2009.11.16 REVIEW

2009.11.16 REVIEW - Lecture 12 Topics Lewis Dot Structures...

Info iconThis preview shows page 1. Sign up to view the full content.

View Full Document Right Arrow Icon
This is the end of the preview. Sign up to access the rest of the document.

Unformatted text preview: Lecture 12 Topics Lewis Dot Structures shows connectivity Octet Rule: atoms share e- to complete octet Resonance movement of electrons Formal Charge charge on atom IF we assume all electrons are equally shared Lewis Dot Structures General Procedure: 1. Count number of valence electrons 2. Less electronegative (or lowest ionization energy) atom in center 3. Arrange other atoms 4. Include the remaining electrons to satisfy octet rule 5. Evaluate formal charges Resonance Resonance is the blending of structures. The electrons in the structure are moving, but connectivity is the same. H H H H H H H H H H H H Formal Charge FC = V - (L + 1/2 B) Formal charge on each atom ASSUMES that all electrons in the bonds are equally shared. Structures that minimize (closer to zero) formal charge contribute more to actual structure More favorable structures have the negative charge on more electronegative atom Lecture 13 Topics Octet Rule Exceptions radicals, ex. CH3 expanded octets, ex. PCl5 electron deficient, BF3 Types of bonds ionic vs. covalent electronegativity, dipole moment polarizability Bond strength/length dissociation energy Octet Rule Exceptions Radicals: unpaired electrons, very reactive, Expanded octets: more than 8 e only elements in period 3 or greater bigger and have empty d orbitals Incomplete octets: less than 8e elements of group 13 ex. BF3 H C H H F Xe F F F H B H H Types of Bonds Ionic NaCl Purely Covalent Cl2 Most compounds have bonds somewhere in the middle; ex. HCl: polar covalent Bond type predicted by electronegativity Electronegativity Electron pulling power of atom: follows Zeff Bond Strength / Length Bond strength affected by: atom size bond order Bond Strength / Length Stronger bonds tend to be shorter bonds Lecture 14 Topics Valence-shell electron-pair repulsion electron-electron repulsions dictate 3D structure electronic vs. molecular shapes Molecular dipole predicted by looking at bond dipoles and 3D structure VSEPR Electronic arrangement Molecular Shapes 1 lone pair, AXnE1 2 lone pairs, AXnE2 Molecular Dipoles Consider arrangement of bond dipoles in space F Xe F F F All bond dipoles cancel in XeF4, so there is no molecular dipole Lecture 15 Topics Valence bond theory bonding described overlap of AOs with unpaired electrons Types of bonds bonds bonds Hybridization AOs on the same atom mix together Valence Bond Theory AOs with unpaired electrons overlap to share electrons forming a bond F F F F and bonds How AOs overlap determines if it is a or a bond. Hybridization valence bond predicts that C makes 2 bonds - 2 unpaired electrons how do we explain 4 bonds? hybridization E 2p 2s sp3 Lecture 16 Topics Valence bond theory cannot explain O2 paramagnetism and electron deficiency Molecular orbital (MO) theory MOs are orbitals on the molecule AOs on different atoms are added to form MOs Homonuclear Diatomics Molecular Orbital Theory mathematical explanation of bonding using AO wavefunctions, molecular orbitals are orbitals on the molecule MOs formed by summing AOs on DIFFERENT atoms Homonuclear Diatomics Li2 N2 O2 F2 energies of MOs are dependent on AOs Lecture 17 Topics MO Theory (continued) heteronuclear diatomics polyatomic molecules, benzene Highest occupied MO vs. lowest unoccupied MO Link to everyday life conductors, insulators colors Heteronuclear Diatomics 4 E2p 2 3 1 E2s 2 1 O2s O2p more electronegative atom, lower E different naming scheme MO Theory: Polyatomics E increases with number of nodes HOMO - LUMO: Everyday Life solution color predicted by HOMO-LUMO gap insulators and conductors explained by E gap Lecture 18 Topics Stoichiometry relationship between quantities of matter in the reaction Balancing chemical reactions Precipitation reactions soluble vs. insoluble electrolytes Stoichiometry 2H2(g) + O2(g) 2H2O(g) Balanced chemical reaction indicate: how much reactant needed how much product formed always in moles! Conservation of mass and charge apply Precipitation Reactions Most ionic compounds are soluble dissolves in solvent NaCl in water Insoluble - doesn't dissolve sand in water Electrolytes - ions in water, conduct electricity Strong - all units form ions Weak - only a fraction form ions non-electrolyte - no ions Ionic, Net Ionic, Spectator ions Precipitation Reactions Some ionic compounds are not soluble Lecture 19 Topics Acids and Bases acids: H+ donor bases: H+ acceptor Strong vs. Weak Acid - Base reactions Acids Acids: H+ donor Strong: All molecules donate H+ Weak: Some fraction of molecules donate H+ Table J.1 HCN = acid Bases Bases: H+ acceptor Strong: All molecules accept H+ Weak: Some fraction of molecules accept H+ Table J.1 OH- = base Acid - Base Reactions Strong acid - strong base reactions form a salt (ionic compound) and water HNO3(aq) + NaOH(aq) NaNO3(aq) + H2O(l) Ionic - explicit about all ions Net ionic - ignores spectator ions, only highlights reaction of interest Lecture 20 Topics Redox reactions oxidation: loss of e reduction: gain of e Oxidation number: charge on atom IF we assume ionic bonding Limiting reactants % yield of reaction Redox Reactions LEO the lions goes GER Loss of electrons: oxidation Gain of electrons: reduction 2Mg(s) + O2(g) 2MgO(s) Oxidation: Mg Mg2+ + 2eReduction: O2 + 4e- 2O2- Oxidation Number This is the charge on atom ASSUMING bonds are ionic Oxidation Rules: know the rules! 2Mg(s) + O2(g) 2MgO(s) Mg is oxidized and the reducing agent O2 is reduced and the oxidizing agent Limiting Reactants Limiting reactant: less reactant than needed ex. 2H2(g) + O2(g) 2H2O(g) if I have 1 mole H2 and 1 mole O2, H2 is the limiting reactant amount of H2O is predicted by H2 % Yield of a Reaction Stoichiometric relations are ideal sometimes competing reactions occur get less product than predicted % Yield = actual yield/theoretical yield Theoretical yield predicted by stoichiometric relation ...
View Full Document

This note was uploaded on 04/29/2010 for the course CHEM CHEM 6A taught by Professor Czarkowski during the Spring '07 term at UCSD.

Ask a homework question - tutors are online