A cylinder of molecular hydrogen is separated by a metal plate from a cylinder of molecular oxygen. The plate is removed, a spark starts the reaction, and gaseous water is formed. Assume the product has had time to cool to ambient temperatures of 400K. 24 g of molecular hydrogen tracts with 76 g of molecular oxygen. The volume of each cylinder is 1.00 L. (Use the ideal gas law, not Van der Waal’s equation of state.)
1.) Write the balanced equation.
2 H2(g) + O2 (g) à 2 H2O (g)
2.) How many moles of molecular hydrogen and molecular oxygen are there originally? What is the limiting reagent?
12 mol H2, 2.375 mol O2
O2 is the limiting reagent.
3.) What is the initial pressure in each cylinder in atm (before mixing the reagents)? Under these estimated pressures, would you expect the gases to behave ideally?
Use PV=nRT. Use R = 0.08206 L atm/mol K so that the units cancel.
P for H2 = 394 atm
P for O2 = 77.96 atm
For pressures above 10 atm, you would not expect the gases to behave ideally. These pressures are far higher than 10 atm, so you would not expect the gases to behave ideally. However, the question indicates that you should use the ideal gas law, although it is likely a poor approximation in this situation.
4.) How many moles of water are formed? How many moles of excess reagent are there?
4.75 mol H2O formed
7.25 mol excess O2
5.) What are the resulting partial pressures of water and the excess reagent in the combined cylinders? What is the total pressure?
Partial pressures: 78.0 atm H2O, 119 atm H2
Total pressure: 197 atm
*Note that each resulting number should have two significant figures. Rounding should be done that the end here, but extra digits were included for comparison.
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