4)Write Lewis Dot Symbols for the Following Ions

CHAPTER 8 Problems: 4, 6, 18, 20, 21, 22, 28, 30, 38, 42, 48, 54, 58, 68, 71, 76, 88, 90, 106 a-c, 115

4)Write Lewis dot symbols for the following ions: a) Li+ b) Cl- c) S2- d) Sr2+ e) N3-

6) Explain what ionic bonding is.

Ionic bonding is the bonding that occurs when electrons are transferred from one atom (or group of atoms) to another atom (or group of atoms). The bonding occurs due to the attractive force that exists between ions of opposite charge.

18) Specify which compound in each of the following pairs of ionic compounds should have a higher lattice energy. Explain your choice.

a) AlN or CaO AlN The ions involved are 3+/3-, while those involved in CaO are 2+/2-.

b) NaF or CsF NaF Both Na and Cs are + ions, but Na+ is smaller than Cs+ and so will have a stronger force of attraction for F-.

- - c) MgCl2 or MgF2 MgF2. Both F and Cl are - ions, but F is smaller than Cl and so will have a stronger force of attraction for Mg2+.

20) Calculate the lattice energy for CaCl2. Use data from Fig. 7.8 and Fig. 7.10, and data from Appendix 2. Note that the second ionization energy for Ca is IE2 = 1145. kJ/mol.

To do this problem we need a Born-Haber cycle.

Ca(s)  Ca(g) Hf(Ca(g))

Cl2 (g)  2 Cl(g) 2 Hf(Cl(g)) + - Ca(g)  Ca (g) + e IE1(Ca) + 2+ - Ca (g)  Ca (g) + e IE2(Ca) 2 Cl(g) + 2 e-  2 Cl-(g) 2 EA(Cl) 2+ - Ca (g) + 2 Cl (g)  CaCl2(s) - Hlattice(CaCl2) ______

Ca(s) + Cl2(g)  CaCl2(s) Hf(CaCl2(s))

1 From Hess’ law

Hf(CaCl2(s)) = Hf(Ca(g)) + 2 Hf(Cl(g)) + IE1(Ca) + IE2(Ca)

+ 2 EA(Cl) - Hlattice(CaCl2(s)) and so

Hlattice(CaCl2(s)) = Hf(Ca(g)) + 2 Hf(Cl(g)) + IE1(Ca) + IE2(Ca)

+ 2 EA(Cl) - Hf(CaCl2(s))

= 179.3 kJ/mol + 2 (121.7 kJ/mol) + 590. kJ/mol + 1145. kJ/mol

+ 2 ( - 349. kJ/mol) - (- 794.96 kJ/mol) = + 2255. kJ/mol

21) An ionic bond is formed between a cation A+ and an anion B-. Based on Coulomb's law

E ~ Q1 Q2 d

How would the energy of the ionic bond be affected by each of the following changes?

a) doubling the radius of A+. d would increases and so E would decrease.

+ b) tripling the charge of A . Q1 would triple, and so E would also triple.

+ - c) doubling the charges on A and B .Both Q1 and Q2 would double, and so E would be 4 times larger (2 x 2).

d) decreasing the radii of A+ and B- to half their original values. If both radii 1 1 decrease to /2 their original value then d will decrease to /2 its original value, and so E will double.

22) Give the empirical formula and name for the compounds formed from each of the pairs of ions a) Rb+ and I- RbI + 2- b) Cs and SO4 Cs2SO4 2+ 3- c) Sr and N Sr3N2 3+ 2- d) Al and S Al2S3

2 28) How many lone pairs of electrons are there on each of the underlined atoms in the following compounds? a) HBr 3 pairs b) H2S 2 pairs c) CH4 none

30) Compare the properties of ionic compounds and covalent compounds.

Ionic Held together by the attractive forces that exist between cations and anions. Usually form between metals and nonmetals. Usually solids at room temperature. Usually have high melting and boiling point temperatures. Usually hard and brittle. Strong electrolytes when dissolved in water.

Covalent Held together by sharing one or more pairs of electrons Usually form between two or more nonmetals. Can be solids, liquids, or gases at room temperature. Usually have low melting and boiling point temperatures. Usually soft. Usually nonelectrolytes when dissolved in water. Exceptions are strong acids, and weak acids and bases.

