Exam 3 Worksheet Answers Chemistry 102

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1 hapter 6 Energy Relationships in hemical Reactions 1. eat a. 125 J b. 110 J c J d kj e J f. 174 J g J 2. Specific heat a b c d e f g eat and final temperature a º b º c º d º e º f º g º 4. eat and final temperature a º b º c º d º e º f º g º 5. eat transfer from a hotter body to water a º b º hemistry 102 1

2 º c º d º e º f º 7. Enthalpy of solution a. Δ soln = kj mol 1 ; T f = 38.2 º b. Δ soln = kj mol 1 ; T f = 19.8 º c. Δ soln = kj mol 1 ; T f = 21.4 º d. Δ soln = 2.33 kj mol 1 ; T f = 22.8 º 8. Enthalpy of combustion a. Δ comb (exp) = 2800 kj mol 1 ; Δ comb (calc) = kj mol 1 ; error = 10.4% b. Δ comb (exp) = 5000 kj mol 1 ; Δ comb (calc) = kj mol 1 ; error = 3.10% c. Δ comb (exp) = 380 kj mol 1 ; Δ comb (calc) = kj mol 1 ; error = 3.43% 9. Enthalpy of combustion a. 278 L b. 349 L c. 156 L d. 178 L e. 168 L f. 109 L 10. Enthalpy of reaction a kj exothermic b. 178 kj endothermic c. 126 kj exothermic d kj exothermic f = kj mol 1 (l 3 (l)) e kj exothermic 11. Enthalpy of reaction ess s Law a. 135 kj exothermic b. 108 kj exothermic c kj exothermic d kj exothermic e kj exothermic 12. Enthalpy of formation a. 463 kj mol 1 exothermic b kj mol 1 endothermic c kj mol 1 exothermic 2

3 d. 188 kj mol 1 exothermic 13. Enthalpy of formation a kj mol 1 exothermic b kj mol 1 exothermic c kj mol 1 exothermic d kj mol 1 exothermic 14. Formation reactions a. 6( s) + 32( g) + 1 2( g) 65 ( l 2 ) b. s ( ) + 2( g) 2( g) c. s ( ) ( g) g 2 ( ) d. 8( s) + 92( g) 818( l) e. N2( g) + 1 2( g) N2( g 2 ) f. N ( g) + 2 ( g) N ( g) hapter 7 Electronic Structure of Atoms 1. Balmer series a. (n 1 = 3) ΔE = J ν = z λ = 656 nm b. (n 1 = 4) ΔE = J ν = z λ = 486 nm c. (n 1 = 5) ΔE = J ν = z λ = 434 nm d. (n 1 = 6) ΔE = J ν = z λ = 410 nm 2. Lyman series a. (n 1 = 2) ΔE = J ν = z λ = 121 nm b. (n 1 = 3) ΔE = J ν = z λ = 103 nm c. (n 1 = 4) ΔE = J ν = z λ = 97.2 nm d. (n 1 = 5) ΔE = J ν = z λ = 94.9 nm 3. Quantum numbers a. (4, 2, ± 2, ± ½); (4, 2, ± 1, ± ½); (4, 2, 0, ± ½) b. (5, 1, ± 1, ± ½); (5, 1, 0, ± ½) c. (3, 0, 0, ± ½); (3, 1, ± 1, ± ½); (3, 1, 0, ± ½) (3, 2, ± 2, ± ½); (3, 2, ± 1, ± ½); (3, 2, 0, ± ½) d. (4, 0, 0, ± ½); (4, 1, ± 1, ± ½); (4, 1, 0, ± ½) (4, 2, ± 2, ± ½); (4, 2, ± 1, ± ½); (4, 2, 0, ± ½) (4, 3, ± 3, ± ½); (4, 3, ± 2, ± ½); (4, 3, ± 1, ± ½); (4, 3, 0, ± ½) e. (2, 0, 0, ± ½) f. (7, 3, ± 3, ± ½); (7, 3, ± 2, ± ½); (7, 3, ± 1, ± ½); (7, 3, 0, ± ½) 3

4 4. v < vi = xi < x < viii < iii < iv < ix < i < ii = vii < xii 5. Quantum number a. F b. G c. E d. F e. G f. G g. E h. F i. F j. F k. E l. G hapter 8 The Periodic Table 1. Electronic configurations a. [Ar] 4s 2, diamagnetic 4s b. [Ar] 4s 2 2, paramagnetic c. [Ar] 4s 0 0, diamagnetic 3p d. [Ar] 4s 1 10, paramagnetic 4s e. [Ar] 4s 0 6, paramagnetic f. [Ar] 4s 0 3, paramagnetic g. [Ar] 4s 0 10, diamagnetic h. [Xe] 6s 2 4f 14 5d 10 6p 2 6 p, paramagnetic i. [Xe] 6s 2 4f 14 5d 10, diamagnetic 5d j. [Ne] 3s 2 3p 4, paramagnetic 3p 4

