Chapter Eight. p328. Bonding: General Concepts

Transcription

1 Chapter Eight p328 Bonding: General Concepts 1

2 Contents

3 8-1 Types of Chemical Bonds p330 Coulomb s law The energy of interaction between a pair of ions can be calculated using Coulomb s law: E 19 Q1Q 2 ( J nm )( r ) where E has units joules, r is the distance between the ion centers in nanometers, and Q 1 and Q 2 are the numerical ion charges. 3

4 p330 For example, the distance between the centers of the Na + and Cl - ions is nm, and the ionic energy pair of ions is E ( J ( 1)( 1) nm )[ ] nm 19 J

5 Questions to Consider p331 What is meant by the term chemical bond? Why do atoms bond with each other to form molecules? 5

6 p331 Figure 8.1 (a) The interaction of two hydrogen atoms.

7 p331 Bond length Figure 8.1(b) Energy profiles as a function between the hydrogen atoms. As the atoms approach each other (right side of graph), the energy decreases until the distances reaches nm and then begins to increase again due to repulsions.

8 Key ideas in bonding p332 Ionic Bonding: Electrons are transferred Covalent Bonding: Electrons are shared equally What about intermediate cases? Polar covalent bond: H F

9 React 1 Chemical bond p333 What is meant by the term chemical bond? Why do atoms bond with each other to form molecules? How do atoms bond with each other to form molecules? 9

10 p Electronegativity Linus Pauling( ) Expected H-X bonding energy = ½ (H-H bond energy + X-X bond energy) = (H-X) act - (H-X) exp If X has a greater electronegativity than H, the shares electron(s) will tend to be closer to the X atom. The molecule will be polar, with charge distribution. The greater is the difference in the electronegativities of the atoms, the greater is the ionic compound and the greater is the value of. 10

11 Electronegativity: the ability of an atom in a molecule to attract shared electrons to itself. p333 Figure 8.2 The effect of an electric field on hydrogen fluoride molecules

12 Polar molecules 12

13 The Pauling electronegativity values p334 Figure 8-3 The Pauling electronegativity values. Electronegativity generally increases across a period and decreases down a group. 13

14 React 2 The general trend for electronegativity p334 What is the general trend for electronegativity across rows and down columns on the periodic table? Explain the trend. 14

15 Ex 8.1 Relative Bond Polarities P335 Order the following bonds according to polarity: H H, O H, Cl H, S H, and F H. 15

16 Table 8.1 The Relationship Between Electronegativity and Bond Type p335 16

17 p Bond polarity and dipole moments 17

18 p336

19 p336

20 p337

21 Ex 8.2 Bond Polarity and Dipole Moment P337 For each of the following molecules, show the direction of the bond polarities and indicate which ones have a dipole moment: HCl, Cl 2, SO 3 (a planar molecule with the oxygen atoms spaced evenly around the central sulfur atom), CH 4 [trtrahedral(see Table 8.2) with the carbon atom at the center], and H 2 S (V-shaped with the sulfur atom at the point). 21

22 Solution: (a) HCl (b) Cl 2 (c) SO 3

23 (d) CH 4 p338 (e) H 2 S

24 8-4 Ions: Electron configurations and sizes p338 Predicting Formulas of Ionic Compounds

25 Sizes of ions p340 Table 8.3 Common ions with noble gas configuration in ionic compounds

26 p341

27 React 3 Choose an alkali metal, an alkaline metal, a noble gas, and a halogen so that they constitute an isoelectronic series when the metals and halogen are written as their most stable ions. What is the electron configuration for each species? Determine the number of electrons for each species. Determine the number of protons for each species. Rank the species according to increasing radius. Rank the species according to increasing ionization energy.

28 Ionic radii 28

29 What we can read from the periodic table: Trends for Atomic size Ion radius Ionization energy Electronegativity Electron configurations Predicting formulas for ionic compounds Ranking polarity of bonds 29

30 Ex 8.4 Relative Lon Size II Choose the largest ion in each of the following groups. P342 a.li +, Na +, K +, Rb +, Cs + b.ba 2+, Cs +, I -, Te 2-

31 Ex 8.3 Relative Lon Size I P342 Arrange the ions Se 2-, Br -, Rb +, and Sr 2+ in order of decreasing size. 31

32 8-5 Energy effects in binary ionic compounds p342

33 p344

34 Formation of an ionic solid p Sublimation of the solid metal M(s) M(g) [endothermic] (For Li(s) is +161 kj.) 2. Ionization of the metal atoms M(g) M + (g) + e [endothermic] (For Li(g) is +520 kj) 3. Dissociation of the nonmetal 1 /2X 2 (g) X(g) [endothermic] (For F is +½ (154 kj) 34

35 Formation of an ionic solid (continued) 4. Formation of X ions in the gas phase: X(g) + e X (g) [exothermic] (For F - is -328 kj/mole) 5. Formation of the solid MX: M + (g) + X (g) MX(s) [quite exothermic] (Corresponding to the lattice energy for LiF, which is kj./mole) 35

36 p345 Figure 8.11 Comparing energy changes

37 Born-Haber cycle for NaCl 37

38 p344

39 Lattice Energy Calculations p344

40 8.6 Partial ionic character of covalent bonds p346 40

41 The relationship between the ionic character of a covalent bond and the electronegativity difference of the bonded atoms. p347 Figure 8.13

42 8-7 The Covalent Chemical Bond: p347 A Model Models Models are attempts to explain how nature operates on the microscopic level based on experiences in the macroscopic world. The Localized Electron Bonding Model 42

