Exam 3 Worksheet Exam 3 Worksheet – Chemistry 102 Chapter 6 – Energy Relationships in Chemical Reactions 1. What is the heat evolved when 12.25 g of the substances decrease in temperature from 75ºC to 55ºC? a. Diamond (c = 6.113 J⋅K−1⋅mol−1) b. Iron (c = 25.10 J⋅K−1⋅mol−1) c. Brass (c = 0.38 J⋅g−1⋅ºC−1) d. Water (c = 4.18 J⋅g−1⋅ºC−1) e. Zinc (c = 25.40 J⋅K−1⋅mol−1) f. Graphite (c = 8.527 J⋅K−1⋅mol−1) g. Gold (c = 25.42 J⋅K−1⋅mol−1) 2. What are the heat capacities of the substances in #1 as J⋅g−1⋅ºC−1 (unless already are)? 3. What is the final temperature of the substances in #1 when 125.5 g samples (of each) evolve 2.50 kJ of heat (initial temperature = 90ºC)? 4. What is the final temperature of the substances in #1 when 125.5 g samples (of each) absorb 2.50 kJ of heat (initial temperature = 10ºC)? 5. If 655 g of the following samples are immersed in 2.5 L of water (assume the density of water = 1.00 g⋅mL−1) initially at room temperature (25 ºC) and each sample is initially at 815 ºC, determine the final temperature of both water and substance. a. Diamond b. Iron c. Brass d. Zinc e. Graphite f. Gold 6. A 0.500 kg block of ice at −15.0 ºC initially is placed in a square room with a wall length of 8.00 feet. If the final temperature of both the water and room are 20.0 ºC, determine the initial temperature of the room. (Assume all heat in the water is absorbed from the room’s air which will behave as an ideal gas ending at a constant pressure of 1 atm). Average molar mass of air is 28.9 g⋅mol−1, cair = 1.0 J⋅g−1⋅ºC−1, cice = 2.06 J⋅g−1⋅ºC−1, cliquid −1
−1
−1
water = 4.18 J⋅g ⋅ºC , and ΔHfus = 6.00 kJ⋅mol . 7. What is ΔHsoln for the compounds listed and, using these, what will the final temperature read in a constant pressure calorimeter (coffee cup calorimeter) if 5.00 g of each substance is dissolved in 75.0 mL of water originally at 22.5 ºC? Assume no heat loss to the surroundings. 1
Exam 3 Worksheet a.
b.
c.
d.
Potassium hydroxide Potassium bromide Sodium chloride Sodium sulfate 8. What is ΔHcomb (in kJ∙mol–1) for each substance if each substance is combusted in a constant volume bomb calorimeter with the water at an initial temperature of 22.02 ºC? The heat capacity of the calorimeter (which includes the water) is 9.87 kJ∙ ºC–1. Determine the error in this value comparing to the real value using appendix 2. Assume all products are gases and the combustion is complete. a. 1.205 g of glucose (C6H12O6) which raises the temperature of water to 23.92 ºC b. 1.750 g of sucrose (C12H22O11) which raises the temperature of water to 24.61 ºC c. 1.850 g of graphite which raises the temperature of water to 27.95 ºC 9. What is the volume at 25 ºC and at 1 atm of pressure of each fuel needed to boil (to completion) 1.50 kg of ice initially at −15.0 ºC? Assume all are complete combustion reactions (in excess oxygen) at 50% efficiency under constant pressure and all fuels behave as ideal gases. ΔHvap(water) = 40.656 kJ⋅mol−1 (plus all other constants in 37). a. Methane (CH4) b. Methanol (CH3OH) c. Ethane (C2H6) d. Acetylene (C2H2) e. Ethylene (C2H4) f. Propane (C3H8) 10. What is the change in enthalpy of reaction for the reactions? Assign whether the process is endothermic or exothermic. a. 2NO2 ( g ) → N2O4 ( g ) b. CaCO3 ( s ) → CaO ( s ) + CO2 ( g ) c. CuO ( s ) + CO ( g ) → Cu ( s ) + CO2 ( g ) d. CH4 ( g ) + 3Cl2 ( g ) → CHCl3 ( l ) + 3HCl ( g ) e. 4NH3 ( g ) + 7O2 ( g ) → 4NO2 ( g ) + 6H2O ( g ) 11. What is the change in enthalpy for the overall reaction given the thermochemical equations? Assign whether the process is endothermic or exothermic. a. BaO ( s ) + H2SO4 ( l ) → BaSO4 ( s ) + H2O ( l )
SO 3 ( g ) + H 2 O ( l ) → H 2SO 4 ( l )
ΔH O = −78.2 kJ
BaO ( s ) + SO 3 ( g ) → BaSO 4 ( s )
ΔH O = −213 kJ
2
Exam 3 Worksheet b. Ca ( OH )2 ( aq ) + 2HCl ( aq ) → CaCl2 ( aq ) + 2H2O ( l )
CaO ( s ) + 2HCl ( aq ) → CaCl2 ( aq ) + H 2O ( l ) ΔH O = −186 kJ
CaO ( s ) + H 2O ( l ) → Ca ( OH )2 ( s )
ΔH O = −65.1 kJ Ca ( OH )2 ( s ) → Ca ( OH )2 ( aq )
ΔH O = −12.6 kJ
c. CuO ( s ) + CO ( g ) → Cu ( s ) + CO2 ( g )
2CO ( g ) + O 2 ( g ) → 2CO 2 ( g )
ΔH O = −566.1 kJ
2Cu ( s ) + O 2 ( g ) → 2CuO ( s )
ΔH O = −310.5 kJ
d. HCl ( g ) + NaNO2 ( s ) → HNO2 ( l ) + NaCl ( s )
2NaCl ( s ) + H 2 O ( l ) → 2HCl ( g ) + Na 2 O ( s ) ΔH O = 507.3 kJ
NO ( g ) + NO 2 ( g ) + Na 2 O ( s ) → 2NaNO 2 ( s ) ΔH O = −427.14 kJ
e.
