Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Chapter 9 Models of Chemical Bonding 9-1 Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Models of Chemical Bonding 9.1 Atomic Properties and Chemical Bonds 9.2 The Ionic Bonding Model 9.3 The Covalent Bonding Model 9.4 Bond Energy and Chemical Change 9.5 Between the Extremes: Electronegativity and Bond Polarity 9-2 Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Figure 9.1 9-3 A general comparison of metals and nonmetals. Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Types of Chemical Bonding 1. Metal with nonmetal: electron transfer and ionic bonding 2. Nonmetal with nonmetal: electron sharing and covalent bonding 3. Metal with metal: electron pooling and metallic bonding 9-4 Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Figure 9.2 9-5 The three models of chemical bonding. Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Lewis Electron-Dot Symbols For main group elements The A group number gives the number of valence electrons. Place one dot per valence electron on each of the four sides of the element symbol. Pair the dots (electrons) until all of the valence electrons are used. Example: Nitrogen, N, is in Group 5A and therefore has 5 valence electrons. . N. . . . N: . . . N. : 9-6 : . : N . . Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Figure 9.4 Lewis electron-dot symbols for elements in Periods 2 and 3. 9-7 Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display. SAMPLE PROBLEM 9.1 PROBLEM: PLAN: Depicting Ion Formation Use partial orbital diagrams and Lewis symbols to depict the formation of Na+ and O2- ions from the atoms, and determine the formula of the compound. Draw orbital diagrams for the atoms and then move electrons to make filled outer levels. It can be seen that 2 sodiums are needed for each oxygen. SOLUTION: O2- Na 2s O 2 Na+ . Na 3s 3p + : O: . Na 9-8 : 2p 2Na+ +: O:2- : 2s Na 2p . 3p . 3s Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Figure 9.5 Three ways to represent the formation of Li+ and Fthrough electron transfer. Electron configurations Li 1s22s1 F 1s22s22p5 + Li+ 1s2 + F- 1s22s22p6 2s 2p Orbital diagrams Li+ Li 1s 2s 1s 2p + + F 1s 2s F1s 2p 2s Lewis electron-dot symbols : 9-9 Li+ + : F: - : + : Li . . :F: 2p Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Periodic Trends in Lattice Energy Coulomb’s Law charge A X charge B electrostatic force a distance2 energy = force X distance therefore charge A X charge B electrostatic energy a distance cation charge X anion charge electrostatic energy a 9-10 cation radius + anion radius a DH0lattice Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Figure 9.7 9-11 Trends in lattice energy. Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Figure 9.8 Electrostatic forces 9-12 and the reason ionic compounds crack. Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Figure 9.9 Solid ionic compound 9-13 Electrical conductance and ion mobility. Molten ionic compound Ionic compound dissolved in water Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Table 9.1 Melting and Boiling Points of Some Ionic Compounds Compound bp (0C) CsBr 636 1300 NaI 661 1304 MgCl2 714 1412 KBr 734 1435 CaCl2 782 >1600 NaCl 801 1413 LiF 845 1676 KF 858 1505 2852 3600 MgO 9-14 mp (0C) Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Figure 9.10 9-15 Covalent bond formation in H2. Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Figure 9.11 9-16 Distribution of electron density of H2. Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Table 9.2 back to previous slide 9-17 Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Figure 9.12 Internuclear distance (bond length) Bond length and covalent radius. Covalent radius Internuclear distance (bond length) 72 pm Internuclear distance (bond length) Covalent radius 100 pm 9-18 Covalent radius 114 pm Internuclear distance (bond length) Covalent radius 133 pm Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Table 9.3 9-19 Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display. SAMPLE PROBLEM 9.2 PROBLEM: Comparing Bond Length and Bond Strength Using the periodic table, but not Tables 9.2 and 9.3, rank the bonds in each set in order of decreasing bond length and bond strength: (a) S - F, S - Br, S - Cl PLAN: (b) C = O, C - O, C O (a) The bond order is one for all and sulfur is bonded to halogens; bond length should increase and bond strength should decrease with increasing atomic radius. (b) The same two atoms are bonded but the bond order changes; bond length decreases as bond order increases while bond strength increases as bond order increases. SOLUTION: (a) Atomic size increases going down a group. (b) Using bond orders we get Bond length: S - Br > S - Cl > S - F Bond length: C - O > C = O > C Bond strength: S - F > S - Cl > S - Br Bond strength: C 9-20 O O>C=O>C-O Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Figure 9.13 Strong forces within molecules and weak forces between them. Strong covalent bonding forces within molecules Weak intermolecular forces between molecules 9-21 Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Figure 9.14 9-22 Covalent bonds of network covalent solids. Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Figure 9.15 9-23 The infrared (IR) spectra of diethyl ether and 2-butanol. Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Figure 9.16 Using bond energies to calculate DH0rxn. Enthalpy, H DH0rxn = DH0reactant bonds broken + DH0product bonds formed BOND BREAKING DH01 = + sum of BE DH02 = - sum of BE BOND FORMATION DH0rxn 9-24 Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Figure 9.17 Using bond energies to calculate DH0rxn of methane. BOND BREAKING Enthalpy,H 4BE(C-H)= +1652kJ 2BE(O2)= + 996kJ DH0(bond breaking) = +2648kJ 2[-BE(C O)]= -1598kJ 4[-BE(O-H)]= -1868kJ DH0(bond forming) = -3466kJ DH0rxn= -818kJ 9-25 BOND FORMATION Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display. SAMPLE PROBLEM 9.3 PROBLEM: Calculating Enthalpy Changes from Bond Energies Use Table 9.2 (button at right) to calculate DH0rxn for the following reaction: CH4(g) + 3Cl2(g) PLAN: CHCl3(g) + 3HCl(g) Write the Lewis structures for all reactants and products and calculate the number of bonds broken and formed. SOLUTION: H Cl H C H + 3 Cl H C Cl + 3 H Cl H bonds broken 9-26 Cl bonds formed Cl Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display. SAMPLE PROBLEM 9.3 Calculating Enthalpy Changes from Bond Energies continued bonds broken 4 C-H bonds formed = 4 mol(413 kJ/mol) = 1652 kJ 3 C-Cl = 3 mol(-339 kJ/mol) = -1017 kJ 3 Cl-Cl = 3 mol(243 kJ/mol) = 729 kJ 1 C-H = 1 mol(-413 kJ/mol) = -413 kJ DH0bonds broken = 2381 kJ 3 H-Cl = 3 mol(-427 kJ/mol) = -1281 kJ DH0bonds formed = -2711 kJ DH0reaction = DH0bonds broken + DH0bonds formed = 2381 kJ + (-2711 kJ) = - 330 kJ 9-27 Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Table 9.4 9-28 Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Figure 9.20 9-29 The Pauling electronegativity (EN) scale. Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display. SAMPLE PROBLEM 9.4 PROBLEM: Determining Bond Polarity from EN Values (a) Use a polar arrow to indicate the polarity of each bond: N-H, F-N, I-Cl. (b) Rank the following bonds in order of increasing polarity: H-N, H-O, H-C. PLAN: (a) Use Figure 9.19(button at right) to find EN values; the arrow should point toward the negative end. (b) Polarity increases across a period. SOLUTION: (a) The EN of N = 3.0, H = 2.1; F = 4.0; I = 2.5, Cl = 3.0 N-H F-N I - Cl (b) The order of increasing EN is C < N < O; all have an EN larger than that of H. H-C < H-N < H-O 9-30 Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Figure 9.21 9-31 Electron density distributions in H2, F2, and HF. Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Figure 9.22 Boundary ranges for classifying ionic character of chemical bonds. 3.0 DEN 2.0 0.0 9-32 Copyright ©The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Figure 9.23 9-33 Properties of the Period 3 chlorides.