Ions and Ionization

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IONS AND IONIZATION

CONCEPT

Atoms have no electric charge; if they acquire one, they are called ions. Ions are involved in a form of chemical bonding that produces extremely strong bonds between metals, or between a metal and a nonmetal. These substances, of which table salt is an example, are called ionic compounds. Ionization is the process whereby electrons are removed from an atom or molecule, as well as the process whereby an ionic substance, such as salt, is dissociated into its component ions in a solution such as water. There are several varieties of ionization, including field ionization, which almost everyone has experienced in the form of static electricity. Ion exchange, or the replacement of one ion by another, is used in applications such as water purification, while chemists and physicists use ions in mass spectrometry, to discover mass and structural information concerning atoms and molecules. Another example of ions at work (and a particularly frightening example at that) is ionizing radiation, associated with the radioactive decay following a nuclear explosion.

HOW IT WORKS

Ions: Positive and Negative

Atoms have no electric charge, because they maintain an equal number of protons (positively charged subatomic particles) and electrons, subatomic particles with a negative charge. In certain situations, however, the atom may lose or gain one or more electrons and acquire a net charge, becoming an ion.

Aluminum, for instance, has an atomic number of 13, which tells us that an aluminum atom will have 13 protons. Given the fact that every proton has a positive charge, and that most atoms tend to be neutral in charge, this means that there are usually 13 electrons, with a negative charge, present in an atom of aluminum. Yet like all metals, aluminum is capable of forming an ion by losing electrons—in this case, three.

CATIONS.

Initially, the aluminum atom had a charge of +13 + (−13) = 0; in other words, its charge was neutral due to the equal numbers of protons and electrons. When it becomes an ion, it loses 3 electrons, leaving behind only 10. Now the charge is +13 + (−10) = +3. Thus the remaining aluminum ion is said to have a net positive charge of 3, represented as +3 or 3+. Chemists differ as to whether they represent the plus sign (or the minus sign, in the case of a negatively charged ion) before or after the number. Because both systems of notation are used, these will be applied interchangeably throughout the course of this essay.

When a neutral atom loses one or more electrons, the result is a positively charged ion, or cation (pronounced KAT-ie-un). Cations are usually represented by a superscript number and plus sign: Al⁺³ or Al³⁺, for instance, represents the aluminum cation described above. A cation is named after the element of which it is an ion: thus the ion we have described is either called the aluminum ion, or the aluminum cation.

ANIONS.

When a neutrally charged atom gains electrons, acquiring a negative charge as a result, this type of ion is known as an anion (AN-ie-un). Anions can be represented symbolically in much the same way as cations: Cl₋, for instance, is an anion of chlorine that forms when it acquires an electron, thus assuming a net charge of −1. Note that the 1 is not represented in the superscript notation, much as people do not write 10¹. In both cases, the 1 is assumed, but any number higher than 1 is shown.

The anion described here is never called a chlorine anion; rather, anions have a special nomenclature. If the anion represents, as was the case here, a single element, it is named by adding the suffix -ide to the name of the original element name: chloride. Such is the case, for instance, with a deadly mixture of carbon and nitrogen (CN⁻), better known as cyanide.

Most often the-ide suffix is used, but in the case of most anions involving more than one element (polyatomic anions), as well as with oxyanions (anions containing oxygen), the rules can get fairly complicated. The general principles for naming anions are as follows:

-ide: A single element with a negative charge. Note, however, that both hydroxide (OH⁻) and cyanide (CN⁻) also receive the-ide suffix, even though they involve more than one element.
-ate: An oxyanion with the normal number of oxygen atoms, a number that depends on the nature of the compound. Examples include oxalate (C₂O₄⁻²) or chlorate (ClO₃⁻).
-ite: An oxyanion containing 1 less oxygen than normal. Examples include chlorite (ClO₂⁻).
hypo____ite: An oxyanion with 2 less oxygens than normal, but with the normal charge. An example is hypochlorite, or ClO⁻.
per____ate: An oxyanion with 1 more oxygen than normal, but with the normal charge. Perchlorate, or ClO₄⁻, is an example.
thio-: An anion in which sulfur has replaced an oxygen. Thus, SO₄⁻² is called sulfate, whereas S₂O₃⁻² is called thiosulfate.

