# Heat capacity ratio

Heat Capacity Ratio for various gases[1][2]
Temp. Gas γ   Temp. Gas γ   Temp. Gas γ
−181°C H2 1.597 200°C Dry Air 1.398 20°C NO 1.400
−76°C 1.453 400°C 1.393 20°C N2O 1.310
20°C 1.410 1000°C 1.365 −181°C N2 1.470
100°C 1.404 2000°C 1.088 15°C 1.404
400°C 1.387 0°C CO2 1.310 20°C Cl2 1.340
1000°C 1.358 20°C 1.300 −115°C CH4 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 H2O 1.330 1000°C 1.195 15°C NH3 1.310
100°C 1.324 20°C CO 1.400 19°C Ne 1.640
200°C 1.310 −181°C O2 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 SO2 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 C2H6 1.220
100°C 1.401 400°C 1.394 16°C C3H8 1.130

The heat capacity ratio or adiabatic index or ratio of specific heats, 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, constant-volume 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.

## 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:

$\gamma\ = 1 + \frac{2}{f}\qquad \mbox{or} \qquad f = \frac{2}{\gamma-1}$

Thus we observe that for a monoatomic 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 (N2) and ~21% oxygen (O2)) 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 P*V 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:

$C_P - C_V \ = \ -T \frac{{\left( {\frac{\part V}{\part T}} \right)_P^2 }} {\left(\frac{\part V}{\part P}\right)_T} \ = \ -T \frac{{ \left( {\frac{\part P}{\part T}} \right) }_V^2} {\left( \frac{\part P}{\part V} \right)_T}$

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.

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.