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Chemistry Appendixes
Appendix A-F_Chem20
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Appendix A-F_Chem20
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contents
Chemistry Appendixes
A. Numerical Answers to Questions
783
B. Scientific Problem Solving
790
B.1
B.2
B.3
B.4
Scientific Problem-Solving Model
Investigation Report Outline
Sample Investigation Report
The Nature of Scientific Research
C. Technological Problem Solving
C.1
C.2
C.3
C.4
Technological Problem-Solving Model
Investigation Reports
Laboratory Equipment
Laboratory Processes
D. STS Problem Solving
D.1 STS Decision-Making Model
D.2 Types of Reports
E. Safety Knowledge and Skills
790
790
793
794
796
796
796
797
802
806
806
806
807
E.1 Laboratory Safety
E.2 Safety Symbols and Information
E.3 Waste Disposal
807
809
810
F. Communication Skills
811
F.1
F.2
F.3
F.4
F.5
Scientific Language
SI Symbols and Conventions
Quantitative Descriptions and Rules
Tables and Graphs
Problem-Solving Methods
811
811
813
815
816
G. Review of Chemistry 20
817
H. Diploma Exam Preparation
823
781
Appendix A-F_Chem20
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I. Data Tables
Thermodynamic Properties of Selected Elements
Thermodynamic Properties of Selected Compounds
Miscellaneous Specific and Volumetric Heat Capacities
Standard Molar Enthalpies of Formation
Relative Strengths of Oxidizing and Reducing Agents
Relative Strengths of Aqueous Acids and Bases
J. Common Chemicals
782
826
826
826
826
827
828
829
830
NEL
Appendix A-F_Chem20
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Appendix A NUMERICAL ANSWERS TO QUESTIONS
Chemistry Review Unit
Unit 2
Are You Ready? (pp. 4–5)
Chapter 4
2. (b) 4.2 g
Section 4.1
Lab Exercise 4.1 (p. 148)
Chapter 1
Section 1.4
Section 1.4 Questions (p. 26)
11. 1, 2, 3, 3–, 2–, 1–
Chapter 2
Section 2.4
Section 2.4 Questions (p. 57)
1. (a) 18.02 g/mol
(b) 44.01 g/mol
(c) 58.44 g/mol
(d) 342.34 g/mol
(e) 252.10 g/mol
2. (a) 4 sig.dig.
(b) 2 sig.dig.
(c) 2 sig.dig.
(d) 3 sig.dig.
(e) 1 sig.dig.
(f) 4 sig.dig.
3. (a) 0.117 mol
(b) 24 g
(c) 50.0 mmol
(d) 12.49 g
4. (a) 1.72 103 g or 1.72 kg
(b) 50.0 L
(c) 1.55 mol/L
(d) 13.6 mL
(e) 2%
(f) 3.94 kJ
5. (a) 0.907 mol
(b) 8.56 mol
(c) 29.21 mol
(d) 1.80 mmol
(e) 2.45 mol
6. (a) 71.9 g
(b) 8.96 g
(c) 1.03 g
(d) 1.49 Mg
(e) 0.10 kg
7. (a) 2.02 g, 70.90 g, 92.92 g
(b) 64.10 g, 96.00 g, 88.02 g, 72.08 g
NEL
1, 5, 3, 2, 4
Practice (p. 150)
2. (a) 0.952 atm, 724 mm Hg
(b) 110 kPa, 1.09 atm
(c) 253 kPa, 1.90 103 mm Hg
Practice (p. 152)
6. 263 kPa
7. 137 L
8. (b) 0.17 L.
9. 149 kPa.
Practice (p. 154)
11.–273 °C, 0 K
12. (a) T (0 273) K 273 K
(b) T (100 272) K 373 K
(c) T (30 273) K 243 K
(d) T (25 273) K 298 K
13. (a) t (0 273) °C –273 °C
(b) t (100 273) °C 173 °C
(c) t (300 273) °C 27 °C
(d) t (373 273) °C 100 °C
Practice (p. 156)
14. (a) 16 mL
15. 0.12 L
16. 26%
17. 79 °C
Practice (p. 159)
20. 87.6 kPa
21. 7.9 L
23. 25 °C
Section 4.1 Questions (pp. 161–162)
1. (a) 8.87 kPa, 66.5 mm Hg
(b) 0.247 atm, 188 mm Hg
(c) 112 kPa, 1.11 atm
2. (a) 298 K
(b) 238 K
(c) 39 °C
(d) 65 °C
3. 384 mm Hg
4. 0.16 L
5. (a) 62 L
(b) 2.3 times larger
6. 231 °C
11. (a) 3.82:1
15. (a) 2.8 L
Section 4.2
Practice (p. 166)
5. 25.0 L
6. 0.60 L
7. (a) 1.5 ML
Section 4.2 Questions (p. 168)
4. (a) 124 kL
(b) 124 kL
5. (a) oxygen, 125 L; nitrogen monoxide,
100 L; water vapour, 150 L
(b) 375 L
(c) 33.3 L
Section 4.3
Section 4.3 Questions (p. 171)
5. 186 L
6. 2.0 mmol
7. 50.4 L
8. 0.18 mol
9. 73 mL
10. 9.80 L
11. 0.539 ML (or 539 kL)
12. 0.727 g
13. 2.58 g
14. (a) 1.8 g/L
Section 4.4
Practice (pp. 174–175)
3. 1.6 MPa
4. 41.0 mmol
5. 34 kL or 34 m3
Lab Exercise 4.B (p. 175)
R, 8.48 kPa⋅L/(mol⋅K)
Section 4.4 Questions (p. 176)
5. 5.6 mol
6. 225 °C
atm•L
8. 0.0821 mol•K
9. (a) 34.0 g/mol
10. 22.4 L, 24.8 L
11. (a) 1.1 g/L
(d) 1.74 g/L
Unit 2 Review (pp. 180–183)
18. (a) 273 K
(b) 294 K
(c) 0 K
Numerical Answers to Questions 783
A
Appendix A-F_Chem20
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19. (a) 0.405 MPa
20.
21.
26.
27.
28.
29.
31.
32.
33.
35.
36.
37.
38.
39.
(b) 102 kPa
(c) 45.6 MPa
(a) 0.21 mol
(b) 0.924 mol
(a) 12.4 kL
(b) 1.4 ML
8.23 L
14 °C
0.33 mol
4.73 L
(a) 1.78 atm
(b) 196 kPa
317 °C
(a) 302 kPa
(b) 30.7 kg
(a) ammonia, 1.00 L; oxygen, 1.25 L
(a) 50 mL
methane, 0.647 g/L; nitrogen, 1.13 g/L
(a) 150 mL
(b) 7.5 mmol
(c) 167 mL
Unit 3
Chapter 5
Section 5.3
Practice (pp. 205–206)
2. 7.5% V/V
3. 32% W/V
4. 4.9% W/W
5. 5.4 ppm
6. 1.8 mol/L
Practice (p. 208)
7. 350 mL
8. 7.5 kg
9. 4.1 mol
10. 0.25 mol
11. 403 mL
12. 54 mL
Practice (p. 210)
13. 15.0 g
14. 0.16 kg
(a) 355 mg
(b) 8.07 mmol/L
(a) 7.83% W/V
(b) 1.34 mol/L
Practice (p. 212)
17. (a) [Na(aq)] 0.82 mol/L;
[S2(aq)] 0.41 mol/L
784
Appendix A
Page 784
(b) [Sr2(aq)] 1.2 mol/L;
[NO3(aq)] 2.4 mol/L
(c) [NH4 (aq)] 0.39 mol/L;
[PO43(aq)] 0.13 mol/L
18. [Fe (aq)] 49.6 mol/L;
3
[Cl(aq)] 149 mmol/L
19. (a) 11.1 g
(b) 18.5 g
Section 5.3 Questions (p. 214)
3. (a) 5% W/V
4. 20 ppm
5. (a) 0.32 mol/L
(b) [NH4(aq)] 0.64 mol/L;
[CO32(aq)] 0.32 mol/L
6. 79.0 g
8. (a) 1.30 g
(b) initial volume 10.0 mL
Chapter 5 Review (pp. 231–233)
22. 51 g, 150 g, 250 g
23. 0.3 L
24. (a) 0.70 mol/L
(b) 0.125 mol/L
(c) 2.0 mol/L
(d) 0.66 mmol/L
25. 12.6 g
26. 42.8 mL
27. 6.6 g
28. (a) 56 mg
32. 28.1 mL
34. (a) [K(aq)] 0.14 mol/L;
[NO3(aq)] 0.14 mol/L
7. 11 mg
8. 4.3 mol/L
9. (a) 0.58 g
(b) 0.75 g
(b) [Ca2(aq)] 0.14 mol/L;
[Cl(aq)] 0.28 mol/L
(c) [NH4(aq)] 0.42 mol/L;
[PO43(aq)] 0.14 mol/L
(c) 1.1 g
10. 0.20 L
Chapter 6
11. (a) 0.143 mol/L
Section 6.2
(b) 0.429 mol/L
Practice (p. 239)
1. (a) 107 mol/L
12. 2.3 g
9
17. (a) 1:10
(b) 1011 mol/L
(b) 0.001 ppm
(c) 102 mol/L
1 g solute
(c) kg solution
(d) 104 mol/L
(d) 30 µg/kg
Section 5.4
Practice (p. 216)
1. 3.10 g
2. 200 g
4. (a) 33.2 g
5. (a) 5.93 g
Practice (pp. 218–219)
6. 42%
7. 22.5 mL
8. (a) 0.250 mmol/L
(b) 0.399 mg
Section 5.4 Questions (p. 219)
3. 3.27 g
4. 43.5 L
(e) 1014 mol/L
2. (a) 3
(b) 5
(c) 7
(d) 10
3. 100
Practice (p. 242)
5. (a) 2.68
(b) 5.0
(c) 6.602
(d) 8.14
6. (a) 5 1012 mol/L
(b) 2.2 103 mol/L
(c) 6 105 mol/L
(d) 1.76 1014 mol/L
Practice (p. 243)
9. 0.65
5. 6.85 g/L
10. 1.6 107 mol/L
6. 1.51 g
Section 6.2 Questions (p. 244)
7. 25.0 mL
6. 5.00
NEL
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Page 785
Appendix A
8. increase of 3 pH units
9. (a) 1.00
(b) 14.80
(b) 5.10
39. 2.82 g
49. 0.012 mol/L
Section 6.3
51. 0
Lab Exercise 6.B (p. 247)
56. (b) 4 103 mol/L
Section 6.5
Practice (p. 255)
4. (a) 0.15 mol/L
(b) 0.82
Chapter 6 Review (pp. 263–264)
14. (a) 2.74 1012 mol/L
(b) 3 104 mol/L
15. 10×
17. (a) between 5.4 and 6.0
(b) 1 106 mol/L to 4 106 mol/L
Unit 3 Review (pp. 265–269)
22. (a) 75.0 ppm
(b) 1.87 mmol/L
23. 1.2 L
24. (a) 0.23 mol/L
(b) 20 mmol
(c) 0.103 L
(d) 0.155 mol/L
(e) 1.21 g
25. 16.9 mg/L
26. (a) 14 L
27. (a) [Na(aq)] 4.48 mol/L;
[S2(aq)] 2.24 mol/L
(b) [Fe2(aq)] 0.44 mol/L;
[NO3(aq)] 0.88 mol/L
(c) [K(aq)] 0.525 mol/L;
[PO43(aq)] 0.175 mol/L
(d) [Co3(aq)] 0.0862 mol/L;
[SO42(aq)] 0.129 mol/L
30. (a) 2.12
(b) 2.60
31. (a) 2.74 1012 mol/L
(b) 3 10–4 mol/L
35. (a) 103 mol/L (fruit juice); 1012
mol/L (cleaning solution)
(b) 109:1
37. (a) 6.3
NEL
Practice (p. 302)
1. 0.537 mol/L
2. 375 mL
3. 90.0 mL
Lab Exercise 7.C (p. 302)
4.4 - 4.8; 6.0 - 6.6; 3.2 - 3.8
Section 6.3 Questions (p. 247)
2. (a) < 4.8
(b) > 12.0
(c) > 5.4
(d) between 6.0 and 7.6
3. (a) between 5.4 and 6.0
(b) 1 106 mol/L
Section 7.4
Unit 4
Are You Ready? (pp. 272–273)
3. (a) 2 105 mol/L
4. (NH4)3PO4(s), 0.2951 mol;
CH3COOH(l), 3.5 g
5. NaOH(aq), 1.10 mol; HCl(aq),
0.00345 L; Na2SO4(aq), 1.13 mol/L
6. CH4(g), 0.611; UF6(g), 0.120; CO2(g),
2.48 103 L; Ar(g) 0.161
Chapter 7
Section 7.2
Practice (p. 290)
9. 12 g
10. 66.2 g
11. 36.3 g
12. 6.11 g
13. 0.307 g
14. 2.32 g
Lab Exercise 7.A (p. 291)
6.68 g
Lab Exercise 7.B (p. 293)
9.12 g
Section 7.2 Questions (p. 293)
6. 3.88 g
8. (a) 2.93 g
(b) 95.9%
10. (a) 11 g
(b) 91%
Section 7.3
Practice (p. 296)
1. 16 L
2. 60.1 kL
3. 150 L
Section 7.3 Questions (p. 298–299)
2. 6.0 L
3. 2.62 ML
4. 232 kg
5. 0.21 kL, or 0.21 m3
6. 0.77 L
7. 1.94 kg
8. 0.57 L
210 mg; 0.20 g
Lab Exercise 7.D (p. 303)
0.351 mol/L
Section 7.4 Questions (p. 303)
1. 0.35 L
2. 17.8 mol/L
3. 23.9 mmol/L
4. (b) 624 mg or 0.624 g
(c) 98.0%
5. 639 g
Chapter 7 Review (p. 309–311)
20. (a) NaHCO3(s), 2.4 mol; Na2CO3(s),
1.2 mol; CO2(g), 1.2 mol; H2O(l),
1.2 mol
(b) 0.631 kg
21. 46 mL
2.93 g
23. (a) 95.0%
26. (a) 14.7 g
(b) 10.7 g
28. 2.08 kg
29. 1.91 kg
30. 1.17 kg
33. 0.668 mol/L
Chapter 8
Section 8.2
Lab Exercise 8.A (p. 317)
5.40 g
Section 8.2 Questions (p. 319)
1. 97 g sodium sulfate
2. 0.135 mol/L
3. 0.210 g predicted; 0.20 g obtained; 5%
4. 0.351 mol/L
Section 8.3
Practice (p. 321)
1. (a) 1.1 g to 1.2 g
2. 33.6 mL
Practice (p. 324)
3. (a) 5.0 mol
(b) 0.55 mol
(c) 0.26 mol
(d) 5.46 mmol
Numerical Answers to Questions 785
A
Appendix A-F_Chem20
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4. (a) 6.25 g
(b) 2.86 g
(c) 36.mg
(d) 4.2 kg
Section 8.3 Questions (p. 327)
4. 3 g
5. (a) 2.77 g
(b) 3.89 g
6. 97.9%
7. (b) 2.3 g BaCl2
(c) 8.4 g 0.036 mol
8. (b) 0.053 mol Zn
(c) 83%
9. (b) 125 mL
(c) 138 mL
Section 8.4
Page 786
27. 1.2 102 g
31. 13.0 mL
33. 2.93 g
34. (b) 1.07 g
36. 1.22 g predicted; 1.27 g obtained;
4.5%
Unit 5
Chapter 10
Chapter 8 Review (pp. 346–348)
13. (a) 0.075 mol lead (II) nitrate
(b) 1 mol propane
(c) 0.50 mol zinc
(d) 50 mmol sulfuric acid
14. 79.7%
15. 50 mL
16. 3.0 g
17. (a) 20 mL, 20 mL, 20 mL, 20 mL
(b) 7, 9, 5, 7
19. 0.0212 mol/L
Unit 4 Review (pp. 349–51)
24. (a) 106%
25. 699 kg
26. (a) 0.20 mol Zn(s)
(b) 5.0 mmol Cl2(aq)
(c) 0.05 mol NaOH(aq)
786
Appendix A
(b) 2.07 MJ
5. 0.242 MJ
6. 64 kJ/mol
7. 26 kJ/mol
Practice (p. 431)
Section 11.3
15. (a) 18 mL
Section 11.3 Questions (p. 501)
(b) 18 mL
(c) 18 mL
40. (a) 0.46 kL
Section 8.5 Questions (p. 339)
8. (a) 1.27 g
9. 0.140 mol/L
(c) 4%
3. (a) 7.8 MJ
Section 10.3
Section 8.5
Practice (p. 336)
2. (b) 7
(b) 2.2 kJ
11. 41.1 kJ/mol
Chapter 10 Review (pp. 466-467)
Analysis
(a) 3.00 105 mol
(b) 2.22 106 mol
(c) 2.778 105 mol
(d) 0.0139 mol/L; BAC 0.064 g/100 mL
2. (a) 2.1 kJ
4. 1.54 g
Section 8.4 Questions (p. 332)
1. [NH3] 20.7 mol/L
4. 0.660 mmol
Web Activity: Web Quest—Blood
Alcohol Content (p. 333)
Section 11.2 Questions (p. 494)
20. 36.8%
Unit 5 Review (pp. 119–120 )
(b) 4.6 103 km
42. 78.2%
48. 96.3%; 94.8%
53. 1990 CO2: 96.3%; 2001 CO: 94.8%
Unit 6
3. (a) 241.8 kJ/mol
(b) 318.0 kJ/mol
(c) 81.6 kJ/mol
(d) 372.8 kJ/mol
4. (a) 114 kJ
(c) 114 kJ/mol
(d) 57 kJ/mol
Section 11.4
Practice in 11.4 (pp. 504–505)
1. 851.5 kJ
2. 131.3 kJ
3. 524.8 kJ
4. 205.7 kJ
Chapter 11
Lab Exercise 11.B (p. 506)
Section 11.1
3488.7 kJ/mol; -3.35 MJ/mol; 4.0%
Case Study, (p. 482)
Lab Exercise 11.C (p. 506)
2. 77.8 kJ
Section 11.1 Questions (pp. 483–484)
1. (a) respectively, for 2003: 8.22%,
35.21%, 43.55%, 3.99%, 9.04%
2. (a) respectively, for 2003: 58%, 63%,
56%, 42%, 8%
3. respectively, for 2003: 30.4%, 29.5%,
2.8%, 17.7%, 1.7%, 17.9%
Section 11.2
Practice (p. 487)
205.9 kJ
Section 11.4 Questions (pp. 508–509)
1. 69.5 kJ
2. (a) 136.4 kJ
(b) 44.2 kJ
3. 121.2 kJ/mol
7. 2.39 MJ
8. 2.20 MJ
9. 250.1 kJ/mol
10. 492 kJ
6. 84 kJ
11. 21.8 kJ
7. 0.39 MJ
Section 11.5
8. 67.7%
Lab Exercise 11.D (p. 513)
Practice (p. 492)
725.9 kJ/mol; 598 kJ/mol; 17.6%
12. 5.00 MJ
Section 11.5 Questions (pp. 514–515)
13. 612 kJ
2. (a) 205.9 kJ
Lab Exercise 11.A (p. 492)
(b) 41.2 kJ
54.6 kJ/mol
(c) 91.8 kJ
NEL
Appendix A-F_Chem20
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Page 787
Appendix A
3. (a) 225.5 kJ/mol
(f) 38 kJ
(b) 58.1 kJ/mol
(g) 38 kJ
(c) 23.6 kJ/mol
(h) 38 kJ
5. –4; –2; 0; 2; 4
Section 13.3 Questions (p. 595)
4. (a) C 2; O –2
4. (a) 145.6 kJ/mol
Unit 6 Review (pp. 547–551)
(b) O 0
5. (a) 349.6 kJ/mol
19. (a) 176.2 kJ
(c) N 3; H 1; Cl 1
20. (b) 128.7 kJ
(d) H 1; P 5; O 2
6. (a) 562.0 kJ
8. (a) 179.2 kJ
(b) 64.5 kJ
(c) 114.7 kJ
(d) 114.7 kJ
9. –114.7 kJ
10. 198.7 kJ/mol
12. 2803.1 kJ/mol; 2.80 MJ/mol;
0.143%
Chapter 11 Review (pp. 519–521)
17. (a) 44 kJ
(b) 36 kJ/g
18. (a) 75.1 kJ
(b) 75.1 kJ
(c) 2.87 MJ/mol
19. 44.4 kJ/mol
20. (a) 21.8 MJ
(b) 393 g
(c) 234 L
(c) 4.02 MJ
21. (a) 1164.8 kJ
(b) 388.3 kJ/mol
(c) 1.51 GJ
22. (a) 136.4 kJ
(b) 311.4 kJ
23. (a) 23.5 kJ
28. (a) 98.0 kJ/mol
29. 890.5 kJ/mol; 1560.4 kJ/mol;
2219.9 kJ/mol
25. 69 kJ
26. (a) 401.0 kJ/mol
27. (a) 27.2 kJ
(b) 2.3 kJ
28. 426 kJ; 110 kJ; 602 kJ
29. (a) 2 219.9 kJ/mol
(b) 2 043.9 kJ/mol
(c) 1 564.2 kJ/mol
Chapter 12
Section 12.3 Questions (p. 70)
(f) Na 1; P 5; O 2
8
5. (a) 3
Section 13.4
Lab Exercise 13.C (p. 598)
0.258 mol/L
Practice (p. 598)
3. 8.5 L
4. 0.325 mol/L
5. 11.4 mmol/L
55.48 MJ/kg; 51.88 MJ/kg;
50.33 MJ/kg
Lab Exercise 13.D (p. 599)
62.31 mol/kg; 66.49 mol/kg;
68.01 mol/kg
Investigation 13.4 (p. 603)
1.31 mol/L
30. (b) 134.6 kJ
Analysis: 2.94%
31. (a) 1366.8 kJ
Section 13.4 Questions (p. 600)
36. 25.6 kJ/mol
22. 1 256.2 kJ
24. 107.4 kJ
(e) Na 1; S 2; O 2
4. 1.92 mol/L
5. Analysis: 74.9 mmol/L
Evaluation: 6.4%
Unit 7
6. 25.9 mmol/L
Chapter 13
7. 29.8 mmol/L
Section 13.3
Chapter 13 Review (pp. 606–609)
Practice (p. 585)
29. 4.8 g
1. (a) 4
(b) 7
30. 56.3 mmol/L
33. Analysis: –33 ºC
(c) 6
(d) 6
Chapter 14
(e) –1
Section 14.1
(f) –1
Practice (p. 614)
2. (a) 1
3. 6
(b) 2
Section 14.2
(b) 95 kJ
(c) 4
Practice (p. 631)
(c) 35 kJ
(d) –3
11. (a) 0.77 V
(d) 35 kJ
(e) –2
(b) 0.45 V
(f) 5
(c) 1.23 V
3. (a) 60 kJ
7. 5 kJ/mol; 6 kJ/mol; 11%
Chapter 12 Review (pp. 545–546)
(g) 0
Practice (p. 633)
16. (a) 52 kJ
(h) –3
12. (a) 0.47 V
NEL
(b) 90 kJ
3. (a) 0
(b) 0.50 V
(c) 22 kJ
(b) 0
(c) 0.77 V
(d) 60 kJ
(c) 4
13. 0.34 V
(e) 38 kJ
(d) 2
14. Cu: 3.38 V; Zn: 2.28 V
Numerical Answers to Questions 787
A
Appendix A-F_Chem20
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Lab Exercise 14.A (p. 633)
Cr2O72–: 1.23 V; Pd2: 0.95 V; Tl:
0.34 V; Ti2: 1.63 V
Section 14.2 Questions (pp. 637–638)
10. (a) 0.48 V
12. 0.28 V
15. 1.56 V
Section 14.3
Practice (p. 644)
5. (a) 0.80 V
(b) 1.23 V
6. (a) 1.48 V
(b) 0.00 V
Practice (p. 645)
12. (a) 0.43 V
(b) 0.29 V
13. 1.30 V
Page 788
Unit 7 Review (pp. 667–669)
41. 5.78 g
