Енталпија

Енталпија је мера тоталне енергије термодинамичког система. Она се састоји од унутрашње енергије, која је енергија неопходна да би се формирао систем, и количине енергије потребне да би се обезбедио простор за систем потискивањем његове околине, и успоставили његова запремина и притисак. It is a state function used in many measurements in chemical, biological, and physical systems at a constant pressure, which is conveniently provided by the large ambient atmosphere. The pressure–volume term expresses the work required to establish the system's physical dimensions, i.e. to make room for it by displacing its surroundings. The pressure-volume term is very small for solids and liquids at common conditions, and fairly small for gases. Therefore, enthalpy is a stand-in for energy in chemical systems; bond, lattice, solvation and other "energies" in chemistry are actually enthalpy differences. As a state function, enthalpy depends only on the final configuration of internal energy, pressure, and volume, not on the path taken to achieve it.

Топлота (енталпија) стварања једињења се односи на један мол супстанце на температури од 25°C и притиску од 101,3 kPa. Под топлотом (енталпијом) стварања подразумева се топлотни ефекат реакције при којој настаје 1 mol сложене супстанце из простих супстанци. Укупна промена енталпије је стална и не зависи од тога да ли се реакција одиграва у једном степену или више степена, ако се полази од истих компонената и добијају исти производи реакције са истим агрегатним стањем. Топлотни ефекат реакције (или промена енталпије реакције) једнак је алгебарском збиру енталпије стварања производа умањеном за суму стварања полазних супстанци (реактаната).

The unit of measurement for enthalpy in the International System of Units (SI) is the joule. Other historical conventional units still in use include the calorie and the British thermal unit (BTU).

The total enthalpy of a system cannot be measured directly because the internal energy contains components that are unknown, not easily accessible, or are not of interest in thermodynamics. In practice, a change in enthalpy is the preferred expression for measurements at constant pressure, because it simplifies the description of energy transfer. When transfer of matter into or out of the system is also prevented and no electrical or shaft work is done, at constant pressure the enthalpy change equals the energy exchanged with the environment by heat.

In chemistry, the standard enthalpy of reaction is the enthalpy change when reactants in their standard states (p = 1 bar; usually T = 298 K) change to products in their standard states. This quantity is the standard heat of reaction at constant pressure and temperature, but it can be measured by calorimetric methods even if the temperature does vary during the measurement, provided that the initial and final pressure and temperature correspond to the standard state. The value does not depend on the path from initial to final state since enthalpy is a state function.

Enthalpies of chemical substances are usually listed for 1 bar (100 kPa) pressure as a standard state. The temperature does not have to be specified, but tables generally list the standard heat of formation at 25 °C (298 K). For endothermic (heat-absorbing) processes, the change ΔH is a positive value; for exothermic (heat-releasing) processes it is negative.

The enthalpy of an ideal gas is independent of its pressure or volume, and depends only on its temperature, which correlates to its thermal energy. Real gases at common temperatures and pressures often closely approximate this behavior, which simplifies practical thermodynamic design and analysis.

Дефиниција

The enthalpy H of a thermodynamic system is defined as the sum of its internal energy and the product of its pressure and volume:

H = U + pV,

where U is the internal energy, p is pressure, and V is the volume of the system.

Enthalpy is an extensive property; it is proportional to the size of the system (for homogeneous systems). As intensive properties, the specific enthalpy h = H/m is referenced to a unit of mass m of the system, and the molar enthalpy H_m is H/n, where n is the number of moles. For inhomogeneous systems the enthalpy is the sum of the enthalpies of the composing subsystems:

$H=\sum _{k}H_{k},$

where

H is the total enthalpy of all the subsystems,

k refers to the various subsystems,

H_k refers to the enthalpy of each subsystem.

A closed system may lie in thermodynamic equilibrium in a static gravitational field, so that its pressure p varies continuously with altitude, while, because of the equilibrium requirement, its temperature T is invariant with altitude. (Correspondingly, the system's gravitational potential energy density also varies with altitude.) Then the enthalpy summation becomes an integral:

$H=\int (\rho h)\,dV,$

where

ρ ("rho") is density (mass per unit volume),

h is the specific enthalpy (enthalpy per unit mass),

(ρh) represents the enthalpy density (enthalpy per unit volume),

dV denotes an infinitesimally small element of volume within the system, for example, the volume of an infinitesimally thin horizontal layer,

the integral therefore represents the sum of the enthalpies of all the elements of the volume.

The enthalpy of a closed homogeneous system is its energy function H(S,p), with its entropy S[p] and its pressure p as natural state variables. A differential relation for it can be derived as follows. We start from the first law of thermodynamics for closed systems for an infinitesimal process:

$dU=\delta Q-\delta W,$

where

𝛿Q is a small amount of heat added to the system,

𝛿W is a small amount of work performed by the system.

