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130 EXERCISE 1 GRAPHING Make a full-page graph of the following data: Put the mass on the vertical axis and the volume on the horizontal axis. VOLUME (ML) 5.0 10.0 5.0 15.0 20.0 25.0 2O .0 25.0 MASS (G) 4.10 15.50 7.70 11.90 31.10 19.70 15.80 39.90 M/V QUESTIONS 1. Could you draw one or two lines that would come close to connecting all the points? 2. For each pair of numbers divide the mass by the volume. Example: The numerals 5.0 and 4.10 are a pair. You would divide 5.0 into 4.10. In the column on the right side of your Data Table enter the results. 3. What regularity exists in your answers for Question 2? 131 EXERCISE 2 DENSITY Formulas for density problems M = mass (g) V = volume (ml) D = density (g/ml) D = M/V V = M/D M = VxD 1. What is the density of a block of wood that weighs 200 grams and has a volume of 400 ml? 2. What volume of liquid bromine will weigh 2000 g? (Density of liquid bromine is 3.10 g/ml.) 3. A rectangular block measures 20 cm x 10 cm x 5.0 cm and has a density of .3 g/ml. What does it weigh? 4. What is the volume of 150 g of nitric acid solution of density 1.35 g/ml? 5. Concentrated sulfuric acid has a density of 1.860 g/ml. How much does 1.00 liter of it weigh? SHOW YOUR WORK. 132 EXERCISE 3 SIGNIFICANT FIGURES CALCULATIONS DIRECTIONS: On a separate sheet of paper, complete the following. A. Determine the number of significant figures in each of the following measurements. 1. 5.432 g 11. 2500 cm 2. 40.319 g 12. 2500.0 cm 3. 146 mm 13. 1.04 x 1014 g 4. 3.256 cm 14. 3.58 x 10-9 L 5. 0.189 in 15. 46.0979 g 6. 439.6 mL 16. 8365.6 g 7. 2873.0 cc 17. 0.002300 mg 8. 99.9 s 18. 7.500 x 108 oz 9. 0.000235 g 19. 3.92 x 10-4 g 10. 144 cL 20. 1.00000 x 103 lbs. B. Add the following measurements: 1. 12 cm + 0.031 cm + 7.969 cm (20. cm) 2. 0.085 cm + 0.062 cm + 0.14 cm (.29 cm) 3. 3.419 g + 3.912 g + 0.00013 g (7.331 g) 4. 30.5 g + 16.82 g + 41.07 g + 85.219 g (173.6 g) 5. 2.3 x 102 g + 4.62 x 102 g + 3.852 x 102 g 133 (1.08 x 103 g) C. Subtract the following measurements: 1. 41.025 cm – 23.28 cm (17.74 cm) 2. 289 g – 43.7 g (245 g) 3. 145.63 mL – 28.9 mL (116.7 mL) 4. 62.47 g – 39.9 g (22.6 g) 5. 40.008 mL – 29.0941 mL (10.914 mL) D. Multiply the following measurements: 1. 2.89 cm x 4.01 cm (11.6 cm2) 2. 17.3 cm x 6.2 cm (110 cm2) 3. 5 cm x 5 cm (20 cm2) 4. 5.0 cm x 5 cm (20 cm2) 5. 5.0 cm x 5.0 cm (25 cm2) 6. 4.8 x 102 m x 2.101 x 103 m (1.0 x 106 m2) 7. 4.218 cm x 6.5 cm (27 cm2) 8. 150.0 m x 4.00 m (6.00 x 102) 9. 282.2 km x 3.0 km (850 km2) 10. 14 x 10-8 m x 3.25 x 10-6 m (4.6 x 10-13) E. Divide the following measurements: 1. 8.071 cm2 / 4.216 cm (1.914 cm) 2. 109.3758 m2 / 5.813 m (18.82 m) 3. 24 789.4 km2 / 43.5 km (5.70 x 102km) 4. 6.058 mm2 / 0.85 mm (7.1 mm) 5. 4.819 cm2 / 9.852 cm (.4891 cm) 6. 139.482 m2 / 68.75 m (2.029 m) 7. 4.23 m2 / 18.941 m (.223 m) 8. 85.621 km2 / 8.05 km (10.6 km) 9. 6.93 x 10-5 cm2 / 4.09 x 10-8 cm (1.69 x 103 cm) 10. 3.005 x 106 m2 / 1.8765 x 104 m (1.601 x 102 m) 134 EXERCISE 4 SUBATOMIC PARTICLES ELEMENT A B C D E F G H I J K L M N O P R S T ATOMIC NUMBER 2 ATOMIC WEIGHT 4 13 ELECTRONS 2 5 PROTONS 2 8 80 76 108 175 31 190 153 16 7 28 59 7 28 28 23 31 12 4 18 93 92 2 8 8 45 98 47 71 NEUTRONS 5 14 28 40 237 146 16 32 135 EXERCISE 5 CHEMICAL FAMILIES 1. A student used the Data Table given below while observing chemical reactions. N.R. = no reaction and R = a reaction. Use the data to answer the questions. Data Table A. D. More observations of Phosphorus: Observations of Potassium: and and and and and and and and and and Argon Chlorine Fluorine Helium Hydrogen Krypton Neon Nitrogen Oxygen Xenon (N.R.) (N.R.) (N.R.) (N.R.) (N.R.) (N.R.) (N.R.) (N.R.) (N.R.) (N.R.) (R) (R) (R) (R) (R) (R) (R) (R) (R) (R) and and and and and Hydrogen (N.R.) Fluorine (N.R.) Chlorine (N.R.) Bromine (N.R.) Iodine (N.R.) (R) (R) (R) (R) (R) E. More observations of Potassium: and Bromine (N.R.) (R) and Iodine (N.R.) (R) F. Observations of Chlorine: B. Observations of Lithium: and and and and and and Nitrogen Argon Helium Krypton Neon Xenon (N.R.) (N.R.) (N.R.) (N.R.) (N.R.) (N.R.) (R) (R) (R) (R) (R) (R) and and and and and Fluorine Argon Helium Krypton Neon Xenon (N.R.) (N.R.) (N.R.) (N.R.) (N.R.) (N.R.) (N.R.) (N.R.) (N.R.) (N.R.) (N.R.) (R) (R) (R) (R) (R) G. Observations of some metals when placed in water Water and: Lithium (N.R.) (R) Sodium (N.R.) (R) Potassium (N.R.) (R) Rubidium (N.R.) (R) Cesium (N.R.) (R) C. Observations of Phosphorus: and and and and and and Lithium Sodium Potassium Rubidium Cesium (R) (R) (R) (R) (R) (R) 136 VI. QUESTIONS 1. Look over your first set of observations (Set A). Which elements do not react with potassium? 2. Look over your second set of observations. Which elements do not react with lithium? 3. Now check your third set of data. Which elements do not react with phosphorus? 4. It should now be apparent to you that there is a regularity in the first three sets of data. The generalization can be made that helium, neon, argon, krypton, and xenon are all nonreactive elements. These gases represent a family of elements that are chemically inert. How many electrons does a neutral atom of each of these have? 5. How many valence electrons do each of these elements have? 6. Examine your data for the observations made of phosphorous. Which gases can be said to be reactive? How many electrons do neutral atoms of these elements have? 7. How many valence electrons do each of these have? 8. Compare these numbers with the numbers of electrons in the inert gas family. Would there be reason to believe that the gases that reacted with phosphorus are also a chemical family? 