Heat Capacity Ratio for various gases^{[1]}^{[2]}

Temp.

Gas

γ


Temp.

Gas

γ


Temp.

Gas

γ

−181 °C

H_{2}

1.597

200 °C

Dry Air

1.398

20 °C

NO

1.400

−76 °C

1.453

400 °C

1.393

20 °C

N_{2}O

1.310

20 °C

1.410

1000 °C

1.365

−181 °C

N_{2}

1.470

100 °C

1.404

2000 °C

1.088

15 °C

1.404

400 °C

1.387

0 °C

CO_{2}

1.310

20 °C

Cl_{2}

1.340

1000 °C

1.358

20 °C

1.300

−115 °C

CH_{4}

1.410

2000 °C

1.318

100 °C

1.281

−74 °C

1.350

20 °C

He

1.660

400 °C

1.235

20 °C

1.320

20 °C

H_{2}O

1.330

1000 °C

1.195

15 °C

NH_{3}

1.310

100 °C

1.324

20 °C

CO

1.400

19 °C

Ne

1.640

200 °C

1.310

−181 °C

O_{2}

1.450

19 °C

Xe

1.660

−180 °C

Ar

1.760

−76 °C

1.415

19 °C

Kr

1.680

20 °C

1.670

20 °C

1.400

15 °C

SO_{2}

1.290

0 °C

Dry Air

1.403

100 °C

1.399

360 °C

Hg

1.670

20 °C

1.400

200 °C

1.397

15 °C

C_{2}H_{6}

1.220

100 °C

1.401

400 °C

1.394

16 °C

C_{3}H_{8}

1.130

In thermal physics and thermodynamics, the heat capacity ratio or adiabatic index or ratio of specific heats or Poisson constant, is the ratio of the heat capacity at constant pressure (C_P) to heat capacity at constant volume (C_V). It is sometimes also known as the isentropic expansion factor and is denoted by \gamma (gamma)(for ideal gas) or \kappa (kappa)(isentropic exponent, for real gas). The former symbol gamma is primarily used by chemical engineers. Mechanical engineers use the Roman letter k.^{[3]}

\gamma = \frac{C_P}{C_V} = \frac{c_P}{c_V}
where, C is the heat capacity and c the specific heat capacity (heat capacity per unit mass) of a gas. Suffix P and V refer to constant pressure and constant volume conditions respectively.
To understand this relation, consider the following thought experiment. A closed pneumatic cylinder contains air. The piston is locked. The pressure inside is equal to atmospheric pressure. This cylinder is heated to a certain target temperature. Since the piston cannot move, the volume is constant. The temperature and pressure will rise. When the target temperature is reached, the heating is stopped. The amount of energy added equals: C_V \Delta T, with \Delta T representing the change in temperature. The piston is now freed and moves outwards, stopping as the pressure inside the chamber equilibrates to atmospheric pressure. We are free to assume the expansion happens fast enough to occur without exchange of heat (adiabatic expansion). Doing this work, air inside the cylinder will cool to below the target temperature. To return to the target temperature (still with a free piston), the air must be heated. This extra heat amounts to about 40% more than the previous amount added. In this example, the amount of heat added with a locked piston is proportional to C_V , whereas the total amount of heat added is proportional to C_P. Therefore, the heat capacity ratio in this example is 1.4.
Another way of understanding the difference between C_P and C_V is that C_P applies if work is done to the system which causes a change in volume (e.g. by moving a piston so as to compress the contents of a cylinder), or if work is done by the system which changes its temperature (e.g. heating the gas in a cylinder to cause a piston to move). C_V applies only if P dV  that is, the work done  is zero. Consider the difference between adding heat to the gas with a locked piston, and adding heat with a piston free to move, so that pressure remains constant. In the second case, the gas will both heat and expand, causing the piston to do mechanical work on the atmosphere. The heat that is added to the gas goes only partly into heating the gas, while the rest is transformed into the mechanical work performed by the piston. In the first, constantvolume case (locked piston) there is no external motion, and thus no mechanical work is done on the atmosphere; C_V is used. In the second case, additional work is done as the volume changes, so the amount of heat required to raise the gas temperature (the specific heat capacity) is higher for this constant pressure case.
Contents

