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Chemistry 1061 Notes ch 8-11 for Exam 3

by: mandygh926

Chemistry 1061 Notes ch 8-11 for Exam 3 chem 1061

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The notes include the main topics for Exam 3: Electron Configuration, Periodic Table Trends, Types of Bonds, Lattice Energy, Lewis Structures, VSEPR, Bond Order, Bond Length, Formal Charge, and Hyb...
Chemical Principles I
Debra Salmon
Chemistry, hybridization
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This 3 page Bundle was uploaded by mandygh926 on Tuesday July 19, 2016. The Bundle belongs to chem 1061 at University of Minnesota taught by Debra Salmon in Fall 2015. Since its upload, it has received 10 views. For similar materials see Chemical Principles I in Chemistry at University of Minnesota.


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Date Created: 07/19/16
Chemistry Notes for Exam 3: Ch. 8­11 on 12/1/15  Corresponding Textbook: Chemistry; The Molecular Nature of Matter and  Change  Includes: Electron Configuration, Periodic Table Trends, Types of Bonds, Lattice Energy, Lewis  Structures, VSEPR, Bond Order, Bond Length, Formal Charge, and Hybridization. Chapter 8: Electron Configuration and Chemical Periodicity 1  Spin quantum number: (m ) s  2  corresponds to direction of electron’s field. For any n value, a lower L value indicates a lower (more stable) sublevel.  Exclusion principle: no two electrons in the same atom can have the same four quantum  #’s  Shielding: reduces the full nuclear charge to an effective nuclear charge making it easier  to remove an electron.   Electron configuration: nƖ  ex. 1s  2s  2p  3s  3p  4s  3d  4p etc…..  Periodic table trends: atomic size←↓(transition elements stay about the same size),  ionization energy↑→, electron affinity↑→, metallic behavior ↓←(memorize periodic  table trends chart). *Anions are larger than parent atoms.  Paramagnetic vs Diamagnetic: Paramagnetic have an unpaired electron which is attracted to an external magnetic field. Diamagnetic have no unpaired electrons, not attracted  (slightly repelled) by magnetic fields. Chapter 9: Models of Chemical Bonding  Ionic bond: between a metal and a non­metal. Large difference in ionization energy.  Forms an ionic solid  Covalent bond: usually between two nonmetals where the electrons are shared. They  have similar ionization energies and forms separate molecules.   Electronegativity: Can be found on the periodic table to determine the bond type between two elements. If the difference in their electronegativities is:  ∆EN > 1.7 Ionic, 0.4 < ∆EN < 1.7 polar covalent, ∆EN < 0.4 nonpolar covalent  Metallic bond: between two metals. De­localization of electrons  Lattice energy: ionization energy plus electron affinity. Affected first by the ion size and  second by the ion charge. atomic¿¿ distancebetweenradi¿ cationchargexanioncharge ∆ H° Lattice ¿  Bond Order: single bonds have B.O. of 1, double bonds have B.O of 2, etc.  Bond Energy: breaking bonds is endothermic and always positive while forming bonds is exothermic and always negative. BREAK = ENDO;    MAKE = EXO ∆ H° =∑ BE −∑ BE rxn reactant bondsbroken product bonds f (BREAK – MAKE)  Bond Length: for groups bond lengths parallel atomic size, for periods atoms closer  together have greater bond lengths Chapter 10: The Shapes of Molecules  Lewis Structure Dot Diagrams: count up valence electrons, distribute evenly, first atom in the formula is usually the central atom, then test for the octet rule. (8 electrons per atom  except H which has 2). Chose the resonance structure with the lowest formal charges.  Exceptions to the octet rule: 1. Electron deficient central atom: Be or B as the central atom 2. Odd­electron molecules: species that contain an odd # of electrons 3. Expanded valence shells: occur with nonmetals from period 3 or higher  Determining the Formal Charge:   1 −¿+ 2 −¿¿ ¿     ¿of unshared valencee of shared valencee −¿−¿ ¿of valencee ¿ ¿of bonding pairs  Calculating Bond Order:  ¿atom pairsheld together  VSEPR: Type of model to describe the bonding of molecules. Written as: AX E  mwnere  A is central atom, X is surrounding atom bond, m is the # of x, E is nonbonding valence  electron group, and n is the # of E.  o (Memorize VSEPR chart of molecular shape names, bond angles, electron group  arrangement, and hybridization for exams. A good chart can be found by googling VSEPR Geometries).   Ideal bonding angles: when all electron pairs are bonded. The real bond angles are less  than ideal when there are lone pairs.  **NOTE: If there is an uneven charge distribution on a molecule, you need to add dipole arrows  on the molecule where the arrow points to the more electronegative atom.  Chapter 11: Theories of Covalent Bonding  Hybridization: the # of hybrid orbitals formed must equal the # of atomic orbitals mixed  Sigma bond: single bonds that form end­to­end overlaps. Allows molecule to spin.  Pi bond: double bonds consist of 1 sigma and 1 pi bond. Form side­to­side overlaps. A  triple bond consists of 2 pi bonds and 1 sigma bond. Restricts movement preventing the  molecule from spinning.  Bond strength: Pi bonds are weaker than sigma bonds. Double bonds are about 2x as  strong as single bonds. −¿∈anti−bonding MO ¿ −¿∈bondingMO−¿of e ¿of e¿  Molecular orbital:   1 MO B.O.= ¿ 2 **NOTE: The Bond order must be greater than zero to form the molecule (B.O. > 0). A higher  bond order tends to have a stronger bond.


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