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.
![A sample of carbon dioxide is contained in a 250.0 mL flask at 0.977 atm and 15.3 °C. How many molecules of gas are in the sample?
[ ] molecules
**Explanation for Educational Use**:
To solve this problem, you would typically use the Ideal Gas Law, expressed as:
\[ PV = nRT \]
where:
- \( P \) is the pressure (0.977 atm)
- \( V \) is the volume (250.0 mL, which needs to be converted to liters)
- \( n \) is the number of moles of gas
- \( R \) is the ideal gas constant (0.0821 L·atm/mol·K)
- \( T \) is the temperature in Kelvin (15.3 °C, which needs to be converted to Kelvin)
You would rearrange the formula to solve for \( n \) (the number of moles):
\[ n = \frac{PV}{RT} \]
After finding the number of moles, you would use Avogadro's number (approximately \( 6.022 \times 10^{23} \) molecules/mol) to calculate the number of molecules.](/v2/_next/image?url=https%3A%2F%2Fcontent.bartleby.com%2Fqna-images%2Fquestion%2F38e21668-df36-4535-9b97-aa4a56ba16f0%2Fb9459a8c-dbcf-404c-925d-5153d5c63a50%2Fsm4bbvr_processed.jpeg&w=3840&q=75)
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