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
Sodium chloride exhibits ionic bonding. The sodium atom 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. 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 (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, 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
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 (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 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 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; 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 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 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 bonds fall into two categories. The first is the case in which the valence electrons are from the sp-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 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, 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 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, 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 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. 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 are crystals as they begin descending but thaw during their fall to Earth. 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.
