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HEAT1=HEAT2 ?



The last two paragraphs are comments on what Leigh wrote in July. But first
I want to elaborate on the observation made by Jim Green, also in July of 1997.

Words *do* matter. Of course, one can use *any* word for anything, but
the object is to be understood not to argue over the word. But if the same
word is used for multiple things, understanding is hampered - if not
destroyed. For example, if the word "heat" is used for what is *inside* of
a system *and* for what is done to the system, no one will understand.
............................................................................
Yes, the word heat has two conflicting meanings in many introductory physics
textbooks. In thermodynamics it refers to that part of a thermal energy change
which is due to a difference of temperatures, dT, between the inside and
outside of a system. It is a path-dependent quantity, usually expressed in
joules. In elementary physics on the other hand, heat is the name given to
the quantity Q when the formula Q=c*m*dT is used to perform simple
calorimetric calculations. Students often say that the temperature of a body
is determined by the amount of heat "it contains". Phrases such as "heat added
to a body" or "heat removed from it" are commonly used by teachers who take
it for granted that "heat" is a synonym for "thermal energy of molecules and
atoms". The experiments of Joule are often interpreted by saying that
"energy in the form of heat" and "mechanical energy" are equivalent. In
thermodynamics, on the other hand, we emphasize that heat is not a state
function, such as energy or enthalpy.

Thus a common statement: "heat is a form of energy" is false in thermodynamics.
This situation was created because the same name, "heat", was given to
different physical quantities expressed in joules. We use it in calorimetry
and we use it in the first law, dE=Q+W. The recognition of this fact is an
important first step toward the elimination of many conceptual conflicts. We
must then decide what to do about this unfortunate situation. Should the term
"heat" be eliminated from elementary calorimetry and replaced by "internal
energy"? I suspect that most phys-L-ers would favor this "modern" approach.
They would disagree with the following formulation, found in Sears and
Zemansky. "The process of combustion releases the internal energy and converts
it into heat. In this form the energy can be utilized for ..."

The other alternative is to retain the traditional meaning of heat (form of
energy) and to invent a new name, for example, "thermal pseudo-energy", for
the path-dependent quantity Q in thermodynamics. This approach could be
defended by observing that "heat" is a common word; introductory courses have
traditionally been structured to quantify common words, such as force, work
and heat. Phrases, such as "energy in the form of heat", "heat released in a
reaction", "heat produced through friction" or "heat flows" are too deeply
rooted to be abandoned. Renaming Q in thermodynamics would be less confusing
than renaming it in elementary physics. Wouldn't you agree with Jim that the
most effective way of confusing students and teachers is to redefine
traditional words with which they are already familiar. Can students learn
thermodynamics before they learn calorimetry?

By the way, I still think that a distinction between thermal energy and
internal energy can be very useful in many problems. Thermal energy is that
part of internal energy which is associated with motions and interactions of
molecules and atoms. The internal energy, on the other hand, is not limited
to its thermal component; part of it may be associated with motions and
interactions of macroscopic components of a system, such as wheels, belts,
springs and pistons. The internal energy does not change when friction slows
down a brick sliding along a horizontal surface. But a conversion of kinetic
energy into thermal energy does take place in the system.
Ludwik Kowalski
............................................................................
On July 13, 1997 Leigh Palmer commented on "thermal energy". He wrote:
This term is not in my lexicon, nor should it be in that of any physics
teacher. It is a source of confusion and ... a barrier to conceptual grasp.
If you will define the term as a function of the parameters which describe
the state of the system ... I will then be able to discuss .... Note that
I am not asking for a semantic clarification here; I am asking for a
physical relation.

The source of confusion is in the meaning of the word heat. Thermal energy,
U, can easily be defined in terms of parameters of a particular system. In
the absence of macroscopic components U=E, where E is the internal energy.
Otherwise U=E-X, where X stands for the sum of all the macroscopic energy
terms, such as 0.5*m*v^2 for the kinetic energy of a sliding block, etc. The
problem, as stated by John Mallinckrodt, is to decide which components are
macroscopic and which are microscopic. I would say that, in general, a
component (subsystem) is macroscopic when the fluctuational part of its
energy is negligibly small in comparison with its non-fluctuating part.
A distinction between macroscopic and microscopic subsystems (and modes) is
useful. Only a macroscopic subsystem can be characterized by a temperature.

By the way, I still do not know why I should reject your verbal description
of the first law, Leigh. You wrote that nobody should be satisfied with it.

The change in thermal energy of a system is a sum of the heat that flows
into it and that part of the work that is done on it that does not appear
as a change in the kinetic energy of the center of mass of the system.

I know that you would prefer "internal energy" instead of "thermal energy"
"heating" instead of "flow of heat", and "working" instead of "work done".
But what really is wrong in the above?
Ludwik Kowalski