1. Chapter 8 Concepts of Chemical Bonding
Chemistry, The Central Science, 11th edition
Theodore L. Brown, H. Eugene LeMay, Jr., and Bruce E. Bursten
Chapter 8 Concepts of Chemical Bonding
John D. Bookstaver
St. Charles Community College
Cottleville, MO
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2. Overview – Chapter 8 Basics of Chemical Bonding
New material:
lattice energy
quantify bond polarity using dipole
moments
relate bond lengths, electronegativity
differences and dipole moments in
molecules
use of formal charge to determine
most likely Lewis structure
resonance structures
bond energies (enthalpies) and
their use in determining DHrxn and
their relation to bond length
Review:
types of chemical bonds – ionic, covalent, metallic
relate bond polarity and the type of bond to differences in electronegativity
Lewis structures of covalently-bonded molecules and polyatomic ions, including exceptions to the octet rule
predict formation of single, double and triple bonds
name molecules
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3. Chemical Bonds Three basic types of bonds Ionic Covalent Metallic
Electrostatic attraction between ions
Covalent
Sharing of electrons
Metallic
Metal atoms bonded to several other atoms
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4. Ionic Bonding
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5. Energetics of Ionic Bonding
As we saw in the last chapter, it takes 495 kJ/mol to remove electrons from sodium.
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6. Energetics of Ionic Bonding
We get 349 kJ/mol back by giving electrons to chlorine.
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7. Energetics of Ionic Bonding
But these numbers don’t explain why the reaction of sodium metal and chlorine gas to form sodium chloride is so exothermic!
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8. Energetics of Ionic Bonding
There must be a third piece to the puzzle.
What is as yet unaccounted for is the electrostatic attraction between the newly-formed sodium cation and chloride anion.
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9. Lattice Energy This third piece of the puzzle is the lattice energy:
The energy required to completely separate a mole of a solid ionic compound into its gaseous ions.
The energy associated with electrostatic interactions is governed by Coulomb’s law:
Eel = 
Q1Q2 d
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10. Lattice Energy
Lattice energy, then, increases with the charge on the ions.
It also increases with decreasing size of ions.
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11. Sample Exercise 8.1 (p. 303)
Without using Table 8.2 (previous slide), arrange the following ionic compounds in order of increasing lattice energy: NaF, CsI, CaO.
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12. Practice Exercise 1 (8.1)
Without looking at Table 8.2, predict which one of the following orderings of lattice energy is correct for these ionic compounds.
NaCl > MgO > CsI > ScN
ScN > MgO > NaCl > CsI
NaCl > CsI > ScN > CaO
MgO > NaCl > ScN > CsI
ScN > CsI > NaCl > MgO
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13. Practice Exercise 1 (8.1)
Without looking at Table 8.2, predict which one of the following orderings of lattice energy is correct for these ionic compounds.
NaCl > MgO > CsI > ScN
ScN > MgO > NaCl > CsI
NaCl > CsI > ScN > CaO
MgO > NaCl > ScN > CsI
ScN > CsI > NaCl > MgO
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14. Practice Exercise 2 (8.1)
Which substance would you expect to have the greatest lattice energy? MgF2, CaF2 or ZrO2?
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15. Sample Exercise 8.2 (p. 306)
Predict the ion generally formed by each of the following atoms:
a) Sr
b) S
c) Al
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16. Practice Exercise 1 (8.2)
Which of these elements is most likely to form ions with a 2+ charge? Li Ca O P Cl
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17. Practice Exercise 1 (8.2)
Which of these elements is most likely to form ions with a 2+ charge? Li Ca O P Cl
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18. Practice Exercise 2 (8.2)
Predict the charges on the ions formed when magnesium reacts with nitrogen.
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19. Energetics of Ionic Bonding
By accounting for all three energies (ionization energy, electron affinity, and lattice energy), we can get a good idea of the energetics involved in such a process.
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20. Energetics of Ionic Bonding
These phenomena also helps explain the “octet rule.”
Metals, for instance, tend to stop losing electrons once they attain a noble gas configuration because energy would be expended that cannot be overcome by lattice energies.
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21. Covalent Bonding In covalent bonds atoms share electrons.
There are several electrostatic interactions in these bonds:
Attractions between electrons and nuclei
Repulsions between electrons
Repulsions between nuclei
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22. Sample Exercise 8.3 (p. 308)
Given the Lewis symbols for the elements nitrogen and fluorine shown in Table 8.1, predict the formula of the stable binary compound (a compound composed of two elements) formed when nitrogen reacts with fluorine, and draw its Lewis structure.
