So2 (5.00 g) and co2 (5.00 g) were placed in a 750.0 ml container at 50.0 °c. the partial pressure of co2 in the container was atm.

4.02 atm. Since we're looking for just the partial pressure of CO2, we can safely ignore the sulfur dioxide since the total pressure in the container will be the sum of all the partial pressures and the partial pressure will be the pressure that the substance would have if it were the only substance in the container. So, let's determine the number of moles of CO2 we have. Start by looking up the atomic weights: Atomic weight carbon = 12.0107 Atomic weight oxygen = 15.999 Molar mass CO2 = 12.0107 + 2 * 15.999 = 44.0087 g/mol Moles CO2 = 5.00 g / 44.0087 g/mol = 0.113613899 mol The ideal gas law is PV = nRT P = Pressure V = Volume n = number of moles R = Ideal Gas Constant (0.082057338 L*atm/K*mol ) T = Absolute temperature Converting from C to K, gives 50.0 + 273.15 = 323.15 K Take the ideal gas law and solve for P, then substitute known values and calculate result. PV = nRT P = nRT/V P = (0.113613899 mol)(0.082057338 L*atm/K*mol)(323.15 K)/(0.75 L) P = (0.009322854 L*atm/K)(323.15 K)/(0.75 L) P = (3.012680306 L*atm)/(0.75 L) P = 4.016907075 atm Since we only have 3 significant digits of precision available, round to 3 significant digits, getting 4.02 atm.

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