In thermodynamics, the specific heat capacity (symbol cp) of a substance is the heat capacity of a sample of the substance divided by the mass of the sample. Informally, it is the amount of energy that must be added, in the form of heat, to one unit of mass of the substance in order to cause an increase of one unit in temperature. The SI unit of specific heat is joule per kelvin and kilogram, J/(K kg). For example, at a temperature of (the specific heat capacity can vary with the temperature), the heat required to raise the temperature of of water by (equivalent to ) is , meaning that the specific heat capacity of water is .
The specific heat often varies with temperature, and is different for each state of matter. Liquid water has one of the highest specific heats among common substances, about 4182 J/(K kg) at 20 °C; but that of ice just below 0 °C is only 2093 J/(K kg). The specific heats of iron, granite, and hydrogen gas are about 449, 790, and 14300 J/(K kg), respectively. While the substance is undergoing a phase transition, such as melting or boiling, its specific heat is technically infinite, because the heat goes into changing its state rather than raising its temperature.
The specific heat of a substance, especially a gas, may be significantly higher when it is allowed to expand as it is heated (specific heat at constant pressure) than when is heated in a closed vessel that prevents expansion (specific heat at constant volume). These two values are usually denoted by and , respectively; their quotient is the heat capacity ratio.
In some contexts, however, the term specific heat capacity (or specific heat) may refer to the ratio between the specific heats of a substance at a given temperature and of a reference substance at a reference temperature, such as water at 15 °C; much in the fashion of specific gravity.
Specific heat relates to other intensive measures of heat capacity with other denominators. If the amount of substance is measured as a number of moles, one gets the molar heat capacity instead (whose SI unit is joule per kelvin per mole, J/(K mol). If the amount is taken to be the volume of the sample (as is sometimes done in engineering), one gets the volumetric heat capacity (whose SI unit is joule per kelvin per cubic meter, J/K/m3).
One of the first scientists to use the concept was Joseph Black, 18th-century medical doctor and professor of Medicine at Glasgow University. He measured the specific heat of many substances, using the term capacity for heat.
The specific heat capacity of a substance, usually denoted by , is the heat capacity of a sample of the substance, divided by the mass of the sample:
where represents the amount of heat needed to uniformly raise the temperature of the sample by a small increment .
Like the heat capacity of an object, the specific heat of a substance may vary, sometimes substantially, depending on the starting temperature of the sample and the pressure applied to it. Therefore, it should be considered a function of those two variables.
These parameters are usually specified when giving the specific heat of a substance. For example, "Water (liquid): = 4185.5 J/K/kg (15 °C, 101.325 kPa)"  When not specified, published values of the specific heat generally are valid for some standard conditions for temperature and pressure.
However, the dependency of on starting temperature and pressure can often be ignored in practical contexts, e.g. when working in narrow ranges of those variables. In those contexts one usually omits the qualifier , and approximates the specific heat by a constant suitable for those ranges.
Specific heat is an intensive property of a substance, an intrinsic characteristic that does not depend on the size or shape of the amount in consideration. (The qualifier "specific" in front of an extensive property often indicates an intensive property derived from it.)
The injection of heat energy into a substance, besides raising its temperature, usually causes an increase in its volume and/or its pressure, depending on how the sample is confined. The choice made about the latter affects the measured specific heat, even for the same starting pressure and starting temperature . Two particular choices are widely used:
The value of is usually less than the value of . This difference is particularly notable in gases where values under constant pressure are typically 30% to 66.7% greater than those at constant volume. Hence the heat capacity ratio of gases is typically between 1.3 and 1.67.
The specific heat can be defined and measured for gases, liquids, and solids of fairly general composition and molecular structure. These include gas mixtures, solutions and alloys, or heterogenous materials such as milk, sand, granite, and concrete, if considered at a sufficiently large scale.
The specific heat can be defined also for materials that change state or composition as the temperature and pressure change, as long as the changes are reversible and gradual. Thus, for example, the concepts are definable for a gas or liquid that dissociates as the temperature increases, as long as the products of the dissociation promptly and completely recombine when it drops.
The specific heat is not meaningful if the substance undergoes irreversible chemical changes, or if there is a phase change, such as melting or boiling, at a sharp temperature within the range of temperatures spanned by the measurement.
The specific heat of a substance is typically determined according to the definition; namely, by measuring the heat capacity of a sample of the substance, usually with a calorimeter, and dividing by the sample's mass . Several techniques can be applied for estimating the heat capacity of a substance as for example fast differential scanning calorimetry.