9. Look over the data for chlorine gas in the sixth set of observations. Is there any regularity? How many electrons do neutral atoms of these elements have? How does this number compare with the number in the inert gas family? Can you make any generalization about these elements? 10. Look over the last set of observations you made. Is there any evidence to support your answer to Question 6? 137 EXERCISE 6 WHAT MAKES A COMPOUND? Predict what ions you would get if you dissolved the following ionic substances. EXAMPLE: MgBr2 A. The + ion is always written first. So you look on the left side of the common ion chart and find the positive ion, which is Mg2+ in this case. B. The - ion is always written last. So you look on the right side of the common ion chart and find the negative ion, which is Br- in this case. Answer: Mg2+ and Br- 1. FeCl3 2. NaI 3. Ba(NO3)2 4. (NH4)3PO4 5. Al2(SO4)3 6. MgSO4 7. Na2CO3 8. CaF2 9. PbI2 10. LiClO 138 EXERCISE 7 FORMULA WRITING DIRECTIONS: On a separate piece of paper, complete the following chart: Bromide Br- Nitrate NO3- Hydroxide OH- Oxide O2- Sulfate SO42- Potassium K+ Silver Ag+ Ammonium NH4+ Mercury (I) Hg22+ Mercury (II) Hg2+ Calcium Ca2+ Copper (II) Cu2+ Iron (II) Fe2+ Iron (III) Fe3+ Antimony(I II) Sb3+ 139 Phosphate PO43- Hydrogen Carbonate HCO3- Chlorate ClO3- Acetate C2H3O2- EXERCISE 8 FORMULA TO NAME DIRECTIONS: On a separate sheet of paper name the following compounds using the stock system. 1.) KCl 2.) NaOH 3.) ZnF2 4.) CuNO2 5.) NaMnO4 6.) KHSO3 7.) Mn (OH)2 8.) CrCl3 9.) (NH4)2SO4 10.) AgCl 11.) PbCrO4 12.) NaOH 13.) NiCl2 14.) Hg2Cl2 15.) Cu(NO3)2 16.) Al2(C2O4)3 17.) Na2Cr2O7 18.) AgClO3 19.) KH2PO4 20.) Cr2S3 21.) NaF 22.) CuNO2 23.) FeBr2 24.) HgCO3 25.) NH4OH 26.) CrCl2 27.) HCl 28.) H2S 29.) Sn(NO3)2 30.) Sn(NO2)4 31.) KOH 32.) Hg(OH)2 33.) KCl 34.) FeCl3 35.) HNO3 36.) 37.) 38.) 39.) 40.) 41.) 42.) 43.) 44.) 45.) 46.) 47.) 48.) 49.) 50.) 51.) 52.) 53.) 54.) 55.) 56.) 57.) 58.) 59.) 60.) 61.) 62.) 63.) 64.) 65.) 66.) 67.) 68.) 69.) 70.) Cu2O Al2(SO4)3 N2O 5 HF Pb(OH)2 NH4NO3 NaHCO3 HgO Zn(NO2)2 H3PO4 CsOH Li2O Ca(OH)2 CaBr2 Fe2O3 H2SO4 FeCO3 SO3 Ba(BrO3)2 Al(OH)3 HClO4 NaC2H3O2 Na2SO3 H2CO3 NH4IO3 LiH CO MgBr2 SnBr2 N2O NH4F AsCl3 KHCO3 K2O ZnO 140 EXERCISE 9 NAME TO FORMULA DIRECTIONS: On a separate sheet of paper, write the correct empirical formula for each of the following: 1. Sodium nitrate 28. Manganese (VII) oxide 2. Aluminum sulfate 29. Mercury (I) oxide 3. Potassium chromate 30. Iron (II) chloride 4. Lithium hydroxide 31. Copper (II) chloride 5. Ammonium phosphate 32. Tin (II) chloride 6. Magnesium nitrate 33. Titanium (IV) iodide 7. Sodium monohydrogen phosphate 34. Nickel (II) fluoride 8. Sodium bicarbonate 35. Manganese (IV) oxide 9. Potassium bisulfate 36. Lead (IV) oxide 10. Hydrogen acetate 37. Nickel (II) oxide 11. Lead (II) chromate 38. Mercury (II) oxide 12. Copper (II) carbonate 39. Cobalt (III) oxide 13. Silver nitrate 40. Copper (II) chlorate 14. Ferrous sulfide 41. Cobalt (III) sulfate 15. Manganous sulfide 42. Manganese (III) sulfate 16. Barium sulfite 43. Iron (III) nitrate 17. Sodium iodide 44. Tin (IV) nitrate 18. Calcium nitrate 45. Cobalt (II) perchlorate 19. Zinc fluoride 46. Chromium (III) sulfate 20. Potassium oxalate 47. Iron (II) hydroxide 21. Cuprous chloride 48. Copper (II) phosphate 22. Manganese (III) chloride 49. Mercury (I) nitrite 23. Iron (III) bromide 50. Lead (II) nitrate 24. Chromium (III) bromide 51. Mercury (I) sulfate 25. Tin (IV) chloride 52. Tin (IV) oxide 26. Chromium (III) oxide 53. Manganese (II) bromide 27. Lead (II) oxide 141 EXERCISE 10 FORMULAS OF COMPOUNDS A student performs an experiment similar to our experiment 5.2. His summary of observations appears below. For convenience he numbered each of his seven precipitates that formed. Ions Present In Each Solution Used NiCl2 BaCl2 Na2SO4 NaOH Ba(OH)2 PPT 1 PPT 2 MgSO4 NiCl2 XX BaCl2 XX XX PPT 3 Na2SO4 XX XX XX NaOH XX XX XX XX Ba(OH)2 XX XX XX XX XX PPT 7 MgSO4 XX XX XX XX XX XX PPT 4 PPT 5 PPT 6 INSTRUCTIONS 1. Fill in the positive and negative ions that are present in each of the six original solutions in the blanks provided in the Data Table--just the kinds of ions, not necessarily the numbers (just OH- not 2OHetc.). 2. Now, recalling that unlike charges attract, list the possible formulas for each of the seven precipitates that formed. Here’s how to do that. For each of the precipitates, write the two chemicals the student mixed to get it. (These are called reactants.) Now exchange the positive ion of each. You need to re-criss-cross the new formulas to get them right. These are called products. There should be two. One of each pair of products is the formula of the precipitate. 3. Go back and eliminate those that can be ruled out by examining the results of the other combinations tried. Lightly cross out those formulas. If you do this carefully, you can reason out the correct formula for each precipitate. In fact, you'll discover that one precipitate is probably a mixture of two solid substances. 