1. General Chemistry M. R. Naimi-Jamal Faculty of Chemistry Iran University of Science & Technology

2. فصل هشتم: پیوند کووالانسی

3. Problems and questions: How is a molecule or polyatomic ion held together? Why are atoms distributed at strange angles? Why are molecules not flat? Can we predict the structure? How is structure related to chemical and physical properties? Chemical Bonding

4. There are 2 extreme forms of connecting or bonding atoms: Ionic—complete transfer of electrons from one atom to another Covalent—electrons shared between atoms Most bonds are somewhere in between. Forms of Chemical Bonds

5. Covalent bond forms by the sharing of VALENCE ELECTRONS, the electrons at the outer edge of the atom. Covalent Bonding The bond arises from the mutual attraction of 2 nuclei for the same electrons.

6. Valence Electrons Electrons are divided between core and valence electrons. Na 1s 2 2s 2 2p 6 3s 1 Core = [Ne] and valence = 3s 1 Br [Ar] 3d 10 4s 2 4p 5 Core = [Ar] 3d 10 and valence = 4s 2 4p 5

7. Chemical Bonding Objectives Objectives are to understand: 1. e - distribution in molecules and ions. 2. molecular structures 3. bond properties and their effect on molecular properties.

8. Electron Distribution in Molecules Electron distribution is depicted with Lewis electron dot structures Electrons are distributed as shared or BOND PAIRS and unshared or LONE PAIRS. G. N. Lewis 1875 - 1946

9. Lewis Symbols A chemical symbol represents the nucleus and the core e -. Dots around the symbol represent valence e -. Si N P As Sb Bi Al Se Ar I.

10. Lewis Structures for Ionic Compounds Ba O O Ba 2+ 2- Mg Cl Cl Cl Mg 2+ - 2 BaO MgCl 2

11. Bond and Lone Pairs Electrons are distributed as shared or BOND PAIRS and unshared or LONE PAIRS. This is called a LEWIS ELECTRON DOT structure.

12. Bond Formation A bond can result from a “head-to-head” overlap of atomic orbitals on neighboring atoms. This type of overlap places bonding electrons in a MOLECULAR ORBITAL along the line between the two atoms and forms a SIGMA BOND (  ).

13. Covalent Bonding

14. Coordinate Covalent Bonds H N H H H N H H H H + Cl Cl -

15. Multiple Covalent Bonds C O O C O O C O O C O O

16. Multiple Covalent Bonds N N N N N N N N

17. Polar Covalent Bonds H Cl δ+δ+δ-δ-

18. Electronegativity Mulliken electronegativity Pauling electronegativity = I + A 2  xA-xB = √ D/23.06 D = 2 E (A-B) – E (A-A) – E (B-B) I = Ionization Energy, A = Electron Affinity

19. Electronegativity

20. Dipole Moments

22. گشتاور دو قطبی‌ گشتاور دو قطبی‌ = فاصله x بار واحد گشتاور دو قطبی‌ = دبی‌ =3.34 x 10 -30 C.m

23. Percent Ionic Character

24. Writing Lewis Structures No. of valence electrons of an atom = group number For groups 1A-4A, no. of bond pairs = group number For groups 5A-7A, BP’s = 8 - gr. no. Except for H (and atoms of 3rd and higher periods), BP’s + LP’s = 4 This observation is called the OCTET RULE Writing Lewis Structures

25. All the valence e - of atoms must appear. Usually, the e - are paired. Usually, each atom requires an octet. –H only requires 2 e -. Multiple bonds may be needed. –Readily formed by C, N, O, S, and P.

26. Skeletal Structure Identify central and terminal atoms. C 2 H 5 OH C H H H H C H H O

27. Skeletal Structure Hydrogen atoms are always terminal atoms. Central atoms are generally those with the lowest electronegativity. Carbon atoms are always central atoms. Generally structures are compact and symmetrical.

28. Building a Dot Structure Ammonia, NH 3 1. Decide on the central atom; Central atom is generally atom of lowest affinity for electrons, but never H, here N is central. 2. Count valence electrons H = 1 and N = 5 Total = (3 x 1) + 5 = 8 electrons / 4 pairs

29. 3.Form a sigma bond between the central atom and surrounding atoms. 4.Remaining electrons form LONE PAIRS to complete octet as needed. 3 BOND PAIRS and 1 LONE PAIR. Note that N has a share in 4 pairs (8 electrons), while H shares 1 pair. Building a Dot Structure

30. Step 1. Central atom = S Sulfite ion, SO 3 2- Step 2. Count valence electrons S = 6 3 x O = 3 x 6 = 18 Negative charge = 2 TOTAL = 26 e- or 13 pairs Step 3. Form sigma bonds

31. 10 pairs of electrons are now left. Sulfite ion, SO 3 2-

32. Remaining pairs become lone pairs, first on outside atoms and then on central atom. OO O S Each atom is surrounded by an octet of electrons.

