Theories of Chemical Bonding

1. Theories of Chemical Bonding Exam #4 (Chapter 8) on 7-December One 3” x 5” notecard Chapter 8 OWL Deadline Tuesday Night Chapter 9 OWL Deadline @ Final Exam Final Exam (review session weekend before) Monday, 12-December @ 8AM in IRC 3 (here) ~50 points just on Chapter 9 (about the same as the other chapters) ~100 points cumulative (Chapters 2-8) One 81/2” x 11” note sheet allowed Final Lab Experiment This Week (let’s look @ WIKI Page) Laboratory Checkout Next Week!

2. What’s Gotten Us Here…. • Electron Configurations • Energy Level Diagrams • Valence Electrons • Lewis Structures • Identify Central Atom, Bonding and Lone Pairs • Formal Charge • Periodic Trends • Size, Electronegativity • Bond Properties • Bond Energy, Length, Order, Polarity

3. Electron Pair Geometry • Clues to the shape of the molecule or ion • Molecular Geometry • Definitive shape of the molecule or ion • Bond Angles… • However, other than identifying ionic and covalent bonds, we haven’t really described how atoms bond together! • Electrons in a covalent bond are shared, but how?

4. Let’s Look at Boron Trifluoride • B has 3 valence e- • Fluorine has 7 valence e- • We know that B follows an exception to the octet rule • Let’s look at the Lewis Structure

5. How do these 3 bonds form? • We have a single s and three p orbitals on the boron atom • We have a single s and three p orbitals on the fluorine atom • How do they overlap to share electrons? • Hybridization!

6. Valence Bond and Molecular Orbital Theories • Share some principles • Bonds are formed when electrons are shared between atoms • The sharing of electrons, and their attraction to the two nuclei of the atoms that are bonded lowers the total energy in the molecule or ion • Two types of bonds can form • sigma (σ) – bonds which lie along the axis between to atoms in a molecule or ion • Atoms can rotate around sigma bonds • pi (π) – bonds which lie in regions outside of the axis between two atoms (but parallel to the axis) • Atoms are fixed and can’t rotate • Multiple bonding

7. Valence Bond Theory • Takes into account • Lewis structure (# of bonds) • VSEPR (overall shape of the molecule or ion) • The available orbitals that can be used for bonding • Quantum Mechanics • Spectroscopy results (like the colorimetry you used in lab) • Thermodynamic data • Helps decide what types of hybrids will form

8. Hybridization (very systematic) • To obtain the bonding description about any atom in a molecule: • 1. Write the Lewis electron-dot formula. • 2. Use VSEPR to determine the electron arrangement about the atom. • 3. From the arrangement, deduce the hybrid orbitals. • 4. Assign the valence electrons to the hybrid orbitals one at a time, pairing only when necessary. • 5. Form bonds by overlapping singly occupied hybrid orbitals with singly occupied orbitals of another atom.

9. Some Tenets of Hybridization • The total number of orbitals you started with (all valence) MUST equal the number of orbitals that you end up with • The number of hybrid orbitals must equal the number of initial orbitals you started with (again all valence) • Two electrons per orbital • The number of sigma bonds equals the number of single bonds • Multiple bonds are sigma + pi bonds • s, p, d orbitals are what we are working with • Valence ONLY!!!

10. Let’s Start with Methane (CH4) • Write the valence electron configurations • Carbon has 4 valence electrons • 4 hydrogen atoms each have 1 valence e- • Draw the Lewis structure • Evaluate electron and molecular structure • Shape helps us understand hybridization • Determine what orbitals will be combined (hybridized) to create the necessary bonds • Describe the hybridization

14. Lone Pairs ? • Lone pairs of electrons can be placed in hybrid orbitals too. • Ammonia • Write electron configurations (valence) • Lewis Structure • Predict electron and molecular geometry • Determine what orbitals might combine and hybridize to describe the bonding • Consider the best way to locate lone pairs (VSEPR) • Describe the hybridization • Let’s do this on the board!

16. OK, Let’s Get Back to BF3 • Lewis Structure • Electron Pair and Molecular Structure (geometry) • Predict Hybridization • Remember-we are normally talking about hybridization about central atoms at this point. • Describe Hybridization

19. 3rd Period Elements (d orbitals) • Hybrids can be made from combinations of s, p and d orbitals. • Only available to elements in the 3rd period or higher (n = 2 only has s and p orbitals available) • Same rules apply • Total number of orbitals are constant • Number of hybrid orbitals equals the number of orbitals combined to make them • Some orbitals may be left unhybridized

20. 5 electron pairs (sp3d)

22. 6 electron pairs (sp3d2)

24. How is this systematic ?

25. Hybrid orbitals are named by using the atomic orbitals that combined: • one s orbital + one p orbital gives two sporbitals • one s orbital + two p orbitals gives three sp2 orbitals • one s orbital + three p orbitals gives four sp3 orbitals • one s orbital + three p orbitals + one d orbital gives five sp3d orbitals • one s orbital + three p orbitals + two d orbitals gives six sp3d2 orbitals

27. What about lone pairs (on the central atom) • We’ve seen NH3, where they go in one of the hybrid orbitals • Usually this is the case, because the hybrid orbital represents the greatest separation of electron pairs (bonding or lone) in 3D space. • They can go in other hybrid orbitals in other cases, for example…..

