1. Matter is composed of either CHEMICAL BONDING (1) Metals (2) Nonmetals (3) Metals and Nonmetals - Atoms - Molecules - Ions - Metallic Bonding - Covalent Bonding - Ionic Bonding 2A-1 (of 15) Chemical bonding involves the valence electrons of atoms

2. 1904 ARNOLD SOMMERFELD Proposed that metal atoms release their valence electrons, and share them between large numbers of metal atoms (2)MATTER COMPOSED OF METALS 2A-2 (of 15)

3. METALLIC BOND – The electrostatic attraction of the shared valence electrons to the nuclei of the many bonding metal atoms Metallic bonding forms crystalline networks containing billions of metal ATOMS that are strongly attracted together 2A-3 (of 15)

4. 1916 GILBERT NEWTON LEWIS Proposed that nonmetal atoms share valence electrons to achieve the electron configurations of Noble Gases (3) MATTER COMPOSED OF NONMETALS 2A-4 (of 15) Diatomic chlorine : Cl : : Cl : : : : Cl – Cl : :: ::

5. LEWIS STRUCTURE – A representation of chemical bonding using electron dot notation 2A-5 (of 15) Cl Cl :: :: BONDING PAIRS: in red LONE PAIRS:in green ::

6. COVALENT BOND – The electrostatic attraction of the shared electrons to the nuclei of the bonding nonmetal atoms Covalent bonding forms individual units called MOLECULES that are weakly attracted to each other 2A-6 (of 15)

7. To draw a proper Lewis Structure for a covalently bonded species: 1 – Add up the valence e - s for all of the atoms in the molecule or ion 2 – Draw a skeletal structure by using pairs of electrons to make bonds 4 – If octets are not produced, make the atoms that have octets share more e - pairs with atoms that do not have octets 3 – Complete octets (or duets for H) for all atoms, outer atoms first, using the remaining valence e - s LEWIS STRUCTURES 2A-7 (of 15)

8. Oxygen difluoride, OF 2 6 + 7 + 7 = 20 valence e - s F O F 2A-8 (of 15)

9. Nitrogen tribromide, NBr 3 5 + 7 + 7 + 7 = 26 valence e - s Br N Br Br 2A-9 (of 15)

10. 1904 RICHARD ABEGG Proposed that atoms gain or lose valence electrons to achieve the electron configurations of Noble Gases (1)MATTER COMPOSED OF METALS AND NONMETALS 2A-10 (of 15)

11. Metal atoms easily lose valence e - s, forming positive ions Al1s 2 2s 2 2p 6 3s 2 3p 1 Fe[Ar]4s 2 3d 6 Nonmetal atoms gain e - s to their valence shells, forming negative ions O1s 2 2s 2 2p 4 Once these ions are formed, they are stable (or unreactive) Al 3+ 1s 2 2s 2 2p 6 Fe 2+ [Ar]3d 6 O 2- 1s 2 2s 2 2p 6 Fe 3+ [Ar]3d 5 2A-11 (of 15)

12. 2A-12 (of 15) IONIC BOND – The electrostatic attraction between positive metal ions and negative nonmetal ions Ionic bonding forms crystalline networks containing billions of positive and negative IONS that are strongly attracted together

13. Sodium chloride Na... Cl :.. Na +.. : Cl : -.. A sodium chloride crystal is a symmetrical array of sodium and chloride ions in a 1:1 ratio EMPIRICAL FORMULA – The simplest whole number ratio of ions of different elements in a compound Empirical Formula: NaCl 2A-13 (of 15)

14. Calcium fluoride Empirical Formula: CaF 2 Ca... F :.. Ca 2+.. : F : -... F :.. : F : -.. 2A-14 (of 15)

15. K3NK3N K.K... N :. K.K. K.K. 2A-15 (of 15) REPRESENTING IONIC BONDING WITH ELECTRON DOT NOTATION K+K+.. : N : 3-.. K+K+ K+K+

17. Fluorine, F 2 7 + 7 = 14 valence e - s F SINGLE BOND – One shared pair of e - s between two atoms 2B-1 (of 15)

18. Oxygen, O 2 6 + 6 = 12 valence e - s O DOUBLE BOND – Two shared pairs of e - s between two atoms 2B-2 (of 15)

19. Nitrogen, N 2 5 + 5 = 10 valence e - s N TRIPLE BOND – Three shared pairs of e - s between two atoms 2B-3 (of 15)

20. BOND ORDER – The number of shared pairs of electrons BOND ENERGY – The energy needed to break a bond BOND LENGTH – The distance between the nuclei of the 2 bonding atoms Bond Order Bond Energy (kJ/mol) Bond Length (nm) F2O2N2F2O2N2 1 154 0.142 2 495 0.121 3 941 0.110 2B-4 (of 15)

