| Topic: |
Science > Physics |
| User: |
"" |
| Date: |
15 Nov 2006 08:36:47 PM |
| Object: |
Thermochemistry basics |
I have some questions about thermochemistry (basics) I've read from a
textbook:
1) Textbook says heat and work are the two ways a therymodynamic
system can transfer energy with surroundings. But isn't heat itself
work at the molecular level (work involved with collisions)? Are they
just making the distinction (implicitly) between macroscopic work and
microscopic work?
2) Heat is defined as energy transferred between system and
surroundings as a result of temperature difference. Temperature is
said to be proportional to thermal energy. Thermal energy is said to
be kinetic energy of the random movement of constituent particles.
Transfer of heat from hot object to cold object makes sense, it's
basically the kinetic energies being transferred via collisions.
Everything fits with the above definitions.
But, what if the cold object is near it's condensation point? Then the
hot object is transfering heat to cold object (as kinetic energies
through collisions - by definition of heat transfer) but the result is
not kinetic energy increases in cold object's particles, but rather
potential energy increases (to overcome attractive forces). I can't
see. to match the definitions and explanation of heat transfer to this
situation (which doesn't involve temp. rise in cold object). Can
anyone explain what's going on?
3) Book says an object has 3 possible types of energy: Kinetic of
whole object, Potential of whole object, and Internal energy
(potential+kinetic of particles)
In HS physics, when we deal with potential and kinetic energy, are we
including the internal energies? Like when collisions between billiard
balls are studied in HS physics, is all that theory ignoring the
internal potential energies?
Any help would be appreciated.
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| User: "" |
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| Title: Re: Thermochemistry basics |
15 Nov 2006 10:44:30 PM |
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<lite.on.beta@gmail.com> wrote in message
news:1163644607.240582.78950@h54g2000cwb.googlegroups.com...
I have some questions about thermochemistry (basics) I've read from a
textbook:
1) Textbook says heat and work are the two ways a therymodynamic
system can transfer energy with surroundings. But isn't heat itself
work at the molecular level (work involved with collisions)? Are they
just making the distinction (implicitly) between macroscopic work and
microscopic work?
Basically, yes, although there is one difference. Heat can be transferred in
different ways--either by thermal contact, or through radiation and
absorption of photons of the appropriate energy (generally, infrared), and
it's hard to imagine how one would call this photon thing "work" in the W =
F.d sense.
2) Heat is defined as energy transferred between system and
surroundings as a result of temperature difference. Temperature is
said to be proportional to thermal energy. Thermal energy is said to
be kinetic energy of the random movement of constituent particles.
Transfer of heat from hot object to cold object makes sense, it's
basically the kinetic energies being transferred via collisions.
Everything fits with the above definitions.
But, what if the cold object is near it's condensation point?
Do you mean boiling point? If not, I'm not sure how to interpret this
sentence.
Then the
hot object is transfering heat to cold object (as kinetic energies
through collisions - by definition of heat transfer) but the result is
not kinetic energy increases in cold object's particles, but rather
potential energy increases (to overcome attractive forces).
No, this is kinetic energy, i.e., motion of molecules. First the molecules
move faster and faster in the liquid phase, then as the T approaches the
b.p., some of them have enough kinetic energy to overcome attractive forces,
and they fly off into the vapor phase. I'm not sure I understand your logic
in calling it potential energy. If you mean that the attractive forces
represent potential energy, yes, they do. However, when a molecule consumes
energy to overcome a potential energy, that kinetic energy is lost, and the
temperature still decreases, i.e., the average kinetic energy still
decreases. Think of rolling a ball up a hill. The kinetic energy of the
ball disappears, as it is consumed and converted to potential energy.
I can't
see. to match the definitions and explanation of heat transfer to this
situation (which doesn't involve temp. rise in cold object). Can
anyone explain what's going on?
Think of it this way. You've got a liquid at its boiling point. Some
molecules have KE sufficient to break free of the attractive forces of the
liquid, and fly off into space. This takes energy with them, and since
their energy was above the average kinetic energy of the molecules in the
liquid, losing those high-energy molecules will decrease the average kinetic
energy of the molecules remaining in the liquid phase. This will decrease
the temperature (this is the reason you sweat), or conversely, you can add
more energy to maintain the temperature. As long as it is at the boiling
point, it is going to lose KE in the molecules that evaporate, and more
energy will need to be input. Strictly speaking, the liquid doesn't even
need to be near its boiling point. Take ethanol or isopropanol at room
temperature (that's about 80 degrees C below their boiling points). The get
cold on your skin as they evaporate. It's just that, near the boiling
point, the rate of heat being carried away is much more noticeable. This is
why evaporating ether (bp 35 C, only about 15 degress C above room
temperature) chills your skin more than ethanol. I should point out that
there is nothing magical about the boilling point--it's just the point at
which the vapor pressure of the liquid is equal to the ambient pressure.
