Crystal


Crystal
/kris"tl/, n.
1. a city in SE Minnesota, near Minneapolis. 25,543.
2. a female given name.

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I
Any solid material whose atoms are arranged in a definite pattern and whose surface regularity reflects its internal symmetry.

Each of a crystal's millions of individual structural units (unit cells) contains all the substance's atoms, molecules, or ions in the same proportions as in its chemical formula (see formula weight). The cells are repeated in all directions to form a geometric pattern, manifested by the number and orientation of external planes (crystal faces). Crystals are classified into seven crystallographic systems based on their symmetry: isometric, trigonal, hexagonal, tetragonal, orthorhombic, monoclinic, and triclinic. Crystals are generally formed when a liquid solidifies, a vapour becomes supersaturated (see saturation), or a liquid solution can no longer retain dissolved material, which is then precipitated. Metals, alloys, minerals, and semiconductors are all crystalline, at least microscopically. (A noncrystalline solid is called amorphous.) Under special conditions, a single crystal can grow to a substantial size; examples include gemstones and some artificial crystals. Few crystals are perfect; defects affect the material's electrical behaviour and may weaken or strengthen it. See also liquid crystal.
II
(as used in expressions)
Crystal Night

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Introduction

      any solid material in which the component atoms are arranged in a definite pattern and whose surface regularity reflects its internal symmetry.

Classification
      The definition of a solid appears obvious; a solid is generally thought of as being hard and firm. Upon inspection, however, the definition becomes less straightforward. A cube of butter, for example, is hard after being stored in a refrigerator and is clearly a solid. After remaining on the kitchen counter for a day, the same cube becomes quite soft, and it is unclear if the butter should still be considered a solid. Many crystals behave like butter in that they are hard at low temperatures but soft at higher temperatures. They are called solids at all temperatures below their melting point. A possible definition of a solid is an object that retains its shape if left undisturbed. The pertinent issue is how long the object keeps its shape. A highly viscous fluid retains its shape for an hour but not a year. A solid must keep its shape longer than that.

Basic units of solids
      The basic units of solids are either atoms (atom) or atoms that have combined into molecules. The electrons (electron) of an atom move in orbits that form a shell structure around the nucleus. The shells are filled in a systematic order, with each shell accommodating only a small number of electrons. Different atoms have different numbers of electrons, which are distributed in a characteristic electronic structure of filled and partially filled shells. The arrangement of an atom's electrons determines its chemical properties. The properties of solids are usually predictable from the properties of their constituent atoms and molecules, and the different shell structures of atoms are therefore responsible for the diversity of solids.

      All occupied shells of the argon (Ar) atom, for example, are filled, resulting in a spherical atomic shape. In solid argon the atoms are arranged according to the closest packing of these spheres. The iron (Fe) atom, in contrast, has one electron shell that is only partially filled, giving the atom a net magnetic moment. Thus, crystalline iron is a magnet. The covalent bond between two carbon (C) atoms is the strongest bond found in nature. This strong bond is responsible for making diamond the hardest solid.

Long- and short-range order
      A solid is crystalline if it has long-range order. Once the positions of an atom and its neighbours are known at one point, the place of each atom is known precisely throughout the crystal. Most liquids lack long-range order, although many have short-range order. Short range is defined as the first- or second-nearest neighbours of an atom. In many liquids the first-neighbour atoms are arranged in the same structure as in the corresponding solid phase. At distances that are many atoms away, however, the positions of the atoms become uncorrelated. These fluids, such as water, have short-range order but lack long-range order. Certain liquids may have short-range order in one direction and long-range order in another direction; these special substances are called liquid crystals. Solid crystals have both short-range order and long-range order.

      Solids that have short-range order but lack long-range order are called amorphous (amorphous solid). Almost any material can be made amorphous by rapid solidification from the melt (molten state). This condition is unstable, and the solid will crystallize in time. If the timescale for crystallization is years, then the amorphous state appears stable. Glasses are an example of amorphous solids. In crystalline silicon (Si) each atom is tetrahedrally bonded to four neighbours. In amorphous silicon (a-Si) the same short-range order exists, but the bond directions become changed at distances farther away from any atom. Amorphous silicon is a type of glass. Quasicrystals are another type of solid that lack long-range order.

      Most solid materials found in nature exist in polycrystalline (polycrystal) form rather than as a single crystal. They are actually composed of millions of grains (small crystals) packed together to fill all space. Each individual grain has a different orientation than its neighbours. Although long-range order exists within one grain, at the boundary between grains, the ordering changes direction. A typical piece of iron or copper (Cu) is polycrystalline. Single crystals of metals are soft and malleable, while polycrystalline metals are harder and stronger and are more useful industrially. Most polycrystalline materials can be made into large single crystals after extended heat treatment. In the past blacksmiths would heat a piece of metal to make it malleable: heat makes a few grains grow large by incorporating smaller ones. The smiths would bend the softened metal into shape and then pound it awhile; the pounding would make it polycrystalline again, increasing its strength.

Categories of crystals
      Crystals are classified in general categories, such as insulators, metals, semiconductors, and molecular solids. A single crystal of an insulator is usually transparent and resembles a piece of glass. Metals are shiny unless they have rusted. Semiconductors are sometimes shiny and sometimes transparent but are never rusty. Many crystals can be classified as a single type of solid, while others have intermediate behaviour. Cadmium sulfide (CdS) can be prepared in pure form and is an excellent insulator; when impurities are added to cadmium sulfide, it becomes an interesting semiconductor. bismuth (Bi) appears to be a metal, but the number of electrons available for electrical conduction is similar to that of semiconductors. In fact, bismuth is called a semimetal. Molecular solids are usually crystals formed from molecules or polymers. They can be insulating, semiconducting, or metallic, depending on the type of molecules in the crystal. New molecules are continuously being synthesized, and many are made into crystals. The number of different crystals is enormous.

Structure
      Crystals can be grown under moderate conditions from all 92 naturally occurring elements (chemical element) except helium, and helium can be crystallized at low temperatures by using 25 atmospheres of pressure. Binary crystals are composed of two elements. There are thousands of binary crystals; some examples are sodium chloride (NaCl), alumina (Al2O3), and ice (H2O). Crystals can also be formed with three or more elements.

The unit cell
 A basic concept in crystal structures is the unit cell. It is the smallest unit of volume that permits identical cells to be stacked together to fill all space. By repeating the pattern of the unit cell over and over in all directions, the entire crystal lattice can be constructed. A cube is the simplest example of a unit cell. Two other examples are shown in Figure 1—>. The first is the unit cell for a face-centred cubic lattice, and the second is for a body-centred cubic lattice. These structures are explained in the following paragraphs. There are only a few different unit-cell shapes, so many different crystals share a single unit-cell type. An important characteristic of a unit cell is the number of atoms it contains. The total number of atoms in the entire crystal is the number in each cell multiplied by the number of unit cells. Copper and aluminum (Al) each have one atom per unit cell, while zinc (Zn) and sodium chloride have two. Most crystals have only a few atoms per unit cell, but there are some exceptions. Crystals of polymers, for example, have thousands of atoms in each unit cell.

Structures of metals
  The elements are found in a variety of crystal packing arrangements. The most common lattice structures for metals are those obtained by stacking the atomic spheres into the most compact arrangement. There are two such possible periodic arrangements. In each, the first layer has the atoms packed into a plane-triangular lattice in which every atom has six immediate neighbours. Figure 2—> shows this arrangement for the atoms labeled A. The second layer is shaded in the figure. It has the same plane-triangular structure; the atoms sit in the holes formed by the first layer. The first layer has two equivalent sets of holes, but the atoms of the second layer can occupy only one set. The third layer, labeled C, has the same structure, but there are two choices for selecting the holes that the atoms will occupy. The third layer can be placed over the atoms of the first layer, generating an alternate layer sequence ABABAB . . ., which is called the hexagonal- closest-packed (hcp) structure. Cadmium and zinc crystallize with this structure. The second possibility is to place the atoms of the third layer over those of neither of the first two but instead over the set of holes in the first layer that remains unoccupied. The fourth layer is placed over the first, and so there is a three-layer repetition ABCABCABC . . ., which is called the face-centred cubic (fcc), or cubic-closest-packed, lattice. Copper, silver (Ag), and gold (Au) crystallize in fcc lattices. In the hcp and the fcc structures the spheres fill 74 percent of the volume, which represents the closest possible packing of spheres. Each atom has 12 neighbours. The number of atoms in a unit cell is two for hcp structures and one for fcc. There are 32 metals that have the hcp lattice and 26 with the fcc. Another possible arrangement is the body-centred cubic (bcc) lattice, in which each atom has eight neighbours arranged at the corners of a cube. Figure 3A—> shows the cesium chloride (CsCl) structure, which is a cubic arrangement. If all atoms in this structure are of the same species, it is a bcc lattice. The spheres occupy 68 percent of the volume. There are 23 metals with the bcc arrangement. The sum of these three numbers (32 + 26 + 23) exceeds the number of elements that form metals (63), since some elements are found in two or three of these structures.

      The fcc structure is also found for crystals of the rare gas (noble gas) solids neon (Ne), argon (Ar), krypton (Kr), and xenon (Xe). Their melting temperatures at atmospheric pressure are: Ne, 24.6 K; Ar, 83.8 K; Kr, 115.8 K; and Xe, 161.4 K.

Structures of nonmetallic elements
 The elements in the fourth row of the periodic table—carbon (carbon group element), silicon, germanium (Ge), and α-tin (α-Sn)—prefer covalent bonding (covalent bond). carbon has several possible crystal structures. Each atom in the covalent bond has four first-neighbours, which are at the corners of a tetrahedron. This arrangement is called the diamond lattice and is shown in Figure 3C—>. There are two atoms in a unit cell, which is fcc. Large crystals of diamond are valuable gemstones. The crystal has other interesting properties; it has the highest sound velocity of any solid and is the best conductor of heat. Besides diamond, the other common form of carbon is graphite, which is a layered material. Each carbon atom has three coplanar near neighbours, forming an arrangement called the honeycomb lattice. Three-dimensional graphite crystals are obtained by stacking similar layers.

