SOLID STATE CHEMISTRY Crystal Structure Solids are divided into 2 categories: I. Crystalline possesses rigid and long-range order; its atoms, molecules or ions occupy specific positions, e.g. ice II. Amorphous lack a well-defined arrangement and long-range molecular order, e.g., glass. Unit cell basic repeating structural unit of a crystalline solid. Each sphere represents an atom, ion or molecule, called a lattice point. CHEM261HC/SS3/01 Unit cell is described by 6 parameters. Axes are a, b and c angles between the axes are, and. ( is the angle between b and c, etc.) The position of an atom in the unit cell is given by fractions of the unit cell axes, (x/a, y/b, z/c), called fractional coordinates. c a b Fig 1 An origin is defined in the unit cell and has position (0,0,0). Other coordinates are given relative to this origin, e.g., an atom in the centre of the unit cell has fractional coordinates of (½,½,½). Unit cells are chosen by a variety of ways, but the most preferable one is the smallest cell that shows greatest symmetry. CHEM261HC/SS3/02
Consider the following figures: Fig 2.1 Fig 2.2 In the 2-d pattern shown above, there are a variety of unit cells, each of which repeats the exact contents of the box. Fig 2.2 is preferred to Fig 2.1 because it is half the area. CHEM261HC/SS3/03 Every crystalline solid can be described in terms of one of seven types of unit cells. simple cube a = b = c α = β = γ = 90 0 tetragonal a = b c α = β = γ = 90 0 orthorhombic a b c α = β = γ = 90 0 rhombohedral a = b = c α = β = γ 90 0 monoclinic a b c triclinic a b c hexagonal γ α = β = 90 0 α β γ 90 0 a = b c α = β = 90 0, γ = 120 o
Note: it is difficult to draw 3-D figures; use 2-D projection along a specific direction. For a simple cube: 1,0 1,0 Fig 4.1 Fig 4.1.1 1,0 Fig 4.1.2 1,0 Simple cubic unit cell contains 1 lattice point. CHEM261HC/SS3/05 More complex lattice types are body-centred and face-centred with 2 and 4 lattice points in each unit cell respectively (Fig 4.2 and 4.3) Fig 4.2 Fig 4.3 For body centered cube (bcc): 1,0 1,0 1,0 1,0 Fig 4.2.1 Fig 4.2.2 CHEM261HC/SS3/06
Lattice point at each corner and one at the centre of the unit cell. i.e., 8 lattice points on corners, each shared by 8 unit cells, ⅛ per unit cell. Total: 8 x ⅛ + 1 = 2 lattice points in bcc. Exercise Show that the face-centered lattice has 4 lattice points. CHEM261HC/SS3/07 Atomic Radius Fig 5.1 scc Fig 5.2 bcc Fig 5.3 fcc For scc ( r - host atom radius; a unit cell length) a = 2r bcc a = 4r (3) ½ fcc a = [(8) ½ ]r CHEM261HC/SS3/08
Gold (Au) crystallizes in a cubic close-packed structure (fcc) and has a density of 19.3 g cm -3. Calculate the atomic radius of gold in picometres. Exercise 1. Metallic iron crystallizes in a simple cubic lattice. The unit cell edge length is 287 pm. The density of iron is 7.87 g cm -3. How many iron atoms are within a unit cell? 2. When silver crystallizes, it forms face-centered cubic cells. The unit cell edge length is 408.7 pm. Calculate the density of silver.
The way spheres are arranged in layers determines the type of unit cell. Simplest case, layer of spheres can be arranged as shown. A 3-dimensional structure can be obtained by placing layers directly above and below this layer. x Focusing on sphere x, see that it is in contact with 6 spheres, 4 around, 1 above and 1 below. Each sphere in this arrangement is said to have a coordination number of 6 because it has 6 immediate neighbours. Coordination number is defined as the number of atoms (or ions) immediately surrounding an atom (or ion) in a crystal lattice. Its value gives us a measure of how tightly the spheres are packed together, the larger the coordination number, the closer the spheres are to each other. The basic, repeating unit in the array of spheres described is called a simple cubic cell (scc). The other types of cubic cells are the body-centred cubic cell (bcc) and the face-centred cubic cell (fcc). Simple cubic A body-centred cubic arrangement differs from a simple cube in that the second layer of spheres fits into the depressions of the first layer and the third layer into the depressions of the second layer. bcc The coordination number of each sphere in this structure is 8 (each sphere is in contact with 4 spheres in the layer above and 4 spheres in the layer below. CHEM261HC/SS3/12
In the face-centred cubic cell, there are spheres at the centre of each of the 6 faces of the cube, in addition to the 8 corner spheres. fcc Most efficient arrangement of spheres. Start with the structure which we call layer A. Focusing on x, you can see that it has 6 immediate neighbours in the layer. x Layer A In the second layer (which we call layer B), spheres are packed into the depressions between the spheres in the first layer so that all the spheres are as close as possible to each other. CHEM261HC/SS3/13 There are 2waysthat a third layer sphere may cover the second layer to achieve closest packing. The spheres may fit into depressions so that each third layer sphere is directly over the first layer sphere. Since there is no difference between the arrangement of the first and third layers, we call the third layer, layer A. The ABAB.. arrangement that is obtained is called an hexagonal close-packing (hcp) structure. Alternatively, the third-layer spheres may fit into the depressions of the second layer. In this case, we call the third layer, layer C. This ABCABC. arrangement is called cubic close-packing (ccp).