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23. Practice Exercise 1 (8.3)
Which of these molecules has the same number of shared electron pairs as unshared electron pairs?
HCl
H2S
PF3
CCl2F2
Br2
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24. Practice Exercise 1 (8.3)
Which of these molecules has the same number of shared electron pairs as unshared electron pairs?
HCl
H2S
PF3
CCl2F2
Br2
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25. Practice Exercise 2 (8.3)
Compare the Lewis symbol for neon with the Lewis structure for methane, CH4.
In what important way are the electron arrangements about neon and carbon alike?
In what important respect are they different?
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26. Polar Covalent Bonds
Though atoms often form compounds by sharing electrons, the electrons are not always shared equally.
Fluorine pulls harder on the electrons it shares with hydrogen than hydrogen does.
Therefore, the fluorine end of the molecule has more electron density than the hydrogen end.
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27. Electronegativity
Electronegativity is the ability of atoms in a molecule to attract electrons to themselves.
On the periodic chart, electronegativity increases as you go…
…from left to right across a row.
…from the bottom to the top of a column.
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28. Sample Exercise 8.4 (p. 311) Which bond is more polar:
a) B-Cl or C-Cl?
b) P-F or P-Cl?
Indicate in each case which atom has the partial negative charge.
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29. Practice Exercise 1 (8.4)
Which of the following bonds is the most polar?
H-F
H-I
Se-F
N-P
Ga-Cl
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30. Practice Exercise 1 (8.4)
Which of the following bonds is the most polar?
H-F
H-I
Se-F
N-P
Ga-Cl
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31. Practice Exercise 2 (8.4) Which of the following bonds is most polar:
S-Cl, S-Br, Se-Cl, or Se-Br?
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32. Polar Covalent Bonds details not required
When two atoms share electrons unequally, a bond dipole results.
The dipole moment, , produced by two equal but opposite charges separated by a distance, r, is calculated:
 = Qr
It is measured in debyes (D).
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33. Polar Covalent Bonds
The greater the difference in electronegativity, the more polar is the bond.
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34. Lewis Structures
Lewis structures are representations of molecules showing all electrons, bonding and nonbonding.
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35. Writing Lewis Structures
Find the sum of valence electrons of all atoms in the polyatomic ion or molecule.
If it is an anion, add one electron for each negative charge.
If it is a cation, subtract one electron for each positive charge.
PCl3
(7) = 26
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36. Writing Lewis Structures
The central atom is the least electronegative element that isn’t hydrogen. Connect the outer atoms to it by single bonds.
Keep track of the electrons:
= 20
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37. Writing Lewis Structures
Fill the octets of the outer atoms.
Keep track of the electrons:
= 20; = 2
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38. Writing Lewis Structures
Fill the octet of the central atom.
Keep track of the electrons:
= 20; = 2; = 0
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39. Writing Lewis Structures
If you run out of electrons before the central atom has an octet…
…form multiple bonds until it does.
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40. Sample Exercise 8.6 (p. 316)
Draw the Lewis structure for phosphorus trichloride, PCl3.
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41. Practice Exercise 1 (8.6)
Which of these molecules has a Lewis structure with a central atoms having no nonbinding electron pairs?
CO2
H2S
PF3
SiF4
More than one of a, b, c, d.
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42. Practice Exercise 1 (8.6)
Which of these molecules has a Lewis structure with a central atoms having no nonbinding electron pairs?
CO2
H2S
PF3
SiF4
More than one of a, b, c, d.
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43. Practice Exercise 2 (8.6)
How many valence electrons should appear in the Lewis structure for CH2Cl2?
Draw the Lewis structure.
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44. Sample Exercise 8.7 (p. 316) Draw the Lewis structure for HCN.
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45. Practice Exercise 1 (8.7)
Draw the Lewis structure(s) for the molecule with the chemical formula C2H3N, where the N is connected to only one other atom. How many double bonds are there in the correct Lewis Structure?
zero
one
two
three
four
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46. Practice Exercise 1 (8.7)
Draw the Lewis structure(s) for the molecule with the chemical formula C2H3N, where the N is connected to only one other atom. How many double bonds are there in the correct Lewis Structure?
zero
one
two
three
four
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47. Practice Exercise 2 (8.7) Draw the Lewis structure for a) NO+ ion;
b) C2H4
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48. Sample Exercise 8.8 (p. 317)
Draw the Lewis structure for the BrO3- ion.
