Ionic Solids

Type of physical science: Chemistry

Field of study: Chemistry of solids

Because ionic solids all consist of crystals held together by the same kinds of attractive forces, they tend to display physical properties that are distinctive. Among these properties are a marked tendency to dissolve in water, interesting electrical characteristics, and moderate to high melting points.

Overview

Ionic solids, or salts, are crystalline materials composed of positive and negative ions held together by the mutual attraction of opposite charges.

Crystalline solids are commonly classified on the basis of the forces that hold their constituent units together. In this scheme, solids are covalent, molecular, metallic, or ionic, and particular solids of a given type tend to exhibit similar physical properties. Thus, metals are lustrous and good conductors; molecular crystals have low melting points and are insulators; covalent solids are extremely rigid and have high melting points.

As a class, ionic solids exhibit similar properties. They tend to have high melting points (though not as high as covalent solids), and most dissolve readily in water. Their crystals tend to be brittle and are readily cleaved to produce smaller crystals with smooth, flat faces at the cleavage plane. In the solid state, these salts are electrical insulators, yet when melted or dissolved, they are very good conductors.

Sodium chloride, also known as table salt or rock salt, is a well-known ionic solid that exhibits typical properties. Examined closely, a sample of this material is seen to consist of small cubes or rectangles, each of which is a single crystal of the salt. For industrial or laboratory applications, larger single crystals are produced. If one of these crystals is firmly tapped by a sharp cutting edge properly aligned, two smaller single crystals will result. Table salt melts at 801 degrees Celsius, a temperature between the low melting point of the molecular crystal sucrose (sugar), 185 degrees Celsius, and the high melting point of the covalent crystal silicon dioxide (quartz), 1,610 degrees Celsius. Table salt is quite soluble in water (36 grams dissolving in 100 millimeters of water at 0 degrees Celsius) and exhibits the expected electrical behavior, conducting in solution the molten state but insulating in the solid state.

The properties of ionic solids depend upon their molecular structures. The crystal lattices of salts arise from a periodic arrangement of positive and negative ions. Positive ions, cations, are formed when an atom or molecule loses one or more of its electrons, one positive charge arising from each electron lost. Thus, sodium atoms tend to lose one electron to form sodium cations with a +1 charge. Calcium atoms tend to lose two electrons to form calcium cations with a +2 charge. Other atoms found toward the left side of the periodic table, such as sodium or calcium, have a tendency to form cations, the magnitudes of their charges ranging from one to three depending upon the atom. In a parallel fashion, negative ions, anions, are formed when an atom or molecule gains one or more excess electrons, -1 charge arising from each electron gained. Thus, chlorine atoms have a tendency to gain one electron to form chloride anions with a -1 charge. Oxygen atoms tend to gain two electrons to form oxide anions with a -2 charge. Other atoms found toward the right side of the periodic table, such as chlorine or oxygen, have a tendency to form anions, the magnitudes of their charges ranging from one to three, depending upon the atom.

In the solid state, cations and anions aggregate to form a periodic arrangement known as an ionic crystal. Because of its positive charge, each cation draws around it negative ions, anions. In like fashion each anion attracts cations as its nearest neighbors. As this process repeats in three dimensions, a regular repeating arrangement is built up. The strength of the mutual electrostatic attractions holding the crystal together is reflected in the size of the lattice energy, the energy required to separate the constituent ions of the crystal completely. Lattice energies depend upon the sizes of the ions in the crystal, the charges of the constituent ions, and the three-dimensional arrangement of the ions.

Not surprisingly, many of the physical properties of ionic solids depend upon their lattice energies. The melting point is the temperature above which ions gain freedom to move off their crystal lattice sites. Salts with higher lattice energies have higher melting points. Thus, sodium chloride has a moderate lattice energy in comparison to calcium fluoride, and the respective melting points of the two salts reflect this difference: 801 degrees Celsius for sodium chloride, as opposed to 1,423 degrees Celsius for calcium fluoride. Once melted or fused, ionic solids are good electrical conductors, because their constituent ions are free to move (and thus transport charge). In contrast, solid salts are insulators because their ions are locked in place at their lattice sites.