Elements and Ion Charges

As one might expect, given the many differences among families of elements on the periodic table, different elements form ions in different ways. Yet precisely because many of these can be grouped into families, primarily according to the column or group they occupy on the periodic table, it is possible to predict the ways in which they will form ions. The table below provides a few rules of thumb. (All group numbers refer to the North American version of the periodic table; see Periodic Table of Elements essay for an explanation of the differences between this and the IUPAC version.)

Alkali metals (Group 1) form 1+ cations. For example, the ion of lithium (Li) is always Li⁺.
Alkaline earth metals (Group 2) form 2+ cations. Thus, beryllium (Be), for instance, forms a Be²⁺ ion.
Most Group 3 metals (aluminum, gallium, and indium) form 3+ cations. The cation of aluminum, thus, is designated as Al₃₊.
Group 6 nonmetals and metalloids (oxygen, sulfur, selenium, and tellurium) form 2− anions. Oxygen, in its normal ionized state, is shown as O²⁻.
Halogens (Group 7) form 1− anions. Fluorine's anion would therefore be designated as Fl⁻.

The metals always form positive ions, or cations; indeed, one of the defining characteristics of a metal is that it tends to lose electrons. However, the many elements of the transition metals family form cations with a variety of different charges; for this reason, there is no easy way to classify the ways in which these elements form cations.

Likewise, it should be evident from the above table that nonmetals, such as oxygen or fluorine, gain electrons to form anions. This, too, is a defining characteristic of this broad grouping of elements. The reasons why these elements—both metals and nonmetals—behave as they do are complex, involving the numbers of valence electrons (the electrons involved in chemical bonding) for each group on the periodic table, as well as the octet rule of chemical bonding, whereby elements typically bond so that each atom has eight valence electrons.

REAL-LIFE APPLICATIONS

Ionic Bonds, Compounds, and Solids

German chemist Richard Abegg (1869-1910), whose work with noble gases led to the discovery of the octet rule, hypothesized that atoms combine with one another because they exchange electrons in such a way that both end up with eight valence electrons. This was an early model of ionic bonding, which results from attractions between ions with opposite electric charges: when they bond, these ions "complete" one another.

As noted, ionic bonds occur when a metal bonds with a nonmetal, and these bonds are extremely strong. Salt, for instance, is formed by an ionic bond between the metal sodium (a +1 cation) and the nonmetal chlorine, a −1 anion. Thus Na⁺ joins with Cl⁻ to form NaCl, or table salt. The strength of the bond in salt is reflected by its melting point of 1,472°F (800°C), which is much higher than that of water at 32°F (0°C).

In ionic bonding, two ions start out with different charges and end up forming a bond in which both have eight valence electrons. In a covalent bond, as is formed between nonmetals, two atoms start out as most atoms do, with a net charge of zero. Each ends up possessing eight valence electrons, but neither atom "owns" them; rather, they share electrons. Today, chemists understand that most bonds are neither purely ionic nor purely covalent; rather, there is a wide range of hybrids between the two extremes. The degree to which elements attract one another is a function of electronegativity, or the relative ability of an atom to attract valence electrons.

IONIC COMPOUNDS AND SOLIDS.

Not only is salt formed by an ionic bond, but it is an ionic compound—that is, a compound containing at least one metal and nonmetal, in which ions are present. Salt is also an example of an ionic solid, or a crystalline solid that contains ions. A crystalline solidis a type of solid in which the constituent parts are arranged in a simple, definite geometric pattern that is repeated in all directions. There are three types of crystalline solid: a molecular solid (for example, sucrose or table sugar), in which the molecules have a neutral electric charge; an atomic solid (a diamond, for instance, which is pure carbon); and an ionic solid.

Salt is not formed of ordinary molecules, in the way that water or carbon dioxide are. (Both of these are molecular compounds, though neither is a solid at room temperature.) The internal structure of salt can be depicted as a repeating series of chloride anions and sodium cations packed closely together like oranges in a crate. Actually, it makes more sense to picture them as cantaloupes and oranges packed together, with the larger chloride anions as the cantaloupes and the smaller sodium cations as the oranges.

If a grocer had to pack these two fruits together in a crate, he would probably put down a layer of cantaloupes, then follow this with a layer of oranges in the spaces between the cantaloupes. This pattern would be repeated as the crate was packed, with the smaller oranges filling the spaces. Much the same is true of the way that chlorine "cantaloupes" and sodium "oranges" are packed together in salt.

