VSEPR

1. 0 VSEPR What shape are your molecules in?

2. Background you need… 0 • Lewis structures • How many bonds do each element make? • What can expand? • Bonding (covalent) • Polarity • Electronegativity and determining bond type • Resonance v. Isomers • Formal charge Let’s review now…..

3. Lewis Structures 0 • Remember that Lewis structures want a full outer shell • Remember that for a given Lewis structure, the number of electrons around the atoms must equal the total number of electrons individually assigned. • Ex: C has 4, H has 1, so CH4 must have 8 total

4. Isomers 0 • Same formula, different arrangement of atoms • Physically break bonds and MOVE atoms

5. Resonance Structures 0 • Have the same alignment of atoms, but different bonding (electrons ONLY are moved, both in bonds and lone pairs)

6. Determining formal charge 0 Formal charge can be determined by: Normal number of electrons in outer shell - [(1/2 the number of bonded electrons) + lone electrons] _____________________________________ = formal charge Example: N in NH4 FC =5- [(1/2 of 8)+ 0]= +1

7. Formal charge and stability 0 • The most “happy” molecules tend to have no formal charges • However, molecules may be “happy” if they have not NET charge on them (if there is 1+ and 1-, so a net of +1 + (-1)=0) • Resonance structures that are the best have a minimal formal charge and a full octet around each atom

8. What is VSEPR? 0 • Valence • Shell • Electron • Pair • Repulsion • Theory

9. Why? 0 • The shape of molecules influences their characteristic: • Physical properties • polarity • boiling point • melting point • state of matter at room temperature • Chemical properties • What it will bond with and energy associated with the bond • Biological properties • Enzyme specificity (induced fit model require close shapes)

10. Valence Bond (VB) Theory 0 • Deals with the overlap of the atomic orbitals (AO) of the participating atoms to form a chemical bond. Due to the overlapping, electrons are found in the bond region. • However, the atomic orbitals for bonding may not be "pure" atomic orbitals directly from the orbitals of the atoms involved. Often, the bonding atomic orbitals have a character of several possible types of orbitals (say s, p, and d). • The methods to get an AO with the proper character for the bonding is called hybridization. The resulting atomic orbitals are called hybridized atomic orbitalsor simply hybrid orbitals.

11. Valence Bond TheoryAnd VSEPR Notation How does Lewis theory explain the bonds in H2 and F2? Sharing of two electrons between the two atoms. Bond Dissociation Energy Bond Length Overlap Of H2 436.4 kJ/mole 74 pm 1s orbitals F2 150.6 kJ/mole 142 pm 2p orbitals Valence bond theory: bonds are formed by sharing of e- from overlapping atomic orbitals.

12. 74 pm Valence bond method • According to this model, the H-H bond forms as a result of the overlap of the 1s orbitals from each atom.

13. Valence bond method • Hybrid orbitals are need to account for the geometry that we observe for many molecules. • Example - Carbon • Outer electron configuration of 2s22px12py1 • We know that carbon will form four equivalent bonds such as in CH4, CH2Cl2 , CCl4. • The electron configuration appears to indicate that only two bonds would form and they would be at right angles -- not tetrahedral angles.

14. 2p 2sp3 energy 2s Unhybridized Hybridized Hybridization • To explain why carbon forms four identical single bonds, we assume the original orbitals will blend together. • This lowers the energy • Lower energy is more favorable

15. The parent geometries: all others come from these 0

16. Some new conventions for shapes….dashes and wedges

17. Steric Number • The number of e- groups, or “things” sprouting off of an atom • These can be either • Bonds • Of any order (1, 2, or 3) Or • Lone pairs of electrons

18. Steric Number Examples • Ex #1: CH4 • There are 4 H’s branching off , so the steric number is 4 • SN=4 • Ex #2: H2O • SN= 4 • Explain why • Ex #3: CO2 • SN= 2 • Explain why

19. General Formulas • All molecules with a shared general formula have a shared geometry • we use them to help note shape • Formulas are typically written with A’s, X’s, and E’s The letters stand for: • A= the central atom • X *= the number of atoms attached to the central atom • E= the number of lone pairs of electrons attached to the central atom • *Some sources use B’s in place of X’s

20. General Formula Examples • Ex #1: CH4 • AX4 • Ex #2: H2O • AX2E2 • Ex #3: CO2 • AX2

21. Linear 0 • AX2

22. Trigonal planar 0 • AX3

23. Tetrahedral 0 • AX4

24. Pyramidal (Trigonal or tetrahedral) 0 • Tetrahedral parent shape • 1 lone pair of electrons • AX3E

25. Bent 0 • Tetrahedral parent shape • 2 lone pair of electrons • AX2E2

26. When determining polarity it is important to look at the dipole moments- do they cancel out? 0

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28. 0

29. Expanded Octets • Atoms with expanded CENTRAL octets are not limited to having only 4 atoms attached. • Shapes that require expanded octets are • Trigonalbipyramidal (TBP) parent • See-saw/ teeter-totter derivative • Octahedral parent • Square pyramidal, square planar, and T- shaped derivatives

30. Trigonal bipyramidal 0 • AX5

31. Seesaw a.k.a. Teeter-totter • Trigonal bipyramidal parent shape • 1 lone pair of electrons • AX4E

32. T-shaped • Trigonal bipyramidal parent shape • 2 lone pair of electrons • AX3E2

33. Linear • Trigonal bipyramidal parent shape • 3 lone pair of electrons • AX2E3

34. Octahedral • AX6

35. Square pyramidal • Octahedral parent shape • 1 lone pair of electrons • AX5E

36. Square planar • Octahedral parent shape • 2 lone pair of electrons • AX4E2

37. T-shaped • Octahedral parent shape • 1 lone pair of electrons • AX3E3

38. See the grid handout for more specifics on each!

39. Summary of shapes

40. ID these VSEPR shapes…

41. Sweet drill and practice web site • Given generic shapes to ID: • http://www.chemistry-drills.com/VSEPR-1.php?q=1 • Given molecules to draw out: • Basic: http://www.chemistry-drills.com/VSEPR-1.php?q=2 • Advanced: http://www.chemistry-drills.com/VSEPR-1.php?q=3