33. Carbon Dioxide, CO 2 1. Central atom = C 2. Valence electrons = 16 or 8 pairs 3. Form sigma bonds. This leaves 6 pairs. This leaves 6 pairs.

34. Carbon Dioxide, CO 2 4. Place lone pairs on outer atoms. BUT C doesn’t obey the octet rule! 5. So that C has an octet, we shall form DOUBLE BONDS between C and O.

35. Carbon Dioxide, CO 2 The second bonding pair forms a pi () bond. The second bonding pair forms a pi (  ) bond. OO C

36. Double and even triple bonds are commonly observed for C, N, P, O, and S H 2 CO SO 3 C2F4C2F4C2F4C2F4

37. Sulfur Dioxide, SO 2 1. Central atom = S 2. Valence electrons = 18 or 9 pairs 3. Form sigma bonds. Sulfur Dioxide, SO 2 OOS OOS This leaves 7 pairs. 4. Place 7 lone pairs on outer atoms.

38. Sulfur Dioxide, SO 2 5. Form pi () bond so that S has an octet — but note that there are two ways of doing this. 5. Form pi (  ) bond so that S has an octet — but note that there are two ways of doing this. Sulfur Dioxide, SO 2

39. This leads to the following structures: These equivalent structures are called: RESONANCE STRUCTURES RESONANCE STRUCTURES The true electronic structure is a HYBRID of the two. Sulfur Dioxide, SO 2

40. Urea, (NH 2 ) 2 CO

41. 1. Central atom = C 2. Number of valence electrons = 24 e - 3. Draw sigma bonds Urea, (NH 2 ) 2 CO

42. 4. Place remaining electron pairs in the molecule. Urea, (NH2)2CO

43. 5. Complete C atom octet with double bond. Urea, (NH 2 ) 2 CO

44. Violations of the Octet Rule Usually occurs with B and elements of higher periods. BF 3 SF 4....

45. Boron Trifluoride Central atom = BCentral atom = B Valence electrons = 24 or electron pairs = 12Valence electrons = 24 or electron pairs = 12 Assemble dot structureAssemble dot structure The B atom has a share in only 6 pairs of electrons (or 3 pairs). B atom in many molecules is electron deficient. The B atom has a share in only 6 pairs of electrons (or 3 pairs). B atom in many molecules is electron deficient.

46. Sulfur Tetrafluoride, SF 4 Central atom = SCentral atom = S Valence electrons = 34 or 17 pairs.Valence electrons = 34 or 17 pairs. Form sigma bonds and distribute electron pairs.Form sigma bonds and distribute electron pairs. 5 pairs around the S atom. A common occurrence outside the 2nd period.

47. Exceptions to the Octet Rule Odd e - species: N=O H—C—H H O—H

48. Exceptions to the Octet Rule Incomplete octets: B F FF

49. Exceptions to the Octet Rule Expanded octets: P Cl Cl Cl Cl S F F F F F F PCl 5 SF 6

50. Formal Charge FC = # valence e- - # lone pair e- - # bond pair e- 2 1

51. Carbon Dioxide, CO 2 4 - (1/2)(8) - 0 = 0 6 - (1/2)(4) - 4 = 0 OO C

52. Which is the predominant resonance structure? Carbon Dioxide, alternative lewis structure OOC 6 - (1/2)(6) - 2 = +1 6 - (1/2)(2) - 6 = -1

53. Boron Trifluoride, BF 3 What if we form a B—F double bond to satisfy the B atom octet?

54. Boron Trifluoride, BF 3 To have +1 charge on F, with its very high affinity for electrons, is not good.To have +1 charge on F, with its very high affinity for electrons, is not good. Negative charges are best placed on atoms with high affinity for electrons.Negative charges are best placed on atoms with high affinity for electrons. F F F B FC = 7 - 2 - 4 = +1 FC = 3 - 4 - 0 = -1 Boron Trifluoride, BF 3

55. Exceptions to the Octet Rule Incomplete octets. B F FF B F FF - + B F FF - +

56. Formal Charges Formal Charges & Lewis Structure Lewis Structure

57. Chapter 7 Questions 6, 8, 18, 21, 31 32, 38, 44, 48 52