The solubility behavior of ionic solids is also related to lattice energy, though in a more complicated fashion. The same factors that tend to make a lattice more resistant to breaking apart, the factors that tend to increase the lattice energy, also tend to stabilize the ions in water once they have been separated. Thus, the solubility of an ionic crystal does not necessarily correlate with decreasing lattice energy. It is clear, however, that any factor that decreases lattice energy without impeding the ability of water molecules to surround and interact with the separated ions will increase the solubility of a salt. One situation that leads to this kind of behavior is found when there is a large mismatch in the sizes of the positive and negative ions of a crystal. Thus, the perchlorate anion is so large in comparison to most common cations that all common perchlorate salts are soluble.

The detailed structure of a salt depends upon properties of its constituent ions. The simplest property, the ratio of cations to anions, is determined by the requirement that the entire crystal must be electrically neutral, that is, the total positive charge must equal the total negative charge. Thus, in sodium chloride, the number of cations (sodium +1 ions) must be equal to the number of anions (chloride -1 ions). In a crystal like calcium fluoride (fluorite), the number of cations (calcium +2 ions) must be half the number of anions (fluoride -1 ions). Because both cations and anions can assume charges from one to three in magnitude, several different cation-to-anion ratios are commonly found.

Other lattice properties are more complicated. The shape of the crystal lattice depends upon both the sizes of the ions and their relative numbers. Sodium chloride crystallizes in a cubic structure. Each sodium ion is surrounded by six chloride ions in a symmetric arrangement; each of these chloride ions is in turn surrounded by six sodium ions similarly arrayed. This particular structure is often referred to as the rock salt structure and is used by many ionic solids. Another prototype lattice is provided by the structure of fluorite. In this arrangement, each calcium is at the center of a cube, the corners of which are occupied by fluoride ions. Each fluoride is in turn surrounded by four calcium ions at the corners of a tetrahedron. Thus, the ratio of cube corners, eight, to tetrahedron corners, four, preserves the ratio of fluoride to calcium ions at two to one, as required by charge neutrality. In addition to these two rather simple ionic structures, other, more complicated ionic lattices are known.

Ionic solids can be prepared in a variety of ways, but two simple methods are widely applicable: direct combination of the elements, and acid-base neutralization. Sodium chloride, for example, can be prepared in spectacular fashion by mixing molten sodium metal with chlorine gas. As this reaction proceeds, a large quantity of heat is evolved, and the sodium-chlorine interface glows intensely. When reaction ceases, a white powder containing crystals of sodium chloride is found. Sodium chloride can also be prepared in less dramatic fashion by reacting appropriate quantities of hydrochloric acid with sodium hydroxide dissolved in water. This reaction also evolves heat, though not nearly so much. The product, sodium chloride, can be recovered by evaporating off the water. In a similar manner, many salts can be produced by the neutralization reactions of appropriate acids and bases.

Applications

The uses to which ionic solids are put are quite varied and depend more upon the chemical properties of individual salts than upon their generic physical properties. Baking soda (sodium bicarbonate) is used in cooking largely because it decomposes to produce carbon dioxide gas, which leavens baked products. Sodium bicarbonate is also used as an antacid, because it will react with and neutralize excess stomach acid; that is, it acts as a base. Similarly, many metal oxides react with water to form bases. Milk of magnesia is a suspension of magnesium oxide (magnesia) in water and is commonly used as an antacid. An important industrial example of such an oxide is calcium oxide, or lime. Sodium cyanide, a violent poison, is used to extract gold or silver from ores because those metals readily dissolve in aqueous solutions of the salt. Potassium iodide is used in the manufacture of photographic emulsions and is added to table salt to prevent goiter. Indeed, these examples provide only a tiny sampling of the variety of uses of ionic solids.

Some idea of the importance of some of these materials can be obtained by considering the fifty most commonly produced industrial chemicals in the United States. Perennially found on this list are materials such as sodium hydroxide (lye, or caustic soda) and calcium oxide (lime), which are used in chemical manufacture and in water treatment; sodium carbonate (soda ash), which is used in the manufacture of glass; ammonium sulfate and ammonium nitrate, which are used as fertilizers; and calcium chloride, which is used to de-ice roads.