142 EXERCISE 11 ASSIGNING OXIDATION NUMBERS Directions: On a separate sheet of paper, assign the oxidation number for each element of the following: 1. H2SeO3 14. OF2 2. Al(NO3)3 15. MnO2 3. BaH2 16. Ca3(PO4)2 4. Na4SiO4 17. CuCrO4 5. NaIO4 18. Na3AsO3 6. Na2O2 19. CaS 7. HBrO 20. KHSO3 8. N2O3 21. CO3-2 9. (NH4)2CO3 22. IO4- 10. K2Cr2O7 23. SCN- 11. Mg(ClO4)2 24. ClO- 12. Sb2O3 25. Hg22+ 13. H3PO4 143 How to Diagram Lewis Structures 1. Find the total number of valence electrons supplied by all the atoms in the structure. The number supplied by each A family element is the same as the group number of the element: a. b. 2. For a negative ion, increase the number by the charge of the ion. For a positive ion, decrease the number by the charge of the ion. Determine the number of electrons that would be required to give two electrons to each hydrogen atom and eight electrons to each of the other atoms individually: Num. e- for individual atoms = 2(num. H atoms) + 8(num. other atoms) 3. The number obtained in step 2 minus the number obtained in step 1 is the number of electrons that must be shared in the final structure: Num. bonding e- = (num. e- for individual atoms) – (total num. e-) 4. One-half the number of bonding electrons (step 3) is the number of covalent bonds in the final structure: Num. e- pair bonds = (num. bonding e-)/2 5. Write the symbols for the atoms present in the structure, arranging them in the way that they are found in the structure. 6. Indicate electron-pair bonds by dashes written between the symbols. Indicate one bond between each pair of symbols, and then use any remaining from the total calculated in step 4 to make multiple bonds. Note that each H atom is limited to one bond. 7. The total number of electrons (step 1) minus the number of bonding electrons (step 3) is the number of unshared electrons: Num. unshared e- = (total num. e-) – (num. bonding e-) Complete the electron octet of each atom (other than the H atoms) by adding dots to represent unshared electrons. 8. Indicate the formal charges of the atoms where appropriate, and evaluate the structure. 144 EXERCISE 12 ELECTRON-DOT SYMBOLS (LEWIS STRUCTURES) DIRECTION: On a separate sheet of paper, construct the electron-dot symbols for each of the following: IONIC COMPOUNDS 1. NaI 5. BaO 2. CaF2 6. K3P 3. AlCl3 7. RbBr 4. Na2S 8. Li3N MOLECULES AND POLYATOMIC IONS 9. HClO3 26. OF2 10. ClO3- 27. H2S 11. ClO2- 28. PCl3 12. ClO- 29. SiO2 13. NH3 30. CO2 14. NH4+ 31. HClO4 15. NH2- 32. N2 16. HNO3 33. H2SO4 17. NO3- 34. PH3 18. HNO2 35. HCN 19. NO2- 36. HClO 20. H2O 37. C2H4 21. H3O+ 38. C2H2 22. OH23. CHI3 24. CH3OH 145 EXERCISE 13 INSTRUCTION Many formulas for substances cannot be explained in terms of ionic bonding. Consider the substance Cl2O. Both the chlorine and the oxygen atom need more electrons for a stable electron population. A model proposed that would allow both atoms to gain electrons is shown in the diagram on this page. The model proposes that atoms overlap and that each atom put one electron into this overlapping region or into an electron "pool." By doing this, each atom appears to gain an electron within its original boundary. For every overlapping region an atom appears to gain one electron. Two overlapping regions, such as for oxygen, will result in the gain of two electrons. This process of overlapping atoms is called covalent bonding. The substance that results from covalent bonding is called a covalent substance. The process of overlapping atoms will keep occurring for a particular atom until it has a stable number of electrons in its original boundary. In general, covalent substances form when both atoms want to gain electrons. Usually both atoms will be found toward the right side of the periodic chart. An important difference between ionic and covalent bonding is that in covalent bonding ions are not formed. The group of atoms that forms as a result of covalent bonding is electrically neutral. As a whole there is the same number of electrons as protons. The group of atoms is referred to as a molecule. Formation of Cl2O 146 EXERCISE 13 COVALENT BONDING WORKSHEET ON COVALENT BONDING 1. Make a diagram similar to those on formation of Cl2O and Br2 for the bonding arrangement in the following molecules: A. Br2O B. NH3 C. CH4 D. PF3 Show on your diagram: 1. The number of electrons involved or not involved with bonding 2. Each covalent bond 147 EXERCISE 14 INFORMATION 148 EXERCISE 14 PREDICTING FORMULAS OF COMPOUNDS Using the Chart "Some Compounds of Third Row Elements" predict the formulas of the compounds produced by combining the following elements: 1. Calcium and Chlorine 2. Boron and Oxygen 3. Nitrogen and Hydrogen 4. Hydrogen and Oxygen 5. Fluorine and Oxygen 6. Selenium and Fluorine 7. Carbon and Bromine 8. Magnesium and Tellurium 9. Gallium and Sulfur 10. Sodium and Iodine 11. Barium and Selenium 12. Germanium and Oxygen 13. Indium and Iodine 14. Arsenic and Hydrogen 15. Antimony and Selenium 149 EXERCISE 15 FORMULA MASS DIRECTIONS: On a sheet of paper, determine the formula mass of each of the following: 1. H3PO4 16. K4Fe(CN)6 2. AlCl3 17. Nd2O3 3. Dy(OH)3 18. Sb(NO3)3 4. K2C4H4O6 19. K3PO4 5. H2SO4 20. Ga2(SO4)3 6. N2O5 21. zinc acetate 7. HNO3 22. copper (I) sulfate 8. NiSO4 23. carbon dioxide 9. Sn(OH)2 24. calcium bicarbonate 10. (NH4)3PO4 25. carbonic acid 11. Fe(C2H3O2)3 26. aluminum nitrate 12. SO2 27. ammonium sulfate 13. AgNO3 28. barium chloride 14. NaIO4 29. iron (II) phosphate 15. Pr(OH)3 30. strontium sulfite 150 EXERCISE 16 MOLE THEORY Directions: On a separate sheet of paper, complete the following table using mole theory. SUBSTANCE 1. BaBr2 GRAM FORMULA MASS QUANTITY IN GRAMS NUMBER OF MOLES 35.0 g 0.120 42.0 g XX 5. H2S 0.223 12.0X1023 6. Bi2O3 XX 5.00 g 8. Cu(OH)2 0.100 XX 3.00X1022 9. FeCl2 10. B2H6 26.7 11. CO2 1.80 XX XX 6.00X1022 12. LiH 13. Na2SO4 XX 6.00X1024 3. C2H6 7. H2O VOLUME IN LITERS AT STP XX 2. AuCl3 4. PbCO3 NUMBER OF MOLECULES 36.9 g XX XX 14. FeCl2 13.2 XX 7.94X1024 15. C6H6 151 EXERCISE 17 EMPIRICAL FORMULAS Directions: On a separate sheet of paper, calculate the empirical formula for each of the following: 1. 88.8 % copper and 11.2 % oxygen 2. 