NO ( g ) + NO 2 ( g ) → N 2 O ( g ) + O 2 ( g )
ΔH O = −42.68 kJ
2HNO 2 ( l ) → N 2 O ( g ) + O 2 ( g ) + H 2 O ( l )
ΔH O = 34.35 kJ
1
1
Cl2 ( g ) + NaBr ( s ) → NaCl ( s ) + Br2 ( l )
2
2
CaO ( s ) + Cl 2 ( g ) → CaOCl2 ( s )
ΔH O = −110.9 kJ
H 2 O ( l ) + CaOCl 2 ( s ) + 2NaBr ( s )
→ 2NaCl ( s ) + Ca ( OH )2 ( s ) + Br2 ( l )
Ca ( OH )2 ( s ) → CaO ( s ) + H 2 O ( l )
ΔH O = −60.2 kJ
ΔH O = 65.1 kJ
12. What is the enthalpy of formation for the substances, given the thermochemical reactions? Assign whether the process is endothermic or exothermic. a. Determine ΔH fO for magnesium nitride. 3Mg ( s ) + 2NH 3 ( g ) → Mg 3 N 2 ( s ) + 3H 2 ( g ) ΔH O = −371 kJ
1
3
N 2 ( g ) + H 2 ( g ) → NH 3 ( g )
ΔH O = −46 kJ
2
2
a. Determine ΔH fO for dinitrogen pentoxide. 2H 2 ( g ) + O 2 ( g ) → 2H 2 O ( l )
ΔH O = −571.5 kJ
N 2 O5 ( g ) + H 2 O ( l ) → 2HNO3 ( l )
ΔH O = −76.6 kJ 1
3
1
N 2 ( g ) + O 2 ( g ) + H 2 ( g ) → HNO3 ( l )
2
2
2
O
b. Determine ΔH f for copper (II) oxide. ΔH O = −174 kJ
2Cu ( s ) + S ( s ) → Cu 2S ( s )
ΔH O = −79.5 kJ
S ( s ) + O2 ( g ) → SO 2 ( g )
ΔH O = −297 kJ Cu 2S ( s ) + 2O 2 ( g ) → 2CuO ( s ) + SO 2 ( g )
ΔH O = −527.5 kJ
3
Exam 3 Worksheet c. Determine ΔH fO for magnesium nitrate. 8Mg ( s ) + Mg ( NO3 )2 ( s ) → Mg3 N 2 ( s ) + 6MgO ( s ) ΔH O = −3884 kJ
Mg3 N 2 ( s ) → 3Mg ( s ) + N 2 ( g )
ΔH O = 463 kJ
2MgO ( s ) → 2Mg ( s ) + O2 ( g )
ΔH O = 1203 kJ
13. What is the enthalpy of formation of the substances given the change in enthalpy of reaction? Assign whether the process is endothermic or exothermic. a. Determine ΔH fO for methanol (CH3OH) given: 2CH3OH ( l ) + 3O2 ( g ) → 2CO2 ( g ) + 4H 2O ( g )
ΔH O = −1199 kJ b. Determine ΔH fO for glycine (C2H5NO2) given: 4C2 H5 NO2 ( s ) + 9O2 ( g ) → 8CO2 ( g ) + 10H2O ( l ) + 2N2 ( g ) which liberates 3896 kJ of heat under constant pressure. c. Determine ΔH fO for formic acid (CH2O2) given HCHO2 ( l ) → CO ( g ) + H2O ( l ) which liberates 33 kJ of heat under constant pressure. d. Determine ΔH fO for baking soda (sodium bicarbonate) given 2NaHCO3 ( s ) → Na 2CO3 ( s ) + H2O ( l ) + CO2 ( g ) which absorbs 85 kJ of heat under constant pressure. 14. What is the formation reaction for each substance? a. C6H5OH(l) b. CO2(g) c. CO(g) d. C8H18(l) e. N2O(g) f. N2H4(g) Chapter 7 – Electronic Structure of Atoms 1. What is the change in energy for the first four lines in the Balmer series for atomic hydrogen (see table 7.1 if needed)? What is the frequency and wavelength of the corresponding photons? 2. What is the change in energy for the first four lines in the Lyman series for atomic hydrogen (see table 7.1 if needed)? What is the frequency and wavelength of the corresponding photons? 3. What are two possible quantum numbers for: a. A 4d electron 4
Exam 3 Worksheet b.
c.
d.
e.
f.
