We’re being asked to calculate the pressure exerted by** one mole of chlorine, Cl**_{2} gas using the Van der Waal’s equation.

The Van der Waals equation is shown below:

$\overline{)\left(\mathbf{P}\mathbf{+}\mathbf{a}\frac{{\mathbf{n}}^{\mathbf{2}}}{{\mathbf{V}}^{\mathbf{2}}}\right)\left(\mathbf{V}\mathbf{-}\mathbf{n}\mathbf{b}\right){\mathbf{=}}{\mathbf{n}}{\mathbf{R}}{\mathbf{T}}}$

▪ **P** = pressure, atm

▪ **V** = volume, L

▪ **n** = # of moles, mol

▪ **R** = gas constant = 0.08206 (L∙atm)/(mol∙K)

▪ **T** = temperature, K

▪ **a** = polarity coefficient

▪ **b **= size coefficient

Let’s first isolate the pressure in the **Van der Waals Equation**:

In Sample Exercise 10.16 in the textbook, we found that one mole of Cl_{2} confined to 22.41 L at 0 ^{o}C deviated slightly from ideal behavior. Calculate the pressure exerted by 1.00 mol Cl_{2} confined to a smaller volume, 6.00 L , at 25 ^{o}C.

Use van der Waals equation for your calculation. (Values for the van der Waals constants are a = 6.49 L^{2}atm/mol^{2}, b = 0.0562 L/mol.)

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