- Thermodynamics questions
- Thermodynamics article
- Specific heat and latent heat of fusion and vaporization
- Zeroth law of thermodynamics
- First law of thermodynamics
- First law of thermodynamics problem solving
- PV diagrams - part 1: Work and isobaric processes
- PV diagrams - part 2: Isothermal, isometric, adiabatic processes
- Second law of thermodynamics
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- What are adibatic isothermal isochoric(3 votes)
- "Adiabatic", "isothermal", and "isochoric" refer to different processes involving manipulation of enclosed ideal gases.
An adiabatic process refers to a situation where a gas is manipulated without losing or gaining heat energy. This mean you can compress or expand the gas, and the pressure and temperature can change, but no heat is transferred into or out of the system. To do this, you must completely insulate your system, and the work done on or by the gas must be done quickly enough so that the heat energy has no time to escape.
An isothermal process is a process where a gas is manipulated in such a way that its temperature stays the same. This means that you can compress or expand the gas to increase or decrease the pressure, but it must be done really slowly so that the temperature will always stay in equilibrium with the gas' surroundings, or otherwise at such a rate that the temperature of the gas stays the same at any point in the process.
An isochoric process is a process where a gas is manipulated in such a way that its volume does not change. This means you can add or remove heat energy to change the temperature of the gas, which in turn changes the pressure of the gas inside the container.
There is also a fourth process called an isobaric process, in which the pressure stays the same. This is achieved by letting the gas expand if the temperature increases, by increasing the volume of the container, or by reducing the volume if the gas' temperature decreases.
I suggest you watch the videos on PV diagrams later in this list if you have yet to do so, to get a better explanation. :)(9 votes)
- In zeorth law of thermodynamics objects 1and 2are made in contact and they are in thermal equilibrium. When second object contacts with third its temperature should decrease why it is same as 1 object(1 vote)
- The Zeroth Law actually states that if object 1 is in thermal equilibrium with objects 2 and 3, then 2 and 3 are also in thermal equilibrium(6 votes)
- is value of U always positive? or is there any process where it can be zero or negative?(2 votes)
- U is always positive, because it's not possible to have less-than-zero-energy. Delta U (change in U) however can be positive, negative, or zero, as it represents the change in energy.(3 votes)
- what happens when work is done on gas molecules?
Also if the collision were to move many gas molecules, could the gas molecules collide into each other eventually cancelling out the energy that all the gas molecules have
Im so confused to what to imagine when work is done on a gas(2 votes)
- you can't cancel the energy of gas because when you try to the other gas molecules will transfer their energy, to that other gas molecule.(1 vote)
- Molecules are also subject to intermolecular interactions (repulsive or attractive forces/energies).....Anyway excellent video...Thank you...(2 votes)
- Is this physics or chemistry?(2 votes)
- this is physics when you refer to plain energy, but chemistry when you talk about thermal energy.(1 vote)
- Why is heat represented by a Q?(1 vote)
- Q came from the calorist's theories of heat being a quality of matter. Q means quantity. Even though the theory of heat was disproven, the term "calorie" and the symbol q remained.In the early days of thermodynamics, the nature of heat was not well understood, and it took some years of experiments and learned discussion to establish exactly what 'heat' was. It is possible that the symbol Q was used to denote a catch-all 'quantity' until its characteristics had been better determined. It appears that, for a very long time, heat was considered to a material that existed in all bodies and materials, with a tendency to flow from areas that were hot to those that were cooler. Before the development of thermodynamics, the then current theory of caloric pictured heat as an elastic, almost weightless, liquid. In discussions of this theory, it is noticeable that the phrase 'the quantity of caloric' occurs frequently, as opposed to other quantities which tended to be referred to simply by their name eg 'the pressure'. This may have led to the adoption of Q for heat energy, particularly if H was then in use for some other quantity.(2 votes)
- Is Q always defined by the system unlike W which can be defined by the surroundings or system?(1 vote)
- Q is the heat of the system so therefore it depends on the system only and this law is made for the internal energy of the system (not surroundings). So it will be positive if flows into the system and it would be negative if it flows out of the system.(2 votes)
- what happens when work is done on gas molecules?