      Another form of crystalline carbon is based on a molecule with 60 carbon atoms called buckminsterfullerene (fullerene) (C60). The molecular shape is spherical. Each carbon is bonded to three neighbours, as in graphite, and the spherical shape is achieved by a mixture of 12 rings with five sides and 20 rings with six sides. Similar structures were first visualized by the American architect R. Buckminster Fuller for geodesic domes. The C60 molecules, also called buckyballs, are quite strong and almost incompressible. Crystals are formed such that the balls are arranged in an fcc lattice with a one-nanometre spacing between the centres of adjacent balls. The similar C70 molecule has the shape of a rugby ball; C70 molecules also form an fcc crystal when stacked together. The solid fullerenes form molecular crystals, with weak binding—provided by van der Waals interactions—between the molecules.

      Many elements form diatomic gases: hydrogen (H), oxygen (O), nitrogen (N), fluorine (F), chlorine (Cl), bromine (Br), and iodine (I). When cooled to low temperature, they form solids of diatomic molecules. Nitrogen has the hcp structure, while oxygen has a more complex structure.

 The most interesting crystal structures are those of elements that are neither metallic, covalent, nor diatomic. Although boron (B) and sulfur (S) have several different crystal structures, each has one arrangement in which it is usually found. Twelve boron atoms form a molecule in the shape of an icosahedron (Figure 4—>). Crystals are formed by stacking the molecules. The β-rhombohedral structure of boron has seven of these icosahedral molecules in each unit cell, giving a total of 84 atoms. Molecules of sulfur are usually arranged in rings; the most common ring has eight atoms. The typical structure is α-sulfur, which has 16 molecules per unit cell, or 128 atoms. In the common crystals of selenium (Se) and tellurium (Te), the atoms are arranged in helical chains, which stack like cordwood. However, selenium also makes eight-atom rings, similar to sulfur, and forms crystals from them. Sulfur also makes helical chains, similar to selenium, and stacks them together into crystals.

Structures of binary crystals
      Binary crystals are found in many structures. Some pairs of elements form more than one structure. At room temperature, cadmium sulfide may crystallize either in the zinc blende or wurtzite structure. alumina also has two possible structures at room temperature, α-alumina (corundum) and β-alumina. Other binary crystals exhibit different structures at different temperatures. Among the most complex crystals are those of silicon dioxide (silica) (SiO2), which has seven different structures at various temperatures and pressures; the most common of these structures is quartz. Some pairs of elements form several different crystals in which the ions have different chemical valences. cadmium (Cd) and phosphorus (P) form the crystals Cd3P2, CdP2, CdP4, Cd7P10, and Cd6P7. Only in the first case are the ions assigned the expected chemical valences of Cd2+ and P3-.

   Among the binary crystals, the easiest structures to visualize are those with equal numbers of the two types of atoms. The structure of sodium chloride (salt) is based on a cube. To construct the lattice, the sodium and chlorine atoms are placed on alternate corners of a cube, and the structure is repeated (Figure 3B—>). The structure of the sodium atoms alone, or the chlorine atoms alone, is fcc and defines the unit cell. The sodium chloride structure thus is made up of two interpenetrating fcc lattices. The cesium chloride lattice (Figure 3A—>) is based on the bcc structure; every other atom is cesium or chlorine. In this case, the unit cell is a cube. The third important structure for AB (binary) lattices is zinc blende (sphalerite) (Figure 3D—>). It is based on the diamond structure, where every other atom is A or B. Many binary semiconductors have this structure, including those with one atom from the third (boron, aluminum, gallium [Ga], or indium [In]) and one from the fifth (nitrogen, phosphorus, arsenic [As], or antimony [Sb]) column of the periodic table (GaAs, InP, etc.). Most of the chalcogenides (O, S, Se, Te) of cadmium and zinc (CdTe, ZnSe, ZnTe, etc.) also have the zinc blende structure. The mineral zinc blende is ZnS; its unit cell is also fcc. The wurtzite structure is based on the hcp lattice, where every other atom is A or B. These four structures comprise most of the binary crystals with equal numbers of cations and anions.

      The fullerene molecule forms binary crystals MxC60 with alkali atoms, where M is potassium (K), rubidium (Rb), or cesium (Cs). The fullerene molecules retain their spherical shape, and the alkali atoms sit between them. The subscript x can take on several values. A compound with x = 6 (e.g., K6C60) is an insulator with the fullerenes in a bcc structure. The case x = 4 is an insulator with the body-centred tetragonal structure, while the case x = 3 is a metal with the fullerenes in an fcc structure. K3C60, Rb3C60, and Cs3C60 are superconductors at low temperatures.

Alloys (alloy)
      Alloys are solid mixtures of atoms with metallic properties. The definition includes both amorphous and crystalline solids. Although many pairs of elements will mix together as solids, many pairs will not. Almost all chemical entities can be mixed in liquid form. But cooling a liquid to form a solid often results in phase separation; a polycrystalline material is obtained in which each grain is purely one atom or the other. Extremely rapid cooling can produce an amorphous alloy. Some pairs of elements form alloys that are metallic crystals. They have useful properties that differ from those exhibited by the pure elements. For example, alloying makes a metal stronger; for this reason alloys of gold, rather than the pure metal, are used in jewelry.

      Atoms (atom) tend to form crystalline alloys when they are of similar size. The sizes of atoms are not easy to define, however, because atoms are not rigid objects with sharp boundaries. The outer part of an atom is composed of electrons in bound orbits; the average number of electrons decreases gradually with increasing distance from the nucleus. There is no point that can be assigned as the precise radius of the atom. Scientists have discovered, however, that each atom in a solid has a characteristic radius that determines its preferred separation from neighbouring atoms. For most types of atom this radius is constant, even in different solids. An empirical radius is assigned to each atom for bonding considerations, which leads to the concept of atomic size. Atoms readily make crystalline alloys when the radii of the two types of atoms agree to within roughly 15 percent.

      Two kinds of ordering are found in crystalline alloys. Most alloys at low temperature are binary crystals with perfect ordering. An example is the alloy of copper and zinc. Copper is fcc, whereas zinc is hcp. A 50-percent-zinc–50-percent-copper alloy has a different structure—β-brass. At low temperatures it has the cesium chloride structure: a bcc lattice with alternating atoms of copper and zinc and a cubic unit cell. If the temperature is raised above 470° C, however, a phase transition to another crystalline state occurs. The ordering at high temperature is also bcc, but now each site has equal probability of having a copper or zinc atom. The two types of atoms randomly occupy each site, but there is still long-range order. At all temperatures, thousands of atoms away from a site, the location of the atom site can be predicted with certainty. At temperatures below 470° C one also knows whether that site will be occupied by a copper or zinc atom, while above 470° C there is an equal likelihood of finding either atom. The high-temperature phase is crystalline but disordered. The disorder phase is obtained through a partial melting, not into a liquid state but into a less ordered one. This behaviour is typical of metal alloys. Other common alloys are steel, an alloy of iron and carbon; stainless steel, an alloy of iron, nickel (Ni), and chromium (Cr); pewter and solder, alloys of tin and lead (Pb); and britannia metal, an alloy of tin, antimony, and copper.

Crystal defects (crystal defect)
      A crystal is never perfect; a variety of imperfections can mar the ordering. A defect is a small imperfection affecting a few atoms. The simplest type of defect is a missing atom and is called a vacancy. Since all atoms occupy space, extra atoms cannot be located at the lattice sites of other atoms, but they can be found between them; such atoms are called interstitials. Thermal vibrations may cause an atom to leave its original crystal site and move into a nearby interstitial site, creating a vacancy-interstitial pair. Vacancies and interstitials are the types of defects found in a pure crystal. In another defect, called an impurity, an atom is present that is different from the host crystal atoms. Impurities may either occupy interstitial spaces or substitute for a host atom in its lattice site.

      There is no sharp distinction between an alloy and a crystal with many impurities. An alloy results when a sufficient number of impurities are added that are soluble in the host metal. However, most elements are not soluble in most crystals. Crystals generally can tolerate a few impurities per million host atoms. If too many impurities of the insoluble variety are added, they coalesce to form their own small crystallite. These inclusions are called precipitates and constitute a large defect.

       germanium is a common impurity in silicon. It prefers the same tetrahedral bonding as silicon and readily substitutes for silicon atoms. Similarly, silicon is a common impurity in germanium. No large crystal can be made without impurities; the purest large crystal ever grown was made of germanium. It had about 1010 impurities in each cubic centimetre of material, which is less than one impurity for each trillion atoms.

      Impurities often make crystals more useful. In the absence of impurities, α- alumina is colourless. Iron and titanium impurities impart to it a blue colour, and the resulting gem-quality mineral is known as sapphire. Chromium impurities are responsible for the red colour characteristic of rubies (ruby), the other gem of α-alumina. Pure semiconductors rarely conduct electricity well at room temperatures. Their ability to conduct electricity is caused by impurities. Such impurities are deliberately added to silicon in the manufacture of integrated circuits. In fluorescent lamps (fluorescent lamp) the visible light is emitted by impurities in the phosphors (luminescent materials).

 Other imperfections in crystals involve many atoms. twinning is a special type of grain boundary defect, in which a crystal is joined to its mirror image. Another kind of imperfection is a dislocation, which is a line defect that may run the length of the crystal. One of the many types of dislocations is due to an extra plane of atoms that is inserted somewhere in the crystal structure. Another type, called an edge dislocation, is shown in Figure 5—>. This line defect occurs when there is a missing row of atoms. In the figure the crystal arrangement is perfect on the top and on the bottom. The defect is the row of atoms missing from region b. This mistake runs in a line that is perpendicular to the page and places a strain on region a.

