General Chemistry

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

2. فصل هشتم: پیوند کووالانسی

3. Chemical Bonding 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?

4. Forms of Chemical Bonds • 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.

5. Covalent Bonding Covalent bond forms by the sharing of VALENCE ELECTRONS, the electrons at the outer edge of the atom. 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 1s2 2s2 2p6 3s1 Core = [Ne] and valence = 3s1 Br [Ar] 3d10 4s2 4p5 Core = [Ar] 3d10 and valence = 4s2 4p5

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

10. • •• 2+ 2- •• Ba O Ba O •• • • • •• •• •• •• Cl •• • •• • 2+ - •• Mg Mg 2 Cl •• • •• •• •• Cl •• • •• Lewis Structures for Ionic Compounds BaO MgCl2

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

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

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

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

17. Polar Covalent Bonds δ+ δ- Cl H

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

19. Electronegativity

20. Dipole Moments

21. 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 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

25. Writing Lewis Structures • 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. H H H C C O H H H Skeletal Structure • Identify central and terminal atoms. • C2H5OH

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, NH3 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. Building a Dot Structure 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.

30. Sulfite ion, SO32- Step 1. Central atom = S Step 3. Form sigma bonds Step 2. Count valence electrons S = 6 3 x O = 3 x 6 = 18 Negative charge = 2 TOTAL = 26 e- or 13 pairs

31. Sulfite ion, SO32- 10 pairs of electrons are now left.

32. • • O • • • • •• • • O S O • • • • •• •• • • Sulfite ion, SO32- Remaining pairs become lone pairs, first on outside atoms and then on central atom. Each atom is surrounded by an octet of electrons.

33. Carbon Dioxide, CO2 1. Central atom = C 2. Valence electrons = 16 or 8 pairs 3. Form sigma bonds. This leaves 6 pairs.

34. Carbon Dioxide, CO2 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. O C O • • • • •• •• Carbon Dioxide, CO2 The second bonding pair forms a pi (p) bond.

36. Double and even triple bonds are commonly observed for C, N, P, O, and S H2CO SO3 C2F4

37. O S O •• •• •• O S O • • • • •• •• Sulfur Dioxide, SO2 Sulfur Dioxide, SO2 1. Central atom = S 2. Valence electrons = 18 or 9 pairs 3. Form sigma bonds. This leaves 7 pairs. 4. Place 7 lone pairs on outer atoms.

38. Sulfur Dioxide, SO2 Sulfur Dioxide, SO2 5. Form pi (p) bond so that S has an octet — but note that there are two ways of doing this.

39. Sulfur Dioxide, SO2 This leads to the following structures: These equivalent structures are called: RESONANCE STRUCTURES The true electronic structure is a HYBRID of the two.

40. Urea, (NH2)2CO

41. Urea, (NH2)2CO Urea, (NH2)2CO 1. Central atom = C 2. Number of valence electrons = 24 e- 3. Draw sigma bonds

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

43. Urea, (NH2)2CO Urea, (NH2)2CO 5. Complete C atom octet with double bond.

44. Violations of the Octet Rule Usually occurs with B and elements of higher periods. . . BF3 SF4

45. Boron Trifluoride • Central atom = B • Valence electrons = 24 or electron pairs = 12 • Assemble 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.

46. Sulfur Tetrafluoride, SF4 • Central atom = S • Valence electrons = 34 or 17 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 •• O—H • H—C—H • ••

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

49. •• •• •• F •• Cl •• •• •• •• •• •• F •• •• Cl F •• Cl •• •• •• •• •• S P •• •• •• F F •• •• Cl Cl •• •• F •• •• •• •• •• •• •• •• Exceptions to the Octet Rule • Expanded octets: PCl5 SF6

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