Atoms

FIELDS OF STUDY: Analytical Chemistry; Spectroscopy; Inorganic Chemistry

ABSTRACT

The basic structure of atoms is defined, and the development of the modern theory of atomic structure is elaborated. Atomic structure is fundamental to all fields of chemistry, but especially to fields that rely on the intrinsic properties of individual atoms.

Visualizing the Atom

An atom is the smallest unit of elemental matter that retains and defines the specific properties of that material. The concept can be visualized by imagining a sample of a pure elemental material being repeatedly divided into ever-smaller portions. Eventually a point would be reached at which the material could no longer be subdivided without destroying its identity. The remaining indivisible portion is an atom of the element. Any further subdivision would require breaking the atom into its component protons, neutrons, and electrons.

Historical Theories of Atomic Structure

In their attempts to comprehend the basic nature of the universe, early Greek philosophers, particularly Democritus (ca. 460 BCE–ca. 370 BCE), followed this kind of logical reasoning to the philosophical concept of the atomos (indivisible), the fundamental thing from which all matter was made. Greek philosophy, however, was based on thought, not experimentation. Two thousand years later, in the Middle Ages, the practice of alchemy arose alongside the practices of metallurgy, and alchemists sought means of transforming materials into other materials, typically through magic and arcane practices. The refinement and working of metals became a very important practical study and gave rise to the first truly scientific book of chemistry, De re metallica (On the nature of metals, 1556), by Georgius Agricola (1494–1555). Alchemy eventually gave way to more scientific study of matter through the use of weights and measures, enabling early scientists to recognize the relationships between matter that formed the foundation on which the modern atomic theory has been constructed.

In the nineteenth century, scientists began to reject the philosophical atomos in favor of the physical atom as the basic building block of matter. John Dalton (1766–1844) conceived of the atom as though it were a billiard ball: a single, hard, uniform spherical object. He based this view on the behavior of gases, some of which he found he could describe mathematically with the billiard model. Other chemical behaviors required a mechanism to account for electrical charge, however, especially in to the case of ionic interactions.

In 1897, J. J. Thomson (1856–1940) discovered that when a polarized electricity source is discharged within a gas-filled tube, the rays emitted from the cathode (the negative electrode), known as cathode rays, obey the mathematics of a stream of charged particles. After calculating the likely mass of the particles, Thomson determined that they were even smaller than atoms, thus proving the existence of subatomic particles. He called these particles "corpuscles" at first, though they soon became known as electrons. Following this discovery, Thomson developed a new atomic model, one in which small, negatively charged particles—electrons—were embedded throughout a positively charged matrix, rather like the plums in a plum pudding. This came to be known as the plum-pudding model.

The Modern Theory of Atomic Structure

In 1909, Ernest Rutherford (1871–1937), aided by Hans Geiger (1882–1945) and Ernest Marsden (1889–1970), conducted the gold-foil experiment, which demonstrated that atoms consist of a nucleus surrounded by primarily empty space. In this experiment, a beam of alpha particles (the nuclei of helium atoms) was directed at a piece of very thin gold foil, which was surrounded by a film strip that would detect alpha particles as they passed through the foil. Rutherford found that while the vast majority of the particles passed directly through the foil as though there was nothing there, some of the particles were deflected in all directions by the foil. This demonstrated that atoms consist of a very small, dense nucleus that can deflect alpha particles, surrounded by a thin, diffuse cloud that is not capable of affecting their movement. Rutherford continued experimenting with alpha particles, and in 1917 he discovered that the collision of an alpha particle with the nucleus of a nitrogen atom released a particle identical to a hydrogen nucleus—in other words, a proton.

Considerations of the energies associated with electrons in atoms led Niels Bohr (1885–1962) to postulate in 1913 that the electrons must be restricted to specific orbital paths around the nucleus, in accordance with the corresponding wavelengths calculated by Louis de Broglie (1892–1987). Measurement of the wavelengths of various emission and absorption spectra indicated that electrons in atoms can only have certain specific energy levels, which is consistent with the idea that electrons can only follow certain stable orbits around a nucleus.