21. H P S Cl I Longest Bond Length? Shortest Bond Length? Highest Bond Energy? Lowest Bond Energy? I2I2 H2H2 P2P2 I2I2 biggest atoms smallest atoms most bonding electrons least bonding electrons, and they are most shielded from the nuclei 2B-5 (of 15)

22. Formaldehyde, CH 2 O 4 + 1 + 1 + 6 = 12 valence e - s H C O H 2B-6 (of 15)

23. Sulfate, SO 4 2- 6 + 4(6) O O S O O + 2 = 32 valence e - s 2- 2B-7 (of 15)

24. NO 3 - 5 + 3(6) + 1 = 24 valence e - s O N O O - O N O O -- O N O O RESONANCE – When more than one Lewis structure can be drawn for a molecule or ion RESONANCE STRUCTURES – The Lewis structures that can be drawn for the molecule or ion The bonding in the real nitrate ion is an average of its resonance structures The average N-O bond order is (1+1+2) / 3 = 1 1 / 3 ↔↔ 2B-8 (of 15)

25. 1932 LINUS PAULING Described how atomic orbitals are involved in covalent bonding 2B-9 (of 15) VALENCE BOND THEORY – Two atoms share electrons by overlapping a valence atomic orbital from each atom, creating a region of space between the nuclei where the electrons reside

26. H atom 1s atomic orbital with 1 valence e - H atom 1s atomic orbital with 1 valence e - H 2 molecule The attraction of the e - s in the molecular orbital to the 2 nuclei bonds the atoms together 2B-10 (of 15) 2 valence e - s in a MOLECULAR ORBITAL

27. Metals – Low EN’s (the most active metals having the lowest EN’s) Nonmetals – High EN’s (the most active nonmetals have the highest EN’s) ELECTRONEGATIVITY – A property developed by Pauling, measuring the attraction of an atom for shared electrons 2B-11 (of 15) Atom with the highest EN? F (4.0) Atom with the lowest EN? Cs(0.7)

28. Nonpolar Covalent0 EN DifferenceBonding Polar CovalentSmall (0.1 – 1.6) Large (1.7 – 3.3)Ionic EN differences between atoms indicates their type of bonding 2B-12 (of 15)

29. 2 atoms with the same EN’s have an EN difference of 0 N – N (EN of N = 3.0) NONPOLAR COVALENT BOND – A bond between 2 atoms in which the electrons are shared evenly 2B-13 (of 15)

30. 2 atoms with close EN’s have an EN difference that is small H – Br (EN of H = 2.1, EN of Br = 2.8) POLAR COVALENT BOND – A bond between 2 atoms in which the electrons are shared unevenly Dipole Moment Arrow 2B-14 (of 15)

31. 2 atoms with extreme EN’s have an EN difference that is large Na – Cl (EN of Na = 0.8, EN of Cl = 3.0) IONIC BOND – A bond between 2 atoms in which the electrons are transferred, creating ions 2B-15 (of 15)

33. FORMAL CHARGE While atoms that covalently bond are not charged, they can be given charges based upon where the bonding electrons are assigned 2C-1 (of 11) FORMAL CHARGE – The charge given to an atom assuming one electron in each bond is assigned to that atom

34. .. 0 0 F S F FSF 0.. S naturally has 6 valence e - s, and now 6  0 F naturally has 7 valence e - s, and now 7  0 2C-2 (of 11) Quick way to determine formal charge: (natural number of valence e - s – 1 e - per bond – each lone pair e - ) S: F: 6 – 2 – 4 = 0 7 – 1 – 6 = 0

35. CO C O +1 2C-3 (of 11) Formal charges are used to determine the validity of a Lewis structure - the most accurate Lewis structures are those with atoms that have formal charges as close to 0 as possible C: O: 4 – 3 – 2 = -1 6 – 3 – 2 = +1

36. S C N thiocyanate, SCN - 6 + 4 + 5 + 1 = 16 valence e - s S C N - -- ↔↔ S: C: N: S: C: N: S: C: N: 6 – 2 – 4 = 0 4 – 4 – 0 = 0 5 – 2 – 4 = -1 6 – 3 – 2 = +1 4 – 4 – 0 = 0 5 – 1 – 6 = -2 6 – 1 – 6 = -1 4 – 4 – 0 = 0 5 – 3 – 2 = 0 2C-4 (of 11)

