Name: __________________________________________________________ Period: ________ AICE Unit 1 Worksheet 1 Directions – Complete the following questions. Show all work for calculations. Be prepared to present and discuss your answers during class. Directions – Based on your reading (or in class lectures and discussion), answer the following questions. 1. The types of masses we use in chemistry are all “relative”. What are they relative to? Explain. 2. What is a mole (in chemistry)? What is Avogadro’s constant? On what is Avogadro’s constant based? Directions – Analyze the mass spectra provided and answer the following questions for #3 - #6 (A) State the number of protons and neutrons in the nucleus for each isotope (B) State the percentage abundance for each isotope (C) Calculate the relative atomic mass to one decimal point of the element shown. 3. Boron 5. Chlorine 6. Copper 4. Magnesium Directions – Use the data provided to calculate the average relative atomic mass to one decimal point for the elements presented. 7. Antimony Isotope Mass Abundance 121 Sb 120.903818 57.21% 123 Sb 122.904216 42.79% 8. Gallium Isotope 69 Ga 71 Ga Mass 68.925581 70.924701 Abundance 60.10% 39.90% Isotope 142 Nd 143 Nd 144 Nd 145 Nd 146 Nd 148 Nd 150 Nd Mass 141.907719 142.909810 143.910083 144.912569 145.913112 147.916889 149.920887 Abundance 27.2% 12.2% 23.8% 8.3% 17.2% 5.7% 5.6% Isotope 50 Cr 52 Cr 53 Cr 54 Cr Mass 49.946050 51.940512 52.940654 53.938885 Abundance 4.345% 83.789% 9.501% 2.365% 9. Neodymium 10. Chromium Directions – Use the information provided to determine the empirical or molecular formulas of the compounds as appropriate. 11. An oxide of copper has the following composition by mass: Cu, 0.635 g; O, 0.0800 g. Calculate the empirical formula of the compound. 12. On complete combustion of 0.400 g of a hydrocarbon, 1.257 g of carbon dioxide and 0.514 g of water were produced. What is the empirical formula of the hydrocarbon? If the molecular mass of the hydrocarbon is 84 g/mol, what is its molecular formula? 13. Combustion of 0.255 g of an alcohol produces 0.561 g CO2, and 0.306 g H2O. Determine the empirical formula of the alcohol. 14. A compound has the following composition by mass: Na, 0.6389 g; C, 0.3333 g; O, 0.8889 g. What is the empirical formula of the compound? If the compound has a mass of 135 ± 5 g/mol, what is it molecular formula and name? Directions – Complete the following questions. Show all work for calculations. 15. Ethene gas (C2H4) burns in diatomic oxygen gas to produce carbon dioxide gas and water vapor. You react 34 grams of ethene with 25 grams of oxygen. If 0.50 moles of water are collected, what was the percent yield of the reaction? 16. Solid iron reacts with diatomic oxygen gas to produce iron (III) oxide. If 15 grams of iron reacts with 45 grams of oxygen, what mass of iron (III) oxide is produced? 17. Methane gas (CH4) burns in diatomic oxygen gas to produce carbon monoxide gas and water vapor. 0.25 moles of methane are reacted in the presence of 0.40 moles of oxygen. If 7.5 grams of water are collected, what is the percent yield? 18. Hydrogen peroxide (H2O2) decomposes to form water and diatomic oxygen gas. If 28 grams of hydrogen peroxide are decomposed, how many moles of water are produced? 19. Hexene (C6H12) burns in diatomic oxygen to produce carbon monoxide and water. If 95 grams of carbon monoxide are produced, what mass of hexene was ignited? 20. Copper (II) oxide reacts with sulfuric acid to produce copper (II) sulfate and water. If you react 0.72 moles of copper (II) oxide with 71 mL of 12 M sulfuric acid, what mass of sulfuric acid is consumed by the reaction? 21. 45.0 g of carbon disulfide is combusted in excess oxygen to form carbon dioxide gas and sulfur dioxide. The resulting gases are bubbled through 2.50 L of a 3.00 M solution of sodium hydroxide. 0.750 L of the final sodium hydroxide solution are reacted with an excess of ammonium chloride solution. What mass of ammonia gas is produced from the final reaction? 22. 150 g of sodium bicarbonate is thermally decomposed to produce solid sodium carbonate, gaseous carbon dioxide, and water vapor. The produced gases are collected and dried by passing them through a drying column of anhydrous calcium chloride. Assuming all of the water is removed, but what mass would the drying column increase? 23. Sodium azide decomposes non-explosively on heating to release nitrogen gas. This provides a convenient method of obtaining pure nitrogen in the laboratory: 2NaN3(s)  2Na(l) + 3N2(g) 3 (A) A student prepared 1.80 dm of pure nitrogen in the lab by this method. This gas volume was measured at room temperature and pressure (r.t.p). How many moles of nitrogen did the student prepare (assume 1 mole of gas molecules occupies 24.0 dm3 at r.t.p.). What mass of sodium azide did the student heat? (B) After cooling, the student obtained 1.15 g of solid sodium. She then carefully reacted this sodium with water to form 25.0 cm3 of aqueous sodium hydroxide. Calculate the concentration, in mol dm-3, of the aqueous sodium hydroxide. 24. A 3.445 g sample of a mixture was analyzed for barium by adding a small excess of sulfuric acid to and aqueous solution of the sample. The resultant reaction produced a precipitate of barium sulfate, which was collected by filtration, washed, dried, and weighed. If 0.2815g of barium sulfate was obtained, what was the mass percentage of barium in the sample? 25. What mass of NaOH is required to precipitate the Cd+2 ions from 35.0 mL of 0.500 M Cd(NO3)2 solution? 26. A solution is made by mixing 12.0 g of NaOH and 75.0 mL of 0.200 M HNO3. Calculate the final concentration of each ion remaining in solution. Is the resulting solution acidic or basic? 27. A 1.248 g sample of limestone rock is pulverized and then treated with 30.00 mL of 1.035 M HCl solution. The excess acid then require 11.56 mL of 1.010 M NaOH for neutralization. Calculate the percent by mass of calcium carbonate in the rock, assuming that it is the substance reacting with the HCl solution.