Question

A 5.0-L flask contains \(0.60 \mathrm{~g} \mathrm{O}_{2}\) at a temperature of \(22^{\circ} \mathrm{C}\). What is the pressure (in atm) inside the flask?

Short answer

So: \(T(K) = 22 + 273.15 = 295.15~K\) #tag_title#Step 2: Calculate moles of gas#tag_content#Next, we need to determine the amount of gas in moles. We know that the molar mass of \(O_2\) is 32.00 g/mol, so we can calculate the moles of gas as follows: \(n = \dfrac{mass}{molar~mass}\) Thus: \(n = \dfrac{0.60~g}{32.00~g/mol} = 0.01875~mol\) #tag_title#Step 3: Apply the Ideal Gas Law#tag_content#Now, we can use the Ideal Gas Law formula (PV = nRT) to find the pressure inside the flask. We know the volume V is 5.0 L and the temperature T is 295.15 K. The Ideal Gas Constant R is 0.0821 L atm/mol K. Solving for P, we get: \(P = \dfrac{nRT}{V}\) Plugging in the values: \(P = \dfrac{0.01875 \times 0.0821 \times 295.15}{5}\) Finally, we calculate the pressure: \(P \approx 0.913~atm\)

Step-by-step solution

Convert temperature to Kelvin

First, we need to convert the temperature from Celsius to Kelvin. The formula for that is: \(T(K) = T(°C) + 273.15\)