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Subatomic Particles
If matter is composed of atoms, what are atoms composed of? Are they the smallest particles, or is there something smaller? There have been several minor but important modifications to Dalton’s atomic theory that shed light to the inner structure of atoms answering just these questions. For one thing, Dalton considered atoms to be indivisible. We know now that atoms not only can be divided, but also are composed of three different kinds of particles, called subatomic particles, with their own properties that are different from the chemical properties of atoms.
It was learned that an atom contains a very small nucleus composed of positively charged protons and uncharged neutrons, surrounded by a much larger volume of space containing negatively charged electrons.
The Electron
The first subatomic particle was identified in in the late 1800s. A number of scientists who were interested in the structure of the atom investigated the electrical discharges that could be produced in low-pressure gases. The most significant discovery was made by an English physicist, J. J. Thomson, using a cathode ray tube in a magnetic field. 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, and the result did not change when different metals were used for the electrodes. His experiments with magnetic fields further revealed the existence of an extremely tiny subatomic particle which was negatively charged. As a result of these experiments, Thomson was able to calculate the charge-to-mass ratio of the cathode ray particles. This subatomic particle was later called the electron, by Irish physicist George Stoney, derived from “electric ion.”.
In 1909, more information about the electron was uncovered by American physicist Robert A. Millikan via his “oil drop” experiments. Millikan created microscopic oil droplets, which could be electrically charged by friction as they formed or by using X-rays. These droplets initially fell due to gravity, but their downward progress could be slowed or even reversed by an electric field lower in the apparatus. By adjusting the electric field strength and making careful measurements and appropriate calculations, Millikan was able to determine the charge on individual drops.
Since the charge of an electron was now known due to Millikan’s research, and the charge-to-mass ratio was already known due to Thomson’s research it was very easy to calculate the mass of an electron, and it was determined that it had a mass of about 9.109 × 10−31 kg.
The Nucleus
Scientists had now established that the atom was not indivisible as Dalton had believed, and due to the work of Thomson, Millikan, and others, the charge and mass of the negative subatomic particles—the electrons—were known.
Between 1909 and 1911, Ernest Rutherford, a Cambridge physicist, and his associates Hans Geiger and Ernest Marsden performed experiments that provided strong evidence concerning the internal structure of an atom. They took a very thin metal foil, such as gold or platinum, and aimed a beam of positively charged particles (called alpha particles, which are combinations of two protons and two neutrons) from a radioactive source toward the foil. Surrounding the foil was a detector—either a scintillator (a material that glows when hit by such particles) or some unexposed film (which is exposed where the particles hit it). The detector allowed the scientists to determine the distribution of the alpha particles after they interacted with the foil. Figure \(\PageIndex{1}\) shows a diagram of the experimental setup. The experiment showed that 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.
(a)
(b)
Figure \(\PageIndex{1}\): a) Geiger and Rutherford fired α particles at a piece of gold foil and detected where those particles went, as shown in this schematic diagram of their experiment. b) Most of the particles passed straight through the foil, but a few were deflected slightly and a very small number were significantly deflected.(CC BY 4.0; Paul Flowers, Klaus Theopold, Richard Langley, William R. Robinson via OpenStax). Detailed Description of Figure \(\PageIndex{1}\).
Analyzing a series of such experiments in detail, Rutherford drew two conclusions:
- The volume occupied by an atom must consist of a large amount of empty space.
- A small, relatively heavy, positively charged body, the nucleus, must be at the center of each atom.
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. Protons and neutrons are concentrated in a central region he called the nucleus (plural, nuclei) of the atom. Electrons are outside the nucleus and orbit about it because they are attracted to the positive charge in the nucleus. Most of the mass of an atom is in the nucleus, while the orbiting electrons account for an atom’s size. As a result, an atom consists largely of empty space. Rutherford called his description the “planetary model” of the atom, Figure \(\PageIndex{2}\).
Isotopes
Another important finding was the discovery of isotopes. During the early 1900s, scientists identified several substances that appeared to be new elements, isolating them from radioactive ores. For example, a “new element” produced by the radioactive decay of thorium was initially given the name mesothorium. However, a more detailed analysis showed that mesothorium was chemically identical to radium (another decay product), despite having a different atomic mass. This result, along with similar findings for other elements, led the English chemist Frederick Soddy to realize that an element could have types of atoms with different masses that were chemically indistinguishable. These different types are called isotopes—atoms of the same element that differ in mass. Soddy was awarded the Nobel Prize in Chemistry in 1921 for this discovery.
The Proton and the Neutron
By 1920, experimental evidence indicated the existence of a second particle. 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. A proton has the same amount of charge as an electron, but its charge is positive, not negative. Another major difference between a proton and an electron is mass. Although still incredibly small, the mass of a proton is 1.673 × 10−27 kg, which is almost 2,000 times greater than the mass of an electron. Because opposite charges attract each other (while like charges repel each other), protons attract electrons (and vice versa).
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.
Key Takeaways
- Atoms are composed of three main subatomic particles: protons, neutrons, and electrons.
- Protons and neutrons are grouped together in the nucleus of an atom, while electrons orbit about the nucleus.
- Isotopes differ in mass because they have different numbers of neutrons, but they are chemically identical because they have the same number of protons.
This page titled 2.4: The Structure of Atoms was last modified on Mon, 18 May 2026 02:32:09 GMT and is shared under a CC BY-NC-SA 4.0 license and was authored, remixed, and/or curated by Burçak Artun, Gorkem Ermut, and Semere Bairu.
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