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What Do Molecules Look Like?

What Do Molecules Look Like?

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What Do Molecules Look Like?

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  1. What Do Molecules Look Like? The Lewis Dot Structure approach provides some insight into molecular structure in terms of bonding, but what about 3D geometry? Recall that we have two types of electron pairs: bonding and lone. Valence-Shell Electron-Pair Repulsion (VSEPR). 3D structure is determined by minimizing repulsion of electron pairs.

  2. Electron pairs (both bonding and lone) are distributed around a central atom such that electron-electron repulsions are minimized.

  3. Electron pairs (both bonding and lone) are distributed around a central atom such that electron-electron repulsions are minimized. 2 electron pairs 3 electron pairs 4 electron pairs Period 1, 2 5 electron pairs 6 electron pairs Period 3 & beyond

  4. H H C H H Arranging Electron Pairs • Must consider both bonding and lone pairs when minimizing repulsion. • Example: CH4 (bonding pairs only) Lewis Structure VSEPR Structure

  5. H H N H Arranging Electron Pairs (cont.) Example: NH3 (both bonding and lone pairs). Lewis Structure VSEPR Structure Note:“electron pair geometry” vs. “molecular shape”

  6. VSEPR Structure Guidelines • The previous examples illustrate the strategy for applying VSEPR to predict molecular structure: • Construct the Lewis Dot Structure • Arrange bonding/lone electron pairs in space such that repulsions are minimized (electron pair geometry). • Name the molecular shape from the position of the atoms. VSEPR Shorthand: 1. Refer to central atom as “A” 2. Attached atoms are referred to as “X” 3. Lone pair are referred to as “E” Examples: CH4: AX4 NH3: AX3E H2O: AX2E2 BF3: AX3

  7. Be F F Be F F VSEPR: 2 electron pairs Experiments show that molecules with multiple bonds can also be linear. Linear (AX2): angle between bonds is 180° Example: BeF2 Multiple bonds are treated as a single effective electron group. F F Be 180° More than one central atom? Determine shape around each.

  8. VSEPR: 3 electron pairs Trigonal Planar (AX3): angle between bonds is 120° Multiple bond is treated as a single effective electron group. Example: BF3 F F 120° B F F F B F

  9. H H C H H VSEPR: 4 electron pairs (cont.) Tetrahedral (AX4): angle between bonds is ~109.5° Example: CH4 109.5° tetrahedral e- pair geometry AND tetrahedral molecular shape

  10. Bonding vs. Lone pairs Bond angle in a tetrahedral arrangement of electron pairs may vary from 109.5° due to size differences between bonding and lone pair electron densities. bonding pair is constrained by two nuclear potentials; more localized in space. lone pair is constrained by only one nuclear potential; less localized (needs more room).

  11. H H N H VSEPR: 4 electron pairs Trigonal pyramidal (AX3E): Bond angles are <109.5°, and structure is nonplanar due to repulsion of lone pair. Example: NH3 107° tetrahedral e- pair geometry; trigonal pyramidal molecular shape

  12. VSEPR: 4 electron pairs (cont.) Classic example of tetrahedral angle shift from 109.5° is water (AX2E2): 104.5o “bent” tetrahedral e- pair geometry; bent molecular shape

  13. VSEPR: 4 electron pairs (cont.) Comparison of CH4 (AX4), NH3 (AX3E), and H2O (AX2E2):

  14. AX2E AX3E AX2E2 1. Refer to central atom as “A” 2. Attached atoms are referred to as “X” 3. Lone pair are referred to as “E”

  15. H H H H Molecular vs. Electron-Pair Geometry F N O C Central Atom Compound Electron-Pair Geometry Molecular Shape Carbon, C CH4 tetrahedral tetrahedral Nitrogen, N NH3 tetrahedral trigonal pyramidal Oxygen, O H2O tetrahedral bent Fluorine, F HF tetrahedral linear

  16. What is the electron-pair geometry and the molecular shape for HCFS? • trigonal planar, bent • trigonal planar, trigonal planar • tetrahedral, trigonal planar • tetrahedral, tetrahedral

  17. VSEPR: Beyond the Octet Systems with expanded valence shells will have five or six electron pairs around a central atom. F Cl F Cl F Cl S P Cl F F Cl F 90° 90° F F F 90° S 120° F F F

  18. F F F F F F F F F F F F VSEPR: 5 electron pairs • • Consider the structure of SF4 (34 e-, AX4E) • What is the optimum arrangement of electron pairs around S? ?? S S S Compare e– pair angles lone-pair / bond-pair: twoat 90o, twoat 120o threeat 90o threeat 90o, three at 120o fourat 90o, one at 120o bond-pair / bond-pair: Repulsive forces (strongest to weakest): lone-pair/lone-pair > lone-pair/bond-pair > bond-pair/bond-pair

  19. VSEPR: 5 electron pairs The optimum structure maximizes the angular separation of the lone pairs. I3- (AX2E3):

  20. AX4E AX3E2 AX2E3 5-electron-pair geometries our previous example

  21. VSEPR: 6 electron pairs Which of these is the more likely structure? See-saw Square Planar

  22. AX5E AX4E2 6-electron-pair geometries our previous example

  23. Molecular Dipole Moments We can use VSEPR to determine the polarity of a whole molecule. • Draw Lewis structures to determine 3D arrangement of atoms. 2. If one “side” of the molecule has more EN atoms than the other, the molecule has a net dipole. Shortcut: completely symmetric molecules will not have a dipole regardless of the polarity of the bonds.

