Perfect Gas/
Ideal Gas and Deviation of real gases from ideality
Definition of ideal gas
ideal gas is
an imaginary gas whose behaviour can be predicted precisely on the basis of
kinetic molecular theory and gas laws.
OR
It is a hypothetical gas which strictly obeys gas laws (Boyle’s
law, Charles’s law etc.) or ideal gas equation (PV = nRT as gas laws contained
in it) for all conditions of temperature and pressure behaving according to the
kinetic molecular theory and does not show deviation from gas laws. It is an
imaginary gas and has no real existence.
OR
Ideal gas is a model gas for which ratio of PV/nRT called
compressibility factor (z) is equal to 1.
In other words,
ideal gas is
a reference gas for which graph between P on x-axis and PV/nRT on y-axis would
be a horizontal line parallel to x-axis.
It is an imaginary gas and has no real existence. Actually
there is no gas which is perfectly ideal. Actually, none of the known gases
exactly follows the ideal gas laws and called as real gases. Some gases
such as H2, N2, O2 do not deviate greatly from
the ideal behaviour at moderate temperatures and pressures because of their smaller
size of molecules and weak forces of attraction between their
molecules. Thus, an ideal gas can only exist at low pressure and high
temperature. Among all real gases helium
acts most likely as an ideal gas at room temperature
Real gases shows deviation from gas laws particularly at high pressure and low
temperature. Therefore, at low temperature, the attractive forces become
significant, eventually the gas liquefies. This property of real gases cause deviates them from
ideal behaviour.
Ideality Sequence
He > H2 > N2 > CO2
> Nh3 > SO2
Conditions for Ideality
1. High temperature
2. Low pressure
Properties of an Ideal Gas in the light of KMT
1. Their molecules occupy no space or occupy negligible
volume and have point mass.
2. Their molecules exert no forces one another.
3. Their molecules
undergo elastic collisions.
4. It cannot be
liquefied due to lack of attractive forces between the molecules.
Graphical Explanation of Deviation
of Real Gases
To understand the attitude of real
gases graphically and their deviation from ideal
behaviour is expressed by introducing a factor known as compressibility factor ‘z’
in the ideal gas equation which is modified as should
PV = z
(nRT) or z = PV/nRT
In case of ideal gas, PV = nRT or z
= 1
In case of real gas, PV ≠ nRt or z ≠ 1 (less
or more than unity)
Thus in case of real gases z can be < 1
or > 1
compressibility factor is a ratio of PV and nRT.
In other word,
The ratio of
volume of real gas, Vreal to the ideal volume of that gas, Vperfect calculated
by ideal gas equation is known as compressibility factor.
z = PVreal/nRT
But from ideal gas equation:
PVperfect = nRT Or Vperfect
= nRT/P
Therefore
z = PVreal/nRT
= Vreal/Vperfect
It is clear from above graphs that the
volume of real gas is more than or less than expected in certain cases.
The graph
plotted between z and P at constant temperature is not a straight line for real gases. There is a significant deviation from the ideal behaviour.
The extent of deviation of real
gases from ideality is based on pressure, temperature and the nature of gases.
Two types of curves are seen. In the
curves for hydrogen and helium, as the pressure increases the value of pV also increases. The second type of
plot is seen in the case of other gases like carbon monoxide and methane. In
these plots, first there is a negative deviation from ideal behaviour, the pV value decreases with increase in
pressure and reaches to a minimum value characteristic of a gas. After that, pV value starts increasing. The curve
then crosses the line for ideal gas and after that shows positive deviation
continuously. It is thus, found that real gases do not follow ideal gas
equation perfectly under all conditions.
The graph plotted between z
and P for an ideal gas and various real gases provides the following
information:
Case-I : If Z=1
All gases reaches a value of z =1 , when the pressure approached to zero. This reveals
that all gases tend to act like ideal gas at very low pressure.
Case-II : If z>1
When Z > 1, it is a positive deviation.
Vreal > Videal
It shows that the gas is less compressible than expected from ideal
behaviour.
The repulsion forces become more significant than the attractive
forces.
The gas cannot be compressed easily.
Usually the z > 1
for so called permanent gases like He, H2 etc.
Case-III: If z
< 1
When Z < 1, it is a negative deviation.
Vreal < Videal
It shows that the gas is more compressible than expected from ideal
behaviour.
The attractive forces are more significant than the repulsive
forces.
The gas can be liquefied easily.
Usually the z < 1
for gases like NH3, CO2, SO2 etc.
