A mixture of gas contains 1.25g of N2, 1.25g of O2, and 1.25g of He in 2.50 L container at 25 C. Find the partial pressure of N2 in the container?
Convert 1.25 g N2 to moles and use PV = nRT. Don't forget to change T to Kelvin.
0.490 mol of argon gas is admitted to an evacuated 50.0 cm^3 container at 60.0 C. The gas then undergoes an isochoric heating to a temperature of 400 C. What is the pressure?
To find the partial pressure of N2 in the container, we need to use the ideal gas law equation:
PV = nRT
Where:
P is the pressure
V is the volume of the container
n is the number of moles of the gas
R is the ideal gas constant (0.0821 L·atm/(K·mol))
T is the temperature in Kelvin
First, let's find the number of moles of N2 in the container. We can use the equation:
moles = mass / molar mass
The molar mass of N2 is 28 g/mol.
moles of N2 = 1.25g / 28 g/mol
= 0.0446 mol
Now let's calculate the partial pressure of N2:
P(N2) = (moles of N2 * R * T) / V
Since the temperature needs to be in Kelvin, we add 273 to the given temperature in Celsius:
T = 25°C + 273
= 298 K
Now we can substitute the values into the equation:
P(N2) = (0.0446 mol * 0.0821 L·atm/(K·mol) * 298 K) / 2.50 L
= 0.492 atm
Therefore, the partial pressure of N2 in the container is 0.492 atm.