38) Classify the following bonds as covalent, polar covalent, or ionic. Explain.

a) The CC bond in H3CCH3 Covalent (both bonded atoms are the same, so EN = 0

b) The KI bond in KI Ionic. Metal + nonmetal. EN = 1.7.

c) The NB bond in H3NBCl3 Two different nonmetals, and so polar covalent. EN = 1.0

d) The CF bond in CF4 Two different nonmetals, and so polar covalent. EN = 1.5

3 42) Draw Lewis structures for the following molecules or ions: - a) NCl3 c) H2O2 e) CN - + b) OCS d) CH3COO f) CH3CH2NH3

Note that structure d has a resonance structure where the single and double CO bonds are interchanged. Also note that the solution manual (which is where I copied the Lewis structures) does not put brackets around ions, as it should.

48) Draw Lewis structures for the following ions. Show formal charge. 2- + a) O2 c) NO 2- + b) C2 d) NH4

- 54) Draw three resonance structures for the chlorate ion (ClO3 ). Show formal charges.

In each of the above resonance structures Cl has a formal charge of +1, and the two single bonded oxygens have a formal charge of -1.

58) Draw three resonance structures for the molecule NNO, where N is the central atom. Indicate formal charges, and rank the resonance structures in order of most important to least important.

4 The structure on the left is the most important resonance structure, as it keeps the formal charges close to zero and puts the negative formal charge on oxygen, the more electronegative atom. The structure in the middle also minimizes formal charges, but is not as good as the first structure as it puts a negative formal charge on a N atom. The structure on the right is the worse resonance structure, as all of the formal charges differ from zero, and one formal charge is -2.

68) Of the noble gases, only Kr, Xe, and Rn are known to form a few compounds with O and/or F. Write Lewis structures for the following molecules. In each case Xe is the central atom. a) XeF2 c) XeF6 e) XeO2F2 b) XeF4 d) XeOF4

71) Write Lewis structures for the reaction

- - AlCl3 + Cl  AlCl4

What kind of bond joins Al and Cl in the product?

Both electrons for one of the Al-Cl bonds come from the Cl atom, and so the bond is a coordinate covalent (dative) bond. Note that AlCl3 is unusual because normally metal + nonmetal leads to an ionic compound, and in AlCl3 the Al atom violates the octet rule by having less than 8 valence electrons.

5 76) From the following data find the average bond enthalpy for an N-H bond

NH3(g)  NH2(g) + H(g) H = 435. kJ/mol

NH2(g)  NH(g) + H(g) H = 381. kJ/mol NH(g)  N(g) + H(g) H = 360. kJ/mol

From Hess’ law, the combination of the above three processes is equivalent to

NH3(g)  N(g) + 3 H(g)

1 So the average N-H bond enthalpy is /3 [ 435. kJ/mol + 381. kJ/mol + 360. kJ/mol ] ; or 392. kJ/mol (close to the value in Table 8.6).

88) Write Lewis structures for BrF3, ClF5, and IF7. Identify the structures where the octet rule is not obeyed.

The octet rule is violated for the central atom in all of the above Lewis structures.

90) Draw two resonance structures for sulfurous acid (H2SO3) - one that obeys the octet rule, the other that minimizes formal charge. Determine the formal charges for each atom in each resonance structure.

The structure on the left obeys the octet rule for all atoms, but has a formal charge of +1 on S and -1 on the O not bonded to an H atom. The S atom in the structure on the right violates the octet rule, but all formal charges are equal to zero. Both resonance structures are likely to be important in describing the bonding in H2SO3.

6 106) Draw Lewis structures for the following organic molecules

a) tetrafluoroethylene, C2F4 b) propane, C3H8 c) butadiene, CH2CHCHCH2

115) Calculate the value for H for the reaction

H2(g) + I2(g)  2 HI(g)

using eq 8.4, and using eq 5.19. Note that Hf(I2(g)) = 61.0 kJ/mol.

From 8.4

H = [ (H-H) + (I-I) ] - [ 2 (H-I) ]

= [ (436.4 kJ/mol) + (151.0 kJ/mol) ] - [ 2(298.3 kJ/mol) ] = - 9.2 kJ/mol

From 5.19

H = [ 2 Hf(HI(g)) ] - [Hf(H2(g)) + Hf(I2(g)) ]

= [ 2(25.9 kJ/mol) ] - [ (0.0 kJ/mol) + (61.0 kJ/mol) ] = - 9.2 kJ/mol

The agreement is not as impressive as it looks. If you examine Table 8.6, you will see all of the bond enthalpies used in the calculation using eq 8.4 are for diatomic molecules, and are therefore exact (and not average) values. So this is not a true test of the accuracy of eq 8.4 for estimating enthalpies of reaction.