  24. Molecular Dipoles The C=O bonds have dipoles of equal magnitude but opposite direction, so there is no net dipole moment. The O-H bonds have dipoles of equal magnitude that do not cancel each other, so water has a net dipole moment.

  25. Molecular Dipoles (cont.) symmetric symmetric asymmetric

  26. Molecular Dipole Example • Write the Lewis dot and VESPR structures for CF2Cl2. Does it have a dipole moment? F 32 e- Cl F Cl Tetrahedral

  27. Advanced VSEPR Application Molecules with more than one central atom… methanol (CH3OH) H C O H H tetrahedral e- pairs tetrahedral shape tetrahedral e- pairs bent shape H

  28. # e- pairs e- Geom. Molec. Geom. The VSEPR Table

  29. # e- pairs e- Geom. Molec. Geom. The VSEPR Table

  30. What is the expected shape of ICl2+? 20 e- AX2E2 A. linear C. tetrahedral D. square planar B. bent

  31. Valence Bond Theory Basic Principle of Localized Electron Model: A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies the region between the two nuclei. Rule 1: Maximum overlap. The bond strength depends on the attraction of nuclei to the shared electrons, so: The greater the orbital overlap, the stronger the bond.

  32. Valence Bond Theory Basic Principle of Localized Electron Model: A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies the region between the two nuclei. Rule 2: Spins pair. The two electrons in the overlap region occupy the same space and therefore must have opposite spins. There may be no more than 2 electrons in a molecular orbital.

  33. Valence Bond Theory • Basic Principle of Localized Electron Model: • A covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies the region between the two nuclei. • Rule 3: Hybridization.To explain experimental observations, Pauling proposed that the valence atomic orbitals in a molecule are different from those in the isolated atoms. We call this concept • Hybridization

  34. What is hybridization? • Atoms adjust to meet the “needs” of the molecule. • In a molecule, electrons rearrange in an attempt to give each atom a noble gas configuration and to minimize electron repulsion. • Atoms in a molecule adjust their orbitals through hybridization in order for the molecule to have a structure with minimum energy. • The source of the valence electrons is not as important as where they are needed in the molecule to achieve a maximum stability.

  35. Example: Methane • 4 equivalent C-H covalent bonds • VSEPR predicts a tetrahedral geometry

  36. How do we explain formation of 4 equivalent C-H bonds? The Valence Orbitals of a Carbon Atom Carbon: 2s22p2

  37. Hybridization: Mixing of Atomic Orbitals to form New Orbitals for Bonding + + – – + – + – + – + – +

  38. Other Representations of Hybridization: y1 = 1/2[(2s) + (2px) + (2py) + (2pz)] y2 = 1/2[(2s) + (2px) - (2py) - (2pz)] y3 = 1/2[(2s) - (2px) + (2py) - (2pz)] y4 = 1/2[(2s) - (2px) - (2py) + (2pz)]

  39. Hybridization is related to the number of valence electron pairs determined from VSEPR: Methane (CH4) VSEPR: AB4 tetrahedral  sp3 hybridized 109.47 º Electron pair geometry determines hybridization, not vice versa!!

  40. Hybridization is related to the number of valence electron pairs determined from VSEPR: Ammonia (NH3) VSEPR: AB3E tetrahedral  sp3 hybridized N H H H 108.1 º

  41. Hybridization is related to the number of valence electron pairs determined from VSEPR: Water (H2O) VSEPR: AB2E2 tetrahedral  sp3 hybridized 105.6 º

  42. sbonding and pbonding • Two modes of bonding are important for • 1st and 2nd row elements: s bonding and p bonding • These two differ in their relationship to the internuclear axis: • s bonds have electron density ALONG the axis • p bonds have electron density ABOVE AND BELOW the axis

  43. Problem: Describe the hybridization and bonding of the carbon orbitals in ethylene (C2H4) VSEPR: AB3 trigonal planar  sp2 hybridized orbitals for s bonding sp2 hybridized orbitals used for s bonding remaining p orbital used for p bonding

  44. Bonding in ethylene (C2H4)

  45. Problem: Describe the hybridization and bonding of the carbon orbitals in Carbon Dioxide (CO2) • VSEPR: AB2 • linear • sp hybridized orbitals for s bonding

  46. Bonding in Carbon Dioxide (CO2)

  47. H H C2 C1 N : N CH3 C H p p sp sp sp sp3 Atoms of the same kind can have different hybridizations Acetonitrile (important solvent and industrial chemical) Bonds s C2: AB4 C1: AB2 2s2 2px2py s p p N: ABE p sp p sp 2s2 2px2py2pz lone pair

  48. What have we learned so far? • Molecular orbitals are combinations of atomic orbitals • Atomic orbitals are “hybridized” to satisfy bonding in molecules • Hybridizationfollows simple rules that can be deduced from the number of chemical bonds in the molecule and the VSEPR model for electron pair geometry

  49. Hybridization • sp3 Hybridization (CH4) • This is the sum of one s and three p orbitals on the carbon atom • We use just the valence orbitals to make bonds • sp3 hybridization gives rise to the tetrahedral nature of the carbon atom

  50. Hybridization • sp2 Hybridization (H2C=CH2) • This is the sum of one s and two p orbitals on the carbon atom • Leaves one p orbital uninvolved – this is free to form a p bond (the second bond in a double bond)