Question

What mass of nitrogen (N2), oxygen (O2), and helium (He) is present in a 1.35 L sample of a gas mixture at 25.0 °C, given that the partial pressures are 236 torr, 155 torr, and 130 torr, respectively?

a) Mass of N2: [answer] g, Mass of O2: [answer] g, Mass of He: [answer] g
b) Mass of N2: [answer] g, Mass of O2: [answer] g, Mass of He: [answer] g
c) Mass of N2: [answer] g, Mass of O2: [answer] g, Mass of He: [answer] g
d) Mass of N2: [answer] g, Mass of O2: [answer] g, Mass of He: [answer] g

Answer 1

To find the mass of each gas in the mixture, calculate the number of moles for N2, O2, and He using the ideal gas law and their respective partial pressures. Multiply the moles by the molar mass of each gas to get the mass.

Step-by-step explanation:

To calculate the mass of nitrogen (N2), oxygen (O2), and helium (He) in a gas mixture using their partial pressures, we can use the ideal gas law in the form PV = nRT, where P is pressure, V is volume, n is the number of moles, R is the ideal gas constant, and T is temperature. First, we must convert temperature to Kelvin (K) and pressure to atmospheres (atm). For 25.0 °C, which is 298.15 K, we use the conversion 1 atm = 760 torr to convert the partial pressures.

For N2, the partial pressure in atm is 236 torr / 760 torr/atm = 0.3105 atm. Using PV = nRT, we calculate n (moles of N2) and then multiply by the molar mass of N2, which is 28.0 g/mol. Repeat the same process for O2 and He. The mass of each gas in the sample is the number of moles multiplied by its respective molar mass.

The molar mass of oxygen (O2) is 32.0 g/mol, and helium (He) is 4.00 g/mol. After calculating the number of moles from the ideal gas law for each gas, we can find the mass for each gas.