42. 8.56 mmol/L
45. (a) 1.22 V
(b) 0.80
(c) 0.00 V
46. (a) 1.90 V
(b) 1.23 V
(c) 1.64 V
47. (b) 1.23 V
48. (b) 2.19 V
49. (a) 1.99 V
(b) 590 s (or 9.84 min)
50. 2.98 kA
51. 20.1 min
52. 1.0 kmol/h
54. (b) 1.20 V
Section 14.4
Practice (p. 653)
1. 45 C
2. 3.89 A
3. 7.13 C
4. 3.91 min
Practice (p. 654)
5. 41 mmol
6. 3.16 h
7. 1.2 A
Section 14.4 Questions (p. 657)
1. 2.80 mmol
2. 0.58 t
3. 82.8 min
4. 52.8 kA
5. (a) 1.63 t
(b) 4.76 t
6. 0.174 mol/L
7. 24.42 g
8. 0.102 g
9. (a) 26 g
(b) 25.4 min
10. 21. 3 g
11. 25.0 A
12. 0.766 g; 0.71 g; 7%
14. (c) 483 g
(d) 41.9 g
Chapter 14 Review (p. 665)
18. (d) 0.93 V
21. (d) 0.00 V
23. 50.3 kg
24. 0.125 A
788
Appendix A
Unit 8
Are You Ready? (pp. 672–673)
5. (a) 5.94 mol/L
(b) 0.220 mol/L
(c) 0.0403 mol/L
(d) 0.0446 mol/L
Chapter 15
Section 15.1
Practice (p. 682)
4. (a) 2.00 mol
(b) 70.0%
6. (c) 1.00 L
(d) 60%
Lab Exercise 15.B (p. 686)
0.46
Section 15.1 Questions (pp. 688–689)
3. (b) 14 mol
5. 51
6. 0.11 mol/L
7. (c) 0.200 mol
(d) 0.80 mol
(e) 0.40 mol
(f) [HBr(g)] = 0.100 mol/L;
[H2(g)] = 0.20 mol/L;
[Br2(g)] = 0.20 mol/L
(g) 4.0
8. 54.1
9. 1.5
10. (a) 0.78 mol/L
(b) 0.39 mol
Chapter 15 Review (p. 705–706)
19. 0.0013 mol/ L
20. 0.0403 mol/ L
21. 0.032
26. (a) 0.068
27. (a) 5.07 mol/L
(b) 7.60 mol/L
28. 1.00 mol/L
32. (b) 0.0029 mol/L
Chapter 16
Section 16.1
Practice (p. 716)
1. 2.3 1012 mol/L
2. 3.3 1011 mol/L
3. 2.5 1013 mol/L
4. 7.2 1013 mol/L
5. 1.4 1014 mol/L
6. 1.8 107 %
Practice (p. 718)
7. (a) 1.8 1012; 2.26; 11.74
4 109; 3 106; 8.4
5.0 104; 3.30; 10.70
4.0 104; 2.5 1011; 3.40
8. 14.64; 0.64
9. 0.09 g
Section 16.1 Questions (p. 721)
4. (b) 7.8 106 mol/L; 5.11
5. 7.7 1012 mol/L
6. 7.40
7. 3 106 mol/L
8. 1000 (103)
9. 0.372
10. 0.25 g
11. 4.27
12. 0.016; 14.02
13. 2.42; 11.58
14. 18 mg
Section 16.3
Practice (p. 743)
2. (a) 0.20 mol/L
(c) 1.9 103 mol/L
(d) 2.64
(f) 4.84
5. (a) 1.16%
(b) 1.36 105
6. (a) 3.7%
(b) 1.35 105; 1.4 104
NEL
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Appendix A
7. (b) 3.5 102 mol/L;
3.5 102 mol/L; 1.45; 1.8%
8. (a) 0.26 mol/L; 0.58; 2.6%
9. (a) 8.2 105
Practice (p. 746)
12. 7.7 1010
13. 4.2 1010
Section 16.3 Questions (p. 750)
1. 10.27
2. 11.23
4. 4.77
5. 1.4 103
6. 1.3 1010
7. 4.2 104 mol/L
10. (a) 9.95
(b) 4.25
(c) 3.85
Section 16.4
Practice (pp. 762–763)
13. (a) < 7
NEL
(b) > 7
(c) ~ 7
(d) < 7
Chapter 16 Review (p. 773)
17. (a) 0.10 mol/L; 1.00
(b) 4.2 103 mol/L; 2.38
(c) 9.4 105 mol/L; 4.03
18. (a) 5.8 107
20. 7; 7
21. 10.0 to 11.4; 2.5 103 to
2.5 103 mol/L
22. 0.019%; 8.46; 5.54; 2.9 106 mol/L
Unit 8 Review (pp. 775779)
26. 0.62 mol/L; 0.62 mol/L; 1.38 mol/L;
0.62 mol/L
29. 5.0 to 5.2; 1 105 mol/L to
6 106 mol/L
35. (c) 25 mL; 50 mL
36. 3.5 mol/L; 3.5 mol/L
37. 2.3
38. 0.882 mol/L
39. (a) (i) 7.00; 10 mL
45.
46.
47.
48.
49.
50.
51.
(ii) 13.00
(iii) 12.60
(v) 11.8
(vi) 1.9
1.24%; 1.6 105
6.5 103 %; 4.2 1010
1.5 104 mol/L; 0.58; 0.30%
(b) 3.1 102 mol/L; 1.51; 1.5%
(c) 2.1 1011 mol/L
0.012%; 8.14; 5.86; 1.4 106 mol/L
7.9 1010
(a) 1.5 103; 5.0 102; 1.9 101
Numerical Answers to Questions 789
A
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Appendix B
Page 790
SCIENTIFIC PROBLEM SOLVING
B.1 Scientific Problem-Solving Model
Scientists ask questions and seek concepts to answer these
questions by applying consistent, logical reasoning to describe,
explain, and predict observations, and by doing experiments
to test hypotheses or predictions from these hypotheses. In this
way science progresses using a general model for solving problems and employing specific processes as part of a problemsolving strategy.
Every investigation in science has a purpose; for example:
• to create a scientific concept (a theory, law, generalization, or definition)
• to test a scientific concept
• to use a scientific concept, e.g., in chemical analysis
Once you know the purpose, you need a problem and a general design. For example, if the purpose is to perform a chemical analysis to determine the quantity of a substance, then
possible designs include distillation and precipitation. Once
you choose a design, there are many specific questions that you
might ask, many possible reactants you might choose, and
many other variables you might need to consider.
B.2 Investigation Report Outline
An investigation report is the final result of your problem
solving. Your report should follow the model outlined in
Figure 1. As a further guide, use the information and instructions for the specific processes listed below. The parts of the
investigation report that you are to provide are indicated in
the text in a checklist (Figure 2).
Report Checklist
Purpose
Problem
Hypothesis
Prediction
Design
Materials
Procedure
Evidence
Analysis
Evaluation (1, 2
and/or 3)
Figure 2
Shaded circles indicate the parts you are expected to complete in
a particular investigation report. One or more parts of an
Evaluation may be required, as indicated by the numbers.
Purpose
Problem
Purpose
Hypothesis
and/or
Prediction
Design
Materials
Problem
Prediction
Procedure
Evidence
Analysis
Appendix B
The Problem is a specific question to be answered in the investigation. If appropriate, you should state the question in terms
of manipulated and responding variables. In most cases, the
problem is chosen for you. Only when creating a concept will
the Purpose and the Problem be the same.
Hypothesis
Evaluation
The hypothesis is an (often untested) empirical or theoretical
concept that provides a possible explanation for a natural or
technological phenomenon. Only some kinds of investigations require a hypothesis, such as investigations that test a
hypothesis using a general question as the Problem.
Synthesis
Prediction
Figure 1
A scientific problem-solving model helps to guide your laboratory
work, but does not illustrate the complexity of the work.
790
Although this is usually provided, you will be expected to identify the purpose of an investigation before, during, and after
your laboratory work. Most often, the purpose is to create, test,
or use a chemistry concept.
The Prediction is the expected answer to the Problem
according to a scientific concept (for example, a hypothesis,
theory, law, or generalization) or another authority (for
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Appendix B
example, a reference source or a label on a bottle). Write your
Prediction using the format, “According to [an authority],
[answer to the Problem].” Include your qualitative and quantitative reasoning (based upon the authority) with the
Prediction.
Design
The Design provides a brief overview of the Procedure to
obtain an answer to the Problem. Included in the Design are
reacting chemicals and, if applicable, brief descriptions of
diagnostic tests, variables, and controls. Write your Design
as a paragraph of one to three sentences.
Materials
This section consists of a complete list of all equipment and
chemicals, in columns, including sizes and quantities.
Appendix C.3, page 797, shows laboratory equipment,
including common sizes.
Procedure
The Procedure is a detailed set of instructions designed to
obtain the evidence needed to answer the Problem. Write a list
of numbered steps in the correct sequence, including any
safety and waste disposal instructions (Appendix E, page 807).
Whenever possible, repeat measurements several times.
(Design, Materials, Procedure, and Skills). Evaluation of Lab
Exercises (simulated investigations) will not usually involve
Part 1. Only if you are confident that no major flaws are
present can you proceed to the second part. In Part 2, you
use the results of the experiment to evaluate the Prediction (if
one was made) and the Hypothesis (if there is one). This
assumes that a prediction is being tested in an experiment.
If it is the experimental design that is being tested, then the
two parts of the evaluation would be reversed and the percent
difference is used to judge the success of the design. The last
part of the Evaluation, Part 3, comes back, full circle, to the
Purpose of the investigation. Was the Purpose fulfilled?
The parts of the Evaluation you will be expected to complete are shown in parentheses after Evaluation in the Report
Checklist.
Write your Evaluation in paragraph form, using the topic
sentences suggested below or an adaptation of them. Some
of the more important criteria for a judgment are listed as
questions; use selected questions to guide your judgments.
Show as much independent, critical, and creative thought as
possible in support of your judgments.
Part 1. Evaluation of the Investigation
•
Evidence
The Evidence includes all qualitative and quantitative observations relevant to answering the Problem. Organize your
evidence in tables whenever possible (Appendix F.4,
page 815). Be as precise as possible in your measurements
and include any unexpected observations that may affect your
answer and its certainty (in significant digits). Scientific honesty demands that you report all evidence collected and not
just the evidence you think is correct or “normal.”
Were you able to answer the Problem using the
chosen experimental design? Are there any obvious
flaws in the design? What alternative designs (better
or worse) are available? As far as you know, is this
design the best available in terms of controls,
variables, efficiency, and safety?
•
Evaluation
Part 1 of the Evaluation of an investigation that you actually
perform usually involves judging the validity of the experiment
NEL
“The materials are judged to be adequate/
inadequate because ...”
Did you have all of the necessary materials? Was the
equipment of reasonable quality? What materials
could be improved to obtain better results?
Analysis
The Analysis includes manipulations, interpretations, and
calculations based on the evidence. Tables and graphs that
include or facilitate interpretations and calculations are
included in the Analysis. You may need to differentiate between
relevant and irrelevant observations. Communicate your
work clearly and logically. Conclude the Analysis with a statement of your experimental answer to the Problem, including
a phrase such as, “According to the evidence gathered in this
experiment, [answer to the Problem statement].”
“The design of the investigation [name or describe
in a few words] is judged to be adequate/
inadequate because …”
•
“The procedure is judged to be adequate/
inadequate because …”
Were the steps that you used in the laboratory
correctly sequenced, and adequate to gather sufficient
evidence? What improvements could be made to the
procedure, such as more trials?
•
“The technological skills are judged to be
adequate/inadequate because …”
Which specialized skills, if any, might have the
greatest effect on the evidence gathered? Was the
evidence from repeated trials reasonably similar?
Scientific Problem Solving 791
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Page 792
How can the evidence gathered be improved?
•
“Based upon my evaluation of the experiment, I am
not/moderately/very certain of my experimental
evidence. The sources of uncertainty or error are …”
State the sources and your confidence in the
experimental evidence. What would be an
acceptable total of the experimental error (1%, 5%,
or 10%)?
Part 2. Evaluation of the Prediction and
Authority Being Tested
•
If applicable, “the percent difference between the
experimental result and the predicted value is…”
% difference experimental value predicted value 100
predicted value
How does this difference compare with your
estimated total uncertainty or experimental error?
•
“The prediction is judged to be verified/
inconclusive/falsified because …”
Does the predicted answer clearly agree with the
experimental answer reported in your analysis? Can
the percent difference be accounted for by the
sources of error (percent error) listed above?
•
“The authority being tested [name the authority or
hypothesis (reasoning)] is judged to be acceptable/
unacceptable in this experiment because …”
Was the prediction verified, inconclusive, or
falsified? How confident do you feel about your
judgment?
Part 3. Revisiting the Purpose
Did you accomplish the Purpose of this investigation?
Is there a need for additional investigations to better
achieve the Purpose?
Notes on Data and Evidence
Data
Data may be found on data sheets, in data tables, and in databases. Data from one of these sources can become evidence
with the purpose to create, test, or use a scientific concept.
Evidence is data with a scientific purpose.
Common Sources of Experimental Error
• conditions (e.g., SATP) not controlled
• impure reactants or products
• any measurement process
• incomplete reaction
• judgment of colour (e.g., indicator)
• loss of solid in a filtration (stuck to glass or passed
through filter)
Experimental Error and Percent Difference
Some people and books use the term “percent error” in place
of “percent difference.” In this textbook, we use two percentages: One is an estimate of the total expected experimental
error (Evaluation, Part 1), and the other is the actual difference (the percent difference) you determined based on your
prediction and analysis (Evaluation, Part 2). The crucial point
is how these two percentages compare. No experiment can
ever be expected to be perfect. For example, if the equipment
you used is only precise to / 5%, then any percent difference you obtain that is less than or equal to 5% is as good
a result as can be expected. If the percent difference is larger
than the reasonably acceptable experimental error, then the
prediction is falsified.
If the percent difference is equal to or less than the experimental error, then the prediction is verified.
• incomplete drying of a product
• manipulative skill
Do not use “human error” as a source of uncertainly or experimental error.
Percent Yield versus Percent Difference
Replication
An authority may be judged unacceptable in one experiment.
This does not mean the authority is immediately discarded.
Replication by independent workers is always required to
refute any accepted theory.