5 k. [Ne] 3s 2 3p 6, diamagnetic 3p l. [Xe] 6s 2 4f 11, paramagnetic 4f m. [Rn] 7s 2 5f 7, paramagnetic 5 f n. [Kr] 5s 0 4d 8, paramagnetic 4d o. [Kr] 5s 0 4d 10, diamagnetic 4d 2. Periodic Trends Largest radii ighest IE a. Si N b. Ge P c. Na Na + d. Br Br e. Se l 3. Elements are placed in the periodic table based on their physical and chemical properties. Elements within a group have more properties in common than elements within the same period. The metallic character of elements increases down a group (from top to bottom) and decreases across a period (from left to right). In the group 4a elements, carbon is a nonmetal, silicon and germanium are metalloids and tin and lead are metals. Tin and lead will react with acid to form hydrogen and a 2+ cation. These metals can also exist in a 4+ oxidation state. For silicon and carbon, the 4+ oxidation state is the most stable. (Think about 2 versus.) Lead(II) is more stable than lead(iv). In period 3, the metallic character decreases from left to right, sodium, magnesium and aluminum are metals, silicon is a metalloid and phosphorus, sulfur, chlorine and argon are nonmetals. The periodicity of the oxides of these elements in water are basic (on the left sodium, etc) to amphoteric (aluminum) to acidic (silicon continuing to the right not including argon). hapter 9 hemical Bonding I The ovalent Bond 1. a. I b. I c. d. I e. f. N g. I 5

6 h. I i. I j. k. N l. m. n. N o. IN p. q. I 2. For the Born aber cycle the standard molar enthalpy of formation (of the solid compound), the enthalpy of sublimation of the metal, the ionization energy(s) of the metal, the electron affinity of the nonmetal, the bond enthalpy of the nonmetal are needed to determine the lattice energy for the compound. a. Lil(s) Li+ (g) + l (g) b. Mgl 2 (s) Mg2+ (g) + 2l (g) c. KF(s) K+ (g) + F (g) d. af 2 (s) a2+ (g) + 2F (g) 3. 6

7 4. Lewis dot structures lone pairs are terminal atoms are not shown for clarity here but must be included on a quiz or exam. a. S, bent, sp 2, polar b., trigonal planar, sp 2, polar c. tetrahedral, sp 3, polar; d. N, linear, sp, polar e. N N, linear, sp, polar bent, sp 3, polar f. l S l F, trigonal pyramidal, sp 3, polar F l g. l h. i., tetrahedral, sp 3, polar, linear, sp, non polar, trigonal planar, sp 2, non polar N N j., trigonal planar, sp 2, nonpolar k. l, trigonal planar, sp 2, nonpolar l. l, trigonal planar, sp 2, polar m., trigonal planar, sp 2, polar 7

8 5. S S these are equivalent structures but based on formal charges, the structure given in 4a is the best. N N N N N N , first structure is most valid since is more electronegative than N. 0-1 l S 0 l l S +1 l, first structure is more valid N N, equivalent to the structure given in 4j 6. All breaking of bonds is endothermic and all making of bonds is exothermic. a. Break 1 mol of,, 3 mol of, and 3 / 2 mol of =; make 2 mol of = and 4 mol of b. Break 1 mol of =, 2 mol of, and 5 / 2 mol of =; make 4 mol of = and 2 mol of c. Break 1 mol of N N, and 3 mol of ; make 6 mol of N d. Break 1 mol of = and ½ mol of =; make 2 mol of = e. Break 1 mol of F F and ; make 2 mol F f. Break 1 mol of l l and ; make 2 mol l g. Break 2 mol of N and N=; make 2 mol of N, N=, and 1 mol of N N a. 637 kj; exothermic b kj; exothermic c. 107 kj; exothermic d. Need e. 543 kj; exothermic f. 185 kj; exothermic g. 193 kj; exothermic a kj b kj c kj d. 283 kj e kj f kj g kj 8

9 9. The differences observed between 7 and 8 can generally be attributed to the methods required for approximating bond enthalpies for non diatomic molecules. The universal application of one bond enthalpy regardless of the environment of the bond (for example, bond energy within molecules varies depending on the molecule) can lead to error. 10. (nly one shown, although the idea is the same for all) 11. This reaction was purposefully chosen. learly, this proposed intermediate state is ridiculous because this reaction occurs through the formation of a N N bond. The idea of bond enthalpies and much better proposed intermediates will be discussed next semester in kinetics. 9