CH. 10 Gases AK - Answer Key for practice worksheet PDF

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Chapter 10: Gases (Answer Key) 1. Name and briefly describe the following gas laws. Be sure to include a mathematical expression that represents each law. a. Boyle’s Law: Pressure and volume (inversely proportional), as pressure increases, volume decreases. P 1V1 = P2V2 b. Charles’ Law: Volume and temperature (directly proportional), as temperature increases, volume increases. V 1/T1 = V2/T2 c. Avogadro’s Law: Volume and moles (directly proportional), the volume of gas maintained at constant temperature and pressure is directly proportional to the number of moles of gas. V 1/n1 = V2/n2

2. Write out the Ideal Gas Law and define its variables. Also define STP PV=nRT P = pressure (atm) V = volume (L) n = # of moles (mol) R = gas constant (0.08206 L-atm/K-mol) T= temperature (K) STP = Standard temperature and pressure (0°C and 1.00 atm) At STP, one mole of an ideal gas occupies 22.41 L

3. Based on the following image, answer the following questions. This figure shows a contained that is sealed at the top by a moveable piston. The container contains an ideal gas at 1.00 atm, 20.0 C, and 1.00 L.

a. What will the change in pressure be if the piston is move to the 2.5 L mark? Assume the temperature will remain constant. P 1V 1 = P 2V 2 1 L  1 atm = 2.5 L  x atm ; x=0.400 atm; ∆P=0.600 atm

b. If the ideal gas contained in the container has a mass of 1.66 grams. What is the identity of the gas? → n=

PV= nRT

(

grams(m ) Molar Mass (MM )

)(

)

( 1 atm ) (1 L )= 1.66 g ∙ 0.08206 L∙ atm ∙ 293 K K ∙ mol MM 0.0415 mol= MM =

1.66 g MM

1.66 g =39.9 g/mol = [Ar] 0.0415 mol

4. Calculate the volume of CO2 at STP produced from the decomposition of 152 g of CaCO3 by the reaction: CaCO3(s) à CaO(s) + CO2(g) 152 g ∙

1mol CaC O 3 100.09 g CaC O 3

=1.52 mol CaC O3

1.52 mol CaC O3 =1.52mol C O 2 1.52 mol C O 2 ∙

22.42 L =34.1 LC O 2 1 mol C O 2

5. A gas is contained in a cylinder with a volume of 5.0 x 102 mL at a temperature of 30.0 oC and pressure of 710.0 torr. The gas is then compressed to a volume of 25 mL and the temperature is raised to 820 oC. What the change in pressure? V1 = 500mL / 1000 = 0.5L T1 = 30.0 °C +273 = 303K P1 = 710.0 torr/ 760 torr = 0.9342 atm n=

V2 = 25mL/1000 = .025L T2 =820°C + 273 = 1093K P2 = ?

(0.5 L ) ( 0.9342 atm ) 0.4671 =0.01878 mol = 24.86 L ∙atm ( 303 K ) 0.08206 K ∙ mol

(

)

L ∙ atm 0.08206 ( 0.01878 mol) ( 1093 K ) ( K ∙mol ) P= =67.39 atm 2

( 0.025 L)

67.39− 0.9342=66.4 ≈ 66 atm

6. Urea (H2NCONH2) is produced commercially from the reaction of ammonia and carbon dioxide. 2NH3(g) + CO2(g) à H2NCONH2(s) + H2O(g) 500.0 L of ammonia gas at 223 oC and 90.0 atm reacts with 600.0 L of carbon dioxide at 223 oC and 45 atm. What mass of urea is produced by this reaction?

7. A solid piece of carbon dioxide (dry ice) with a mass of 7.8 g is placed in a 4.0 L container at 27 oC that is already filled with air at 740 torr. After the carbon dioxide vaporizes, what is its partial pressure and what is the total pressure? • • •

Calculate moles of CO2 Calculate partial pressures Add partial pressures PV=nRT

8. The mole fraction of nitrogen in the air is 0.7808. Calculate the partial pressure of nitrogen in air at 760 torr. Mole Fraction = P1/ Ptotal

9. A sample of solid potassium chlorate (KClO3) was heated in a test tube and decomposed by the following reaction: 2KClO3(s) à 2KCl(s) + 3O2(g) The oxygen produced was collected by displacement of water at 22oC at a total pressure of 754 torr. The total volume of gas collected was 0.650 L and the vapor pressure of water at 22oC is 21 torr. Calculate the partial pressure of O2 in the gas collected and the mass of KClO3 in the sample that was decomposed. • • •

Find partial pressure of O2 (Ptotal = P1 + P2) Find moles of O2 (PV = nRT) Find moles, grams of KClO3 (mole ratio)...