In most experiments, a percent difference is a measure of
accuracy. In some experiments in which a product is
collected and measured, a percent yield is used instead of a
percent difference.
actual quantity obtained
% yield predicted (maximum) quantity 100
792
Appendix B
NEL
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Appendix B
B.3 Sample Investigation Report
The Reaction of Hydrochloric Acid
with Zinc
Purpose
The purpose of this investigation is to test one of the ideas
of the collision–reaction theory.
Report Checklist
Purpose
Problem
Hypothesis
Prediction
Design
Materials
Procedure
Evidence
Analysis
Evaluation (1, 2, 3)
Problem
2. Carefully place a piece of Zn(s) into the hydrochloric
How does changing the concentration of hydrochloric acid
affect the time required for the reaction of hydrochloric acid
with zinc?
3. Measure and record the time required for all of the
Prediction
According to the collision–reaction theory, if the concentration of hydrochloric acid is increased, then the time required
for the reaction with zinc will decrease. The reasoning that
supports the prediction is that a higher concentration produces
more collisions per second between the hydrochloric acid
entities and the zinc atoms. More collisions per second would
produce more reactions per second and, therefore, a shorter
time required to consume the zinc.
Design
Different known concentrations of excess hydrochloric acid
react with the same quantity of zinc metal. The time for the
zinc to completely react is measured for each concentration
of acid solution. The concentration of hydrochloric acid is
the manipulated variable and time is the responding variable. The temperature, mass, and surface area of zinc, and
volume of acid are the controlled variables.
acid solution and note the starting time of the
reaction.
zinc to react.
4. Repeat steps 1 to 3 using 1.5 mol/L, 1.0 mol/L, and
0.5 mol/L HCl(aq).
5. Neutralize the acid with a weak base and then pour it
down the sink with the water running.
Evidence
Gas bubbles formed immediately on the surface of the zinc
strip when it was placed into the hydrochloric acid solution.
The bubbles appeared to form more rapidly when the concentration of the acid was higher.
The Reaction of HCl(aq) with Zn(s)
Concentration of HCl(aq)
(mol/L)
Materials
Procedure
1. Transfer 10 mL of 2.0 mol/L HCl(aq) into an
18 150 mm test tube. Avoid contact with skin, eyes,
clothing, or the desk. If you spill this acid on your
skin, wash immediately with lots of cool water.
NEL
2.0
70
1.5
80
1.0
144
0.5
258
Analysis
The Reaction of
HCl(aq) with Zn(s)
300
Time (s)
lab apron
eye protection
(4) 10 mL graduated cylinders
(4) 18 150 mm test tubes and test-tube rack
clock or watch (precise to the nearest second)
four pieces of a zinc metal strip (5 mm 5 mm)
HCl(aq): 2.0 mol/L, 1.5 mol/L, 1.0 mol/L, 0.5 mol/L
a weak base (e.g., baking soda)
Time for reaction
(s)
200
100
0
0.5
1.0
1.5
2.0
Concentration of HCl(aq) (mol/L)
Scientific Problem Solving 793
B
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Page 794
According to the evidence obtained, increasing the concentration of hydrochloric acid decreases the time required for
the complete reaction of a fixed quantity of zinc.
Evaluation
The design, reacting zinc with excess hydrochloric acid, is
judged to be adequate because this experiment produced the
type of evidence needed to answer the problem with a high
degree of certainty. There are no obvious flaws in this design.
In my judgment, the design is efficient and safe, and all necessary variables are controlled.
The materials appear to be adequate because the quality of
the evidence was sufficient to give a clear answer to the problem.
Using a pipette would provide better precision than a graduated cylinder but this would not change the overall result.
The procedure is also judged adequate because it produced
sufficient evidence. Possible improvements include extending
the range of concentrations and performing more trials for
each concentration. The reaction could be done in a container that allowed better mixing.
The technological skills are judged adequate because no
specialized skills were involved. Timing the start of the reac-
tion may have some uncertainty but the lack of multiple trials
makes this difficult to judge.
Based upon my evaluation of the experiment, I am very
certain about the experimental results. Sources of uncertainty
in this investigation include: measurement errors for volume
and time, the purity and surface area of the zinc metal strip,
the concentration of the acid, and a little uncertainty in estimating when the last bit of zinc had reacted. An estimate of
the total effect of all experimental uncertainties is about 5%.
The prediction is judged to be verified because the qualitative observations and the graph clearly indicate that the
reaction time decreases as the concentration increases. There
is little deviation from a smooth curve in the graphed results.
The collision–reaction theory is judged to be acceptable in
this experiment because the prediction was clearly verified. I
am quite confident in this judgment because other groups
in the class obtained similar results.
The purpose of this investigation—to test one idea of the
collision–reaction theory—was accomplished but only for
one reaction. Replication with many other reactions would
need to be investigated to have a more valid test.
B.4 The Nature of Scientific Research
Citizens in a democratic society are often required to read
and interpret media reports of scientific research. Health and
environment research reports are, for example, commonly
portrayed in the media. Sometimes the research reports appear
to contradict each other and sometimes the reports promote
more uncertainty than certainty. Understanding the terminology and concepts for describing a research study is increasingly important for responsible citizenry. Listed below are
some of the terms and concepts that will help you both answer
questions in this textbook and understand and critique media
reports of research.
Types of Studies
correlational study—the connection or degree of agreement
(e.g., –0.3, 1.0) is sought between two variables, often
without controlling for other variables; correlational
studies often lead to cause-and-effect studies
cause-and-effect study—one variable is manipulated and all
other variables, other than the responding variable, are
controlled
control experiment—see cause-and-effect study
clinical trial—a controlled study involving people; usually a
final-stage, double-blind study
794
Appendix B
Design Factors
term of study—the duration of the experiment e.g.,
observations over 5 s, 30 min, 3 mon or 15 a; long-term
studies are most often preferred
sample size—the number of entities or people in a study;
generally large sample sizes are preferred
random sample—one chosen randomly from the
population of entities (to reduce bias)
replication—repetition of a study, generally, by an
independent research group
placebo—in medicinal experiments, an inactive item (e.g.,
sugar pill) or treatment given to the control group
placebo effect—a beneficial effect arising from a patient’s
expectations; present in both the control group and the
experimental group
single blind—the subject (e.g., patient) does not know
whether she/he has received the treatment or a placebo,
but the experimenter knows
double blind—neither the subject (e.g., patient) nor the
directly involved experimenter knows whether the
subject has received the treatment or a placebo
control—a standard or comparison value, or procedure
(e.g., leaving one of several identical samples unaltered
for comparison), or a placebo
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Appendix B
control group—a comparison group that does not receive
the experimental treatment
experimental group—a group that receives the experimental
treatment
ways of knowing—methods used to obtain knowledge or
information; examples include traditional (Aboriginal),
empirical, theoretical, referenced, and memorized
Scientific Attitudes
Nature of Evidence and Results
anecdotal—based upon personal experience or hearsay
reliable—reproducible or consistent time after time
valid—judged to be supported by adequate designs,
materials, procedures, and skills
accurate—judged to be true or agreeing with the accepted
value
precise—closely related or very similar; related to
reproducibility of results
statistical bias—a sampling or testing error caused by
systematically favouring some outcomes over others
random result—a result that could be expected on the basis
of probability (e.g., 50% heads and tails when flipping a
large number of coins)
coincidence—a result that is accidental and irrelevant to the
study
significant difference—a difference that is greater than could
be randomly expected when an experimental group and
a control group are compared
certainty—the degree to which something is accepted by an
individual or community (e.g., the evidence may have a
high or low degree of certainty); measured by, for
example, counting significant digits
scientific attitude—a disposition or demonstration of
feelings or thoughts (e.g., honesty, objectivity,
willingness to change, respect for evidence, critical
mindedness, suspended judgment, open-mindedness,
and questioning predisposition)
tolerance of uncertainty—the degree to which people and
institutions tolerate uncertainty (without claiming
absolute certainty), although they strive for greater and
greater certainty
COMMUNICATION example 1
Create an experimental design to test a new drug to promote weight loss.
Solution
Randomly selected control and experimental groups of 1000
volunteers are studied over two years. Both receive pills: The
experimental group receives the drug; the control group
unknowingly receives a placebo. Technicians, who do not
know which group each person is in, record the weight of
the subjects every month. Experimenters, who never meet or
see the volunteers, analyze the evidence gathered.
Reporting Research
refereed journal—an academic journal for which research
papers are sent to subject experts to determine whether
the report is of sufficient quality to publish; also called
peer-reviewed journal
abstract—a short summary describing the research
processes and results
Science–Technology–Society (STS) Issues
risk–benefit analysis—a process of gathering and analyzing
evidence that leads to decision making (and to an
evaluation of the process itself)
stewardship—actively supervising and managing an entity
or event (e.g, the environment)
perspective—a point of view or way of analyzing an object
or event
multiperspective—based upon positive and negative
evidence and arguments from many perspectives (e.g.,
scientific, technological, economic, environmental,
political, legal, ethical, social, and emotional)
NEL
COMMUNICATION example 2
Act as a referee (peer-reviewer) to critique the following
experimental design including, if necessary, suggestions for
improvement.
Ten volunteers are provided with their horoscope for the test
day. The volunteers orally respond in a group to the
question: Does this horoscope describe your personality and
life situation?
Solution
This experimental design is very inadequate because:
•
ten volunteers is a very small sample size; the number
needs to be at least 100 or more, randomly selected from
thousands
•
to control the horoscope variable, all subjects should be
provided with the same horoscope
•
to control subject interaction and influence, subjects
should be isolated from one another
•
to provide for accurate reporting of the subjects’
responses, investigators should make audio or audiovisual
recordings
Scientific Problem Solving 795
B
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TECHNOLOGICAL PROBLEM SOLVING
The goal of technological problem solving is to solve practical
problems by developing or revising a product or a process.
The product or the process must fulfill its function, but it is not
essential to understand why or how it works. Products are
evaluated based on criteria such as simplicity, reliability, and
cost. Technological processes are also evaluated by their efficiency. Technological products and processes may have both
intended and unintended consequences. Therefore, it is important that various perspectives, such as ecological, economic, and
political, are used in any assessment. For example, chlorofluorocarbons may be simple and inexpensive to make, and they
may be useful for a particular function, but their effect on the
ozone layer in the upper atmosphere must also be considered.
Processes such as the chlorine bleaching of wood pulp may
be efficient, but they may adversely affect an ecosystem. We
often look to technological fixes as solutions for problems.
Ecological, economic, political, legal, ethical, and/or social
efforts by individuals or groups can often lead to more sustaining solutions than quick technological fixes.
Chemistry has always been closely associated with technology. Part of technology is the laboratory equipment,
processes, and procedures used in both chemical and technological research and development. In modern chemistry,
simple equipment and processes, such as beakers and filtration, are still used but chemistry also depends on sophisticated technology, such as computers, to store and manipulate
the evidence collected.
C.1 Technological Problem-Solving Model
Technological problem solving is similar in some ways to scientific problem solving but its purpose differs. A characteristic of technological problem solving is a systematic,
trial-and-error manipulation of variables (Figure 3). Variables
are predicted and tested and the results are evaluated. When
the cycle is repeated many times, the most effective set of conditions can be determined. Compare this model with the scientific problem-solving model in Figure 1, on page 790.
Technological problem-solving contexts include industrial,
commercial, and consumer. The general process for technological problem solving is similar in all of these contexts.
Technological problem solving is very common to us in our
everyday lives. Learning more about this systematic trial-anderror approach can help us on an everyday basis.
Technological
Problem
Prediction
of Variables
Evaluation
Product/Process
Design
Product/Process
Analysis
Evidence
Synthesis
Figure 3
A technological problem-solving model
C.2 Investigation Reports
Investigation reports for technological problem-solving investigations can use the same headings as those for scientific
problem-solving investigations (Appendix B.2, page 790).
Some key differences between these two types of reports are
Purpose
The purpose will be to solve a specific, practical problem
by developing or revising a product or process.
796
Appendix C
Evaluation
Evaluation criteria can be many and varied. In some cases,
it will be sufficient to demonstrate that the product or
process works for the materials used. You may also be asked
to judge the simplicity, reliability, and efficiency of the
product or process, recognize its value and limitations, and
evaluate it from a variety of perspectives (Appendix D.2,
page 806).
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C.3 Laboratory Equipment
10 mL
25 mL
50 mL
100 mL
500 mL
1000 mL
13 × 100 mm
18 × 150 mm
25 × 200 mm
meniscus finder
pipette bulb
mortar and pestle
U-tube
test tube
1 mL
10 mL
25 mL
1 mL
10 mL
C
100 mL
250 mL
500 mL
1000 mL
2000 mL
beaker
50 mL
100 mL
150 mL
250 mL
400 mL
600 mL
1000 mL
graduated cylinder
dropper
volumetric
pipette
graduated
pipette
125 mL
250 mL
500 mL
1000 mL
volumetric
flask
Erlenmeyer
flask
50 mL
burette
dropper
bottles
watch glass
burette (utility) clamp
funnel
thermometer
wash bottle
evaporating dish
clamp holder
extension clamp
test-tube clamp
funnel rack
well plate
(microplate)
crucible tongs
wire gauze
laboratory scoop
beaker tongs
Figure 4
Common lab equipment
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weighing boat
Glassware is breakable and should always be handled with care.
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Using a Laboratory Burner
The procedure outlined below should be practised and memorized. Note the safety caution. You are responsible for your
safety and the safety of others near you.
1. Turn the air and gas adjustments to the off position
(Figure 5).
2. Connect the burner hose to the gas outlet on the
bench.
3. Turn the bench gas valve to the fully on position.
4. If you suspect that there may be any gas leaks, replace
the burner. (Give the leaky burner to your teacher.)
5. While holding a lit match above and to one side of
the barrel, open the burner gas valve until a small
yellow flame results (Figure 6). If a striker is used
instead of matches, generate sparks over the top of
the barrel (Figure 7).
6. Adjust the air flow and obtain a pale blue flame with
a dual cone (Figure 8). In most common types of
laboratory burners, rotating the barrel adjusts the air
intake. Rotate the barrel slowly. If too much air is
added, the flame may go out. If this happens,
immediately turn the gas flow off and relight the
burner following the procedure outlined above. If
your burners have a different kind of air adjustment,
revise the procedure accordingly.
7. Adjust the gas valve on the burner to increase or
decrease the height of the blue flame. The hottest part
of the flame is the tip of the inner blue cone. Usually
a 5 to 10 cm flame, which just about touches the
object heated, is used.
barrel
air valve
gas supply
gas
valve
8. Laboratory burners, when lit, should not be left un-
attended. If the burner is on but not being used,
adjust the air and gas intakes to obtain a small yellow
flame. This flame is more visible and, therefore, less
likely to cause problems.
When lighting or using a laboratory burner, never
position your head or fingers directly above the
barrel. Tie back long hair and sleeves.
Figure 5
The parts of a common laboratory burner
Figure 6
A yellow flame is a relatively cool flame
and is easier to obtain than a blue flame
when lighting a burner. A yellow flame is
not used for heating objects because it
contains a lot of black soot.
798
Appendix C
Figure 7
To generate a spark with a striker, pull up
and across on the side of the handle
containing the flint.
Figure 8
A pale, almost invisible flame is much
hotter than a yellow flame. The hottest
point is at the tip of the inner blue cone.
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Appendix C
Using a Laboratory Balance
There are two types of balances: electronic (Figure 9)
and mechanical (Figure 10).
All balances must be handled
carefully and kept clean.
Always place chemicals into
a container such as a beaker
or plastic boat to avoid contamination and corrosion of
the balance pan. To avoid
error due to convection cur- Figure 9
An electronic balance
rents in the air, allow hot or
cold samples to return to
room temperature before placing them on the balance. Always
record masses showing the correct precision. On a centigram
balance, mass is measured to the nearest hundredth of a gram
(0.01 g). When it is necessary to move a balance, hold the
instrument by the base and steady the beam. Never lift a balance by the beams or pans.
To avoid contaminating a whole bottle of reagent, do not
scoop diectly from the original container of a chemical. Pour
a quantity of the chemical into a clean, dry beaker or bottle,
from which samples can be taken. Another acceptable technique for dispensing a small quantity of chemical is to rotate
or tap the chemical bottle.
displayed. Air currents or the high sensitivity of the
balance may cause the last digit to vary.
4. Remove the container and sample.
There is a video demonstration of this technique on the
Nelson Web site.
www.science.nelson.com
GO
C
Using a Mechanical Balance
Different kinds of mechanical balances are shown in
Figures 10(a) and (b). Some general procedures apply to most
of them.
1. Clean and zero the balance. (Turn the zero
adjustment screw so that the beam is balanced when
the instrument reads 0 g and no load is on the pan.)
2. Place the container on the pan.
3. Move the largest beam mass one notch at a time until
the beam drops, and then move the mass back one
notch.
4. Repeat this process with the next smaller mass and
continue until all masses have been moved and the
beam is balanced. If you are using a dial type balance,
the final step will be to turn the dial until the beam
balances, as shown in Figure 10(c).
5. Record the mass of the container.
Using an Electronic Balance
Electronic balances are sensitive to small movements and
changes in level; do not lean on the counter when using the
balance.
1. Place a container or weighing paper on the balance.
2. Reset (tare) the balance so the mass of the container
registers as zero.
(a)
3. Add chemical until the desired mass of chemical is
(b)
6. Set the masses on the beams to correspond to the
total mass of the container plus the desired sample.
7. Add the chemical until the beam is once again
balanced.
8. Remove the sample from the pan and return balance to
the zero position.
(c)
Figure 10
(a) On this type of mechanical balance, the sample is balanced by moving masses on several beams.
(b) Another type of mechanical balance has beams for the larger masses and a dial for the final adjustment.
(c) The dial reading on this balance with a vernier scale is 2.34 g. To read the hundredth of a gram, look below the zero on the vernier,
and then look for the line on the vernier that lines up best with a line on the dial.
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Using a Multimeter
5. Rinse the probes with pure water before testing
A multimeter (Figure 11) is a device that measures a variety
of electrical quantities, such as resistance, voltage, and current.
6. Shut off the meter by using either the on/off switch or
(a)
(b)
another sample.
by turning the dial to any setting other than
“Resistance.”
Voltage Measurements of Cells and
Batteries
1. Set the dial to the appropriate value on the direct
current volts (DCV) scale; for example, 3 V.
2. The black lead (labelled negative or COM) is
normally connected to the anode, and the red lead
(positive) is connected to the cathode of a voltaic cell.
3. Make a firm contact between each metal probe and
an electrode of the cell. (Press firmly with the pointed
probe or use leads with an alligator clip.)
Figure 11
(a) An analog meter has a needle that moves in front of a
labelled scale.
(b) A digital meter gives a direct reading with appropriate units.
4. On analog meters (those with a needle), read the scale
Conductivity Measurements of Solutions
or a digital meter registers a negative number, then
switch the connections to the cell.
1. Set the dial on the meter to one of the higher values on
the ohm (Ω) scale; for example, R 100 or R 1 K.
2. Touch the two metal probes together to check the
battery. If the needle does not deflect significantly
(more than one-half scale), have your teacher adjust
the meter or replace the battery.
3. Test a sample of pure water as a control and note the
movement of the needle.
4. Test your aqueous sample and record the deflection
of the needle according to your teacher’s
instructions.
corresponding to the meter value you set in step 1.