      Dislocations are formed when a crystal is grown, and great care must be taken to produce a crystal free of them. Dislocations are stable and will exist for years. They relieve mechanical stress. If one presses on a crystal, it will accommodate the induced stress by growing dislocations at the surface, which gradually move inward. Dislocations make a crystal mechanically harder. When a metal bar is cold-worked by rolling or hammering, dislocations and grain boundaries are introduced; this causes the hardening.

Determination of crystal structures
      Crystal structures are determined by scattering experiments using a portion of the crystal as the target. A beam of particles is sent toward the target, and upon impact some of the particles scatter from the crystal and ricochet in various directions. A measurement of the scattered particles provides raw data, which is then computer-processed to give a picture of the atomic arrangements. The positions are then inferred from the computer-analyzed data.

 Max von Laue (Laue, Max von) first suggested in 1912 that this measurement could be done using X rays (X-ray), which are electromagnetic radiation of very high frequency. High frequencies are needed because these waves have a short wavelength. Von Laue realized that atoms have a spacing of only a few angstroms (1 angstrom [Å] is 10−10 metre, or 3.94 × 10−9 inch). In order to measure atomic arrangements, the particles scattering from the target must also have a wavelength of a few angstroms. X rays are required when the beam consists of electromagnetic radiation. The X rays only scatter in certain directions, and there are many X rays associated with each direction. The scattered particles appear in spots corresponding to locations where the scattering from each identical atom produces an outgoing wave that has all the wavelengths (wavelength) in phase. Figure 6—> shows incoming waves in phase. The scattering from atom A2 has a longer path than that from atom A1. If this additional path has a length (AB + BC) that is an exact multiple of the wavelength, then the two outgoing waves are in phase and reinforce each other. If the scattering angle is changed slightly, the waves no longer add coherently and begin to cancel one another. Combining the scattered radiation from all the atoms in the crystal causes all the outgoing waves to add coherently in certain directions and produce a strong signal in the scattered wave. If the extra path length (AB + BC) is five wavelengths, for example, the spot appears in one place. If it is six wavelengths, the spot is elsewhere. Thus, the different spots correspond to the different multiples of the wavelength of the X ray. The measurement produces two types of information: the directions of the spots and their intensity. This information is insufficient to deduce the exact crystal structure, however, as there is no algorithm by which the computer can go directly from the data to the structure. The crystallographer must propose various structures and compute how they would scatter the X rays. The theoretical results are compared with the measured one, and the theoretical arrangement is chosen that best fits the data. Although this procedure is fast when there are only a few atoms in a unit cell, it may take months or years for complex structures. Some protein molecules, for instance, have hundreds of atoms. Crystals of the proteins are grown, and X rays are used to measure the structure. The goal is to determine how the atoms are arranged in the protein, rather than how the proteins are arranged in the crystal.

      Beams of neutrons (neutron) may also be used to measure crystal structure. The beam of neutrons is obtained by drilling a hole in the side of a nuclear reactor. The energetic neutrons created in nuclear fission escape through the hole. The motion of elementary particles is governed by quantum, or wave, mechanics. Each neutron has a wavelength that depends on its momentum. The scattering directions are determined by the wavelength, as is the case with X rays. The wavelengths for neutrons from a reactor are suitable for measuring crystal structures.

      X rays and neutrons provide the basis for two competing technologies in crystallography. Although they are similar in principle, the two methods have some differences. X rays scatter from the electrons (electron) in the atoms so that more electrons result in more scattering. X rays easily detect atoms of high atomic number, which have many electrons, but cannot readily locate atoms with few electrons. In hydrogen-bonded crystals, X rays do not detect the protons at all. Neutrons, on the other hand, scatter from the atomic nucleus. They scatter readily from protons and are excellent for determining the structure of hydrogen-bonded solids. One drawback to this method is that some nuclei absorb neutrons completely, and there is little scattering from these targets.

      Beams of electrons (electron beam) can also be used to measure crystal structure, because energetic electrons have a wavelength that is suitable for such measurements. The problem with electrons is that they scatter strongly from atoms. Proper interpretation of the experimental results requires that an electron scatter only from one atom and leave the crystal without scattering again. Low-energy electrons scatter many times, and the interpretation must reflect this. Low-energy electron diffraction (LEED) is a technique in which a beam of electrons is directed toward the surface. The scattered electrons that reflect backward from the surface are measured. They scatter many times before leaving backward but mainly leave in a few directions that appear as “spots” in the measurements. An analysis of the varied spots gives information on the crystalline arrangement. Because the electrons are scattered strongly by the atoms in the first few layers of the surface, the measurement gives only the arrangements of atoms in these layers. It is assumed that the same structure is repeated throughout the crystal. Another scattering experiment involves electrons of extremely high energy. The scattering rate decreases as the energy of the electron increases, so that very energetic electrons usually scatter only once. Various electron microscopes (electron microscope) are constructed on this principle.

Types of bonds (chemical bonding)
      The properties of a solid can usually be predicted from the valence and bonding preferences of its constituent atoms. Four main bonding types are discussed here: ionic, covalent, metallic, and molecular. Hydrogen-bonded solids, such as ice, make up another category that is important in a few crystals. There are many examples of solids that have a single bonding type, while other solids have a mixture of types, such as covalent and metallic or covalent and ionic.

Ionic bonds (ionic bond)
      Sodium chloride (salt) exhibits ionic bonding. The sodium atom (sodium) has a single electron in its outermost shell, while chlorine needs one electron to fill its outer shell. Sodium donates one electron to chlorine, forming a sodium ion (Na+) and a chlorine ion (Cl). Each ion thus attains a closed outer shell of electrons and takes on a spherical shape. In addition to having filled shells and a spherical shape, the ions of an ionic solid have integer valence. An ion with positive valence is called a cation. In an ionic solid the cations are surrounded by ions with negative valence, called anions (anion). Similarly, each anion is surrounded by cations. Since opposite charges attract, the preferred bonding occurs when each ion has as many neighbours as possible, consistent with the ion radii. Six or eight nearest neighbours are typical; the number depends on the size of the ions and not on the bond angles. The alkali halide crystals are binaries of the AH type, where A is an alkali ion (lithium [Li], sodium, potassium, rubidium, or cesium) and H is a halide ion (fluorine, chlorine, bromine, or iodine). The crystals have ionic bonding, and each ion has six or eight neighbours. Metal ions in the alkaline earth series (alkaline-earth metal) (magnesium [Mg], calcium [Ca], barium [Ba], and strontium [Sr]) have two electrons in their outer shells and form divalent cations in ionic crystals. The chalcogenides (oxygen group element) (oxygen, sulfur, selenium, and tellurium) need two electrons to fill their outer p-shell. (Electron shells are divided into subshells, designated as s, p, d, f, g, and so forth. Each subshell is divided further into orbitals.) Two electrons are transferred from the cations to the anions, leaving each with a closed shell. The alkaline earth chalcogenides form ionic binary crystals such as barium oxide (BaO), calcium sulfide (CaS), barium selenide (BaSe), or strontium oxide (SrO). They have the same structure as sodium chloride, with each atom having six neighbours. oxygen can be combined with various cations to form a large number of ionically bonded solids.

Covalent bonds (covalent bond)
       silicon, carbon, germanium, and a few other elements form covalently bonded solids. In these elements there are four electrons in the outer sp-shell, which is half filled. (The sp-shell is a hybrid formed from one s and one p subshell.) In the covalent bond an atom shares one valence (valence electron) (outer-shell) electron with each of its four nearest neighbour atoms. The bonds are highly directional and prefer a tetrahedral arrangement. A covalent bond is formed by two electrons—one from each atom—located in orbitals between the ions. Insulators, in contrast, have all their electrons within shells inside the atoms.

      The perpetual spin of an electron is an important aspect of the covalent bond. From a vantage point above the spinning particle, counterclockwise rotation is designated spin-up, while clockwise rotation is spin-down. A fundamental law of quantum physics is the Pauli exclusion principle, which states that no two electrons can occupy the same point in space at the same time with the same direction of spin. In a covalent bond two electrons occupy the same small volume of space (i.e., the same orbital) at all times, so they must have opposite spin: one up and one down. The exclusion principle is then satisfied, and the resulting bond is strong.

      In graphite the carbon atoms are arranged in parallel sheets, and each atom has only three near neighbours. The covalent bonds between adjacent carbons within each layer are quite strong and are called σ (sigma bond) bonds. The fourth valence electron in carbon has its orbital perpendicular to the plane. This orbital bonds weakly with the similar orbitals on all three neighbours, forming π (pi bond) bonds. The four bonds for each carbon atom in the graphite structure are not arranged in a tetrahedron; three are in a plane. The planar arrangement results in strong bonding, although not as strong as the bonding in the diamond configuration. The bonding between layers is quite weak and arises from the van der Waals interaction (van der Waals forces); there is much slippage parallel to the layers. Diamond and graphite form an interesting contrast: diamond is the hardest material in nature and is used as an abrasive, while graphite is used as a lubricant.

      Besides the elemental semiconductors, such as silicon and germanium, some binary crystals are covalently bonded. gallium has three electrons in the outer shell, while arsenic lacks three. Gallium arsenide (GaAs) could be formed as an insulator by transferring three electrons from gallium to arsenic; however, this does not occur. Instead, the bonding is more covalent, and gallium arsenide is a covalent semiconductor. The outer shells of the gallium atoms contribute three electrons, and those of the arsenic atoms contribute five, providing the eight electrons needed for four covalent bonds. The centres of the bonds are not at the midpoint between the ions but are shifted slightly toward the arsenic. Such bonding is typical of the III–V semiconductors—i.e., those consisting of one element from the third column of the periodic table and one from the fifth column. Elements from the third column (boron, aluminum, gallium, and indium) contribute three electrons, while the fifth-column elements (nitrogen, phosphorus, arsenic, and antimony) contribute five electrons. All III–V semiconductors are covalently bonded and typically have the zinc blende structure with four neighbours per atom. Most common semiconductors (semiconductor) favour this arrangement.

