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AP Chemistry

AP Chemistry. Chapters 9. Vocab (Ch 9). VSEPR- Valence Shell e- Pair Repulsion bonding pair non bonding pair – lone pair of electrons electron domain – regions around a central atom where e- are likely to be found. molecular geometry- the arrangement of atoms in space

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AP Chemistry

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  1. AP Chemistry Chapters 9

  2. Vocab (Ch 9) • VSEPR- Valence Shell e- Pair Repulsion • bonding pair • non bonding pair – lone pair of electrons • electron domain – regions around a central atom where e- are likely to be found. • molecular geometry- the arrangement of atoms in space • electron domain geometry- the arrangement of e- domains about the central atom of a molecule • The Molecular geometry is a derivation of the Electron-Domain geometry • See Table 9.2 (page 309)

  3. Electron Domains • The e- in a multiple bond constitute a single e- domain. • # of e- domains = (# of atoms bonded to the central atom) + (# of non bonding pairs on the central atom) Page 306

  4. Molecular Shapes Website • See VSEPR table handout for molecular shapes • http://www.molecules.org/VSEPR_table.html • See B& L page 307-309

  5. Effect of Non bonding e- and multiple bonds on Bond Angles • e- domains for non-bonding e- pairs exert greater repulsive forces on adjacent e- domains and thus tend to compress the bond angles • e- domains for multiple bonds exert a greater repulsive force on adjacent e- domains than do single bonds. Page 310

  6. Molecules with Expanded Valence Shells • These shapes generally contain axial and equatorial positions • See B&L pg. 312 • Variations of the trigonal bipyramidal shape show lone electron pairs in the equatorial position • Variations of the octahedral shape show lone electron pairs in the axial positions • Page 311

  7. Molecules With More Than One Central Atom • You can use the VSEPR theory for molecules with more than one central atom, such as, CH3NH2. • Pages 313-314

  8. Bond Polarity • Bond Polarity is a measure of how equally the e- in a bond are shared between the 2 atoms of the bond. • Polarity is used when talking about covalently bonded molecules. • If the molecule has only 2 different atoms, such as, HF or CCl4 you can calculate the electronegativity difference and determine the type of covalent bond (polar or non-polar).

  9. Polarity and Bond Type

  10. Dipole Moment • Dipole Moment – the measure of the amount of charge separation in the molecule. • For a molecule with more than 2 atoms, the dipole moment depends on both the polarities of the individual bonds and the geometry of the molecule. • The overall dipole moment of a polyatomic molecule is the sum of its bond dipoles. • See B&L page 315 figure 9.9

  11. Dipole Moment • For each bond in the molecule, consider the bond dipole (the dipole moment due only to the 2 atoms in that bond) • The dipole “arrow” should point toward the more electronegative atom in the bond • The overall dipole moment of a polyatomic molecule is the sum of its bond dipoles. (Consider the magnitude and direction of the bond dipoles)

  12. Different Theories • VSEPR Theory (using Lewis Dot Structures) • Valence Bond Theory (using hybridization) • Molecular Orbital Theory (shows allowed states for e- in molecules) • Go to the following web-site for a compare and contrasting of the 3 different theories • http://www.chem.ufl.edu/~chm2040/Notes/Chapter_12/theory.html

  13. sp Hybrid Orbitals • See section 9.5 • pages 318-320

  14. sp2 and sp3 Hybrid Orbitals • See section 9.5 • pages 320-322

  15. d Orbital Hybridization • See section 9.5 • pages 322

  16. Multiple Bonds and Hybridization • See section 9.6 • pages 324-326

  17. Delocalized π Bonding • See section 9.6 • pages 327-330

  18. Sigma Bonds ( σ ) • Sigma bonds occur when the e- density is concentrated between the 2 nuclei. • These are single covalent bonds. • Sigma bonds can form from the overlap of an s orbital with another s orbital, an s orbital with a p orbital, or a p orbital with a p orbital.

  19. Pi Bonds ( π ) • Overlap of 2 “p” orbitals oriented perpendicularly to the inter-nuclear axis • This overlap results in the sharing of electrons. • The shared electron pair of a pi bond occupies the space above and below the line that represents where the two atoms are joined together.

  20. Hybridization • An atom in a molecule may adopt a different set of atomic orbitals (called hybrid orbitals) than those it has in the free state. • See B&L pages 319-322 for explanation and diagrams of electron promotion

  21. Multiple Bond • A multiple bond consists of one sigma bond and at least one pi bond. • A double bond consists of one sigma bond and one pi bond. • A triple bond consists of one sigma bond and two pi bonds. • A pi bond always accompanies a sigma bond when forming double and triple bonds.

  22. Sigma and Pi • Most of the bonding that you have seen so far has bonding e- that are localized. • σ and π bonds are associated with the 2 atoms that form the bond (and NO other atoms) • Delocalized bonding can occur in molecules that have π bonds and more than one resonance structure.

  23. Molecular Orbital Diagram • Energy Level Diagram (Molecular Orbital Diagram) • The H2 molecule is the easiest molecule to plot on the molecular orbital diagram • Whenever 2 atomic orbitals overlap, 2 molecular orbitals form (one is a bonding orbital and one is an anti-bonding orbital). This is not on the AP exam

  24. Paramagnetism • Molecules with one or more unpaired electrons are attracted into a magnetic field • The more unpaired electrons in species, the stronger the force of attraction • This behavior is called paramagnetism

  25. Diamagnetism • Substances with no unpaired electrons are weakly repelled from a magnetic field • This property is called diamagnetism • Diamagnetism is much weaker than paramagnetism

  26. Problems to Try • Ch 9 • # 5-13, 15, 23, 27, 31, 32, 34, 40, 43, 44 (a and c), 63 • AP Exam Problems to Try 1999 # 8 2000 # 7 (last section) 2002 # 6 2003 # 8 2004 # 7 & # 8

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