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CHEMICAL BONDING: ORBITALS

CHEMICAL BONDING: ORBITALS. Chapter 9. A review: views a molecule as a collection of atoms bound together by sharing electrons between their atomic orbitals.

chase-williamson
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CHEMICAL BONDING: ORBITALS

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  1. CHEMICAL BONDING: ORBITALS Chapter 9

  2. A review: • views a molecule as a collection of atoms bound together by sharing electrons between their atomic orbitals. • The arrangement of valence electrons is represented by the Lewis structure and the molecular geometry can be predicted from the VSEPR model. LOCALIZED ELECTRON (LE) MODEL

  3. Using the 2p and the 2s orbitals from carbon in methane would result in 2 different types of bonds when they overlap with the 1s from hydrogen. [three 2p/1s bonds and one 2s/2p bond] However, experiments show that methane has FOUR IDENTICAL bonds. • Since the 3p orbitals occupy the x, y and z-axes, you would expect those overlaps of atomic orbitals to be at bond angles of 90°. All 4 angles are 109.5°. PROBLEMS WITH THEORY

  4. an extension of the LE model—it’s all about hybridization. Two atoms form a bond when both of the following conditions occur: • There is orbital overlap between two atoms. • A maximum of two electrons, of opposite spin, can be present in the overlapping orbitals. VALENCE BOND THEORY

  5. Because of orbital overlap, the pairof electrons is found within a region influenced by both nuclei. Both electrons are attracted to both atomic nuclei and this leads to bonding. • As the extent of overlap increases, the strength of the bond increases. The electronic energy drops as the atoms approach each other but increases when they become too close. This means there is an optimum distance, the observed bond distance, at which the total energy is at a minimum. VALENCE BOND THEORY

  6. sigma (σ) bond - overlap of two s orbitals or an s and a p orbital or head-to-head p orbitals. Electron density of a sigma bond is greatest along the axisof the bond. • maximum overlap - forms the strongest possible bond; two atoms are arranged to give the greatest possible orbital overlap. Tricky with p orbitals since they are directional. • hybrid orbitals - a blending of atomic orbitals to create orbitals of intermediate energy. VALENCE BOND THEORY

  7. A carbon atom – no bonding. VALENCE BOND THEORY

  8. sp3 Hybridization • Once the carbon atom begins to bond with an atom, the atomic orbitals HYBRIDIZE which changes their shape considerably. • There is an energy payoff which explains why they behave this way. VALENCE BOND THEORY

  9. The original 2s and three 2p atomic orbitals characteristic of a free carbon atom are combined to form a new set of four sp3 orbitals. VALENCE BOND THEORY

  10. Another way to look at it. VALENCE BOND THEORY

  11. Describe the bonding in the ammonia molecule using the LE model. • Draw the Lewis structure • Determine arrangement of electron pairs. • Determine hybridized orbitals needed in the bonding. PRACTICE ONE

  12. Ammonia molecule with sp3 hybridization.

  13. It is helpful to think of electron pairs as “sites” of electron density that can be occupied by either a lone pair or a shared pair. • If there are 4 “sites” then 4 orbitals need to hybridize so use one s and three p’s to make FOUR equivalent orbitals named sppp or sp3 orbitals [1 s + 3 p = 4 sp3 orbitals]. HELPFUL HINT

  14. The tetrahedral set of four sp3orbitals of the carbon atom are used to share electron pairs with the four 1s orbitals of the hydrogen atoms to form the four equivalent C-H VALENCE BOND THEORY

  15. Lowers the number of hybridizing orbitals since unhybridized orbitals are necessary to form the pi bonds • Pi (π) bonds - come from the sideways overlap of p atomic orbitals; the region above and below the internuclearaxis. NEVER occur without a sigma bond first! • may form only if unhybridizedp orbitals remain on the bonded atoms • occur when spor sp2hybridization is present on central atom NOT sp3hybridization. • Carbon often does this with N, O, P, and S since they have p orbitals MULTIPLE BONDING

  16. sp2 Hybridization • The set of p’s that are unhybridized are vertical. • This is the formation of a sp2 set of orbitals at angles of 120°. Whenever an atom is surrounded by 3 effective e- pairs, a set of sp2 hybrid orbitals is required. This molecule would contain a double bond like ethylene. The molecule reserves a set of p’s to form the π bond. VALENCE BOND THEORY

  17. The hybridization of the s, px, and pyatomic orbitals results in the formation of three sp2 orbitals centered in the xyplane. VALENCE BOND THEORY

  18. The hybridization of the s, px, and py atomic orbitals results in the formation of three sp2 orbitals centered in the xy plane VALENCE BOND THEORY

  19. sigmabond pi bond VALENCE BOND THEORY

  20. carbon #1 carbon #2 Overlapping unhybridized p orbitals that form the pi bond VALENCE BOND THEORY

  21. sp Hybridization • This is the formation of a spset of orbitals at 180°. • This molecule would contain a triple bond like ethyne or the double-double arrangement in carbon dioxide. • The molecule reserves TWO sets of unhybridizedp’s to form the 2 π bonds. VALENCE BOND THEORY

  22. When one s orbital and one p orbital are hybridized, a set of two sp orbitals orientated at 180 degrees results. VALENCE BOND THEORY

  23. At right is a picture of the two unhybridizedp’s on the C atom that are about to make a triple bond. The one labeled at the top is in the plane of this page, the other plain p is in a plane perpendicular to this page. VALENCE BOND THEORY

  24. Look at the CO2 Lewis diagram. The carbon has 2 sites of electron density, each occupied by a double bond, and is therefore sphybridized while the oxygens have 3 sites (2 lone pairs and a double bond). The oxygen’s have sp2hybridization. VALENCE BOND THEORY

  25. Describe the bonding in the N2 molecule. PRACTICE TWO

  26. dsp3 Hybridization • Hybridization when there are five pairs of electrons around the central atom. This requires a trigonalbipyramidal arrangement which in turn requires the dsp3 hybridization. VALENCE BOND THEORY

  27. A set of dsp3 hybrid orbitals on phosphorus atom VALENCE BOND THEORY

  28. The orbitals used to form the bonds in PCl5 VALENCE BOND THEORY

  29. Describe the bonding in I3-, the triiodide ion. PRACTICE THREE

  30. d2sp3 Hybridization • Hybridization when there are six pairs of electrons around the central atom. This requires an octahedral arrangement which in turn requires the d2sp3 hybridization. VALENCE BOND THEORY

  31. An octahedral set of d2sp3 orbitals on sulfur VALENCE BOND THEORY

  32. An octahedral set of d2sp3 orbitals on xenon VALENCE BOND THEORY

  33. Why is the xenon atom in XeF4 hybridized? PRACTICE FOUR

  34. Draw the Lewis structure for benzene: PRACTICE FIVE

  35. The sigma bond formations The pi bond formations BENZENE

  36. For each of the following molecules or ions, predict the hybridization of each atom, and describe the molecular structure. a. CO b. BF4-c. XeF2 PRACTICE SIX

  37. a. CO PRACTICE SIX

  38. b. BF4- PRACTICE SIX

  39. c. XeF2 PRACTICE SIX

  40. Three distinct steps: • Draw the Lewis structure • Determine the arrangements of electron pairs using the VSEPR model. • Specify the hybrid orbitals needed to accommodate the electron pairs. • Remember that the tendency to minimize energy is more important than the maintenance of the characteristics of atoms as they exist in the free state. LE MODEL SUMMARY

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