2. Find the mass in grams of \( 2.00 \times 10^{23} \) molecules of \( F_{2} \).

To find the mass of \( 2.00 \times 10^{23} \) molecules of \( F_2 \), we first need to know how many moles this number corresponds to using Avogadro's number (\( 6.022 \times 10^{23} \) molecules/mol). Now, \( \frac{2.00 \times 10^{23} \text{ molecules}}{6.022 \times 10^{23} \text{ molecules/mol}} \approx 0.332 \text{ moles} \). Next, we calculate the molar mass of \( F_2 \). The atomic mass of fluorine is approximately \( 19.00 \text{ g/mol} \), so for \( F_2 \) it’s \( 19.00 \times 2 = 38.00 \text{ g/mol} \). Now we can calculate the mass: Mass = moles × molar mass = \( 0.332 \text{ moles} \times 38.00 \text{ g/mol} \approx 12.616 \text{ grams} \). So the mass of \( 2.00 \times 10^{23} \) molecules of \( F_2 \) is approximately \( 12.62 \text{ grams} \).