Electron Affinity

Type of physical science: Chemistry

Field of study: Chemistry of the elements

The electron affinity measures how strongly neutral atoms and molecules attract an additional electron. It shows the periodic behavior typical of various chemical properties and plays an important part in the analysis of the energies of many chemical changes.

Overview

Negative and positive ions are of equal importance in the chemistry of solutions. Any substance that dissolves to form ions must have equal amounts of negative and positive charge on these ions. In gases, positive ions are more familiar than negative ions, and the former have been easier to make and study in the laboratory.

Negative gaseous ions are formed when a gaseous atom or molecule captures a free electron. The attachment of an electron to an atom or molecule typically creates an excess of energy, and the newly created negative ion must dispose of this energy. This might cause the breaking of a bond between two atoms in a molecule. One molecular fragment will then carry a negative charge. Lone atoms that capture an electron cannot dispose of the excess energy in this way.

Another method of forming negative gaseous ions is by the breaking of a bond in such a way that one newly formed fragment holds both the electrons from the bond; this fragment will then have a negative charge. The other fragment will have a positive charge. Lightning puts much energy into the molecules of the air and forms many gaseous ions, both positive and negative. Once a negative ion is formed, it can react with a neutral molecule, and the reaction products will include a new negative ion.

The simplest negative ions are made from atoms. Given, for example, the importance of the chloride ion in solution, it is logical to make and study it as a gas. The neutral chlorine atom attracts an additional electron; if it is to hold the electron, it must give off the energy of attraction. Conversely, energy must be put into the typical negative ion in order to remove an electron. This energy is the electron affinity; the energy needed to remove an electron from a negative gaseous ion is clearly a measure of the ability of the neutral atom or molecule to hold an electron to itself. The more tightly an electron is held, the more energy will be required to remove it. If the electron affinity is zero, the atom cannot hold the added electron that it otherwise would have held alongside its own electrons.

Many gaseous molecules also attract an electron to become negative ions. Molecules generally have an easier time disposing of the energy of attraction than do atoms. This is particularly true as the number of atoms in the molecule increases.

Positive ions are formed by removing electrons from neutral atoms or molecules. In the case of atoms, the number of electrons removed depends on the amount of energy available.

Thus, ions having a charge of plus two or plus three or higher can be made; however, one generally cannot make negative ions with a charge greater than minus one, which corresponds to the addition of one electron to the neutral atom. This can be understood by taking into account two factors. The first is that like charges repel one another. The second is that the electrons, carriers of unit negative charge, are in constant motion in the atom and free to move away from one another. If an electron is to be added to a negative gaseous ion, that electron will feel the repulsion of the negative charge already on the ion. For atomic and small molecular ions, this repulsion is greater than the attraction, so that a second electron cannot be added, and small gaseous ions of charge 2 are not observed. The oxide ion of charge 2 can exist only in ionic crystals, where positive charges are arranged around it. The attraction of oppositely charged ions holds the crystal together by making its energy lower.

The electron affinity is a measure of how tightly a neutral atom or molecule is able to hold an additional electron. This makes it of great interest in analyzing the changes that ionic substances undergo. Electron affinities are commonly reported in electronvolts, a unit that gives values in the range of 0 to 4 for most atoms and molecules. A few atoms have negative electron affinities, which means that they repel rather than attract an additional electron. This change of sign has no parallel among ionization energies, which are always positive.

Electron affinities are measured by various methods. One, called photodetachment, uses light as an energy source to remove the added electron from the gaseous ion. The electron affinity is equal to the energy of the longest-wavelength light that is able to remove the electron.

A second method, known as photoelectron spectroscopy, is similar to the first, except that the electron is removed by light of convenient energy, and the extra energy carried off by the electron is measured. The electron affinity is the difference between the energy supplied by the light and the energy carried off by the electron. When other necessary information is at hand, electron affinities can also be found by analyzing the energy involved in reactions between negative gaseous ions and atoms or molecules.

Although the importance of electron affinities was recognized long before there were good techniques for measuring them, they are conceptually simple. A negative ion is made from a gaseous atom or molecule, and an electron is then removed to make the neutral atom or molecule again in an experiment that yields the required energy.

The concept of electron affinity presents some difficulties. The definition given above is clear, and it is consistent with other definitions. Note, however, that it assumes the existence of the negative ion. Some scientists prefer to think of electron affinity in terms of a process that starts with the neutral atom or molecule, to which an electron is added. As discussed above, the atom or molecule must dispose of the energy it gains as it attaches the electron to itself. This makes a negative change in the energy, and electron affinities are negative in some tabulations.

But there is reluctance to define a set of mostly negative values. To circumvent this, scientists define the electron affinity as the energy given off, making it a positive quantity for most substances.

Some physicists define electron affinity as the attraction a metal has for the electrons that conduct electricity in it. This establishes electron affinity values for some solids. But the definition is not universally used among physicists.

