How many liters of oxygen measured at 295 K and 763 torr are consumed in the complete combustion of 1.55 L of C2H4 measured at STP? C2H4 (g) +3 O2 (g) --> 2 CO2 (g) + 2 H20 (1)
Ideal and Real Gases
Ideal gases obey conditions of the general gas laws under all states of pressure and temperature. Ideal gases are also named perfect gases. The attributes of ideal gases are as follows,
Gas Laws
Gas laws describe the ways in which volume, temperature, pressure, and other conditions correlate when matter is in a gaseous state. The very first observations about the physical properties of gases was made by Robert Boyle in 1662. Later discoveries were made by Charles, Gay-Lussac, Avogadro, and others. Eventually, these observations were combined to produce the ideal gas law.
Gaseous State
It is well known that matter exists in different forms in our surroundings. There are five known states of matter, such as solids, gases, liquids, plasma and Bose-Einstein condensate. The last two are known newly in the recent days. Thus, the detailed forms of matter studied are solids, gases and liquids. The best example of a substance that is present in different states is water. It is solid ice, gaseous vapor or steam and liquid water depending on the temperature and pressure conditions. This is due to the difference in the intermolecular forces and distances. The occurrence of three different phases is due to the difference in the two major forces, the force which tends to tightly hold molecules i.e., forces of attraction and the disruptive forces obtained from the thermal energy of molecules.
![**Title: Calculating Partial Pressure and Mole Fraction of Oxygen**
Oxygen is collected over water at 22°C and a barometric pressure of 0.985 atm.
**a) What is the partial pressure of oxygen (in torr)?**
**b) What is the mole fraction of oxygen in the container?**
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**Note**: To solve part (a), use the fact that water vapor exerts a pressure that depends on temperature. You will need to subtract the vapor pressure of water at 22°C from the total barometric pressure to find the partial pressure of oxygen.
To solve part (b), once the partial pressure of oxygen is found, use it to calculate the mole fraction using the formula:
\[ \text{Mole fraction of oxygen} = \frac{\text{Partial pressure of oxygen}}{\text{Total pressure}} \]](/v2/_next/image?url=https%3A%2F%2Fcontent.bartleby.com%2Fqna-images%2Fquestion%2Fb7c14283-9a82-48f7-90b1-831241c404f8%2Feb4a43a0-6d5b-443a-b09f-9d2680b4a8e5%2Ft6617xg_processed.jpeg&w=3840&q=75)
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