# Specific heat capacity

The specific heat capacity 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 its temperature. The SI unit of specific heat is joule per kelvin and kilogram, J/(K kg).[1] [2] For example, at a temperature of 25 °C (the specific heat capacity can vary with the temperature), the heat required to raise the temperature of 1 kg of water by 1 K (equivalent to 1 °C) is 4179.6 joules, meaning that the specific heat of water is 4179.6 J·kg−1·K−1.[3]

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, oak wood, and hydrogen gas are about 449, 790, 2400, 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 ${\displaystyle c_{P}}$ and ${\displaystyle c_{V}}$, respectively; their quotient ${\displaystyle \gamma =c_{P}/c_{V}}$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;[4] 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.[5]

## Definition

The specific heat capacity of a substance, usually denoted by ${\displaystyle c}$, is the heat capacity ${\displaystyle C}$ of a sample of the substance, divided by the mass ${\displaystyle M}$ of the sample:[6]

${\displaystyle c={\frac {C}{M}}={\frac {1}{M}}\cdot {\frac {\mathrm {d} Q}{\mathrm {d} T}}}$

where ${\displaystyle \mathrm {d} Q}$ represents the amount of heat needed to uniformly raise the temperature of the sample by a small increment ${\displaystyle \mathrm {d} T}$.

Like the heat capacity of an object, the specific heat of a substance may vary, sometimes substantially, depending on the starting temperature ${\displaystyle T}$ of the sample and the pressure ${\displaystyle p}$ applied to it. Therefore, it should be considered a function ${\displaystyle c(p,T)}$ of those two variables.

These parameters are usually specified when giving the specific heat of a substance. For example, "Water (liquid): ${\displaystyle c_{p}}$ = 4185.5 J/K/kg (15 °C, 101.325 kPa)" [7] When not specified, published values of the specific heat ${\displaystyle c}$ generally are valid for some standard conditions for temperature and pressure.

However, the dependency of ${\displaystyle c}$ 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 ${\displaystyle (p,T)}$, and approximates the specific heat by a constant ${\displaystyle c}$ 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.[8])

### Variations

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 ${\displaystyle p}$ and starting temperature ${\displaystyle T}$. Two particular choices are widely used:

• If the pressure is kept constant (for instance, at the ambient atmospheric pressure), and the sample is allowed to expand, the expansion generates work as the force from the pressure displaces the enclosure or the surrounding fluid. That work must come from the heat energy provided. The specific heat thus obtained is said to be measured at constant pressure (or isobaric), and is often denoted ${\displaystyle c_{p}}$, ${\displaystyle c_{\mathrm {p} }}$, etc.
• On the other hand, if the expansion is prevented — for example by a sufficiently rigid enclosure, or by increasing the external pressure to counteract the internal one — no work is generated, and the heat energy that would have gone into it must instead contribute to the internal energy of the sample, including raising its temperature by an extra amount. The specific heat obtained this way is said to be measured at constant volume (or isochoric) and denoted ${\displaystyle c_{V}}$, ${\displaystyle c_{v}}$ ${\displaystyle c_{\mathrm {v} }}$, etc.

The value of ${\displaystyle c_{V}}$ is usually less than the value of ${\displaystyle c_{p}}$. 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.[9]

### Applicability

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.

## 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.[10][11]

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.

## Units

### International system

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.

The specific heat of a substance (per unit of mass) has dimension L2·T−2·Θ−1, or (L/T)2/Θ. Therefore, the SI unit J/K/kg is equivalent to metre squared per second squared per kelvin (m2 K−1 s−2).

### English (Imperial) engineering units

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 = 0.45459237 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),[12][13] 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.

### Calories

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:

• the "small calorie" (or "gram-calorie", "cal") is 4.184 J, exactly. It was originally defined so that the specific heat of liquid water would be 1 cal/C°/g.
• The "grand calorie" (also "kilocalorie", "kilogram-calorie", or "food calorie"; "kcal" or "Cal") is 1000 small calories, that is, 4184 J, exactly. It was originally defined so that the specific heat of water would be 1 Cal/C°/kg.

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

1 cal/°C/g ("small calorie") = 1 Cal/°C/kg = 1 kcal/°C/kg ("large calorie") = 4184 J/K/kg.

In either unit, the specific heat of water is approximately 1. The combinations cal/°C/kg = 4.144 J/K/kg and kcal/°C/g = 4184,000 J/K/kg do not seem to be widely used.

## Physical basis of specific heat

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.