      The factor that determines whether a binary crystal will act as an insulator or a semiconductor is the valence of its constituent atoms. Ions (ion) that donate or accept one or two valence electrons form insulators. Those that have three to five valence electrons tend to have covalent bonds and form semiconductors. There are exceptions to these rules, however, as is the case with the IV–VI semiconductors such as lead sulfide. Heavier elements from the fourth column of the periodic table (germanium, tin, and lead) combine with the chalcogenides from the sixth row to form good binary semiconductors such as germanium telluride (GeTe) or tin sulfide (SnS). They have the sodium chloride structure, where each atom has six neighbours. Although not tetrahedrally bonded, they are good semiconductors.

      Filled atomic shells with d-orbitals have an important role in covalent bonding. Electrons in atomic orbits have angular momentum (L), which is quantized in integer (n) multiples of Planck's constant h: L = nh. Electron orbitals with n = 0 are called s-states, with n = 1 are p-states, and with n = 2 are d-states. silver and copper ions have one valence electron outside their closed shells. The outermost filled shell is a d-state and affects the bonding. Eight binary crystals are formed from the copper and silver halides. Three (AgF, AgCl, AgBr) have the sodium chloride structure with six neighbours. The other five (AgI, CuF, CuCl, CuBr, CuI) have the zinc blende structure with four neighbours. The bonding in this group of solids is on the borderline between covalent and ionic, since the crystals prefer both types of bonds. The alkali metal halides exhibit somewhat different behaviour. The alkali metals are also monovalent cations, but their halides are strictly ionic. The difference in bonding between the alkali metals on the one hand and silver and copper on the other hand is that silver and copper have filled d-shells while the alkalis have filled p-shells. Since the d-shells are filled, they do not covalently bond. This group of electrons is, however, highly polarizable, which influences the bonding of the valence electrons. Similar behaviour is found for zinc and cadmium, which have two valence electrons outside a filled d-shell. They form binary crystals with the chalcogenides, which have tetrahedral bonding. In this case the covalent bonding seems to be preferred over the ionic bond. In contrast, the alkaline earth chalcogenides, which are also divalent, have outer p-shells and are ionic. The zinc and cadmium chalcogenides are covalent, as the outer d-shell electrons of the two cations favour covalent bonding.

Metallic bonds (metallic bond)
      Metallic bonds fall into two categories. The first is the case in which the valence electrons (valence electron) are from the p-shells of the metal ions; this bonding is quite weak. In the second category the valence electrons are from partially filled d-shells, and this bonding is quite strong. The d-bonds dominate when both types of bonding are present.

      The simple metals (metal) are bonded with sp-electrons. The electrons of these metal atoms are in filled atomic shells except for a few electrons that are in unfilled sp-shells. The electrons from the unfilled shells are detached from the metal ion and are free to wander throughout the crystal. They are called conduction electrons, since they are responsible for the electrical conductivity of metals. Although the conduction electrons may roam anywhere in the crystal, they are distributed uniformly throughout the entire solid. Any large imbalance of charge is prevented by the strong electrical attraction between the negative electrons and the positive ions, plus the strong repulsion between electrons. The phrase electron correlation describes the correlated movements of the electrons; the motion of each electron depends on the positions of neighbouring electrons. Electrons have strong short-range order with one another. Correlation ensures that each unit cell in the crystal has, on the average, the number of electrons needed to cancel the positive charge of the cation so that the unit cell is electrically neutral.

      Cohesive energy is the energy gained by arranging the atoms in a crystalline state, as compared with the gas state. Insulators and semiconductors have large cohesive energies; these solids are bound together strongly and have good mechanical strength. Metals with electrons in sp-bonds have very small cohesive energies. This type of metallic bond is weak; the crystals are barely held together. Single crystals of simple metals such as sodium are mechanically weak. At room temperature the crystals have the mechanical consistency of warm butter. Special care must be used in handling these crystals, because they are easily distorted. Metals such as magnesium or aluminum must be alloyed or polycrystalline to have any mechanical strength. Although the simple metals are found in a variety of structures, most are in one of the three closest-packed structures: fcc, bcc, and hcp. Theoretical calculations show that the cohesive energy of a given metal is almost the same in each of the different crystal arrangements; therefore, crystal arrangements are unimportant in metals bound with electrons from sp-shells.

      A different type of metallic bonding is found in transition metals (transition element), which are metals whose atoms are characterized by unfilled d-shells. The d-orbitals are more tightly bound to an ion than the sp-orbitals. Electrons in d-shells do not wander away from the ion. The d-orbitals form a covalent bond with the d-orbitals on the neighbouring atoms. The bonding of d-orbitals does not occur in a tetrahedral arrangement but has a different directional preference. In metals the bonds from d-orbitals are not completely filled with electrons. This situation is different from the tetrahedral bonds in semiconductors, which are filled with eight electrons. In transition metals the covalent bonds formed with the d-electrons are much stronger than the weak bonds made with the sp-electrons of simple metals. The cohesive energy is much larger in transition metals. Titanium, iron, and tungsten, for example, have exceptional mechanical strength. Crystal arrangements are important in the behaviour of the transition metals and occur in the close-packed fcc, bcc, or hcp arrangements.

Molecular binding
      The Dutch physicist Johannes D. van der Waals (Waals, Johannes Diederik van der) first proposed the force that binds molecular solids. Any two atoms or molecules have a force of attraction (F) that varies according to the inverse seventh power of the distance R between the centres of the atoms or molecules: F = −C/R7, where C is a constant. The force, known as the van der Waals force (van der Waals forces), declines rapidly with the distance R and is quite weak. If the atoms or molecules have a net charge, there is a strong force whose strength varies according to Coulomb's law (Coulomb force) as the inverse second power of the separation distance: F = −C′/R2, where C′ is a constant. This force provides the binding in ionic crystals and some of the binding in metals. Coulomb's law does not apply to atoms or molecules without a net charge. Molecules with a dipole moment, such as water, have a strong attractive force owing to the interactions between the dipoles. For atoms and molecules with neither net charges nor dipole moments, the van der Waals force provides the crystal binding. The force of gravity also acts between neutral atoms and molecules, but it is far too weak to bind molecules into crystals.

      The van der Waals force is caused by quantum fluctuations. Two neighbouring atoms that are each fluctuating can lower their joint energy by correlating their fluctuations. The van der Waals force arises from correlations in their dipole fluctuations. Electrons bound in atomic orbits are in constant motion around the nucleus, and the distribution of charges in the atom changes constantly as the electrons move, owing to quantum fluctuations. One fluctuation might produce a momentary electric dipole moment (i.e., a separation of charges) on an atom if a majority of its electrons are on one side of the nucleus. The dipole moment creates an electric field on a neighbouring atom; this field will induce a dipole moment on the second atom. The two dipoles attract one another via the van der Waals interaction. Since the force depends on the inverse seventh power of the distance, it declines rapidly with increasing distance. Atoms have a typical radius of one to three angstroms. The van der Waals force binds atoms and molecules within a few angstroms of each other; beyond that distance the force is negligible. Although weak, the van der Waals force is always present and is important in cases where the other forces are absent.

       hydrogen is rarely found as a single atom. Instead it forms diatomic molecules (H2), which are gaseous at room temperature. At lower temperatures the hydrogen becomes a liquid and at about 20 K turns into a solid. The molecule retains its identity in the liquid and solid states. The solid exists as a molecular crystal of covalently bound H2 molecules. The molecules attract one another by van der Waals forces, which provide the crystal binding. helium, the second element in the periodic table, has two electrons, which constitute a filled atomic shell. In its liquid and solid states, the helium atoms are bound together by van der Waals forces. In fact, all the rare gases (helium, neon, argon, krypton, and xenon) are molecular crystals with the binding provided by van der Waals forces.

      Many organic molecules form crystals where the molecules are bound by van der Waals forces. In methane (CH4), a central carbon makes a covalent bond with each hydrogen atom, forming a tetrahedron. In crystalline methane the molecules are arranged in the fcc structure. benzene (C6H6) has the carbon atoms in a hexagonal ring; each carbon has three coplanar σ bonds, as in graphite, where two bonds are with neighbouring carbon atoms and the third bond is with a hydrogen atom. Crystalline benzene has four molecules per unit cell in a complex arrangement. Fullerene and the rare gas atoms are spherical, and the crystalline arrangement corresponds to the closest packing of spheres. Most organic molecules, however, are not spherical and display irregular shapes. For odd-shaped molecules, the van der Waals interaction depends on the rotational orientation of the two molecules. In order to maximize the force, the molecules in the crystal have unusual arrangements, as in the case of benzene.

Hydrogen bonding
      Hydrogen bonding is important in a few crystals, notably in ice. With its lone electron, a hydrogen atom usually forms a single covalent bond with an electronegative atom (electronegativity). In the hydrogen bond the atom is ionized to a proton. The proton sits between two anions and joins them. Hydrogen bonding occurs with only the most electronegative ions: nitrogen, oxygen, and fluorine. In water the hydrogen links pairs of oxygen ions. Water is found in many different crystal structures, but they all have the feature that the hydrogen atoms sit between pairs of oxygen. Another hydrogen-bonded solid is hydrogen fluoride (HF), in which the hydrogen atom (proton) links pairs of fluorines.

