Atom

C

Rutherford’s Nuclear Atom

Scientists realized that if all atoms contain electrons but are electrically neutral, atoms must also contain an equal quantity of positive charge to balance the electrons’ negative charge. Furthermore, if electrons are indeed much less massive than even the lightest atom, then this positive charge must account for most of the mass of the atom. Thomson proposed a model by which this phenomenon could occur: He suggested that the atom was a sphere of positive charge into which the negative electrons were imbedded, like raisins in a loaf of raisin bread. In 1911 British scientist Ernest Rutherford set out to test Thomson’s proposal by firing a beam of charged particles at atoms.

Rutherford chose alpha particles for his beam. Alpha particles are heavy particles with twice the positive charge of a proton. Alpha particles are now known to be the nuclei of helium atoms, which contain two protons and two neutrons. If Thomson’s model of the atom was correct, Rutherford theorized that the electric charge and the mass of the atoms would be too spread out to significantly deflect the alpha particles. Rutherford was quite surprised to observe something very different. Most of the alpha particles did indeed change their paths by a small angle, and occasionally an alpha particle bounced back in the opposite direction. The alpha particles that bounced back must have struck something at least as heavy as themselves. This led Rutherford to propose a very different model for the atom. Instead of supposing that the positive charge and mass were spread throughout the volume of the atom, he theorized that it was concentrated in the center of the atom. Rutherford called this concentrated region of electric charge the nucleus of the atom.

In the span of 100 years, from Dalton to Rutherford, the basic ideas of atomic structure evolved from very primitive concepts of how atoms combined with one another to an understanding of the constituents of atoms—a positively charged nucleus surrounded by negatively charged electrons. The interactions between the nucleus and the electrons still required study. It was natural for physicists to model the atom, in which tiny electrons orbit a much more massive nucleus, after a familiar structure such as the solar system, in which planets orbit around a much more massive Sun. Rutherford’s model of the atom did indeed resemble a tiny solar system. The only difference between early models of the nuclear atom and the solar system was that atoms were held together by electromagnetic force, while gravitational force holds together the solar system.

D

The Bohr Model

Danish physicist Niels Bohr used new knowledge about the radiation emitted from atoms to develop a model of the atom significantly different from Rutherford’s model. Scientists of the 19th century discovered that when an electrical discharge passes through a small quantity of a gas in a glass tube, the atoms in the gas emit light. This radiation occurs only at certain discrete wavelengths, and different elements and compounds emit different wavelengths. Bohr, working in Rutherford’s laboratory, set out to understand the emission of radiation at these wavelengths based on the nuclear model of the atom.

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Using Rutherford’s model of the atom as a miniature solar system, Bohr developed a theory by which he could predict the same wavelengths scientists had measured radiating from atoms with a single electron. However, when conceiving this theory, Bohr was forced to make some startling conclusions. He concluded that because atoms emit light only at discrete wavelengths, electrons could only orbit at certain designated radii, and light could be emitted only when an electron jumped from one of these designated orbits to another. Both of these conclusions were in disagreement with classical physics, which imposed no strict rules on the size of orbits. To make his theory work, Bohr had to propose special rules that violated the rules of classical physics. He concluded that, on the atomic scale, certain preferred states of motion were especially stable. In these states of motion an orbiting electron (contrary to the laws of electromagnetism) would not radiate energy.

At the same time that Bohr and Rutherford were developing the nuclear model of the atom, other experiments indicated similar failures of classical physics. These experiments included the emission of radiation from hot, glowing objects (called thermal radiation) and the release of electrons from metal surfaces illuminated with ultraviolet light (the photoelectric effect). Classical physics could not account for these observations, and scientists began to realize that they needed to take a new approach. They called this new approach quantum mechanics (see Quantum Theory), and they developed a mathematical basis for it in the 1920s. The laws of classical physics work perfectly well on the scale of everyday objects, but on the tiny atomic scale, the laws of quantum mechanics apply.

E

Quantum Theory of Atoms

The quantum mechanical view of atomic structure maintains some of Rutherford and Bohr’s ideas. The nucleus is still at the center of the atom and provides the electrical attraction that binds the electrons to the atom. Contrary to Bohr’s theory, however, the electrons do not circulate in definite planet-like orbits. The quantum-mechanical approach acknowledges the wavelike character of electrons and provides the framework for viewing the electrons as fuzzy clouds of negative charge. Electrons still have assigned states of motion, but these states of motion do not correspond to fixed orbits. Instead, they tell us something about the geometry of the electron cloud—its size and shape and whether it is spherical or bunched in lobes like a figure eight. Physicists called these states of motion orbitals. Quantum mechanics also provides the mathematical basis for understanding how atoms that join together in molecules share electrons. Nearly 100 years after Faraday’s pioneering experiments, the quantum theory confirmed that it is indeed electrical forces that are responsible for the structure of molecules.

Two of the rules of quantum theory that are most important to explaining the atom are the idea of wave-particle duality and the exclusion principle. French physicist Louis de Broglie first suggested that particles could be described as waves in 1924. In the same decade, Austrian physicist Erwin Schrödinger and German physicist Werner Heisenberg expanded de Broglie’s ideas into formal, mathematical descriptions of quantum mechanics. The exclusion principle was developed by Austrian-born American physicist Wolfgang Pauli in 1925. The Pauli exclusion principle states that no two electrons in an atom can have exactly the same characteristics.

The combination of wave-particle duality and the Pauli exclusion principle sets up the rules for filling electron orbitals in atoms. The way electrons fill up orbitals determines the number of electrons that end up in the atom’s valence shell. This in turn determines an atom’s chemical and physical properties, such as how it reacts with other atoms and how well it conducts electricity. These rules explain why atoms with similar numbers of electrons can have very different properties, and why chemical properties reappear again and again in a regular pattern among the elements.

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