This close packing of positive and negative charges helps to form a tight bond, and therefore salt must be heated to a high temperature before it will melt. Solid salt does not conduct electricity, but when melted, it makes for an extremely good conductor. When it is solid, the ions are tightly packed, and thus there is no freedom of movement to carry the electric charge along; but when the structure is disturbed by melting, movement of ions is therefore possible.

Ionization and Ionization Energy

Water is not a good conductor, although it will certainly allow an electric current to flow through it, which is why it is dangerous to operate an electrical appliance near water. We have already seen that salt becomes a good conductor when melted, but this can also be achieved by dissolving it in water. This is one type of ionization, which can be defined as the process in which one or more electrons are removed from an atom or molecule to create an ion, or the process in which an ionic solid, such as salt, dissociates into its component ions upon being dissolved in a solution.

The amount of energy required to achieve ionization is called ionization energy or ionization potential. When an atom is at its normal energy level, it is said to be in a ground state. At that point, electrons occupy their normal orbital patterns. There is always a high degree of attraction between the electron and the positively charged nucleus, where the protons reside. The energy required to move an electron to a higher orbital pattern increases the overall energy of the atom, which is then said to be in an excited state.

The excited state of the atom is simply a step along the way toward ionizing it by removing the electron. "Step" is an appropriate metaphor, because electrons do not simply drift along a continuum from one energy level to the next, the way a person walks gently up a ramp. Rather, they make discrete steps, like a person climbing a flight of stairs or a ladder. This is one of the key principles of quantum mechanics, a cutting-edge area of physics that also has numerous applications to chemistry. Just as we speak of a sudden change as a "quantum leap," electrons make quantum jumps from one energy level to another.

Due to the high attraction between the electron and the nucleus, the first electron to be removed is on the outermost orbital—that is, one of the valence electrons. This amount of energy is called the first ionization energy. To remove a second electron will be considerably more difficult, because now the atom is a cation, and the positive charge of the protons in the nucleus is greater than the negative charge of the electrons. Hence the second ionization energy, required to remove a second electron, is much higher than the first ionization energy.

IONIZATION ENERGIES OF ELEMENTS AND COMPOUNDS.

First and second ionization energy levels for given elements have been established, though it should be noted that hydrogen, because it has only one electron, has only a first ionization energy. In general, the figures for ionization energy increase from left to right along a period or row on the periodic table, and decrease from top to bottom along a column or group.

The reason ionization energy increases along a period is that nonmetals, on the right side of the table, have higher ionization energies than metals, which are on the left side. Ionization energies decrease along a group because elements lower on the table have higher atomic numbers, which means more protons and thus more electrons. It is therefore easier for them to give up one of their electrons than it is for an element with a lower atomic number—just as it would be easier for a millionaire to lose a dollar than it would be for a person earning minimum wage.

For molecules in compounds, the ionization energies are generally related to those of the elements whose atoms make up the molecule. Just as elements with fewer electrons are typically less inclined to give up one, so are molecules with only a few atoms. Thus the ionization energy of carbon dioxide (CO₂), with just three atoms, is relatively high. Conversely, in larger molecules as with larger atoms, there are more electrons to give up, and therefore it is easier to separate one of these from the molecules.

IONIZATION PROCESSES.

A number of methods are used to produce ions for mass spectrometry (discussed below) or other applications. The most common of these methods is electron impact, produced by bombarding a sample of gas with a stream of fast-moving electrons. Though easier than some other methods, this one is not particularly efficient, because it supplies more energy than is needed to remove the electron. An electron gun, usually a heated tungsten wire, produces huge amounts of electrons, which are then fired at the gas. Because electrons are so small, this is rather like using a rapid-fire machine gun to kill mosquitoes: it is almost inevitable that some of the mosquitoes will get hit, but plenty of the rounds will fire into the air without hitting a single insect.

Another ionization process is field ionization, in which ionization is produced by subjecting a molecule to a very intense electric field. Field ionization occurs in daily life, when static electricity builds up on a dry day: the small spark that jumps from the tip of your finger when you touch a doorknob is actually a stream of electrons. In field ionization as applied in laboratories, fine, sharpened wires are used. This process is much more efficient than electron impact ionization, employing far less energy in relation to that required to remove the electrons. Because it deposits less energy into the ion source, the parent ion, it is often used when it is necessary not to damage the ionized specimen.