5. If the needle attempts to move to the left off the scale
Using a Pipette
A pipette is a specially designed glass tube used to measure precise volumes of liquids. There are two types of pipettes and a
variety of sizes for each type. A volumetric pipette (Figure 12)
transfers a fixed volume, such as 10.00 mL or 25.00 mL, accurate to within 0.04 mL. A graduated pipette (Figure 13) measures a range of volumes, just as a graduated cylinder does. A
10 mL graduated pipette delivers volumes accurate to within
0.1 mL. There is a video demonstration of this technique on
the Nelson Web site.
www.science.nelson.com
GO
Figure 12
A volumetric pipette delivers the volume printed on the label if the temperature is near room temperature.
Figure 13
To use a graduated pipette, you must be able to start and stop the flow of the liquid.
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Appendix C
1. Rinse the pipette with small volumes of distilled water
using a wash bottle, and then with the sample
solution.
A clean pipette has no visible residue or liquid drops
clinging to the inside wall. Rinsing with aqueous
ammonia and scrubbing with a pipe cleaner might be
necessary to clean the pipette.
2. Hold the pipette with your thumb and fingers near
the top. Leave your index finger free.
3. Place the pipette in the sample solution, resting the tip
on the bottom of the container if possible. Be careful
that the tip does not hit the sides of the container.
4. Squeeze the bulb into the palm of your hand and
place the bulb firmly and squarely on the end of the
pipette (Figure 14) with your thumb across the top of
the bulb.
5. Release your grip on the bulb until the liquid has
risen above the calibration line.
This may require bringing the level up in stages:
remove the bulb, put your finger on the pipette,
squeeze the air out of the bulb, re-place the bulb, and
continue the procedure.
7. Wipe all solution from the outside of the pipette
using a paper towel.
8. While touching the tip of the pipette to the inside of a
waste beaker, gently roll your index finger (or squeeze
the valve of the dispensing bulb) to allow the liquid
level to drop until the bottom of the meniscus reaches
the calibration line (Figure 16).
To avoid parallax errors, set the meniscus at eye level.
Stop the flow when the bottom of the meniscus is on
the calibration line. Use the bulb to raise the level of
the liquid again if necessary.
9. While holding the pipette vertically, touch the pipette
tip to the inside wall of a clean receiving container.
Remove your finger or adjust the valve and allow the
liquid to drain freely until the solution stops flowing.
10. Finish by touching the pipette tip to the inside of the
container held at about a 45° angle (Figure 17). Do
not shake the pipette. The delivery pipette is
calibrated to leave a small volume in the tip.
Never use your mouth to draw a liquid up a pipette.
Always use a pipette bulb.
6. Remove the bulb, placing your index finger over the
top.
If you are using a dispensing bulb (Figure 15), it
remains attached to the pipette.
Figure 14
Release the bulb slowly.
Pressing down with your thumb
placed across the top of the
bulb maintains a good seal.
Setting the pipette tip on the
bottom slows the rise or fall of
the liquid.
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Figure 15
A dispensing pipette bulb
uses a small valve in the side
stem to control the flow of
liquid in a pipette.
Figure 16
To allow the liquid to drop slowly
to the calibration line, it is
necessary for your finger and the
pipette top to be dry. Also keep
the tip on the bottom to slow
down the flow.
Figure 17
A vertical volumetric pipette is
drained by gravity and then the
tip is placed against the inside
wall of the container. A small
volume is expected to remain in
the tip.
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C.4 Laboratory Processes
The processes or experimental procedures listed below are
part of common designs used in scientific or technological
laboratories.
Crystallization
Crystallization is used to separate a solid from a solution by
evaporating the solvent or lowering the temperature.
Evaporating the solvent is useful for quantitative analysis of
a binary solution; lowering the temperature is commonly
used to purify and separate a solid whose solubility is temperature-sensitive. Chemicals that have a low boiling point or
decompose on heating cannot be separated by crystallization
using a heat source.
Filtration
In filtration, solid is separated from a mixture using a porous
filter paper. The more porous papers are called qualitative
filter papers. Quantitative filter papers allow only invisibly
small particles through the pores of the paper. There is a video
demonstration of this technique on the Nelson Web site.
www.science.nelson.com
GO
1. Set up a filtration apparatus (Figure 19): stand,
funnel holder, filter funnel, waste beaker, wash bottle,
and a stirring rod with a flat end for scraping.
Figure 19
The tip of the
funnel should
touch the inside
wall of the
collecting beaker.
1. Measure the mass of a clean beaker or evaporating
dish.
2. Place an accurate volume of the solution in the
container.
3. Set the container aside to evaporate the solution
slowly, or warm the container gently on a hot plate or
with a laboratory burner.
4. When the contents appear dry, measure the mass of
the container and solid (Figure 18).
2. Fold the filter paper along its diameter, and then fold
it again to form a cone. A better seal of the filter paper
on the funnel is obtained if a small piece of the
outside corner of the filter paper is torn off
(Figure 20).
Figure 18
When the substance has crystallized, it may appear dry but
small quantities of water may still be present. To be certain
the solid is dry, it must be heated until the mass becomes
constant.
(a)
(b)
(c)
(d)
Figure 20
To prepare a filter paper, fold it in half twice, and then remove the
outside corner as shown.
3. Measure and record the mass of the filter paper after
5. Heat the solid with a hot plate or burner, cool it, and
measure the mass again.
6. Repeat step 5 until the final mass remains constant.
(Constant mass indicates that all of the solvent has
evaporated.)
802
Appendix C
removing the corner.
4. While holding the open filter paper in the funnel, wet
the entire paper and seal the top edge firmly against
the funnel.
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Appendix C
5. With the stirring rod touching the spout of the
Figure 22
Raise the meniscus
finder along the back
of the neck of the
volumetric flask until
the meniscus is
outlined as a sharp,
black line against a
white background.
beaker, decant most of the solution into the funnel
(Figure 21). Transferring the solid too soon clogs the
pores of the filter paper. Keep the level of liquid about
two-thirds up the height of
the filter paper. The
stirring rod should be
rinsed each time it is
removed.
Figure 21
Pouring along the stirring rod
prevents drops of liquid from
going down the outside of the
beaker when you stop
pouring.
1. (Prelab) Calculate the required mass of solute from
the volume and concentration of the solution.
2. Measure the required mass of solute in a clean, dry
beaker or weighing boat. (Refer to “Using a
Laboratory Balance” on page 799.)
3. Pour less than one-half of the final volume of pure
6. When most of the solution has been filtered, pour the
remaining solid and solution into the funnel. Use the
wash bottle and the flat end of the stirring rod to
clean any remaining solid from the beaker.
7. Rinse the stirring rod and the beaker.
8. Wash the solid two or three times to ensure that no
solution is left in the filter paper. Direct a gentle
stream of water around the top of the filter paper.
9. When the filtrate has stopped dripping from the
funnel, remove the filter paper. Press your thumb
against the thick (three-fold) side of the filter paper
and slide the paper up the inside of the funnel.
water into a beaker. Transfer the solute to the water.
Stir to dissolve.
4. Transfer the solution and all water used to rinse the
equipment into a clean volumetric flask. (The beaker
and any other equipment should be rinsed two or
three times with pure water.)
5. Add pure water, using a medicine dropper for the
final few millilitres while using a meniscus finder to
set the bottom of the meniscus on the calibration
line.
6. Stopper the flask and mix the solution by slowly
inverting the flask several times.
10. Transfer the filter paper from the funnel onto a
labelled watch glass and unfold the paper to let the
precipitate dry.
11. Determine the mass of the filter paper and dry
precipitate.
Preparing a Standard Solution by
Dilution
There is a video demonstration of this technique on the
Nelson Web site.
www.science.nelson.com
Preparation of Standard Solutions
Laboratory procedures often call for the use of a solution of
specific, accurate concentration. The apparatus used to prepare such a solution is a volumetric flask. A meniscus finder
is useful in setting the bottom of the meniscus on the calibration line (Figure 22).
GO
1. (Prelab) Calculate the volume of concentrated
reagent required.
2. Add approximately one-half of the final volume of
pure water to the volumetric flask.
3. Measure the required volume of stock solution using
a pipette. (Refer to “Using a Pipette” on page 800).
Preparing a Standard Solution from a
Solid Reagent
4. Transfer the stock solution slowly into the volumetric
There is a video demonstration of this technique on the
Nelson Web site.
5. Add pure water, and then use a medicine dropper and
www.science.nelson.com
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GO
flask while mixing.
a meniscus finder to set the bottom of the meniscus
on the calibration line (Figure 22).
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6. Stopper and mix the solution by slowly inverting the
flask several times.
If water is added directly to some solids or
concentrated liquids, there may be boiling or
splattering. Always add a solid solute or
concentrated liquids to water.
Titration
0.1 mL. Avoid parallax errors by reading volumes at
eye level with the aid of a meniscus finder.
4. Pipette a known volume of the solution of unknown
concentration into a clean Erlenmeyer flask. Place a
white piece of paper beneath the Erlenmeyer flask to
make it easier to detect colour changes.
5. Add an indicator if one is required. Add the smallest
Titration is used in the volumetric analysis of an unknown
concentration of a solution. Titration involves adding a solution (the titrant) from a burette to another solution (the
sample) in an Erlenmeyer flask until a recognizable endpoint, such as a colour change, occurs. (See the video on the
Nelson Web site.)
www.science.nelson.com
3. Record the initial burette reading to the nearest
GO
1. Rinse the burette with small volumes of pure water
using a wash bottle. Using a burette funnel, rinse with
small volumes of the titrant (Figure 23). (If liquid
droplets remain on the sides of the burette after
rinsing, scrub the burette with a burette brush. If the
tip of the burette is chipped or broken, replace the tip
or the whole burette.)
quantity necessary (usually 1 to 2 drops) to produce a
noticeable colour change in your sample.
6. Add the solution from the burette quickly at first, and
then slowly, drop-by-drop, near the endpoint
(Figure 24). Stop as soon as a drop of the titrant
produces a permanent colour change in the sample
solution. A permanent colour change is considered to
be a noticeable change that lasts for 10 s after
swirling.
7. Record the final burette reading to the nearest
0.1 mL.
8. The final burette reading for one trial becomes the
initial burette reading for the next trial. Three trials
with results within 0.2 mL are normally required for a
reliable analysis of an unknown solution.
9. Drain and rinse the burette with pure water. Store the
burette upside down with the stopcock open.
Figure 23
A burette should be rinsed with water
and then the titrant before use.
Figure 24
Near the endpoint, continuous gentle
swirling of the solution is particularly
important.
2. Using a small burette funnel, pour the titrant solution
into the burette until the level is near the top. Open
the stopcock for maximum flow to clear any air
bubbles from the tip and to bring the liquid level
down to the scale.
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Appendix C
Diagnostic Tests
The tests described in Table 1 are commonly used to detect the
presence of a specific substance. All diagnostic tests include
a brief procedure, some expected evidence, and an interpretation of the evidence obtained. This is conveniently communicated using the format, “If [procedure] and [evidence],
then [analysis].” Diagnostic tests can be constructed using
any characteristic empirical property of a substance. For
example, diagnostic tests for acids, bases, and neutral substances can be specified in terms of the pH of the solutions.
For specific chemical reactions, properties of the products
that the reactants do not have, such as the insolubility of a
precipitate, the production of a gas, or the colour of ions in
aqueous solutions, can be used to construct diagnostic tests.
If possible, you should use a control to illustrate that the test
does not give the same results with other substances. For
example, in the test for oxygen, inserting a glowing splint into
a test tube that contains only air is used to compare the effect
of air on the splint with a test tube in which you expect oxygen
has been collected.
Communication of Diagnostic Tests
The procedure, evidence, and analysis information for a diagnostic test can be communicated in three different formats:
• “If ... and ... then ...” statement
• table
• flowchart
Table 1 Some Standard Diagnostic Tests
Substance tested
Diagnostic test
water
If cobalt(II) chloride paper is exposed to a liquid or vapour, and the paper turns from blue to pink, then water
is likely present.
oxygen
If a glowing splint is inserted into the test tube, and the splint glows brighter or re-lights, then oxygen gas is
likely present.
hydrogen
If a flame is inserted into the test tube, and a squeal or pop is heard, then hydrogen is likely present.
carbon dioxide
If the unknown gas is bubbled into a limewater solution, and the limewater turns cloudy, then carbon dioxide
is likely present.
halogens
If a few millilitres of a hydrocarbon solvent is added, with shaking, to a solution in a test tube,
and the colour of the solvent appears to be
• light yellow-green, then chlorine is likely present
• orange, then bromine is likely present
• purple, then iodine is likely present
acid
If strips of blue and red litmus paper are dipped into the solution, and the blue litmus turns red, then an acid
is present.
base
If strips of blue and red litmus paper are dipped into the solution, and the red litmus turns blue, then a base
is present.
neutral solution
If strips of blue and red litmus paper are dipped into the solution, and neither litmus changes colour, then
only neutral substances are likely present.
neutral ionic
solution
If a neutral solution is tested for conductivity with a multimeter, and the solution conducts a current, then a
neutral ionic substance is likely present.
neutral molecular
solution
If a neutral solution is tested for conductivity with a multimeter, and the solution does not conduct a current,
then a neutral molecular substance is likely present.
There are thousands of diagnostic tests. You can create some of these using data from the periodic table
(on the inside front cover of this book), and from the data tables in Appendix I and on the inside back cover.
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Appendix D
Page 806
STS PROBLEM SOLVING
Science is a human endeavour, technology has a social purpose, and both have always been part of society. Science and
technology together affect society in a myriad of ways. Society
also affects science and technology by placing controls on
them and expecting solutions to societal problems.
When controversial issues related to science and technology
arise in our society, there is often heated debate among var-
ious special-interest groups. Often, little progress is made
because different parties in the debate generally recognize
only a single perspective on the issue. Many people now realize
that an informed multi-perspective view is more defensible.
The following model represents one possible procedure for
making an informed decision on a social issue related to science and technology.
D.1 STS Decision-Making Model
1. Identify an STS (science–technology–society) issue.
Newspapers, magazines, and news broadcasts are
sources of current STS issues. However, some issues
like acid rain have been current for some time and
rarely appear in the news. When identifying an issue
for debate, it is convenient to state the issue as a
resolution (e.g., “Be it resolved that the use of fossil
fuels for heating homes should be eliminated.” ).
2. Design a plan to address the STS issue. Possible
designs include individual research, a debate, a townhall meeting (or role-playing), or participation in an
actual hearing or on a committee.
3. Identify and obtain relevant information on as many
perspectives as possible. An STS issue will always have
scientific and technological perspectives. Common
perspectives are shown in Table 2.
Table 2 Perspectives on STS Issues
scientific
ethical
technological
social
ecological
militaristic
economic
esthetic
political
mystical
legal
emotional
D.2 Types of Reports
There are many ways to communicate the results of an investigation of an STS issue (Table 3). All methods will require
some research about the issue and perspectives on the issue
(including positive and negative viewpoints). Some methods
can also include alternative solutions and the evaluation of
these solutions. Working within a group and brainstorming
is a useful process. No matter how the issue will be presented
and reported, you need to be well prepared.
806
Appendix D
Another perspective is the world view or perspective
of Aboriginal peoples. In general, Aboriginal peoples
believe that we are an integral part of our
environment. Their holistic view includes not only a
physical interdependence but also a spiritual one.
Information from different perspectives can be
obtained from references and through group
discussions. There are many sides to every issue. There
can be positive and negative viewpoints about the
resolution from every perspective.
4. Generate a number of alternative solutions to the STS
problem. Some obvious solutions will arise from the
resolution. Other creative solutions often arise from a
brainstorming session within a group.
5. Evaluate each solution and decide which is best. One
method is to rank the value of a particular solution
from each perspective. For example, a solution might
have little economic advantage and be ranked as 1 on
a scale of 1 to 5; the solution might have a significant
ecological benefit and be ranked as 5, for a total of 6.
A different solution might be judged as 3 from the
economic perspective and 1 from the ecological
perspective, for a total of 4. The solution with the
highest total is likely to be chosen. Although
simplistic, this method facilitates evaluation and
illustrates the tradeoffs that occur in any real issue.
Table 3 STS Investigations and Reports
Plan
Reporting suggestions
individual or group
research
• written report or poster
• multimedia presentation
debate
• research notes
• videotape of the debate
role-playing (e.g.,
town hall meeting)
• research notes
• videotape of the meeting
survey
• survey form with tables and graphs
newspaper article
• published article
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SAFETY KNOWLEDGE AND SKILLS
Appendix E
E.1 Laboratory Safety
Safety is always important in a laboratory or in other settings
that feature chemicals or technological devices. It is your
responsibility to be aware of possible hazards, to know the
rules—including ones specific to your classroom—and to
behave appropriately. Always alert the teacher in case of any
accident.
Alberta Education has an extensive document,“Safety in the
Science Classroom,” that deals with hazards and safety.
www.science.nelson.com
Safety in the laboratory is an attitude and a habit more than
it is a set of rules. It is easier to prevent accidents than to deal
with the consequences of an accident. Most of the following
rules are common sense:
•
•
•
•
•
•
•
•
•
•
•
•
•
NEL
•
•
•
GO
General Safety Rules
•
•
Read all directions before doing any laboratory work,
and follow all verbal instructions.
Know the potential hazards, including the contents
and location of MSDS, and the location of all safety
equipment.
Wear eye protection and lab aprons/coats.
Behave responsibly. Avoid sudden or rapid motion
that may interfere with someone carrying or working
with chemicals.
Wear closed shoes (not sandals or bare feet) when
working in the laboratory.
Place your books, bags, and purses away from the
work area.
Do not chew gum, eat, drink, or taste anything in the
laboratory.
Ask for assistance when you are not sure how to do a
procedural step.
Inform your teacher immediately if any problem or
accident occurs.
Never attempt any unauthorized or unsupervised
experiments.
Never handle any chemical with your hands. Use a
laboratory scoop or spoon for solids.
Never use the contents of a bottle that has no label or
an illegible label. Always double check the label to
ensure that you are using the chemical needed. Always
pour from the side opposite the label.
When leaving chemicals in containers, ensure that they
are labelled.
Do not take any more chemical than needed and never
return excess chemicals to their original container.
•
•
•
•
Hold larger bottles with both hands; one hand on the
base.
Do not inhale any vapours directly from any container.
If smell is to be tested, fan the vapours toward your
nose, keeping the container away from underneath
your nose.
Always use a pipette bulb, and never pipette by mouth.
When heating a test tube over a burner, use a test-tube
holder with the test tube at an angle, facing away from
you and others. Gently move the test tube backwards
and forwards through the flame.
Clean up all spills, even spills of water, immediately.
Clean up your work area at the end of an experiment.
Dispose of chemicals appropriately as directed by your
teacher.
Always wash your hands with soap and water before
you leave the laboratory.
Do not forget safety procedures when you leave the
laboratory. These same rules also apply at home or at
work.
Glass Safety and Cuts
• Never use glassware that is cracked or chipped. Give
•
•
•
such glassware to your teacher or dispose of it as
directed. Do not put the item back into circulation.
Never pick up broken glassware with your fingers. Use
a broom and dustpan.
Do not put broken glassware into garbage containers.
Dispose of glass fragments in special containers
marked “broken glass.”
If you cut yourself, inform your teacher immediately.
Imbedded glass or continued bleeding requires
medical attention.