Also if the collision were to move many gas molecules, could the gas molecules collide into each other eventually cancelling out the energy that all the gas molecules have(1 vote)
- What are the different kinds of energy that a gas molecule can have? Well, since the gas molecule can move around we know it can have just regular kinetic energy. Sometimes we call that translational kinetic energy and if it's a molecule like this. Look at, a diatomic molecule. It can rotate about some center point. And so, if we had a molecule built out of multiple atoms it can rotate and because of htis it can also have rotational kinetic energy. This would not be true if it was just monotomic. In other words, if it was just a single atom, the rotation of a single atom it turns out is not a meaningful significant contribution to the energy that a gas can have. But if it's diatomic, or triatomic, or any multi-atomic atom it can have a rotational kinetic energy as well. And you might think that's it. It can rotate. It can move around. What other kind of energy can it have? Turns out it can have one more. Again, if it's a diatomic atom like this the atoms that make up the molecule are bonded. They can oscillate kind of like two masses on a spring and because of this you can get an oscillation form of energy. This degree of freedom, we call it, is another place that the energy can go. So, if you add energy to a gas those gas molecules are going to either start moving around faster, start rotating faster, or start oscillating faster, or some combination of all of those. Those are the three ways that energy shows up when it gets added to a gas. The three ways that energy manifests itself when you put energy into a gas. And physicists created a name for all that energy. If you add up all that energy for a gas we call it the internal energy and we give it the letter U. So, U is the internal energy of the gas, and by internal energy of a gas we mean all the energy, the kinetic, the rotational kinetic, the vibrational energy. All of that energy added up is what we mean by the internal energy of a gas. But why am I telling you all of this? I'm tell you this because I want to talk to you about the first law of thermodynamics and the first law of thermodynamics is really an answer to the question: how do you change the internal energy of a gas? How do you increase the internal energy or decrease the internal energy? We know it's going to happen once you increase or decrease it. The gas molecules are either gonna speed up or slow down, rotate faster, rotate slower, and so on. But how do you get the energy in there? That's usually formulated, this first law of thermodynamics is usually formulated in the context of a gas that's contained in an enclosed container. Usually some sort of cylinder as the way it's shown. And in equation form the first law looks like this. We want to know how you change the internal energy of a gas. So, it looks like this Delta, which represents the change in the final value minus the initial value of the internal energy equals. So, this is the equation representation of the first law. What's gonna go on the right-hand side? All the ways you can change the internal energy of a gas. One way you can do that, well, just stick a fire underneath this container. Let's say this container's closed up. And you put a fire underneath. That heat's gonna enter into the gas and that gas will start moving around faster, and start vibrating faster, rotating faster, depending on the temperature, and this heat is the first way you can change the internal energy of a gas. So, if heat enter this enclosed container the internal energy will go up. So, if you add heat, and that's represented with the letter Q. Q is the letter we choose to represent the heat energy. If I had 100 Joules of heat energy that's 100 Joules that can go into increasing the internal energy of a gas. But it doesn't have to be fire that you use to add heat energy to a gas. You can imagine just submerging this enclosed container in some sort of heat reservoir. Maybe some boiling hot water or just warm water, and that would also add heat to the gas. You might object. You might say, "Wait, it's enclosed. That means nothing can get in or out. "How can heat get in?" Well, by enclosed we mean no particles, no molecules can get in or out. But heat's just a form of energy. So, what's really happening is this heat is causing the sides of this container, the atoms and molecules that make up the container to start vibrating faster back and forth. What happens is when this molecule collides with that faster vibrating side it gets a kick, a boost, for every time it hits one of those faster moving molecules it gets a boost. That's how energy's entering but just energy's entering. There's no molecules actually entering into this gas. So, it really is enclosed. Okay, so heat is the first way we talked about internal energy of a gas changing. But there's another common way to add internal energy to a gas. And we need another term over here. And the additional way to add internal energy to a gas is imagine instead of having a container where none of the sides can move, right, where this container's completely rigid. Nothing can move. Imagine the top of this container being such that is has a tightly fitted piston. And this piston, imagine, can move up and down. So, this piston can change the volume in which this gas gets to play in here. So, since this piston can move up or down. Well, what can happen? How can we add energy? You can just exert a big force downward and compress the gas into a smaller and smaller region. We said that the gas when it hits the faster moving molecules in wall gets a kick. Well, the same is true