Chemical ionization employs a method similar to that of electron impact ionization, except that instead of electrons, a beam of positively charged molecular ions is used to bombard and ionize the sample. The ions used in this bombardment are typically small molecules, such as those in methane, propane, or ammonia. Nonetheless, the molecular ion is much larger than an electron, and these collisions are highly reactive, yet they tend to be much more efficient than electron impact ionization. Many mass spectrometers use a source capable of both electron impact and chemical ionization.

Ionization can also be supplied by electromagnetic radiation, with wavelengths shorter than those of visible light—that is, ultraviolet light, x rays, or gamma rays. This process, called photoionization, can ionize small molecules, such as those of oxygen (O₂). Photoionization occurs in the upper atmosphere, where ultraviolet radiation from the Sun causes ionization of oxygen and nitrogen (N₂) in their molecular forms.

In addition to these other ionization processes, ionization can be produced in a fairly simple way by subjecting atoms or molecules to the heat from a flame. The temperatures, however, must be several thousand degrees to be effective, and therefore specialized flames, such as an electrical arc, spark, or plasma, are typically used.

MASS SPECTROMETRY.

As not-ed, a number of the ionization methods described above are used in mass spectrometry, a means of obtaining structure and mass information concerning atoms or molecules. In mass spectrometry, ionized particles are accelerated in a curved path through an electromagnetic field. The field will tend to deflect lighter particles from the curve more easily than heavier ones. By the time the particles reach the detector, which measures the ratio between mass and charge, the ions will have been separated into groups according to their respective mass-to-charge ratios.

When molecules are subjected to mass spectrometry, fragmentation occurs. Each molecule breaks apart in a characteristic fashion, and this makes it possible for a skilled observer to interpret the mass spectrum of the particles generated. Mass spectrometry is used to establish values for ionization energy, as well as to ascertain the mass of substances when that mass is not known. It can also be used to determine the chemical makeup of a substance. Mass spectrometry is applied by chemists, not only in pure research, but in applications within the environmental, pharmaceutical, and forensic (crime-solving) fields. Chemists for petroleum companies use it to analyze hydrocarbons, as do scientists working in areas that require flavor and fragrance analysis.

Ion Exchange

The process of replacing one ion with another of a similar charge is called ion exchange. When an ionic solution—for example, salt dissolved in water—is placed in contact with a solid containing ions that are only weakly bonded within its crystalline structure, ion exchange is possible. Throughout the exchange, electric neutrality is maintained: in other words, the total number of positive charges in the solid and solution equals the total number of negative charges. The only thing that changes is the type of ion in the solid and in the solution.

There are natural ion exchangers, such as zeolites, a class of minerals that contain aluminum, silicon, oxygen, and a loosely held cation of an alkali metal or alkaline earth metal (Group 1 and Group 2 of the periodic table, respectively). When the zeolite is placed in an ionic solution, exchange occurs between the loosely held zeolite cation and the dissolved cation. In synthetic ion exchangers, used in laboratories, a charged group is attached to a rigid structural framework. One end of the charged group is permanently fixed to the frame, while a positively or negatively charged portion kept loose at the other end attracts other ions in the solution.

RESINS AND SEMIPERMEABLE MEMBRANES.

Anion resins and cation resins are both solid materials, but in an anion resin, the positive ions are tightly bonded, while the negative ions are loosely bonded. The negative ions will exchange with negative ions in the solution. The reverse is true for a cation resin, in which the negative ions are tightly bonded, while the loosely bonded positive ions exchange with positive ions in the solution. These resins have applications in scientific studies, where they are used to isolate and collect various types of ionic substances. They are also used to purify water, by removing all ions.

Semipermeable membranes are natural or synthetic materials with the ability to selectively permit or retard the passage of charged and uncharged molecules through their surfaces. In the cells of living things, semipermeable membranes regulate the balance between sodium and potassium cations. Kidney dialysis, which removes harmful wastes from the urine of patients with kidneys that do not function properly, is an example of the use of semipermeable membranes to purify large ionic molecules. When seawater is purified by removal of salt, or otherwise unsafe water is freed of its impurities, the water is passed through a semipermeable membrane in a process known as reverse osmosis.

Ionizing Radiation

We have observed a number of ways in which ionization is useful, and indeed part of nature. But ionizing radiation, which causes ionization in the substance through which it passes, is extremely harmful. It should be noted that not all radiation is unhealthy: after all, Earth receives heat and light from the Sun by means of radiation. Ionizing radiation, on the other hand, is the kind of radiation associated with radioactive fallout from nuclear warfare, and with nuclear disasters such as the one at Chernobyl in the former Soviet Union in 1986.

