Physics · Ch 12 — Atoms
Introduction
Introduction
12.1 Introduction
By the close of the nineteenth century, enough experimental evidence had built up in favour of the atomic hypothesis of matter — the idea that all matter is made up of tiny, indivisible particles called atoms. A crucial breakthrough came in 1897, when the English physicist J. J. Thomson carried out experiments on electric discharge through gases. His results revealed that atoms of every element contain identical, negatively charged constituents — what we now call electrons. Yet atoms as a whole are electrically neutral, so Thomson reasoned that an atom must also contain some positive charge that exactly cancels the negative charge of its electrons.
This raised a fundamental question: how are the positive charge and the electrons actually arranged inside an atom? In other words, what is the structure of an atom?
Thomson's Plum Pudding Model
In 1898, J. J. Thomson proposed the first model of the atom to answer this question. In his picture, the positive charge of the atom is spread uniformly throughout its entire volume, and the negatively charged electrons sit embedded within it — much like seeds scattered through a watermelon. Because of this imagery, the model came to be known as the plum pudding model of the atom.
The name comes from a traditional English dessert, plum pudding, in which raisins (the "plums") are distributed throughout a soft cake. In Thomson's model, the electrons play the role of the plums and the uniformly spread positive charge plays the role of the pudding.
As this chapter will show, later experiments found that the actual distribution of electrons and positive charge inside an atom is very different from what Thomson's model proposed.
Continuous and Discrete Spectra
Condensed matter — solids and liquids — as well as dense gases, emit electromagnetic radiation at all temperatures, and this radiation contains a continuous spread of many wavelengths, though with differing intensities. Such radiation is understood to arise from the oscillations of atoms and molecules as they interact with their neighbours.
By contrast, light emitted from a rarefied (low-density) gas — heated in a flame, or excited electrically in a glow tube such as a neon sign or a mercury vapour lamp — contains only certain discrete wavelengths. Viewed through a spectroscope, this light appears as a series of sharp, bright lines. In such gases the average spacing between atoms is large, so the radiation emitted can be attributed to individual, non-interacting atoms rather than to interactions between neighbouring atoms or molecules.
It was established early in the nineteenth century that every element has its own characteristic spectrum — a fixed set of spectral lines at fixed relative positions. Hydrogen, for instance, always produces the same set of lines. This pointed to a close relationship between an atom's internal structure and the spectrum of radiation it emits.
Balmer's Formula for Hydrogen
In 1885, Johann Jakob Balmer obtained a simple empirical formula giving the wavelengths of a group of spectral lines emitted by atomic hydrogen. Since hydrogen is the simplest of all elements, this chapter studies its spectrum in detail as a way of understanding atomic structure in general.
Rutherford's Proposal and the Birth of the Nuclear Model
Ernst Rutherford, a former research student of J. J. Thomson, was engaged in experiments on the alpha-particles emitted by certain radioactive elements. In 1906 he proposed a classic experiment: scatter these alpha-particles off atoms as a way of probing atomic structure. This experiment was carried out around 1911 by Hans Geiger and Ernst Marsden (then a twenty-year-old student who had not yet earned his degree) — its details, and the model it led to, are taken up in Section 12.2.
The explanation of the scattering results gave birth to Rutherford's planetary model of the atom (also called the nuclear model). According to this model, the entire positive charge of the atom, and almost all of its mass, is concentrated in a small central volume called the nucleus, with the electrons revolving around it much as planets revolve around the Sun.
The Problem That Remained
Rutherford's nuclear model was a major advance in how we picture the atom, but it left one serious question unanswered: it could not explain why atoms emit light of only discrete wavelengths. How could an atom as simple as hydrogen — a single electron bound to a single proton — produce a spectrum made up of several distinct wavelengths? In the classical picture, the electron revolves around the nucleus much as a planet revolves around the Sun, but as we will see, this classical picture runs into serious difficulties.
The inability of the classical, Rutherford-style picture to account for discrete atomic spectra is exactly the puzzle that Niels Bohr's ideas — introduced later in this chapter — were designed to resolve.