40.0 % carbon; 6.67 % hydrogen and 53.33 % oxygen 3. 92.3 % carbon; and 7.7 % hydrogen 4. 70.0 % iron and 30.0 % oxygen 5. 5.88 % hydrogen and 94.12 % oxygen 6. 79.90 % copper and 20.10 % oxygen 7. 56.4 % potassium; 8.7 % carbon and 34.9 % oxygen 8. 10.04 % carbon; 0.84 % hydrogen and 89.12 % chlorine 9. 42.50 % chromium and 57.50 % chlorine 10. 15.8 % carbon and 84.2 % sulfur 11. 30.43 % nitrogen and 69.57 % oxygen 12. 82.4 % nitrogen and 17.60 % hydrogen 13. 12.5 % hydrogen; 37.5 % carbon and 50.0 % oxygen 14. 75.0 % carbon and 25.0 % hydrogen 15. 29.40 % calcium; 23.56 % sulfur and 47.04 % oxygen 152 EXERCISE 18 TYPICAL MOLE PROBLEMS 1. Determine the number of moles of each kind of atom in each of the following: 5 moles C3H5Br 7 moles Ni(OH)2 1 mole CO2 2 moles (NH4)2SO4 2. Find the molecular weight of these: H2O F2 NaNO3 KH2PO4 3. Find the number of moles of each compound: 36g H2O 4. 38g F2 Given Ba(NO3)2 + NaCl BaCl2 + NaNO3 How many moles of NaCl react with: 1 mole Ba(NO3)2 5 moles Ba(NO3)2 5. How many moles of copper atoms are in 2 moles of CuSO4? 6. How many moles of molecules would there be in 10.g of H2O CaCO3 7. If you had 5 moles of CH3CCl3 molecules, how many grams would they weigh? 8. How many moles of particles are in 300. x 1023 particles? 9. Which would contain more molecules : 20. g of NH4Cl or 20. g of H20? 10. How many moles of oxygen atoms are in 10 moles of Fe3(PO4)2 molecules? 153 EXERCISE 19 BALANCING EQUATIONS Directions: On a separate sheet of paper, balance the following equations by inspection: 1. Na + Br2 2. Al 3. P + O2 4. N2O4 5. HClO3 6. Al2O3 + 7. C12H22O11 8. NH4NO3 9. H3PO3 10. Sb2S3 + 11. (NH4)2Cr2O7 12. C4H10 13. TiCl4 + H2O 14. H2SO3 15. Fe 16. CaO 17. Sc2O3 + 18. Pb + H 2O + O2 Pb(OH)2 19. Ag + H 2S + O2 Ag2S 20. P4 + S8 21. NaF 22. Al 23. N2 + 24. Zn 25. C12H26 + S Al2S3 P2O5 NO2 ClO2 + Cl2 + PH3 + H2S Cr2O3 + H2O + N2 CO2 + H2O TiO2 + HCl SO2 + H2O Fe3O4 + SO3 HNO3 O2 H2 Ca3(PO4)2 Sc2(SO4)3 + H 2O P4S10 FeCl3 + SbCl3 + P2O5 H2 + CO H2O H3PO4 + O2 HCl AlCl3 H2O H2O + + O2 + + C N2 O HCl + H2O C+ + + NaBr Na3FeF6 + AlCl3 + NaCl H2 NH3 Zn(NO3)2 + CO2 + H2O 154 H2 EXERCISE 20 COMPLETE, BALANCE AND TYPE OF REACTION Directions: On a separate sheet of paper, complete; balance and indicate the type of reaction for the following: 1. iron + copper (II) sulfate 2. potassium chloride + mercury (I) sulfate 3. sulfur + oxygen 4. sulfuric acid + cadmium 5. zinc + silver sulfate 6. aluminum sulfite + sodium hydroxide 7. iron (III) oxide + hydrogen 8. sodium sulfate + barium bromide 9. mercury (II) nitrate + potassium sulfate 10. nitrogen + hydrogen 11. silver nitrate + sodium carbonate 12. mercury + oxygen 13. sodium hydroxide + barium bromide 14. silicon + sulfur 15. antimony (III) sulfide + iron 16. potassium chloride + aluminum chloride 17. copper + oxygen 18. calcium sulfide + hydrochloric acid 19. antimony + sulfur 20. iron (III) oxide + aluminum 21. mercury (II) oxide 22. iron (III) bromide + barium hydroxide 23. magnesium + lead (II) sulfate 24. aluminum + nitrogen 25. hydrosulfuric acid + iron (II) chloride 155 EXERCISE 21 COMPLETE AND BALANCE 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26. 27. 28. 29. 30. 31. 32. 33. 34. 35. 36. 37. 38. 39. 40. 41. 42. 43. 44. 45. SODIUM CARBONATE LITHIUM + NITRIC ACID CARBON + SULFUR SILVER ACETATE + POTASSIUM CHLORATE ACETIC ACID + POTASSIUM CARBONATE OXYGEN + TITANIUM LEAD (II) PHOSPHATE + SODIUM BROMIDE C7H16 BURNS CESIUM OXALATE + GOLD (III) NITRITE IRON (III) CHLORIDE CALCIUM + SILICON ALUMINUM PHOSPHATE + MOLYBDENUM (VI) ACETATE HYDROGEN SULFIDE + AMMONIUM HYDROXIDE MANGANESE + SILVER SELENIUM + CHLORINE PHOSPHORUS BURNS POTASSIUM CARBONATE + HYDROCHLORIC ACID VANADIUM (V) OXIDE COBALT (III) SULFIDE + CARBONIC ACID POTASSIUM CHLORATE HYDROGEN PERCHLORATE + AMMONIUM HYDROXIDE CARBON + IODINE COPPER (II) HYDROGEN SULFATE + HYDROGEN HYDROXIDE C21H44 BURNS MANGANESE (II) CARBONATE + SULFURIC ACID OXYGEN + FLUORINE POTASSIUM THIOCYANATE + IRON (III) CHLORIDE TUNGSTEN (VI) OXIDE + SODIUM LEAD + COBALT (III) NITRITE BARIUM HYDROXIDE + HYDROGEN SULFITE CALCIUM + WATER CARBON DIOXIDE + NICKEL (III) OXIDE (ONE PRODUCT) SULFUR BURNS PHOSPHORUS (V) BROMIDE MAGNESIUM PHOSPHATE + COPPER (I) ACETATE AMMONIUM CARBONATE (THREE PRODUCTS) ZINC + CARBONIC ACID CARBON DIOXIDE + WATER (ONE PRODUCT) BORON + FLUORINE MERCURY (I) NITRATE + STRONTIUM IODIDE MANGANESE OXIDE SULFURIC ACID + CESIUM SULFURIC ACID (TWO PRODUCTS) MANGANESE (II) HYDROXIDE + HYDROCHLORIC ACID SELENIUM + SULFUR 156 EXERCISE 22 CHEMICAL EQUATIONS AND MOLES DIRECTIONS: On a separate sheet of paper, do the following: (Note: Equations are unbalanced.) 1. How many moles of cadmium nitrate will react with 6.678 moles of sulfuric acid? Cd(NO3)2 + H2SO4 CdSO4 + HNO3 (6.678 mole) 2. How many atoms of iron will react with an excess of sulfur to form 2.0 moles of iron (II) product? Fe + S FeS (1.2 x 1024 atoms) 3. How many moles of silver nitrate are needed to react with 0.34654 mole of sodium (0.34654 mole) bromide? AgNO3 + NaBr AgBr + NaNO3 4. Aluminum reacts with oxygen to form the oxide. If we want to produce 45,670 mol of the oxide how much of each is needed? (91.340 mol Al) (68.505 mol O2) Al + O2 Al2O3 5. How many moles of mercury are formed from the decomposition of 6 908 moles of mercury (II) oxide? HgO Hg + O2 (6 908 mol) 6. How many liters of nitrogen at STP are needed to react with 9.000 7g of aluminum? (3.73 L) Al + N2 AlN 7. How many grams of sodium hydroxide are needed to react with 2.22kg of barium bromide? NaOH + BaBr2 NaBr + Ba(OH)2 (598g) 8. How many molecules of hydrogen are needed to react with 55.55g iron (III) oxide? H2 + (6.29x1023 mcls) Fe2O3 H2O + Fe 9. How many grams of zinc are needed to react with .000 006 785 grams of silver sulfate? (1.42x10-6g) Zn + Ag2SO4 ZnSO4 + Ag 10. How many grams of silicon are needed to react with 23.90g of sulfur? Si + S SiS2 (10.5g) 11. How much silver phosphate is produced if 10.0g of silver acetate is reacted with an excess of sodium phosphate? AgC2H3O2 + Na3PO4 Ag3PO4 + NaC2H3O2 (4.05g) 12. What mass of sodium hydroxide is needed to completely react with 25.0g of sulfuric (20.4g) acid? NaOH + H2SO4 Na2SO4 + H2O 13. Sulfur (IV) oxide reacts with water to form sulfurous acid, how much sulfurous acid is produced from 2.75g of sulfur (IV) oxide and an excess of water? SO2 + H2O H2SO3 (3.52g) 14. 