Atomic structure is an abstract concept that proves difficult for many students. Several of the activities included here simplify the structure of the atom to a point that arguably introduces errors; however, they are meant to help students conceptualize the ideas of the atom in a manner they can grasp. Even before describing to students what we believe to be in an atom, it is important for them to understand how and why we believe what we do. (NSES, 177) A short historical development of the current conception of the atom follows.
Atoms were first suggested by Democritus of Abdera, a Greek who lived during the fourth century B.C. His ideas, however, were not based on any experimental evidence. John Dalton, in the late 1700's is credited with first relating observed chemical phenomena to the idea of individual atoms. His theory stemmed from his study of combining substances through chemical reactions to make new substances. He found substances combined in fixed and predictable proportions, which led to his theory that:
Much of Dalton's theory remains accepted; however, further study led scientists to understand that atoms are in fact composed of smaller sub-atomic particles and not all atoms of the same element are identical. Though even more recent experimentation has revealed fundamental particles beyond electrons, protons, and neutrons, the two models mentioned in the Virginia Science Standards of Learning for Physical Science, require familiarity only with these three particles.
Electrons were first discovered by scientists who were more interested in electricity than atomic structure. Sir Joseph J. Thomson (1856-1940) was one of these scientists. In 1897, he found cathode rays could be deflected by magnets or electrically charged plates. A cathode ray is a beam of particles that travels from the cathode (negatively charged plate) to the anode (positively charged plate) inside a glass tube with a vacuum or low-pressure gas. Through his experiment, Thomson showed cathode rays are collections of small negatively charged particles, which he named electrons. He further found they were nearly 2000 times lighter than a hydrogen atom, the lightest element. Because he found that all cathode rays were composed of electrons, he concluded that electrons must be a fundamental part of the atom. He formulated a model, which he named the "Plum Pudding Model", to describe what he felt the atom must be like. In the model, electrons were distributed throughout an atomic material much like the fruit in a plum pudding or chocolate chips in a cookie.
Atoms are normally electrically neutral. Because the electrons were found as negatively charged, scientists deduced that some positive charge source must be present to off-set the negative. Experimental evidence soon led to the discovery of the proton which is a positively charged subatomic particle that is approximately 1840 times heavier than an electron. It was not until 1932 that Sir James Chadwick found and confirmed evidence of the neutron, or the electrically neutral subatomic particle.
The discovery of the nucleus is one example of an ingenious indirect measurement technique that students can readily appreciate. Ernest Rutherford (1871-1937) wanted to test the accepted structure of an atom at the time. To test this theory, Rutherford set up what came to be known as the "Gold Foil Experiment" where he directed a beam of alpha particles (helium ions which are positively charged, relatively heavy, and fast moving) at a very thin sheet of gold. A fluorescent screen was positioned around the gold foil set-up. The expected result was that the alpha particles would travel through the gold foil unabated, strike the screen, and cause a bright dot to appear at the point of impact. To Rutherford's surprise, however, some of the alpha particles were actually scattered by the gold foil.
His results led Rutherford to conclude that the atom contained some dense and positively charged region that he termed the nucleus. He proposed that the remainder of the atom must be empty space containing the much smaller electrons. Later experiments led scientists to discover a third subatomic particle, the neutron, located in the nucleus with the protons.
Niels Bohr (1885-1962) was a young associate of Rutherford. In 1913, he proposed another model of the atom where electrons are arranged in circular paths about the nucleus. His model closely related to the movement of the planets around the sun and is sometimes referred to as the "Planetary Model." In this model, electrons in a particular path (or radius) contain a fixed amount of energy that corresponds to that path. As long as an electron continues to have that energy, it would remain in that path or energy level. (Energy level is a region around the nucleus where an electron is likely to be moving.) From experiments, he knew that atoms could under some circumstances absorb energy ; under other conditions, atoms can emit energy. Emitted energy was in the form of light of discrete and predictable frequencies seen as colors. He deduced then, that when energy was added, electrons were pushed to new and higher energy levels that were discrete like rungs on a ladder. He concluded that electrons could not exist at just any level within the atom-only at the defined energy levels which, unlike a ladder, were not evenly spaced around the nucleus. He termed these packets of energy "quanta" (singular quantum or the amount of energy required to move an electron from its present energy level to the next higher one.)
Thirteen years after Bohr developed his model, Austrian Erwin Schroedinger went another step in the development of atomic models with the quantum mechanical model. Where the earlier models were physical in nature, Schroedinger's model was very mathematical. Like the Bohr model, the quantum model includes quantized energy levels for the electrons, but it does not define exact paths for the electrons. Instead, the model includes the likely path or probable path of an electron about the nucleus-like a blurry cloud of negative charge. It is also analogous to locating the position of a single fan blade when it is rotating. At a given instant, the fan blade is at a certain location but there are many possible locations available to it. Over a period of time, these positions all blend to form a blur where there is a high probability of finding the fan blade. The electron cloud is similarly a time-averaged view of the electron.
The model further reveals information about the arrangement of electrons within the energy levels. Within the first energy level closest to the nucleus, two electrons can be located. In the second energy level, up to eight electrons can be found. The third level can hold up to eighteen electrons, and the fourth can hold up to thirty-two. The atomic orbitals are more complicated than this description implies, but this offers a glimpse at their arrangement.
All atoms follow the basic arrangement, yet atoms of different substances are different. Scientists currently know of approximately 118 different atoms. Each is unique chemically and physically. Elements are identified by the number of protons in the nucleus, which is given as the atomic number on the periodic table. Neutral atoms contain an equal number of electrons around the nucleus as protons in the nucleus. The electrons are arranged in energy levels as described in the quantum model. As one moves across the periodic table of elements, one finds a single proton and a single electron are added to the preceding atom. The number of neutrons is not always as straightforward as the number of protons. In a single atom, the number of neutrons is found by subtracting the atomic number from the atomic mass. Usually, atomic masses on the periodic table are not integral values because the reported values are weighted averages of all known isotopes or forms of the atom. (Isotopes-atoms of the same element with different numbers of neutrons in the nucleus.)
National Science Education Standards (1996). National Academy Press, Washington, D.C., p. 177.
Tocci, S., Viehland, C. (1996). Chemistry Visualizing Matter, Holt, Rinehart, and Winston, New York, pp. 81-96.
Wilbraham, A.C., Staly, D. D., Simpson, C. J., Matta, M. S. (1987). Chemistry, Addison-Wesley Publishing Company, Reading MA, pp. 67-79, 245-249.