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IB Chem HL2 Bonding (Topic 4and 14)

IB Chem HL2 Bonding (Topic 4and 14). Review (Topic 4). Predict whether or not a molecule is polar from its molecular shape and bond polarities

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IB Chem HL2 Bonding (Topic 4and 14)

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  1. IB Chem HL2Bonding (Topic 4and 14)

  2. Review (Topic 4) • Predict whether or not a molecule is polar from its molecular shape and bond polarities • Predict whether a compound of two elements would be ionic from the position of the elements in the periodic table or from their electronegativity values • State the formula of common polyatomic ions formed by non-metals in periods 2 and 3 • Describe the lattice structure of ionic compounds (5.2.12.B.1) • Describe and compare the structure and bonding in the three allotropes of carbon • Describe the structure and bonding in silicon and silicon dioxide

  3. Review (Topic 4) • Describe the types of intermolecular forces (attraction between molecules that have temporary dipoles, permanent dipoles, or hydrogen bonds) and explain how they arise from the structural features of the molecules • Describe and explain how intermolecular forces affect the boiling points of substances (5.2.12.C.2) • Describe the metallic bond as the electrostatic attraction between a lattice of positive ions and delocalized electrons • Explain the electrical conductivity and malleability of metals (5.2.12.C.2) • Compare and explain the properties of substance resulting from different types of bonding (5.2.12.C.2)

  4. IB Standards • 4.2.7 Predict the shape and bond angles for species with 4, 3, and 2 negative charge centers using VSEPR. http://www.youtube.com/watch?v=qvPKWcXdo7Y • 14.1.1 Predict the shape and bond angles for species with 5 and 6 negative charge centres using VSEPR. (5.1.A1, 5.2.A2, 5.3.C1) http://www.youtube.com/watch?v=keHS-CASZfc&feature=em-subs_digest-vrecs http://www.youtube.com/watch?v=Ip8v87vxSok

  5. # of atoms bonded tocentral atom # lone pairs on central atom Arrangement ofelectron pairs Molecular Geometry Class linear linear B B Valence shell electron pair repulsion (VSEPR) model: Predict the geometry of the molecule from the electrostatic repulsions between the electron (bonding and nonbonding) pairs. AB2 2 0

  6. Cl 0 lone pairs on central atom Be Cl 2 atoms bonded to central atom

  7. BF3

  8. Methane

  9. Phosphorus Pentachloride

  10. Sulfur Hexafluoride

  11. lone-pair vs. lone-pair repulsion lone-pair vs. bonding- pair repulsion bonding-pair vs. bonding- pair repulsion > > Take note of this comparison:

  12. # of atoms bonded tocentral atom # lone pairs on central atom Arrangement ofelectron pairs Molecular Geometry bent Class trigonal planar VSEPR trigonal planar trigonal planar AB3 3 0 AB2E 2 1

  13. # of atoms bonded tocentral atom # lone pairs on central atom trigonal pyramidal Arrangement ofelectron pairs Molecular Geometry Class tetrahedral VSEPR tetrahedral tetrahedral AB4 4 0 AB3E 3 1

  14. # of atoms bonded tocentral atom # lone pairs on central atom trigonal pyramidal Arrangement ofelectron pairs Molecular Geometry AB3E 3 1 tetrahedral Class bent tetrahedral VSEPR tetrahedral tetrahedral AB4 4 0 AB2E2 2 2

  15. # of atoms bonded tocentral atom # lone pairs on central atom trigonal bipyramidal distorted tetrahedron Arrangement ofelectron pairs Molecular Geometry Class VSEPR trigonal bipyramidal trigonal bipyramidal AB5 5 0 AB4E 4 1

  16. # of atoms bonded tocentral atom # lone pairs on central atom trigonal bipyramidal distorted tetrahedron/seesaw Arrangement ofelectron pairs Molecular Geometry AB4E 4 1 Class trigonal bipyramidal T-shaped VSEPR trigonal bipyramidal trigonal bipyramidal AB5 5 0 AB3E2 3 2

  17. Strategy for Predicting Molecular GeometrySee page 130 of textbook • Draw Lewis structure for molecule. • Count number of lone pairs on the central atom and number of atoms bonded to the central atom. • Use VSEPR to predict the geometry of the molecule.

  18. Practice Exercise #1 Use the VSEPR model, predict the geometry of the following molecules and ions and briefly explain your answer: AsH3 OF2 C2H4

  19. a) The sequence of steps in determining molecular geometry is:

  20. Answers to Practice 3(molecular/ion geometry)

  21. Hybridization IB Standard • 14.2.2 Explain hybridization in terms of the mixing of atomic orbitals to form new orbitals for bonding. Include sp, sp2 and sp3.(5.1.A1, 5.2.A1, 5.3.C1) • 14.2.3 Identify and explain the relationships between Lewis structures, molecular shapes and types of hybridization (sp, sp2 and sp3).(5.1.A1, 5.2.A1,5.2.B3, 5.3.C1,5.6.A7)

  22. Hybridization – mixing of two or more atomic orbitals to form a new set of hybrid orbitals • Mix at least 2 nonequivalent atomic orbitals (e.g.s and p). Hybrid orbitals have very different shape from original atomic orbitals. • Number of hybrid orbitals is equal to number of pure atomic orbitals used in the hybridization process. • Covalent bonds are formed by: • Overlap of hybrid orbitals with atomic orbitals • Overlap of hybrid orbitals with other hybrid orbitals

  23. Formation of sp3 Hybrid Orbitals

  24. Formation of Covalent Bonds in CH4

  25. Predict correct bond angle sp3-Hybridized N Atom in NH3

  26. Formation of sp Hybrid Orbitals

  27. Formation of sp2 Hybrid Orbitals

  28. sp2 Hybridization of Carbon

  29. Unhybridized 2pz orbital (gray), which is perpendicular to the plane of the hybrid (green) orbitals.

  30. Sigma bond (s) – electron density between the 2 atoms Pi bond (p) – electron density above and below plane of nuclei of the bonding atoms Bonding in Ethylene, C2H4

  31. Sigma (s) and Pi Bonds (p) 1 sigma bond Single bond 1 sigma bond and 1 pi bond Double bond Triple bond 1 sigma bond and 2 pi bonds

  32. How do I predict the hybridization of the central atom? • Draw the Lewis structure of the molecule. • Count the number of lone pairs AND the number of atoms bonded to the central atom # of Lone Pairs + # of Bonded Atoms Hybridization Examples 2 sp BeCl2 3 sp2 BF3 4 sp3 CH4, NH3, H2O 5 sp3d PCl5 6 sp3d2 SF6

  33. Practice 2: Do Now Indicate the hybridization of carbon in each of the following compounds: • CH3COOH • CO2

  34. Topic: Delocalization of ElectronsIB Standard • 14.3.1 Describe the delocalization of pi electrons and explain how this can account for the structures of some species.

  35. Delocalized electrons are not confined between two adjacent bonding atoms, but actually extend over three or more atoms. Example: Benzene, C6H6 Delocalized p orbitals

  36. - - + + O O O O O O A resonance structure is one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure. Example: ozone, O3

  37. Delocalized electrons give special properties to the structures in which they are found. • Intermediate bond lengths and strengths. • Greater stability because delocalization spreads the pi electrons which minimizes repulsion between them.

  38. Practice 3: Do Now Draw the resonance structures of • Nitrite, NO2- • Carbonate, CO32-

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