1. Daniel L. Reger Scott R. Goode David W. Ball http://academic.cengage.com/chemistry/reger Chapter 9 Chemical Bonds

2. Chemical Bonds Chemical bonds are the forces that hold the atoms together in substances. This chapter discusses two limiting types of bonding. Ionic bonding Covalent bonding

3. Lewis Electron-dot Symbols A Lewis electron-dot symbol consists of the symbol for the element surrounded by dots, one for each valence electron.

4. Lewis Symbols for Cations Cations of most representative elements have no valence shell electrons shown in the Lewis symbol. Na   Na + + e -  Ca   Ca 2+ + 2e -

5. Lewis Symbols for Anions The Lewis symbols of most monatomic anions show eight valence electrons. Cl + e - → Cl - Se + 2e - → Se 2-

6. Ionic Bonding Ionic bonding results from the electrostatic attraction between cations and anions. Formation of an ionic bond can be viewed as a transfer of electrons. Na + F → Na + + F - (or NaF)

7. Lattice energy is the energy required to separate one mole of ionic solid into its gaseous ions. Lattice Energy

8. The magnitude of the lattice energy is described by: where Q 1 and Q 2 are the ionic charges, and r is the distance between the ions in the solid. Lattice Energies

9. The lattice energy is strongly influenced by the charges of the ions because the charge on each ion can double or triple. Lattice Energy and Ionic Charge

10. The effect of distance (the sum of ionic radii) is usually small compared to charge effects because size differences are usually small. Lattice Energy and Distance

11. Test Your Skill Arrange the following ionic compounds in order of increasing lattice energy. MgO, NaCl, MgCl 2, CaO

12. A covalent bond result from the sharing of two electrons between two atoms, as shown here for H 2. Covalent Bonding

13. Orbital Overlap Bond length is the distance between the nuclei of two bonded atoms. At this distance, the partially occupied valence orbitals of the atoms overlap. Two electrons occupy both of the overlapping orbitals of the bonded atoms.

14. Two hydrogen atoms become more stable as their orbitals, each containing one electron, overlap. Orbital Overlap

15. Lewis Structures Lewis structures represent covalent bonding by showing how the valence electrons are present in a molecule. Bonding pairs are shared between two atoms and are represented by lines. Lone pairs are entirely on one atom and are represented by two dots. H Cl Bonding Pair Lone Pair

16. Octet Rule Octet Rule: atoms share electrons until each atom is surrounded by eight. Single Bond - sharing one pair of electrons Double Bond - sharing two pairs of electrons Triple Bond - sharing three pairs of electrons F O C O N

17. Bond Length 147 pm 125 pm 110 pm N NN NN N

18. The bond order is the number of electron pairs shared between two atoms. The skeleton structure shows which atoms are bonded to each other. A central atom is bonded to two or more other atoms. A terminal atom is bonded to only one other atom. Definitions

19. 1. Write the skeleton structure. 2. Sum the valence electrons. 3. Subtract two electrons for each bond in the skeleton structure. 4. Count the number of electrons needed to satisfy octet rule for each atom. If the number of electrons needed equals the number remaining, go to 5. If fewer electron remain, add one bond for every two additional electrons needed. 5. Place remaining electrons as lone pairs to satisfy the octet rule for each atom (not H). Writing Lewis Structures

20. Write the Lewis structure of fromaldehyde, H 2 CO.

21. Test Your Skill Write the Lewis structure of N 2 H 2.

22. In I 2 the sharing of the electrons in the covalent bond is equal; in ClF it is not. Dipole moment is a measure of the unequal sharing of electrons. The unequal sharing leads to a polar covalent bond that is indicated with the symbol  followed by a sign to show partial charges.   Cl-F Bond Polarity

23. Measuring Dipole Moments Molecules oriented randomly when field is off. Molecules oriented with field when field is on.

