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Practice Problems

Practice Problems. Try to do the following problems in your head. 9.5 Suppose you have 0.10 mole of uranium, U, atoms. (a) How many uranium atoms do you have? (b) What is the mass in grams of this much uranium?. Practice Problems. Try to do the following problems in your head.

regina-clements
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Practice Problems

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  1. Practice Problems Try to do the following problems in your head. 9.5 Suppose you have 0.10 mole of uranium, U, atoms. (a) How many uranium atoms do you have? (b) What is the mass in grams of this much uranium?

  2. Practice Problems Try to do the following problems in your head. 9.6 Suppose you have 120.11 g of carbon atoms. (a) How many moles of carbon atoms do you have? (b) How many carbon atoms do you have?

  3. Practice Problems 9.11 Suppose you have 100.0 g of H2O. (a) How many moles of H2O molecules do you have? (b) How many H2O molecules do you have? Answer: (a) or (b)

  4. Practice Problems 9.13 Suppose you have 5.000 × 1024 molecules of methane, CH4. (a) How many moles of methane do you have? (b) How many grams of methane do you have?

  5. Practice Problems 9.14 (a) Express the balanced chemical equation CH4 + 2 O22 H2O + CO2 in words, using the word mole(s) wherever appropriate. (b) To produce 1 mole of CO2 from this reaction, how many grams of CH4 and O2 must you combine? (c) What is the theoretical yield in grams of H2O for this reaction? (d) If you recover 30.0 g of H2O, what is the percent yield? Answer: (a) One mole of methane molecules and 2 moles of oxygen molecules react to give 2 moles of water molecules and 1 mole of carbon dioxide molecules. (b) The molar mass of methane is (12.011 g/mol C) + (4 × 1.0079 g/mol H) = 16.043 g/mol CH4. This is the mass of 1 mole of methane, which is what the reaction calls for. The molar mass of oxygen, O2, is (2 × 15.999 g/mol O) = 31.998 g/mol. This is the mass of 1 mole of O2. Because the reaction calls for 2 moles, we multiply by 2 to get 63.996 g of O2 needed. (c) The most we can hope to form is 2 moles of water. The molar mass of water is (2 × 1.0079g/mol H) + (15.999 g/mol O) = 18.015 g/mol H2O. This is the mass of 1 mole of water. So the theoretical yield is just twice this, which is 36.030 g of H2O. (d) % yield = (30.0 g/36.030 g) × 100 = 83.3%

  6. Workpatch 9.3 This concept is so important that you need to practice it now. (a) Express the balanced equation in words, using the word moles for every reactant and product. (b) Write every conversion factor you can think of from the balanced chemical equation. (c) Write all the conversion factors from the formula Al2O3.

  7. Practice Problems 9.21 How many O atoms are there in 10.7 g of oxygen, O2, molecules? Answer: Don’t be afraid to string out conversion factors. They may come from a balanced equation, a molecular formula, or your knowledge of the mole, depending on the question you are trying to answer.

  8. Practice Problems 9.22 How many aluminum atoms are there in 10.0 g of aluminum oxide, Al2O3?

  9. Practice Problems 9.23 How many molecules of water are there in 10.0 g of water?

  10. Practice Problems In Practice Problems 9.24–9.26, use the method of dividing moles by coefficients to determine whether the reaction is being run in a balanced or limiting fashion. If it is being run in a limiting fashion and you are asked for theoretical yield, remember to use only the limiting reactant in your calculation. 9.26 Chemical treatment of zinc sulfide, ZnS, with oxygen, O2, gives zinc oxide, ZnO, and sulfur dioxide gas, SO2. (a) Write a balanced equation for the reaction. (b) If 10.0 g of ZnS is combined with 10.0 g of O2, what is the theoretical yield of each product, in grams? (c) How much of the excess reactant is left over, in grams? (d) Suppose only 7.50 g of ZnO is recovered. What is the percent yield of ZnO?

  11. Practice Problems 9.27 A compound is known to contain carbon and hydrogen. It might also contain oxygen. A sample of the compound is burned. The results of the combustion analysis are 74.9% C and 21.5% H. (a) What is the chemical formula for this compound? (b) Write the balanced combustion reaction (reaction with O2) for this compound. Answer (not worked out on purpose—you do it): (a) CH4(b) CH4 + 2 O22 H2O + CO2

  12. Practice Problems 9.28 A compound is found to contain 52.4% carbon and 13.1% hydrogen. It might also contain oxygen. (a) What is the chemical formula for this compound? (b) Write the balanced combustion reaction (reaction with O2) for this compound.

  13. Practice Problems 9.29 A compound is known to contain carbon and hydrogen and might also contain oxygen. A sample is burned yielding 54.6 % C and 9.16 % H. (a) What is the empirical formula of the compound? (b) The molar mass of the compound is 132.159 g/mol. What is the molecular formula? (c) Write the balanced combustion reaction (reaction with O2) for the compound. Answer (not worked out on purpose—you do it): (a) C2H4O (b) C6H12O3 (c) To balance the equation, look at C and H first. Then balance the elemental substance (O2) last: We make O2 provide 15 O atoms by multiplying it by 7.5 (that is, ): This is a perfectly correct balanced equation, but if you prefer the balancing coefficients to be whole numbers, you can multiply all of them by some number that makes them whole (in this case, 2): 2 C6H12O3 + 15 O212 CO2 + 12 H2O

  14. Practice Problems 9.30 A compound is known to contain carbon and hydrogen and might also contain oxygen. A sample is burned yielding 85.62 % C and 14.37 % H. (Be careful here. Remember, there is always a little inaccuracy associated with measured numbers). (a) What is the empirical formula for the compound? (b) The molar mass is 28.054 g/mol. What is the molecular formula? (c) Write the balanced combustion reaction.

  15. Practice Problems 9.31 What is the mass percent of each element in hydrogen peroxide, H2O2? Answer: 5.93% H, 94.07% O

  16. Practice Problems 9.32 What is the mass percent of each element in trinitrotoluene (TNT), C7H5N3O6?

  17. Practice Problems 9.33 A compound is found to have the following elemental mass percents: Cl = 89.09%, C = 10.06%, H = 0.84%. The molar mass of the compound is 119.378 g/mol. What are the empirical and molecular formulas?

  18. Practice Problems 9.34 A compound is known to contain C and H, and might also contain O. It is analyzed for C and H only, yielding the mass percents C = 54.53% and H = 9.15%. The molar mass of the compound is 88.106 g/mol. What are the empirical and molecular formulas?

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