43 Fundamental Properties of Models p A model does not equal reality. 2. Models are oversimplifications, and are therefore often wrong. 3. Models become more complicated as they age. 4. We must understand the underlying assumptions in a model so that we don t misuse it. 43

44 p348

45 8-8 Covalent Bond Energies and Chemical Reactions p350 45

46 p351

47 Ex 8.5 H from Bond Energies p352 Using the bond energies listed in Table 8.4, calculate H for the reaction of methane with chlorine and fluorine to give Freon-12(CF 2 Cl 2 ). CH 4( g) 2 Cl2( g) 2 F2 ( g) CF2Cl 2( g) 2 HF( g) 2 HCl( g)

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49 8-9 The Localized Electron Bonding Model p354 A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms. 49

50 Localized Electron Model 1. Description of valence electron arrangement (Lewis structure). 2. Prediction of geometry (VSEPR model). 3. Description of atomic orbital types used to share electrons or hold long pairs. 50

51 p Lewis Structure Shows how valence electrons are arranged among atoms in a molecule. Reflects central idea that stability of a compound relates to noble gas electron configuration. 51

52 Lewis Structures p Sum the valence electrons. 2. Place bonding electrons between pairs of atoms. 3. Atoms usually have noble gas configurations. 52

53 Ex 8.6 Writing Lewis Structures P357 Give the Lewis structure for each of the following. a. HF, b. N 2, c. NH 3, d. CH 4, e. CF 4, f. NO +

54 p357

55 8-11 Exceptions to the Octet Rule p358

56 Ex 8.7 Lewis Structures for Molecules That Violate the Octet Rule I Write the Lewis structure for PCl 5. P360

57 Ex 8.8 Lewis Structures for Molecules That Violate the Octet Rule II P361 Write the Lewis structure for each molecule or ion. a. ClF 3 b. XeO 3 c. RnCl 2 d. BeCl 2 e. ICl 4-57

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59 8-12 Resonance p362

60 Ex 8.9 Resonance Structures P363 Describe the electron arrangement in the nitrite anion (NO 2- ) using the localized electron model.

61 Rules Governing Formal Charge p366

62 Ex 8.10 Formal Charges Give possible Lewis structures for XeO 3, an explosive P366 compound of xenon. Which Lewis structure or structures are most appropriate according to the formal charges?

63 8-13 Molecular Structure: The VSEPR Model p367 VSEPR Model The structure around a given atom is determined principally by minimizing electron pair repulsions. 63

64 p369

65 Predicting a VSEPR Structure 1. Draw Lewis structure. 2. Put pairs as far apart as possible. 3. Determine positions of atoms from the way electron pairs are shared. 4. Determine the name of molecular structure from positions of the atoms. 65

66 Ex 8.11 Prediction of Molecular Structure I Describe the molecular structure of the water molecule. P369 Solution The Lewis structure for water is There are four pairs of electrons: two bonding pairs and two nonbonding pairs. To minimize repulsions, these best arrangement in a tetrahedral array, as shown in Fig Figure 8.17

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68 p370

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70 p371

71 Ex 8.12 Prediction of Molecular Structure II When phosphorus reacts with excess chlorine gas, the compound phosphorus pentachloride (PCl 5 ) is formed. In the gaseous and liquid states, this substance consists of PCl 5 molecules, but in the solid state it consists of a 1 : 1 mixture of PCl 4+ and PCl 6- ions. Predict the geometric structures of PCl 5, PCl 4+, and PCl 6-. P373

72 Solution The Lewis structure for PCl 5 is shown. Five pairs of electrons around the phosphorous atom require a trigonal bipyramidal arrangement (see Table 8.6). p373 The Lewis structure for the PCl 4+ ions (5+4(7) -1 = 32 valence electrons) is shown. There are four pairs of electrons surrounding the phosphorus atom in the PCl 4+ ion, which requires a tetrahedral arrangement of the pairs. The Lewis structure for PCl 6- (5 + 6(7) + 1 = 48 valence electrons) is shown. Since each electron pair is shared with a chlorine atom, an octahedral PCl 6- anion is predicted.

73 Ex 8.13 Prediction of Molecular Structure III P373 Because the noble gases have filled s and p valence orbitals, they were not expected to be chemically reactive. In fact, for many years these elements were called insert gases because of this supposed inability to form any compounds. However. In the early 1960s several compounds of krypton, xenon, and radon were synthesized. For example, a team at the Argonne National Laboratory produced the stable colorless compound xenon tetrafluoride (XeF 4 ). Predict its structure and whether it has a dipole moment. 73

74 Solution The Lewis structure for XeF 4 is p374 The xenon atom in this molecule is surrounded by six pairs of electrons, which means an octahedral arrangement. The arrangement in Fig. 8.20(b) is preferred, and the molecular structure is predicted to be square planar. There is an octahedral arrangement of electron pairs, but the atoms form a square planar structure. Although each Xe-F bond is polar, their structure causes the polarities to cancel. Thus XeF 4 has no dipole moment as shown in the margin.

75 Ex 8.14 Structures of Molecules with Multiple Bonds Predict the molecular structure of the sulfur dioxide P376 molecule. Is this molecule expected to have a dipole moment? Solution We must determine the Lewis structure for the SO 2 molecule, which has 18 valence electrons. The expected resonance structures are The structure of the SO 2 molecule expected to be V- shaped, with a 120-degree bond angle. The molecule has a dipole moment as shown:

76 Molecules Containing No Single Central Atom p377

77 p379 The VSEPR Model- How Well Does It Work?

78 VSEPR 78

79 VSEPR: Two Electron Pairs 79

80 VSEPR: Three Electron Pairs 80

81 VSEPR: Four Electron Pairs 81

82 VSEPR: Iodine Pentafluoride 82