In both the structures, each sphere has a coordination number of 12 (each sphere is in contact with 6 spheres in its own layer, 3 spheres in the layer above and 3 spheres in the layer below). The unoccupied space in a close-packed structure amounts to 26% of the total volume. Close-packed single layer hexagonal close-packing (hcp) ABAB.structure cubic close-packing (ccp) ABCABC.structure Calculating the unoccupied space in a close-packed array Since the space occupied by hard spheres is the same for ccp and hcp arrays, choose the geometrically simpler structure, ccp, for the calculation. Consider the Figure 1: Figure 1:
Exercise Calculate the fraction of space occupied by identical spheres in a primitive cubic unit cell. Calculate the fraction of space occupied by identical spheres in a primitive cubic unit cell Solution
Holes in close-packed structures 2 types of holes or unoccupied space between the spheres: (i) An octahedral hole lies between two planar triangle of spheres in adjoining layers.(fig A) Fig A For a crystal consisting of N spheres, there are N octahedral holes. CHEM261HC/SS3/18 Calculating the size of an octahedral hole Calculate the maximum radius of a sphere that may be accommodated in an octahedral hole in a close-packed solid composed of spheres of radius r. Answer CHEM261HC/SS3/19
(ii) A tetrahedral hole, T, shown in the Fig B below, is formed by a planar triangle of touching spheres capped by a single sphere lying in the dip between them. Fig B The tetrahedral hole in any close-packed solid can be divided into two sets: in one, the apex of the tetrahedron is directed up (T). in the other, the apex points down (T ). In an arrangement of N close-packed spheres, there are N tetrahedral holes in each set, and 2N tetrahedral holes in all. In a close-packed structure of spheres of radius r, a tetrahedral hole can accommodate another hard sphere of radius no greater than 0.225 r. Ionic Solids In general, anions tend to be larger than cations. + - - + + - - + + The exact type of structure adopted depends on the relative sizes of the anions and the cations and the charges on each. Each type of ionic solid has one thing in common. The ions are in fixed positions and therefore ionic solids are poor conductors of electricity. CHEM261HC/SS3/21
Simple Ionic Crystals The formation of ions The basic process involved in the formation of an ionic compound is the transfer of one or more electrons from one type of atom to another. The resulting ions are held together by electrostatic attraction. The arrangement of the ions in the solid is the one which gives the highest electrostatic energy. To see what factors determine such an arrangement, let us consider the process of bringing up successive anions around a given cation. If there are already n anions surrounding the cation, the addition of a further anion produces a number of repulsions between its charge and the charges on the n anions already present. CHEM261HC/SS3/22 Thus, there are two opposing tendencies: One is to increase the attractive forces by making the coordination number of the cation as large as possible. Balance this by addition of more and more anions to increase the repulsive forces. When the two tendencies balance, the final structure results. The coordination numbers in a solid of a given formula, such as AB, depends on the number of the larger ions which may be packed around the smaller one. The stoichiometry in this example, 1:1, determines the coordination number of the larger ion. Generally, cations are smaller in size than anions and it is the number of anions which can pack around the cation that determines the coordination numbers and structures. CHEM261HC/SS3/23
For example, sodium chloride crystallizes in a structure where the sodium ion is surrounded by six chloride ions and the chloride ion has six sodium ions around it. The larger cesium cation allows a coordination number of eight in the structure of cesium chloride. NaCl CsCl ZnS The number of anions which can pack around a given cation may be determined from the ratio of the radii of the cation and anion. The radius ratio, r + /r -, may be used to give an indication of the likely coordination number for a salt of a given formula type. For example, in six coordination, a cross section through a site in the lattice is: CHEM261HC/SS3/25