Crystal growth
      The earliest crystal grower was nature. Many excellent crystals of minerals formed in the geologic past are found in mines and caves throughout the world. Most precious and semiprecious stones are well-formed crystals. Early efforts to produce synthetic crystals were concentrated on making gems. Synthetic ruby was grown by the French scientist Marc Antoine Augustin Gaudin in 1873. Since about 1950 scientists have learned to grow in the laboratory crystals of quality equal or superior to those found in nature. New techniques for growth are continually being developed, and crystals with three or more atoms per unit cell are continually being discovered.

Vapour growth
      Crystals can be grown from a vapour when the molecules of the gas attach themselves to a surface and move into the crystal arrangement. Several important conditions must be met for this to occur. At constant temperature and equilibrium conditions, the average number of molecules in the gas and solid states is constant; molecules leave the gas and attach to the surface at the same rate that they leave the surface to become gas molecules. For crystals to grow, the gas-solid chemical system must be in a nonequilibrium state such that there are too many gaseous molecules for the conditions of pressure and temperature. This state is called supersaturation. Molecules are more prone to leave the gas than to rejoin it, so they become deposited on the surface of the container. Supersaturation can be induced by maintaining the crystal at a lower temperature than the gas. A critical stage in the growth of a crystal is seeding, in which a small piece of crystal of the proper structure and orientation, called a seed, is introduced into the container. The gas molecules find the seed a more favourable surface than the walls and preferentially deposit there. Once the molecule is on the surface of the seed, it wanders around this surface to find the preferred site for attachment. Growth proceeds one molecule at a time and one layer at a time. The process is slow; it takes days to grow a small crystal. Crystals are grown at temperatures well below the melting point to reduce the density of defects. The advantage of vapour growth is that very pure crystals can be grown by this method, while the disadvantage is that it is slow.

      Most clouds in the atmosphere are ice crystals that form by vapour growth from water molecules. Most raindrops (rain) are crystals as they begin descending but thaw during their fall to Earth. Seeding (cloud seeding) for rain—accomplished by dropping silver iodide crystals from airplanes—is known to induce precipitation. In the laboratory, vapour growth is usually accomplished by flowing a supersaturated gas over a seed crystal. Quite often a chemical reaction at the surface is needed to deposit the atoms. Crystals of silicon can be grown by flowing chlorosilane (SiCl4) and hydrogen (H2) over a seed crystal of silicon. Hydrogen acts as the buffer gas by controlling the temperature and rate of flow. The molecules dissociate on the surface in a chemical reaction that forms hydrogen chloride (HCl) molecules. Hydrogen chloride molecules leave the surface, while silicon atoms remain to grow into a crystal. Binary crystals such as gallium arsenide (GaAs) are grown by a similar method. One process employs gallium chloride (GaCl) as the gallium carrier. Arsenic is provided by molecules such as arsenous chloride (AsCl3), arsine (AsH3), or As4 (yellow arsenic). These molecules, with hydrogen as the buffer gas, grow crystals of gallium arsenide while forming gas molecules such as gallium trichloride (GaCl3) and hydrogen chloride. Trimethylgallium, (CH3)3Ga, is another molecule that can be used to deliver gallium to the surface.

Growth from solution
      Large single crystals may be grown from solution. In this technique the seed crystal is immersed in a solvent that contains typically about 10–30 percent of the desired solute. The choice of solvent usually depends on the solubility of the solute. The temperature and pH (a measure of acidity or basicity) of the solution must be well controlled. The method is faster than vapour growth, because there is a higher concentration of molecules at the surface in a liquid as compared to a gas, but it is still relatively slow.

Growth from the melt
  This method is the most basic. A gas is cooled until it becomes a liquid, which is then cooled further until it becomes a solid. Polycrystalline (polycrystal) solids are typically produced by this method unless special techniques are employed. In any case, the temperature must be controlled carefully. Large crystals can be grown rapidly from the liquid elements using a popular method invented in 1918 by the Polish scientist Jan Czochralski and called crystal pulling. One attaches a seed crystal to the bottom of a vertical arm such that the seed is barely in contact with the material at the surface of the melt. A modern Czochralski apparatus is shown in Figure 7A—>. The arm is raised slowly, and a crystal grows underneath at the interface between the crystal and the melt. Usually the crystal is rotated slowly, so that inhomogeneities in the liquid are not replicated in the crystal. Large-diameter crystals of silicon are grown in this way for use as computer chips (integrated circuit). Based on measurements of the weight of the crystal during the pulling process, computer-controlled apparatuses can vary the pulling rate to produce any desired diameter. Crystal pulling is the least expensive way to grow large amounts of pure crystal. A photograph of a single crystal of stainless steel grown by the Czochralski method is shown in Figure 7B—>. The original seed is on the right tip. Binary crystals can also be pulled; for example, synthetic sapphire crystals can be pulled from molten alumina. Special care is required to grow binary and other multicomponent crystals; the temperature must be precisely controlled because such crystals may be grown only at a single, extremely high temperature. The melt has a tendency to be inhomogeneous, since the two liquids may try to separate by gravity.

      The Bridgman method (named after the American scientist Percy Williams Bridgman) is also widely used for growing large single crystals. The molten material is put into a crucible, often of silica, which has a cylindrical shape with a conical lower end. Heaters maintain the molten state. As the crucible is slowly lowered into a cooler region, a crystal starts growing in the conical tip. The crucible is lowered at a rate that matches the growth of the crystal, so that the interface between crystal and melt is always at the same temperature. The rate of moving the crucible depends on the temperature and the material. When done successfully, the entire molten material in the crucible grows into a single large crystal. One disadvantage of the method is that excess impurities are pushed out of the crystal during growth. A layer of impurities grows at the interface between melt and solid as this surface moves up the melt, and the impurities become concentrated in the higher part of the crystal.

       epitaxy is the technique of growing a crystal, layer by layer, on the atomically flat surface of another crystal. In homoepitaxy a crystal is grown on a substrate of the same material. Silicon layers of different impurity content, for example, are grown on silicon substrates in the manufacture of computer chips. Heteroepitaxy, on the other hand, is the growth of one crystal on the substrate of another. Silicon substrates are often used since they are readily available in atomically smooth form. Many different semiconductor crystals can be grown on silicon, such as gallium arsenide, germanium, cadmium telluride (CdTe), and lead telluride (PbTe). Any flat substrate can be used for epitaxy, however, and insulators such as rock salt (NaCl) and magnesium oxide (MgO) are also used.

      Molecular-beam epitaxy, commonly abbreviated as MBE, is a form of vapour growth. The field began when the American scientist John Read Arthur reported in 1968 that gallium arsenide could be grown by sending a beam of gallium atoms and arsenic molecules toward the flat surface of a crystal of the molecule. The amount of gas molecules can be controlled to grow just one layer, or just two, or any desired amount. This method is slow, since molecular beams have low densities of atoms. Chemical vapour deposition (CVD) is another form of epitaxy that makes use of the vapour growth technique. Also known as vapour-phase epitaxy (VPE), it is much faster than MBE since the atoms are delivered in a flowing gas rather than in a molecular beam. Synthetic diamonds (synthetic diamond) are grown by CVD. Rapid growth occurs when methane (CH4) is mixed with atomic hydrogen gas, which serves as a catalyst. Methane dissociates on a heated surface of diamond. The carbon remains on the surface, and the hydrogen leaves as a molecule. Growth rates are several micrometres (1 micrometre is equal to 0.00004 inch) per hour. At that rate, a stone 1 centimetre (0.4 inch) thick is grown in 18 weeks. CVD diamonds are of poor quality as gemstones but are important electronic materials. Because hydrogen is found in nature as a molecule rather than as a single atom, making atomic hydrogen gas is the major expense in growing CVD diamonds. Liquid-phase epitaxy (LPE) uses the solution method to grow crystals on a substrate. The substrate is placed in a solution with a saturated concentration of solute. This technique is used to grow many crystals employed in modern electronics and optoelectronic devices, such as gallium arsenide, gallium aluminum arsenide, and gallium phosphide.

 An important concern in successful epitaxy is matching lattice distances. If the spacing between atoms in the substrate is close to that of the top crystal, then that crystal will grow well; a small difference in lattice distance can be accommodated as the top crystal grows. When the lattice distances are different, however, the top crystal becomes deformed, since structural defects such as dislocations appear (see Figure 5—>). Although few crystals share the same lattice distance, a number of examples are known. Aluminum arsenide and gallium arsenide have the same crystal structure and the same lattice parameters to within 0.1 percent; they grow excellent crystals on one another. Such materials, known as superlattices, have a repeated structure of n layers of GaAs, m layers of AlAs, n layers of GaAs, m layers of AlAs, and so forth. Superlattices represent artificially created structures that are thermodynamically stable; they have many applications in the modern electronics industry. Another lattice-matched epitaxial system is mercury telluride (HgTe) and cadmium telluride (CdTe). These two semiconductors form a continuous semiconductor alloy CdxHg1 − xTe, where x is any number between 0 and 1. This alloy is used as a detector of infrared radiation and is incorporated in particular in night-vision goggles.

Dendritic growth
      At slow rates of crystal growth, the interface between melt and solid remains planar, and growth occurs uniformly across the surface. At faster rates of crystal growth, instabilities are more likely to occur; this leads to dendritic growth. Solidification releases excess energy in the form of heat at the interface between solid and melt. At slow growth rates, the heat leaves the surface by diffusion. Rapid growth creates more heat, which is dissipated by convection (liquid flow) when diffusion is too slow. Convection breaks the planar symmetry so that crystal growth develops along columns, or “fingers,” rather than along planes. Each crystal has certain directions in which growth is fastest, and dendrites grow in these directions. As the columns grow larger, their surfaces become flatter and more unstable. This feather or tree structure is characteristic of dendritic growth. Snowflakes are an example of crystals that result from dendritic growth.