739g of bromine reacts with potassium iodide in excess. How much of each product is formed in the reaction? Br2 + KI KBr + I2 (1 170g I2) (1.10x103g KBr) 15. 27.0g of silver oxide decomposes into silver and oxygen. How many liters of oxygen are (1.30L O2) produced? Ag2O Ag + O2 157 EXERCISE 23 INFORMATION There are two factors to consider when expressing the energy change of a substance: 1. The temperature change of the substance must be considered. 2. The amount of substance that changes temperature must be considered. Both these factors are combined in a unit of energy called the calorie. A calorie is a unit of energy. It is defined as the amount of energy it takes to heat one gram of water 1 °C. This definition can be expressed by three equations: 1. No. of calories = (Change in temperature) x (no. of grams of water) 2. t = _____no. of calories________ no. of grams of water 3. No. of grams of water = no. of calories t Typical problems involving the calorie: 1. 400 grams of water are heated from 250C to 400C. required? Solution: Use equation 1. How many calories are The number of grams of water = 400 The change in temperature (t) = 400 - 250 = 150 Thus: Number of calories = (400) x (15) = 6,000 calories 2. 200 calories of energy are added to 50 grams of water. What temperature change occurs in the water? Solution: Use equation 2. The number of grams of water = 50 The number of calories = 200 Thus: t = 200/50 = 4oC 3. 1000 calories of energy are added to a water sample. The temperature of the water changes from 400C to 650C. How much did the water sample weigh? Solution: Use equation 3. The number of calories Change in temperature t Weight of water sample = 1000/25 = 40 g 158 = 1000 = 250 EXERCISE 23 INFORMATION, CONTINUED In the experiment you made an important discovery: A substance doesn't have to change temperature in order for it to gain or lose heat. You found that when moth ball crystals froze and changed from a liquid to a solid, there was no change in temperature. In fact, you could use the procedure followed in this experiment to determine the freezing point of many substances. Phase change energy is the name given to the energy necessary to change a liquid to a solid or a solid to a liquid. A molecule does not change its molecular formula when it undergoes a phase change. If a substance changes from a solid to a liquid, heat must be added. This is called endothermic change. If a substance changes from a liquid to a solid, heat must be released. This is called exothermic change. How would you report this exothermic or endothermic energy change? In other words, what units would you use to report this energy? There are two basic ways to express the energy change of a substance: 1. One is calories per gram. 2. The other is calories per mole. The two basic formulas are: calories per gram = no. of calories_____ no. of grams of substance calories per mole = (calories per gram) x (molecular weight) An example: When 40 grams of a substance are frozen, this causes 220 grams of water to rise in temperature from 250C to 330C. Find the number of calories per mole if the molecular weight is 120 g. Solution: To determine the calories per mole we must find the following: 1. No. of calories 2. No. of calories per gram 3. No. of calories per mole Step 1. No. of calories = (220) x (330-250) = (220) x (8) = 1760 calories Step 2. No. of calories per gram = 1760 calories 40 grams Step 3. No. of calories per mole = 44 calories gram = (44 cal. per gram) x (120 g/mol) = 5,280 cal/mol 159 EXERCISE 23-CALORIES Show your work on each problem. 1. How many calories would be required to heat 220 grams of water from 200C to 350C? 2. 30 grams of a substance are burned, and the heat produced is used to warm 2000 grams of water from 220C to 30C. A. How many calories are released by the substance? B. How many calories would be released if only 1.0 g of the substance were burned? C. If this substance has a molecular weight of 60 grams per mole, how many calories are released by one mole? 3. Two grams of candle are burned. Each gram of candle releases 7000 cal when burned. If all the heat from the burning candle goes into a sample of water, what is the weight of the water? The water temperature increases 250C. 4. If 400 grams of water at 250C are heated by adding 4000 calories to it, A. What is t? B. What is the final temperature of the water? 5. Find the laboratory data in your notebook on Experiment 11.1. Use the data to answer these questions: A. Determine the temperature change of a hot-water sample. B. Determine the weight of the hot-water sample. C. How many calories of energy were lost by the hot water? D. Determine the temperature change of the cold-water sample. E. Determine the weight of the cold-water sample. F. How many calories of energy were gained by the cold-water sample? G. Should there be a similarity in the answers for Questions C and F? Explain. H. What might account for the difference in the answers for Questions C and F? 160 EXERCISE 24 CALORIMETRY DIRECTIONS: On a separate sheet of paper, complete the following. 