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49. Practice Exercise 1 (8.8)
How many nonbonding electron pairs are there in the Lewis structure of the peroxide ion, O22-? 7 6 5 4 3
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50. Practice Exercise 1 (8.8)
How many nonbonding electron pairs are there in the Lewis structure of the peroxide ion, O22-? 7 6 5 4 3
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51. Practice Exercise 2 (8.8) Draw the Lewis structure for a) ClO2- ion
b) PO43- ion
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52. Writing Lewis Structures Formal Charges
“Formal charges” help us choose which Lewis Structure is most likely, from a group of Lewis Structures that all fulfill the octet rule.
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53. Writing Lewis Structures
After drawing suitable Lewis Structures, assign formal charges.
For each atom, count the electrons in lone pairs and half the electrons it shares with other atoms.
Subtract that from the number of valence electrons for that atom: the difference is its formal charge.
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54. Writing Lewis Structures
The best Lewis structure…
…is the one with the fewest charges.
…puts a negative charge on the most electronegative atom.
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55. Sample Exercise 8.9 (p. 318)
The following are three possible Lewis structures for the thiocyanate ion, NCS-:
a) Determine the formal charges of the
atoms in each structure.
b) Which Lewis structure is the preferred
one?
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56. Practice Exercise 1 (8.9)
Phosphorus oxychloride has the chemical formula POCl3, with P as the central atom. To minimize formal charge, how many bonds does phosphorus make to the other atoms in the molecule?
(Count each single bond as one, each double bond as two, and each triple bond as three.) 3 4 5 6 7
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57. Practice Exercise 1 (8.9)
Phosphorus oxychloride has the chemical formula POCl3, with P as the central atom. To minimize formal charge, how many bonds does phosphorus make to the other atoms in the molecule?
(Count each single bond as one, each double bond as two, and each triple bond as three.) 3 4
5 3 single bonds, 1 double bond 6 7
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58. Practice Exercise 2 (8.9)
The cyanate ion (NCO-), like the thiocyanate ion, has three possible Lewis structures.
a) Draw these three Lewis structures and
assign formal charges to the atoms in
each structure;
b) Which Lewis structures should be the
preferred one?
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59. Resonance This is the Lewis structure we would draw for ozone, O3. - +
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60. Resonance
But this is at odds with the true, observed structure of ozone, in which…
…both O-O bonds are the same length.
…both outer oxygens have a charge of -1/2.
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61. Resonance
One Lewis structure cannot accurately depict a molecule like ozone.
We use multiple structures, resonance structures, to describe the molecule.
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62. Resonance Just as green is a synthesis of blue and yellow…
…ozone is a synthesis of these two resonance structures.
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63. Sample Exercise 8.10 (p. 321)
Which is predicted to have the shorter sulfur-oxygen bonds, SO3 or SO32-?
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64. Practice Exercise 1 (8.10)
Which of these statements about resonance is true?
When you draw resonance structures, it is permissible to alter the way atoms are connected.
The nitrate ion has one long N-O bond and two short N-O bonds.
“Resonance” refers to the idea that molecules are resonating rapidly between different bonding patterns.
The cyanide ion has only one dominant resonance structure.
All of the above are true.
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65. Practice Exercise 1 (8.10)
Which of these statements about resonance is true?
When you draw resonance structures, it is permissible to alter the way atoms are connected.
The nitrate ion has one long N-O bond and two short N-O bonds.
“Resonance” refers to the idea that molecules are resonating rapidly between different bonding patterns.
The cyanide ion has only one dominant resonance structure.
All of the above are true.
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66. Practice Exercise 2 (8.10)
Draw two equivalent resonance structures for the formate ion, HCO2-.
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67. Resonance
In truth, the electrons that form the second C-O bond in the double bonds below do not always sit between that C and that O, but rather can move among the two oxygens and the carbon.
They are not localized; they are delocalized.
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68. Resonance
The organic compound benzene, C6H6, has two resonance structures.
It is commonly depicted as a hexagon with a circle inside to signify the delocalized electrons in the ring.
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69. Exceptions to the Octet Rule
There are three types of ions or molecules that do not follow the octet rule:
Ions or molecules with an odd number of electrons
Ions or molecules with less than an octet
Ions or molecules with more than eight valence electrons (an expanded octet)
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70. Odd Number of Electrons
Though relatively rare and usually quite unstable and reactive, there are ions and molecules with an odd number of electrons.
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71. Fewer Than Eight Electrons
Consider BF3:
Giving boron a filled octet places a negative charge on the boron and a positive charge on fluorine.
This would not be an accurate picture of the distribution of electrons in BF3.
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72. Fewer Than Eight Electrons
Therefore, structures that put a double bond between boron and fluorine are much less important than the one that leaves boron with only 6 valence electrons.
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73. Fewer Than Eight Electrons
The lesson is: if filling the octet of the central atom results in a negative charge on the central atom and a positive charge on the more electronegative outer atom, don’t fill the octet of the central atom.
