1. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 1 Chemical Bonds Chapter Nine

2. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 2 Forces called chemical bonds hold atoms together in molecules and keep ions in place in solid ionic compounds. Chemical bonds are electrostatic forces; they reflect a balance in the forces of attraction and repulsion between electrically charged particles. Chemical Bonds: A Preview

3. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 3 Electrostatic Attractions and Repulsions Nuclei repel other nuclei Nuclei attract electron(s) Electrons repel other electrons Both attractions and repulsions occur; what’s the net effect?? Answer: it depends on the distance between nuclei …

4. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 4 Energy of Interaction Closer together; attraction increases. At 74 pm, attractive forces are at a maximum, energy is at a minimum. Closer than 74 pm, repulsion increases. Hydrogen nuclei far apart; just a little attraction.

5. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 5 Figure 13.1: (a) The interaction of two hydrogen atoms (b) Energy profile as a function of the distance between the nuclei of the hydrogen atoms.

6. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 6 Valence electrons play a fundamental role in chemical bonding. When metals and nonmetals combine, valence electrons usually are transferred from the metal to the nonmetal atoms, giving rise to ionic bonds. In combinations involving only nonmetals, one or more pairs of valence electrons are shared between the bonded atoms, producing covalent bonds. In losing, gaining, or sharing electrons to form chemical bonds, atoms tend to acquire the electron configurations of noble gases. The Lewis Theory of Chemical Bonding: An Overview

7. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 7 In a Lewis symbol, the chemical symbol for the element represents the nucleus and core electrons of the atom. Dots around the symbol represent the valence electrons. In writing Lewis symbols, the first four dots are placed singly on each of the four sides of the chemical symbol. Dots are paired as the next four are added. Lewis symbols are used primarily for those elements that acquire noble-gas configurations when they form bonds. Lewis Symbols

8. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 8 Example 9.1 Give Lewis symbols for magnesium, silicon, and phosphorus.

9. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 9 When atoms lose or gain electrons, they may acquire a noble gas configuration, but do not become noble gases. Because the two ions formed in a reaction between a metal and a nonmetal have opposite charges, they are strongly attracted to one another and form an ion pair. The net attractive electrostatic forces that hold the cations and anions together are ionic bonds. The highly ordered solid collection of ions is called an ionic crystal. Ionic Bonds and Ionic Crystals

10. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 10 Formation of a Crystal of Sodium Chloride Sodium reacts violently in chlorine gas. Na donates an electron to Cl … … and opposites attract.

11. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 11 A computer representation of K 3 C 60, a superconducting substance formed by reacting potassium with buckminster fullerine (C 60 ) Source: Photo Researchers, Inc.

12. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 12 Lewis symbols can be used to represent ionic bonding between nonmetals and: the s-block metals, some p-block metals, and a few d-block metals. Instead of using complete electron configurations to represent the loss and gain of electrons, Lewis symbols can be used. Using Lewis Symbols to Represent Ionic Bonding

13. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 13 Example 9.2 Use Lewis symbols to show the formation of ionic bonds between magnesium and nitrogen. What are the name and formula of the compound that results?

14. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 14 From the data above, it doesn’t appear that the formation of NaCl from its elements is energetically favored. However … … the enthalpy of formation of the ionic compound is more important than either the first ionization energy or electron affinity. The overall enthalpy change can be calculated using a step- wise procedure called the Born–Haber cycle. Energy Changes in Ionic Compound Formation Na(g)  Na + (g) + e – I 1 = +496 kJ/mol Cl(g) + e –  Cl – (g)EA 1 = –349 kJ/mol

15. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 15 The Born–Haber cycle is a hypothetical process, in which ΔH f is represented by several steps. What law can be used to find an enthalpy change that occurs in steps?? Energy Changes in Ionic Compound Formation (cont’d) ΔH f ° for NaCl is very negative because … … ΔH 5 —the lattice energy—is very negative.

16. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 16 Example 9.3 Use the following data to determine the lattice energy of MgF 2 (s): enthalpy of sublimation of magnesium, +146 kJ/mol; I 1 for Mg, +738 kJ/mol; I 2 for Mg, +1451 kJ/mol; bond-dissociation energy of F 2 (g), +159 kJ/mol F 2 ; electron affinity of F, –328 kJ/mol F; enthalpy of formation of MgF 2 (s), –1124 kJ/mol.

17. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 17 A Lewis structure is a combination of Lewis symbols that represents the formation of covalent bonds between atoms. In most cases, a Lewis structure shows the bonded atoms with the electron configuration of a noble gas; that is, the atoms obey the octet rule. (H obeys the duet rule.) The shared electrons can be counted for each atom that shares them, so each atom may have a noble gas configuration. Lewis Structures of Simple Molecules

18. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 18 The shared pairs of electrons in a molecule are called bonding pairs. In common practice, the bonding pair is represented by a dash (—). The other electron pairs, which are not shared, are called nonbonding pairs, or lone pairs. Lewis Structures (cont’d) Each chlorine atom sees an octet of electrons.

19. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 19 Some Illustrative Compounds Note that the two-dimensional Lewis structures do not necessarily show the correct shapes of the three-dimensional molecules. Nor are they intended to do so. The Lewis structure for water may be drawn with all three atoms in a line: H–O–H. We will learn how to predict shapes of molecules in Chapter 10.

20. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 20 The covalent bond in which one pair of electrons is shared is called a single bond. Multiple bonds can also form: Multiple Covalent Bonds In a double bond two pairs of electrons are shared. Note that each atom obeys the octet rule, even with multiple bonds. In a triple bond three pairs of electrons are shared.

21. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 21 The Lewis structure commonly drawn for oxygen is But oxygen is paramagnetic, and therefore must have unpaired electrons. Lewis structures are a useful tool, but they do not always represent molecules correctly, even when the Lewis structure is plausible. The Importance of Experimental Evidence

22. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 22 Liquid oxygen poured into the space between the poles of a strong magnet remains there until it boils away. Source: Donald Clegg

23. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 23 Electronegativity (EN) is a measure of the ability of an atom to attract its bonding electrons to itself. EN is related to ionization energy and electron affinity. The greater the EN of an atom in a molecule, the more strongly the atom attracts the electrons in a covalent bond. Polar Covalent Bonds and Electronegativity Electronegativity generally increases from left to right within a period, and it generally increases from the bottom to the top within a group.

24. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 24 Pauling’s Electronegativities It would be a good idea to remember the four elements of highest electronegativity: N, O, F, Cl.

25. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 25 Figure 13.3: The Pauling electronegativity values as updated by A.L. Allred in 1961. (cont’d)

26. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 26 Example 9.4 Referring only to the periodic table inside the front cover, arrange the following sets of atoms in the expected order of increasing electronegativity. (a) Cl, Mg, Si (b) As, N, Sb (c) As, Se, Sb

27. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 27 Identical atoms have the same electronegativity and share a bonding electron pair equally. The bond is a nonpolar covalent bond. When electronegativities differ significantly, electron pairs are shared unequally. The electrons are drawn closer to the atom of higher electronegativity; the bond is a polar covalent bond. With still larger differences in electronegativity, electrons may be completely transferred from metal to nonmetal atoms to form ionic bonds. Electronegativity Difference and Bond Type

28. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 28 Electronegativity Difference and Bond Type No sharp cutoff between ionic and covalent bonds. C—H bonds are virtually nonpolar. Cs—F bonds are “so polar” that we call the bonds ____.

29. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 29 Depicting Polar Covalent Bonds In nonpolar bonds, electrons are shared equally. Unequal sharing in polar covalent bonds. Polar bonds are often depicted using colors to show electrostatic potential (blue = positive, red = negative). Polar bonds are also depicted by partial positive and partial negative symbols … … or with a cross-based arrow pointing to the more electronegative element.

30. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 30

31. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 31

32. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 32 Figure 13.4: (a) The charge distribution in the water molecule. (b) The water molecule in an electric field.

33. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 33 The polar bonds result in dipole moment

34. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 34 The HCL molecule has a dipole moment

35. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 35 Sulfur has a partial positive charge

36. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 36 Hydrogen atoms and a small partial negative charge on the carbon

37. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 37 Hydrogen atoms have a partial positive charge

38. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 38 Example 9.5 Use electronegativity values to arrange the following bonds in order of increasing polarity: Br—Cl, Cl—Cl, Cl—F, H—Cl, I—Cl

39. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 39 The skeletal structure shows the arrangement of atoms. Hydrogen atoms are terminal atoms (bonded to only one other atom). The central atom of a structure usually has the lowest electronegativity. In oxoacids (HClO 4, HNO 3, etc.) hydrogen atoms are usually bonded to oxygen atoms. Molecules and polyatomic ions usually have compact, symmetrical structures. Writing Lewis Structures: Skeletal Structures

40. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 40 1.Determine the total number of valence electrons. 2.Write a plausible skeletal structure and connect the atoms by single dashes (covalent bonds). 3.Place pairs of electrons as lone pairs around the terminal atoms to give each terminal atom (except H) an octet. 4.Assign any remaining electrons as lone pairs around the central atom. 5.If necessary (if there are not enough electrons), move one or more lone pairs of electrons from a terminal atom to form a multiple bond to the central atom. Writing Lewis Structures: A Method

41. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 41 Example 9.6 Write the Lewis structure of nitrogen trifluoride, NF 3. Example 9.7 Write a plausible Lewis structure for phosgene, COCl 2. Example 9.8 Write a plausible Lewis structure for the chlorate ion, ClO 3 –.

42. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 42 Formal charge is the difference between the number of valence electrons in a free (uncombined) atom and the number of electrons assigned to that atom when bonded to other atoms in a Lewis structure. Formal charge is a hypothetical quantity; a useful tool. Usually, the most plausible Lewis structure is one with no formal charges. When formal charges are required, they should be as small as possible. Negative formal charges should appear on the most electronegative atoms. Adjacent atoms in a structure should not carry formal charges of the same sign. Formal Charge

43. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 43 Formal Charge Illustrated

44. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 44 Example 9.9 In Example 9.7, we wrote a Lewis structure for the molecule COCl 2, shown here as structure (a). Show that structure (a) is more plausible than (b) or (c).

45. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 45 When a molecule or ion can be represented by two or more plausible Lewis structures that differ only in the distribution of electrons, the true structure is a composite, or hybrid, of them. The different plausible structures are called resonance structures. The actual molecule or ion that is a hybrid of the resonance structures is called a resonance hybrid. Electrons that are part of the resonance hybrid are spread out over several atoms and are referred to as being delocalized. Resonance: Delocalized Bonding Three pairs of electrons are distributed among two bonds.

46. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 46 Example 9.10 Write three equivalent Lewis structures for the SO 3 molecule that conform to the octet rule, and describe how the resonance hybrid is related to the three structures.

47. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 47 Molecules with an odd number of valence electrons have at least one of them unpaired and are called free radicals. Some molecules have incomplete octets. These are usually compounds of Be, B, or Al; they generally have some unusual bonding characteristics, and are often quite reactive. Some compounds have expanded valence shells, which means that the central atom has more than eight electrons around it. A central atom can have expanded valence if it is in the third period or lower (i.e., S, Cl, P). Molecules that Don’t Follow the Octet Rule

48. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 48 Example 9.11 Write the Lewis structure for bromine pentafluoride, BrF 5. Example 9.12 A Conceptual Example Indicate the error in each of the following Lewis structures. Replace each by a more acceptable structure(s).

49. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 49 Bond order is the number of shared electron pairs in a bond. A single bond has BO = 1, a double bond has BO = 2, etc. Bond length is the distance between the nuclei of two atoms joined by a covalent bond. Bond length depends on the particular atoms in the bond and on the bond order. Bond Order and Bond Length

50. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 50 Example 9.13 Estimate the length of (a) the nitrogen-to-nitrogen bond in N 2 H 4 and (b) the bond in BrCl.

51. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 51 Bond-dissociation energy (D) is the energy required to break one mole of a particular type of covalent bond in a gas-phase compound. Energies of some bonds can differ from compound to compound, so we use an average bond energy. Bond Energy The H—H bond energy is precisely known … … while the O—H bond energies for the two bonds in H 2 O are different.

52. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 52

53. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 53 The higher the order (for a particular type of bond), the shorter and the stronger (higher energy) the bond. A N=N double bond is shorter and stronger than a N–N single bond. There are four electrons between the two positive nuclei in N=N. This produces more electrostatic attraction than the two electrons between the nuclei in N–N. Trends in Bond Lengths and Energies

54. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 54 For the reaction N 2 (g) + 2 H 2 (g)  N 2 H 4 (g) to occur … Calculations Involving Bond Energies … plus 2(436 kJ), to break bonds of the reactants. … we must supply 946 kJ … When the bonds of the product form, 163 kJ plus 4(389 kJ) of energy is liberated. ΔH = (+946 kJ) + 2(+436 kJ) + (–163 kJ) + 4(–389 kJ)

55. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 55 Example 9.14 Use bond energies from Table 9.1 to estimate the enthalpy of formation of gaseous hydrazine. Compare the result with the value of ΔH° f [N 2 H 4 (g)] from Appendix C.

56. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 56 Hydrocarbons with double or triple bonds between carbon atoms are called unsaturated hydrocarbons. Alkenes are hydrocarbons with one or more C=C double bonds. Simple alkenes have just one double bond in their molecules. The simplest alkene is C 2 H 4, ethene (ethylene). Alkynes are hydrocarbons that have one or more carbon– carbon triple bonds. The simplest alkyne is C 2 H 2, ethyne (acetylene). Alkenes and Alkynes

57. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 57 Molecular Models of Ethene and Ethyne

58. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 58 Polymers are compounds in which many identical molecules have been joined together. Monomers are the simple molecules which join together to form polymers. Often, the monomers have double or triple bonds. The process of these molecules joining together is called polymerization. Many everyday products and many biological compounds are polymers. Polymers

59. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 59 Formation of Polyethylene Another ethylene molecule adds to a long chain formed from more ethylene molecules.

60. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 60 The plastics shown here were manufactured with ethylene. Source: Comstock - Mountainside, NJ

61. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 61

62. Hall © 2005 Prentice Hall © 2005 General Chemistry 4 th edition, Hill, Petrucci, McCreary, Perry Chapter Nine 62 Cumulative Example Use data from this chapter and elsewhere in the text to estimate a value for the standard enthalpy of formation at 25 °C of HNCO(g). Assess the most likely source of error in this estimation.