1.) How many calories are needed to raise 145 g of water from 40.0OC to 65.0OC? (3620 cal) 2.) What temperature change is produced when 898 cal. are absorbed by 134 g of water? (6.70OC) 3.) How many grams of water can be heated from 34.5OC to 89.6OC using 3566 calories? (64.7g) 4.) How many kilocalories are required to change the temperature of 321 g of water 45.6OC? (14.6 kcal) 5.) How many grams of water can be heated 65.0OC by 4.56 kcal? (70.2g) 6.) What is the final temperature after 90.0 calories are absorbed by 11.0 g of water at 25.0OC? (33.2OC) 7.) How many calories are released when 344 g of water-cools from 70.0OC to 30.0OC? (13 800 cal) 8.) What temperature change is produced when 666 g of water gives off 9.88 kcal? (14.8OC) 9.) What is the final temperature when 245 calories are given off by 45.5 g of water at 46.0OC? (40.6OC) 10.) How many calories are needed if the temperature of 34.6 g of water changes 3.67OC? (127 cal) 161 EXERCISE 25 ENTHALPY DIRECTIONS: On a separate sheet of paper do the following: 1.) Calculate the change in enthalpy for 9279 g of oxygen combining with hydrogen. Product is a liquid. ( -166 000 kJ) 2.) Calculate the change in enthalpy for the burning of 3 940 kg of methane CH4. All substances are gases. (-1.98X108kJ) 3.) Calculate the change in enthalpy for the formation of 0.000 297 5g of SO3 from sulfur and oxygen. (-0.00164 kJ) 4.) Calculate the change in enthalpy for the formation of 28.54 g of sliver chloride from sliver and chlorine. (-25.2 kJ) 5.) Calculate the change in enthalpy for the formation of CuI from 354 g of copper and an excess of iodine. (-378 kJ) Use the following heats of formation to solve the above H2O (l) = - 286 kJ/mol CO2 = - 393.5 kJ/mol H2O(g) = - 242 kJ/mol CH4 = -74.8 kJ/mol SO3 = - 441 kJ/mol AgCl = -127 kJ/mol CuI = - 67.8 kJ/mol 162 EXERCISE 26 GAS LAWS DIRECTIONS: On a separate sheet of paper, answer the following questions. 1.) 842 L of a gas is at 132oC, what is its volume at 152oC? (884 L) 2.) If 496 ml of a gas is at 3.21 atm, what is its pressure at 694 mL? (2.29 atm.) 3.) If 62.7 mL of gas is at 307 K, what is its temp at 37.8 mL? (185 K) 4.) If a gas is at 444 mm Hg and 0.277 L, what is its volume at 684 mm Hg? (0.180 L) 5.) If 697 L of a gas is at 2.11 atm and 419 K, what is its temp. at 3.11 atm. and 546 L? (484K) 6.) If 0.0677 L of a gas is at 432 PSI and 267 K, what is the pressure of the gas at 342 K and .148 L? (253 PSI) 7.) What is the volume of a gas at STP if 4 944 ft3 is at 760 mm and 86.0OC? (3 760 ft3) 8.) 89 246 cm3 of a gas is at 0.0178 atm, what is the pressure of the gas at 422 cm3? (3.76 atm.) 9.) A sample of gas occupies 75.0 ml at 725 mm Hg and 18OC. Calculate its volume at 801 mm Hg and 298 K? (69.5 mL) 10.) A sample of gas occupies 75.0 mL at 1.15 atm. and 25OC. At what temp. will it occupy 4.00 L at 1.00 atm?: (13 800 K) 11.) A sample of gas occupies 8.25 L at STP. What will its volume become at 735 mm Hg and 20O0C? (9.16 L) 12.) A sample of gas occupies 0.575 L at 1.50 atm. and 125OC. What pressure will it have if its volume changes to 0.300 L and its temperature to 200OC? (3.42 atm.) 13. A sample of gas occupies 2.30 L at 825 mmHg and 70oC. What is its volume at STP? (1.99 L) 163 EXERCISE 27 MORE GAS LAW CALCULATIONS 1. Calculate the volume of 1.00 mole of CO2 at STP. 2. At sea level the average lung capacity for humans is about 4.0 liters. At 37o C (body temperature) how many grams of oxygen gas could your lungs hold? 3. If 40.0 grams of methane, CH4, is confined to 2500 mL at 200oC, what pressure does it exert? 4. Calculate the relative rates of diffusion of methane and sulfur dioxide. 5. What volume tank would be needed to contain 3 900 g of helium at 3.45 atm. and 45.0 o C? 6. What temp. would be on a 34.0 L container with 856 g of oxygen at 45 atm? 7. What is the diffusion rate of carbon dioxide to Radon? 8. What pressure is exerted on a 35 678 mL container, if it has 3 899 g of chlorine at 34.9 o C? 9. What is the diffusion rate of Argon to Krypton? 10. If 5.03 x 1024 atoms of neon are in a 21.9 L tank at 765K, What is the pressure? 11. If gas A has a pressure of 235mm and gas B has a pressure of 567 mm and gas C has a pressure of 242 mm, what is the total pressure on the tank? 164 EXERCISE 28 SOLUTION CHEMISTRY DIRECTIONS: On a separate sheet of paper, do the following: A. Calculate the molarity of the following solution. 1. 316g of MgBr2 in 859 mL of solution. (2.00M) 2. 4.67g of Ca(C2H3O2)2 in 465 mL of solution. ( .0635M) 3. 4.009g of LiClO3 in 297 mL of solution. ( .149M) B. Calculate the molality of the following solutions: 1. 199g of NiBr2 in 599g of water. (1.52m) 2. 67.0g of RbI in 2 007g of water. ( .157m) 3. .00987g of AgClO2 in .333kg of water. ( .000169m) C. Calculate the mass of solute needed to make the following solutions: 1. 1 206mL of .0464M Sc(NO3)3 (12.9g) 2. 134mL of 1.345M VBr3 (52.4g) 3. 5.166L of .397M IrCl4 (685g) D. Calculate the volume of solution that can be made with the given information. 1. 90.0g of NH4Cl to make 0.750M solution. (2.24L) 2. 356g of NaNO3 to make 0.00467M solution. (897L) 3. .0493g of AgNO3 to make 0.00000111M solution. (261L) 165 EXERCISE 29 RATE OF REACTION DIRECTIONS: On a separate sheet of paper, do the following: 1. Write the rate equation for the following: a. 2Q + 3Z -- 4Y b. H2 + Cl2 -- 2HCl c. 2 C18H38 + 55 O2 -- 36 CO2 + 38 H2O 2. List the five factors that affect the rate of a reaction and explain in terms of kinetic theory. 