Ideal gas relations 1

Relation with degrees of freedom 1.1

Real gas relations 2

Thermodynamic expressions 3

Adiabatic process 4

See also 5

References 6
Ideal gas relations
For an ideal gas, the heat capacity is constant with temperature. Accordingly we can express the enthalpy as H = C_P T and the internal energy as U = C_V T. Thus, it can also be said that the heat capacity ratio is the ratio between the enthalpy to the internal energy:

\gamma = \frac{H}{U}
Furthermore, the heat capacities can be expressed in terms of heat capacity ratio ( \gamma ) and the gas constant ( R ):

C_P = \frac{\gamma n R}{\gamma  1} \qquad \mbox{and} \qquad C_V = \frac{n R}{\gamma  1},
where n is the amount of substance in moles.
It can be rather difficult to find tabulated information for C_V, since C_P is more commonly tabulated. The following relation, can be used to determine C_V:

C_V = C_P  nR
Relation with degrees of freedom
The heat capacity ratio ( \gamma ) for an ideal gas can be related to the degrees of freedom ( f ) of a molecule by:

\gamma\ = 1 + \frac{2}{f}\qquad \mbox{or} \qquad f = \frac{2}{\gamma1}
Thus we observe that for a monatomic gas, with three degrees of freedom:

\gamma\ = \frac{5}{3} \approx 1.67,
while for a diatomic gas, with five degrees of freedom (at room temperature: three translational and two rotational degrees of freedom; the vibrational degree of freedom is not involved except at high temperatures):

\gamma = \frac{7}{5} = 1.4.
E.g.: The terrestrial air is primarily made up of diatomic gases (~78% nitrogen (N_{2}) and ~21% oxygen (O_{2})) and at standard conditions it can be considered to be an ideal gas. The above value of 1.4 is highly consistent with the measured adiabatic indices for dry air within a temperature range of 0 to 200 °C, exhibiting a deviation of only 0.2% (see tablation above).
Real gas relations
As temperature increases, higher energy rotational and vibrational states become accessible to molecular gases, thus increasing the number of degrees of freedom and lowering \gamma. For a real gas, both C_P and C_V increase with increasing temperature, while continuing to differ from each other by a fixed constant (as above, C_P = C_V + nR) which reflects the relatively constant PV difference in work done during expansion, for constant pressure vs. constant volume conditions. Thus, the ratio of the two values, \gamma, decreases with increasing temperature. For more information on mechanisms for storing heat in gases, see the gas section of specific heat capacity.
Thermodynamic expressions
Values based on approximations (particularly C_P  C_V = nR) are in many cases not sufficiently accurate for practical engineering calculations such as flow rates through pipes and valves. An experimental value should be used rather than one based on this approximation, where possible. A rigorous value for the ratio \frac{C_P}{C_V} can also be calculated by determining C_V from the residual properties expressed as:

C_P  C_V \ = \ T \frac \right)_P^2 }} {\left(\frac{\part V}{\part P}\right)_T} \ = \ T \frac \right) }_V^2} {\left( \frac{\part P}{\part V} \right)_T}
Values for C_P are readily available and recorded, but values for C_V need to be determined via relations such as these. See here for the derivation of the thermodynamic relations between the heat capacities.
The above definition is the approach used to develop rigorous expressions from equations of state (such as Peng–Robinson), which match experimental values so closely that there is little need to develop a database of ratios or C_V values. Values can also be determined through finite difference approximation.
Adiabatic process
This ratio gives the important relation for an isentropic (quasistatic, reversible, adiabatic process) process of a simple compressible calorically perfect ideal gas:

PV^\gamma = \text{constant}
where P is the pressure and V is the volume.
See also
References

^ White, Frank M.: Fluid Mechanics 4th ed. McGraw Hill

^ Lange's Handbook of Chemistry, 10th ed. page 1524

^ Fox, R., A. McDonald, P. Pritchard: Introduction to Fluid Mechanics 6th ed. Wiley
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