Oxygen gas, from the decomposition of KClO3, was collected by water displacement. The pressure and temperature in the lab during the experiment were 451.0 torr and 20.0 C. What was the partial pressure of oxygen?
Look up the vapor pressure of H2O at 20 C. This is pH2O.
Then Ptotal = pO2 + pH2O
You know Ptotal and pH2O, solve for pO2.
To find the partial pressure of oxygen, we need to apply Dalton's law of partial pressures, which states that the total pressure of a mixture of gases is equal to the sum of the partial pressures of each individual gas.
In this case, the oxygen gas is the only gas collected, so its partial pressure is equal to the total pressure in the lab.
Therefore, the partial pressure of oxygen is 451.0 torr.
To calculate the partial pressure of oxygen, we need to use the ideal gas law equation, which states that:
PV = nRT
Where:
P is the pressure of the gas
V is the volume of the gas
n is the number of moles of gas
R is the ideal gas constant
T is the temperature in Kelvin
In this case, we know the pressure (451.0 torr) and the temperature (20.0 °C), but we need to convert the temperature to Kelvin.
To convert Celsius to Kelvin, we use the equation:
T(K) = T(°C) + 273.15
So, let's convert the temperature from Celsius to Kelvin:
T(K) = 20.0 + 273.15
T(K) = 293.15 K
Now that we have the temperature in Kelvin, we can rearrange the ideal gas law equation to solve for the number of moles of gas (n):
n = PV / RT
We need to use the molar volume of an ideal gas at standard temperature and pressure (STP), which is 22.4 L/mol, to calculate the volume of the gas collected. However, since the gas was collected by water displacement and the question does not provide the volume measurement, we cannot determine the number of moles directly.
Therefore, without additional information, we cannot calculate the partial pressure of oxygen.