3. Differentiate between a heterogeneous and a homogeneous catalyst. 4. How does a catalyst affect reaction rate? 5. Diagram and label a potential energy curve for a catalyzed and uncatalyzed exothermic reaction. 166 EXERCISE 30 REACTION RATES 1. What is the rate expression for the following reaction? What is the value of the rate constant for the reaction? H2O2 + 2HI 2H2O + I2 Trial 1 2 3 [H2O2] 0.1M 0.1M 0.2 M [HI] 0.1M 0.2M 0.1 M Rate ((mol/L)/s) 0.0076 0.0152 0.0152 2. Assume that NO(g) and H2(g) react according to the rate expression: rate = k[NO]2[H2] . How does the rate change if: a. the concentration of H2 is doubled? b. The volume of the enclosing vessel is suddenly halved? c. The temperature is decreased? 3. For the reaction: H2(g) + I2(g) 2HI(g) , the following data were obtained: Experime nt Initial [H2] Initial [I2] 1 2 3 1.0M 1.0M 2.0M 1.0M 2.0M 2.0M Initial rate of formation of HI ((mol/L)/s) 0.20 0.40 0.80 a. Write the rate expression for this reaction. b. What would be the initial rate of formation of the hydrogen iodide if the initial concentrations of the hydrogen gas and the iodine gas were each 0.50M? 167 EXERCISE 30 CONTINUED 4. The reaction CH3COCH3 + I2 CH3COCH2I + HI is run under carefully controlled conditions in the presence of an excess of acid. Write the rate expression for the reaction using the following data: Initial concentration of CH3COCH3 0.100M 0.0500M 0.0500M Initial concentration of Initial rate I2 ((mol/L)/s) 1.16x10-7 5.79x10-8 5.78x10-8 0.100M 0.100M 0.500M 5. NO will react with bromine according to the reaction 2NO + Br2 2NOBr. Experimentally, the following data were obtained: Trial Initial [NO] Initial [Br2] 1 2 3 1.0M 1.0 M 2.0 M 1.0 M 2.0 M 2.0 M Initial rate of formation of NOBr ((mol/L)/s) 0.80 1.60 6.40 a. Write the rate expression for this reaction. b. Calculate the rate constant, k. c. What would be the initial rate of formation of NOBr if the initial concentrations of NO and Br2 were 0.60M and 0.25M, respectively? d. If the initial concentration of NO were 1.4M, what initial concentration of Br2 would produce NOBr at an initial rate of 0.75(mol/L)/s? e. Note that concentrations and rates are all specified “initial”. How would rate differ if measured later in the reaction? Why would it differ? 168 EXERCISE 31 LeCHATELIER For these questions predict what changes will take place resulting from the changes in concentration, pressure, or heat. Le Chatelier's Principle will be helpful in making these predictions. Le Chatelier's Principle: If a system at equilibrium is changed, the system will partially counteract that change. Predict the result of these changes on the following equilibrium system(s): 1. Cu2+ (light blue) + 4NH3 (colorless) Cu(NH3)42+ (dark blue) a. Add Cu2+ solution b. Add NH3 solution c. Remove NH3 d. Add Cu(NH3)42+ 2. 2CO + 02 2C02 + heat a. Add CO2 b. Remove O2 c. Add CO d. Add heat e. Remove heat 3. N2 + 3H2 2NH3 + heat a. Remove NH3 b. Add H2 c. Remove N2 d. Add heat e. Remove heat (cool) 4. Cr2O72- (orange) + OH- (colorless) 2Cr2O72- (yellow) + H+ (colorless) a. Add H+ b. Remove OHc. Add Ba2+ (reacts with CrO42- but not with Cr2O72-) d. Remove Cr2O72- 169 EXERCISE 32 EQUILIBRIUM DIRECTIONS: On a separate sheet of paper, do the following. 1.) Write the expressions for the equilibrium constant for each of the following reversible reactions. a. H2O + CO b. N2 + 3H2 c. 4NH3 + 5O2 CO2 + H2 2NH3 4NO + 6H2O 2.) Determine the equilibrium constant for the following reaction if the concentration of [N2O4] = .00150 and [NO2] = .571 N2O4 2NO2 3.) For the reaction 2SO2 + O2 2SO3, the concentrations of the sulfur oxides are [SO2] =2.00 and [SO3] = 10.0. What is the concentration of oxygen when the Keq = 800.0 for the reaction? 4.) For the gaseous reactions below, indicate what happens to the equilibrium position (shift to left or right) when the indicated stress on the system occurs. a. remove NH3 b. decrease pressure N2 + 3H2 c. decrease temp. d. add a catalyst CO2 + H2 + heat e. increase [SO2] f. increase temp. 2SO2 + O2 g. increase pressure h. increase [CO] CO2(g) + C(s) + heat i. decrease pressure j. remove N2O4 N2O4(g) + heat k. increase [H2] l. increase pressure H2(g) + Cl2(g) 2HCl(g) + heat m. decrease [O2] n. add a catalyst N2 + O2 + heat o. increase [Cl2] p. decrease pressure 2NH3 + heat CO + H2O 2SO3 + heat 2CO(g) 2NO2(g) 2NO PCl3(s) + Cl2(g) PCl5(s) + heat 170 EXERCISE 33 SOLUBILITY EQUILIBRIA Directions: On a separate sheet of paper, do the following: 1. 2. Write the solubility product equilibrium constant for the following: a. lead (II) sulfate b. silver chloride c. calcium phosphate In a saturated solution of silver oxalate, the concentration of the silver ion is .000 340 M. Calculate the value for Ksp. 3. What is the concentration of the silver ion in a saturated solution of silver chloride? Ksp = 1.80 x 10-10 4. In a saturated solution of silver chromate, the concentration of silver ion is 4.00 x 10-6 M and the chromate ion concentration is .0990 M. What is the value for the Ksp? 5. The Ksp for lead (II) chloride is 1.60 x 10-5. If the concentration of the lead ion is .0178 M, what is the concentration of chloride ion? 6. If the Ksp for AgBr is 4.90 x 10-13, and the concentration of silver and bromide ions are each 9.1 x 10-5 M in a solution of silver bromide, will a precipitate form? 7. The Ksp for HgS is 3x 10-53. Explain what this means in terms of solubility. 