2. STEREOCHEMISTRY CHEMISTRY 1

3. Stereochemistry The study of shapes of molecules is called stereochemistry. It is a very important concept in biochemistry because often the body is able to absorb one particular shape of molecule but not another.

4. VSEPR Theory To determine the shape of the molecule, you must determine how many areas of electrons are present. Each of these areas will repel each other, thus forming the shape of the molecule.

5. VSEPR Theory Shapes can be estimated by several different methods, but the most commonly accepted is VSEPR theory. This stands for Valence Shell Electron Pair Repulsions.

6. VSEPR Theory The VSEPR Theory is based on the premise that the valence electrons of each peripheral atom will repel each other strongly and therefore cause the peripheral atoms to move as far from each other as possible.

7. VSEPR Theory An area of electrons is either 1.Single bond 2.Double bond 3.Triple bond 4.Pair of unshared electrons

8. # bonding sites on CA ExDrawingShapeBond Angle 1H2H2 Linear180 2BeCl 2 Linear180 Linked Example: Linear Pairs

9. # bonding sites on CA ExDrawingShapeBond Angle 3BF 3 Trigonal Planar 120 4CCl 4 H 2 O Tetrahedral 109.5 Linked example: Trigonal Planar

10. # bonding sites on CA ExDrawingShapeBond Angle 5PCl 5 SF 4 Trigonal Bipyramidal 120 90 180 6SF 6 BrF 5 Octahedral 90 180 Linked Example: Trigonal Bipyramidal

11. VSEPR Theory To determine the shape using VSEPR, we must first determine the Lewis Dot Structure of the molecule. Let’s review………

12. Lewis Dot Structures The Lewis Dot Structures for atoms show ALL valence electrons ONLY. Valence electrons are considered to be ALL outer- shell electrons; be careful that you show whether they are paired or not.

13. Lewis Dot Structures Drawing the Lewis dot structures for ions follows the same pattern. Remember that metals will lose all of their valence electrons to form a positively-charged particles called a cation.

14. Lewis Dot Structures Metals lose electrons from “outside” to “inside”.

15. Lewis Dot Structures Non-metals will gain enough electrons to completely fill their “p” sublevel and form a negatively-charged particle called an anion. Monatomic ions will ALWAYS have four pairs of electrons showing in the dot structure.

16. Lewis Dot Structures “Gained” electrons should be indicated with an “x” or “o” rather than a dot in order to distinguish them from the atom’s own electrons.

17. Lewis Dot Structures How to draw Lewis dot structures for ionic compounds:

18. Lewis Dot Structures 1.Draw the Lewis dot structure for the positive ion (with charge). If there is more than one positive ion, be sure to use subscripts.

19. Lewis Dot Structures 2.Draw the Lewis dot structure for the negative ion (with charge) very close by. If there is more than one negative ion, be sure to use subscripts. 3.That's IT!!!

20. Lewis Dot Structures Example 3 Draw the Lewis structure for the following ionic compounds: 1.Sodium Chloride 2.Magnesium Fluoride 3.Aluminum Oxide

21. Lewis Dot Structures How to draw Lewis structures for covalent substances. 1.Select a reasonable (symmetrical) “skeleton” for the molecule or polyatomic ion.

22. Lewis Dot Structures 2.Decide which atom is the central atom. This is the atom present in the fewest number, or, if there is the same number of all atoms, it is the LEAST electronegative element. Hydrogen can never be a central atom.

23. Lewis Dot Structures 3.Remember NASA: Calculate N, the number of valence electrons needed by ALL atoms in the molecule or ion to achieve noble gas configurations. Most of the time, each atom will need 8 electrons each. We’ll talk about exceptions later…

24. Lewis Dot Structures 4.Calculate A 1, the number of electrons available in the valence shells of all the atoms. The number of available is found by looking at the column number on the periodic table.

25. Lewis Dot Structures 5.Calculate S, the total number of electrons shared in the molecule or ion, using the relationship S = N – A 1.

26. Lewis Dot Structures 6.Place the S electrons into the skeleton as shared pairs. Use double and triple bonds only when necessary. Show the shared pairs as dashes. Each dash represents a pair of electrons.

27. Lewis Dot Structures 7.Calculate A 2, the number of additional electrons needing to be represented; A 2 = A 1 – S. Place the additional electrons into the skeleton as unshared pairs to fill the octet of the peripheral atoms. Place any electrons which are left over on the central atom in pairs. Remember that you must show all of the A electrons.

28. Lewis Dot Structures 8.Check to see if the central atom has at least an octet. If the particle is an ion, be sure to place it in square brackets and put the charge outside the bracket.

29. Lewis Dot Structure Example 4 Draw the Lewis structure for the following compounds: 1.Carbon Tetrachloride 2.Phosphorus Triflouride

30. APPLYING TH E VSEPR THEORY NOW WE APPLY THE VSEPR USING THE LEWIS DOT STRUCTURE

31. VSEPR THEORY Determine how many bonds, lone pairs, or unpaired electrons there are from the Lewis Dot Structure. Look at the notes and determine from the data on VSEPR which geometry applies to the molecule.

32. VSEPR IN ACTION Now, let’s divide into groups of four. Take a Styrofoam modeling kit, one per group. Come pick a molecule out of the jar on the lab table. Create the molecule using the modeling kit.

33. VSEPR IN ACTION THAT’S IT. LET’S GET TO WORK! Please raise your hand to get my attention, and I will come help your group.