Valence Shell

FIELDS OF STUDY: Organic Chemistry; Inorganic Chemistry; Physical Chemistry

ABSTRACT

The characteristics of the valence electron shells of atoms are discussed. The valence shell holds the valence electrons of an atom. The number of valence electrons determines the chemical behavior of the atom and the number and type of chemical bonds it can form with other atoms.

The Modern Atomic Model

The electron was first identified as a charged particle by British physicist J. J. Thomson (1856–1940) in 1897, and the proton was discovered in 1917 by British physicist Ernest Rutherford (1871–1937). These discoveries revealed that atoms are not the indivisible absolutes that they were previously believed to be; rather, they are composite entities containing smaller, "subatomic" particles. The existence of the third essential subatomic particle, the neutron, was demonstrated by a third British physicist, James Chadwick (1891–1974), in 1932. With this last discovery, the observed behavior of the atom seemed to agree with the theoretical calculations of the developing field of quantum mechanics.

In simple terms, quantum mechanics describes the behavior of subatomic particles, such as electrons, on an infinitesimal scale. At the time, electrons were envisioned as being bound to the nucleus in the same way that the moon is bound to Earth, and terms such as "orbit" and "orbital" are still used to describe the relationship of electrons to the nucleus. However, it is now understood that electrons, like light, exhibit wave-particle duality; that is, they behave as both physical, massive (in the sense of having mass) particles and massless waves. This is perhaps the most difficult aspect of the modern atomic model to grasp, but it is fundamental to the behavior of electrons in atoms.

Early observations showed that electrons in an atom can have only very specific energies and that the transition of an electron from a lower energy level to a higher one can only occur when the electron acquires a specific amount of energy. Conversely, when the electron drops from a higher energy level to a lower one, it has to give up exactly that same amount of energy. The different energy levels that electrons can occupy in an atom are termed electron shells, and each one is identified by its principal quantum number (n), so that the first electron shell is the n = 1 shell, the second is the n = 2 shell, and so on. Each shell is divided into subshells that can hold only a specific number of one type of orbital (s, p, d, or f), and each orbital can contain no more than two electrons at any time.

In any neutral atom, the number of electrons is equal to the atomic number of the element, which is simply the number of protons in the nucleus; in other words, the numbers of electrons and protons are equal. The electrons are ordered in pairs in the orbitals of successive electron shells. The outermost shell of most of the noble gases (helium, neon, argon, krypton, xenon, and radon) contains a full complement of eight electrons; the exception is helium, which has only two electrons total. All other atoms have fewer than eight electrons in their outermost shells. The number of electrons in this shell determines the maximum number of bonds that a particular atom can form with other atoms. This is known as the valence of the atom, and the outermost (in most cases) electron shell is correspondingly called the valence shell—or, more accurately, its outermost electrons are known as valence electrons. (In most elements, valence electrons only occupy the outermost electron shell, but for the group of elements known as transition metals, they may also be found in inner shells.)

Valence Electrons and the Periodic Table

The number of electrons in the valence shell of an atom determines the chemical behavior of that atom. An atom is at an energy minimum, and therefore at its most stable, when the valence shell is completely filled with the eight electrons it is allowed to contain. This is the basis of the octet rule, and it is also why the noble gases are the most chemically inert, or nonreactive, elements. Elements that have six or seven electrons in their valence shell readily accept extra electrons to achieve the electron configuration of the next noble-gas element in the periodic table, while elements with only one or two electrons in their valence shell readily donate those electrons to achieve the electron configuration of the noble-gas element that precedes them. The more electrons that an atom has in its valence shell, the more electronegative it is said to be, meaning that it easily accepts extra electrons from other atoms. Conversely, atoms that easily give up their valence electrons are said to be electropositive. The energy required to remove an electron from the valence shell to produce a cation, or positively charged ion, is called the ionization energy. The energy necessary to remove the first electron from the neutral atom is called the first (or initial) ionization energy, the energy required to remove a second electron is the second ionization energy, and so on. Removing two electrons from the neutral atom at once requires energy equivalent to the sum of the first and second ionization energies.

The elements of the periodic table are arranged in vertical columns called groups and horizontal rows called periods. The periods are arranged in order of increasing principal quantum number, while the groups are arranged in order of their valence-shell electrons. The valence shell of the first two groups is the s subshell of the electron shell identified by the corresponding principal quantum number. (Although the letters s, p, d, and f technically identify the individual orbitals, each of which can hold only two electrons, they are often used to refer to the subshell containing the orbitals as well.) Group 1 elements (the lithium group) have just one electron in this orbital and very readily give it up to form the corresponding cation, which has a 1+ charge. When this happens, the valence shell of the ion has the same electron configuration as the noble-gas element immediately before it. For example, the electron distribution in the valence shell of the sodium cation (Na⁺), which has the atomic number 11, is the same as that of the neon atom, which has the atomic number 10. The group 1 elements have the lowest ionization energy of all the elements. Group 2 elements (the beryllium group) have two electrons in their valence s orbital, which they can give up almost as easily to form cations with a 2+ charge.

The last six groups (13–18) have their valence electrons in the p subshell, which contains three p orbitals for a maximum possible total of six electrons. The elements in groups 13, 14, and 15 (the boron, carbon, and nitrogen groups) neither accept nor donate electrons readily to form ions. These elements most commonly form covalent bonds with other atoms by sharing unpaired electrons, which are single valence electrons that inhabit an orbital alone and are not part of a chemical bond. An unpaired electron can form a covalent bond by sharing the orbital of an unpaired electron on another atom (at which point, of course, they cease to be called unpaired electrons). The group 16 elements (the oxygen group) readily accept two electrons to form anions, or negatively charged ions, with a 2− charge. The group 17 elements (the halogens, also called the fluorine group) are the most electronegative of the elements and very quickly accept an electron from another atom to form a halide ion with a 1− charge. Finally, apart from helium, the noble-gas elements (group 18) have two s electrons and six p electrons, making a complete octet. In all cases, whether through the formation of ions or by sharing electrons covalently, an atom will achieve a complete valence-shell electron configuration.

In between these extremes are the transition metals, including the lanthanides and the actinides. The valence electrons of these elements are distributed between their d and f orbitals. These orbitals are close together in energy, much more numerous than s and p orbitals (the subshells contain five and seven orbitals, respectively), and do not conform to the octet rule.

Oxidation States

A useful method for tracking valence electrons is to determine each atom’s specific oxidation state. The oxidation state of an atom indicates the number of electrons it has either donated to or accepted from other atoms in order to form chemical bonds. This is especially useful in balancing redox (reduction-oxidation) reactions.

PRINCIPAL TERMS

• atomic model: a theoretical representation of the structure and behavior of an atom based on the nature and behavior of its component particles.
• electron: a fundamental subatomic particle with a single negative electrical charge, found in a large, diffuse cloud around the nucleus.
• electronegative: describes an atom that tends to accept and retain electrons to form a negatively charged ion.
• ionization energy: the amount of energy required to remove an electron from an atom in a gaseous state.
• octet rule: the tendency of atoms when bonding to either accept or donate electrons in such a way that they end up with eight electrons in the outermost electron shell.
• oxidation state: a number that indicates the degree to which an atom or ion in a chemical compound has been oxidized or reduced.
• unpaired electron: a valence electron that occupies an orbital by itself and is not involved in the formation of a chemical bond.

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