The ideal gas law states that a gas sample will be composed of molecules which are totally independent of one another and therefore will behave "ideally". In actual fact, the Ideal Gas Law is actually an equation:
PV = nRT
where R = universal gas law constant
P= Pressure
V= Volume
n= number of moles
T= Temperature
The Universal Gas Law Constant R, will have a constant value associated with a set of units. Quite often the value for R=.0821 liter-atm/mol-Kelvin. This sets the units that the other variables must be expressed in. Using this R value means that Pressure must be in atmospheres, Volume must be in liter units, Temperature must be in Kelvin units and the amount of gas, n, must be in moles units.
If a gas behaves exactly as the ideal gas laws would predict it to behave in terms of volume, pressure, moles, and temperature, then the gas is said to be an ideal gas.
If, on the other hand, the gas deviates from Ideal Gas behavior, then the gas is said to be acting like a "real gas".
As an example, If I compute theoretically using the Ideal Gas Law equation the volume of a gas sample of known moles (1), temperature (0 C or 273 K), and pressure (1 atmosphere), the volume would be 22.412 liters and that would be a theoretical value based on the Ideal Gas Law assuming that the molecules of gas acted independently of one another and did not have so called "intermolecular forces" acting upon them.
Now if I then go to the laboratory and in an experimental set-up, I reproduce the exact theoretical conditions of moles (1), temperature ( 0 C or 273 K), and pressure (1 atmosphere) in the laboratory. I then proceed to measure the volume of the gas sample under those real lab conditions, and the volume is determined to be experimentally 22.207 liters. If I compare my theoretical volume (22.412 liters) and my experimental volume (22.207 liters), I find that they do not match, but they deviate. The volume that the gas should be occupying under the stated conditions under ideal conditions is not the same as the volume of the gas as measured in the laboratory. We say that the gas in the lab is deviating in its volume behavior from what is expected ideally. In other words the gas is behaving as a "real" gas and deviates from "ideality".
On the other hand if I were to lower the pressure and/or raise the temperature of the gas sample and then proceed to recalculate the ideal (calculated) and real (measured) volumes. I would find that there was not as much difference between the two volume values. The gas is said to tend toward greater ideality. If the two values were the same, then we could say that at those conditions of temperature and pressure the gas is acting as an ideal gas and the deviation (between the theoretical calculation and the experimental measured value) is zero.
Using the "molecular view" we could say that the lowering of the pressure allowed the gas molecules to spread farther out from one another. Since the molecules are farther apart, there were fewer intermolecular forces acting upon them and the intermolecular forces are approaching zero. (At zero intermolecular forces the gas would behave as an ideal gas).
Raising the temperature of the gas sample in the lab would increase the Average Kinetic Energy of the gas molecules and cause them to speed up in their motion to one another. This increase in speed will overcome any intermolecular forces acting between the molecules and the gas would behave more as an ideal gas.
The question becomes what kinds of "intermolecular forces" would be present that would cause gas molecules to be interdependent instead of independent of one another?
Intermolecular forces are electrical forces that exist between molecules that would cause one molecule to influence another. Such forces will cause a gas to deviate from Ideal Gas behavior.
Intermolecular forces are called Van Der Waals forces . There are three such types of Van Der Waals forces:
Molecules can be polar or non-polar. What does this mean?
Polar molecules are molecules that have a concentration of negative charge on one end of the molecule and a concentration of positive charge on the other end. This is brought about by having the negative electrons within the atoms of the molecule shift toward those atoms that are most capable of attracting them. This shift produces an electrical force called a "molecular dipole". Other molecules that are polar will be attracted since each of them, in turn, have a positive and negative end. We know that positive attracts negative so these molecules will line up so that this statement is observed.
Non-polar molecules are molecules that do NOT have this concentration of positive or negative charge. That is because the shift of electrons occurs in a symmetrical (balanced) way so that there is, net wise, no shift of electrons within the molecule hence no concentration of negative electrons on one end of the molecule.
You can do a simple demonstration to show that molecules that are polar do have positive and negative areas within the molecule. If you take some burets and fill them with different liquids such as water (polar), ethyl alcohol(polar), carbon tetrachloride (non-polar). Then take an amber rod and rub it with animal fur to impart a negative charge upon the rod. Open up the buret so that the liquid can begin flowing aout of the buret in a stream. Bring the negatively charged rod close to the liquid stream and see what happens. What do you suppose will happen if the molecules are polar when the negatively charged rod is brought very close to the stream of liquid without actually touching it? What about the non-polar carbon tetrachloride?
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R. H. Logan, Instructor of Chemistry, Dallas County Community College
District, North Lake College.
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All contents copyrighted (c) 1996 R.H. Logan, Instructor of Chemistry,DCCCD All Rights reservedRevised: 7/10/97
Original Date of Creation: 11/19/95
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