The specific heat of gases can be measured at constant volume, by enclosing the sample in a rigid container. On the other hand, measuring the specific heat at constant volume can be prohibitively difficult for liquids and solids, since one often would need impractical pressures in order to prevent the expansion that would be caused by even small increases in temperature. Instead, the common practice is to measure the specific heat at constant pressure (allowing the material to expand or contract as it wishes), determine separately the coefficient of thermal expansion and the compressibility of the material, and compute the specific heat at constant volume from these data according to the laws of thermodynamics.
The SI unit for specific heat is joule per kelvin per kilogram (J/K/kg, J/(kg K), J K-1 kg-1, etc.). Since an increment of temperature of one degree Celsius is the same as an increment of one kelvin, that is the same as joule per degree Celsius per kilogram (J/°C/kg). Sometimes the gram is used instead of kilogram for the unit of mass: 1 J/K/kg = 0.001 J/K/g.
Professionals in construction, civil engineering, chemical engineering, and other technical disciplines, especially in the United States, may use the so-called English Engineering units, that include the Imperial pound (lb = 0.45359237 kg) as the unit of mass, the degree Fahrenheit or Rankine (°F = 5/9 K, about 0.555556 K) as the unit of temperature increment, and the British thermal unit (BTU ? 1055.06 J), as the unit of heat.
In those contexts, the unit of specific heat is BTU/°F/lb = 4177.6 J/K/kg. The BTU was originally defined so that the average specific heat of water would be 1 BTU/°F/lb.
In chemistry, heat amounts were often measured in calories. Confusingly, two units with that name, denoted "cal" or "Cal", have been commonly used to measure amounts of heat:
While these units are still used in some contexts (such as kilogram calorie in nutrition), their use is now deprecated in technical and scientific fields. When heat is measured in these units, the unit of specific heat is usually
In either unit, the specific heat of water is approximately 1. The combinations cal/°C/kg = 4.184 J/K/kg and kcal/°C/g = 4184,000 J/K/kg do not seem to be widely used.
The temperature of a sample of a substance reflects the average kinetic energy of its constituent particles (atoms or molecules) relative to its center of mass. However, not all energy provided to a sample of a substance will go into raising its temperature, exemplified via the equipartition theorem.
Quantum mechanics predicts that, at room temperature and ordinary pressures, an isolated atom in a gas cannot store any significant amount of energy except in the form of kinetic energy. Thus, heat capacity per mole is the same for all monoatomic gases (such as the noble gases). More precisely, 12.5 J/K/mol and 21 J/K/mol, where 8.31446 J/K/mol is the ideal gas unit (which is the product of Boltzmann conversion constant from kelvin microscopic energy unit to the macroscopic energy unit joule, and Avogadro's number).
Therefore, the specific heat (per unit of mass, not per mole) of a monoatomic gas will be inversely proportional to its (adimensional) atomic weight . That is, approximately,
For the noble gases, from helium to xenon, these computed values are
On the other hand, a polyatomic gas molecule (consisting of two or more atoms bound together) can store heat energy in other forms besides its kinetic energy. These forms include rotation of the molecule, and vibration of the atoms relative to its center of mass.
These extra degrees of freedom or "modes" contribute to the specific heat of the substance. Namely, when heat energy is injected into a gas with polyatomic molecules, only part of it will go into increasing their kinetic energy, and hence the temperature; the rest will go to into those other degrees of freedom. In order to achieve the same increase in temperature, more heat energy will have to be provided to a mol of that substance than to a mol of a monoatomic gas. Therefore, the specific heat of a polyatomic gas depends not only on its molecular mass, but also on the number of degrees of freedom that the molecules have.
Quantum mechanics further says that each rotational or vibrational mode can only take or lose energy in certain discrete amount (quanta). Depending on the temperature, the average heat energy per molecule may be too small compared to the quanta needed to activate some of those degrees of freedom. Those modes are said to be "frozen out". In that case, the specific heat of the substance is going to increase with temperature, sometimes in a step-like fashion, as more modes become unfrozen and start absorbing part of the input heat energy.
For example, the molar heat capacity of nitrogen at constant volume is 20.6 J/K/mol (at 15 °C, 1 atm), which is 2.49. That is the value expected from theory if each molecule had 5 degrees of freedom. These turn out to be three degrees of the molecule's velocity vector, plus two degrees from its rotation about an axis through the center of mass and perpendicular to the line of the two atoms. Because of those two extra degrees of freedom, the specific heat of (736 J/K/kg) is greater than that of an hypothetical monoatomic gas with the same molecular mass 28 (445 J/K/kg), by a factor of 5/3.
