# Ch 7 The Crystalline Solid State

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Ch 7 The Crystalline Solid State
Formulas and Structures Simple Structures Unit Cell = simplest repeating unit of the regular crystalline array Bravais Lattices = 14 possible basic crystal structure unit cell types

Atoms on corners and edges are shared between unit cells
Rectangular corners shared by 8 unit cells (8 x 1/8 = 1 total atom in cell) Other corners shared unequally, but still contribute 1 atom to unit cell Edges shared by 4 cells, contribute (4 x ¼ = 1 atom to unit cell) Faces shared by 2 cells, contribute ½ atom to each Angles and Dimensions can vary: triclinic has all different lengths and angles Lattice Points = positions of atoms needed to generate the whole crystal Body Centered Cubic: (0,0,0) = origin and (½ , ½, ½) = center All other atoms can be generated from these 2 by moving them exactly one cell length Cubic Structures Primitive Cubic is the simplest type To fully describe: length of side, 90o angles, (0,0,0) lattice point 8 x 1/8 = 1 atom in the unit cell each atom surrounded by 6 others (Coordination Number = CN = 6) Not efficiently packed; only 52.4% of volume is occupied (74.1% max) Vacant space at center with CN = r sphere would fit here

Body-Centered Cubic (bcc)
One more atom is added to the center of the cube Size of unit cell must increase over simple cubic Diagonal across center = 4r (r = radius of one atom) Corner atoms not in contact with each other due to cell size expansion Side = 2.31r (Calculate this in Ex. 7-2) Unit cell contains 1(1) + 8(1/8) = 2 atoms Lattice Points: (0,0,0) and ( ½ , ½ , ½ ) Close-Packed Structures Spheres will arrange to take up the least space, not cubic or bcc Two structures have almost identical packing efficiency (74.1%) Hexagonal Close Packing = (hcp) Cubic Close Packing = (ccp) = Face-Centered Cubic = (fcc) Both have CN = 12 for each atom: 6 in its layer, 3 above, and 3 below HCP has 3rd layer in identical position as first ABA; second layer above holes; CCP has 3rd layer displaced so above holes in first layer ABC Both have 2 tetrahedral (Td) holes per atom (CN = 4) formed by 3 atoms in one layer and one atom above/below Both have 1 octahedral (Oh) hole per atom (CN = 6) formed by 3 atoms in one layer and 3 atoms above/below 4r l l

HCP unit cell is smaller than the hexagonal prism
Take four touching atoms in 1st layer and extend lines up to 3rd layer 8(1/8) = 2 atoms in unit cell Dimensions = 2r, 2r, and 2.83r Angles = 120o, 90o, and 90o Lattice Points = (0,0,0,) and ( 1/3, 2/3, 1/2)

CCP unit cell One corner from layer 1, opposite corner from layer 4 6 atoms each from layers 2 and 3 in the unit cell Unit cell is fcc 8(1/8) + 6(1/2) = 4 atoms in the unit cell Lattice Points = (0,0,0), ( ½ , 0 , ½), ( ½ , ½, 0), (0 , ½ , ½)

Holes 2 Td and 1 Oh hole per atom in both (hcp) and (ccp) If ionic compound, small cations can fill these holes Td hole = 0.225r Oh hole = 0.414r NaCl has Cl- in (ccp) with Na+ (ccp) but also in the Oh holes Na+ = 0.695r, so it forces the Cl- to be farther apart, but still CN = 6 Metallic Crystals Most metals crystallize in (bcc), (ccp), or (hcp) structures Changing pressure or temperature can interchange these forms for a metal Must consider bonding, not just geometric packing only Soft and malleable metals usually have (ccp) structure (copper) Harder and more brittle metals usually have the (hcp) structure (zinc) Most metals can be bent due to non-directional bonding Weak bonding to all neighbors, not strong bonding to any single one Atoms can slide past each other and then realign into crystal form Dislocations = imperfections in lattice; make it easier to bend Impurities = other elements; allow slippage of layers Work Hardening = hammer until impurities are together Heat: can soften by dispersing impurities or harden if control cooling