### Monoatomic gases

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, ${\displaystyle c_{V,\mathrm {m} }=3R/2\approx {}}$12.5 J/K/mol and ${\displaystyle c_{P,\mathrm {m} }=5/2\approx {}}$21 J/K/mol, where ${\displaystyle R\approx {}}$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 ${\displaystyle A}$. That is, approximately,

${\displaystyle c_{V}\approx {}}$12470 J/K/kg${\displaystyle /A\quad \quad \quad c_{p}\approx {}}$20785 J/K/kg${\displaystyle /A}$

For the noble gases, from helium to xenon, these computed values are

 Gas ${\displaystyle A}$ ${\displaystyle c_{V}}$ (J/K/kg) ${\displaystyle c_{p}}$ (J/K/kg) He Ne Ar Kr Xe 4.00 20.17 39.95 83.80 131.29 3118 618.3 312.2 148.8 94.99 5197 1031 520.3 248.0 158.3

### Polyatomic gases

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.[14][15] [16]

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 quantum 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 N
2
at constant volume is ${\displaystyle c_{V,\mathrm {m} }={}}$ 20.6 J/K/mol (at 15 °C, 1 atm), which is 2.49${\displaystyle R}$.[17] 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 ${\displaystyle c_{V}}$ of N
2
(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 ${\displaystyle c_{V}}$ 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.[18] The last value corresponds almost exactly to the predicted value for 7 degrees of freedom per molecule.

## Thermodynamic derivation

In theory, the specific heat of a substance can also be derived from its abstract thermodynamic modeling by an equation of state and an internal energy function.

### State of matter in a homogeneous sample

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 ${\displaystyle M}$. Assume that the evolution of the system is always slow enough for the internal pressure ${\displaystyle P}$ and temperature ${\displaystyle T}$ be considered uniform throughout. The pressure ${\displaystyle P}$ 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 ${\displaystyle T}$, the pressure ${\displaystyle P}$, and its specific volume ${\displaystyle V={\boldsymbol {\mathrm {V} }}/M}$, where ${\displaystyle {\boldsymbol {\mathrm {V} }}}$ is the volume of the sample. (This quantity is the reciprocal ${\displaystyle 1/\rho }$ of the material's density ${\displaystyle \rho =M/{\boldsymbol {\mathrm {V} }}}$.) Like ${\displaystyle T}$ and ${\displaystyle P}$, the specific volume ${\displaystyle V}$ 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: ${\displaystyle F(T,P,V)=0.}$ The function ${\displaystyle F}$ depends on the material under consideration. The specific internal energy stored internally in the sample, per unit of mass, will then be another function ${\displaystyle U(T,P,V)}$ of these state variables, that is also specific of the material. The total internal energy in the sample then will be ${\displaystyle MU(T,P,V)}$.

For some simple materials, like an ideal gas, one can derive from basic theory the equation of state ${\displaystyle F=0}$ and even the specific internal energy ${\displaystyle U}$ In general, these functions must be determined experimentally for each substance.

### Conservation of energy

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 ${\displaystyle M\mathrm {d} U}$ in the total internal energy ${\displaystyle MU}$ must be matched by the net flow of heat energy ${\displaystyle \mathrm {d} Q}$ into the sample, plus any net mechanical energy provided to it by enclosure or surrounding medium on it. The latter is ${\displaystyle -P\mathrm {d} {\boldsymbol {\mathrm {V} }}}$, where ${\displaystyle \mathrm {d} {\boldsymbol {\mathrm {V} }}}$ is the change in the sample's volume in that infinitesimal step.[19] Therefore

${\displaystyle \mathrm {d} Q-P\mathrm {d} {\boldsymbol {\mathrm {V} }}=M\mathrm {d} U}$

hence

${\displaystyle {\frac {\mathrm {d} Q}{M}}-P\mathrm {d} V=\mathrm {d} U}$

If the volume of the sample (hence the specific volume of the material) is kept constant during the injection of the heat amount ${\displaystyle \mathrm {d} Q}$, then the term ${\displaystyle P\mathrm {d} V}$ is zero (no mechanical work is done). Then, dividing by ${\displaystyle \mathrm {d} T}$,

${\displaystyle {\frac {\mathrm {d} Q}{M\mathrm {d} T}}={\frac {\mathrm {d} U}{\mathrm {d} T}}}$

where ${\displaystyle \mathrm {d} T}$ is the change in temperature that resulted from the heat input. The left-hand side is the specific heat at constant volume ${\displaystyle c_{V}}$ of the material.

For the heat capacity at constant pressure, it is useful to define the specific enthalpy of the system as the sum ${\displaystyle H(T,P,V)=U(T,P,V)+PV}$. An infinitesimal change in the specific enthalpy will then be

${\displaystyle \mathrm {d} H=\mathrm {d} U+V\mathrm {d} P+P\mathrm {d} V}$

therefore

${\displaystyle {\frac {\mathrm {d} Q}{M}}+V\mathrm {d} P=\mathrm {d} H}$

If the pressure is kept constant, the second term on the left-hand side is zero, and

${\displaystyle {\frac {\mathrm {d} Q}{M\mathrm {d} T}}={\frac {\mathrm {d} H}{\mathrm {d} T}}}$

The left-hand side is the specific heat at constant pressure ${\displaystyle c_{P}}$ of the material.

