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How is enthalpy not directily measurable?

  1. Apr 14, 2015 #1
    I keep hearing that enthalpy is not directly measurable and that on it's own it carries no physical signifigance. But if you have a gas in a container for example, it has some internal energy which I'm assuming is measurable (a least in principle), and you can also measure its pressure as well as its volume. So how is it not measuable? Also, I've been told that enthalpy basically gives you the total energy possesed by a system at any point in time and that it's equivalent to the energy required to create the system. So there is some physical signifigance to enthalpy on its own? I'm so confused
  2. jcsd
  3. Apr 14, 2015 #2
    Good question (meaning I don't know the answer but wish I did)...I believe it has to do with the fact it is dependent on the "internal energy" of the system, which has to be defined relative to some reference.

    It's surely complicated, but it would be nice to have description if one were possible of the basic issue: My guess trying to read through the wiki's is that because the "state functions" of a system, which try to apply a "measure" of energy content to the internals of a thing, don't really have an obvious reference of measure, other than at some very distant point - the flow of energy over time in the universe. So we define it in terms of that flow, or "change in state over time".

    Hopefully someone more knowledgeable will chime in and correct that.

    From Wikipedia, the free encyclopedia
    Enthalpy is defined as a thermodynamic potential, designated by the letter "H", that consists of the internal energy of the system (U) plus the product of pressure(p) and volume (V) of the system:[1]


    Since U, p and V are all functions of the state of the thermodynamic system, enthalpy is a state function.

    The unit of measurement for enthalpy in the International System of Units (SI) is the joule, but other historical, conventional units are still in use, such as theBritish thermal unit and the calorie.

    The total enthalpy, H, of a system cannot be measured directly. The same situation exists in classical mechanics: only a change or difference in energy carries physical meaning. Enthalpy itself is a thermodynamic potential, so in order to measure the enthalpy of a system, we must refer to a defined reference point; therefore what we measure is the change in enthalpy, ΔH. The change ΔH is positive in endothermic reactions, and negative in heat-releasing exothermicprocesses.

    Internal energy
    From Wikipedia, the free encyclopedia
    In thermodynamics, the internal energy is one of the two cardinal state functions of the state variables of a thermodynamic system. It refers to energy contained within the system, while excluding the kinetic energy of motion of the system as a whole and the potential energy of the system as a whole due to external force fields. It keeps account of the gains and losses of energy of the system.

    The internal energy of a system can be changed by (1) heating the system, or (2) by doing work on it, or (3) by adding or taking away matter.[1] When matter transfer is prevented by impermeable walls containing the system, it is said to be closed. Then the first law of thermodynamics states that the increase in internal energy is equal to the total heat added and work done on the system by the surroundings. If the containing walls pass neither matter nor energy, the system is said to be isolated. Then its internal energy cannot change.

    The internal energy of a given state of a system cannot be directly measured. It is determined through some convenient chain of thermodynamic operations andthermodynamic processes by which the given state can be prepared, starting with a reference state which is customarily assigned a reference value for its internal energy. Such a chain, or path, can be theoretically described by certain extensive state variables of the system, namely, its entropy, S, its volume, V, and its mole numbers, {Nj}. The internal energy, U(S,V,{Nj}), is a function of those. Sometimes, to that list are appended other extensive state variables, for example electric dipole moment. For practical considerations in thermodynamics and engineering it is rarely necessary or convenient to consider all energies belonging to the total intrinsic energy of a system, such as the energy given by the equivalence of mass. Typically, descriptions only include components relevant to the system and processes under study. Thermodynamics is chiefly concerned only with changes in the internal energy.
  4. Apr 14, 2015 #3
    Your assumption that the internal energy can be directly measured is not correct. The internal energy represents the total energy of the material, including random kinetic energy of the molecules, individual molecular rotations and vibrations, and attraction/repulsion energy of the molecules with one another. The internal energy can be measured only relative to a change from one thermodynamic equilibrium state to another. Often, the internal energy is referenced to a specified reference state of T and P, T and V, or P and V (where V is specific volume), and is taken to be zero in that reference state.
    This is trying to ascribe to the enthalpy some fundamental characteristics that it does not possess. The important fundamental physical property is the internal energy U (which we just discussed above). The enthalpy, defined as U + PV, is just a convenient parameter to work with in many types of thermodynamic analyses. It is particularly useful for systems at constant pressure, and for open systems in which material enters and/or leaves. As you noted, like the internal energy U, the enthalpy H is a function of state.

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