Facts of the Matter
Understanding the tiny atom is a giant task
FROM THE TIME of the discovery that atoms were composed of smaller particles, chemists and physicists tried to visualize how an atom looks.
The earliest model was J.J. Thompson's 1897 "plum pudding" atom, which consisted of a positive nucleus with negative electrons embedded in it.
This model was displaced by a miniature "solar system" model when Ernest Rutherford and his team discovered in 1911 that the atom is mostly empty space.
The problem was that Maxwell's electromagnetic equations predicted that the electrons orbiting the nucleus should emit radiation as they lost energy and spiraled into the nucleus.
One year later, Niels Bohr made the startling proposition that electron orbits are quantized, allowing the electrons to exist in only certain specific energy levels.
This led to the "electron shell" model that we might vaguely remember from high school chemistry.
The Bohr atom revolutionized physics and opened another door into the world of the quantum that had already been set in motion by Max Planck and Albert Einstein.
Upon closer inspection, the electron shells revealed a structure that would ultimately lead to an explanation of the nature of chemical bonds by Linus Pauling in the 1930s, thereby bringing physics and chemistry under the same physical science umbrella.
Chemical reactions are really nothing more than rearrangements of atoms.
They are the result of electrons moving between or among atoms due to differences in atoms' affinity for electrons.
ANY STABLE arrangement of chemically bonded atoms represents an energetically favorable state, although not necessarily the lowest possible energy state.
Most substances are in a state of metastable, neutral equilibrium, like a ball resting high on a staircase.
Nearly all spontaneous chemical reactions take place in such a way as to leave the products of the reaction in a lower state of energy than the reactants. For this reason most chemical reactions give off heat that represents the loss of energy.
The atoms of each chemical element differ by the number of positive charges in their nuclei. Although atoms of the same element can have a different number of neutrons, it is the positively charged protons that distinguish and atom of one chemical element from another. Extra neutrons slow the rate of reaction but do not affect the overall chemical properties of the element.
But it is the orbiting electrons, which are equal in number to nuclear protons, that give the atom its chemical properties.
Bohr's model of electrons existing in shells around the nucleus allows for easy visualization of an atom but does not adapt to explain the chemical properties of the elements.
Negatively charged electrons, driven by the laws of quantum mechanics, have unusual properties when they are confined within the electric force field of the positively charged atomic nucleus.
Because of the uncertainty principle, it is not possible to precisely locate the electron, although they are in constant motion around the nucleus. Their motion is random but confined into regions of predictable shape that define the likelihood of finding it in the vicinity of the nucleus.
The interaction between electricity and magnetism is ultimately responsible for the geometry of an electron's motion as described by its wave equation.
The quantum state of an electron in an atom depends upon four factors: its overall energy, its angular momentum, the way its energy changes in response to an external magnetic field and the direction of its spin.
Spin is independent of the other three states. It is either "up" or "down" and refers to one of two possible directions of the electron's angular momentum.
Each of these factors is described by a quantum number, the value of which follows specific rules that are defined by natural laws.
Each quantum state is called an "orbital" that is uniquely defined by the values of three quantum numbers. Electrons can have only two possible spin states, so an orbital can contain only two electrons. (Each set of three quantum numbers corresponds to a uniquely shaped orbital.)
The electronic structure of the atom is due to a principle first recognized by Wolfgang Pauli in 1925, and for which he was awarded the 1945 Nobel Prize in physics. Pauli was among the founders of quantum theory during the first third of the 20th century.
The Pauli exclusion principle states that no two electrons in an atom can have the same four quantum numbers. This forces electrons to "pile up" in an atom to occupy successively higher energy levels, corresponding to the "shells" of the classical planetary model of the atom.
Energy levels and sublevels (also called shells and subshells) are defined by the quantum numbers. For example, electrons with the same energy are in the same shell. Electrons with the same energy and angular momentum are in the same subshell.
Shells are identified by the principal quantum number from one to seven (or K, L, M, etc.). Subshells are identified by the letters s, p, d, f.
IF ALL OF THIS sounds confusing, then it is easy to understand why it took a century and a half of intense study to figure it out.
As confusing as it might seem at first, it is both logical and quantitative, and it is essential for understanding the chemical properties of atoms.
There are two fundamental factors that control an atom's chemical properties. The first is the tendency for an atom to be electrically neutral, containing an equal number of positive and negative charges.
The second factor is that every atom is less reactive when the valence, or outermost shell, is full. This is what allows each type of atom to display its unique chemical properties.
Atoms that have the same number of electrons in the valence shell have similar chemical properties in the sense that they form compounds containing the same number of atoms.
For example, the noble gases (helium, neon, argon, krypton, xenon, radon) have valence shells that are filled and do not react chemically under normal conditions.
Other atoms react with one another in an attempt to fill the valence shells.
Electrons might migrate from one atom to another, or two atoms might share their electrons. In either case the result is strong chemical bonds that produce stable molecules.
picks up where your high school science teacher left off. He is a professor of science at Honolulu Community College, where he teaches earth and physical science and investigates life and the universe. He can be reached by e-mail at firstname.lastname@example.org