Detailed Explanation
1. If
z = 1 then line would be parallel
to x-axis.
2. If z
< 1 then the line obtained will below the line of an ideal gas, which means
that there is larger decrease in volume of the gas than predicted by general
gas equation due to the appearance of attractive forces present among the
molecules.
3. If z
> 1 then the line obtained will above the line of an ideal gas, which means
that there is less decrease in the volume of the gas than predicted by general
gas equation due to the appearance of repulsive forces present among the
molecules.
Upto what pressure a gas will follow the
ideal gas law, depends upon nature of the gas and its temperature. The
temperature at which a real gas obeys ideal gas law over an appreciable range
of pressure is called Boyle temperature or Boyle point. Boyle
point of a gas depends upon its nature.
Above their Boyle point, real gases show
positive deviations from ideality and z
values are greater than one. The forces of attraction between the molecules are
very feeble. Below Boyle temperature real gases first show decrease in z value with increasing pressure, which
reaches a minimum value. On further increase in pressure, the value of z increases continuously. Above
explanation shows that at low pressure and high temperature gases show ideal
behaviour. These conditions are different for different gases.
Definition of real gas
The gases which do not obey gas laws strictly at all temperatures
and pressures and show deviation from ideal behaviour under all or certain
conditions of temperatures and pressures is called Real or Non-Ideal Gas. It
exists in the universe e.g. CO2, SO2, Cl2, F2
etc.
Conditions for Ideality for Real Gases
Real gases follow the ideal gas equation of state only at
sufficiently low densities. At atmospheric pressure, the ideal gas law is quite
well satisfied for most gases, but for some (e.g. water vapours, ammonia, SO2
etc.) there are deviations of 1 to 2%. e.g.
a). Boyle’s law (PV = K), is no longer satisfied at high pressure.
b). Charle’s law (V/T = K), begins to break down at low temperature
c). Avogadro’s law (V/n = K), does not hold strictly for real gases at
moderate pressure.
Deviation of Real Gases from Ideal Behaviour
A Real Gas is one, which do not obey gas laws strictly and show
deviation from ideal behaviour under all or certain conditions of temperature
and pressure.
Conditions
of Deviation of Real Gases from Ideal Behaviour/Conditions for Non-Ideality
Real gases show considerable deviations from their ideal behaviour
at very high pressure and at very low temperature. Thus, there are two
conditions under which real gases show considerable deviation from their ideal
behaviour:
1. At very high pressure.
2. At very low temperature.
Causes of Deviation
To analyze why real gases deviate from ideal behaviour, we should
know the basics of formulation of the ideal gas equation, which was obtained
from certain faulty assumptions of kinetic molecular theory. The deviation of
real gases from the ideality is due to following to erroneous assumptions of kinetic molecular theory:
1. Actual volume of gas
molecules is negligibly small as compared to total volume of gas.
2. Gas Molecules have
neither attractive nor repulsive forces.
Contrary to these faulty assumptions, molecules of all gases do
exert some attractive force on one another and they do occupy some space in the
total volume of gas. The assumption 1 is valid at low pressure while assumption
2 is effective at high temperature. The postulate 1 is responsible for
deviation at very high pressure and postulate 2 for the deviation at low
temperature.
The reason is that the attractive forces diminish rapidly
as the distance between molecules increases. Thus at high pressure and low
temperature intermolecular forces becomes significant because molecules tend to
be close together.
The main reason of deviation of real gases from ideality is the
presence of intermolecular forces and un-negligible volume of gas molecules.
Reason of deviation
at high pressure
At high pressure, the gas volume decreases and gas molecules come
closer (thereby decreasing the empty spaces among molecules) due to which
volume occupied by the actual gas molecule does not remain negligible as
compared to total volume of gas. Moreover, attractive forces also arise among
molecules at high pressure. Hence at
high pressure gases deviate from ideal behaviour.
Reason of deviation
at low temperature
At low temperature, the kinetic energy of gas molecules decreases
and intermolecular attractive forces become more effective. Hence at low
temperature, gases deviate from ideal behaviour. There are two types of
intermolecular forces.
1. van der Waal’s Forces
It is the force between gaseous molecules arising due to transit or
temporary polarity due to unsymmetrical distribution of electronic charge cloud
around the nucleus. e.g. London dispersion forces.
2. Dipole-Dipole
Interaction
It is the force between gaseous molecules arising due to permanent
polarity because of unequal distribution of electronic charge cloud around the
nucleus e.g. Hydrogen bonding; H–Cl, NH3, H2S etc.
Ideal Vs Real Gas
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