Electric properties

      The German physicist Georg Simon Ohm discovered the basic law of electric (electricity) conduction, which is now called Ohm's law. His law relates the voltage (V, measured in volts), the current (I, in amperes), and the resistance (R, in ohms) according to the formula V = RI. A current I through a solid induces a voltage V; the resistance R is the constant of proportionality. The value of R is an important factor in the design of electrical circuits. It is determined by the shape of the resistor: a long narrow object has more resistance than a short wide one of the same material. For solids, the important parameter is the resistivity ρ, which is given in units of ohm-metres. It is the resistivity per volume unit and is independent of shape. The relationship between R and ρ is R = ρL/A, where A is the area of the resistor and L is the length. These dimensions are measured in the direction of the current: L is the length of the current path, and A is the cross-sectional area. The resistance of a copper bar depends on its shape, but at a given temperature every piece of pure copper has the same resistivity. Thus the resistivity is a fundamental parameter of a material and is investigated by scientists. Resistivities of solids span a wide range of values. Certain metals have zero resistivity at low temperatures; they are called superconductors (superconductivity). At the other extreme, very good insulators such as sulfur and polystyrene have resistivities larger than one quadrillion ohm-metres. At room temperature, the metal with the lowest value of resistivity is silver, with ρ = 1.6 × 10−8 ohm-metre; the second best conductor is copper, with ρ = 1.7 × 10−8 ohm-metre. Copper, rather than silver, is used in household wires because of the high cost of silver.

Conduction through ion hopping
      Electrical conductivity σ is the inverse of resistivity and is measured in units of ohm-metre−1. Electrical current (electric current) is produced by the motion of charges. In crystals, electrical current is due to the motion of both ions (ion) and electrons. Ions move by hopping occasionally from site to site; all solids can conduct electricity in this manner. When the voltage is zero, there is no net current because the ions hop randomly in all directions. The imposition of a small voltage causes the ions to slightly favour one direction of motion, which leads to a net flow of charge in that direction; this constitutes an electrical current. The electricity conducted by this process is quite small and is usually negligible compared with that carried by the electrons. When an ion hops, it must migrate to a vacant site, which could be either an interstitial or a vacancy. Ionic conductivity can occur because hopping ions cause vacancies to move through the solid. An ion hops to the vacant site, thereby filling the vacancy, while creating a new one at the ion's former site. Repeating this process causes the position of the vacancy to migrate through the crystal. The motion of the vacancy arises from the motion of ions, which carry charge and contribute to electrical conductivity.

      Ion hops are induced by thermal fluctuations. Most of the ions move within their lattice site, vibrating around this point. temperature is defined as the average energy of this vibrational (vibration) motion; the more the ions move, the higher the temperature. An individual ion at times moves slowly and at times vibrates quite rapidly but usually has an energy near the average value. Each ion shares its vibrational energy with its neighbouring ions. An ion typically has some neighbours with small vibrations and others with large ones. The average energy shared with the neighbours is close to the average energy of all the atoms. As a random process, however, it occasionally happens that all neighbours of an ion may have large vibrations, in which case the ion will acquire an unusually high energy. This energy may be high enough to cause it to leave its site and hop to a neighbouring site. A thermal fluctuation is the rare process in which the energy at a local site may be much higher or lower than the average energy in the crystal. Probability theory shows that the higher the temperature, the more frequent are these thermal fluctuations. Ions therefore hop more often at high temperature.

      A few solids conduct electricity better by ion motion than by electron motion. These unusual materials are technologically important in making batteries (battery). All batteries have two electrodes separated by an electrolyte, which is a material that conducts ions better than electrons. An example of a crystal electrolyte is β-alumina, which readily conducts monovalent cations such as silver (Ag+) and sodium (Na+). Among all ions, silver has the largest value of ionic conductivity in many different electronic insulators. The copper ion (Cu+) forms the same type of chemical bonds as does the silver ion, but the copper ion, because of its smaller radius, does not migrate as well within an electrolyte. Silver ions fit perfectly into the interstitial sites of the crystal lattices of several electrolytes, while the smaller copper ions permit the neighbouring ions to collapse around them, inhibiting further hopping. There are a few good conductors of the inexpensive copper ion that can be used as solid electrolytes in batteries. Silver is too costly and heavy to use in large-volume batteries such as those found in automobiles, but it is used in the smaller batteries that power devices such as hearing aids.

Conduction electrons
      Electrons (electron) carry the basic unit of charge e, equal to 1.6022 × 10−19 coulomb. They have a small mass and move rapidly. Most electrons in solids are bound to the atoms in local orbits, but a small fraction of the electrons are available to move easily through the entire crystal. These so-called conduction electrons carry the electrical current. Solids with many conduction electrons are metals, while those with a few are semimetals or semiconductors. In insulators (insulator), nearly all the electrons are bound, and very few electrons are capable of carrying current. A typical metal has one or more conduction electrons in each atomic unit cell, a semiconductor may have only one conduction electron for each thousand unit cells, and an insulator may have one conduction electron per one million or one trillion unit cells.

      The bonding properties of the individual atoms of a solid determine the behaviour of the bulk solid. The electrical properties of a solid can usually be predicted from the valence and bonding preferences of its atoms. In the argon atom, for example, all atomic shells are filled with electrons. The electrons of solid argon remain in the atomic shells; none are conduction electrons, and the electrical resistivity is therefore high. Solid argon, like all the rare gas solids, is a good insulator. A few conduction electrons are contributed by impurities, and so the conductivity, though small, is not zero. These conduction electrons move quite readily through the solid. The term mobility is used to describe how well a conduction electron moves through the solid in response to a voltage. Conductivity is the product of mobility, the electrical charge e, and the number N of conduction electrons per unit volume: σ = Neμ, where σ is the conductivity and μ is the mobility. The mobility of the rare gas solids is high, but their conductivity is nonetheless low because there is a small number of conduction electrons.

Electrical insulators
      Like the rare gas solids, most ionic solids are electrical insulators. In sodium chloride (salt), for example, each sodium atom donates its single valence electron to a chlorine atom, thus forming a solid composed of Na+ and Cl ions. All electrons are in filled shells at low temperature, and in a perfect crystal there are no conduction electrons. Sodium chloride is thus an insulator with a very high resistivity. Some conduction electrons are provided by impurities or thermal excitations. At high temperatures large ion vibrations from thermal fluctuations may knock an electron out of a filled shell, upon which it becomes a conduction electron and contributes to the conductivity. The number of conduction electrons created by thermal excitations is small for most insulators. Although defects can be responsible for producing conduction electrons, they can also destroy the conducting ability of electrons by trapping them. The defects have local orbitals that provide a lower energy state for the electron than the one occupied in the conduction state. A conduction electron becomes bound at the defect, ceasing to contribute to the conductivity. This process is very efficient in insulators, so the few conduction electrons provided by impurities and thermal fluctuations are usually trapped at other defects. By definition, an insulator is a solid that does not provide a stable environment for conduction electrons.

Conductivity of metals
      Metals have a high density of conduction electrons. The aluminum atom has three valence electrons in a partially filled outer shell. In metallic aluminum the three valence electrons per atom become conduction electrons. The number of conduction electrons is constant, depending on neither temperature nor impurities. Metals conduct electricity at all temperatures, but for most metals the conductivity is best at low temperatures. Divalent atoms, such as magnesium or calcium, donate both valence electrons to become conduction electrons, while monovalent atoms, such as lithium or gold, donate one. As will be recalled, the number of conduction electrons alone does not determine conductivity; it depends on electron mobility as well. Silver, with only one conduction electron per atom, is a better conductor than aluminum with three, for the higher mobility of silver compensates for its fewer electrons.

      In metals such as sodium and aluminum, the atoms donate all their valence electrons to the conduction band. The resulting ions are small, occupying only 10–15 percent of the volume of the crystal. The conduction electrons are free to roam (free-electron model of metals) through the remaining space. A simple model, which often describes well the properties of the conduction electrons, treats them as interacting neither with the ions nor with each other. The electrons are approximated as free particles wandering easily through the crystal. This concept was first proposed by the German scientist Arnold Johannes Wilhelm Sommerfeld (Sommerfeld, Arnold). It works quite well for those metals, known as simple metals, whose conduction electrons are donated from sp-shells—for example, aluminum, magnesium, calcium, zinc, and lead. They are called simple because they are aptly described by the simple theory of Sommerfeld.

      The transition metals (transition element) are found in three rows of the periodic table: the first row consists of scandium through nickel, the second row is yttrium through palladium, and the third row is lanthanum plus hafnium through platinum. Within these rows, as the atomic number increases, the electrons fill d-states in the outer shell of the atom. In crystal form the transition metal atoms are metals with interesting properties. The d-electrons are more tightly bound to the ion centre than are sp-electrons. While the sp-valence electrons become conduction electrons that move freely through the crystal, the d-electrons tend to stay localized near the ion. Neighbouring ions may covalently bond d-electrons. In most cases, these d-states are only partially filled. Electrons in these d-states can conduct as well as those in the sp-states, but the electron motion in the d-states is not well approximated by the Sommerfeld model of free particles. Instead, the electrons move from ion to ion through the shared covalent bonds of the d-electrons. These metals have some conduction electrons donated from sp-states and others from d-states; therefore, some electrons move freely according to the Sommerfeld model, while others move through the bonds. Each electron switches back and forth between these two modes of conduction, resulting in electron motion that is quite complicated.

      An applied voltage causes the electrons of metals to accelerate and contribute to the electric current. The electrons scatter occasionally from imperfections in the crystal, and the rate of scattering determines the mobility. The electrons do not scatter from the ions in the crystal that are located at the expected site in the crystal lattice. The electrons move to accommodate the host ions rather than scatter from them. If an ion is missing, misplaced, or of a different species, however, the electron will scatter from this defect. Ions vibrate around their lattice site, with the amplitude vibration increasing with temperature. The vibration may cause the ion to be displaced from its crystal site, providing a defect from which an electron will scatter. The resistivity of metals increases at high temperature, owing to the increase in vibrations of the ions in the crystal and the resulting increase in scattering.

