1 / 61

Theories of Chemical Bonding

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 F inal Exam (review session weekend before) Monday, 12-December @ 8AM in IRC 3 (here) ~50 points just on Chapter 9

ocean
Download Presentation

Theories of Chemical Bonding

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript


  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

  11. 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!

  12. 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

  13. 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

  14. 5 electron pairs (sp3d)

  15. 6 electron pairs (sp3d2)

  16. How is this systematic ?

  17. 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

  18. 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…..

  19. 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.

  20. 5 total pairs (4 structural, one lone)

  21. 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

  22. 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)

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

  24. 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)

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

  26. 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

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

  28. 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

  29. 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.

  30. 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

  31. 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)

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

  33. 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.

  34. 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)

More Related