37. thiocyanate, SCN - 6 + 4 + 5 + 1 = 16 valence e - s S C N - -- ↔↔ The best Lewis structures have (1)formal charges for the most atoms as close to 0 as possible (2)negative formal charges go on the atom with the greatest EN 2C-5 (of 11) S: C: N: S: C: N: S: C: N: 6 – 2 – 4 = 0 4 – 4 – 0 = 0 5 – 2 – 4 = -1 6 – 3 – 2 = +1 4 – 4 – 0 = 0 5 – 1 – 6 = -2 6 – 1 – 6 = -1 4 – 4 – 0 = 0 5 – 3 – 2 = 0

38. COVALENT COMPOUNDS THAT DO NOT OBEY THE OCTET RULE (1)Molecules with hypovalent central atoms Covalent compounds with Band Be BeH 2 2 + 1 + 1 = 4 valence e - s H Be H (atoms with less than 4 valence electrons) 2C-6 (of 11)

39. BF 3 3 + 7 + 7 + 7 = 24 valence e - s F B F F NO! F B F F F is too electronegative to share more than 1 pair of e - s 2C-7 (of 11)

40. (2)Molecules with hypervalent central atoms (atoms that have empty d orbitals in their outer shell) Nonmetal atoms in the 3 rd, 4 th, 5 th, or 6 th Periods PF 5 5 + 5(7) = 40 valence e - s FPFP FF FF P can make 5 bonds using empty d orbitals in its outer shell 3s 3p ↑↓ ↑ ↑ ↑ ___ ___ ___ ___ ___ ___ ___ 3d 2C-8 (of 11)

41. ClF 3 7 + 3(7) = 28 valence e - s F F Cl F Only 26 valence e - s F F Cl F 2C-9 (of 11)

42. Sulfate, SO 4 2- 6 + 4(6) O O S O O + 2 = 32 valence e - s 2- Experimental data shows the S-O bonds are stronger than single bonds Reducing the formal charge on atoms that can exceed the octet rule can produce a more accurate Lewis structure S must make 2 double bonds to reduce its formal charge to 0 S: O: 6 – 4 – 0 = +2 6 – 1 – 6 = -1 2C-10 (of 11)

43. Sulfate, SO 4 2- 6 + 4(6) O O S O O + 2 = 32 valence e - s 2- Experimental data shows the S-O bonds are stronger than single bonds Reducing the formal charge on atoms that can exceed the octet rule can produce a more accurate Lewis structure S must make 2 double bonds to reduce its formal charge to 0 S: O: 6 – 1 – 6 = -1 6 – 2 – 4 = 0 + 5 other resonance structures 6 – 6 – 0 = 0 2C-11 (of 11)

45. MOLECULAR SHAPE VSEPR THEORY (Valence Shell Electron Pair Repulsion) – All atoms and lone pairs attached to a central atom will spread out as far as possible to minimize repulsion A Lewis structure must be drawn to use the VSEPR Theory 2D-1 (of 15)

46. CO 2 4 + 6 + 6 O C O = 16 valence e - s STERIC NUMBER (SN) – The sum of the bonded atoms and lone pairs on a central atom The steric number of carbon is 2 (SN = 2): 2 bonded atoms and 0 lone pairs Linear Bond angle is 180° O C O 2D-2 (of 15)

47. H B H H SN = 3 3 bonded atoms and 0 lone pairs Trigonal Planar Bond angle is 120° BH 3 3 + 1 + 1 + 1= 6 valence e - s H HH B 2D-3 (of 15)

48. SO 2 6 + 6 + 6 O S O = 18 valence e - s SN = 3 2 bonded atoms and 1 lone pairs Bent Bond angle is 120° O O S 2D-4 (of 15)

49. H H C H H SN = 4 4 bonded atoms and no lone pairs Tetrahedral Bond angle is 109.5° HC HC H H H CH 4 2D-5 (of 15)

50. H N H H SN = 4 3 bonded atoms and 1 lone pairs Trigonal Pyramidal Bond angle is 108° N H H H NH 3 2D-6 (of 15)

51. H2OH2O SN = 4 2 bonded atoms and 2 lone pairs Bent Bond angle is 105° O H H.. H – O : H 2D-7 (of 15)

52. F F P F F F SN = 5 5 bonded atoms and no lone pairs Trigonal Bipyramidal 3 Equatorial F’s in a plane, 120° apart 2 Axial F’s 180° apart, 90° from the plane F P F F F F PF 5 2D-8 (of 15)