Whereas thermal radiation from the Sun can be harmful if one is exposed to it for too long, ionizing radiation is much more so, and involves much greater quantities of energy deposited per area per second. In ionizing radiation, an ionizing particle knocks an electron off an atom or molecule in a living system (that is, human, animal, or plant), freeing an electron. The molecule left behind becomes a free radical, a highly reactive group of atoms with unpaired electrons. These can spawn other free radicals, inducing chemical changes that can cause cancer and genetic damage.

WHERE TO LEARN MORE

"The Atom." Thinkquest (Web site). <http://library.thinkquest.org/17940/texts/atom/atom.html> (May 18, 2001).

"Explore the Atom" CERN—European Organization for Nuclear Research (Web site). <http://public.web.cern.ch/Public/SCIENCE/Welcome.html> (May 18, 2001).

Gallant, Roy A. The Ever-Changing Atom. New York: Benchmark Books, 1999.

"Ion Exchange Chromatography" (Web site). <http://ntri.tamuk.edu/fplc/ion.html> (June 1, 2001).

"Just After Big Bang, Colliding Gold Ions Exploded." UniSci: Daily University Science News (Web site). <http://unisci.com/stories/20012/0501012.htm> (June 1, 2001).

"A Look Inside the Atom" (Web site). <http://www.aip.org/history/electron/jjhome.htm> (May 18, 2001).

"Portrait of the Atom" (Web site). <http://www.inetarena.com/~pdx4d/snelson/Portrait.html> (May 18, 2001).

"A Science Odyssey: You Try It: Atom Builder." PBS—Public Broadcasting System (Web site). <http://www.pbs.org/wgbh/aso/tryit/atom/> (May 18, 2001).

"Table of Inorganic Ions and Rules for Naming Inorganic Compounds." Augustana College (Web site.) <http://inst.augie.edu/~dew/242mat.htm> (June 1, 2001).

"Virtual Chemistry Lab." University of Oxford Department of Chemistry (Web site). <http://www.chem.ox.ac.uk/vrchemistry/complex/menubar.html> (June 1, 2001).

KEY TERMS

ANION:

The negatively charged ion thatresults when an atom gains one or more electrons. An anion (pronounced "AN-ie-un") of a single element is named by adding the suffix -ide to the name of the original element—hence, "chloride." Other rules apply for more complex anions.

CATION:

The positively charged ion thatresults when an atom loses one or more electrons. A cation (pronounced KAT-ieun) is named after the element of which it is an ion and, thus, is called, for instance, the aluminum ion or the aluminum cation.

CRYSTALLINE SOLID:

A type of solid in which the constituent parts have a simple and definite geometric arrangement that is repeated in all directions. Types of crystalline solids include atomic solids, ionic solids, and molecular solids.

ELECTRON:

Negatively charged particles in an atom. The number of electrons and protons in an atom is the same, thus canceling out one another. When an atom loses or gains one or more electrons, however—thus becoming an ion—it acquires a net electric charge.

EXCITED STATE:

A term describing thecharacteristics of an atom that has acquired excess energy.

GROUND STATE:

A term describing the characteristics of an atom that is at its ordinary energy level.

ION:

An atom or group of atoms that has lost or gained one or more electrons, and thus has a net electric charge. There are two types of ions: anions and cations.

ION EXCHANGE:

The process of replacing one ion with another of similar charge.

IONIC BONDING:

A form of chemical bonding that results from attractions between ions with opposite electric charges.

IONIC COMPOUND:

A compound (two or more elements chemically bonded to one another), in which ions are present. Ionic compounds contain at least one metal and nonmetal, joined by an ionic bond.

IONIC SOLID:

A form of crystalline solid that contains ions. When mixed with a solvent such as water, ions from table salt—an example of an ionic solid—move freely throughout the solution, making it possible to conduct an electric current.

IONIZATION:

A term that refers to two different processes. In one kind of ionization, one or more electrons are removed from an atom or molecule to create an ion. A second kind of ionization occurs whenan ionic solid, such as salt dissociates into its component ions upon being dissolved in a solution.

IONIZATION ENERGY:

The amount of energy required to achieve ionization in which one or more electrons are removed from an atom or molecule. Another term for this is ionization potential.