here. If we take this gas molecule and imagine we're pushing the piston down. When it collides up here with this piston that's moving downward, again, it gets a kick and it starts moving faster. It gains energy and that can manifest itself as translational kinetic energy. Rotational kinetic energy. Vibrational, regardless, it's all internal energy and so, well, we're exerting a force. This force is exerted over a certain amount of distance. If we're pushing this thing down and we know what force times distance is. We're doing work. That's how we're adding energy into this gas is work is being done by pushing the piston down, and if we do work on the gas that means we're adding internal energy to the gas and the value of the work is the amount of energy we're trying to add to the internal energy of the gas. So, these are the two common ways that you can add energy, internal energy, to a gas and this is the formula version of the first law of thermodynamics. So, this is it. These are the two ways you can add energy. Q is the heat, W is the work done on the gas. Now, important point here. This is the work done on the gas. And that's why I put a plus sign here because we're, by doing work on the gas, we're adding energy to the internal energy. We're adding energy. The gas is gaining energy because we're doing work on it. Now, some textbooks and other resources, you'll see the first law written like this except instead of a plus sign you'll see a minus sign and there'll be a minus sign here instead of a plus sign because they'll be talking about the work done instead of on the gas, they'll be talking about the work by the gas. Which is to say if the gas were to push the piston up it would be losing energy. It would be doing work on its surrounding environment out here, losing that energy, that would be energy that's getting taken away from the internal energy of the gas, because we're talking about the gas in the piston here. Inside the cylinder. So, you've gotta be careful. I've formulated in the version with the plus sign but you can also see this with the minus sign. And in problems with questions you've gotta be really careful. You've gotta look at is it asking you for the work done on the gas or it asking you for the work done by the gas? You might think, "Oh my gosh. "This is gonna be really hard. "How do I, how do I figure out one? "If I know the work done on the gas "how do I figure out the work done by the gas?" Well, that's really easy. The work done on the gas is equal to negative the work done by the gas. And vice versa. The work done by the gas is equal to negative the work done on the gas because if the gas does 100 Joules of work that's, like, someone doing negative 100 Joules of work on the gas. Because if you do negative work on an object you're actually taking energy away from it. So, that's the first law of thermodynamics. It tells you how to add internal energy to a gas. Now, if you're clever you might object at this point and say, "Wait a minute. This isn't a new law of physics. "This is just conservation of energy." It says the energy you add into a gas shows up as the energy in the gas, and what I have to say to that is you're right. This is just conservation of energy. It's not really a new law at all. Why do we give it a special name? Well, the reason is, when physicists were developing this law it wasn't clear what heat was. It wasn't clear that heat was just another form of energy. It took a while to realize that heat is just another form of energy. For a while, there was talk about heat being some sort of weird fluid that could actually enter a substance and now we laugh at that. We're like, "Pfft, ha-ha-ha-ha, what a bunch of dopes." But this is a microscopic property and it's hard to see unless you have an ability to talk about these things microscopically which took a while to do. It's not obvious what we should think about this material of heat as being. So, that's why it got a special name and the name stuck. We like it. It's the first law of thermodynamics. Now, the way you use it in problems. Gotta be careful. Heat can't just enter. It can also exit. So, if heat enters, Q is a positive number. So, is 100 Joules of heat enters you'd plug in a positive 100 Joules of heat for Q. But if heat leaves, if you stick this whole thing in an ice bath, and a certain amount of heat leaves, maybe 100 Joules of heat leaves, you'd have to plug in negative 100 because this Q represents the heat added to the gas, and if you're taking heat away, well, that's like negative heat being added to the gas. So, if heat leaves your gas that's a negative Q. If heat's added, that's a positive Q. How about work? There's work done on the gas. How do you know if this is positive or negative? Well, if you're pushing the piston down. Well, it doesn't matter. If the piston moves down, whether you're pushing it or not, work is being done on the gas. So, this is positive work done. So, if the piston's moving down, positive work is being done on the gas, that means energy is entering. That's a positive value of W. Now, if the piston expands. That is to say if the gas expands and the piston moves up, well, that's gas doing work on the piston and the environment. That's negative work done on the gas. Remember, we're talking about work done on the gas. So, if piston moves down, positive work done on the gas. Gotta plug in a positive value here. If the piston moves up, that's negative value of the work done on the gas. You'd have to plug in a negative value for the work here. So, be very careful. This is a way to calculate the internal energy. It's the first law of thermodynamics and one of the most fundamental, and most often used equations in all of thermodynamics.