### Connection to equation of state

In general, the infinitesimal quantities ${\displaystyle \mathrm {d} T,\mathrm {d} P,\mathrm {d} V,\mathrm {d} U}$ are constrained by the equation of state and the specific internal energy function. Namely,

${\displaystyle \left\{{\begin{array}{lcl}\displaystyle \mathrm {d} T{\frac {\partial F}{\partial T}}(T,P,V)+\mathrm {d} P{\frac {\partial F}{\partial P}}(T,P,V)+\mathrm {d} V{\frac {\partial F}{\partial V}}(T,P,V)&=&0\\[2ex]\displaystyle \mathrm {d} T{\frac {\partial U}{\partial T}}(T,P,V)+\mathrm {d} P{\frac {\partial U}{\partial P}}(T,P,V)+\mathrm {d} V{\frac {\partial U}{\partial V}}(T,P,V)&=&\mathrm {d} U\end{array}}\right.}$

Here ${\displaystyle (\partial F/\partial T)(T,P,V)}$ denotes the (partial) derivative of the state equation ${\displaystyle F}$ with respect to its ${\displaystyle T}$ argument, keeping the other two arguments fixed, evaluated at the state ${\displaystyle (T,P,V)}$ 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.

This analysis also holds no matter how the energy increment ${\displaystyle \mathrm {d} Q}$ is injected into the sample (by heat conduction, irradiation, electromagnetic induction, radioactive decay, etc.

### Relation between heat capacities

For any specific volume ${\displaystyle V}$, denote ${\displaystyle p_{V}(T)}$ the function that describes how the pressure varies with the temperature ${\displaystyle T}$, as allowed by the equation of state, when the specific volume of the material is forcefully kept constant at ${\displaystyle V}$. Analogously, for any pressure ${\displaystyle P}$, let ${\displaystyle v_{P}(T)}$ be the function that describes how the specific volume varies with the temperature, when the pressure is kept constant at ${\displaystyle P}$. Namely, those functions are such that

${\displaystyle F(T,p_{V}(T),V)=0\quad \quad {}}$and${\displaystyle {}\quad \quad F(T,P,v_{P}(T))=0}$

for any values of ${\displaystyle T,P,V}$. In other words, the graphs of ${\displaystyle p_{V}(T)}$ and ${\displaystyle v_{P}(T)}$ are slices of the surface defined by the state equation, cut by planes of constant ${\displaystyle V}$ and constant ${\displaystyle P}$, respectively.

Then, from the fundamental thermodynamic relation it follows that

${\displaystyle c_{P}(T,P,V)-c_{V}(T,P,V)=T\left[{\frac {\mathrm {d} p_{V}}{\mathrm {d} T}}(T)\right]\left[{\frac {\mathrm {d} v_{P}}{\mathrm {d} T}}(T)\right]}$

This equation can be rewritten as

${\displaystyle c_{P}(T,P,V)-c_{V}(T,P,V)=VT{\frac {\alpha ^{2}}{\beta _{T}}},}$

where

${\displaystyle \alpha }$ is the coefficient of thermal expansion,
${\displaystyle \beta _{T}}$ is the isothermal compressibility,

both depending on the state ${\displaystyle (T,P,V)}$.

The heat capacity ratio, or adiabatic index, is the ratio ${\displaystyle c_{P}/c_{V}}$ of the heat capacity at constant pressure to heat capacity at constant volume. It is sometimes also known as the isentropic expansion factor.

## References

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13. Cardarelli, Francois (2012). Scientific Unit Conversion: A Practical Guide to Metrication. M.J. Shields (translation) (2 ed.). Springer. p. 19. ISBN 9781447108054.
14. Feynman, R., Lectures in Physics, vol. I, chapter 40, pp. 7-8
15. Reif, F. (1965). Fundamentals of statistical and thermal physics. McGraw-Hill. pp. 253–254.
16. Charles Kittel; Herbert Kroemer (2000). Thermal physics. Freeman. p. 78. ISBN 978-0-7167-1088-2.
17. Steven T. Thornton and Andrew Rex (1993): Modern Physics for Scientists and Engineers, Saunders College Publishing, 1993
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• Emmerich Wilhelm & Trevor M. Letcher, Eds., 2010, Heat Capacities: Liquids, Solutions and Vapours, Cambridge, U.K.:Royal Society of Chemistry, ISBN 0-85404-176-1. A very recent outline of selected traditional aspects of the title subject, including a recent specialist introduction to its theory, Emmerich Wilhelm, "Heat Capacities: Introduction, Concepts, and Selected Applications" (Chapter 1, pp. 1–27), chapters on traditional and more contemporary experimental methods such as photoacoustic methods, e.g., Jan Thoen & Christ Glorieux, "Photothermal Techniques for Heat Capacities," and chapters on newer research interests, including on the heat capacities of proteins and other polymeric systems (Chs. 16, 15), of liquid crystals (Ch. 17), etc.