1. CHAPTER 8 Basic Concepts in Chemical Bonding

2. Introduction Attractive forces that hold atoms together in compounds are called chemical bonds. The electrons involved in bonding are usually those in the outermost (valence) shell.

3. Introduction Chemical bonds are classified into two types: Ionic bonding results from electrostatic attractions among ions, which are formed by the transfer of one or more electrons from one atom to another. Covalent bonding results from sharing one or more electron pairs between two atoms.

4. Lewis Dot Representations of Atoms The Lewis dot representation (or Lewis dot formulas) convenient bookkeeping method for valence electrons electrons that are transferred or involved in chemical bonding

5. Lewis Dot Representations of Atoms

6. elements in same group have same Lewis dot structures

7. Lewis Dot Representations of Atoms elements in same group have same Lewis dot structures

8. Ionic Bonding We can also use Lewis formulas to represent the neutral atoms and the ions they form.

9. Ionic Bonding Li + ions contain two electrons same number as helium F - ions contain ten electrons same number as neon Li + ions are isoelectronic with helium F - ions are isoelectronic with neon Isoelectronic species contain the same number of electrons.

10. Ionic Bonding potassium reacts with bromine equation for the reaction

11. Ionic Bonding potassium reacts with bromine 2 K(s) + Br 2 ( l )  + Br - (s) electronically this is occurring 4s 4p K [Ar]  Br [Ar]   becomes

12. Ionic Bonding 4s 4p K + [Ar] Br -   [Kr] Lewis dot formulas

13. Ionic Bonding cations become isoelectronic with preceding noble gas anions become isoelectronic with following noble gas

14. Covalent Bonding covalent bonds form when atoms share electrons share 2 electrons - single covalent bond share 4 electrons - double covalent bond share 6 electrons - triple covalent bond attraction is electrostatic in nature lower potential energy when bound

15. Covalent Bonding

17. Lewis dot representation H molecule formation HCl molecule formation

18. Covalent Bonding extremes in bonding pure covalent bonds electrons equally shared by the atoms pure ionic bonds electrons are completely lost or gained by one of the atoms most compounds fall somewhere between these two extremes

19. Lewis Dot Formulas for Molecules and Polyatomic Ions homonuclear diatomic molecules hydrogen, H 2 fluorine, F 2 nitrogen, N 2

20. Lewis Dot Formulas for Molecules and Polyatomic Ions heteronuclear diatomic molecules hydrogen halides hydrogen fluoride, HF hydrogen chloride, HCl hydrogen bromide, HBr

21. Lewis Dot Formulas for Molecules and Polyatomic Ions water, H 2 O

22. Lewis Dot Formulas for Molecules and Polyatomic Ions ammonia molecule, NH 3

23. Lewis Dot Formulas for Molecules and Polyatomic Ions ammonium ion, NH 4 +

24. The Octet Rule Elements try to a achieve noble gas configuration in most of their compounds. Lewis dot formulas are based on the octet rule.

25. N - A = S rule N = # of electrons needed to be noble gas usually 8 2 for H A = # of electrons available in outer shells equal to group # 8 for noble gases

26. N - A = S rule for ions add one e - for each negative charge subtract one e - for each positive charge central atom in a molecule or polyatomic ion is determined by: atom that requires largest number of electrons for two atoms in same group - less electronegative element is central

27. Drawing Lewis Dot Formulas Write Lewis dot and dash formulas for hydrogen cyanide, HCN. N = 2 + 8 + 8 = 18 A = 1 + 4 + 5 = 10 S = 8

28. Drawing Lewis Dot Formulas Write Lewis dot and dash formulas for the sulfite ion, SO 3 2-. N = 8 + 3(8)= 32 A = 6 + 3(6) + 2 = 26 S = 6

29. Drawing Lewis Dot Formulas Write Lewis dot and dash formulas for sulfur trioxide, SO 3. N = 8 + 3(8) = 32 A = 6 + 3(6) = 24 S =8

30. Resonance three possible structures for SO 3 two or more dot structures necessary to show bonding invoke resonance Double-headed arrows are used to indicate resonance formulas.

31. Resonance Structures Some molecules are not described by Lewis Structures very well. Typically, structures with multiple bonds can have similar structures with the multiple bonds between different pairs of atoms.

