Answer:
To calculate the pressure exerted by 0.25 moles of chlorine gas in a volume of 5.00 L at 37°C using the ideal gas equation, we can use the following formula: PV = nRT
Where P is the pressure, V is the volume, n is the number of moles, R is the gas constant, and T is the temperature in Kelvin. We first need to convert the temperature from Celsius to Kelvin by adding 273.15 to the temperature:
T = 37°C + 273.15 = 310.15 K
Substituting the values given in the problem, we get:
P(5.00 L) = (0.25 mol)(0.08206 L·atm/mol·K)(310.15 K)
Solving for P, we get: P = 3.15 atm
Therefore, the pressure exerted by 0.25 moles of chlorine gas in a volume of 5.00 L at 37°C using the ideal gas equation is 3.15 atm.
To calculate the pressure using the van der Waals equation, we need to use the following formula: (P + a(n/V)^2)(V - nb) = nRT
Where a and b are constants specific to the gas, and n/V is the molar density. For chlorine gas, a = 6.49 L^2·atm/mol^2 and b = 0.0562 L/mol.
We can calculate n/V as follows: n/V = 0.25 mol / 5.00 L = 0.05 mol/L
Substituting the values given in the problem, we get: (P + 6.49 L^2·atm/mol^2 (0.05 mol/L)^2)(5.00 L - 0.0562 L/mol (0.25 mol)) = (0.25 mol)(0.08206 L·atm/mol·K)(310.15 K)
Simplifying and solving for P, we get:
P = 3.11 atm
Therefore, the pressure exerted by 0.25 moles of chlorine gas in a volume of 5.00 L at 37°C using the van der Waals equation is 3.11 atm.
Step-by-step explanation:
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