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Course: High school chemistry > Unit 6
Lesson 1: Thermal energy and equilibriumFirst and second laws of thermodynamics
The first law of thermodynamics states that if a system gains/loses some amount of energy, the surroundings must lose/gain the same amount of energy. The second law of thermodynamics states that entropy does not spontaneously decrease. Entropy is a measure of how spread out energy is. As an example, if two systems at different temperatures are in contact, heat will flow from the warmer system to the cooler system until they reach thermal equilibrium (equal temperatures). Created by Mahesh Shenoy.
Want to join the conversation?
- can the first law of thermodynamics be violated.(2 votes)
- could the second law just be thought of as like osmosis in biology? like for example is entropy just meaning like concentration? with low entropy meaning highly concentrated and opposite with low entropy?(1 vote)
- Mr. Shenoy's description seems to define entropy as the quality or state of being spread-out. Yes, this can be compared to osmosis. Think of it like this: objects "seek" to be the same temperature as their surroundings. This is just like osmosis: both cells and kinetic energy get less concentrated over time as they cross their "borders", leading to equilibrium.
(It is important to note heat capacity, which dictates how much effort it takes to raise the temperature of an object. This is why aluminum foil is much cooler than the average baking sheet in the oven)
I hope it helps!(1 vote)
Video transcript
- [Presenter] If you
take a very hot coffee, say, in a thermo flask, and keep it in a room, then you know that that coffee will automatically start
cooling down all by itself until it reaches its
room temperature, right? But my question is, why
can't the reverses happen? Why can't a room-temperature coffee automatically, all by itself, heat up? I know that sounds like a silly question, because we don't see that happening, but why not, why doesn't this happen? It's questions like these
that led us to discovering one of the most profound
principles of thermodynamics, and that's exactly what we're
gonna learn in this video. Now, before we start, because there's going to be a lot of energy transfer happening, it's important for us to be sure that we are talking about the same thing, and one of the ways to do that is to define something called a system and a surrounding, okay? So what exactly is a system? A system is basically a group
of stuff that interests us. For example, if I draw a
boundary around this coffee and call my coffee as a system, then all the molecules of that coffee now become part of my system, and everything that is
outside of that boundary, like, say, the molecules of
air inside the thermo flask, the molecules of the thermo flask itself, and the table, and the
room, and everything else, becomes the surrounding. So you get it? So you have a system, you have a boundary, and then everything
outside of that boundary is the surrounding. The boundary is imaginary, of course, but we can define our
system however we want. For example, another way I
could have defined my system is by considering a boundary over here. Now I would say that, look, everything inside the thermo flask, the coffee molecules, the
molecules of the air over here, including the thermo flask itself, is now my system, that's
my objects of interest, and everything outside of that now becomes the surrounding, okay? So we are completely free to
define what our system is. And just to take another example, I could have also defined my system as a part of coffee and a part of the thermo
flask, as I've shown over here, and again, everything else
becomes the surrounding. Now, this may not be the
most useful one for us, but we can choose whatever we want, okay? Now, for the rest of this video, let us define this to be our
system and the surrounding. Let's consider the thermo flask and all the molecules
inside of the thermo flask, all of that will be our system, and everything else
will be our surrounding. All right, now let's look
at what really happened when our coffee cooled down. Well, initially, our system
is at very high temperature. This means the particles of that system, basically, the molecules of the coffee and all of these over here, they have a very high
average kinetic energy. Remember, that's what
temperature is, right? But later on, our system cooled down, our system's temperature decreased. That means the average kinetic
energy of the particles must have become smaller. This means this system lost some energy. Where did it go? Well, you can probably guess it. That must have flown into the surrounding. And here's the thing. If our system lost some amount of energy, the surrounding must have gained exactly that same amount of energy, right, because energy can neither
be created nor destroyed, isn't it? Guess what? That itself is what we call the first law of thermodynamics
(chuckling), okay? First law of thermodynamics is basically energy conservation. You cannot destroy or create energy, energy must always be accounted for. So in our case, when it comes
to system and surrounding, that basically means that energy lost or gained by the system should exactly equal the
energy gained or lost by the surrounding. Now, before we move ahead, one question that we could have is when the system lost energy, its temperature dropped significantly. Now, the surrounding has gained
the same amount of energy, so its temperature should
shoot up significantly, right, so why doesn't the room
just become hotter? Well, the short answer is, yes, the surrounding has gained
the same amount of energy, but remember, the surrounding
has way more molecules compared to the system. You have so many more molecules, and because, remember,
the surrounding represents the air molecules over here,
the molecules of this room, and so much more, since our surrounding
has way more molecules compared to the system, the energy, when you distribute it over all of these molecules, the average kinetic energy gained, that is almost negligible. As a result of that, the temperature of the
surrounding hardly changes. All right (chuckling), now that brings us to our main question. If you have a high temperature
system, like our hot coffee, it can automatically cool down by transferring energy to its surrounding. Why can't the reverse happen? Why can't energy go from the surrounding to our room temperature system and increase its temperature? The question to really ponder upon is, does this violate the first
law of thermodynamics? Why don't you pause the video
and just think about it. All right, let's see. As long as we made sure that the energy gained by my system is exactly equal to the energy
lost by the surrounding, we are done. The first law of thermodynamics
has no problems with it. So you see, if we just think from the first law of thermodynamics, a hot coffee can cool down all by itself, and the reverse must also be possible. That's why we say this
