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Lecture 20

Lecture 20. VSEPR. Valence Bond Theory – Hybridization. One of the issues that arises is that the number of partially filled or empty atomic orbitals did not predict the number of bonds or orientation of bonds

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Lecture 20

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  1. Lecture 20 VSEPR

  2. Valence Bond Theory – Hybridization • One of the issues that arises is that the number of partially filled or empty atomic orbitals did not predict the number of bonds or orientation of bonds • C = 2s22px12py12pz0 would predict two or three bonds that are 90° apart, rather than four bonds that are 109.5° apart • To adjust for these inconsistencies, it was postulated that the valence atomic orbitals could hybridizebefore bonding took place • one hybridization of C is to mix all the 2s and 2p orbitals to get four orbitals that point at the corners of a tetrahedron Tro: Chemistry: A Molecular Approach, 2/e

  3. Molecular Orbitals • When the wave functions combine constructively, the resulting molecular orbital has less energy than the original atomic orbitals – it is called a Bonding Molecular Orbital • s, p • most of the electron density between the nuclei • When the wave functions combine destructively, the resulting molecular orbital has more energy than the original atomic orbitals – it is called an Antibonding Molecular Orbital • s*, p* • most of the electron density outside the nuclei • nodes between nuclei Tro: Chemistry: A Molecular Approach, 2/e

  4. VSEPR Theory • Electron groups around the central atom will be most stable when they are as far apart as possible – we call this valence shell electron pair repulsion theory • because electrons are negatively charged, they should be most stable when they are separated as much as possible • The resulting geometric arrangement will allow us to predict the shapes and bond angles in the molecule Tro: Chemistry: A Molecular Approach, 2/e

  5. Refresher: Lewis Structures of Molecules • Lewis theory allows us to predict the distribution of valence electrons in a molecule • Useful for understanding the bonding in many compounds • Allows us to predict shapes of molecules • Allows us to predict properties of molecules and how they will interact together Tro: Chemistry: A Molecular Approach, 2/e

  6. C B N O F Lewis Structures • Generally try to follow the common bonding patterns • C = 4 bonds & 0 lone pairs, N = 3 bonds & 1 lone pair, O= 2 bonds & 2 lone pairs, H and halogen = 1 bond, Be = 2 bonds & 0 lone pairs, B = 3 bonds & 0 lone pairs • often Lewis structures with line bonds have the lone pairs left off • their presence is assumed from common bonding patterns • Structures that result in bonding patterns different from the common may have formal charges Tro: Chemistry: A Molecular Approach, 2/e

  7. O H O N O Example: Writing Lewis structures of molecules, HNO3 • Write skeletal structure • H always terminal • in oxyacid, H outside attached to O’s • make least electronegative atom central • N is central • not H • Count valence electrons • sum the valence electrons for each atom • add one electron for each − charge • subtract one electron for each + charge N = 5 H = 1 O3 = 36 = 18 Total = 24 e− Tro: Chemistry: A Molecular Approach, 2/e

  8. Example: Writing Lewis structures of molecules, HNO3 • Attach atom together with pairs of electrons, and subtract from the total • don’t forget, a line represents 2 electrons Electrons Start 24 Used 8 Left 16 Tro: Chemistry: A Molecular Approach, 2/e

  9. Example: Writing Lewis structures of molecules, HNO3 • Complete octets, outside-in • H is already complete with 2 • 1 bond and re-count electrons N = 5 H = 1 O3 = 36 = 18 Total = 24 e− Electrons Start 24 Used 8 Left 16 Electrons Start 16 Used 16 Left 0 Tro: Chemistry: A Molecular Approach, 2/e

  10. Example: Writing Lewis structures of molecules, HNO3 • If all octets complete, give extra electrons to the central atom • If central atom does not have octet, bring in electrons from outside atoms to share Tro: Chemistry: A Molecular Approach, 2/e

  11. Practice – Draw Lewis Structures of the Following CO2 SeOF2 NO2− H3PO4 SO32− P2H4 Tro: Chemistry: A Molecular Approach, 2/e

  12. Practice – Lewis Structures CO2 SeOF2 NO2− H3PO4 SO32− P2H4 16 e− 32 e− 26 e− 26 e− 18 e− 14 e− Tro: Chemistry: A Molecular Approach, 2/e

