Week 7 Chem131
Week 7 Chem131 CHEM131
Popular in Chemistry I - Fundamentals of General Chemistry
Popular in Chemistry
This 2 page Class Notes was uploaded by Maya Silver-Isenstadt on Friday October 14, 2016. The Class Notes belongs to CHEM131 at University of Maryland - College Park taught by John Ondov in Fall 2016. Since its upload, it has received 23 views. For similar materials see Chemistry I - Fundamentals of General Chemistry in Chemistry at University of Maryland - College Park.
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Date Created: 10/14/16
Chapter Seven 7.1) ● Paramagnetic = unpaired electrons ● Liquid oxygen is magnetic (diatomic) ● Lewis structure fails to predict magnetic character 7.2) ● Valence bond theory→ chemical bond is overlap between two half-filled atomic orbitals (AO’s) ● Hybridized atomic orbitals→ s,p,d,f orbitals ● Interaction energy is calculated as a function of internuclear distance. ● The interaction energy is usually negative (stabilizing) when the interacting atomic orbitals contain a total of two electrons that can spin-pair (orient with opposing spins) ● The geometry of overlapping orbitals determines shape 7.3) ● Orbitals in a molecule aren’t the same as orbitals in an atom. ● Hybridization → standard atomic orbitals combined to form hybrid orbitals. ● Greater the overlap, the lower the energy and the stronger the bond is. ● Hybrid orbitals minimize energy by maximizing orbital overlap. ● The number of standard atomic orbitals = the number of hybrid orbitals 3 ● sp orbital ● **Nitrogen orbitals in ammonia are sp hybrids. 2 ● sp → has one s and two p orbitals and one unhybridized p orbital. Trigonal planar. ● When p orbitals overlap side to side it makes a π bond. When they overlap end to end they make a sigma σ bond. ● Single = sigma ● Double = one sigma and one pi ● Triple = two pi and one sigma ● Compounds with same molecular formula but different structures are ISOMERS. ○ 7.4) ● Linear combination of atomic orbitals = a weighted linear sum- analogous to a weighted average- of the valence atomic orbitals of the atoms in the molecule. ● Bonding orbital = lower energy and increased electron density ● Antibonding= raises energy and decreases electron density. ● Different phases result in destructive interference. ● Bond Order = [# electrons in bonding molecular orbital - # in antibonding molecular orbital]/2 ● Order = 7.5) ● Electrons delocalize over the entire molecule 7.6) ● Electron sea→ When metal atoms donate greater than or equal to one electron. ○ Conductivity ○ Malleability ○ Ductile ● Semiconductors and Band Theory ○ Spacing gets small and it is just a band of energy levels ○ Bond Gap is large in insulators ○ Doping→ “holes” in valence bonds ○ Silicon and Phosphorus → n-type semiconductor ○ P-type → holes move opposite to electrons ○ P-n → junctions = diodes and amplifiers
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