Unit 1 Gases

1.5 Exercises

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Section 1.5 Exercises

  1. Calculate the density of Freon 12, CF2Cl2, at 30.0 °C and 0.954 atm.
  2. A cylinder of O2(g) used in breathing by emphysema patients has a volume of 3.00 L at a pressure of 10.0 atm. If the temperature of the cylinder is 28.0 °C, what mass of oxygen is in the cylinder?
  3. What is the molar mass of a gas if 0.281 g of the gas occupies a volume of 125 mL at a temperature 126 °C and a pressure of 777 torr?
  4. The density of a certain gaseous fluoride of phosphorus is 3.93 g/L at STP. Calculate the molar mass of this fluoride and determine its molecular formula.
  5. A 36.0–L cylinder of a gas used for calibration of blood gas analyzers in medical laboratories contains 350 g CO2, 805 g O2, and 4,880 g N2. At 25 degrees C, what is the pressure in the cylinder in atmospheres?
  6. A sample of gas isolated from unrefined petroleum contains 90.0% CH4, 8.9% C2H6, and 1.1% C3H8 at a total pressure of 307.2 kPa. What is the partial pressure of each component of this gas? (The percentages given indicate the percent of the total pressure that is due to each component.)
  7. Most mixtures of hydrogen gas with oxygen gas are explosive. However, a mixture that contains less than 3.0 % O2 is not. If enough O2 is added to a cylinder of H2 at 33.2 atm to bring the total pressure to 34.5 atm, is the mixture explosive?
  8. A sample of carbon monoxide was collected over water at a total pressure of 756 torr and a temperature of 18 °C. What is the pressure of the carbon monoxide? (See Table 1 for the vapor pressure of water.)
  9. Joseph Priestley first prepared pure oxygen by heating mercuric oxide, HgO:
    [latex]2 \text{HgO}(s) \longrightarrow 2\text{Hg}(l) + \text{O}_2(g)[/latex]
    (a) Outline the steps necessary to answer the following question: What volume of O2 at 23 °C and 0.975 atm is produced by the decomposition of 5.36 g of HgO?
    (b) Answer the question.
  10. The chlorofluorocarbon CCl2F2 can be recycled into a different compound by reaction with hydrogen to produce CH2F2(g), a compound useful in chemical manufacturing:
    [latex]\text{CCl}_2 \text{F}_2(g) + 4 \text{H}_2(g) \longrightarrow \text{CH}_2 \text{F}_2(g) + 2\text{HCl}(g)[/latex]
    (a) Outline the steps necessary to answer the following question: What volume of hydrogen at 225 atm and 35.5 °C would be required to react with 1 ton (1.000 × 103 kg) of CCl2F2?

    (b) Answer the question.

  11. Lime, CaO, is produced by heating calcium carbonate, CaCO3; carbon dioxide is the other product.

    (a) Outline the steps necessary to answer the following question: What volume of carbon dioxide at 875°C and 0.966 atm is produced by the decomposition of 1 ton (1.000 × 103 kg) of calcium carbonate?
    (b) Answer the question.

  12. Calculate the volume of oxygen required to burn 12.00 L of ethane gas, C2H6, to produce carbon dioxide and water, if the volumes of C2H6 and O2 are measured under the same conditions of temperature and pressure.
  13. Consider the following questions:

    (a) What is the total volume of the CO2(g) and H2O(g) at 600 °C and 0.888 atm produced by the combustion of 1.00 L of C2H6(g) measured at STP?
    (b) What is the partial pressure of H2O in the product gases?

  14. What volume of oxygen at 423.0 K and a pressure of 127.4 kPa is produced by the decomposition of 129.7 g of BaO2 to BaO and O2?
  15. Ethanol, C2H5OH, is produced industrially from ethylene, C2H4, by the following sequence of reactions:
    [latex]3 \text{C}_2 \text{H}_4 + 2\text{H}_2 \text{SO}_4 \longrightarrow \text{C}_2 \text{H}_5 \text{HSO}_4 + (\text{C}_2 \text{H}_5)_2 \text{SO}_4[/latex]
    [latex]\text{C}_2 \text{H}_5 \text{HSO}_4 + (\text{C}_2 \text{H}_5)_2 \text{SO}_4 + 3\text{H}_2 \text{O} \longrightarrow 3\text{C}_2 \text{H}_5 \text{OH} + 2\text{H}_2 \text{SO}_4[/latex]
    What volume of ethylene at STP is required to produce 1.000 metric ton (1000 kg) of ethanol if the overall yield of ethanol is 90.1%?
  16. A sample of a compound of xenon and fluorine was confined in a bulb with a pressure of 18 torr. Hydrogen was added to the bulb until the pressure was 72 torr. Passage of an electric spark through the mixture produced Xe and HF. After the HF was removed by reaction with solid KOH, the final pressure of xenon and unreacted hydrogen in the bulb was 36 torr. What is the empirical formula of the xenon fluoride in the original sample? (Note: Xenon fluorides contain only one xenon atom per molecule.)

Solutions

  1. 4.64 g L−1
  2. 38.8 g
  3. 72.0 g mol−1
  4. 88.1 g mol−1; PF3
  5. 141 atm
  6. CH4: 276 kPa; C2H6: 27 kPa; C3H8: 3.4 kPa
  7. Yes
  8. 740 torr
  9. (a) Determine the moles of HgO that decompose; using the chemical equation, determine the moles of O2 produced by decomposition of this amount of HgO; and determine the volume of O2 from the moles of O2, temperature, and pressure.
    (b) 0.308 L
  10. (a) Determine the molar mass of CCl2F2. From the balanced equation, calculate the moles of H2 needed for the complete reaction. From the ideal gas law, convert moles of H2 into volume.
    (b) 3.72 × 103 L
  11. (a) Balance the equation. Determine the grams of CO2 produced and the number of moles. From the ideal gas law, determine the volume of gas.
    (b) 7.43 × 105 L
  12. 42.00 L
  13. (a) 18.0 L
    (b) 0.533 atm
  14. 10.57 L O2
  15. 5.40 × 105 L
  16. [latex]\begin{array}{llclclcl} & \text{XeF}_x(\text{g}) &+& \dfrac{x}{2}\;\text{H}_2(\text{g}) & \longrightarrow & \text{Xe}(\text{g}) & + &x\;\text{HF}(\text{g}) \\ \text{Initial Pressure:} &18 \text{ torr} &&54 \text{ torr} &&0 \text{ torr} &&0 \text{ torr} \\ \text{New Pressure:} &0 \text{ torr} &&18 \text{ torr} &&18 \text{ torr} &&(\text{HF removed with KOH}) \end{array}[/latex]The partial pressure of H2 consumed is, by stoichiometry, x/2 of the partial pressure of XeFx.
    [latex]P_{\text{H}_2} \text{ consumed}= \left(18\text{ torr XeF}_x \right)\left(\frac{x/2\; \text{mol H}_2}{1\;\text{mol XeF}_x}\right)[/latex]
    The new total pressure is 36 torr of which 18 torr is Xe. The remainder H2 partial pressure is 18 torr equalling the initial H2 partial pressure subtract the partial pressure of H2 consumed.[latex]54\;\text{torr} - (18\;\text{torr XeF}_x)\left(\frac{x/2\; \text {mol H}_2}{1\;\text{mol XeF}_x}\right)= 18\;\text{torr} \\ x = 4[/latex]The formula is XeF4.

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