CHEM 1212 1st Edition Lecture 3 Outline of Last Lecture I. Properties of Liquidsa. Vaporization and Condensationb. Vapor Pressurec. Vapor Pressure, Enthalpy of Vaporization, and the Clausius-Clapeyron EquationII. Crystal Lattices and Unit CellsOutline of Current Lecture I. Structures and Formulas of Ionic SolidsII. Bonding in Metals and SemiconductorsIII. SemiconductorsIV. Bonding in Ionic Compounds: Lattice EnergyV. The Solid State: Other types of solid materialsVI. Phase Changes Involving SolidsVII. Phase DiagramsCurrent LectureI. Structures and Formulas of Ionic Solidsa. 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 holesb. 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 saltThese 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.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  CsClii. The formula for an ionic compound is always reflected in the composition of its unit celliii. An ion in a unit cell that is surrounded by 6 oppositely charged ions is said to be in an octahedral holeiv. An ion in a unit cell that is surrounded by 4 oppositely charged ions is said to be in a tetrahedral hole1. There are 8 tetrahedral holes in a face-centered cubic unit cellv. 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 Xn- lattice. Example – CsCl 2. Mn+ ions in all the octahedral holes in a face-centered cubicXn- lattice. Example – NaCla. 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 metals3. Mn+ ions occupying half of the tetrahedral holes in a face-centered cubic Xn- lattice. Example – ZnS II. Bonding in Metals and Semiconductorsa. The molecular orbital (MO) theory can be used to describe metallic bondingb. 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 samplec. Example: Lithium i. A mole of lithium has 1 mol of valence electrons and they occupy the lower-energy bonding orbitalsii. 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 atomsd. Band theory: theory of metallic bonding stated abovee. 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 spini. There are not enough electrons to fill all molecular orbitals in metalsii. The lowest possible energy for a system occurs when all electrons are in the orbitals of the lowest possible energy1. This is only reached at 0 K2. At 0 K, the highest filled level is called the Fermi level3. 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 level4. 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 levela. These negative electrons and positive holes are what cause electrical conductivity in metalsb. Electrical conductivity occurs from the movement of the negative electrons and hole when an electric field is applied to the systemi. In the presence of the electric field, the negative electrons move toward the positive side, while positive holes move to the negative sidef. 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 statesi. When electrons move back down to their original energy level, they emit a photonii. 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 characteristicsi. Example: most metals are malleable and ductile (can be rolled intosheets 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 electronsiv. In contrast – solids like diamonds are rigid because they have localized bonds which put the atoms in fixed positions (movementwould require breaking the bonds)III. Semiconductorsa. 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, theon switch give it the energy to conduct electricity)b. The band theory used for metals helps us understand semiconductors as wellc. Bonding in semiconductors: the band gapi. Band Gap: unlike metals, with all molecular orbitals close togetherin a “band”, semiconductors have two distinguishable “bands” of molecular orbitals1. A lower-energy valence band and a higher energy conduction band2. In Group 4A: carbon (in diamonds), silicon, and germanium, the valence band is filled while the conductionband is empty3. 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 conductionband1. 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 band1. Since diamond can not create positive holes, it is