Atomic Structure and Chemical Laws: Key Discoveries
**Guy-Lussac’s Law**
The French chemist, Guy-Lussac, studied the behavior of gases when their temperature changed. His studies extended to all gases and he found that, regardless of the nature of the gas in question, the behavior was similar. Guy-Lussac’s law, also known as one of the two laws on the expansion of gases, states that the pressure of a gas is directly proportional to its temperature. These studies later led him to investigate the chemical reactions of gaseous substances, discovering the law of combining volumes, or Guy-Lussac’s Law: When chemical reactions produce gaseous substances, the volumes of the substances involved, measured under the same conditions of pressure and temperature, maintain simple and natural number ratios.
**Lavoisier’s Law**
When Lavoisier studied combustion, he used a hermetically sealed container and found that the mass before and after combustion did not change. The same happens in any other chemical reaction: although things may seem to disappear, the mass does not change before and after the transformation. In the case of combustion, an interplay occurs between the components of the fuel and the oxygen in the air, producing mainly carbon dioxide and water, both in gaseous states. This is why the fuel seems to almost completely disappear. Thanks to these and other experiments, Lavoisier enunciated the law that bears his name, or the law of conservation of mass: In a chemical reaction, mass is neither created nor destroyed, only transformed.
**Atomic Spectra**
That chemical elements emit different colored light when heated was a familiar fact to manufacturers of fireworks. By mixing different compounds in gunpowder, they obtained explosions of diverse colors. The German chemist, Robert Bunsen, applied this property and demonstrated that the study of the flame of a given chemical compound could help determine its composition. In the mid-nineteenth century, with the collaboration of the physicist Gustav Kirchhoff, Bunsen discovered atomic spectra. When a chemical element is heated, it emits a bright light, but with different properties than sunlight. If sunlight is decomposed by passing it through a prism, a continuous spectrum of colors is obtained, similar to a rainbow. Passing the light emitted by an element through a prism does not produce a continuous spectrum of light, but a series of colored bands on a black background. The emission spectrum of an element is characteristic; no other element presents the same colored bands. Furthermore, it is independent of the compound the element forms. For example, the spectrum of sodium always presents the same pattern, regardless of whether it is in the form of chloride, sulfate, or oxide. The importance of atomic spectra in chemical analysis was quickly accepted, and thanks to them, new elements were discovered, such as helium in the Sun, even before its detection on Earth.
**Absorption Spectra**
If a beam of white light is passed through the vapors of an element, an absorption spectrum is obtained. When white light is decomposed, a series of dark bands appear, precisely in the same locations where the heated element produced an emission spectrum. Since the emission and absorption spectra coincide and do not depend on the compound being studied, but on the element, it seems clear that atomic spectra are related to atoms, which must have an internal structure that accounts for both types of spectra.
**Thomson and Rutherford Models**
Thomson, the discoverer of the electron, suggested an atomic structure similar to a plum pudding. The atom was a spongy sphere positively charged with embedded electrons, as many as were necessary to compensate for the atom’s charge and make it electrically neutral. To test this model, Rutherford designed a series of experiments in which he bombarded a very thin gold foil with alpha particles, which are radioactive particles with a positive charge. If Thomson’s atomic model corresponded to reality, the particles would pass through the atoms without altering their trajectory. Although Rutherford observed that most of the particles passed through the foil as the theory predicted, a few practically bounced back. In Rutherford’s own words: “It was as surprising as if you fired 15-inch bullets at a sheet of paper, and some of them bounced back.” To explain this experiment, Rutherford proposed his atomic model, in which the atom is practically an empty space with a nucleus in the center, which contains almost all the mass of the atom and the positive charge. Electrons revolve around the nucleus, forming a neutral atom.
**Bohr Model**
Rutherford’s atomic model was incomplete, as it did not explain the atomic spectra and, moreover, was unstable: an electron following circular orbits would emit energy and eventually fall into the nucleus. The Danish physicist Niels Bohr proposed a new atomic model based on four postulates, among which was the stability of orbitals:
- The atom is formed by a nucleus, with a positive charge that contains most of the atom’s mass, and a cloud where the electrons move.
- Electrons move in circular orbits around the nucleus.
- Only those orbits in which the electron’s angular momentum is a multiple of Planck’s constant are possible. In these orbits, the electron neither emits nor absorbs energy.
- The transition of an electron from one orbit to another involves the absorption or emission of radiation. The atom absorbs or emits radiation only when an electron transitions from one orbit to another.
Bohr’s atomic model allowed for the explanation of the atomic spectrum of hydrogen, but it could not predict the spectra of the remaining elements of the periodic table, so it was abandoned in favor of quantum mechanics.
**Quantum Numbers**
Today, the atomic model is based on quantum mechanics, and the electron is described by a wave function. This function is defined by four quantum numbers: n, l, and m. The first three describe the region of space where the electron is found, which is called an orbital.
- The principal quantum number, n, can take any natural number except zero: 1, 2, 3, 4… It indicates the size of the orbital and its energy. The larger it is, the larger the orbital and the greater its energy.
- The azimuthal quantum number, l, is related to the angular momentum of the electron and the shape of the orbital.
- The magnetic quantum number, m, is related to the spatial orientation of the orbital.