A 5p electron An electron in the third energy level An electron in the fourth energy level A 2s electron A 7f electron 4. Rank the electrons by relative energy within an atom of mercury. i. (5, 0, 0, +½) ii. An electron in a 4d subshell. iii. (4, 0, 0, ‐½) iv.
(3, 2, ‐2, ‐½) v.
(1, 0, 0, ‐½) vi.
(2, 1, ‐1, +½) vii.
(4, 2, 1, ‐½) viii.
An electron in a 3p subshell. ix. (4, 1, 0, +½) x. (3, 0, 0, ‐½) xi. (2, 1, 1, ‐½) xii. (4, 3, ‐3, +½) 5. Are the quantum number sets for an electron in an atom of tin describe the electron in a ground state (G), excited state (E), or forbidden state (F)? a. (6, 3, ‐4, ½) b. (3, 2, 0, ½) c. (4, 3, ‐3, ½) d. (2, 2, 0, ½) e. (1, 0, 0, ½) f. (5, 1, ‐1, ½) g. (5, 2, ‐1, ½) h. (3, 2, ‐3, ½) i. (2, 1, 0, 1) j. (4, 0, ‐1, ½) k. (5, 2, 0, ½) l. (5, 0, 0, ½) Chapter 8 – The Periodic Table 1. What is the ground state electronic configuration for the species, what is the orbital diagram for the highest energy (partially filled in most cases) sublevel, and is the species paramagnetic or diamagnetic? a. Ca b. Ti c. Ti4+ d. Cu 5
Exam 3 Worksheet e.
f.
g.
h.
i.
j.
k.
l.
m.
n.
o.
Fe2+ Cr3+ Zn2+ Pb Pb2+ S S2‐ Ho Am Pd2+ Ag+ 2. For the combinations select the species with the largest radii and the highest ionization energy. a. C, N, and Si b. Ge, P, and Si c. Na+ and Na d. Br and Br– e. S, Cl, and Br 3. Describe in general the periodicity of the periodic table of the elements. Use one group to describe the periodicity within groups (select for instance, group 4) and use one period (3) to describe periodicity within periods. Chapter 9 – Chemical Bonding I: The Covalent Bond 1. For the compounds and elements listed, designate the type of bonding present: ionic (I), covalent (C), both ionic and covalent (IC), or neither (N). a. Na2O b. MgCO3 c. N2O d. NaNO2 e. NO2 f. Al g. K2SO4 h. K2S i. Al2O3 j. SO3 k. Pd l. PCl5 m. Cl2 n. Ne o. NaClO4 p. ICl3 6
Exam 3 Worksheet 2.
3.
4.
5.
6.
7.
q. FeCl3 Give the reaction corresponding to the lattice energy for each ionic compound listed and describe (in general) what is needed to determine the lattice energy in a Born‐
Haber cycle. a. LiCl b. MgCl2 c. KF d. CaF2 Give the Born‐Haber cycle for magnesium chloride. The molar enthalpy of sublimation of magnesium is 147.70 kJ∙mol–1. What is the lattice energy for magnesium chloride? Give the Lewis dot structure, the shape (at the central atom, unless otherwise stated), the hybridization, and the polarity of the following molecules. a. SO2 b. COH2 c. CH3OH (give shape, etc. for carbon and oxygen) d. HCN e. N2O f. SOCl2 g. CF2Cl2 h. HCCH i. H2CCH2 j. N2O4 k. H2C2O4 (this is oxalic acid – the hydrogen atoms are bonded to oxygen) l. Cl2CCH2 m. H2CO3 Give resonance structures for any applicable molecules in number 4, using formal charges – comment on the validity of the structures. Determine the number and type of bonds broken versus formed for the reactions. Assign which are exothermic and which are endothermic. a. Combustion of 1 mol of methanol (CH3OH) b. Combustion of 1 mol of acetylene (C2H2) c. Formation of 2 mol of ammonia d. Combustion of 1 mol of CO e. Formation of 2 mol of hydrogen fluoride f. Formation of 2 mol of hydrogen chloride g. Dimerization of NO2 to N2O4 (making 1 mol of N2O4) Using bond enthalpies, assign whether the reactions in number 6 will be exothermic or endothermic. 7
Exam 3 Worksheet 8. Calculate the reaction enthalpies for the reactions in number 6 using standard molar enthalpies of formation. 9. Comment on any differences between the values in number 7 to number 8. 10. Plot the reactions in number 7 in an enthalpy diagram showing the “intermediate” where all bonds are broken. 11. Do you think any “intermediate” you designated in number 10 is a true intermediate for these reactions? 8