53. SN = 5 4 bonded atoms and 1 lone pair SF 4 ←2 close 90º interactions 3 close 90º interactions → ←most stable configuration e - pair in equatorial position e - pair in axial position 2D-9 (of 15) : F : F S F : F :

54. SN = 5 4 bonded atoms and 1 lone pair SF 4 : F : F S F : F : FPFFPF F F See-Saw e - pairs always go in equatorial positions to minimize repulsion 2D-10 (of 15)

55. SN = 5 3 bonded atoms and 2 lone pairs ClF 3 F Cl F F Cl F F T-Shape 2D-11 (of 15)

56. SN = 5 2 bonded atoms and 3 lone pairs XeF 2 F Xe F F Xe F Linear 2D-12 (of 15)

57. SN = 6 6 bonded atoms and no lone pairs SF 6 FSFFSF FSFFSF F F Octahedral 90º and 180º F F F F F F 2D-13 (of 15)

58. SN = 6 5 bonded atoms and 1 lone pair IF 5 FIFI FI FI F F Square Pyramidal F F F F F F 2D-14 (of 15)

59. SN = 6 4 bonded atoms and 2 lone pairs XeF 4 Xe F F Square Planar F F F F Xe F F 2D-15 (of 15)

61. MOLECULAR POLARITY A BOND is polar if it has a positive end and a negative end A MOLECULE is polar if it has a positive end and a negative end To determine if a molecule is polar or nonpolar: 1)Draw the correct Lewis structure 2)Draw its correct shape 3)Use EN’s to determine if the BONDS in the molecule are polar or nonpolar 4)For the polar bonds, label the positive and negative ends with δ + and δ - 5)If a line can be drawn separating all δ +’s from all δ -’s, the molecule is polar, if not its nonpolar 2E-1 (of 13)

62. .. H – O : H O H H δ+δ+ δ-δ- δ+δ+ δ-δ- EN’s: O = 3.5, H = 2.1 3.5 – 2.1 = 1.4  the O-H BONDS are polar All of the δ+ ’s can be separated from all of the δ- ’s,  the H 2 O MOLECULE is polar 2E-2 (of 13)

63. δ+δ+ δ-δ- δ+δ+ δ-δ- EN’s: N = 3.0, H = 2.1 3.0 – 2.1 = 0.9  the N-H BONDS are polar All of the δ+ ’s can be separated from all of the δ- ’s,  the NH 3 MOLECULE is polar H N H H N H H H δ+δ+ δ-δ- 2E-3 (of 13)

64. F F C F F FC FC F F F EN’s: C = 2.5, F = 4.0 4.0 – 2.5 = 1.5  the C-F BONDS are polar All of the δ+ ’s cannot be separated from all of the δ- ’s,  the CF 4 MOLECULE is nonpolar δ+δ+ δ-δ- δ-δ- δ+δ+ δ-δ- δ+δ+ δ-δ- δ+δ+ 2E-4 (of 13)

65. A more exact way to determine if a molecule is polar or nonpolar: 1)Draw the correct Lewis structure 2)Draw its correct shape 3)Use EN’s to determine if the BONDS in the molecule are polar or nonpolar 4)For the polar bonds, draw a DIPOLE MOMENT ARROW pointing toward the negative end of the bond 5)If the dipole moments are symmetrical the molecule is NONPOLAR 2E-5 (of 13)

66. Dipole moments of equal magnitude are symmetrical if: 1)there are 2 dipole moments that are linear Y X Y 2E-6 (of 13)

67. Dipole moments of equal magnitude are symmetrical if: 2)there are 3 dipole moments that are trigonal planar YXYX Y Y 2E-7 (of 13)

68. Dipole moments of equal magnitude are symmetrical if: 3)there are 4 dipole moments that are tetrahedral YXYX Y Y Y 2E-8 (of 13)

69. O C O Symmetrical dipole moments  the CO 2 MOLECULE is nonpolar 2E-9 (of 13)

70. .. H – O : H O H H Assymmetrical dipole moments  the H 2 O molecule is POLAR 2E-10 (of 13)

71. H N H H N H H H Assymmetrical dipole moments  the NH 3 molecule is POLAR 2E-11 (of 13)

72. F F C F F FC FC F F F Symmetrical dipole moments  the CF 4 molecule is NONPOLAR 2E-12 (of 13)

73. Cl F C F F Cl C F F F Assymmetrical dipole moments because the C-Cl dipole moment is smaller than the C-F dipole moments  the CClF 3 molecule is POLAR 2E-13 (of 13)