Similar calculations may be carried out for all the coordination numbers. The results of which indicate the range of values for the radius ratio within which different coordination numbers should be stable. These are shown below: r + / r - : 0.155 to 0.23 to 0.41 to 0.73 to higher values C.N. : 3 4 6 8 CHEM261HC/SS3/26 The validity of this simple method of predicting the coordination number may be assessed by examining the structures of some AB and AB 2 compounds. Ionic solids of the formula AB generally adopt either the sodium chloride or cesium chloride structure. The common AB 2 structures are those of rutile (titanium dioxide) in which the coordination is 6:3 or of fluorite (calcium fluoride) where the coordination is 8:4. Rutile fluorite
A number of other structures adopted by crystals of more complex stoichiometry are: oxygen rhenium Spinel structure Rhenium oxide oxide ion metal ion in O h site metal ion in T h site oxygen calcium perovskite titanium Wurtzite structure (ZnS) CHEM261HC/SS3/28 Lattice enthalpy and the Born-Haber cycle The lattice enthalpy, H Lo, is the standard molar enthalpy change accompanying the formation of a gas of ions from the solid: MX(s) M + (g) + X - (g) H L o Because lattice disruptions is always endothermic, lattice enthalpies are always positive. Lattice enthalpies are determined from enthalpy data by using a Born-Haber cycle, a closed path of steps that includes lattice formation as one stage, such as that shown: CHEM261HC/SS3/29
K + (g) + e - (g) + Cl(g) 122 K + (g) + e - (g) + ½Cl 2 (g) -355 kj mol -1 425 K + (g) + Cl - (g) K(g) + ½Cl 2 (g) 89 K(s) + ½Cl 2 (g) x 438 KCl(s) The Born-Haber cycle for KCl, lattice enthalpy (kj mol -1 ) = -x The standard enthalpy of decomposition of a compound into its elements is the negative of its standard enthalpy of formation, f H o : M(s) + X(s) MX(s) f H o MX(s) M(s) + X(s) - f H o Likewise, the standard enthalpy of lattice formation from the gaseous ions is the negative of the lattice enthalpy: MX(s) M + (g) + X - (g) H o L M + (g) + X - (g) MX(s) - H o L For a solid element, the standard enthalpy of atomization, atom H o, is the standard enthalpy of sublimation, e.g. the process below K(s) K(g) sub H o =+89kJmol -1 CHEM261HC/SS3/31
For a gaseous element, the standard enthalpy of atomization is the standard enthalpy of dissociation, asin Cl 2 (g) 2Cl(g) dis H o = +244 kj mol -1 The standard enthalpy for the formation of ions from their neutral atoms is the enthalpy of ionization (for the formation of cations) and the electron-gain enthalpy (for anions). Two examples are K(g) K + (g) + e - ion H o = +425 kj mol -1 Cl(g) + e - Cl - (g) eg H o = -355 kj mol -1 The value of the lattice enthalpy, theonly unknown in a well chosen cycle, is found from the requirement that the sum of the enthalpy changes round a complete cycle is zero. Example Calculate the lattice enthalpy of KCl(s) using a Born-Haber cycle and the following information: H/(kJ mol -1 ) Sublimation of K(s) +89 Ionization of K(g) +425 Dissociation of Cl 2 (g) +244 Electron gain by Cl(g) -355 Formation of KCl(s) -438
Solution Exercise Calculate the lattice enthalpy of magnesium bromide from the data shown below: H/(kJ mol -1 ) Sublimation of Mg(s) +148 Ionization of Mg(g) to Mg 2+ +2187 Vaporization of Br 2 (l) +31 Dissociation of Br 2 (g) +193 Electron gain by Br(g) -331 Formation of MgBr 2 (s) -524
Solution Calculation of lattice enthalpies To calculate the lattice enthalpy of an ionic solid, we need to take into account several contributions to its enthalpy. These include the attractions and repulsions between the ions. This calculation yields the Born-Mayer equation for the lattice enthalpy at T =0: 2 N A z A z B e H L od 1 4 o d d o A where d o =r + +r - (the distance between centres of neighbouring cations and anions) N A is the Avogadro s constant
z A and z B the charge numbers of the cation and anion e is the fundamental charge o the vacuum permittivity d is a constant (normally 34.5 pm) used to represent the repulsion between ions at short range A is called the Madelung constant Madelung constants Structural type A Caesium chloride 1.763 Fluorite 2.519 Rock salt 1.748 Rutile 2.408 Sphalerite 1.638 Wurtzite 1.642 Example Using the Born-Mayer equation, estimate the lattice enthalpy of sodium chloride. Solution CHEM261HC/SS3/37
Exercise for the Idle Mind Estimate the lattice enthalpy of CsCl using the experimental d o = 356 pm CHEM261HC/SS3/38