1. 1 Chapter 1: Covalent Bonding and Shapes of Molecules Aspartame (Nutrasweet ® )

2. 2 I.Lewis Structures A.Lewis symbols of elements B.Ionic, covalent, and polar bonds C.Lewis structures D.Formal charge E.Resonance structures II.Molecular Shapes A.VSEPR theory B.Polarity of molecules III.Valence Bond Model A.Atomic and molecular orbitals B.Hybrid atomic orbitals IV.Functional Groups

3. 3 I. Lewis Structures A. Lewis symbols of elements Periodic Table Review:periods; principle quantum numbers s-block, p-block, d-block; groups s and p atomic orbitals rules for filling atomic orbitals core and valence electrons Electron ConfigurationLewis Symbol H C N O F Cl

4. 4 I. Lewis Structures B. Ionic, covalent, and polar bonds ionic bonds: transfer of electrons covalent bonds: sharing of electrons electronegativity,  : relative attraction for electrons in a bond - increases going up and to the right in the periodic table - Pauling electronegativity scale (arbitrary): Table 1.5

5. 5 I. Lewis Structures B. Ionic, covalent, and polar bonds H—H  = 2.1 2.1  = 0  equal sharing of electrons Cl—Cl= nonpolar covalent bond  = 3.0 3.0 H—Cl  = 0.9  unequal sharing of electrons  = 2.1 3.0 = polar covalent bond Na + Cl –  = 2.1  transfer of electrons  = 0.9 3.0 = ionic bond generally: when  < 1.9  covalent > 1.9  ionic  +  – nonmetal + nonmetal metal + nonmetal

6. 6 I. Lewis Structures C. Lewis structures 1. Count all the valence electrons; add one for each – charge subtract one for each + charge 2. Draw single bonds between the atoms (the connectivity of the atoms is determined experimentally and is usually given in the problem). 3. Using the remaining valence electrons, place octets on all atoms (exception H), in order of decreasing electronegativity. 4. If atoms do not have octets, use lone pair electrons on adjacent atoms to form double or triple bonds to complete the octets.

7. 7 I. Lewis Structures C. Lewis structures CCl 4 CH 2 O C 2 H 2 CH 3 OH CH 3 CHCH 2 HCN

8. 8 I. Lewis Structures D. Formal charge (Use the silly, complex formula in the textbook, or use this easier method:) 1. Divide the electrons in each bond equally between the two atoms sharing them. 2. Count the number of electrons each atom now has and compare this number to its normal valence. more electrons than normal valence  negative formal charge fewer electrons than normal valence  positive formal charge H 3 O + CH 3 O – CH 3 + CON 3 –

9. 9 I. Lewis Structures D. Formal charge When two or more nonequivalent Lewis structures are possible, the better (more stable) one is the one with: 1. fewer formal charges 2. more octets 3. a – charge on a more electronegative atom, or a + charge on a more electropositive atom COCl 2 BF 3 (CH 3 ) 2 SOHOCN In decreasing order of importance

10. 10 I. Lewis Structures E. Resonance structures -two or more equivalent Lewis structures -nuclei remain in fixed positions, but electrons arranged differently neither of these accurately describes the formate ion actual species is an average of the two (resonance hybrid) delocalized electrons

11. 11 I. Lewis Structures E. Resonance structures more stable major contributor less stable minor contributor Draw resonance structures for the following species. If the structures are not equivalent, indicate which would be the major contributor. CH 3 NO 2 CH 2 CHO –

12. 12 II. Molecular Shapes A. VSEPR theory e – pairmolecular formulaLewis structuregeometrygeometry (angle) C 2 H 2 CH 2 O CH 4 HCN NH 3 H 2 O

13. 13 II. Molecular Shapes B. Polarity of molecules CCl 4 CHCl 3 CH 2 OCO 2 If the individual dipole moments in a molecule do not exactly cancel, then the molecule will have a net dipole moment and be a polar molecule.

14. 14 III. Valence Bond Model A. Atomic and molecular orbitals H + H  H—H Two electrons in  1s are lower energy than in the separate atoms  covalent bond

15. 15 III. Valence Bond Model B. Hybrid atomic orbitals 1. sp 3 hybridization CH 4 facts: tetrahedral, 4 equivalent bonds

16. 16 III. Valence Bond Model B. Hybrid atomic orbitals 1. sp 3 hybridization

17. 17 III. Valence Bond Model B. Hybrid atomic orbitals 2. sp 2 hybridization C 2 H 4 facts: all six atoms lie in same plane trigonal planar = sp 2

18. 18 III. Valence Bond Model B. Hybrid atomic orbitals 2. sp 2 hybridization

19. 19 III. Valence Bond Model B. Hybrid atomic orbitals 3. sp hybridization C 2 H 2 facts:linear = sp

20. 20 III. Valence Bond Model B. Hybrid atomic orbitals 3. sp hybridization

21. 21 III. Valence Bond Model B. Hybrid atomic orbitals What is the hybridization of each indicated atom in the following compound? 17-ethynylestradiol (“The Pill”)

22. 22 IV. Functional Groups Atoms or groups of atoms that behave similarly, regardless of the structure to which they are attached. methanolethanol  -phenethyl alcohol (lilacs) geraniol retinol (vit A)

23. 23 IV. Functional Groups Hydrocarbons (C & H only)Heteroatomic compounds aliphaticaromatic alkanes alkenes alkynes cyclic compounds alcohols ethers aldehydes ketones carboxylic acids esters amines amides