Question

Oxygen difluoride can be synthesized via the balanced equation below. If excess sodium

hydroxide is used, how many grams of molecular fluorine must a researcher add to the
reaction in order to collect 50.0 mL of oxygen difluoride at a pressure of 1.10 atm and
a temperature of 24.5 ◦C? Note that R = 0.08206 L·atm/K·mole.
2F2 + 2NaOH −→ OF2 + 2NaF + H2O

Answer 1

To calculate the grams of molecular fluorine needed to produce 50.0 mL of oxygen difluoride, we can use the ideal gas law and stoichiometry.

In order to determine the amount of molecular fluorine needed to produce 50.0 mL of oxygen difluoride, we need to use the ideal gas law.

The equation is balanced, so we can use the stoichiometry to convert between the volume of oxygen difluoride and the moles of molecular fluorine.

Then, we can convert moles of molecular fluorine to grams using its molar mass.

First, let's convert the volume of oxygen difluoride to moles using the ideal gas law: PV = nRT. We're given the pressure, temperature, and volume, so we can solve for moles of oxygen difluoride.

We'll use the molar volume (22.4 L/mol) to convert from moles to liters.

Next, we can use the stoichiometry of the balanced equation to convert from moles of oxygen difluoride to moles of molecular fluorine.

The stoichiometry tells us that 2 moles of molecular fluorine are required for every 1 mole of oxygen difluoride.

Finally, we can convert moles of molecular fluorine to grams using its molar mass. The molar mass of molecular fluorine is 38.0 g/mol.