Burns
• In a laboratory where burners or hot plates are being
•
used, never pick up a glass object without first
checking the temperature by lightly and quickly
touching the item. Glass items that have been heated
stay hot for a long time but do not appear to be hot.
Metal items such as ring stands and hot plates can also
cause burns; take care when touching them.
Before using a laboratory burner, make sure that long
hair is always tied back. Do not wear loose clothing.
(Wide long sleeves should be tied back or rolled up.)
Safety Knowledge and Skills
807
E
Appendix A-F_Chem20
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•
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Page 808
Do not use a laboratory burner near wooden shelves,
flammable liquids, or any other item that is
combustible. Know how to use the type of burner in
your laboratory. (See Using a Laboratory Burner,
page 798)
Never look down the barrel of a laboratory burner.
Always pick up a burner by the base, never by the
barrel.
Never leave a lighted bunsen burner unattended.
If you burn yourself, immediately run cold water over
the burned area and inform your teacher.
Eye Safety
• Always wear approved eye protection in a laboratory,
•
•
•
•
no matter how simple or safe the task appears to be.
Keep the safety glasses over your eyes, not on top of
your head. For certain experiments, full face protection may be necessary.
Never look directly into the opening of flasks or test
tubes.
If, in spite of all precautions, you get a solution in your
eye, quickly use the eyewash station or nearest running
water. Continue to rinse the eye with water for at least
15 minutes. This is a very long time—have someone
time you. Unless you have a plumbed eyewash system,
you will also need assistance in refilling the eyewash
container. Have another student inform your teacher of
the accident. The injured eye should be examined by a
doctor.
It is recommended that you do not wear contact lenses
in the laboratory. If you wear contact lenses in the
laboratory, there is a danger that a chemical might get
behind the lens where it cannot be rinsed out with
water. If you must wear contact lenses in the chemistry
laboratory, be extra careful. Tell your teacher if you are
wearing contact lenses in the laboratory. Whether or
not you wear contact lenses, do not touch your eyes
without first washing your hands.
If a piece of glass or other foreign object enters an eye,
immediate medical attention is required.
Fire Safety
Immediately inform your teacher of any fires. Very small fires
in a container may be extinguished by covering the container
with a wet paper towel or a ceramic square, which would cut
off the supply of air. If anyone’s clothes or hair catch fire, the
fire can be extinguished by smothering the flames with a
blanket or a piece of clothing. Larger fires require a fire extin-
808
Appendix E
guisher. (Know how to use the fire extinguisher that is in your
laboratory.) If the fire is too large to approach safely with an
extinguisher, vacate the location and sound the fire alarm.
(School staff will inform the fire department.)
If you use a fire extinguisher, direct the extinguisher at the
base of the fire and use a sweeping motion, moving the extinguisher nozzle back and forth across the front of the fire’s
base. You must use the correct extinguisher for the kind of
fire you are trying to control. Each extinguisher is marked
with the class of fire for which it is effective. The fire classes
are outlined below. Most fire extinguishers in schools are of
the ABC type.
•
•
•
•
•
Class A fires involve ordinary combustible materials
that leave coals or ashes, such as wood, paper, or cloth.
Use water or dry chemical extinguishers on Class A
fires. (Carbon dioxide extinguishers are not satisfactory as carbon dioxide dissipates quickly and the hot
coals can reignite.)
Class B fires involve flammable liquids such as gasoline
or solvents. Carbon dioxide or dry chemical
extinguishers are effective on Class B fires. (Water is
not effective on a Class B fire since the water splashes
the burning liquid and spreads the fire.)
Class C fires involve live electrical equipment, such as
appliances, photocopiers, computers, or laboratory
electrical apparatus. Carbon dioxide or dry chemical
extinguishers are recommended for Class C fires.
Carbon dioxide extinguishers are much cleaner than
the dry chemical variety. (Using water on live electrical
devices can result in severe electrical shock.)
Class D fires involve burning metals, such as sodium,
potassium, magnesium, or aluminium. Sand or salt are
usually used to put out Class D fires. (Using water on a
metal fire can cause a violent reaction.)
Class E fires involve a radioactive substance. These
involve special considerations at each site.
Electrical Safety
Water or wet hands should never be used near electrical equipment. When unplugging equipment, remove the plug gently
from the socket (do not pull on the cord). Do not use any
devices with electric motors when flammable liquids are
present unless the area is well ventilated.
NEL
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Appendix E
E.2 Safety Symbols and Information
Educational, Commercial, and
Industrial Information
Class B:
Flammable and
combustible
material
Class A:
Compressed gas
Class C:
Oxidizing
material
Class D: Poisonous and Infectious Materials
Division 1
Division 2
Division 3
Materials causing
immediate and
serious toxic effect
Materials
causing other
toxic effects
Biohazardous
infectious
material
The Workplace Hazardous Materials Information System
(WHMIS) provides workers and students with complete and
accurate information regarding hazardous products. All chemical products supplied to schools, businesses, and industry
must contain standardized labels and be accompanied by
Material Safety Data Sheets (MSDS) providing detailed information about the product. Clear and standardized labelling
is an important component of WHMIS (Figure 25). These
labels must be present on the product’s original container or
be added to other containers if the product is transferred.
Although MSDS must be supplied with every product sold,
current MSDS can also be obtained at several Internet sites,
which are useful for researching information about chemicals.
www.science.nelson.com
Figure 25
WHMIS symbols
Class F:
Dangerously reactive
material
Class E:
Corrosive
material
Poison
Danger
E
Flammable
Explosive
Warning
Corrosive
Caution
GO
Consumer Information
The Canadian Hazardous Products Act requires manufacturers of consumer products containing chemicals to include
a symbol specifying both the nature and degree of the primary hazard, and to note any secondary hazards, first aid
treatment, storage, and disposal. The symbols show the hazard
by an illustration and the degree of the hazard by the type of
border surrounding the illustration (Figure 26).
Figure 26
Household Hazardous Product Symbols
NEL
Safety Knowledge and Skills
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E.3 Waste Disposal
Disposal of chemical wastes at home, at school, or at work is
a societal issue. We all need to be stewards of our planet; in
other words, to behave as custodians or keepers. Some governments, institutions, and industries have begun to implement product stewardship programs. This is an environmental
management plan based on the principle that whoever designs,
produces, sells, or uses a product should take responsibility for
minimizing the product’s environmental impact over its complete life cycle. Governments have regulations for the handling, transportation, and disposal of chemicals, but each of
us needs to take responsibility for the wastes we produce at
home and at school.
Most laboratory waste can be washed down the drain, or,
if it is in solid form, placed in ordinary garbage containers.
However, some waste must be treated more carefully.
Throughout this textbook, special waste disposal problems
are noted, but it is your responsibility to dispose of waste in
the safest possible manner.
Heavy Metal Solutions
Heavy metal compounds (for example, lead, mercury, or cadmium compounds) should not be flushed down the drain.
These substances are cumulative poisons and should be kept
out of the environment. A special container is kept in the laboratory for heavy metal solutions. Pour any heavy metal waste
into this container. Remember that paper towels used to wipe
up solutions of heavy metals, as well as filter papers with
heavy metal compounds imbedded in them, should be treated
as solid toxic waste.
Disposal of heavy metal solutions is usually accomplished
by precipitating the metal ion (for example, as lead(II) silicate)
and disposing of the solid. Disposal may be by elaborate
means such as deep well burial, or by simpler but accepted
means such as delivering the substance to a landfill. Heavy
metal compounds should not be placed in school garbage
containers. Usually, waste disposal companies collect materials
that require special disposal and dispose of them as required
by law.
Flammable Substances
Flammable liquids should not generally be washed down the
drain. (The exceptions to this rule are aqueous solutions of
non-toxic flammables such as alcohol–water solutions: they
can safely be flushed.) Special fire-resistant containers are
used to store flammable liquid waste. Waste solids that pose
a fire hazard should be stored in fireproof containers. Care
must be taken not to allow flammable waste to come into
contact with any sparks, flames, other ignition sources, or
oxidizing materials. The particular method of disposal depends
on the nature of the substance.
Corrosive Solutions
Solutions that are corrosive but not toxic, such as acids, bases,
or oxidizing agents, can usually be washed down the drain, but
care should be taken to ensure that they are properly neutralized and diluted.
To neutralize diluted waste acids, use diluted waste bases,
and vice versa. Or, use sodium bicarbonate for neutralizing the
acid and use dilute hydrochloric acid for neutralizing the
base. Oxidizing agents, such as potassium permanganate,
should also be diluted with a 10% aqueous solution of sodium
thiosulfate (reducing agent) before washing them into the
drain.
Use large quantities of water and continue to pour water
down the drain for a few minutes after all the substance has
been washed away.
810
Appendix E
Toxic Substances
Solutions of toxic substances, such as oxalic acid, should not
be poured down the drain, but should be disposed of in the
same manner as heavy metal solutions. Solid toxic substances
are handled similarly to precipitates of heavy metal. Chemicals
should be stored in their original containers, with their labels
clearly visible.
Appendix A-F_Chem20
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Appendix F
Page 811
COMMUNICATION SKILLS
Communication is essential in science. The international
scope of science requires that quantities, chemical symbols, and
mathematical tools such as numbers, operations, tables, and
graphs, be understood by scientists in different countries with
different languages. The way in which scientific knowledge
is expressed also reflects the nature of scientific knowledge, and
in particular, the certainty of the knowledge.
F.1 Scientific Language
Science deals with two types of knowledge: empirical (observable “facts”) and theoretical (non-observable ideas). Directly
observable knowledge is generally considered to be more certain than interpretations or theoretical concepts. Theories
are subject to change and, therefore, are less certain than the
observations upon which they are based.
When observations are interpreted or explained, the language used should reflect some uncertainty or tentativeness.
Use phrases such as
•
•
•
•
•
The evidence suggests that…
According to the theory of…
It appears likely that…
Scientists generally believe that…
One could hypothesize that…
Avoid the use of the word “prove.” Scientific concepts cannot
be proven. The evidence may be extensive and reliable, but a
concept to explain the evidence will never be 100% certain.
In general, the language that you use should reflect the
certainty of the information (observations are more certain
than scientific concepts), and it should refer to the evidence
available to you.
DID YOU KNOW
?
Confidence in Empirical versus Theoretical
Knowledge
A candle does not burn unless air is present. In a closed container, a candle flame is extinguished after a short period of
time. These are simple and relatively certain facts that can be
directly stated. At one time, scientists believed that burning
releases a substance called phlogiston, which was absorbed
by the air until it could hold no more phlogiston; this is what
stopped the burning. This theory, which was firmly believed by
many chemists until the 1800s, was eventually replaced by the
oxygen theory of combustion. The facts (evidence) remained
the same but the idea (theory) completely changed.
F.2 SI Symbols and Conventions
The International System of Units, known as SI from the
French name, Système international d’unités, is the measurement and communication system used internationally by scientists; it is also the legal measurement system in Canada and
most countries in the world. Physical quantities are ultimately
expressed in terms of seven fundamental SI units, called base
units, which cannot be expressed as combinations of simpler
units (Table 4). Although the base unit for mass is the kilogram (kg), it is more common in a chemistry laboratory to use
the gram (g). Similarly, although the base unit for temperature (T) is kelvin (K), the common temperature (t) unit is
degree Celsius (°C).
All other quantities can be expressed in terms of these seven
fundamental quantities. For convenience, a unit derived from
a combination of base units may be assigned a symbol of its
own. Table 5 lists a few of the physical quantities and derived
units most commonly encountered in chemistry.
NEL
Table 4 Quantities and Fundamental Base Units
Quantity
Symbol
Unit
Symbol
length
l
metre
m
time
t
second
s
mass
m
kilogram
kg
chemical amount
n
mole
mol
temperature
T
kelvin
K
electric current
I
ampere
A
luminous intensity
Iv
candela
cd
Quantities and their SI base units are listed in Table 5, on
the next page. These are the units most widely used by scientists. For convenience, however, units such as tonne (T)
for mass and annum (a) for year are sometimes used to represent quantities that would be inconveniently large when
expressed as base units.
Communication Skills
811
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Table 5 Quantities and Base Units
Quantity
Symbol
Unit
Symbol
molar mass
M
grams per mole
g/mol
volume
V
litre
L
amount concentration
c
moles per litre
mol/L
pressure
P
pascal
Pa
energy
E
joule
J
heat capacity
C
joules per degree Celsius
J/°C
specific heat capacity
c
joules per gram per degree Celsius
J/(g•°C)
volumetric heat capacity
c
megajoules per cubic metre
per degree Celsius
MJ/(m3•°C)
molar enthalpy
rHm
kilojoules per mole
kJ/mol
enthalpy change
rH
kilojoules
kJ
electric charge
Q
coulomb
C
electric potential difference (voltage)
E
volt (joules per coulomb)
V
SI Prefixes
Next to universality, the most important feature of any system
of units is convenience. SI has been designed to maximize
convenience in a number of ways. A given quantity is always
measured in the same base unit regardless of the context in
which it is measured. For example, all forms of energy,
including energy in food, are measured in joules. When a unit
is too large or too small for convenient measurement, the
unit is adjusted in size with a prefix. (See Table 6.) Prefixes
allow units to be changed in size by multiples of ten. However,
except for the use of “centi” in centimetre, we commonly use
only prefixes that change the unit in multiples of a thousand.
Table 6 Some SI Prefixes
Prefix
Symbol
Fact
tera
T
1012
giga
G
109
mega
M
106
kilo
k
103
milli
m
103
micro
m
106
nano
n
109
pico
p
1012
Scientific Notation
the following numbers are expressed in regular notation and
scientific notation:
Regular notation
Scientific notation
1200 L
0.000 000 998 mol/L
1.200 103 L
–
9.98 10 7 mol/L
On some calculators, the F e E key or the FSE key changes
the number in the display into or from scientific notation. To
enter a value in scientific notation in your calculator, the
EXP or EE key is used to enter the power of ten. Note that the
base 10 is not keyed into the calculator. For example, to enter
1.200 × 10 3 press
1
•
2
EXP
3
9.98 × 10
9
•
9
8
EXP
–7
press
7
+/–
All mathematical operations and functions (such as , –, ×,
÷, log) can be carried out with numbers in scientific notation.
Scientific notation is useful in calculations because it simplifies the cancellation of units and the totalling of powers of
ten. However, scientific notation is sometimes overused. SI
recommends that, wherever possible, prefixes be used to report
measured values. Scientific notation should be reserved for
situations where no prefix exists, or where it is essential to
use the same unit (for example, comparing a wide range of
energy values in kilojoules per gram). A reported value should
use a prefix or scientific notation, but not both, unless you
are comparing values. Scientific notation should usually use
the base unit.
Scientific notation is a convenient method for expressing
either a very large value or a very small value as a number
between 1 and 10 multiplied by a power of 10. For example,
812
Appendix F
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Appendix F
F.3 Quantitative Descriptions and Rules
Quantities that have exact values are either defined quantities
(for example, 1 t is defined as exactly 1000 kg, and the SI
prefix kilo, k, is exactly 1000), or quantities obtained by
counting (for example, 32 people in a class or any coefficient
in a balanced chemical equation). You can be certain about
such quantities; there will be a small degree of uncertainty
when counting very large numbers.
On the other hand, most quantities are measured by a
person using some measuring instrument (for example, measuring the mass of a chemical using a balance). Since every
instrument has its limitations and no one can perfectly
measure a quantity, there is always some uncertainty about the
number obtained. This uncertainty depends on the size of
the sample measured, the particular instrument used, and
the technological skill of the person doing the measurement.
Accuracy
Accuracy is an expression of how close an experimental value
is to the accepted value. The comparison of the two values is
often expressed as a percent difference. For example, the accuracy of a prediction based on some authority can be expressed
as the absolute value of the difference divided by a predicted
value and converted to a percent.
% difference experimental value predicted value 100
predicted value
This expression of accuracy is often used in the Evaluation
section of investigation reports.
(a)
(b)
Precision
Accuracy is an expression of how close a value is to the
accepted, expected, or predicted value, whereas precision is a
measure of the reproducibility or consistency of a result
(Figure 27). Accuracy is generally attributed to an error in
the system (a systematic error); precision is associated with a
random error of measurement. For example, if you used a
balance without zeroing it, you might obtain measurements
that have high precision (reproducibility) but low accuracy.
The systematic error might be high (low accuracy), but the
random error of the measurement is low (high precision).
Scientists define precision as the closeness of the agreement between independent measurements. We make the
assumption that, if an instrument produces a certain decimal fraction (like a tenth of a unit), then all repeated measurements would be the same except for that last digit.
As long as an instrument is read correctly, for simplicity, we
will assume that precision is the place value of the last measurable digit and is determined by the instrument. A mass of
17.13 g is more precise than 17.1 g. The precision is determined by the particular system or instrument used; for
example, a centigram balance versus a decigram balance.
You may not know how uncertain the last measured digit
is. On a centigram balance, the error of measurement in the
last digit is usually 0.01 g. Measurements such as 12.39 g,
12.40 g, and 12.41 g all have the same precision (hundredths),
and may all be equally correct masses for the same object.
The precision with which you read a thermometer might be
0.2 °C (for example, 21.0 °C, 21.2 °C or 21.4 °C) and a ruler
might be read to 0.5 mm; you must decide, for example,
whether to record 11.0 mm, 11.5 mm, or 12.0 mm.
(c)
Figure 27
The positions of the darts in each of these figures are analogous to measured or calculated results in a laboratory setting.
The results in (a) are precise and accurate, in (b) they are precise but not accurate, and in (c) they are neither precise nor accurate.
NEL
Communication Skills
813
F
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Page 814
Precision Rule for Calculations
A result obtained by adding or subtracting measured values is
rounded to the same precision (number of decimal places) as
the least precise value used in the calculation. For example,
12.6 g 2.07 g 0.142 g totals to 14.812 g on your calculator.
This value is rounded to one-tenth of a gram and reported as
14.8 g because the first measurement limits the precision of
the final result to tenths of a gram and the rounding rule suggests leaving the 8 as is. The final result is reported to the least
number of decimal places in the values added or subtracted.
the decimal point. For example, 6.20 mL (3 significant digits)
has the same number of significant digits as 0.00620 L.
For each of the following measurements, the certainty
(number of significant digits) is stated beside the measured
or calculated value:
0.41 mL
a certainty of 2 significant digits
700 mol
a certainty of 3 significant digits
0.020 50 km
a certainty of 4 significant digits
a certainty of 1 significant digit
2 1040 m
Certainty Rule for Calculations
Precision Rule for pH
The precision rule for pH and pOH is a special case, although
the logic is consistent with that for values expressed in
scientific notation. A hydrogen or hydronium ion concentration of 1.0 107 mol/L converts to a pH of 7.00. Just as
the 7 is not counted as a significant digit when communicating the scientific notation value, the 7 is also not counted
when communicating the pH value. Therefore, the rule is
• The number of digits following the decimal point in a pH
or pOH value is equal to the number of significant digits
in the corresponding hydronium or hydroxide ion concentration.
and
• The number of significant digits in a hydronium or hydroxide
ion concentration is equal to the number of digits following
the decimal place in the corresponding pH or pOH.
Certainty
How certain you are about a measurement depends on two
factors—the precision of the instrument and the value of the
measured quantity. More precise instruments give more certain values; for example, 15.215 °C as opposed to 15 °C.