Conducting properties of semiconductors (semiconductor)
      Semiconductors have conducting properties intermediate to those of insulators and metals. In some cases the semiconductors are insulators, while in others they are metals. Semiconductors share with insulators the property that they have no conduction electrons in a perfect crystal without thermal fluctuations. Conduction electrons are provided by electrons from impurities or by thermal fluctuation of electrons from atomic shells. The important difference between insulators and semiconductors is in the nature of the traps (trap). A trap is a local electron energy state at a defect. Although the traps in insulators bind conduction electrons tightly, those in semiconductors only weakly bind the electrons. A trapped conduction electron in a semiconductor can be kicked back to the conduction band by thermal fluctuations. At room temperature, the majority of extra electrons are found in the conduction band rather than in traps. The inability of traps to keep electrons is the main difference between semiconductors and insulators. A semiconductor at room temperature has a sufficient number of conduction electrons to provide good electrical conductivity. Since the mobility of electrons in many semiconductors is exceptionally high, even a small number of conduction electrons is generally sufficient to allow high conductivity.

       phosphorus has five valence electrons, while silicon has four. When a phosphorus atom substitutes for an atom in a silicon crystal lattice, four of its five valence electrons enter covalent bonds. The fifth one is extra, sitting in a shallow trap around the phosphorus site. It is easily excited, however, to the conduction band by thermal fluctuations. At room temperature, there is nearly one conduction electron in silicon for each phosphorus impurity. By controlling the number of impurities, it is possible to control the conductivity of silicon. Other substitutional atoms such as arsenic and antimony also serve as electron donors to the conduction band of silicon.

      If a sufficient number of conduction electrons are added to a semiconductor through the introduction of impurities, the electrical properties become metallic. There is a critical concentration of impurities Nc, which depends on the type of impurity. For impurity concentrations less than the critical amount Nc, the conduction electrons become bound in traps at extremely low temperatures, and the semiconductor becomes an insulator. For a concentration of impurities higher than Nc, the conduction electrons are not bound in traps at low temperatures, and the semiconductor exhibits metallic conduction. For phosphorus impurities in silicon, Nc = 2 × 1018 impurities per cubic centimetre. Although this number seems large, it represents about one phosphorus atom for each 100,000 silicon atoms. On a percentage basis, a small number of phosphorus atoms will change silicon from an insulator to a metallic conductor. Other semiconductors have similar properties. In gallium arsenide the critical concentration of impurities for metallic conduction is 100 times smaller than in silicon.

      Gallium atoms (gallium), like those of phosphorus, can be used as substitutional impurities in silicon. Each atom contributes three electrons to covalent bonds. Since four electrons are needed to complete a tetrahedral arrangement, there is one electron absent per gallium atom from a full set of covalent bonds. The missing electron is called a hole. Holes can move around the crystal in a process similar to the motion of ion vacancies, except in this case there is an electron vacancy. An electron from a nearby covalent bond can jump over and fill the empty electron state, thereby moving the hole to the neighbouring bond. The hole contributes a positive charge, since it is the absence of an electron. The mobility of holes in response to an external voltage is almost as high as the mobility of conduction electrons. A semiconductor may have a high density of impurities that cause holes, and a high electrical conductivity is created by their motion. A p-type semiconductor is one with a preponderance of holes; an n-type semiconductor has a preponderance of conduction electrons. The symbols p and n come from the sign of the charge of the particles: positive for holes and negative for electrons.

      Thermal fluctuations can excite an electron out of a covalent bond, making it a conduction electron. The bond is left with a missing electron, which constitutes a hole. Thermal fluctuations thus make electron-hole pairs. Usually the electron and hole separate in space, and each wanders away. The Swiss-American scientist Gregory Hugh Wannier first suggested that the electron and hole could bind together weakly. This bound state, called a Wannier exciton, does exist; the hole has a positive charge, the electron has a negative charge, and the opposites attract. The exciton is observed easily in experiments with electromagnetic radiation. It lives for only a short time—between a nanosecond and a microsecond—depending on the semiconductor. The short lifetime is due to the preference for the electron to reenter a covalent bond state, thereby eliminating both the hole and the conduction electron. This recombination of electron and hole is easily accomplished from the exciton state, since the two particles are spatially nearby. If the electron and hole escape the exciton state by thermal fluctuation, they travel away from each other. Recombination is then less probable, since it occurs only when the wandering particles pass close to one another again. Recombination also can occur at defect sites. First, one particle becomes bound to the defect, followed by the second particle. The electron and hole are again close to one another, and the electron can reoccupy the covalent bond.

      As in metals, the mobility of electrons in semiconductors is limited by electron scattering. For crystals with few defects, the mobility is limited by defect scattering at the lowest temperatures and by ion vibrations at moderate and high temperatures. Since semiconductors with few defects have a small number of conduction electrons, the resistivity is high. The number of conduction electrons is increased in semiconductors by adding impurities. Unfortunately, this also increases the scattering from impurities, which reduces the mobility. Figure 8 shows the resistivity of silicon at room temperature (T = 300 K) as a function of the concentration of impurities. The two curves represent conduction by electrons and by holes. Each grid mark on the graph is a factor of 10. The resistivity varies by a factor of one million from the lowest to the highest concentration of impurities.

      Semiconductors with few impurities are good photoconductors. photoconductivity is the phenomenon in which the electrical conductivity of a solid is increased by exposing it to light. Light is electromagnetic radiation within a specific narrow band of frequencies. The quanta of light are absorbed by the semiconductor, creating electron-hole pairs that provide the electrical conduction. More intense light produces more electron-hole pairs and gives rise to better conductivity. Each semiconductor absorbs light over a specific frequency range, so different semiconductors are used as photoconductors for different ranges of frequency.

      Zinc oxide (ZnO) is an interesting material with respect to conductivity. It crystallizes in the wurtzite structure, and its bonding is a mix of ionic and covalent. High-purity single crystals are insulators. Zinc oxide is the most piezoelectric (piezoelectricity) of all materials and is widely used as a transducer in electronic devices. (Piezoelectricity is the property of a crystal to become polarized when subjected to pressure.) Zinc oxide is a good semiconductor when aluminum impurities are included in the crystal. Polycrystalline ceramics of semiconducting zinc oxide conduct well and obey Ohm's law. The addition of small amounts of other oxides, such as those of barium and chromium, causes zinc oxide ceramics to have very nonohmic electrical properties; the electrical current (electric current) in such ceramics is the most nonlinear of any known material. The current I becomes proportional to a power of the voltage Vn, where the exponent n has values of more than 100 in certain ranges of voltage. This material is called a varistor, which is a contraction of the words variable and resistor. Zinc oxide varistors are widely used as circuit elements to protect against voltage surges. Figure 9 shows a graph of current versus voltage for a zinc oxide varistor used in household electronics. There is little current until a critical voltage of about 330 volts is reached, at which point the current rises steeply in a nonlinear fashion. Another interesting application of zinc oxide was its former use as a white pigment in paint. It has been replaced by titanium dioxide (TiO2), however, which is whiter.

Explanation of magnetism
      Electrons (electron) are perpetually rotating, and, since the electron has a charge, its spin produces a small magnetic moment. Magnetic moments are small magnets with north and south poles. The direction of the moment is from the south to the north pole. In nonmagnetic materials the electron moments cancel, since there is random ordering to the direction of the electron spins. Whenever two electrons have their moments aligned in opposite directions, their effects tend to cancel. Magnets (magnet) are formed when a large number of the electrons align their individual moments in the same direction. Only a small percentage of crystals are magnetic. The forces that tend to align the electron spins are subtle. There are three separate parts to the explanation of magnetism. They are as follows:

      1. Most magnets are composed of atoms (atom) whose valence electrons (valence electron) are in d- or f-shells. Atomic shell notation refers to angular momentum, where s has zero unit, p has one, d has two, and f has three. Electrons in d-shells tend to be bound to the ion, and those in f-shells are bound even more tightly.

      2. Each electron orbital can be occupied by two electrons—one with spin up and one with spin down. The d-shell has five orbital states and 10 electrons when filled; the f-shell has seven orbital states and 14 electrons when filled. Electrons are added one at a time to the d-states according to the empirical rule that the electrons arrange themselves in the state with the maximum spin and the maximum magnetic moment. If the first electron has spin up, the next four will also have spin up. A maximum of five electrons with spin up are allowed in the d-shell, so the sixth must have spin down. Similarly, the f-shell (shell atomic model) accepts seven electrons with the same direction of spin before taking electrons with the opposite spin orientation. The order in which electrons fill atomic shells is described by Hund's rules, of which the first is maximizing the total spin. Atoms with electrons in partially filled d- and f-shells usually have a nonzero total spin and thus a net magnetic moment. These magnetic ions (ion) are the building blocks for magnetic crystals.

      Hund's first rule is due to a phenomenon called electron exchange. As discussed above, a fundamental rule of quantum mechanics, the Pauli exclusion principle, states that no two electrons with the same direction of spin can occupy the same point in space at the same time. Electrons have charge and repel one another. If two electrons come close together, a large amount of repulsive energy is produced. Physical systems prefer the state of lowest energy, and so electrons avoid such close approach. When their spins are parallel, electrons avoid each other because of the Pauli principle. Electrons in the same shell thus prefer to have their spins parallel, since this configuration keeps the electrons apart and thereby reduces the amount of repulsive energy. The concept of electron exchange is the basis of magnetism. It explains why ions such as iron have large magnetic moments. In divalent iron (Fe2+) the six d-electrons are arranged to achieve maximum electron spin and magnetic moment.