Applications

Electron affinities have an intrinsic importance in chemistry; the attraction of an atom or molecule for an electron is fundamental. However, the primary interest in electron affinities has long been because of their place in a Born-Haber cycle analysis. Briefly, this analysis seeks to understand the energy of a crystal in terms of all the energies that contribute to it. When some process or change is followed around a cycle, that is, through a series of steps that ends where it started, the sum of all the energy changes must be zero. This is a property of the energy, and not of a particular cycle or kind of analysis.

In the case of ionic crystals, the cycle can start with the elements that react to make the crystal, and this energy change can be measured in the laboratory. The crystal consists of alternating positive and negative ions. They are strongly held together by the attraction of opposite charges, and the energy of charged bodies separated by a known distance can be calculated quite well, if not perfectly. The next step in the cycle is to suppose that all the ions move far from one another, so that the attraction between them is effectively zero. This step cannot be carried out in the laboratory, so the energy is calculated. With the ions far from one another, the next step is to move electrons from the negative ions to the positive ions to make neutral atoms. The energy needed to remove an electron from a gaseous negative ion is the electron affinity. The energy change when an electron is attached to a gaseous positive ion is simply the negative of the ionization energy. The next step around the cycle is to return these gaseous atoms to their standard forms as elements. For the metal in the crystal, this will mean converting it from gaseous atoms to atoms in a solid. For the simplest crystals, the other element is a gas such as chlorine or oxygen, and the energy corresponding to its combination into diatomic molecules must be known.

The steps described above make a complete cycle, in which the elements that react at first are formed again at the end. Because the sum of all energies around the cycle must be zero, add the known or calculated energies and any departure from zero indicates an error. Note, however, that the electron affinity or the ionization energy of an element will be the same in all such Born-Haber cycles. Thus, data concerning the elements can be taken from one cycle to another, and suspected errors can be checked by insertion into other sets of cyclic data.

If the least certain quantity among the energies around the cycle is the energy needed to separate the ions in the crystal from one another, that value can be adjusted so that the sum of all values around the cycle is zero. This assumes that the electron affinities are known quite accurately, and the demand for such values has created a long-standing interest in their measurement.

Broadly considered, the ability of an atom to attract electrons increases in going up a column of the periodic table, and in going from left to right along a row of the chart. It is in these directions that, with some exceptions, the elements become smaller, and electrons near the atom are closer to the positively charged nucleus that attracts them. The electron affinities of the elements follow this trend roughly, with some interesting exceptions. For example, values for all elements starting with beryllium (atomic number 4) at the top of the column, for all the elements starting with zinc (atomic number 30) at the top of the column, and for all elements starting with helium (atomic number 2) are zero or even slightly negative according to the definition stated above. Each of the atoms in these two columns has no trouble holding its own electrons. The total lack of attraction for an extra electron is explained by noting that each element in these two groups must put the added electron into an orbital at a substantially higher energy than those used for the electrons of the neutral atom. But attractive energies are low, that is, negative, energies. Thus, none of these atoms can attract an extra electron strongly enough to hold it.

The electron affinity of hydrogen is 0.754 electronvolt. The element is sometimes placed at the top of the first column of the periodic table. Its electron affinity is greater than that of any other element in that column. This is consistent with standard chemical reasoning about the periodicity of chemical properties; hydrogen is less metallic than any other element in that column. Note, however, that each of the metals in the first column has at least some attraction for an electron, even though these elements all form ions in solution in which the neutral atom loses an electron.

Hydrogen is sometimes placed at the top of the second column from the right in the periodic table. Considered there, its electron affinity is much less than that of any of the elements directly under it. This, too, is consistent with the behavior of the elements in that column, fluorine, chlorine, bromine, and iodine; all of them act as strong attractors throughout chemistry, and their electron affinities are all above 3 electronvolts. The attraction of the neutral hydrogen atom for an electron is thus intermediate in comparison to other elements.

Nitrogen, with seven electrons in the neutral atom, has zero electron affinity. This surprising result is understood by noting that in the neutral atom, the three 2p orbitals each have one electron in them, and the addition of another electron to any of these half-filled orbitals is not easy in the lone atom. This point is discussed below where the electron affinity is compared to the electronegativity.

Oxygen has an electron affinity of 1.46 electronvolts, which places it well below the strongly electron attracting atoms fluorine, chlorine, and so forth, discussed above. Oxygen is one of the commonest of the elements and forms compounds with most other elements.

Electron affinities of most metals are less than one electronvolt, but a few metals, notably nickel, copper, and silver, have values above 1; gold and platinum have values above 2 electronvolts. This fits well with their general chemical behavior; they hold onto their electrons more tightly than do most metals.

The simplest molecules consist of two atoms of the same element. When the electron affinities of atoms are compared to those of such molecules, the atom usually has the greater electron affinity. Hydrogen is an extreme case. The electron affinity of the atom is 0.754 electronvolt. For the pair of hydrogen atoms that make up the molecule, no value is given, and it is assumed to be zero. Like the atoms of the last column in the periodic table, the hydrogen molecule has no place to put another electron at any energy that could hold the electron. The electron affinities of most two-atom molecules are less than the values for the lone atoms.