  13. Formal Charge • During bonding, atoms may end with more or fewer electrons than the valence electrons they brought in order to fulfill octets • This results in atoms having a formal charge FC = valence e−− nonbonding e−− ½ bonding e− left O FC = 6 − 4 − ½ (4) = 0 S FC = 6 − 2 − ½ (6) = +1 right O FC = 6 − 6 − ½ (2) = −1 • Sum of all the formal charges in a molecule = 0 • in an ion, total equals the charge Tro: Chemistry: A Molecular Approach, 2/e

  14. Writing Lewis Formulas of Molecules (cont’d) • Assign formal charges to the atoms • fc = valence e−− lone pair e−−½ bonding e− • or follow the common bonding patterns −1 0 +1 −1 0 0 +1 0 Tro: Chemistry: A Molecular Approach, 2/e

  15. Exceptions to the Octet Rule • Expanded octets • elements with empty d orbitals can have more than eight electrons • Odd number electron species e.g., NO • will have one unpaired electron • free-radical • very reactive • Incomplete octets • B, Al Tro: Chemistry: A Molecular Approach, 2/e

  16. Practice – Assign formal charges CO2 SeOF2 NO2− H3PO4 SO32− P2H4 Tro: Chemistry: A Molecular Approach, 2/e

  17. Practice - Assign formal charges CO2 SeOF2 NO2− H3PO4 SO32− P2H4 all 0 P = +1 rest 0 S = +1 Se = +1 all 0 Tro: Chemistry: A Molecular Approach, 2/e

  18. Drawing Resonance Structures • Draw first Lewis structure that maximizes octets • Assign formal charges • Move electron pairs from atoms with (−) formal charge toward atoms with (+) formal charge • If (+) fc atom 2nd row, only move in electrons if you can move out electron pairs from multiple bond • If (+) fc atom 3rd row or below, keep bringing in electron pairs to reduce the formal charge, even if get expanded octet −1 −1 +1 −1 −1 +1 Tro: Chemistry: A Molecular Approach, 2/e

  19. Drawing Resonance Structures −1 • Draw first Lewis structure that maximizes octets • Assign formal charges • Move electron pairs from atoms with (−) formal charge toward atoms with (+) formal charge • If (+) fc atom 2nd row, only move in electrons if you can move out electron pairs from multiple bond • If (+) fc atom 3rd row or below, keep bringing in electron pairs to reduce the formal charge, even if get expanded octet +2 − 1 Tro: Chemistry: A Molecular Approach, 2/e

  20. Evaluating Resonance Structures • Better structures have fewer formal charges • Better structures have smaller formal charges • Better structures have the negative formal charge on the more electronegative atom Tro: Chemistry: A Molecular Approach, 2/e

  21. Practice – Identify Structures with Better or Equal Resonance Forms and Draw Them CO2 SeOF2 NO2− H3PO4 SO32− P2H4 all 0 P = +1 rest 0 S = +1 Se = +1 all 0 Tro: Chemistry: A Molecular Approach, 2/e

  22. Practice – Identify Structures with Better or Equal Resonance Forms and Draw Them CO2 SeOF2 NO2− H3PO4 SO32− P2H4 none −1 +1 none Tro: Chemistry: A Molecular Approach, 2/e

  23. VSEPR Theory • Electron groups around the central atom will be most stable when they are as far apart as possible – we call this valence shell electron pair repulsion theory • because electrons are negatively charged, they should be most stable when they are separated as much as possible • The resulting geometric arrangement will allow us to predict the shapes and bond angles in the molecule Tro: Chemistry: A Molecular Approach, 2/e

  24. there are three electron groups on N one lone pair one single bond one double bond • • • • • • • • O N O • • • • Electron Groups • The Lewis structure predicts the number of valence electron pairs around the central atom(s) • Each lone pair of electrons constitutes one electron group on a central atom • Each bond constitutes one electron group on a central atom • regardless of whether it is single, double, or triple Tro: Chemistry: A Molecular Approach, 2/e

  25. Electron Group Geometry There are five basic arrangements of electron groups around a central atom • based on a maximum of six bonding electron groups • though there may be more than six on very large atoms, it is very rare • http://intro.chem.okstate.edu/1314f97/chapter9/VSEPR.html Tro: Chemistry: A Molecular Approach, 2/e

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