Because of its electrical neutrality, the neutron remained an elusive theoretical construct until 1932, when an experiment by James Chadwick (1891–1974) demonstrated their existence. Chadwick used alpha particles to bombard a beryllium target, causing it to release radiation that would strike a paraffin target, which would in turn release protons from hydrogen nuclei into an ionization chamber. He determined that the radiation released from the beryllium was actually a neutral particle of approximately the same mass as a proton, now known to be a neutron.

These discoveries about the nature of the various subatomic particles formed the foundation of a new model of atomic structure, based on the principles of quantum mechanics. In this model, the energy levels of electrons can only change by discrete steps, or quanta, and thus electrons can occupy only specific regions of space about the nucleus, which came to be called atomic orbitals. These orbitals have specific shapes and orientations around the nucleus of the atom, which determine how atoms interact to form chemical bonds. The physical orientation of those bonds determines the shapes and properties of both the atoms and the molecules that they form.

The Periodic Table

The periodic table of the elements displays the relationships of the elements according to the quantum mechanical model of atomic structure. Chemists had tried to sequentially order the known elements by their common properties since the late eighteenth century. Dmitri Mendeleev (1834–1907) is credited with developing the first modern periodic table, which he published in 1869, though much information was missing. The periodic table of the elements contains all ninety-eight naturally occurring elements, as well as eighteen more that have been (or are claimed to have been) produced artificially. Each square in the table corresponds to a single element. In a standard periodic table, the internationally accepted symbol of the element is in the center of the square, with its atomic number at the top and its atomic mass at the bottom. Some periodic tables also include the normal oxidation states in which each element is found and its electron configuration according to the quantum mechanical model of atomic structure.

The identity of each atom is determined by the number of protons in its nucleus, meaning that all atoms with the same number of protons in their respective nuclei are atoms of the same element. However, the number of neutrons and electrons in an atom can change. Neutral atoms must have the exact same number of electrons orbiting the nucleus as there are protons within the nucleus; those atoms with more or fewer electrons than protons are ions, or atoms with a net electrical charge. Atoms with different numbers of neutrons in their nuclei are different isotopes of the same element. The atomic mass of an atom is the total mass of all protons, neutrons, and electrons in the atom; however, because electrons are so small and have such little mass, the atomic mass can be approximated by adding the number of neutrons and protons together. Atomic mass is measured in unified atomic mass units (u), one of which is approximately equal to the mass of a proton or neutron.

PRINCIPAL TERMS

• atomic mass: the total mass of the protons, neutrons, and electrons in an individual atom.
• atomic number: the number of protons in the nucleus of an atom, used to uniquely identify each element.
• electron: a fundamental subatomic particle with a single negative electrical charge, found in a large, diffuse cloud around the nucleus.
• element: a form of matter consisting only of atoms of the same atomic number.
• isotope: an atom of a specific element that contains the usual number of protons in its nucleus but a different number of neutrons.
• neutron: a fundamental subatomic particle in the atomic nucleus that is electrically neutral and about equal in mass to the mass of one proton.
• nucleus: the central core of an atom, consisting of specific numbers of protons and neutrons and accounting for at least 99.98 percent of the atomic mass.
• proton: a fundamental subatomic particle with a single positive electrical charge, found in the atomic nucleus.

Bibliography

Agricola, Georgius. De re metallica. Trans. Herbert Clark Hoover and Lou Henry Hoover. New York: Dover, 1950. Print.

The Britannica Guide to the 100 Most Influential Scientists. Introd. John Gribbin. London: Constable, 2008. Print.

Gribbin, John. Science: A History, 1543–2001. London: Lane, 2002. Print.

Johnson, Rebecca L. Atomic Structure. Minneapolis: Lerner, 2008. Print.

Kragh, Helge. Niels Bohr and the Quantum Atom: The Bohr Model of Atomic Structure, 1913–1925. Oxford: Oxford UP, 2012. Print.

Winter, Mark J. The Orbitron: A Gallery of Atomic Orbitals and Molecular Orbitals on the WWW. U of Sheffield, 2002. Web. 4 Apr. 2014.