32. Resonance flaw in our representations of molecules no single or double bonds in SO 3 all bonds are the same best picture incorrect shape

33. Limitations of the Octet Rule for Lewis Formulas species that violate the octet rule N - A = S rule does not apply covalent compounds of Be, B, N covalent compounds with expanded octets species containing an odd number of electrons

34. Limitations of the Octet Rule for Lewis Formulas In cases where the octet rule does not apply, the substituents attached to the central atom nearly always attain noble gas configurations. The central atom does not.

35. Limitations of the Octet Rule for Lewis Formulas Write dot and dash formulas for BBr 3.

36. Limitations of the Octet Rule for Lewis Formulas Write dot and dash formulas for AsF 5.

37. Formal Charges It is possible to draw more than one Lewis structure with the octet rule obeyed for all the atoms. To determine which structure is most reasonable, we use formal charge.

38. Formal Charges The charge on an atom that it would have if all the atoms had the same electronegativity. To calculate formal charge, electrons are assigned as follows: All nonbonding electrons are assigned to the atom on which they are found. Half the bonding electrons are assigned to each atom in a bond.

39. Formal Charges valence electrons – (½ number of bonds + lone pair electrons) Consider:

40. Formal Charges For C: There are 4 valence electrons. In the Lewis structure there are 2 nonbonding electrons and 3 from the triple bond. There are 5 electrons from the Lewis structure. Formal charge: 4 - 5 = -1

41. Formal Charges Consider: For N: There are 5 valence electrons. In the Lewis structure there are 2 nonbonding electrons and 3 from the triple bond. There are 5 electrons from the Lewis structure. Formal charge = 5 - 5 = 0 We write:

42. Formal Charges The most stable structure has: the smallest formal charge on each atom, the most negative formal charge on the most electronegative atoms.

43. Polar and Nonpolar Covalent Bonds polar covalent bonds unequally shared electrons assymmetrical charge distribution different electronegativities

44. Electronegativity

45. Polar and Nonpolar Covalent Bonds

46. Electron density map of HF blue areas - low electron density red areas - high electron density polar molecules have separation of centers of negative and positive charge

47. Polar and Nonpolar Covalent Bonds

48. Electron density map of HI blue areas - low electron density red areas - high electron density notice that the charge separation is not as big as for HF HI is only slightly polar

49. Polar and Nonpolar Covalent Bonds

50. Continuous Range of Bonding Types all bonds have some ionic and some covalent character HI is about 17% ionic greater the electronegativity differences the more polar the bond

51. Continuous Range of Bonding Types molecules whose centers of positive and negative charge do not coincide are polar have a dipole moment dipole moment has the symbol   is the product of the distance,d, separating charges of equal magnitude and opposite sign, and the magnitude of the charge, q

52. Dipole Moments molecules that have a small separation of charge have small  molecules that have a large separation of charge have a large  example HF and HI

53. Dipole Moments in molecules some nonpolar molecules have polar bonds 2 conditions must hold for a molecule to be polar

54. Dipole Moments There must be at least one polar bond present or one lone pair of electrons. The polar bonds, if there are more than one, and lone pairs must be arranged so that their dipole moments do not cancel one another.

55. Polar and Nonpolar Covalent Bonds nonpolar covalent bonds electrons are shared equally symmetrical charge distribution must be the same element to share exactly equally H 2 N 2

56. Bond Energies Bond energy - amount of energy required to break the bond and separate the atoms in the gas phase

57. Bond Energies Table of average bond energies MoleculeBond Energy (kJ/mol) F 2 159 Cl 2 243 Br 2 192 O 2 (double bond)498 N 2 (triple bond)946

58. Bond Energies Bond energies can be calculated from other  H o 298 values

59. Bond Energies Example: Calculate the bond energy for hydrogen fluoride, HF.

60. Bond Energies Example: Calculate the bond energy for hydrogen fluoride, HF.

61. Bond Energies Example: Calculate the average N-H bond energy in ammonia, NH 3.

62. Bond Energies

63. In gas phase reactions  H o values may be related to bond energies of all species in the reaction.

64. Bond Energies Example: Use the bond energies listed in Table 8.4 to estimate the heat of reaction at 25 o C for the reaction below.

65. Bond Energies Example: Use the bond energies listed in Table 8.4 to estimate the heat of reaction at 25 o C for the reaction below.