24. Electronegativity is a measure of the ability of an atom to attract the shared electrons in a chemical bond. Electronegativity

25. Electronegativity Trends

27. Example: Electronegativity Select the most polar bond. Cl-FO-FP-F

28. Properties of Compounds

29. Formal charge is a charge assigned to atoms in Lewis structures by assuming the shared electrons are divided equally between the bonded atoms. Complete Lewis structure of CO with formal charges (shown in circles) is: Formal Charges C O → C - + O + C O

30. Formal Charges Add formal charges to the Lewis structure of HNO 3 shown below. O N O O H

31. Test Your Skill Add formal charges to the Lewis structure of HNO 3 shown below. O N O O H

32. Lewis structures that show the smallest formal charges are favored. Lewis structures that have adjacent atoms with formal charges of the same sign are much less favorable. Lewis structures that place negative formal charges on the more electronegative atoms are favored. Formal charges of opposite sign are usually on adjacent atoms. Structure Stability

33. Test Your Skill Of the two structures shown for HNO 3, use the stability rules to predict which will be more favored. O N O O H O O H

34. Resonance in Lewis Structures Resonance structures differ only in the distribution of the valence electrons. All resonance structures follow the rules for writing Lewis structures. Resonance structures are indicated by a double headed arrow. O N O O H O O H ↔

35. Draw the third possible resonance structure for HNO 3, the first two are below. Drawing Resonance Structures O N O O H O O H ↔

36. HNO 3 Resonance Forms A total of three resonance forms can be written for HNO 3. The first and last structures are equally favored because of fewer formal charges; the middle structure less favored. O N O O H O O H ↔ ↔ O O H

37. Test Your Skill Write all resonance structures, including formal charges, for O 3

38. No resonance structure is correct by itself; the correct structure is an average of all resonance structures. Average Structure

39. Contribution of Resonance Structures Equivalent resonance structures, such as the two for O 3, contribute equally to the average structure. Bond order in O 3 is the average of a double bond and a single bond = 1.5. O ↔ O O OO O

40. Test Your Skill (a) Draw all resonance structures of N 2 O. (b) How much does each contribute to the final average structure?

41. Central atoms from Groups 2A and 3A do not have enough valence electrons to complete an octet - the Lewis structures are electron-deficient. Central beryllium atom can have only two bonds because at total of 4 electrons is available. Electron Deficient Molecules H Be H

42. Any species with an odd number of electrons must violate the octet rule. Write Lewis structures for odd electron species with 7 electrons around one atom. (radicals) Resonance structure on left is favored since it has no formal charges. Odd Electron Molecules N O ↔ N O

43. Expanded Valence Shell Molecules Expanded valence shell molecules have more than eight electrons about an atom in a Lewis structure. A large class of expanded valence shell compounds have the general formula YF n where Y = P, S, Cl, As, Se, Br, Te, I, or Xe.

44. For expanded valence shell species, place the extra electrons around the central atom (the central atom must be from the third or later row of the periodic table). The sulfur has an expanded valence shell of 10 electrons. Lewis Structure of SF 4 S F F F F

45. Example: Expanded Valence Shell Draw the Lewis structure of XeF 2.

46. Test Your Skill Draw the Lewis structure of IF 3.

47. Many oxides of the third and later period elements have expanded valence shell structures that minimize formal charges, for example ClO 2 -. The two expanded valence shell structures on the right are favored because of fewer formal charges. Oxides of Heavier Elements Cl OO ↔ OO ↔ OO

48. Bond Energies The bond dissociation energy or bond energy (D) is the energy required to break one mole of bonds in a gaseous species. H 2 (g)  H(g) + H(g)  H = D(H-H) = 436 kJ/mol Bond energies are always endothermic; it takes energy to break a bond.

49. Bond energies, except those of diatomic molecules, are average energies. Bond energies can be used to calculate approximate enthalpies of reaction.  H reaction =  (bond energies of bonds broken) -  (bond energies of bonds formed) Lewis structures need to be written to determine the types and numbers of bonds broken and formed. Enthalpies of Reaction

50. Enthalpy Diagram

51. Example: Enthalpy Change Calculate an approximate enthalpy change for the reaction 2H 2 (g) + O 2 (g)  2H 2 O(g) Bond Energies H-H 436 kJ/mol O=O 498 kJ/mol O-H 463 kJ/mol

52. Test Your Skill Calculate an approximate enthalpy change for the reaction C 2 H 4 (g) + H 2 (g)  C 2 H 6 (g) Bond Energies C-H 414 kJ/mol H-H 436 kJ/mol C-C 348 kJ/mol C=C 611 kJ/mol