Exceptions are carbon and iron: For each of them, a pair of atoms holds an electron much more tightly than does one atom.

Some trends in electron affinities of molecules are not easily explained, but others are clearly related to the values from the atoms in them. As an example, consider magnesium, which has zero electron affinity. If a hydrogen atom is added to it, the value is 1.05 electronvolts; if, instead, one magnesium and one chlorine combine in a molecule, the electron affinity is 1.589.

The increase parallels that of the atoms hydrogen and oxygen.

Some data exist that make it possible to compare electron affinities of atoms with those of two-atom and three-atom molecules of the same element. For oxygen alone, and with two and three atoms joined, the value is least for two atoms. Sulfur acts similarly, but carbon shows the greatest value of the electron affinity when two atoms are joined. It is the task of scientists who study atomic structure to explain such differences in behavior.

Context

It is instructive to compare ionization energies and electron affinities for the same atom. One rule holds in every case: The ionization energy is always greater than the electron affinity for a given atom or molecule. This is a consequence of the charges in an atom acting on one another. The electrons, carriers of negative charge, are held to the atom by the attraction of the positive charge on the nucleus; opposite charges attract. Electrons in an atom repel one another; like charges repel. Thus, when an electron is attached to an atom to make a negative ion, the electrons present in the neutral atom repel it. Exactly the same argument holds when an atom loses an electron to become a positively charged ion, but this time the total number of electrons is one less, and the repulsions among them are less. This line of reasoning is sufficient to explain the observation that ionization energies are always greater than electron affinities for the same atom. Another reason has to do with the arrangement of electrons in atoms, as discussed in the next paragraph.

A concept that is not directly measurable is the electronegativity. It is the attraction an atom has for a pair of electrons that it shares with a neighboring atom. Typically, one of the pair of electrons comes from each of the atoms involved. An atom such as nitrogen has a quite high electronegativity, though its electron affinity is zero. This apparent contradiction is resolved by noting that nitrogen has a considerable attraction for its own electrons, as shown by its high ionization energy. Moreover, nitrogen easily shares electrons with several neighboring atoms simultaneously, and this alters the arrangement and the energies of the electrons around it.

Ions in solution are more extensively studied than ions among gaseous molecules. The latter have the advantage of conceptual simplicity, but the former are easier to make. Many substances, when dissolved in water, separate into ions. This is the case because water molecules arrange themselves around the ions in ways that make the ions stable in solution. The action of water with respect to ions dissolved in it makes the environments and energies of dissolved ions far more complicated than those of ions among gases.

Principal terms

ATOM: the smallest and simplest unit of an element that can exist

BORN-HABER CYCLE: a sequence of changes that brings an ionic substance back to its original state; for such a cycle, some energies will be positive and some negative, and they must add to zero

ELECTRON: the smallest unit of negative electrical charge; an atom or molecule that has lost an electron will have a positive charge; an atom or molecule that has gained an electron will have a negative charge

ELECTRON AFFINITY: the change in energy when an electron is removed from a gaseous negative ion; the measured energy change is usually regarded as the electron affinity of the neutral atom or molecule that results when the electron is removed

ELECTRONVOLT: the energy gained by an electron when it falls through a voltage difference of 1 volt; a secondary unit in the SI system

ION: an atom or molecule that has gained or lost an electron

IONIZATION ENERGY: the change in energy when an electron is removed from a gaseous atom or molecule; "ionization potential" is an older synonym

MOLECULE: two or more atoms that are held together strongly enough to be identified and studied as a unit

PHOTODETACHMENT: the removal of an electron from a gaseous negative ion as a result of the absorption of light

Bibliography

Bailar, John C., Jr., et al. CHEMISTRY. San Diego: Harcourt Brace Jovanovich, 1989. A brief presentation that includes a detailed figure and statements on systematic aspects of electron affinities. Again, the sign convention is opposite to the one used herein.

Cotton, F. Albert, and Geoffrey Wilkinson. ADVANCED INORGANIC CHEMISTRY: A COMPREHENSIVE TEXT. 5th ed. New York: Wiley, 1988. All editions of this landmark book discuss electron affinities and show their application to the analysis of energy cycles in crystals.

Kotz, John C., and K. F. Purcell. CHEMISTRY AND CHEMICAL REACTIVITY. Philadelphia: Saunders College Publishing, 1987. These authors define electron affinity so as to give it a negative sign, but the figures are helpful and the context is good.

Myers, R. Thomas. "The Periodicity of Electron Affinity." JOURNAL OF CHEMICAL EDUCATION 67 (April, 1990): 307-308. This article gives data for the majority of elements and pays particular attention to those that are difficult to understand. The author boldly suggests that a few inexplicable data might be erroneous.

Weast, Robert C., ed. CRC HANDBOOK OF CHEMISTRY AND PHYSICS. 1st Student ed. Boca Raton, Fla.: CRC Press, 1988. The sixty-seventh and later editions offer tables of electron affinity data for atoms and molecules that far surpass the atomic data found in most texts.

Electrons and Atoms

The Periodic Table and the Atomic Shell Model