes with most of the air pumped out. He found that the rays bent towards a positive plate and in a magnetic field, so they were made of negatively charged particles. By measuring how much they bent, he found their charge-to-mass ratio, and it was the same whatever metal the electrodes were made of. These particles were far lighter than any atom. They were later called electrons.

Since atoms are neutral, Thomson pictured the atom as a sphere of positive charge with electrons scattered through it, often called the "plum pudding" model (Canadians might picture blueberries in a muffin). In 1909 Robert Millikan's oil-drop experiment measured the charge on one electron: about 1.60 × 10⁻¹⁹ C. Combined with Thomson's ratio, this gave the electron's mass: about 9.11 × 10⁻³¹ kg, roughly 1/1836 of the mass of a hydrogen atom.

Rutherford: the nuclear atom (1911)

Ernest Rutherford had already won the 1908 Nobel Prize in Chemistry for his work on radioactivity, much of it done at McGill University in Montréal. Later, in Manchester, his assistants Hans Geiger and Ernest Marsden fired alpha (α) particles (small, fast, positively charged particles from a radioactive source) at gold foil only a few hundred atoms thick, and watched for flashes on a glowing screen.

Observations:

• Most alpha particles went straight through the foil.
• A few were deflected through small angles.
• A very small fraction bounced back towards the source.

Rutherford said it was "almost as incredible as if you fired a 15-inch shell at a piece of tissue paper and it came back and hit you." With Thomson's spread-out positive charge, every alpha particle should have gone almost straight through.

Inferences:

• Most of the atom is empty space (most particles passed through).
• The positive charge and almost all of the mass are concentrated in a tiny, dense nucleus (only a direct hit on something small, heavy and positive can push an alpha particle back).
• Electrons move in the space around the nucleus.

The nucleus is about 10⁻¹⁵ to 10⁻¹⁴ m across, while the atom is about 10⁻¹⁰ m across. Rutherford later identified the proton (1919), and James Chadwick discovered the neutron in 1932, which explained why atoms are heavier than their protons alone.

The problem with Rutherford's model

Rutherford pictured electrons orbiting the nucleus like planets around the Sun. Physicists immediately saw two problems:

1. Collapse. According to classical physics, a charged particle moving in a circle is accelerating and should give off electromagnetic radiation, lose energy and spiral into the nucleus in a tiny fraction of a second. Real atoms are stable.
2. Spectra. A spiralling electron would give off a smooth rainbow of all colours. But heated hydrogen gives off only a few specific colours of light (a line spectrum).

These failures led Niels Bohr to a new model in 1913, which is the subject of the next lesson. Notice the pattern: the model changed because evidence demanded it.