In a homogeneous system in which only reversible processes or pure heat transfer are considered, the second law of thermodynamics gives 𝛿Q = T dS, with T the absolute temperature and dS the infinitesimal change in entropy S of the system. Furthermore, if only pV work is done, 𝛿W = p dV. As a result,

$dU=T\,dS-p\,dV.$

Adding d(pV) to both sides of this expression gives

$dU+d(pV)=T\,dS-p\,dV+d(pV),$

or

$d(U+pV)=T\,dS+V\,dp.$

So

$dH(S,p)=T\,dS+V\,dp.$

Карактеристичне функције

The enthalpy, H(S[p], p, {N_i}), expresses the thermodynamics of a system in the energy representation. As a function of state, its arguments include both one intensive and several extensive state variables. The state variables S[p], p, and {N_i} are said to be the natural state variables in this representation. They are suitable for describing processes in which they are determined by factors in the surroundings. For example, when a virtual parcel of atmospheric air moves to a different altitude, the pressure surrounding it changes, and the process is often so rapid that there is too little time for heat transfer. This is the basis of the so-called adiabatic approximation that is used in meteorology.

Conjugate with the enthalpy, with these arguments, the other characteristic function of state of a thermodynamic system is its entropy, as a function, S[p](H, p, {N_i}), of the same list of variables of state, except that the entropy, S[p], is replaced in the list by the enthalpy, H. It expresses the entropy representation. The state variables H, p, and {N_i} are said to be the natural state variables in this representation. They are suitable for describing processes in which they are experimentally controlled. For example, H and p can be controlled by allowing heat transfer, and by varying only the external pressure on the piston that sets the volume of the system.

Физичко тумачење

The U term is the energy of the system, and the pV term can be interpreted as the work that would be required to "make room" for the system if the pressure of the environment remained constant. When a system, for example, n moles of a gas of volume V at pressure p and temperature T, is created or brought to its present state from absolute zero, energy must be supplied equal to its internal energy U plus pV, where pV is the work done in pushing against the ambient (atmospheric) pressure.

In physics and statistical mechanics it may be more interesting to study the internal properties of a constant-volume system and therefore the internal energy is used. In chemistry, experiments are often conducted at constant atmospheric pressure, and the pressure–volume work represents a small, well-defined energy exchange with the atmosphere, so that ΔH is the appropriate expression for the heat of reaction. For a heat engine, the change in its enthalpy after a full cycle is equal to zero, since the final and initial state are equal.

Однос са топлотом

In order to discuss the relation between the enthalpy increase and heat supply, we return to the first law for closed systems, with the physics sign convention: dU = δQ − δW, where the heat δQ is supplied by conduction, radiation, Joule heating, or friction from stirring by a shaft with paddles or by an externally driven magnetic field acting on an internal rotor (which is surroundings-based work, but contributes to system-based heat). We apply it to the special case with a constant pressure at the surface. In this case the work is given by p dV (where p is the pressure at the surface, dV is the increase of the volume of the system). Cases of long range electromagnetic interaction require further state variables in their formulation, and are not considered here. In this case the first law reads:

$dU=\delta Q-p\,dV.$

Now,

$dH=dU+d(pV).$

So

$dH=\delta Q+V\,dp+p\,dV-p\,dV$

$\,\,=\delta Q+V\,dp.$

If the system is under constant pressure, dp = 0 and consequently, the increase in enthalpy of the system is equal to the heat added or given off:

$dH=\delta Q.$

This is why the now-obsolete term heat content was used in the 19th century.

Литература

• Dalton, J.P. (1909). „Researches on the Joule–Kelvin effect, especially at low temperatures. I. Calculations for hydrogen” (PDF). KNAW Proceedings. 11: 863—873. Bibcode:1908KNAB...11..863D. 
• Haase, R. (1971). Jost, W., ур. Physical Chemistry: An Advanced Treatise. New York: Academic. стр. 29. 
• Gibbs, J. W. The Collected Works of J. Willard Gibbs, Vol. I (1948 изд.). New Haven, CT: Yale University Press. стр. 88. 
• Howard, I. K. (2002). „H Is for Enthalpy, Thanks to Heike Kamerlingh Onnes and Alfred W. Porter”. J. Chem. Educ. 79 (6): 697—698. Bibcode:2002JChEd..79..697H. doi:10.1021/ed079p697. 
• Laidler, K. (1995). The World of Physical Chemistry. Oxford: Oxford University Press. стр. 110. 
• Kittel, C.; Kroemer, H. (1980). Thermal Physics. New York: S. R. Furphy & Co. стр. 246. 
• DeHoff, R. (2006). Thermodynamics in Materials Science. CRC Press. ISBN 9780849340659. 
• Guggenheim, E. A. (1959). Thermodynamics. Amsterdam: North-Holland Publishing Company. 
• Moran, M. J.; Shapiro, H. N. (2006). Fundamentals of Engineering Thermodynamics (5th изд.). John Wiley & Sons. стр. 511. ISBN 9780470030370. 

Спољашње везе

• Енталпија
• Први закон термодинамике
• Пример израчунавања енталпије