1. Chapter 9 Chemical Bonding I: Lewis Theory
Outline
Lewis Theory
Types of Chemical Bonds
Ionic
Born-Haber Cycle
Lattice Energy
Covalent
Electronegativity
Lewis Structures
Bond Energy
Bond Length
Metallic

2. Gilbert Newton Lewis

3. Potential Energy versus Distance

4. Where is the electrostatic potential most energetically favorable?

5. How are bonds classified?
Figure: 09-03
Title:
Ionic, Covalent, and Metallic Bonding
Caption:
Examples of the three types of bonding explored in Chapter 9.

6. How are bonds classified?

7. Figure: 09-07
Title:
Intermolecular and Intramolecular Forces
Caption:
The covalent bonds between atoms of a molecule are much stronger than the interactions between molecules. To boil a molecular substance, you simply have to overcome the relatively weak intermolecular forces, so molecular compounds generally have low boiling points.

10. FIGURE 8.6 The trends in the electronegativities of the main group elements follow those of ionization energies: both tend to (a) increase with increasing atomic number across a row and (b) decrease with increasing atomic number within a group.

11. FIGURE 8.6 The trends in the electronegativities of the main group elements follow those of ionization energies: both tend to (a) increase with increasing atomic number across a row and (b) decrease with increasing atomic number within a group.

12. Bond Polarity
NaCl
HCl
Cl-Cl

14. Bond Polarity ENCl = 3.0 ENH = 2.1 3.0 – 2.1 = 0.9 Polar Covalent
3.0 − 3.0 = 0
Pure Covalent
ENCl = 3.0
ENNa = 0.9
3.0 – 0.9 = 2.1
Ionic
Tro: Chemistry: A Molecular Approach, 2/e

16. When is a compound considered ionic?

17. How does an electric field affect polar molecules?
Figure: UN
Title:
Bond polarity in HF
Caption:
Hydrogen and fluorine do not share their electrons equally; the electrons are around F more than H.

18. What are some exceptions to the octet rule?
FIGURE 8.17 The chain structure of solid beryllium chloride, with bridging chlorine atoms, changes to discrete BeCl2 molecules in the gas phase. Each beryllium atom in the solid has a completed octet. In the gas phase, each Be atom is surrounded by only four electrons.
FIGURE 8.16 Aluminum(III) chloride completes the octet on aluminum by sharing a pair of electrons with one Cl atom form a second molecule of AlCl3.

19. What is the optimum bond length of a hydrogen molecule?
FIGURE 8.1 The electrostatic potential energy (Eel) profile of a H–H bond: (a) Two H atoms that are far apart do not interact, so Eel = 0. (b) As the atoms approach each other, mutual attraction between their positive nuclei and negative electrons causes Eel values to decrease. (c) At 74 pm Eel reaches a minimum as the two atoms form a H–H bond. (d) If the atoms are even closer together, repulsion between their nuclei causes Eel to increase and destabilizes the bond.

22. How can the energy of a reaction be determined?
FIGURE 8.20 The combustion of 1 mole of methane requires that 4 moles of C–H bonds and 2 moles of O– –O bonds be broken. These processes require an enthalpy change of about kJ. In the formation of 4 moles of O–H bonds and 2 moles of C– – –O bonds, there is an enthalpy change of about –3450 kJ. The overall reaction is exothermic:2642 kJ – 3450 kJ = –808 kJ.

23. Example – Bond Energy
Approximate the ΔHrxn for the production of ammonia by the Haber process: N2 (g) + 3 H2 (g)  2 NH3 (g)

24. Example – Bond Energy
Approximate the ΔHrxn for the combustion of methane: CH4 (g) + 2 O2 (g)  CO2 (g) + 2 H2O (g)

25. Example – Bond Energy
Approximate the ΔHrxn for the halogenation of acetylene gas: C2H2 (g) + 2 Cl2 (g)  C2H2Cl4 (g)

26. What happens when sodium metal and chlorine gas are placed in the same reaction flask?
Tro: Chemistry: A Molecular Approach, 2/e

27. What is the Born-Haber Cycle?
FIGURE 11.6 The reaction between sodium metal and chlorine gas releases more than 400 kJ of energy per mole of NaCl produced. The Born–Haber cycle shows that the most exothermic step is the combination of free sodium ions Na+(g) and free chloride ions Cl–(g) to form NaCl(s).

28. What is lattice energy? Figure: 09-05 Title: Lattice Energy Caption:
The lattice energy of an ionic compound is the energy associated with forming a crystalline lattice of the compound from the gaseous ions.

29. How does atomic size affect lattice energy?
Metal Chloride
Lattice Energy
LiCl
-834 kJ/mol
NaCl
-787 kJ/mol
KCl
-701 kJ/mol
CsCl
-657 kJ.mol
Figure: UN
Title:
Group I metal chlorides with radii
Caption:
The relative sizes of the ions will determine the lattice energy.

30. How does ionic charge affect lattice energy?
Ionic Compound
Lattice Energy
NaF
-910 kJ/mol
CaO
-3414 kJ/mol
Figure: UN
Title:
Group I metal chlorides with radii
Caption:
The relative sizes of the ions will determine the lattice energy.

31. What are some lattice energy values?

32. Figure: 09-06
Title:
Born-Haber Cycle for Sodium Chloride
Caption:
The sum of the steps is the formation of NaCl from elemental Na and Cl. The enthalpy change of the last step is the lattice energy.

33. Example – Formation of Ionic Compounds
Calculate the enthalpy of formation of sodium chloride from it’s elements. Given:
Na (s) → Na (g) kJ/mol
Na (g) → Na+(g) + 1 e kJ/mol
½ Cl2 (g) → Cl (g) kJ/mol
Cl (g) + 1 e- → Cl- (g) kJ/mol
Na+ (g) + Cl- (g) → NaCl (s) kJ/mol

35. Example – Formation of Ionic Compounds
Calculate the energy released in kJ/mol when sodium iodide is formed.
Na (s) + ½ I2 (s) → NaI (s)
The energy of vaporization of elemental sodium is 107 kJ/mol. The ionization energy of sodium is 496 kJ/mol. The sum of the enthalpies of dissociation and subimation of elemental iodine is 214 kJ/mol and the electron affinity of iodine is -295 kJ/mol. The lattice energy of sodium iodide is -704 kJ/mol.

36. Example – Formation of Ionic Compounds
Calculate the energy released in kJ/mol when lithium hydride is formed. The heat of vaporization of elemental lithium is 161 kJ/mol, the ionization energy of lithium is 520 kJ/mol. The dissociation energy of hydrogen gas is 436 kJ/mol and the electron affinity of a gaseous hydrogen atom is -73 kJ/mol. The lattice energy of lithium hydride is -917 kJ/mol.

37. Example – Formation of Ionic Compounds
Determine the energy of formation of magnesium bromide. Given:
Mg (s) → Mg (g) kJ/mol
Mg (g) → Mg+(g) + 1 e kJ/mol
Mg+(g) → Mg2+ (g) + 1 e ,450.7 kJ/mol Br2 (g) → 2 Br (g) kJ/mol
Br (g) + 1 e- → Br- (g) kJ/mol
Mg2+ (g) + 2 Br- (g) → MgBr2 (s) -2,440 kJ/mol

38. Metal Bonding