Consider two measurements with the same precision, 0.4 g and
12.8 g. If the balance used is precise to 0.2 g, the value
0.4 g could vary by as much as 50%. However, 0.2 g of 12.8 g
is a variation of less than 2%. For both factors—precision of
instrument and value of the measured quantity—the more
digits in a measurement, the more certain you are about the
measurement. We communicate the certainty of any measurement by the number of significant digits. In a measured or calculated value, significant digits are all those digits that are
certain plus one estimated (uncertain) digit. Significant digits
include all digits correctly reported from a measurement,
except leading zeros. Leading zeros are the zeros at the
beginning of a decimal fraction and are written only to locate
814
Appendix F
Significant digits are primarily used to determine the certainty of a result obtained from calculations using several
measured values. A result obtained by multiplying or dividing
measured values is rounded to the same certainty (number
of significant digits) as the least certain value used in the
calculation. For example, 0.024 89 mol 6.94 g/mol is displayed as 0.1727366 g on a calculator. This is correctly reported
as 0.173 g or 173 mg because the second value used (6.94)
limits the final result to a certainty of three significant digits.
Research in the News
News media often quote the results of surveys, such as the
percentage of people who would vote for a certain political
party. What does it mean when we hear, “The results were
Yes 52%, No 42%, Undecided 5%, with a margin of error of
3% 19 times out of 20”? The margin of error (3%) is usually
calculated as the reciprocal of the square root of the sample
size. A larger sample size would therefore produce a smaller
margin of error. The confidence level (19 times out of 20) is like
the precision. If the survey were repeated 20 times, the result
would be within the percent error 19 times and 1 time it
would be very different. The pollster, in effect, claims to be
(accurate) within 3% of the “real” answer 19 out of 20 times.
Rounding
When completing calculations that involve more than one step,
there are two rules that are used for answers in this textbook:
• Never round off partial answers in your calculator.
• Always round off when communicating partial
answers on paper.
When chained calculations involve both multiplication/division and addition/subtraction, you may be required to store
the partial answers in your calculator memory or to use the
bracket function on your calculator.
NEL
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Appendix F
Rounding Rule
Calculations are usually based on measurements (for example,
in the Analysis section of a report). To report a calculated
result correctly, follow this procedure. Check the first digit
following the digit that will be rounded. If this digit is less
than 5, it and all following digits are discarded. If this digit is
5 or greater, it and all following digits are discarded, and the
preceding digit is increased by one.
F.4 Tables and Graphs
Tables
The Reaction of
HCl(aq) with Zn(s)
F
300
Time (s)
Both tables and graphs are used to summarize information
and to illustrate patterns or relationships. Preparing tables
and graphs requires some knowledge of accepted practice
and some skill in designing the table or graph to best describe
the information.
200
100
1. Write a descriptive title (Table 7).
2. The row or column with the manipulated variable
usually precedes the row or column with the
responding variable.
3. Label all rows and/or columns with a heading,
including units in parentheses where necessary. Units
are not usually written in the main body of the table.
Table 7 The Reaction of HCl(aq) with Zn(s)
Concentration of HCl(aq)
(mol/L)
Time for reaction
(s)
2.0
70
1.5
80
1.0
144
0.5
258
Graphs
1. Write a descriptive title on the graph and label the
axes (Figure 28).
• Label the horizontal (x) axis with the name of the
manipulated variable and the vertical (y) axis with
the name of the responding variable.
• Include the unit in parentheses on each axis label,
for example, “Time (s).”
2. Assign numbers to the scale on each axis.
• As a general rule, the data points should be spread
out so that at least one-half of the graph paper is
used.
NEL
0
0.5
1.0
1.5
2.0
Concentration of HCl(aq) (mol/L)
Figure 28
A sample graph
• Choose a scale that is easy to read and has equal
divisions. Each division (or square) must represent
a small simple number of units of the variable; for
example, 0.1, 0.2, 0.5, or 1.0.
• It is not necessary to have the same scale on each
axis or to start a scale at zero.
• Do not label every division line on the axis. Scales
on graphs are labelled in a way similar to the way
scales on rulers are labelled.
3. Plot the data points.
• Locate each data point by making a small dot in
pencil. When all points are drawn and checked,
draw an X over each point, or circle each point in
ink.
• Be suspicious of a data point that is obviously not
part of the pattern. Double-check the location of
such points, but do not eliminate the point from
the graph if it does not align with the rest.
4. Draw the best-fitting line.
• Using a sharp pencil, draw a line that best
represents the trend shown by the collection of
points. Do not force the line to go through each
point. Uncertainty of experimental measurements
may cause some of the points to be misaligned.
Communication Skills
815
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• If the collection of points appears to fall in a
straight line, use a ruler to draw the line. Otherwise,
draw a smooth curve that best represents the
pattern of the points.
•
•
• Since the data points are in ink and the line is in
pencil, it is easy to change the position of the line if
your first curve does not fit the points to your
satisfaction.
•
Using Your Graph
Interpolation is used to find values between measured
points on the graph.
Extrapolation is used to find values beyond the
measured points on a graph. A dotted line on a graph
indicates an extrapolation.
The scattering of points gives a visual indication of the
uncertainty in the experiment. A point that is
obviously not part of the pattern may require a
remeasurement to check for an error or may indicate
the influence of an unexpected variable.
Although a graph is constructed using a limited number of
measured values, the pattern may be used to extend the empirical information.
F.5 Problem-Solving Methods
Definition of Terms
2 1.5 mol
• ratio: comparison of two numbers; e.g., , 3
1L
• proportion: an equality of two ratios; e.g.
nNH3
1.5 mol
2. 0 L
1L
• conversion factor: a specific type of ratio that is
used to convert a quantity from one unit to another
unit; e.g.,
1000 m
1 km 40.00 g
1 mol
or , or 1 km
1000 m 1 mol
40.00 g
• formula: a mathematical statement of a
relationship using SI and IUPAC symbols and
format; e.g., m nM (Note: m n M is not
acceptable.)
Basic Problem-Solving Methods
A stoichiometry example is used to compare the three basic
methods.
Copper metal is used to recover silver from a silver
nitrate solution. Predict the mass of silver obtained from
the complete reaction of 50.0 g of copper.
Cu(s) 2 AgNO3(aq) → 2 Ag(s) Cu(NO3)2(aq)
50.0 g
63.55 g/mol
m
107.87 g/mol
nAg
2
0.787 mol
1
nAg 1.57 mol
mAg
107.87 g
1.57 mol
1 mol
mAg 170 g
Formula Method
m
50.0 g
nCu 0.787 mol
M
63. 55 g/mol
nrequired
2
nAg 0.787 mol 0.787 mol ngiven
1
1.57 mol
mAg nM 1.57 mol 107.87 g/mol 170 g
Conversion Factor Methods
(a) Step method
1 mol
nCu 50.0 g 0.787 mol
63. 55 g
2
nAg 0.787 mol 1.57 mol
1
107.87
g
mAg 1.57 mol 170 g
mol
(b) Full method (“Factor Label”)
Proportion Method
nCu
1 mol
50. 0 g
63. 55 g
816
Appendix F
nCu 0.787 mol
1 mol Cu
2 mol Ag
mAg 50.0 g Cu 63.55 g Cu
1 mol Cu
107.87 g Ag
170 g
1 mol Ag
NEL
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Appendix G
Page 817
REVIEW OF CHEMISTRY 20
Your review of Chemistry 20 for Chemistry 30 will be more successful if you study the highlighted Summaries, Sample Problems,
and Communication Examples in each chapter. By answering the following questions you will find out where you need to check your
understanding before starting Chemistry 30.
Unit 1 Chemical Bonding
(Chapter 3)
1. Distinguish between the two important types of
scientific knowledge.
2. Identify the characteristics of acceptable scientific
theories.
3. Explain the octet rule and how it relates to chemical
reactivity.
4. Copy and complete the following table.
Table 1 Theoretical Descriptions of Selected Elements
Element
name
Lewis Group Number
Number Number
symbol number of valence of lone of bonding
electrons pairs
electrons
calcium
aluminium
arsenic
(b) Ionic compounds are electrical conductors in
molten and aqueous states.
8. Using Lewis symbols and formulas, write the
formation equation for each of the following
compounds.
(a) potassium bromide
(b) sodium oxide
(c) calcium fluoride
9. Why are chemical formulas for ionic compounds
always based the simplest whole number ratio of
ions? Is the simplest whole number ratio also used for
molecular formulas? Why or why not?
10. Compare ionic and covalent bonds, including how
they are formed, according to theory and the nature
of the bond.
11. For each of the following molecular formulas, draw
the Lewis, structural, and stereochemical formulas,
and state the shape around the central atom.
oxygen
(a) OCl2
(d) HCN
bromine
(b) SiH4
(e) CH2O
neon
(c) NCl3
5. (a) State the types of elements expected to react to
form compounds containing covalent bonds.
(b) State the types of elements expected to react to
form compounds containing ionic bonds.
(c) Explain your answers to (a) and (b) using the
concept of electronegativity.
(d) Why is it difficult to predict the type of bonding
in some compounds using only
electronegativities?
6. The two major types of compounds are ionic and
molecular.
(a) Compare the naming of these compounds.
(b) Given the name of an example of each
compound, outline how the chemical formula is
obtained. Use specific examples in your answer.
7. Theories are created to explain observations. For each
of the following properties of ionic compounds, write
a brief theoretical explanation.
(a) Ionic compounds are hard solids with high
melting and boiling points.
NEL
12. Classify each of the molecules represented in the
previous question as polar or nonpolar. Justify your
answer using the molecular shape and bond dipoles
(charge distributions).
13. Methylisocyanate is a toxic pesticide that is
manufactured using the following chemical reaction.
CS2 CH3NH2 → CH3NCS H2S
Rewrite this chemical equation using structural
formulas for all reactants and products.
14. Define the three types of intermolecular forces. For
each type of force, state how you would know if this
type of force is likely present among molecules of a
substance.
15. Each of the following four substances is either a
liquid at SATP or converted to a liquid by changing
the conditions: C3H7F, C3H5(OH)3, C3H7NH2, C3H8
(a) Construct a table to summarize the types of
intermolecular forces believed to be present
among molecules of each of these substances.
(b) Predict the order of boiling points from lowest to
highest. Justify your answer.
Review of Chemistry 20 817
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16. Why are boiling points often used as an indirect
measure of the strength of intermolecular forces
among molecules of a substance?
17. Explain each of the following observations in terms
of the characteristics of molecules and intermolecular
forces.
(a) The boiling point of fluorine is significantly less
than that of chlorine.
(b) Drops of ethanol are attracted to an electrically
charged strip.
(c) Ice has a regular hexagonal structure.
18. A simple, but useful, distinction that is often made is
to classify the water on Earth as either fresh water (as
in most lakes and streams) or salt water (as in the
oceans).
(a) Contrast these two terms from a scientific
perspective.
(b) How is this distinction useful from a
technological perspective?
19. Describe an example in which scientific research led
to the development of a new technology.
Unit 2 Gases (Chapter 4)
1. List seven ways by which empirical knowledge is
communicated.
2. List the three characteristics of acceptable scientific
laws and generalizations.
3. Describe one natural phenomenon and one
technological product that each depend on the
properties of gases.
4. Complete the following statements.
(a) At a constant temperature and chemical amount
of gas, as the pressure increases, the volume
________.
(b) At a constant pressure and chemical amount of
gas, as the temperature decreases, the volume
________.
(c) At a constant volume and temperature, if the
chemical amount of gas inside a container is
increased, the pressure ________.
818
Appendix G
5. Choose one of the statements in question 4 and write
a general design for an experiment to test the
statement. Include the identification of all important
variables.
6. For each statement in question 4, sketch a graph of
the relationship between the manipulated and
responding variables.
7. Convert 95.8 kPa into units of millimetres of mercury
and atmospheres.
8. A 1.5 L volume of gas is compressed at a constant
temperature from 1.0 atm to 5.0 atm. Calculate the
final volume.
9. A balloon can hold 800 mL of air before breaking. A
balloon at 4.0 °C containing 750 mL of air is allowed
to warm up. Assuming a constant pressure inside the
balloon, determine the minimum Celsius
temperature when the balloon breaks.
10. A sample of argon gas at 101 kPa and 22.0 °C
occupies a volume of 150 mL. If the volume doubles
at a temperature of 150 °C, determine the new
pressure.
11. Using the kinetic molecular theory, explain Boyle’s
and Charles’ laws.
12. Illustrate the law of combining volumes using a
simple example. Describe the theory used to explain
this law.
13. Many people use propane barbeques for outdoor
cooking. Predict the volume of carbon dioxide
produced when 15 L of propane completely burns at
SATP.
14. Describe and compare the behaviour of real and ideal
gases using the kinetic molecular theory.
15. Predict the volume that 25.0 g of oxygen gas would
occupy at 22.0 °C and 98.1 kPa.
16. Compare the volume that 0.278 mol of hydrogen
would occupy at STP and SATP.
17. An average bungalow requires about 400 kmol of
methane per year for space heating.
(a) Determine the volume of methane at SATP used
in one year.
(b) Predict the volume of methane used if the
pressure is 98.5 kPa and the temperature is
12.7 °C.
NEL
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Appendix G
Unit 3 Solutions, Acids
and Bases (Chapters 5 & 6)
1. For each of the following perspectives write a brief
statement describing the focus or concern of that
point of view.
• scientific
• technological
10. Explain, in terms of breaking and forming bonds,
why the dissolving of substances in water can be
either exothermic or endothermic.
11. Compounds may be ionic or molecular and may also
be acids, bases, or neutral compounds.
• economic
(a) Design an experiment to classify the solute in
each of a number of different solutions.
• ecological
(b) Outline the expected results.
• political
2. List three topics that are current STS issues.
3. Classify each of the following statements using one of
the issue perspectives listed in question 1. All of the
statements concern sulfur dioxide emissions.
(a) An industry spokesman reported that emissions
of sulfur dioxide were within the limits set by
environmental legislation.
(b) Laboratory research has provided evidence that
sulfur dioxide from the combustion of fossil fuels
is converted to sulfur trioxide in the presence of
oxygen.
(c) The cost of ending sulfur dioxide pollution of the
atmosphere will be high. The longer we delay
facing the problem, the greater will be the cost.
(d) Sulfur oxides and their related dissolved acids are
particularly damaging to soil microbes, water life
forms, plants, building materials, and people.
(e) One of the most promising scrubbers to remove
sulfur dioxide gas from a smoke stack is the
limestone−dolomite process.
4. Compare the goals of science and technology.
5. Describe a homogeneous mixture and provide several
examples.
6. Define the two main parts of a solution. State an
example using a chemical formula and identity the
two parts in words.
7. In the exploration of outer space, scientists usually
look for the presence of water as a strong indication
of the existence of living things. Briefly explain this
statement in terms of solutions and reactions.
8. List at least six examples of manufactured solutions
found in the home and six examples of natural
solutions found in the environment.
NEL
9. Distinguish between electrolytes and non-electrolytes,
including examples of types of substances in each
category.
12. Write dissociation or ionization equations for the
following pure substances dissolving in water.
(a) lithium phosphate solid
(b) hydrogen chloride gas
(c) aluminium sulfate solid
13. For each of the following pure substances, write the
formulas for the entities present when each substance
is placed in water.
(a) Sr(OH)2(s)
(d) CH3COOH(l)
(b) HNO3(l)
(e) AgCl(s)
(c) C3H8(g)
(f) CH3OH(l)
14. List the three advantages of solutions for
technological applications.
15. Suppose you are given four unlabelled beakers, each
containing a colourless aqueous solution of one
solute. The possible solutions are NaCl(aq), HCl(aq),
BaCl2(aq), and CH3Cl(aq). Write a series of
diagnostic tests to distinguish each solution from the
others.
16. Compare the ways in which solution concentrations
are expressed in chemistry labs, consumer products,
and environmental studies.
17. A household cleaner contains 12.5 g of sodium
hypochlorite in 500 mL of solution. Determine the
percentage mass by volume concentration of this
solution.
18. A drain cleaner contains 2.75 mol/L sodium
hydroxide. Calculate the volume of solution that
contains 0.375 mol of sodium hydroxide.
19. A windshield washer solution was prepared by
dissolving 100 g of methanol in water to form 2.00 L
of solution. Calculate the amount concentration of
the solution.
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20. A 0.251 mol/L calcium chloride solution is required
for an experiment.
(a) Calculate the mass of calcium chloride that needs
to be measured.
(b) Write a specific procedure for an untrained
laboratory technician to prepare this solution.
21. (a) Predict the volume of concentrated, 14.6 mol/L
phosphoric acid required to prepare 250 mL of a
0.375 mol/L solution.
(b) Write a specific procedure to prepare this
solution.
28. Use the modified Arrhenius theory to write chemical
equations explaining the following evidence.
(a) A vinegar solution is acidic.
(b) A baking soda (sodium hydrogen carbonate)
solution has a pH of 8.
(c) Some muriatic (hydrochloric) acid is neutralized
with a lye (sodium hydroxide) solution.
29. A simple window cleaning solution containing
0.25 mol/L ammonia has a pOH of 2.5.
(a) Convert the pOH into an amount concentration
of hydroxide ions.
22. Calculate the amount concentration of each ion in a
2.1 mol/L solution of iron(III) chloride?
(b) Write a balanced chemical equation to explain
this basic solution.
23. How does the solubility of solids and gases change as
the temperature increases?
(c) Is ammonia a strong or weak base? Justify your
answer.
24. Excess copper(II) sulfate is added to water in a closed
system until no more solute dissolves at a constant
temperature.
30. Write a design for an experiment to identify strong
and weak acids. Include three different diagnostic
tests and identify important controlled variables.
(a) Describe some empirical properties of this
mixture.
(b) Provide a brief theoretical explanation of these
properties.
25. Write the acid formula for each of the following
substances.
(a) aqueous hydrogen bromide
(b) aqueous hydrogen sulfite
(c) hydrofluoric acid
(d) sulfuric acid
26. Copy and complete the following table.
Table 2 Hydroxide Concentrations and pHs
[H3O+(aq)] (mol/L)
pH
Acidic/basic/neutral
7
1.0 10
8
3.7
6.23 109
27. The pH of pure water is 7, of carbonated water about
5, and of a cola drink about 3. How many times more
acidic is a cola drink than carbonated water and pure
water?
31. Polyprotic acids and bases occur naturally and are
manufactured for a variety of purposes.
(a) Distinguish between monoprotic and polyprotic
acids and bases.
(b) Using boric acid (aqueous hydrogen borate) as an
example, write a series of chemical equations
showing successive reactions with water.
32. Most scientists agree that the increasing emission of
carbon dioxide into the atmosphere from the burning
of fossil fuels is the prime cause of global warming.
This problem might be even worse if it were not for
the fact that approximately half of the carbon dioxide
produced is absorbed by the world’s oceans. However,
recent research has shown that this is making the
oceans more acidic—about 30% more acidic over the
past two hundred years.
(a) Use the modified Arrhenius theory to write a
chemical equation explaining the increased
acidity of the world’s oceans.
(b) Scientists are not certain what effect the increased
acidity will have. If we assume there will be a
problem in the oceans, describe some solutions to
reduce the addition of carbon dioxide to the
oceans.
