Polarity, Lewis Structures, and Resonance

1. Polarity, Lewis Structures, and Resonance Sections 8.4-8.6

2. Bond Polarity • Ionic and covalent bonding is not black and white • Sharing is not usually equal • It’s an electron tug of war • Must look at electronegativity to determine how equally electrons are shared

3. Electronegativity • The ability of an atom in a molecule to attract electrons to itself • Developed by Linus Pauling • Gave values for all elements based on thermochemical data Linus Carl Pauling (1901-1994)

4. Electronegativity: • On the periodic chart, electronegativity increases as you go… • …from left to right across a row. • …from the bottom to the top of a column.

5. Electronegativities of the Elements Cs (EN = 0.7) is least electronegative element F with EN = 4.0 is most electronegative element Au is at the peak of an island of electronegativity, and is most electronegative metal

6. Electronegativity and Bond Polarity Electronegativity tells us what kind of bonding we have, i.e. whether it is ionic or covalent. The greater the difference in EN between the two elements forming the bond, the more ionic is the bond. Typical ranges for EN differences are: EN difference bonding type Example EN difference range __________________________________________________________________________ > 2.0 Ionic LiF 4.0-1.0 = 3.0 0.5-2.0 polar covalent HF 4.0-2.1 = 1.9 <0.5 nonpolar covalent F-F 4.0-4.0 = 0.0

7. Polar Covalent Bonds • When two atoms share electrons unequally, a bond dipole results. • The dipole moment, , produced by two equal but opposite charges separated by a distance, r, is calculated:  = Qr • It is measured in debyes (D).

8. Polar Covalent Bonds The greater the difference in electronegativity, the more polar is the bond.

9. Bond Types and Nomenclature Naming ionic and covalent compounds

10. Naming Ionic Compounds 1. Name cation first, then anion • Cation = name of the element Ca2+ = calcium • Anion = root + -ide Cl = chloride CaCl2 = calcium chloride

11. Naming Ionic Compounds(continued) • Transition metals may form more than one cation • Use Roman numeralin name PbCl2 Pb2+is cation PbCl2 = lead (II) chloride If the metal is a transition metal:

12. Naming Ionic Compounds(continued) • If a polyatomic ion is present just use its name • Ex: NaCN = sodium cyanide

13. Naming Molecular Compounds • 1st elementin the formula is named first • 2nd elementnamed as if it were an anion • Greek prefixes denote how many atoms of element are present • Do not use mono- for the first element P2O5 = diphosphorus pentoxide

15. 11 hendeca-12 dodeca-13 triskaideka-14 tetradeca-15 pentadeca- 16 hexadeca-17 heptadeca-18 octadeca-19 enneadeca-20 icosa- Greek prefixes cont.

16. Lewis Structures Representations of molecules showing all electrons, bonding and nonbonding.

17. Find the sum of valence electrons of all atoms - For anions, add 1 electron for each negative charge. For cations, subtract 1 electron for each positive charge. PCl3 Writing Lewis Structures 5 + 3(7) = 26

18. Writing Lewis Structures • The central atom is the least electronegative that isn’t hydrogen. Connect the outer atoms to it by single bonds. Keep track of the electrons: 26  6 = 20

19. Writing Lewis Structures • Fill the octets of the outer atoms. Keep track of the electrons: 26  6 = 20  18 = 2

20. Writing Lewis Structures • Fill the octet of the central atom. Keep track of the electrons: 26  6 = 20  18 = 2  2 = 0

21. Writing Lewis Structures • If you run out of electrons before the central atom has an octet… …form multiple bonds until it does.

22. Formal Charge • Then assign formal charges. • For each atom, count the electrons in lone pairs and half the electrons it shares with other atoms. • Subtract that from the number of valence electrons for that atom: The difference is its formal charge.

23. Writing Lewis Structures • The best Lewis structure… • …is the one with the fewest charges. • …puts a negative charge on the most electronegative atom.

24. Resonance This is the Lewis structure we would draw for ozone, O3. + -

25. Resonance • But this is at odds with the true, observed structure of ozone, in which… • …both O—O bonds are the same length. • …both outer oxygens have a charge of 1/2.

26. Resonance • One Lewis structure cannot accurately depict a molecule such as ozone. • We use multiple structures, resonance structures, to describe the molecule.

27. Resonance Just as green is a synthesis of blue and yellow… …ozone is a synthesis of these two resonance structures.

28. Resonance • In truth, the electrons that form the second C—O bond in the double bonds below do not always sit between that C and that O, but rather can move among the two oxygens and the carbon. • They are not localized, but rather are delocalized.

29. Resonance • The organic compound benzene, C6H6, has two resonance structures. • It is commonly depicted as a hexagon with a circle inside to signify the delocalized electrons in the ring.