3. Given the Following Information, Determine the Electron Affinity (EA1) of Fluorine

Chem 31 SI 11/15/09

Chapter 9 Material

Warm Up:

1. Draw a dot structure for Magnesium and Sulfur, and visually represent how electrons are exchanged to form Magnesium Sulfide.

2. Without looking up the actual ∆H values, map out the strategy you would use (using a Born – Haber cycle) to determine the ∆Hlattice for MgS, assuming that the ∆Hfo, ∆Hsub, ionization energies and electron affinities are all available.

Bonus Fun:

3. Given the following information, determine the electron affinity (EA1) of fluorine:

K+ (g)+ F- (g) → KF (s)∆Hlattice = -821 kJ/mol

K(s) + ½ F2 → KF (s)∆Hfo = -569 kJ/mol (published)

K (s) →K (g)∆Hsub K = 90 kJ/mol

K (g) →K+ (g)∆H IE1 K = 419kJ/mol

½ F2 (g) → F (g)∆H ½ BE F2 = 79.5 kJ/mol

F (g) → F- (g)∆H EA1 F= ??? kJ/mol

Answer: ∆H EA1 = -336.5 kJ/mol

4. Reproduce Coulombs law as we learned it during lecture.

5. Based on this, identify the two important trends in the magnitude of lattice energy:

1)

2)

6. Use the trends to predict which of the following has the most exothermic ∆Hlattice:

LiCl vs CsClNaF vs. CaO

7. Use the trends to place the following ionic compounds in order of increasingly exothermic ∆Hlattice.

KClSrORbBrCaO

8. Polar covalent bonds are essentially an intermediate between pure covalent bonds and ionic bonds. What atomic property accounts for unequal sharing? Visually represent the polar covalent bonds in HF and in H20:

9. On the periodic table above, use arrows to represent the trends in Electronegativity.

10. Reproduce table 9.1 from the textbook (page 379)

11. Based on the above table and figure 9.1 on page 378, predict whether a bond between the following pairs of atoms will be Covalent, Polar-Covalent, or Ionic:

I and ICs and BrH and I

12. What is the difference between electronegativity and electron affinity?

13. Rank the members of each set of compounds in order of increasing ionic character of their bonds:

A) HBr, HCl, HIB) SCl2, PCl3, SiCl4C) H2O, CH4, HF

Formal Charge:

Reproduce the general formula for calculating the formal charge on an element:

14. In order for Hydrogen to complete it’s doublet and have a formal charge of zero, it tends to form ______bond(s) and have ______lone pair(s).

15. In order for Carbon to complete it’s octet and have a formal charge of zero, it tends to form ______bond(s) and have ______lone pair(s).

16. In order for Nitrogen to complete it’s octet and have a formal charge of zero, it tends to form ______bond(s) and have ______lone pair(s).

17. In order for Oxygen to complete it’s octet and have a formal charge of zero, it tends to form ______bond(s) and have ______lone pair(s).

18. In order for Flourine to complete it’s octet and have a formal charge of zero, it tends to form ______bond(s) and have ______lone pair(s).

19. Assign formal charges to all elements in the following lewis structures and circle the preferred structure:

Lewis Structures

20. Write out the rules for Completing Lewis Structures from page 382

21. Draw Viable Lewis Structures for:

CO2N2

NO2Br (N is the central atom)NH4+

CH2OCO

(note: in order to draw lewis structures for incomplete and expanded octets, you will be successful if you focus on whether the formal charges are zero instead of whether the octet rule is met)

Incomplete Octets

22. What is the most significant element when drawing lewis structures with incomplete octets?______

23. Draw viable lewis structures for the following?

BF3BH3

Odd-Electron Species:

24. Draw the best possible lewis structure for the following:

NO2NO

Expanded Octets:

25. What determines whether or not an element can expand its octet?

26. Draw viable lewis structures for the following:

XeF3+ClF4-

SO2Cl2 (S is the central atom)H3PO4 (P is the central atom)

Additional Lewis Structure Practice: (assign formal charges to elements in the ionic compounds)

ClF5Cl3PO (P is Central)

AsF6-BrO2

CN-ClO2-

IF5OF2

SO3XeF2

PO43-NO

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