CHEM 1212 1st Edition Lecture 3 Outline of Last Lecture I Properties of Liquids a Vaporization and Condensation b Vapor Pressure c Vapor Pressure Enthalpy of Vaporization and the Clausius Clapeyron Equation II Crystal Lattices and Unit Cells Outline of Current Lecture I Structures and Formulas of Ionic Solids II Bonding in Metals and Semiconductors III Semiconductors IV Bonding in Ionic Compounds Lattice Energy V The Solid State Other types of solid materials VI Phase Changes Involving Solids VII Phase Diagrams Current Lecture I Structures and Formulas of Ionic Solids a An ionic compound s unit cell is made up of a primitive or face centered cubic lattice of ions with oppositely charged ions filling in the holes b The choice of lattice and number and location of holes filled with oppositely charged ions help us understand the relationship between the lattice structure and the formula of the salt These notes represent a detailed interpretation of the professor s lecture GradeBuddy is best used as a supplement to your own notes not as a substitute II i Example CsCl Cesium Chloride has a primitive cubic unit cell of Cl ions 1 4 of a Cl ion in each corner of the cube and a positively charged Cs ion is in the center of the unit cell in total there is 1 cesium ion and 1 chloride ion in the unit cell which gives us the formula ration CsCl ii The formula for an ionic compound is always reflected in the composition of its unit cell iii An ion in a unit cell that is surrounded by 6 oppositely charged ions is said to be in an octahedral hole iv An ion in a unit cell that is surrounded by 4 oppositely charged ions is said to be in a tetrahedral hole 1 There are 8 tetrahedral holes in a face centered cubic unit cell v Compounds with a MX formula 1 1 ratio commonly form 1 of 3 possible crystal structures 1 Mn ions occupying the cubic hole in a primitive cubic Xnlattice Example CsCl 2 Mn ions in all the octahedral holes in a face centered cubic Xn lattice Example NaCl a This structure is used by all the alkali metal halides except CsCl CsBr and CsI and all the oxides and sulfides of the alkaline earth metals 3 Mn ions occupying half of the tetrahedral holes in a facecentered cubic Xn lattice Example ZnS Bonding in Metals and Semiconductors a The molecular orbital MO theory can be used to describe metallic bonding b A metal is a supermolecule made up of many atoms and valence orbitals and to understand the bonds in metals you must look at all the atoms in a given sample c Example Lithium i A mole of lithium has 1 mol of valence electrons and they occupy the lower energy bonding orbitals ii The bonding is delocalized meaning that the electrons are associated with all the atoms in the crystal and not just with a specific bond between two atoms d Band theory theory of metallic bonding stated above e All molecular orbitals within a metallic bond are so close in energy level that they are indistinguishable from one another each molecular orbital can have two electrons of the opposite spin i There are not enough electrons to fill all molecular orbitals in metals ii The lowest possible energy for a system occurs when all electrons are in the orbitals of the lowest possible energy 1 This is only reached at 0 K 2 At 0 K the highest filled level is called the Fermi level 3 When the temperature of the system goes above 0 K the added energy will cause some of the electrons to occupy orbitals above the Fermi level 4 For each electron that moves to a higher energy level there become two singly occupied levels an orbital above the Fermi level with only 1 electron and a positive hole caused by the absence of an electron below the Fermi level a These negative electrons and positive holes are what cause electrical conductivity in metals b Electrical conductivity occurs from the movement of the negative electrons and hole when an electric field is applied to the system i In the presence of the electric field the negative electrons move toward the positive side while positive holes move to the negative side f Because the energy gaps between levels in metals are very small they can absorb energy of almost any wavelength causing electrons to jump to higher energy states i When electrons move back down to their original energy level they emit a photon ii It is this constant absorption and reemission of light that gives metal its shiny reflective appearance g The picture of molecular orbitals in metals explain many of its characteristics i Example most metals are malleable and ductile can be rolled into sheets or made into wires ii The reason this is possible is because of the delocalization of electrons not bonded in a particular place iii The atoms can easily slide past each other and remain bonded because of the coulombic attractions between the nuclei and electrons III iv In contrast solids like diamonds are rigid because they have localized bonds which put the atoms in fixed positions movement would require breaking the bonds Semiconductors a Semiconductor materials are given their name because they do not conduct electricity easily but can be prompted to do so by the input of energy all electronics contain semiconductors with on off switches the on switch give it the energy to conduct electricity b The band theory used for metals helps us understand semiconductors as well c Bonding in semiconductors the band gap i Band Gap unlike metals with all molecular orbitals close together in a band semiconductors have two distinguishable bands of molecular orbitals 1 A lower energy valence band and a higher energy conduction band 2 In Group 4A carbon in diamonds silicon and germanium the valence band is filled while the conduction band is empty 3 The gap acts as a barrier to keep electrons from entering the higher energy state this is unlike metals where electrons excite easily ii The reason semiconductors are able to conduct electricity is because thermal energy gives the electrons the energy needed to promote from the valence band across the gap to the conduction band 1 Conduction occurs when the electrons in the conduction band move in one direction while the positive holes in the valence band move in the other direction iii Diamond s band gap is so large that electrons can not make the jump from the valence band to the conduction band 1 Since diamond can not create positive holes it is an insulator nonconductor iv Intrinsic semiconductors naturally occurring property of the pure material i e pure silicon or germanium