1 / 26

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

beaumont-dilan
Télécharger la présentation

AP Chemistry

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  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

More Related