Diamond Each atom is tetrahedrally bonded to 4 other carbon atoms Directional single bonds, unlike metals, cause hardness Binary Compound Structures Simplest structures just have 2nd element in holes of the first element’s lattice Small cations can fit in Td/Oh holes of large anions Large cations may only be able to fit in Oh holes Even larger cations may cause changes in structure if they don’t fit in holes Relative Number of Cations/Anions M2X won’t allow close-packing of anion lattice with cations in Oh holes because there are more cations than holes Alternatives: cations in Td holes, vacancies, not close-packed

Sodium Chloride Structure: NaCl
Na+ in (fcc) and Cl- in (fcc) Offset by ½ unit cell length Na+ are centered in edges of Cl- lattice (or vice versa) Most alkali halides have this same structure Large size difference of ions facilitate this structure Each Cl- has CN = 6 Na+ Each Na+ has CN = 6 Cl- Cesium Chloride Structure: CsCl Cs in simple cubic structure with Cl- in center (or vice versa) Cl- = 0.83Cs+ size (0.73r in center is ideal) Rare structure, need big cation (Cs, Tl only cations known with this structure)

Zinc Blende Structure: ZnS
Same as diamond structure with alternating Zn and S atoms Alternate: Zn and S each in (fcc) lattices combined so each ion is in a Td hole of the other lattice Stoichiometry: only ½ of the Td holes are occupied and ½ are vacant Wurtzite Structure: ZnS Rarer than Zinc Blende structure for ZnS; formed at higher Temperatures Zn and S each in (hcp) lattices combined so each ion is in a Td hole of the other lattice Again, ½ of the Td holes are vacant

Fluorite Structure: CaF2
Ca2+ in (ccp) lattice with 8 F- surrounding each and occupying all Td holes Alternate: F- in simple cubic lattice with Ca2+ in alternate body centers Nearly perfect radius fits for this structure Antifluorite Structure Reverse stoichiometry compounds like Na2O Every Td hole in the anion lattice is occupied by a cation

Nickel Arsenide Structure: NiAs
As atoms in close packed layers exactly above each other Ni atoms in all the Oh holes Both Ni and As have CN = 6 Alternate: Ni atoms occupy all Oh holes of (hcp) As lattice Usual for MX compounds where X = Sn, As, Bi, S, Se, Te Rutile Structure: TiO2 Distorted TiO6 octahedra forming columns by sharing edges Ti CN = 6; O CN = 3 Adjacent columns connected by sharing corners of octahedra Unit cell has Ti at corners and in the body center, 4 O in the faces, and 2 O in the plane of the body center Ti MgF2, ZnF2 are other examples

Complex Structures Substitution of cation or anion in a given structure is possible, giving a range of compounds with essentially the same structure. When charge or size is different in the substitution, the structure must change Balance charge by leaving vacancies Lattice can spread/contract out to accommodate larger/smaller ions Complex, non-spherical anions may distort the basic lattice shape Large cations may require increased coordination number (> 6) Calcite Stucture: CN = 6 LiNO3 NaNO3 MgCO3 CaCO3 FeCO3 InBO3 YBO3 Aragonite Stucture: CN = 9 CaCO3 KNO3 SrCO3 LaBO3 Hexagonal, CN = 6 O per metal Orthorhombic, CN = 9 O per metal

The Radius Ratio We can crudely predict CN by using the ratio r+/r- This assumes atoms are just packing as hard spheres (not really all that occurs) Examples: NaCl r+/r- = 113/167 (CN = 4) = 0.667 r+/r- = 116/167 (CN = 6) = Fits best with CN = 6 ZnS r+/r- = 74/170 (CN = 4) = Fits best with CN = 4 r+/r- = 88/170 (CN = 6) = CN = 4 in actual structure 4) Example: CaF2 has fluoride ions in a simple cubic array and calcium ions in alternate body centers, with r+/r- = What are the coordination numbers of the two ions predicted by r+/r- ? What are the coordination numbers observed? Predict coordination numbers of Ca2+ in CaCl2 and CaBr2.