171 EXERCISE 34 ACIDS AND BASES 1. Which of the following are acids, bases, both, or neither? a. Fb. H2C204 c. S032d. HC4H406e. H2S f. HC032g. K+ 2. For the pairs of reactants given below: A. Show with a chemical equation what products could form from the equilibrium mixture. B. Label which reactant is the acid and which is the base. C. Predict whether the equilibrium mixture favors the reactants or the products. CH3COOH + HS- NH4+ + F- HSO4- + NH3 HNO2 + S2- NH3 + HPO42- HPO42- + S2- C6H5COO- + HClO H2CO3 + OH- 172 EXERCISE 35 MOLARITY CALCULATIONS 1. How many moles of NaOH are dissolved in 5OO ml of .10 M solution? 2. If 4 moles of substance are dissolved in 500 ml of solution, what in the molarity? 3. If you had .3 moles of a substance and you needed to make a solution with a concentration of .6 M, what volume of solution could you make? 4. If you had 15 g of acetic acid (CH3COOH) and added this amount to enough water to make 2.5 liters of solution, what would be the molarity of the solution? 5. How many liters of a .5 M NaCl solution could you make, using 29 grams of NaCl? 173 EXERCISE 36 ACID-BASE CHEMISTRY Directions: On a separate sheet of paper, do the following: 1. Calculate the Ka for each of the following weak electrolytes. a. A .00100 M acetic acid solution with [H3O+]= 1.27 x 10-4 M. b. (1.85 X 10-5) A .0070 M NH4OH solution with [OH-]= 3.46 x 10-4 M. c. (1.80 x 10-5) A .100 M HCN solution with [H3O+] = 7.85 x 10-6M (6.16 x 10-10) 2. Calculate the Ka for the following weak electrolytes: a. .100M HF b. .0500M HC2H3O2 1.8 % ionized c. .200M HAsO2 3. Calculate the pH and pOH of each of the following: a. .0889M HCl (1.051 and 12.949) b. .333M Ba(OH)2 (13.823 and .177) c. 35.6 g of HNO3 in 355 mL of solution. (-.202 and 14.202) d. 45.6 g of HC2H3O2 in 794 mL of solution. 8.50 % ionized .00550 % ionized (7.90 x 10-4) (1.80 x 10-5) (6.05 x 10-10) Ka= 1.80 x 10-5 (2.38 and 11.62) e. .345M H2X (.161 and 13.839) f. 1.09M C6H5NH2 Kb= 4.20 x 10-10 174 (9.330 and 4.670) exercise 36 continued 4. 5. Complete the following tables: [H+] pH a. 3.4 x 10-8 _____ _____ b. 1.44 x 10-14 _____ _____ c. _____ 5.667 _____ d. _____ _____ 1.229 e. _____ -.087 _____ f. _____ _____ 10.876 pOH What is the pH of a buffer containing .125M HCN and .0117M NaCN? Ka= 4.00 x 10-10 6. (8.369) Using the buffer system in problem #4, what concentration of NaCN is needed to set up a buffer with a pH of 7.000 if we start with 0.1000 HCN? (4.00X10-4) 175 EXERCISE 37 WRITING REDOX REACTIONS Write the equation for the cell reaction that will occur spontaneously in a cell that contains the two electrodes given. 1. Zn2+/Zn and I2/I- 2. Br2/Br- and Pb2+/Pb 3. Ca2+/Ca and Cu2+/Cu 4. Mg2+/Mg and Cu2+/Cu 5. K+/K and Pb2+/Pb 6. Al3+/Al and Br2/Br- Write the two half reactions and the final balanced equation for each of the following, and state whether or not the products or the reactants are favored: 7. Zn + Ag+ 8. F2 + Fe 9. MnO4- + H2O + Fe2+ 10. AgCl + H2O 11. Zn + OH- + Ag2O2 + H+ 12. Mg2+ + Hg 13. Br2 + Ce3+ 14. H3O+ + Pb 15. Cr2O72- + H3O+ + Cd 176 EXERCISE 38 REDOX BALANCING Directions: On a separate sheet of paper, balance the following equations using the bracket method: Cu(NO3)2 + NO + H2O 1. Cu + HNO3 2. Fe(NO3)2 + HNO3 3. Zn + HNO3 4. Sb + H2SO4 5. H2S + H2SO3 6. Ag + HClO3 + HCl 7. HNO3 + H2SO4 + Hg 8. CO + I2O5 9. Zn + HNO3 10. FeSO4 + KMnO4 + H2SO4 Fe2(SO4)3 + MnSO4 + K2SO4 + H2O 11. K2Cr2O7 + KBr + H2SO4 12. Cu + H2SO4 13. Tl2O3 + Zn(NO3)2 + NO2 + H2O Sb2(SO4)3 + SO2 + H2O S + H 2O AgCl + H2O Hg2SO4 + NO + H2O CO2 + I2 HCl Fe(NO3)3 + NO + H2O Zn(NO3)2 + NH4NO3 + H2O Br2 + Cr2 (SO4)3 + K2SO4 + H2O CuSO4 + SO2 + H2O TlCl + Cl2 177 + H2O EXERCISE 38, PAGE 2 14. As2O3 + HNO3 + H2O 15. Pt + HCl + HNO3 H2PtCl6 + NO2 + H2O 16. H2SO3 + I2 + H2O 17. SbCl5 + KI 18. CdS + I2 + HCl 19. AuCl3 + KI 20. Ti2(SO4)3 + Fe2(SO4)3 21. BrO3- + Br- + H+ Br2 + H2O 22. Cl2 + OH- 23. Fe + Cr2O72- + H+ 24. Ce4+ + C2O42- 25. I- + Cr2O72- H2SO4 + HI SbCl3 + KCl + I2 CdCl2 + HI + S AuCl + KCl + I2 H3AsO4 + NO2 Ti(SO4)2 + FeSO4 Cl- + ClO3- + H2O Cr3+ + Fe3+ + H2O Ce3+ + CO2 + H+ Cr3+ + H2O + I2 178 EXERCISE 39 NUCLEAR CHEMISTRY Balancing nuclear equations is somewhat more complicated than balancing chemical equations because atoms do not have to be conserved in a nuclear reaction. There are two rules to observe when balancing nuclear reactions: 1. Charges on the product and reactant sides of the equation must be equal. 2. Total mass on the product and reactant sides must be equal. This rule ignores the small defect in mass. EXAMPLE Conservation of charge 92U + On = 56Ba + 36Kr + 30n Conservation of mass 235 U + 1n = 141Ba + 92Kr + 31n The complete reaction 235 92U + 1 On 141 56Ba + 92 36Kr +3 1 0n For the reactions below fill in the missing number (e = electron and n = neutron). 1. 238 2. 9 92U 4Be + 3. 7 + 4. 239 5. 3Li 92U 226 1 + ??? 1 88Ra 0n 2He 1H 239 ??? = 4 1 0n 2He 93Np ??? 92U + ???Rn + 12 ???C + ___?___ ??? -1e + 4 (electron) 2He Find the symbol for the element by using the atomic number. 179 EXERCISE 40 NUCLEAR EQUATIONS AND HALF-LIFE DIRECTIONS: On a sheet of paper, do the following: 1. a. Complete and balance the following nuclear equations. 14 14 6C 7N 218 b. c. d. e. f. g. h. 84Po + 4 2He 22 11Na k. 86Rn + ___?___ -1e + l. ___?___ m. 140 __?__ 230 88Ra j. 10Ne + ___?___ 0 140 59Pr + ___?___ 222 234 90Th i. 22 226 88Ra ___?___ 58Ce + ___?___ 24 12Mg + 4 2He n. 0 -1e 234 91Pa 234 92U 9 5B 1 1H 8 1 13 3 2He 16 8O + 7 2He 6C + 0n ___?___ + ___?___ + ___?___ 7N + ___?___ 2He + ___?___ 12 1H ___?___ 3 2. The half-life of iodine-131 is 8 days. Calculate the amount remaining after 8 days, 16 days, 24 days and 32 days if we start with 21.8 grams of the radioisotope. + ___?___ 3. What does the half-life indicate about an unstable isotope? 180 181