A few ionic solids form crystal lattices resistant enough to fracture and dissolution by water that these salts can function as construction materials. Marble is largely calcium carbonate, as is limestone. Gypsum is a hydrated calcium sulfate. Bone, a natural structural material, is largely composed of carbonates and phosphates of calcium. These examples are somewhat exceptional, however. Typical ionic solids, such as sodium chloride or calcium fluoride, are far too brittle to be generally useful in construction.

A rather interesting if somewhat specialized use to which some ionic solids are applied arises in the measurement and production of infrared radiation, a form of light invisible to the human eye. While most solids absorb infrared radiation, many ionic crystals exhibit high transparency in the infrared region of the spectrum. Consequently, materials such as potassium bromide, sodium chloride, calcium fluoride, and several other ionic solids find use as materials of construction for optical components used in the infrared. Thus, windows, prisms, beamsplitters, and lenses are often fashioned of these materials for use in devices that produce or measure infrared radiation.

Ionic solids are employed in a wide variety of applications, ranging from uses that exploit the properties of individual salts to those that depend upon the characteristics of most ionic crystals. In one form or another, these materials are found in kitchens and on highways, in hospitals and in manufacturing plants, in research laboratories and in living organisms.

Context

Even though salts have been known since prehistoric times, an understanding of their molecular-level structure and the physical basis of their properties did not emerge until the early twentieth century, when theories of solution behavior and chemical bonding came into general acceptance.

In the first half of the nineteenth century, around 1830, Michael Faraday, the famous English chemist, undertook an extensive series of experiments in which he passed electrical current through solutions of a variety of materials. To explain the electrical conductance he observed, Faraday suggested the existence of submolecular particles bearing charge. Some time later, these particles came to be called ions. Faraday proposed that ions were formed and subsequently discharged by the application of an electrical current.

About a half century later, as part of his 1883 doctoral thesis, Svante August Arrhenius proposed that ions are formed by the action of dissolving salts in water. Arrhenius' electrolytic, or ionic, theory, as it came to be called, generated considerable controversy in spite of the fact that it seemed to explain a wide range of puzzling experimental results. The anomalous behavior of salt solutions with regard to the osmotic pressure, studied prominently by Jacobus van't Hoff, and similar anomalies in the freezing points and vapor pressures of salt solutions, investigated by Francois-Marie Raoult, were largely explained by Arrhenius' theory. Moreover, measurements of ionic mobility in solution, performed by Johann Hittorf and later refined by Friedrich Kohlrausch, were compatible with the ionic theory of Arrhenius. Nevertheless, opposition to the ionic theory persisted well into the twentieth century, until the theoretical work of Peter Debye and Erich Huckel (later extended by Lars Onsager) was able to extend Arrhenius' ideas into a comprehensive and quantitative theory of solution behavior.

About the same time that the ionic theory of solution behavior was coming to be widely accepted, the foundations for a modern theory of chemical bonding were being put down.

Between the years 1916 and 1919, the separate efforts of Walther Kossel, Irving Langmuir, and Gilbert Newton Lewis resulted in a successful, if not quite thoroughly modern, model of the forces that hold molecules together. Lewis, in particular, proposed the idea that one type of bond--an ionic, or electrovalent, bond--results from the transfer of one or more electrons from one atom to another. This transfer produces ions of opposite charge, which are mutually attracted to one another. Thus, Faraday's idea that ions were formed by an electrical current, modified by Arrhenius to the view that ions were produced upon dissolution, was changed once again: In Lewis' approach, ions are produced at the time any compound held together by electrovalent bonds is formed. Hence, the ions of such a solid compound must exist before any electrical current is passed, or even before the compound is dissolved in water. The dissolution process merely frees the ions; it imparts mobility (and the ability to conduct electrical current) to the ions that already existed in the solid compound.

Experimental support for the interpretation that arises from the Lewis model of ionic bonding was at hand by the early 1920's. By that time, workers in several laboratories had demonstrated that the electrical conductivity of molten salts was the result of the transport of ionic charge. Even more convincing, perhaps, were the results of X-ray studies initiated in 1912 by William Henry and Lawrence Bragg.