When this happens, pockets of this pressure build up in the middle of the
liquid, bubbles form, and the liquid boils. Even if a liquid isn't above
its boiling point, the effect of evaporation is the same--it cools the
liquid, or it requires the input of thermal energy to maintain constant
temperature.
Does this make sense?
3) Book says an object has 3 possible types of energy: Kinetic of
whole object, Potential of whole object, and Internal energy
(potential+kinetic of particles)
In HS physics, when we deal with potential and kinetic energy, are we
including the internal energies?
Generally, no, unless you're specifically measuring temperature changes,
inelasticity of collisions, that sort of thing where whole-object kinetic
energy is transformed into internal energy.
Like when collisions between billiard
balls are studied in HS physics, is all that theory ignoring the
internal potential energies?
Yes, generally. I can't remember if it was in HS or college, but we did an
experiment colliding balls, measuring the KE of the rebounding balls, and
calculating the amount of energy lost as internal energy (i.e, heat). But
unless you're specifically measuring something like that, you're right in
general, macroscopic observations of motion generally neglect internal
energy.
Eric Lucas
Any help would be appreciated.
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| User: "Marvin" |
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| Title: Re: Thermochemistry basics |
16 Nov 2006 11:24:29 AM |
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wrote:
I have some questions about thermochemistry (basics) I've read from a
textbook:
1) Textbook says heat and work are the two ways a therymodynamic
system can transfer energy with surroundings. But isn't heat itself
work at the molecular level (work involved with collisions)? Are they
just making the distinction (implicitly) between macroscopic work and
microscopic work?
Don't confuse clasical thermodynamics and statistical
thermodynamics. The people who developed classical
thermodynamics knew nothing or atoms and molecules, and they
developed a purely macroscopic model.
In Einstein's "wonder year" of 1905, his paper on Brownian
motion showed that macroscopic processes can be described by
interactions between molecules and particles large enough
to see with a microscopic. His is a statistical model. It
is rather like the statistical thermodynamic model.
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| User: "Andy Resnick" |
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| Title: Re: Thermochemistry basics |
16 Nov 2006 12:29:52 PM |
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wrote:
I have some questions about thermochemistry (basics) I've read from a
textbook:
1) Textbook says heat and work are the two ways a therymodynamic
system can transfer energy with surroundings. But isn't heat itself
work at the molecular level (work involved with collisions)? Are they
just making the distinction (implicitly) between macroscopic work and
microscopic work?
Not really. The first law of thermodynamics is simply conservation of
energy, and is dU = Q-W. That is, an isolated system's energy can
change by two different mechanisms: by work W (sometimes called "P-V
energy") or by the flow of heat Q. "Heat" is energy that is unable to
do any mechanical or chemical work, and is related to entropy. While
it's simple to explain the molecular basis for W, I'm not sure there's a
simple, clean, explanation for the molecular basis of Q. Statistical
mechanics is great for large collections of molecules and leads to a
microscopic description of Q.
2) Heat is defined as energy transferred between system and
surroundings as a result of temperature difference. Temperature is
said to be proportional to thermal energy. Thermal energy is said to
be kinetic energy of the random movement of constituent particles.
Yeah, sure... that's close enough.
Transfer of heat from hot object to cold object makes sense, it's
basically the kinetic energies being transferred via collisions.
Everything fits with the above definitions.
Heat can flow either accompanied by material flow or not; keeping this
mental picture of billiard balls bumping around isn't always the best
idea. And besides, kinetic energy can be used to perform work (via the
increase in pressure), so it's not really heat.
But, what if the cold object is near it's condensation point? Then the
hot object is transfering heat to cold object (as kinetic energies
through collisions - by definition of heat transfer) but the result is
not kinetic energy increases in cold object's particles, but rather
potential energy increases (to overcome attractive forces). I can't
see. to match the definitions and explanation of heat transfer to this
situation (which doesn't involve temp. rise in cold object). Can
anyone explain what's going on?