This value for the specific heat of nitrogen is practically constant from below -150 °C to about 300 °C. In that temperature range, the two additional degrees of freedom that correspond to vibrations of the atoms, stretching and compressing the bond, are still "frozen out". At about that temperature, those modes begin to "un-freeze", and as a result starts to increase rapidly at first, then slower as it tends to another constant value. It is 35.5 J/K/mol at 1500 °C, 36.9 at 2500 °C, and 37.5 at 3500 °C. The last value corresponds almost exactly to the predicted value for 7 degrees of freedom per molecule.
To apply the theory, one considers the sample of the substance (solid, liquid, or gas) for which the specific heat can be defined; in particular, that it has homogeneous composition and fixed mass . Assume that the evolution of the system is always slow enough for the internal pressure and temperature be considered uniform throughout. The pressure would be equal to the pressure applied to it by the enclosure or some surrounding fluid, such as air.
The state of the material can then be specified by three parameters: its temperature , the pressure , and its specific volume , where is the volume of the sample. (This quantity is the reciprocal of the material's density .) Like and , the specific volume is an intensive property of the material and its state, that does not depend on the amount of substance in the sample.
Those variables are not independent. The allowed states are defined by an equation of state relating those three variables: The function depends on the material under consideration. The specific internal energy stored internally in the sample, per unit of mass, will then be another function of these state variables, that is also specific of the material. The total internal energy in the sample then will be .
For some simple materials, like an ideal gas, one can derive from basic theory the equation of state and even the specific internal energy In general, these functions must be determined experimentally for each substance.
The absolute value of this quantity is undefined, and (for the purposes of thermodynamics) the state of "zero internal energy" can be chosen arbitrarily. However, by the law of conservation of energy, any infinitesimal increase in the total internal energy must be matched by the net flow of heat energy into the sample, plus any net mechanical energy provided to it by enclosure or surrounding medium on it. The latter is , where is the change in the sample's volume in that infinitesimal step. Therefore
If the volume of the sample (hence the specific volume of the material) is kept constant during the injection of the heat amount , then the term is zero (no mechanical work is done). Then, dividing by ,
where is the change in temperature that resulted from the heat input. The left-hand side is the specific heat at constant volume of the material.
For the heat capacity at constant pressure, it is useful to define the specific enthalpy of the system as the sum . An infinitesimal change in the specific enthalpy will then be
If the pressure is kept constant, the second term on the left-hand side is zero, and
The left-hand side is the specific heat at constant pressure of the material.
In general, the infinitesimal quantities are constrained by the equation of state and the specific internal energy function. Namely,
Here denotes the (partial) derivative of the state equation with respect to its argument, keeping the other two arguments fixed, evaluated at the state in question. The other partial derivatives are defined in the same way. These two equations on the four infinitesimal increments normally constrain them to a two-dimensional linear subspace space of possible infinitesimal state changes, that depends on the material and on the state. The constant-volume and constant-pressure changes are only two particular directions in this space.
For any specific volume , denote the function that describes how the pressure varies with the temperature , as allowed by the equation of state, when the specific volume of the material is forcefully kept constant at . Analogously, for any pressure , let be the function that describes how the specific volume varies with the temperature, when the pressure is kept constant at . Namely, those functions are such that
for any values of . In other words, the graphs of and are slices of the surface defined by the state equation, cut by planes of constant and constant , respectively.
Then, from the fundamental thermodynamic relation it follows that
This equation can be rewritten as
both depending on the state .
The heat capacity ratio, or adiabatic index, is the ratio of the heat capacity at constant pressure to heat capacity at constant volume. It is sometimes also known as the isentropic expansion factor.
The path integral Monte Carlo method is a numerical approach for determining the values of heat capacity, based on quantum dynamical principles. However, good approximations can be made for gases in many states using simpler methods outlined below. For many solids composed of relatively heavy atoms (atomic number > iron), at non-cryogenic temperatures, the heat capacity at room temperature approaches 3R = 24.94 joules per kelvin per mole of atoms (Dulong-Petit law, R is the gas constant). Low temperature approximations for both gases and solids at temperatures less than their characteristic Einstein temperatures or Debye temperatures can be made by the methods of Einstein and Debye discussed below.
this equation reduces simply to Mayer's relation:
The differences in heat capacities as defined by the above Mayer relation is only exact for an ideal gas and would be different for any real gas. -->
The adjective specific before the name of an extensive quantity is often used to mean divided by mass.