Using the information on the table of van der Waals Constants of Some Common gases, calculate the pressure exerted by 4.37 moles of chlorine gas confined in a volume of 2.45 L at 38 degrees Celsius. Compare the pressure with that calculated using the ideal gas equation.
As I see it this is a case of looking up the numbers in the table and substituting them into the van der Waals equation, then comparing with that obtained with PV = nRT
How can we help you do that?
To calculate the pressure exerted by chlorine gas using the van der Waals equation, we need the van der Waals constants for chlorine (Cl2) from the table. Since the table is not provided, I cannot give you the exact values. However, I can guide you on how to solve the problem using the van der Waals equation and the ideal gas equation.
The van der Waals equation for a real gas is given by:
(P + a(n/V)^2)(V - nb) = nRT
where P is the pressure, V is the volume, n is the number of moles, R is the ideal gas constant (0.0821 L·atm/mol·K), T is the temperature in Kelvin, a is the van der Waals constant for attractive forces between molecules, and b is the van der Waals constant for molecular volume.
To calculate the pressure using the van der Waals equation, you'll need to know the values of a and b for chlorine gas. Once you have those values, you can substitute them into the equation along with the given values of n, V, and T.
To compare the pressure calculated using the van der Waals equation with that calculated using the ideal gas equation, we can use the ideal gas equation:
PV = nRT
Rearranging this equation, we get:
P = (n/V)RT
Plugging in the given values of n, V, and T, and the ideal gas constant R, you can calculate the pressure using the ideal gas equation.
By comparing the pressures calculated using both equations, you can determine the difference between the real behavior of chlorine gas and the behavior predicted by the ideal gas equation.