1. CHEMISTRY 161 Chapter 9

2. Periodic Table of the Elements chemical reactivity - valence electrons ns 1 ns 2 ns 2 np x

3. atoms combine to form compounds in an attempt to obtain a stable noble gas electron configuration THE OCTET RULE ns 2 np 6 Iso electronic

4. 2. ELECTRON SHARING 1. ELECTRON FULLY TRANSFERED IONIC BONDING COVALENT BONDING A + B → AB 2 Na(s) + Cl 2 (g)  2 NaCl(s) 2 H 2 (g) + O 2 (g)  2 H 2 O(l) EXP I

5. LEWIS MODEL OF BONDING Gilbert Lewis (1875-1946) LEWIS DOT SYMBOL DOT represents one valence electron H.H.

6. with the exception of He, the main group number represents number of ‘dots’.......... only valence electron are considered

7. electron transfer IONIC BONDING Lewis Symbol 1s 2 2s 2 2p 6 3s 1 Ne core implied in symbol Na

8. Lewis Symbol 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 3s 2 3p 5 Ne core implied in symbol Cl Na

9. Na +  1s22s22p61s22s22p6 1s 2 2s 2 2p 6 3s 2 3p 6 the loss or gain of electrons(dots) until IONIC BONDING the formation of ionic bonds is represented in terms of Lewis symbols both species have reached an octet of electrons Cl Na Cl

10. [Ne] 3s 2 3p 6 represents one orbital (Pauli: 2 electrons)

11. ions stack together in regular crystalline structures ionic solids typically 1. high melting and boiling points 2. brittle 3. form electrolyte solutions if they dissolve in water electrostatic interaction

12. Li(s) + ½ F 2 (g) → LiF(s) lattice energy (up to few 1000 kJmol -1 ) enthalpy of formation Hess’s Law Li + (g) + F - (g) → LiF(s) Born-Haber Cycle

13. Li(s) + ½ F 2 (g) → LiF(s) Li(s) + ½ F 2 (g) Li(g) + F(g) Li + (g) + F - (g) LiF(s) ΔHoRΔHoR ΔHo1ΔHo1 ΔHo2ΔHo2 ΔHo3ΔHo3 ΔHo4ΔHo4 ΔHo5ΔHo5 ΔH o R = Σ ΔH o i i=1 5

14. Mg(s) + ½ O 2 (g) → MgO(s) Mg(s) + ½ O 2 (g) Mg(g) + O(g) Mg 2+ (g) + O 2- (g) MgO(s) ΔHoRΔHoR ΔHo1ΔHo1 ΔHo2ΔHo2 ΔHo3ΔHo3 ΔHo4ΔHo4 ΔHo5ΔHo5 ΔH o R = Σ ΔH o i i=1 7 Mg + (g) + O - (g) ΔHo6ΔHo6 ΔHo7ΔHo7

15. FF COVALENT BONDING sharing electrons (electron pair) FF electronic configuration of F is 1s 2 2s 2 2p 5 THE OCTET RULE

16. FF + FFFF non-bonding, or lone pair of electrons bonding pair of electrons

17. H 2 H 2 is the simplest covalent molecule H + HHH + the bond length of H 2 is the distance where the total energy of the molecule is minimumbond length

19. EXAMPLES CO 2 H2OH2OH2OH2O NH 3 CH 4 single bonds HX C2H4C2H4C2H4C2H4 double bonds O2O2O2O2 triple bonds C2H2C2H2C2H2C2H2 N2N2N2N2 HCN

20. few 100 kJ/mol

22. electronegativity difference between two atoms involved in the bond IONIC OR COVALENT

23. Electronegativity increases F > O > N, Cl > Br > I, C, S …….. Na, Ba, Ra > K, Rb > Cs, Fr ELECTRONEGATIVITY is the tendency of an atom in a bond to attract shared electrons to itself F Cl I ONC SNa K Rb Cs FrRa Ba Li Br energies of the atomic orbital with the unpaired electron

24. Se Electronegativity increases F > O > N, Cl > Br > I, C, S …….. Na, Ba, Ra > K, Rb > Cs, Fr ELECTRONEGATIVITY F is the most electronegative Li Br P I NO Cl FC SNa K Rb Cs FrRa Ba H has an electronegativity about the same a P

25. bonds are neither completely ionic nor covalent (only in homonuclear molecules) IONIC VERSUS COVALENT BONDS

26. compounds composed of elements with large difference in ELECTRONEGATIVITY significant ionic character in their bonding A B B has greater electronegativity

27. IONIC VERSUS COVALENT BONDS A B B has a greater share

28. H + + FFH HYDROGEN FLUORIDE Fluorine is more electronegative than hydrogen.

29. H + + F ++ –– FH HYDROGEN FLUORIDE This is a polar covalent bond (dipole moment). The bond has a partly ionic and partly covalent nature. Fluorine is more electronegative than hydrogen.

30. Microwave Spectroscopy molecules need a dipole moment

31. Variation of ionic character with electronegativity.

32. IONIC COMPOUDS COVALENT COMPOUNDS ELECTRONEGATIVITY LEWIS SYMBOLS

33. OHH non-bonding, or lone pair of electrons bonding pair of electrons Lewis considers only valence electrons H2OH2O single – double – triple

34. 1. concept of resonances LEWIS STRUCTURES 2. exceptions to the octet rule

35. 1. RESONANCES OOOOOO NO 3 - N: 1s 2 2s 2 2p 3 O: 1s 2 2s 2 2p 4 plus one extra electron for negative charge

36. + - -

37. experiment shows all three bonds are the same any one of the structures suggests one is different! OO O N bond angles 120 0 128 pm

38. modify the description by blending the structures blending of structures is called resonance OO O N bond angles 120 0 128 pm

39. use a double headed arrow between the structures electrons involved are said to be DELOCALIZED over the structure. blended structure is a RESONANCE HYBRID RESONANCE N OO O OO O NN OO O

40. We use a double headed arrow between the structures.. RESONANCE N OO O OO O NN OO O OO O N

41. 1. more than 8 electrons around central atom 3. molecules with unpaired electrons 2. less than an octet around central atom 2. Exceptions to the octet rule

42. elements in rows 3 and following can exceed octet rule F S F F F F F SF 6 S F F F F F F 1. more than 8 electrons around central atom participation of d electrons

43. Lewis structure for SF 6 F has seven 1s 2 2s 2 2p 5 S has six 1s 2 2s 2 2p 6 3s 2 2p 4 SF 2 SF 4 SF 6 PF 3 PF 5 NF 3 NF 5 ClO 4 - SO 4 2- I3-I3-I3-I3-

44. 2. less than an octet around central atom BeH 2 AlF 3 BF 3 resonances NH 3 (dative bond) Lewis acids Lewis base

45. 3. molecules with unpaired electrons FREE RADICALS NO but not NO -