      3. Individual ions with fixed magnetic moments may cooperatively align their moments, resulting in the presence of magnetic properties of the crystal as a whole. Ferromagnet (ferromagnetism) crystals have the magnetic moments from all their constituent ions aligned in the same direction; the magnetic moment of the crystal is the summation of the individual moments of the ions. There must be a magnetic force between the different ions that causes them to cooperatively align their moments. This force is also due to electron exchange. The d-orbitals from neighbouring ions overlap weakly into covalent bonds. The d-electrons on the separate ions are shared with the neighbour through covalent bonding. The electron exchange will tend to align the spins on the two neighbours. Aligning all pairs of neighbours aligns all ions. The exchange force between neighbours is much weaker than the force within the atomic shell of one ion. Although weak, the force is sufficient to cause ferromagnetism.

Ferromagnetic materials
       iron is a typical ferromagnet. Not all bars of iron are magnets; the existence of magnetism is determined by the nature of the domains within the bar. A domain is a region of a crystal in which all the ions are ferromagnetically aligned in the same direction. A bar may be composed of many domains, each having a different magnetic orientation. Such a bar would not appear to be magnetic. Each piece of the bar is magnetic, but the domains have moments that point in different directions, so the bar has no net moment. If the bar of iron is placed in a strong magnetic field, however, the bar becomes magnetic. The field causes the bar to become a single domain with all moments aligned along the external field. The domains do not rotate their moments; instead, the walls between domains move. The domain with a moment along the field grows, while the others become smaller. If removed from the magnetic field, the iron bar will remain magnetized for a considerable time period. Nearly all bars of iron are polycrystalline: they have many small grains (grain) of single crystals, which are packed together with random orientation. A grain could be a single domain, a domain could include many grains, or a large grain could have several domains.

      Ferromagnetic materials change their magnetic ordering at a characteristic temperature Tc called the Curie temperature (Curie point). The Curie temperatures for three common ferromagnetics—iron, cobalt, and nickel—are 1,043 K; 1,394 K; and 631 K, respectively. For temperatures below Tc the magnetic moments of the ions are aligned and the crystal is magnetic. For temperatures above Tc the crystal is not ferromagnetic, since the individual atomic moments are no longer aligned. Above Tc the moments have short-range order but not long-range order. Short-range order means there is local ordering. If a moment points in one direction, its neighbours have a tendency to point in the same direction. This tendency is maintained over several lattice sites but is not maintained for long distances. Long-range order is the tendency for moments to align for large distances. For temperatures a few degrees below Tc the moments have strong short-range order but only a small amount of long-range order, so the bar is not very magnetic. The tendency for long-range order increases at lower temperature. The Curie temperature is the point where long-range order begins as the temperature is lowered.

      If an iron bar is heated to a temperature above Tc, the bar is no longer magnetic. If the bar is then cooled to a temperature below Tc, the grains become magnetic, but they orient their moments in random directions, so the bar as a whole is not magnetic. A bar can be demagnetized by heating the bar and then cooling it. By inserting it in a large magnetic field, the bar can be remagnetized.

      Ferromagnetism is found in many insulators (insulator) as well as metals. Chromium bromide (CrBr3) is an insulator since chromium is trivalent and a bromine atom needs one electron to complete its outer shell. The trivalent chromium atoms each have a moment, and these align ferromagnetically below the Curie temperature of 37 K. Gadolinium chloride (GdCl3; Tc = 2.2 K) and europium oxide (EuO; Tc = 77 K) are two other examples among many.

Antiferromagnetic materials (antiferromagnetism)
 Many crystals have magnetic ions that are ordered in arrangements other than ferromagnetic. In antiferromagnetic ordering, the moments pointing in one direction are balanced by others pointing in the opposite direction, with the result that the substance has no net magnetization. The exchange interaction between ions in this case has the opposite sign and favours the alternate arrangements of spins. The sign of the exchange interaction between ions depends on the length of the covalent bond and the bonding angles; it may have either orientation. The characteristic temperature associated with antiferromagnetism is called the Néel temperature TN. Below TN the ions are antiferromagnetically ordered, while above this temperature there is no long-range antiparallel order. Some examples of antiferromagnetic crystals are manganese oxide (MnO; TN = 116 K), manganese sulfide (MnS; TN = 160 K), and iron oxide (FeO; TN = 198 K). Manganese oxide is an insulator since manganese atoms are divalent and oxygen atoms accept two electrons. The manganese ion has a fixed magnetic moment. The crystal structure of manganese oxide is the same as that of sodium chloride shown in Figure 3B—>. Below the Néel temperature the atomic unit cell doubles in size to include two atoms of each type of ion. This is necessary because below TN neighbouring manganese atoms have moments in the opposite direction and are no longer equivalent; the unit cell must therefore include one moment in each of the two directions. Fluorides such as manganese fluoride (MnF2), iron (II) fluoride (FeF2), cobalt fluoride (CoF2), and nickel fluoride (NiF2) are other crystals that exhibit antiferromagnetic ordering of the transition metal ions.

Ferrimagnetic materials (ferrimagnetism)
      Ferrimagnetism is another type of magnetic ordering. In ferrimagnets the moments are in an antiparallel alignment, but they do not cancel. The best example of a ferrimagnetic mineral is magnetite (Fe3O4). Two iron ions are trivalent, while one is divalent. The two trivalent ions align with opposite moments and cancel one another, so the net moment arises from the divalent iron ion. The historic lodestone that was the first magnetic material discovered was a form of magnetite. Another class of ferrimagnets has the garnet structure; Y3Fe5O12 is one such crystal. Only the iron ions are magnetic. Three point in one direction and two in the other, so there is a net magnetic moment of one iron ion in the unit cell. However, if a rare-earth ion (rare-earth element) such as gadolinium (Gd) is substituted for yttrium (Y), then the rare-earth ion also contributes to the ferrimagnetism.

      Ferrites (ferrite) are oxides of iron with the formula MFe2O4, where M is a divalent ion such as nickel, zinc, cadmium, manganese, or magnesium. The ferric iron ion is trivalent, and the oxygen ion accepts two electrons. Actually M can also be divalent iron, forming magnetite (Fe3O4). The crystal structure is called spinel, which is the mineral name for MgAl2O4. Ferrites are electrical insulators with magnetic ordering. Their insulating quality makes them useful as magnetic cores. When metallic ferromagnetic materials are exposed to alternating magnetic fields, significant heating losses occur from eddy currents. Ferrite magnets greatly reduce such heat losses because of their high resistivity. They also absorb electromagnetic radiation of very long wavelengths. This property is unusual, since metals reflect such radiation while insulators transmit it. In small crystallites of ferrites, the electromagnetic radiation causes the magnetic moment to rotate at the frequency of light. The absorption frequency of the light depends on the detailed shape of the crystal grain. A polycrystalline material, with a range of grain sizes and shapes, absorbs over a broad range of radar frequencies. Ferrite crystals are a major ingredient in snooper paint, which makes stealth airplanes undetectable by radar. Regular airplanes reflect radar, and the reflected signals can be used to locate the aircraft. Stealth planes absorb the radar signals and so cannot be located in this way.

The Kondo effect
      Magnetic ions have interesting properties when they are found as impurities in nonmagnetic crystals. They usually retain their magnetic moment, so small magnets are distributed randomly throughout the crystal. If the host crystal is a metal, the magnetic impurities make an interesting contribution to the electrical resistivity. The conduction electrons scatter from the magnetic impurity. Since the conduction electron and the impurity both have spin, they can mutually flip spins while scattering. The spin-flip scattering is strong at low temperatures and actually increases slightly as temperature decreases. This phenomenon is called the Kondo effect after the Japanese theoretical physicist Jun Kondo, who first explained the increase in resistivity resulting from magnetic impurities. There is a characteristic temperature, called the Kondo temperature, which depends on the impurity and on the metallic host. The resistivity increases at low temperature, starting near the Kondo temperature. A typical example of a Kondo system is iron impurities in copper; (copper) the system's Kondo temperature is 24 K. The solid line in Figure 10 shows the resistivity in copper at low temperature when there are 110 iron impurities per 1,000,000 copper atoms. The dashed line a is the resistivity in the absence of impurities. It increases at higher temperature because the electron scatters from ion vibrations. The dashed line b is the resistivity from spin-flip scattering. Nearly all transition metal atoms are found as magnetic impurities in copper or gold. Each system has a different Kondo temperature, which varies from 1,000 K to a fraction of 1 K. The spin-flip part of the electrical resistivity is unique in that it is large at low temperatures and decreases at high temperatures; most contributions to the resistivity increase at high temperatures.

Gerald D. Mahan

Additional Reading
Works on solids in general include Lawrence H. Van Vlack, Elements of Materials Science and Engineering, 6th ed. (1989), an elementary textbook; Charles Kittel, Introduction to Solid State Physics, 8th ed. (2005), a standard college textbook; and Linus Pauling, The Nature of the Chemical Bond and the Structure of Molecules and Crystals, 3rd ed. (1960, reissued 1989), a classic work on chemical bonding.For crystalline solids in particular, Alan Holden and Phylis Morrison, Crystals and Crystal Growing (1960, reissued 1982), is a readable illustrated treatment; Richard P. Feynman, The Feynman Lectures on Physics: From Crystal Structure to Magnetism, vol. 3 (1999), is also highly accessible as a sound recording.Richard Dalven, Introduction to Applied Solid State Physics, 2nd ed. (1990), is an intermediate-level textbook on semiconductors and semiconducting devices.Books on magnetism include David Jiles, Introduction to Magnetism and Magnetic Materials, 2nd. ed. (1998); and Robert C. O'Handley, Modern Magnetic Materials: Principles and Applications (2000).Fullerenes are reviewed in Wanda Andreoni, The Physics of Fullerene-Based and Fullerene-Related Materials (2000); and Mildred S. Dresselhaus, Gene Dresselhaus, and Phaedon Avouris (eds.), Carbon Nanotubes: Synthesis, Structure, Properties, and Applications (2001).

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Universalium. 2010.

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