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74. More Than Eight Electrons
The only way PCl5 can exist is if phosphorus has 10 electrons around it.
It is allowed to expand the octet of atoms on the 3rd row or below.
Presumably d orbitals in these atoms participate in bonding.
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75. More Than Eight Electrons
Even though we can draw a Lewis structure for the phosphate ion that has only 8 electrons around the central phosphorus, the better structure puts a double bond between the phosphorus and one of the oxygens.
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76. More Than Eight Electrons
This eliminates the charge on the phosphorus and the charge on one of the oxygens.
The lesson is: when the central atom in on the 3rd row or below and expanding its octet eliminates some formal charges, do so.
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77. Sample Exercise 8.11 (p. 324) Draw the Lewis structure for ICl4-.
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78. Practice Exercise 1 (8.11)
In which of these molecules or ions is there only one lone pair of electrons on the central sulfur atom?
SF4
SF6
SOF4
SF2
SO42-
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79. Practice Exercise 1 (8.11)
In which of these molecules or ions is there only one lone pair of electrons on the central sulfur atom?
SF4
SF6
SOF4
SF2
SO42-
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80. Practice Exercise 2 (8.11)
a) Which of the following atoms is never found with more than an octet of electrons around it: S, C, P, Br?
b) Draw the Lewis structure for XeF2.
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81. Covalent Bond Strength end of Thermodynamics Unit
Most simply, the strength of a bond is measured by determining how much energy is required to break the bond.
This is the bond enthalpy.
The bond enthalpy for a Cl-Cl bond, D(Cl-Cl), is measured to be 242 kJ/mol.
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82. Average Bond Enthalpies
This table lists the average bond enthalpies for many different types of bonds.
Average bond enthalpies are positive, because bond breaking is an endothermic process.
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83. Average Bond Enthalpies
NOTE: These are average bond enthalpies, not absolute bond enthalpies; the C-H bonds in methane, CH4, will be a bit different than the C-H bond in chloroform, CHCl3.
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84. Enthalpies of Reaction
Yet another way to estimate H for a reaction is to compare the bond enthalpies of bonds broken to the bond enthalpies of the new bonds formed.
In other words,
Hrxn = (bond enthalpies of bonds broken) -
(bond enthalpies of bonds formed)
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85. Enthalpies of Reaction
CH4 (g) + Cl2 (g) 
CH3Cl (g) + HCl (g)
In this example, one C-H bond and one Cl-Cl bond are broken; one C-Cl and one H-Cl bond are formed.
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86. Enthalpies of Reaction
H = [D(C-H) + D(Cl-Cl)] - [D(C-Cl) + D(H-Cl)]
= [(413 kJ) + (242 kJ)] - [(328 kJ) + (431 kJ)]
= (655 kJ) - (759 kJ)
= -104 kJ
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87. Sample Exercise 8.12 (p. 328)
Using Table 8.4, estimate ΔH for the reaction in the figure (where we explicitly show the bonds involved in the reactants and products). (A: kJ)
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88. Practice Exercise 1 (8.12)
Using Table 8.4, estimate DH for the “water splitting reaction”: H2O(g)  H2(g) + ½ O2(g).
242 kJ
417 kJ
5 kJ
– 5 kJ
– 468 kJ
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89. Practice Exercise 1 (8.12)
Using Table 8.4, estimate DH for the “water splitting reaction”: H2O(g)  H2(g) + ½ O2(g).
242 kJ
417 kJ
5 kJ
– 5 kJ
– 468 kJ
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90. Practice Exercise 2 (8.12)
Using Table 8.4, estimate ΔH for the reaction shown.
(– 468 kJ)
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91. Bond Enthalpy and Bond Length
We can also measure an average bond length for different bond types.
As the number of bonds between two atoms increases, the bond length decreases.
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92. Sample Integrative Exercise 8
Phosgene, a substance used in poisonous gas warfare in World War I, is so named because it was first prepared by the action of sunlight on a mixture of carbon monoxide and chlorine gases. Its name comes from the Greek words phos (light) and genes (born of). Phosgene has the following elemental composition: % C, 16.17% O, and 71.69% Cl by mass. Its molar mass is 98.9 g/mol.
a) Determine the molecular formula of this compound.
b) Draw three Lewis structures for the molecule that satisfy the octet rule for each atom. (The Cl and O atoms bond to C.)
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93. Sample Integrative Exercise 8
c) Using formal charges, determine which Lewis structure is the most important one.
Using average bond enthalpies,
estimate DH for the formation of
gaseous phosgene from CO(g) and Cl2(g).
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