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CHEM 230 (5)
Lecture

Class Notes 5 Summer 2012.pdf

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Department
Chemistry
Course
CHEM 230
Professor
Jennifer Warriner
Semester
Summer

Description
Main Group Chemistry • Chemistry of the s- and p-block elements • Only part of periodic table containing non-metals? • (Metals) / (Semi-metals or Metalloids) / (Non-metals) • Trend from metals to non-metals • Electrical Properties: Metals Metalloids Non-metals • Bonding in metals: Localized metal-metal bonding is not possible • Understood via Band Theory: 85 • Conductors – high mobility of electrons • Bonding in Li metal– overlap of 2s orbitals • 2s orbitals of same energy • Not all MOs populated, but very close in E • Electrons can easily move into vacant orbitals • Can also include valence p-orbitals • s and p orbital energies impact energy separation of MOs • Insulators: large difference in s and p orbital energies 86 • Difficult to populate unfilled MOs • High χ results in valence electrons tightly bound • Metalloids: Intermediate between two extremes • Semiconductors – conductivity increases with increasing temperature • foundation of modern electronics • Band gap – energy needed to promote and electron C (5.39eV); Si(1.10eV); Ge(0.66eV); αSn(0.08eV) • n- and p-type semiconductors – doping with Group 13 or 15 • p-type - positive hole, doped with Group 13 (e.g. Ga doped Si) • n-type – negative hole, 87 The Chemistry of Hydrogen • Brief look at properties and reactivity • Great importance in theoretical chemistry rd • Most abundant element in universe, 3 on earth (O and Si) Properties –A colourless odourless gas at 298 K • Isotopes: Deuterium ( ) and Tritium ( ) The Hydrogen ion (proton) • Ionization energy: • Forms the oxonium ion in solution: The Hydride ion ∆H EA : • Overall energy to form H - = +145 kJ/mol (g) • compare to F -(g)= -249 kJ/mol and Cl -(g)= -228 kJ/mol ∴ Ionic hydrides are less stable– more reactive Formation of H 2 • H has many industrial applications Haber process (forming NH ) 2 3 • Formed by the Water-gas shift reaction: 88 • Reactivity: H is2not very reactive under ambient conditions: • Radical chain reaction with O : 2 • Radical chain reactions also occur with the halogens • Fuel Cells • Overall reaction: • Major problem is transport and utilization of H 2 89 Group 1 Elements • Alkali metals, ground state: Properties – Low ionization energies, form ionic compounds 1+ 2+ • Salts with M , noble gas configuration (M to high in E) Industrial Uses: • NaCl and NaOH • K salts as fertilizers (95% of demand) • Potash (Canada!) • Li – medicine, batteries Reactivity: (a) Vigorous reaction with H O 2 • Reactivity increase with H Odown the group 2 (b) React with halides (ionic lattices) (c) Formation of NaOH (the chloralkali process) • Electrolysis of NaCl (aq) 90 (d) Reducing agents – e.g. LiH, NaH, NaBH , L4AlH 4 Biological Aspects: • Na and K ions are essential to life • Charge balance of negative protein units • Maintenance of osmotic pressure in cells, membrane potentials Group 2 Elements • Alkali earth metals: • Be behaves differently - Small cation, very high charge density Properties: • Soft, silver coloured metals • M , form ionic solids (except Be) • Salts used as drying agents Industrial Uses: • Be and Mg – improve steel properties • CaO as a binding agent Reactivity: With O 2 91 Biological Aspects: • Mg in chlorophyll – 2+ 2+ • Mg uptake in cells, Ca in body fluids Group 13 Elements • 3+ oxidation state common • High abundance of Al in earth’s crust Properties: • Boron is a non-metal • B ; extended structures, allotropes • no trend in MPs, elements form different structures • M very high charge density Industrial Uses: • B in borosilicate glass • Al greatest commercial importance • GaAs - semiconductor Reactivity: • ExpectAl to be very reactive with O ! 2 • Why does this not happen to our cans? 92 • Neutral hydrid
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