1. Molecular Geometry: Lewis Structures & VSEPR

2. Our Goal… To determine the shape and polarity of a molecule using Lewis structures.

3. Background There are several different bonding theories; the one used depends on the question asked.  How are the electrons distributed? Lewis Structure  What is the shape of the molecule? VSEPR theory  How are the molecules put together? Hybridization theory

4. Lewis Structures: Rules of the Game 1.Determine the total number of valence e’s available from all of the bonding atoms. 2.Calculate the total number of valence e’s that would be needed to give every Hydrogen a “duet” and all other atoms an “octet.” 3.Subtract (step 2) - (step 1) = Number of electrons to share. 4.Divide : step 3 / 2 = Number of covalent bonds to include (Since e’s are always shared in pairs) 5.Draw a “skeleton structure” with the least common atom in the center, and other atoms arranged to give a symmetryical molecule. For acids, show each hydrogen attached to an oxygen.For acids, show each hydrogen attached to an oxygen. 6.Subtract (step 1) – (step 3) = number of unshared electrons to arrange around the atoms to give each “non-hydrogen” atom an octet.

5. Lewis structures PCl 3

6. Step 1 Total available e’s is: P: 1 X 5 = 5 Cl: 3 X 7 = 21 Total = 26 Lewis structures

7. PCl 3 Step 2 Number of e’s needed to give each atom an octet 4 Atoms X 8 e- = 32 e’s

8. Lewis structures PCl 3 Step 3 Number of shared electrons is: (step 2) – (step 1) = 32 – 26 = 6

9. Lewis structures PCl 3 Step 4 Number of bonds = (step 3) / 2 = 6/2 = 3

10. Lewis structures PCl 3 Step 5 Draw the symmetrical molecule Cl—P—Cl Cl

11. Lewis structures PCl 3 Step 6 Number of unshared e’s : (step 1) – (step 3) = 26 – 6 = 20 Cl—P—Cl Cl.

12. Lewis structures CO 2

13. Step 1 Total available e’s is: C: 1 X 4 = 4 O: 2 X 6 = 12 Total = 16 Lewis structures

14. CO 2 Step 2 Number of e’s needed to give each atom an octet 3 Atoms X 8 e- = 24 e’s

15. Lewis structures CO 2 Step 3 Number of shared electrons is: (step 2) – (step 1) = 24 – 16 = 8

16. CO 2 Step 4 Number of bonds = (step 3) / 2 = 8/2 = 4 Lewis structures

17. CO 2 Step 5 Draw the symmetrical molecule. Note: The get four bonds when only two atoms are attached to the central atom, we’ll need double bonds! Lewis structures

18. O = C = O. CO 2 Step 6 Number of unshared e’s : (step 1) – (step 3) = 16 – 8 = 8

19. Valence Shell Electron Pair Repulsion Theory (VSEPR) repulsionpairsvalence shell electrons - VSEPR theory is based on the idea that repulsion between pairs of valence shell electrons forces atoms to assume positions as far apart as possible. VSEPR Theory

20. VSEPR “AXE” notation A A = Represents the central atom X X = Number of atoms directly attached to the central atom E E = Number of lone electron pairs on the central atom So, PCl 3 has the VSEPR formula: VSEPR Theory: An axe to grind

21. VSEPR “AXE” notation A = The central atom X = Number of atoms directly attached to the central atom E = Number of lone electron pairs on the central atom So, PCl 3 has the VSEPR formula:...... : Cl – P – Cl : ˙ ˙ | ˙ ˙ : Cl : ˙ ˙ VSEPR Theory AX 3 E 1

22. PCl 3 has the Lewis Structure:...... : Cl – P – Cl : ˙ ˙ | ˙ ˙ : Cl : ˙ ˙ This means PCl 3 has: -Four electron clouds, 3 shared pairs, 1 lone pair VSEPR Theory

23. VSEPR “AXE” notation for CO 2 : A A = Represents the central atom X X = Number of atoms directly attached to the central atom E E = Number of lone electron pairs on the central atom So, CO 2 has the VSEPR formula: VSEPR Theory O = C = O.

24. VSEPR Theory VSEPR “AXE” notation for CO 2 : A = Represents the central atom X = Represents the atoms directly attached to the central atom E = Number of lone electron pairs on the central atom (if there are none, don’t show E) So, CO 2 has the VSEPR formula: AX 2 VSEPR Theory O = C = O.

25. CO 2 has the Lewis Structure: This means it has: -Two electron clouds, two shared and no lone pairs VSEPR Theory O = C = O.

26. Geometries and Hybridizations

28. -To determine if a molecule is polar, you need to know two things: - polarity of the bonds in a molecule - polarity of the bonds in a molecule - how the bonds are arranged - how the bonds are arranged -A molecule is considered polar if its center of negative and positive charge do not coincide. Polarity

29. -To determine if a molecule is polar, you need to know two things: - polarity of the bonds in a molecule - polarity of the bonds in a molecule - how the bonds are arranged - how the bonds are arranged -A molecule is considered polar if its center of negative and positive charge do not coincide. -Polar molecules have a dipole (a vector quantity) Polarity

30. -To determine if a molecule is polar, you need to know two things: - polarity of the bonds in a molecule - polarity of the bonds in a molecule - how the bonds are arranged - how the bonds are arranged -A molecule is considered polar if its center of negative and positive charge do not coincide. -Polar molecules have a dipole (a vector quantity) -If these dipoles act equally and in opposition to each other, the dipoles cancel out and the molecule is considered nonpolar. Polarity

31. Two simple rules to help determine molecular polarity (most of the time) If there are lone pairs on the central atom – the molecule is polar.If there are lone pairs on the central atom – the molecule is polar. If there is more than one “type” of atom directly attached to the central atom – the molecule is polar.If there is more than one “type” of atom directly attached to the central atom – the molecule is polar. Polarity

32. Polarity  Question: Why is water polar, but carbon dioxide is not?  Question: Is BF 3 polar? Defend your answer!  Question: What is the capitol of Argentina? What does this have to do with molecular shape?

33. Used to predict the correct Lewis Structure. 1)Half of the electrons in a bond are assigned to each atom in the bond. 2)Both electrons of an unshared pair of electrons are assigned to the atoms to which the unshared pair belong. 3)The formal charge of an atom is equal to the valence electrons minus the number of electrons assigned to each atom. Formal Charge = (group number) – (assigned electrons) 4)The sum of the formal charges will equal the charge on the molecule or polyatomic ion. Formal Charge

34. 1)A Lewis structure in which all formal charges are equal to zero is preferred. 2)If a Lewis structure has non-zero formal charges, the structure with the fewest nonzero formal charges is preferred. 3)A Lewis structure with one large formal charge is preferable to a structure with several small charges. 4)A Lewis structure with adjacent formal charges should have opposite signs. 5)When choosing between several Lewis structures, the structure with negative formal charge on the more electronegative atom is preferable. Formal Charge