28. 4 total pairs (2 structural, 2 lone) • SeH2 • Valence Electron Configuration • Lewis Dot Structure • Predict Electron Pair and Molecular Geometry • Different-two lone pairs. • Predict Hybridization • Draw with electron pairs in the appropriate orbitals and describe the hybridization • What happens to the tetrahedral bond angles? • Let’s do this on the board.

29. 5 total pairs (4 structural, one lone)

30. pi (π) bonds • sigma (σ) bonds lie along the internuclear axis between atoms • pi (π) bonds are parallel to the internuclear axis, but lie separated from it in space • Created from p orbitals not used in hybridization • Responsible for multiple bonds • We will concern ourselves with pi bonds in carbon containing (organic) molecules only

31. Double Bonds (one σ, one π) • Ethene (ethylene): • C2H4 • 3 x sp2 hybrid orbitals on each carbon • 2 C-H sigma bonds on each carbon • 1 C-C sigma bond • 1 C-C pi bond • Double bond C-C (one sigma plus one pi)

33. Note that the geometry of the sp2 hybrids (trigonalplanar) controls geometry about each C atom.

35. Triple Bonds (one σ, two π) • Acetylene • C2H2 • 2 sp hybrid orbitals on each carbon • 1 C-H sigma bonds on each carbon • 1 C-C sigma bond • 2 C-C pi bonds • Triple bond C-C (one sigma plus two pi)

37. Note that the geometry of the sphybrids (linear) controls the geometry about each C atom.

38. What’s different about multiple bonds (sigma + pi) vs. single (sigma) bonds? • Location of the bonds • Along the internuclear axis versus parallel to it • Single bonds (sigma only) can rotate • Multiple bonds (sigma + pi) are rigid • Leads to structural isomers • cis vs. trans forms • More than one multiple (sigma + pi) in a molecule can lead to different resonance forms

39. Do we have cisvesus trans 1, 2 dichloroethane? Let’s draw on the board.

40. Benzene (C6H6) / Resonance • 6 carbons, each with 3 sp2 hybrid orbitals • Each results in 3 sigma bonds • 1 sigma bond to H, 2 sigma bonds to C • One pi bond on each carbon atom

43. sp2 hybridization still controls geometry on each carbon atom • pi bonds make the molecule “flat” • Two resonance forms describe possible location of pi bonds • How does this influence bond order, bond length, bond energy? • VERY stable-electrons are distributed throughout the molecule.

44. Molecular Orbital (MO) Theory • Considers orbitals to be more spread out (delocalized) than Valence Bond Theory • Spectroscopy results are more supportive of MO theory • Is better at describing why certain molecules do not form (why bonding does not happen in some cases) • Can be used in a complimentary fashion with Valence Bond Theory to describe what is happening in molecules. • We still have sigma (σ) and pi (π) bonding. • We add non-bonding and antibondingvariables

45. How to conceptualize MO theory? • Instead of thinking of how hybrids form on a central atom…. • Think about how orbitals on two separate atoms in a bond can come together and form bonds in a molecule • Combining orbitals (MO) instead of overlapping orbitals (VBT)

46. What is some of this terminology? What do the symbols mean?

47. Final Exam & Reminders • OWL Deadline (Chapter 9) is the start of the final exam • Review Sessions (IRC 3-here) • Sunday 10-11:30 AM • Sunday 12:30-2 PM • Final Exam (150 points) • 45 points Chapter 9 (short answer) • 128 points Chapters 2-8 (multiple choice) • Pen/Pencils, Calculator, Ruler, 81/2” x 11” note sheet • I will provide thermodynamic and other constants • Planck’s constant, speed of light, bond enthalpies, etc.

50. Some important principles of MO Theory • The number of molecular orbitals formed equals the number of atomic orbitals that have combined! • Electrons are spread out over the entire molecule • As opposed to being more localized in specific bonds as we see with VBT • The total energy of the molecular orbitals equals the sum of the energy of the atomic orbitals • It is just that some orbitals don’t contain electrons (antibonding)