33. Using pesticides as an example, summarize the
intended and unintended consequences of this
chemical technology.
820
Appendix G
NEL
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Appendix G
Unit 4 Quantitative
Relationships
(Chapters 7 & 8)
1. Compare the fields of chemistry and chemical
technology.
2. Describe two examples of chemical technologies,
used by consumers, that are based on the
stoichiometry of chemical reactions.
3. Distinguish between qualitative and quantitative
chemical analysis and provide an example of each
type of analysis.
4. For each of the following mixtures, write a balanced
net ionic equation and identify all spectator ions. All
reactant solutions are assumed to be at least
0.10 mol/L in concentration.
(a) sodium hydroxide and cobalt(II) chloride
solutions
(b) silver nitrate and calcium iodide solutions
(c) silver nitrate solution and zinc metal
(d) hydrochloric acid and solid calcium hydroxide
(e) the precipitation of aluminium hydroxide in
qualitative analysis
5. In your own words, describe the meaning of
stoichiometry.
6. List the three types of stoichiometry and describe
how each type is recognized.
7. In general, how do chemical industries use the
principles of stoichiometry to maximize yields and
minimize waste?
8. In the steel industry, carbon reacts with iron(III)
oxide (from iron ore) to produce molten iron and
carbon dioxide.
(a) Write a complete balanced chemical equation for
this reaction.
(b) Translate this chemical equation into an English
sentence including all chemical amounts and
states of matter.
(c) Using the coefficients, calculate the mass of each
reactant and product in this balanced chemical
equation.
9. Predict the mass of lead(II) iodide precipitate that
forms when 2.93 g of potassium iodide in solution
reacts with excess lead(II) nitrate.
10. In a hard water analysis, a calcium chloride solution is
reacted with excess aqueous sodium oxalate to
produce 0.452 g of calcium oxalate precipitate.
Determine the mass of calcium chloride present in
the original solution.
11. Analysis for sulfate ions is usually done by first
precipitating barium sulfate from a sample. The filter
paper containing the barium sulfate precipitate is
then ignited. Carbon from the burnt filter paper then
reacts with the barium sulfate as shown in the
balanced chemical equation below.
BaSO4(s) + 2 C(s) → BaS(s) + 2 CO2(g)
(a) Predict the mass of carbon required to react with
1.50 g of barium sulfate precipitate.
(b) List the assumptions you have made in this
calculation.
12. In a test of the stoichiometric method, an excess of
sodium hydroxide solution is reacted with a solution
containing 1.50 g of aluminium sulfate.
(a) Predict the mass of precipitate expected in this
reaction.
(b) If the actual yield in this experiment was 0.96 g of
precipitate, calculate the percent yield.
(c) Outline at least three possible reasons for the
discrepancy between the theoretical (predicted)
yield and the actual yield.
13. Powdered aluminium metal is one of the fuels used in
the solid rocket boosters for the NASA Space Shuttle.
What volume of oxygen at SATP is required to react
completely with 100 kg of aluminium?
14. A portable hydrogen generator uses the reaction of
solid calcium hydride and water to form calcium
hydroxide and hydrogen. Determine the volume of
hydrogen at 96.5 kPa and 22 °C that can be produced
from a 50 g cartridge of CaH2(s).
15. A volumetric analysis shows that it takes 32.0 mL of
2.12 mol/L NaOH(aq) to completely react with
10.0 mL of sulfuric acid from a car battery. Calculate
the amount concentration of sulfuric acid in the
battery solution.
(d) How does the total mass of reactants compare
with the total mass of products? What principle
does this illustrate?
NEL
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Page 822
16. In a laboratory, silver metal can be recycled to
produce silver nitrate by the following reaction.
22. Titration curves are useful in studying the progress of
a reaction, such as an acid−base reaction.
3 Ag(s) 4 HNO3(aq) →
3 AgNO3(aq) NO(g) 2 H2O(l)
Predict the volume of 15.4 mol/L nitric acid required
to react with 1.68 kg of silver metal.
17. Distinguish between limiting and excess reagents.
18. Describe the purpose of using an excess reagent in a
quantitative analysis?
19. Calcium carbonate is commonly used in simple
antacid products to counteract acidity in the
stomach. Suppose you add a 750 mg tablet of calcium
carbonate to 200 mL of 0.10 mol/L hydrochloric acid
(representing the stomach acid).
(a) Sketch a general curve for the titration of a strong
base with a strong acid. Label the axes and
provide a title for the graph. No numbers are
required.
(b) Place an “X” on the curve where the reaction is
complete. At what pH should this occur?
(c) Identify a suitable indicator for any strong base−
strong acid titration and justify your answer.
(d) Would your answers to (a), (b), and (c) change if
a strong acid were titrated with a strong base?
Note any differences.
(a) Which reactant is in excess and by how much?
Give your answer in moles.
(b) Predict the mass of calcium chloride formed in
this reaction.
20. Complete the Materials and Analysis of the following
lab report.
Problem
What is the amount concentration of an unknown
sodium carbonate solution?
Design
Samples of sodium carbonate solution were titrated
with a standardized hydrochloric acid solution using
methyl orange as the indicator.
Evidence
Table 3 Titration of 25.0 mL Samples of Na2CO3(aq) with
0.352 mol/L HCl(aq)
Trial
1
2
3
4
Final burette reading (mL)
16.5
31.8
47.0
16.4
Initial burette reading (mL)
0.6
16.5
31.8
1.2
822
Appendix G
NEL
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Appendix H
Page 823
DIPLOMA EXAM PREPARATION
You have been preparing for the Diploma Exam throughout
your high school career. In your final year, as you work through
the Chemistry 30 course, here are some tips that will help
you perform as well as you possibly can in the Diploma Exam.
• Involve Yourself in Class: Attend class regularly
and be active in your learning by asking questions
and completing assignments. If you work steadily,
there will be no need to try to learn everything just
before the exam.
• Keep Up-to-Date with Chemistry 30 Material:
Schedule a regular review time every week and use
this time to organize your notes, review the
material, and ask yourself questions about what
you have learned. Use the Self Quizzes, Chapter
Summaries, and other study aids.
• Read and Understand the Scoring Criteria for
Diploma Exams: The full scoring criteria for the
different types of questions are available in the
Biology 30 Information Bulletin found online.
Read these criteria carefully and make sure you
understand what they mean.
www.science.nelson.com
GO
• Practice Writing Old Exams: Simulate the
conditions of the exam to get used to sitting
through an entire exam and the time constraints of
writing the exam. You will also get used to the types
of questions on the exam and, afterward, be able to
compare your answers to the scoring criteria.
• Read the Instructions: Make sure you read the
instructions, directions, and questions very
carefully.
• Become Familiar with the Types of Questions:
Read the information below and practice answering
each type of question.
There are three types of questions on the Diploma Exam:
multiple choice, numerical response, and written response.
Multiple Choice Questions
Multiple choice questions are a large part of the diploma
exam. Most of the multiple choice questions on the diploma
exam are context-dependent. The others are called “discrete.”
Context-dependent multiple choice questions use information provided in addition to the actual question. Examples
of this type of question include questions 10 and 11 in the
Unit 2 Review.
NEL
Use this information to answer questions 9 to 11.
The empirical study of gases provided a number of laws
that formed the basis for important developments in chemistry such as atomic theory and the mole concept.
Statements
1. The volume of a gas varies inversely with the pressure
on the gas.
2. Volumes of reacting gases are always in simple, whole
number ratios.
3. The volume of a gas varies directly with the absolute
temperature of the gas.
4. The volume of a gas varies directly with the absolute
temperature and inversely with the pressure.
10. Which statements require that the temperature be a
controlled variable?
A. 1, 2, 3, and 4
B. 1, 3, and 4 only
C. 1 and 2 only
D. 3 and 4 only
11. Identify the statement that is best explained by
Avogadro’s theory.
A. 1
B. 2
H
C. 3
D. 4
Discrete multiple choice questions have no additional
information or directions, such as questions 1 and 2 in the
Chapter 7 Review.
1. A main goal of technology is to
A. advance science
B. identify problems
C. explain natural processes
D. solve practical problems
2. In the reaction of aqueous solutions of sodium
sulfide and zinc nitrate in a chemical analysis, the
spectator ions are
A. sodium and nitrate ions
B. sulfide and zinc ions
C. sodium and zinc ions
D. sulfide and nitrate ions
Diploma Exam Preparation
823
Appendix G-I_Chem20
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Page 824
When answering multiple choice questions:
• Try to answer the question before looking at the
choices.
• Eliminate any choices that are incorrect by crossing
them out.
• Stay alert for key words: most, least, not one of the
following, etc. Negative terms (“Which of the
following is not a physical property of a gas?”) will
be italicized. Look out for these.
• Choose the correct answer on the question sheet
and then fill in the corresponding circle on the
answer sheet. It is important that you stay aware of
time, so that you don’t run out of time to
transcribe your answers from the question sheet to
the answer sheet.
Numerical Response Questions
Numerical response questions on the Diploma Exam are
clearly indicated with the heading “NUMERICAL
RESPONSE.” Examples of these types of questions are clearly
marked with the icon “NR” in this textbook. There are four
types of numerical response questions on the Diploma Exam.
They are
• calculation of numerical values,
• calculation of numerical values expressed in
scientific notation,
• selecting numerical responses from diagrams or
lists, and
• determining the sequence of listed events.
Specific instructions for recording the answer to each type
of numerical response are given in the instructions of the
Diploma Exam, as well as with each question. Read the instructions carefully.
• Numerical calculations: This category is fairly
straightforward. You have to use the provided data
to calculate an answer. The answer is a numerical
response with a maximum of four digits (including
the decimal point). The first digit of your answer
goes in the left-hand box on the answer sheet.
Depending on the number of digits in your answer,
there may be unfilled boxes to the right. The
decimal point, if there is one, occupies one of the
boxes. If an answer has a value between 0 and 1, for
example 0.25, make sure you record the ‘0’ before
the decimal point.
824
Appendix H
• Calculations requiring an answer in scientific
notation: This category is similar to regular
numerical calculations, except that the four digits
of the response come from an answer in the form
a.b × 10cd. When you have completed your
calculation, just write the four digits represented by
a, b, c, and d in order.
• Numerical responses from diagrams or lists: This
category involves selecting numbers (usually
representing a term or item from several provided)
and writing them in the correct order.
• Sequence of numbered events or data: These
questions ask you to rearrange variables, events, or
data into a specified order. Pay particular attention
to the instructions, which might specify, for
example, “in order of increasing melting
temperature.”
Written Response Questions
There are two written response questions on the Chemistry
30 Diploma Exam. One written response question is a closedresponse question (which has only one correct response) and
the other is an open-response question (which has more than
one correct response).
Learn to determine which type of question is being asked.
Closed response questions have specific questions that must
be directly answered. These questions are presented as sections and subsections (question 1. a, b, c, etc.). Open response
questions, or unified response questions, typically begin with
the phrase “Write a unified response...” The question is asked
as a series of bullets, and the answers are written in full sentences. Each bullet must be addressed and combined or “unified” into the answer.
When answering written response questions
• Carefully read the information box and make sure
you fully understand the material and all of the
question parts before beginning to answer.
• Identify each key piece of information and make
notes about the meaning and implications of that
information. If it helps, mark key words and
phrases. Identify which unit of Chemistry 30 is
being addressed. This will help you focus your
attention to the correct material.
• Identify any irrelevant information.
NEL
Appendix G-I_Chem20
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Appendix H
• Identify the directing words in the question. These
are usually highlighted in bold in the question. The
directing words have specific meanings and are
indicators of what the graders expect for an answer.
Examples of directing words include illustrate,
analyze, explain, and predict. A complete list of
directing words and their meanings can be found
online. These words are also included and defined
in the Glossary. In your preparation, refer to the list
of meanings for directing words. Make sure that
you know what is expected for each directing word.
www.science.nelson.com
GO
• Read the question carefully and ask yourself what
you are being asked to do. Write the question out in
your own words if there are any doubts. Remember,
if you don’t understand the question, you will
probably not be able to answer it correctly!
• Summarize your answers on scrap paper before
writing them on the test answer page.
• Once you have answered the question, review your
answer and make sure you have addressed all parts
of the question. This is especially important for the
open response question.
Answering Closed Response Questions
Closed response questions are often based on a summary
of current research or a scenario, and data may be given in
graph or table format. There are several parts to the question, but the number of parts depends on the context of the
question. An example of a closed response question follows.
This one is taken from the Chapter 13 Review.
35. Vanadium is a very versatile element in terms of its
reactivity. Vanadium metal reacts with fluorine to
form VF5, with chlorine to form VCl4, with bromine
to form VBr3, with iodine to form VI2, with oxygen to
form V2O5, and with hydrochloric acid to form VCl2.
(a) Identify the oxidation states of vanadium in each
of these compounds.
(b) What interpretation can be made about the
oxidizing power of the chemicals that react with
vanadium metal?
(c) Describe how the oxidation state of vanadium
relates to the colours of the compounds formed.
(d) Report on some technological applications of
vanadium and its compounds.
Each section of the question must be answered completely
for full marks. The number of marks for the question is given
in parentheses. Use this as a guide for how detailed your
answer should be.
Answering Open Response Questions
Open response questions are often based on a situation or
scenario. You are generally asked to write an essay-type
response, guided by the directing word in the question.
Following the initial question are several points. Your responses
to these points should be integrated into your answer.
34. For the production of pulp from wood, a variety of
methods are used, including mechanical and
chemical processes. These have advantages and
disadvantages that have been widely debated. Prepare
an argument for or against the following statement:
“The immediate economic value of using technology
to produce a product far outweighs any possible
future adverse effects.”
Your response should also include
• researched information about a variety of
mechanical and chemical processes
• an evaluation of these processes from
technological, economic, and ecological
perspectives
• reference to redox chemistry
You will do best answering these questions if you refer to current scientific advances in your answer, as you address the
technological and societal aspects of the question. Try to stay
up-to-date on current events by reading the newspaper, science magazines, or reliable science Web sites on a regular
basis. Make sure that your answer includes both scientific
and technology and society aspects.
Write your answer out in full. When you have completed
your answer, recheck it against the bullets, making sure all
parts of the question have been addressed.
The open response questions are scored against two separate scales: a science scale, and a technology and society scale.
The scores are: 0 (insufficient), 1 (poor), 2 (limited),
3 (satisfactory), 4 (good) and 5 (excellent). The highest score
(5) is given for clear, complete answers that address all of the
directing words and give more than one example or piece of
information for each bullet in the question. The lowest
response (0) is given if the answer does not address the questions presented, or is too brief to assess.
Complete examples of closed-response questions are online.
