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English schoolteacher John Dalton helped to revolutionize chemistry with his hypothesis that the behavior of matter could be explained using an atomic theory. First published in 1807, many of Dalton’s hypotheses about the microscopic features of matter are still valid in modern atomic theory.
Dalton’s atomic theory provides a microscopic explanation of the many macroscopic properties of matter that you’ve learned about. For example, if an element, such as copper, consists of only one kind of atom, then it cannot be broken down into simpler substances, that is, into substances composed of fewer types of atoms. And if atoms are neither created nor destroyed during a chemical change, then the total mass of matter present when matter changes from one type to another will remain constant (the law of conservation of matter).
French chemist Joseph Proust had demonstrated that all samples of a pure compound contain the same elements in the same proportion by mass, which is known as the law of definite proportions or the law of constant composition. The suggestion that the numbers of atoms of the elements in a given compound always exist in the same ratio is consistent with these observations.
Dalton also used data from Proust, as well as results from his own experiments, to formulate another interesting law. The law of multiple proportions states that when two elements react to form more than one compound, a fixed mass of one element will react with masses of the other elements in a ratio of small, whole numbers.
This apparatus consisted of a sealed glass tube from which almost all the air had been removed; the tube contained two metal electrodes. When high voltage was applied across the electrodes, a visible beam called a cathode ray appeared between them. This beam was deflected toward the positive charge and away from the negative charge. It was produced in the same way with identical properties when different metals were used for the electrodes.
In similar experiments, the ray was simultaneously deflected by an applied magnetic field, and measurements of the extent of deflection and the magnetic field strength allowed Thomson to calculate the charge-to-mass ratio of the cathode ray particles. The results of these measurements indicated that these particles were much lighter than atoms.
The images above show J. J. Thomson produced a visible beam in a cathode ray tube (a). An early cathode ray tube (b), invented in 1897 by Ferdinand Braun. In the diagram of a cathode ray (c), the beam (shown in yellow) comes from the cathode and is accelerated past the anode toward a fluorescent scale at the end of the tube. Simultaneous deflections by applied electric and magnetic fields permitted Thomson to calculate the mass-to-charge ratio of the particles composing the cathode ray.
Although controversial at the time, Thomson’s idea was gradually accepted, and his cathode-ray particle is what we now call an electron, a negatively charged, subatomic particle with a mass more than one thousand times less that of an atom. The term “electron” was coined in 1891 by Irish physicist George Stoney, from “electric ion.”
Scientists had now established that the atom was not indivisible as Dalton had believed, and due to the work of Thomson and others, the charge and mass of the negative, subatomic particles—the electrons—were known. However, the positively charged part of an atom was not yet well understood.
In the following image, William Thomson suggested that atoms resemble plum pudding (a), an English dessert consisting of moist cake with embedded raisins (“plums”). (b) Nagaoka proposed that atoms resembled the planet Saturn, with a ring of electrons surrounding a positive “planet.”
The next major development in understanding the atom came from Ernest Rutherford, a physicist from New Zealand who largely spent his scientific career in Canada and England. He performed a series of experiments using a beam of high-speed, positively charged alpha particles (α particles) that were produced by the radioactive decay of radium. α particles consist of two protons and two neutrons (you will learn more about radioactive decay in the chapter on nuclear chemistry).
Rutherford and his colleagues Hans Geiger (later famous for the Geiger counter) and Ernest Marsden aimed a beam of α particles, the source of which was embedded in a lead block to absorb most of the radiation, at a very thin piece of gold foil and examined the resultant scattering of the α particles using a luminescent screen that glowed briefly where hit by an α particle.
What did they discover? Most particles passed right through the foil without being deflected at all. However, some were diverted slightly, and a very small number were deflected almost straight back toward the source. Rutherford described finding these results: “It was quite the most incredible event that has ever happened to me in my life. 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.”
IN CONTEXT
Here is what Rutherford deduced:
- Because most of the fast-moving α particles passed through the gold atoms undeflected, they must have traveled through essentially empty space inside the atom.
- Alpha particles are positively charged, so deflections arise when they encountered another positive charge (like charges repel each other).
- Since like charges repel one another, the few positively charged α particles that changed paths abruptly must have hit, or closely approached, another body that also had a highly concentrated, positive charge.
- Since the deflections occurred a small fraction of the time, this charge only occupied a small amount of the space in the gold foil.
This analysis led Rutherford to propose a model in which an atom consists of a very small, positively charged nucleus, in which most of the mass of the atom is concentrated, surrounded by the negatively charged electrons so that the atom is electrically neutral.
After many more experiments, Rutherford also discovered that the nuclei of other elements contain the hydrogen nucleus as a “building block,” and he named this more fundamental particle the proton, the positively charged, subatomic particle found in the nucleus. With one addition (the neutron), this nuclear model of the atom, proposed over a century ago, is still used today.
One puzzle remained: The nucleus was known to contain almost all of the mass of an atom, with the number of protons only providing half, or less, of that mass. Different proposals were made to explain what constituted the remaining mass, including the existence of neutral particles in the nucleus.
As you might expect, detecting uncharged particles is very challenging, and it was not until 1932 that James Chadwick found evidence of neutrons, uncharged, subatomic particles with a mass approximately the same as that of protons. The existence of the neutron also explained isotopes. They differ in mass because they have different numbers of neutrons, but they are chemically identical because they have the same number of protons.
Source: THIS TUTORIAL HAS BEEN ADAPTED FROM OPENSTAX “CHEMISTRY: ATOMS FIRST 2E”. ACCESS FOR FREE AT Chemistry: Atoms First 2e. LICENSE: CREATIVE COMMONS ATTRIBUTION 4.0 INTERNATIONAL