Earlier in 1912, Max von Laue had suggested that X rays could be used to investigate the structures of crystalline solids. Within the year, the Braggs were able to put von Laue's idea to use. They collected and analyzed the X-ray diffraction pattern of sodium chloride and demonstrated convincingly that the distance from a sodium ion to each of its nearest-neighbor chloride ions is the same. This result showed that the crystal could not be composed of aggregated molecules, sodium bound to chlorine. Were that the case, one sodium-to-chlorine distance would be shorter than the others. Rather, the crystal consisted of sodium units surrounded by chloride units, which in turn were surrounded by sodium units, and so on.

With the acceptance of Arrhenius' electrolytic theory, the development of the Lewis ionic theory of bonding, and the emergence of X-ray crystallography, an understanding of ionic solids was placed on a firm foundation. Further refinements, to be sure, were introduced when quantum theory was developed later in the twentieth century. Our understanding of solid-state electrical, magnetic, and optical properties, in particular, continues to grow, and forms the basis of much research in materials science, solid-state chemistry, and solid-state physics.

Nevertheless, in broad outline, the properties of ionic crystals can be understood in terms of ideas proposed by chemists almost a century ago.

The model of an ionic crystal that had emerged by the second decade of the twentieth century is the model still in use today. Nevertheless, the extent to which any particular salt adheres to the model is in many cases an open question. It is well known that many chemical bonds possess partial ionic character. Electrons involved in such bonds are neither completely transferred nor evenly shared. For solids formed from substances bonded in such a fashion, properties intermediate in character between those of an ionic solid and a covalent solid, or between an ionic solid and a metal, might be expected and are indeed observed. Nevertheless, the concept of an ionic solid remains a useful idea that successfully rationalizes the properties of many common solids.

Principal terms

ATOM: the smallest part of an element; atoms are the fundamental building blocks of chemistry

CRYSTAL: a solid in which external surfaces are flat and make definite angles with one another, crystals exhibit long-range order at the molecular level

ELECTRICAL CONDUCTOR: a substance that allows the ready flow of current; conductors possess mobile charge carriers; in a metal, these carriers are electrons

ELECTROLYTE: a substance that dissolves in water to produce a solution that conducts electricity; in solution, current is carried by the movement of ions

ELECTRON: a fundamental particle and the smallest unit of negative charge; electrons are found within atoms

ION: a charged particle; starting from a neutral molecule, positive ions result from removing one or more electrons, negative ions, from adding one or more electrons

LATTICE: a repeating spatial pattern; the long-range order of crystals consists of structural units in a three-dimensional lattice

Bibliography

Cotton, F. A., Geoffrey Wilkinson, and P. L. Gaus. BASIC INORGANIC CHEMISTRY. 2d ed. New York: John Wiley & Sons, 1987. Chapter 4 of this work discusses ionic solids and describes the structures of several common ionic lattices.

Galwey, Andrew K. CHEMISTRY OF SOLIDS. London: Chapman and Hall, 1967. This excellent little book starts at an introductory level and proceeds rapidly to cover many interesting and important aspects of solid-state chemistry. The author assumes the reader has some previous knowledge of chemical bonding.

Jolly, William L. MODERN INORGANIC CHEMISTRY. New York: McGraw-Hill, 1984. This is a standard and well-regarded work that treats a broad subject. Chapter 11 discusses classification of solids.

Mackay, K. M., and R. A. Mackay. INTRODUCTION TO MODERN INORGANIC CHEMISTRY. 4th ed. Englewood Cliffs, N.J.: Prentice-Hall, 1989. This book covers a broad area. Contains a large number of figures that are quite helpful in visualizing the three-dimensional structures of solids. Chapter 5 treats the solid state.

Mortimer, Charles E. CHEMISTRY. 6th ed. Belmont, Calif.: Wadsworth, 1986. This book is one of many general chemistry texts, almost all of which discuss ionic solids in a chapter devoted to liquids and solids (chapter 11). Chapter 7 describes ionic bonding. Recommended for its clear writing and the author's willingness to avoid unnecessary detail.

Wells, A. F. STRUCTURAL INORGANIC CHEMISTRY. 5th ed. Oxford, England: Oxford University Press, 1984. This work is considered the standard reference for structures of inorganic solids.

Concentrations in Solutions

Electrons and Atoms