www.science.nelson.com
NEL
GO
Diploma Exam Preparation
825
H
Appendix G-I_Chem20
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10:43 AM
Page 826
DATA TABLES
Appendix I
THERMODYNAMIC PROPERTIES OF SELECTED
ELEMENTS*
Formula
∆fusHm0
(kJ/mol)
∆vapHm0
(kJ/mol)
aluminium
Al
10.79
294
argon
Ar
1.18
beryllium
Be
7.90
Name
6.43
THERMODYNAMIC PROPERTIES OF SELECTED
COMPOUNDS*
c
(J/(g°C))
Name
Formula
0.897
ice
H2O(s)
c
(J/(g°C))
—
2.00
0.520
water
H2O(l)
—
1.825
steam
H2O(g)
—
480
1.026
ammonia
NH3(g)
5.66
23.33
2.06
0.474
methanol
CH3OH(l)
3.22
35.21
2.53
B
50.2
bromine
Br2
10.57
carbon
(graphite)
C
chlorine
Cl2
chromium
Cr
21.0
cobalt
Co
16.06
377
0.421
copper
Cu
12.93
300.4
0.385
fluorine
F2
0.51
gallium
Ga
5.58
254
0.371
germanium
Ge
36.94
334
0.320
gold
Au
12.72
324
0.129
Substance
air
6.40
6.01
∆vapHm0
(kJ/mol)
297
boron
117
∆fusHm0
(kJ/mol)
29.96
—
20.41
339.5
6.62
0.709
0.479
40.65
4.19
—
2.02
ethanol
C2H5OH(l)
4.93
38.56
2.44
Freon-12
CCl2F2(g)
4.14
20.1
0.60
*at 101.325 kPa (1 atm)
0.449
0.824
MISCELLANEOUS SPECIFIC AND VOLUMETRIC HEAT
CAPACITIES
Specific heat
capacity, c
(J/(g°C))
Volumetric heat
capacity, c
(MJ/(m3°C))
1.01
0.0012
4.19
helium
He
0.014
0.08
5.193
hydrogen
H2
0.12
0.90
14.304
water
4.19
0.214
wood
1.26
—
0.449
glass
0.84
—
0.248
polystyrene
0.30
—
0.129
brick/rock
—
1.023
concrete
—
2.1
0.479
ethylene glycol (50%)
—
3.7
0.140
aluminium
0.897
—
1.030
copper
0.385
—
0.444
tin
0.228
—
iodine
I2
15.52
iron
Fe
13.81
krypton
Kr
1.64
lead
Pb
4.78
magnesium
Mg
8.48
manganese
Mn
12.91
mercury
Hg
2.29
neon
Ne
0.33
41.57
340
9.08
179.5
128
221
59.1
1.71
nickel
Ni
17.04
nitrogen
N2
0.71
5.57
oxygen
O2
0.44
6.82
phosphorus
P4
0.66
platinum
Pt
22.17
radon
Rn
3.25
scandium
Sc
selenium
Se
6.69
silicon
Si
50.21
359
0.705
silver
Ag
11.28
258
0.235
sulfur
S8
1.72
45
0.710
tin
Sn
7.17
296.1
0.228
14.1
377.5
12.4
469
18.10
332.7
95.48
1.040
0.918
0.769
0.133
0.094
0.568
0.321
titanium
Ti
14.15
425
0.523
tungsten
W
52.31
806.7
0.132
uranium
U
9.14
417.1
0.116
vanadium
V
459
0.489
xenon
Xe
2.27
zinc
Zn
7.07
21.5
12.57
123.6
1.9
0.158
0.388
* molar enthalpies at 101.325 kPa (1 atm) and specific heat capacities for standard
state at SATP
826
Appendix I
NEL
Appendix G-I_Chem20
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Page 827
Appendix I
STANDARD MOLAR ENTHALPIES OF FORMATION
∆fH m°
∆fH m°
NEL
Chemical name
Formula
acetone
aluminium oxide
ammonia
ammonium chloride
ammonium nitrate
barium carbonate
barium chloride
barium hydroxide
barium oxide
barium sulfate
benzene
bromine (vapour)
butane
calcium carbonate
calcium chloride
calcium hydroxide
calcium oxide
calcium sulfate
carbon dioxide
carbon disulfide
carbon monoxide
chloroethene
chromium(III) oxide
copper(I) oxide
copper(II) oxide
copper(I) sulfide
copper(II) sulfide
1,2-dichloroethane
dinitrogen tetraoxide
ethane
ethane-1,2-diol
ethanoic (acetic) acid
ethanol
ethene (ethylene)
ethyne (acetylene)
glucose
hexane
hydrogen bromide
hydrogen chloride
hydrogen fluoride
hydrogen iodide
hydrogen perchlorate
hydrogen peroxide
hydrogen sulfide
iodine (vapour)
iron(II) oxide
iron(III) oxide
iron(II, III) oxide
lead(II) bromide
lead(II) chloride
lead(II) oxide
lead(IV) oxide
magnesium carbonate
magnesium chloride
magnesium hydroxide
magnesium oxide
magnesium sulfate
(CH3)2CO(l)
Al2O3(s)
NH3(g)
NH4Cl(s)
NH4NO3(s)
BaCO3(s)
BaCl2(s)
Ba(OH)2(s)
BaO(s)
BaSO4(s)
C6H6(l)
Br2(g)
C4H10(g)
CaCO3(s)
CaCl2(s)
Ca(OH)2(s)
CaO(s)
CaSO4(s)
CO2(g)
CS2(l)
CO(g)
C2H3Cl(g)
Cr2O3(s)
Cu2O(s)
CuO(s)
Cu2S(s)
CuS(s)
C2H4Cl2(l)
N2O4(g)
C2H6(g)
C2H4(OH)2(l)
CH3COOH(l)
C2H5OH(l)
C2H4(g)
C2H2(g)
C6H12O6(s)
C6H14(l)
HBr(g)
HCl(g)
HF(g)
HI(g)
HClO4(l)
H2O2(l)
H2S(g)
I2(g)
FeO(s)
Fe2O3(s)
Fe3O4(s)
PbBr2(s)
PbCl2(s)
PbO(s)
PbO2(s)
MgCO3(s)
MgCl2(s)
Mg(OH)2(s)
MgO(s)
MgSO4(s)
(kJ/mol)
–248.1
–1675.7
–45.9
–314.4
–365.6
–1213.0
–855.0
–944.7
–548.0
–1473.2
+49.1
+30.9
–125.7
–1207.6
–795.4
–985.2
–634.9
–1434.5
–393.5
+89.0
–110.5
+37.3
–1139.7
–168.6
–157.3
–79.5
–53.1
–126.9
11.1
–84.0
–454.8
–484.3
–277.6
+52.4
+227.4
–1273.3
–198.7
–36.3
–92.3
–273.3
+26.5
–40.6
–187.8
–20.6
+62.4
–272.0
–824.2
–1118.4
–278.7
–359.4
–219.0
–277.4
–1095.8
–641.3
–924.5
–601.6
-1284.9
Chemical name
Formula
(kJ/mol)
manganese(II) oxide
manganese(IV) oxide
mercury(II) oxide (red)
mercury(II) sulfide (red)
methanal (formaldehyde)
methane
methanoic (formic) acid
methanol
methylpropane
nickel(II) oxide
nitric acid
nitrogen dioxide
nitrogen monoxide
nitromethane
octane
ozone
pentane
phenylethene (styrene)
phosphorus pentachloride
phosphorus trichloride
(liquid)
phosphorus trichloride
(vapour)
potassium bromide
potassium chlorate
potassium chloride
potassium hydroxide
propane
silicon dioxide (a–quartz)
silver bromide
silver chloride
silver iodide
sodium bromide
sodium chloride
sodium hydroxide
sodium iodide
sucrose
sulfur dioxide
sulfur trioxide (liquid)
sulfur trioxide (vapour)
sulfuric acid
tin(II) chloride
tin(IV) chloride
tin(II) oxide
tin(IV) oxide
2,2,4-trimethylpentane
urea
water (liquid)
water (vapour)
zinc oxide
zinc sulfide
MnO(s)
MnO2(s)
HgO(s)
HgS(s)
CH2O(g)
CH4(g)
HCOOH(l)
CH3OH(l)
C4H10(g)
NiO(s)
HNO3(l)
NO2(g)
NO(g)
CH3NO2(l)
C8H18(l)
O3(g)
C5H12(l)
C6H5CHCH2(l)
PCl5(s)
–385.2
–520.0
–90.8
–58.2
–108.6
–74.6
–425.0
–239.2
–134.2
–240.6
–174.1
+33.2
+91.3
–113.1
–250.1
+142.7
–173.5
+103.8
–443.5
PCl3(l)
–319.7
PCl3(g)
KBr(s)
KClO3(s)
KCl(s)
KOH(s)
C3H8(g)
SiO2(s)
AgBr(s)
AgCl(s)
AgI(s)
NaBr(s)
NaCl(s)
NaOH(s)
NaI(s)
C12H22O11(s)
SO2(g)
SO3(l)
SO3(g)
H2SO4(l)
SnCl2(s)
SnCl4(l)
SnO(s)
SnO2(s)
C8H18(l)
CO(NH2)2(s)
H2O(l)
H2O(g)
ZnO(s)
ZnS(s)
–287.0
-393.8
–397.7
–436.5
–424.6
–103.8
–910.7
–100.4
–127.0
–61.8
–361.1
–411.2
–425.6
–287.8
–2226.1
–296.8
–441.0
–395.7
–814.0
-325.1
-511.3
–280.7
–577.6
–259.2
–333.5
–285.8
–241.8
–350.5
–206.0
I
• Standard molar enthalpies (heats) of formation are measured at SATP (25 °C
and 100 kPa). The values were obtained from The CRC Handbook of Chemistry
and Physics.
• The standard molar enthalpies of elements in their standard states are defined
as zero.
Data Tables
827
Appendix G-I_Chem20
11/1/06
10:44 AM
Page 828
RELATIVE STRENGTHS OF OXIDIZING AND REDUCING AGENTS
SOA
Strongest
Oxidizing
Agents
DECREASING STRENGTH OF OXIDIZING AGENTS
F2(g)
–
+
PbO2(s) + SO42 (aq) + 4 H (aq)
–
+
MnO4 (aq) + 8 H (aq)
3+
Au (aq)
–
+
ClO4 (aq) + 8 H (aq)
Cl2(g)
+
2 HNO2(aq) + 4 H (aq)
+
2–
Cr2O7 (aq) + 14 H (aq)
+
O2(g) + 4 H (aq)
+
MnO2(s) + 4 H (aq)
–
+
2 IO3 (aq) + 12 H (aq)
Br2(l)
+
Hg2 (aq)
–
ClO (aq) + H2O(l)
+
Ag (aq)
–
+
2 NO3 (aq) + 4 H (aq)
3+
Fe (aq)
+
O2(g) + 2 H (aq)
–
MnO4 (aq) + 2 H2O(l)
I2(s)
+
Cu (aq)
O2(g) + 2 H2O(l)
+
Cu2 (aq)
+
2–
SO4 (aq) + 4 H (aq)
4+
Sn (aq)
+
Cu2 (aq)
+
S(s) + 2 H (aq)
AgBr(s)
+
2 H (aq)
+
Pb2 (aq)
+
Sn2 (aq)
AgI(s)
+
Ni2 (aq)
+
Co2 (aq)
+
H3PO4(aq) + 2 H (l)
PbSO4(s)
+
Se(s) + 2 H (aq)
+
Cd2 (aq)
+
Cr3 (aq)
2+
Fe (aq)
NO2 (aq) + H2O(l)
Ag2S(s)
+
Zn2 (aq)
+
Te(s) + 2 H (aq)
2 H2O(l)
+
Cr2 (aq)
Se(s)
SO42–(aq) + H2O(l)
+
Al3 (aq)
+
Mg2 (aq)
+
Na (aq)
2+
Ca (aq)
+
Ba2 (aq)
+
K (aq)
+
Li (aq)
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
–
2e
–
2e
–
5e
–
3e
–
8e
–
2e
–
4e
–
6e
–
4e
–
2e
–
10 e
–
2e
–
2e
–
2e
–
e
–
2e
–
e
–
2e
–
3e
–
2e
–
e
–
4e
–
2e
–
2e
–
2e
–
e
–
2e
–
e
–
2e
–
2e
–
2e
e–
2 e–
2 e–
2 e–
2 e–
2 e–
2 e–
e–
2 e–
e2 e–
2 e–
2 e–
2 e–
2 e–
2 e–
2 e–
3 e–
2 e–
e–
2 e–
2 e–
e–
e–
Reducing agents
–
2 F (aq)
PbSO4(s) + 2 H2O(l)
+
Mn2 (aq) + 4 H2O(l)
Au(s)
–
Cl (aq) + 4 H2O(l)
–
2 Cl (aq)
N2O(g) + 3 H2O(l)
+
2 Cr3 (aq) + 7 H2O(l)
2 H2O(l)
+
Mn2 (aq) + 2 H2O(l)
I2(s) + 6 H2O(l)
–
2 Br (aq)
Hg(l)
–
–
Cl (aq) + 2 OH (aq)
Ag(s)
N2O4(g) + 2 H2O(l)
+
Fe2 (aq)
H2O2(l)
MnO2(s) + 4 OH–(aq)
–
2 I (aq)
Cu(s)
–
4 OH (aq)
Cu(s)
H2SO3(aq) + H2O(l)
+
Sn2 (aq)
+
Cu (aq)
H2S(aq)
–
Ag(s) + Br (aq)
H2(g)
Pb(s)
Sn(s)
Ag(s) + I–(aq)
Ni(s)
Co(s)
H3PO3(aq) + H2O(l)
Pb(s) + SO42–(aq)
H2Se(aq)
Cd(s)
+
Cr2 (aq)
Fe(s)
NO(g) + 2 OH–(aq)
2 Ag(s) + S2–(aq)
Zn(s)
H2Te(aq)
H2(g) + 2 OH–(aq)
Cr(s)
Se2–(aq)
SO32–(aq) + 2 OH–(aq)
Al(s)
Mg(s)
Na(s)
Ca(s)
Ba(s)
K(s)
Li(s)
E°r (V)
+2.87
+1.69
+1.51
+1.50
+1.39
+1.36
+1.30
+1.23
+1.23
+1.22
+1.20
+1.07
+0.85
+0.84
+0.80
+0.80
+0.77
+0.70
+0.60
+0.54
+0.52
+0.40
+0.34
+0.17
+0.15
+0.15
+0.14
+0.07
0.00
–0.13
–0.14
–0.15
–0.26
–0.28
–0.28
–0.36
–0.40
–0.40
–0.41
–0.45
-0.46
–0.69
–0.76
–0.79
–0.83
–0.91
–0.92
–0.93
–1.66
–2.37
–2.71
–2.87
–2.91
–2.93
–3.04
DECREASING STRENGTH OF REDUCING AGENTS
Oxidizing agents
SRA
Strongest
Reducing
Agents
• 1.0 mol/L solutions at 25 °C and 1 atm
• Values in this table are taken from The CRC Handbook of Chemistry and Physics.
828
Appendix I
NEL
Appendix G-I_Chem20
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Page 829
Appendix I
RELATIVE STRENGTHS OF AQUEOUS ACIDS AND BASES
Name
Formula
Formula
Name
perchloric acid
HClO4(aq)
ClO4–(aq)
perchlorate ion
–
very large
hydroiodic acid
HI(aq)
I (aq)
iodide ion
very large
hydrobromic acid
HBr(aq)
Br–(aq)
bromide ion
very large
hydrochloric acid
HCl(aq)
Cl–(aq)
chloride ion
very large
sulfuric acid
H2SO4(aq)
HSO4–(aq)
hydrogen sulfate ion
–
nitric acid
HNO3(aq)
NO3 (aq)
hydronium ion
H3O+(aq)
H2O(l)
water
5.4 x 10–2
oxalic acid
HOOCCOOH(aq)
HOOCCOO–(aq)
hydrogen oxalate ion
1.4 x 10–2
sulfurous acid (SO2 + H2O)
H2SO3(aq)
HSO3–(aq)
hydrogen sulfite ion
1.0 x 10
–2
–
nitrate ion
2–
hydrogen sulfate ion
HSO4 (aq)
SO4 (aq)
sulfate ion
6.9 x 10–3
phosphoric acid
H3PO4(aq)
H2PO4–(aq)
dihydrogen phosphate ion
5.6 x 10–3
nitrous acid
HNO2(aq)
NO2–(aq)
nitrite ion
7.4 x 10–4
citric acid*
H3C6H5O7(aq)
H2C6H5O7–(aq)
dihydrogen citrate ion*
–4
–
S T R E N G T H
6.3 x 10
hydrofluoric acid
HF(aq)
F (aq)
1.8 x 10–4
methanoic acid
HCOOH(aq)
HCOO–(aq)
methanoate ion
1.5 x 10–4
hydrogen oxalate ion
HOOCCOO–(aq)
OOCCOO2–(aq)
oxalate ion
9.1 x 10–5
ascorbic acid
H2C6H6O6(aq)
HC6H6O6–(aq)
hydrogen ascorbate ion
–5
fluoride ion
–
6.3 x 10
benzoic acid
C6H5COOH(aq)
C6H5COO (aq)
benzoate ion
1.8 x 10–5
ethanoic (acetic) acid
CH3COOH(aq)
CH3COO–(aq)
ethanoate (acetate) ion
1.7 x 10–5
dihydrogen citrate ion*
H2C6H5O7-
HC6H5O72-
hydrogen citrate ion*
4.5 x 10–7
carbonic acid (CO2 + H2O)
H2CO3(aq)
HCO3–(aq)
hydrogen carbonate ion
–7
2-
3-
O F
4.0 x 10
hydrogen citrate ion*
HC6H5O7
C6H5O7
citrate ion*
8.9 x 10–8
hydrosulfuric acid
H2S(aq)
HS–(aq)
hydrogen sulfide ion
6.3 x 10–8
hydrogen sulfite ion
HSO3–(aq)
SO32–(aq)
sulfite ion
6.2 x 10–8
dihydrogen phosphate ion
H2PO4–(aq)
HPO42–(aq)
hydrogen phosphate ion
A C I D S
–8
–
4.0 x 10
hypochlorous acid
HClO(aq)
ClO (aq)
hypochlorite ion
6.2 x 10–10
hydrocyanic acid
HCN(aq)
CN–(aq)
cyanide ion
5.8 x 10–10
boric acid
H3BO3(aq)
H2BO3–(aq)
dihydrogen borate ion
5.6 x 10–10
ammonium ion
NH4+(aq)
NH3(aq)
ammonia
1.0 x 10
–10
–
phenol
C6H5OH(aq)
C6H5O (aq)
phenoxide ion
4.7 x 10–11
hydrogen carbonate ion
HCO3–(aq)
CO32–(aq)
carbonate ion
2.2 x 10–12
hydrogen peroxide
H2O2(aq)
HO2–(aq)
hydrogen peroxide ion
2.0 x 10–12
hydrogen ascorbate ion
HC6H6O6–(aq)
C6H6O62–(aq)
ascorbate ion
4.8 x 10–13
hydrogen phosphate ion
HPO42–(aq)
PO43–(aq)
phosphate ion
hydrogen sulfide ion
HS (aq)
S2–(aq)
sulfide ion
1.0 x 10–14
water (55.5 mol/L)
H2O(l)
OH–(aq)
hydroxide ion
very small
hydroxide ion
OH–(aq)
O2–(aq)
oxide ion
1.3 x 10
–13
–
B A S E S
D E C R E A S I N G
very large
1.0
O F
Strongest
Acid
very large
Conjugate base
S T R E N G T H
SA
Acid
D E C R E A S I N G
Equilibrium
constant, Ka
(mol/L)
SB
Strongest
Base
* The molecular formula representing (triprotic) citric acid has been compressed here to its simplest form for ease of use when writing
proton transfer equations.
Values in this table are taken from Lange’s Handbook of Chemistry for 25 °C.
NEL
Data Tables
829
I
Appendix G-I_Chem20
11/1/06
Appendix J
10:44 AM
Page 830
COMMON CHEMICALS
You live in a chemical world. As one bumper sticker asks, “What in the world isn’t chemistry?” Every natural and technologically produced substance around you is composed of chemicals. Many of these chemicals are used to make your
life easier or safer, and some of them have life-saving properties. Following is a list of selected common chemicals. The
chemicals marked with an asterisk are to be memorized.
Common name
Recommended name
Formula
Common use/source
acetic acid*
acetone*
acetylene*
ASA (Aspirin®)
baking soda*
battery acid*
bleach
bluestone
brine*
citric acid
CFC
charcoal/graphite*
dry ice*
ethylene*
ethylene glycol*
freon-12
Glauber’s salt
glucose*
grain alcohol*
gypsum
lime (quicklime)*
limestone*
lye (caustic soda)*
malachite
methyl hydrate*
milk of magnesia
MSG
muriatic acid*
natural gas*
PCBs
potash*
road salt*
rotten-egg gas*
rubbing alcohol
sand (silica)
slaked lime*
soda ash*
sugar*
table salt*
washing soda*
vitamin C
ethanoic acid
propanone
ethyne
acetylsalicylic acid
sodium hydrogen carbonate
sulfuric acid
sodium hypochlorite
copper(II) sulfate—(1/5)-water
aqueous sodium chloride
2-hydroxy-1,2,3-propanetricarboxylic acid
chlorofluorocarbon
carbon
carbon dioxide
ethene
ethane-1,2-diol
dichlorodifluoromethane
sodium sulfate—(1/10)-water
D-glucose; dextrose
ethanol (ethyl alcohol)
calcium sulfate—water
calcium oxide
calcium carbonate
sodium hydroxide
copper(II) hydroxide carbonate
methanol (methyl alcohol)
magnesium hydroxide
monosodium glutamate
hydrochloric acid
methane
polychlorinated biphenyls
potassium chloride
calcium chloride or sodium chloride
hydrogen sulfide
propan-2-ol
silicon dioxide
calcium hydroxide
sodium carbonate
sucrose
sodium chloride
sodium carbonate—(1/10)-water
ascorbic acid
CH3COOH(aq)
(CH3)2CO(l)
C2H2(g)
C6H4COOCH3COOH(s)
NaHCO3(s)
H2SO4(aq)
NaClO(s)
CuSO4 •5 H2O(s)
NaCl(aq)
C3H4OH(COOH)3
CxClyFz(l) ; e.g.,C2Cl2F4(l)
C(s)
CO2(g)
C2H4(g)
C2H4(OH)2(l)
CCl2F2(l)
Na2SO4 •10 H2O(s)
C6H12O6(s)
C2H5OH(l)
CaSO4 •2 H2O(s)
CaO(s)
CaCO3(s)
NaOH(s)
Cu(OH)2•CuCO3(s)
CH3OH(l)
Mg(OH)2(s)
NaC5H8NO4(s)
HCl(aq)
CH4(g)
(C6HxCly)2 ; e.g., (C6H4Cl2)2(l)
KCl(s)
CaCl2(s) or NaCl2(s)
H2S(g)
CH3CHOHCH3(l)
SiO2(s)
Ca(OH)2(s)
Na2CO3(s)
C12H22O11(s)
NaCl(s)
Na2CO3 •10 H2O(s)
H2C6H6O6(s)
vinegar
nail polish remover
cutting/welding torch
for pain relief medication
leavening agent
car batteries
bleach for clothing
algicide/fungicide
water-softening agent
in fruit and beverages
refrigerant
fuel/lead pencils
“fizz” in carbonated beverages
for polymerization
radiator antifreeze
refrigerant
solar heat storage
in plants and blood
beverage alcohol
wallboard
masonry
chalk and building materials
oven/drain cleaner
copper mineral
gas-line antifreeze
antacid (for indigestion)
flavour enhancer
in concrete etching
fuel
in transformers
fertilizer
melts ice
in natural gas
for massage
in glass making
limewater
in laundry detergents
sweetener
seasoning
water softener
vitamin
830
Appendix J
NEL