Phase transitions.... like I said, the mental picture of billiard balls
isn't always the best idea. Think in terms of energy flow, and
potential energy in attractive forces, surface tension, hydrogen bonding,
3) Book says an object has 3 possible types of energy: Kinetic of
whole object, Potential of whole object, and Internal energy
(potential+kinetic of particles)
I like to think that an object (a protein, or maybe a smaller molecule,
or even a lecture hall) has "configurational energy"- the energy due to
how things are arranged. That way, I don't need to know all the
details- keeping track of all the amino acid bonds, for example. All I
know is that a given arrangement of constituent parts of the object has
a certain energy associated with it. Move things around, the
configurational energy changes. An example- a room with a billiard ball
somewhere in it. Depending on where the billiard ball is, the ball has
different amounts of (gravitational) potential energy. Or a room filled
with an ideal gas. Depending on where the molecules of gas are placed
in the room, the entropy takes different values. This is like your
"internal energy" concept.
In HS physics, when we deal with potential and kinetic energy, are we
including the internal energies? Like when collisions between billiard
balls are studied in HS physics, is all that theory ignoring the
internal potential energies?
Yes, no... it depends. Internal energy can be used to describe spin of
particles, dipole moments, etc. For your collision experiments, the
internal configuration of each ball remained (mostly) unchanged before
and after the collision, so that energy component was implicitly
subtracted out and not relevant. Contrast that with a receptor-ligand
reaction, where a protein may undergo a large change in configuration
("shape") in response to binding. There, the internal energy changes
are important to keep track of.
--
Andrew Resnick, Ph.D.
Department of Physiology and Biophysics
Case Western Reserve University
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| User: "" |
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| Title: Re: Thermochemistry basics |
16 Nov 2006 12:52:25 AM |
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In article <1163644607.240582.78950@h54g2000cwb.googlegroups.com>, writes:
I have some questions about thermochemistry (basics) I've read from a
textbook:
1) Textbook says heat and work are the two ways a therymodynamic
system can transfer energy with surroundings. But isn't heat itself
work at the molecular level (work involved with collisions)? Are they
just making the distinction (implicitly) between macroscopic work and
microscopic work?
In fact, yes. You've a system with an insanely large number of
degrees of freedom, you separate the few large scale degrees of
freedom you care about and deal with them explicitly, while throwing
everything else into a large bag labeled "heat". Of course, it helps
that the disparity of magnitudes between the "macro" and "micro" is
such that the separation is not very sensitive to where you set the
boundary. But, "macro" and "micro" is a matter of context. When you
happen to model galactic collisions, you'll treat the translational
and rotational motion of a galaxy as a whole as "mcroscopic" while
energy transfer at the level of individual solar systems of even star
clusters will fall under "heat". Of course, the denizens of said
solar systems will view it differently:-)
2) Heat is defined as energy transferred between system and
surroundings as a result of temperature difference. Temperature is
said to be proportional to thermal energy. Thermal energy is said to
be kinetic energy of the random movement of constituent particles.
No, not just kinetic. That's a common (way to common) error in
most introductory texts, stemming no doubt from the fact that the most
asic system dealt with in stat mechanics is the "ideal gas" where the
only microscopic degrees of freedom are kinetic. But thermal energy
encompasses the energy stored in all microscopic degrees of freedom,
so in general kinetic as well as potential are involved.
Transfer of heat from hot object to cold object makes sense, it's
basically the kinetic energies being transferred via collisions.
Everything fits with the above definitions.
But, what if the cold object is near it's condensation point? Then the
hot object is transfering heat to cold object (as kinetic energies
through collisions - by definition of heat transfer) but the result is
not kinetic energy increases in cold object's particles, but rather
potential energy increases (to overcome attractive forces). I can't
see. to match the definitions and explanation of heat transfer to this
situation (which doesn't involve temp. rise in cold object). Can
anyone explain what's going on?
I hope the above answers it.
3) Book says an object has 3 possible types of energy: Kinetic of
whole object, Potential of whole object, and Internal energy
(potential+kinetic of particles)
In HS physics, when we deal with potential and kinetic energy, are we
including the internal energies?
In general, no, except for cases where it clearly changes.
Like when collisions between billiard
balls are studied in HS physics, is all that theory ignoring the
internal potential energies?
Yes. And internal kinetic as well.
Mati Meron | "When you argue with a fool,
meron@cars.uchicago.edu | chances are he is doing just the same"
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