is not a silly question. Now, since we know that
this does not happen, that means there must be
something else that's going on. There must be another law that could be preventing
that from happening. What is that? That is the second law of thermodynamics, and that says that entropy
cannot decrease spontaneously. Now, I know this again
brings up a lot of questions, so let's first try to
understand what entropy is. There are many ways to
think about entropy, but the way I like to think about it is it's a measure of how much
your energy has spread out. So what does this mean? To understand that, again,
let's consider a new system. This time, let's consider the entire room and
everything inside of it as our system, and let's assume that
this system is isolated from its surrounding. Everything outside the room
now becomes the surrounding, and let's assume that it's isolated. That means, let's say that, you know, there is no energy transfer between the system and the surrounding, we have insulated the whole thing, which is not really possible, but let's assume that, okay? This means whatever
happens inside my system, the energy must stay within
the system, all right? The energy cannot escape anywhere. That's what we are assuming. All right, so now if we go
back to initial conditions where the coffee was very hot, yeah, this particular coffee was very hot, then look, that energy was
concentrated over here, there was a lot of
concentration of energy, and therefore we would say the entropy of our system was high. But then, then, once
the coffee cooled down, the total energy of my new
system, this entire system, has stayed the same, right, because the energy just went from the coffee and the thermo flask into the room. The energy has not changed, but the energy is now
more spread out, right? Therefore, now, the entropy has increased. So again, this means we
started with low entropy, because in the beginning, we had some concentration of energy, and then we went towards high entropy because the energy got more spread out. More spread out the energy gets, higher, the entropy becomes. Okay, so now, what is
the second law saying? The second law says the entropy cannot decrease spontaneously. In other words, the energy cannot get
concentrated spontaneously. Energy can spread out spontaneously, that can happen, but it cannot get
concentrated spontaneously. That's what our second law
of thermodynamics says. So let's see if we can
apply this over here now. So right now, if you
consider the situation, and if you look at the entropy right now, the energy is pretty spread out. Compared to that, or initially, the energy was slightly more
concentrated in the coffee. As a result, we started with a slightly lower entropy system, and as time passed by,
the entropy increased, the energy got more spread out. This happens until the temperature stays, temperature becomes equal. Once that happens, we say
thermal equilibrium has reached. That means after that, there won't be any
significant energy flowing in or any energy flowing out, at least at a macroscopic scale. So we have reached a thermal equilibrium, and when that has happened,
the energy has spread out, as a result, the entropy has increased. That can happen according to the second
law of thermodynamics. But why can't the reverse happen? Why can't the energy from our room just enter into our coffee? Well, if that happened, then look, the energy would get more concentrated, and therefore, the energy would now, the entropy would now decrease, and the second law says that
cannot happen spontaneously. Entropy cannot decrease spontaneously. That's why the coffee, a
room-temperature coffee, cannot spontaneously become hotter. What I find fascinating is
that we take this for granted. I mean, we know that this cannot happen, but the reason is entropy. It's not that straightforward (chuckling). That is pretty cool if
you think about it, right? And what's also cool is that this is the reason
why heat always flows from a hot body to a
cold body, spontaneously, because that's how the entropy increases, and that's allowed, but the reverse cannot happen. If the heat were to flow spontaneously from a cold body to a hot body, then that would violate the
second law of thermodynamics. The entropy would decrease. That is not allowed. But anyways, this brings
us to the last question. What about refrigerators? If you think about it, the inside of a
refrigerator is pretty cold, and since the heat is flowing
out of the refrigerator all the time, that means the heat is
flowing from the cold region to a warm region, warmer
region, a hotter region, so that's exactly
opposite of what we said. So does that break the
second law of thermodynamics? No, because the key word
over here is spontaneously. We said this cannot happen spontaneously. A refrigerator does not
do this spontaneously, it does it by using electrical energy. There is a heat pump over there
which runs on electricity, and from using electricity,
it is pumping the heat out. So look, it's not happening all by itself. It's using electricity to do
it, therefore, that's fine. It's not a spontaneous process, so it's not breaking the second law. You can imagine, if the power goes off, then the heat will start flowing back in and all your food will get spoiled. But now you might say,
well, wait a second, what about entropy of the system? Well, now we need to be careful, because look, since we
are drawing electricity, that means this system is no
longer an isolated system. So an easier way to
think about it would be instead of plugging it to our socket, let's say the refrigerator
was hooked up to a generator, and now we can make sure we still have a completely
isolated system, we don't need any energy from the outside. This is an electric generator, let's say, it uses diesel and it
converts into electricity. Now, what would happen? Well, now think about it. Before we switch on the generator, the generator has some diesel in it. Diesel contains some energy. This is not thermal energy,
it contains chemical energy, but it is concentrated energy. That means, initially, there was a pocket of low entropy in this entire system, right? Okay, now what happens once we
start running the generator? What happens to that low
entropy chemical energy inside the diesel? Well, it eventually goes out as heat and ends up as thermal energy of, thermal energy of all
the particles over here, which means the energy has spread out, and as a result, look, the
entropy has eventually increased. Therefore, if you correctly account for all the energy sources, all of the energy within a system, you will find that the
entropy will never decrease. Sure, it's possible to
actively pump out heat and make sure that the entropy
of one part of a system becomes lower, but as a consequence, you will find that the entropy of some
other part of the system will always increase to ensure that the entropy
of the total system never becomes smaller, it will only, it can stay the same, or
it can only become larger. There is no way (chuckling) to violate the second
law of thermodynamics.