Skip to main content
Loading Explore guides…
Thermodynamics · Acedra
Home
Explore
Chemistry
Thermodynamics
Sign in to save
Chat
If YouTube asks you to sign in, open the video on YouTube.
Watch on YouTube
Reload player
Transcript
79 segments
Copy
Hide video
0:01
- Hi I'm Jeff Yarger. I'm a Professor of Chemistry, Biochemistry, and Physics, and this video is gonna cover Physical Chemistry for the Life Sciences edition number two, and today we're gonna be looking at Chapter 1 in this textbook. So specifically, the first part of this book is going through biochemical thermodynamics. And today we're gonna be looking at Chapter 1 which is the first law of thermodynamics or the conservation of energy and its implication in biochemistry.
0:38
So this chapter is broken up into numerous components,
0:46
three major components, which is to look at the conservation of energy, the first law of thermodynamics,
0:53
to look at it specifically from an energy perspective, the internal energy of the system, and defining what an energy,
1:05
the enthalpy which is another form of energy, and how we often use that in chemistry and biochemistry, and then finally looking at specific chemical and physical changes and how that changes the energy typically through the enthalpy. And we'll discuss why the enthalpy is so critical in chemistry and biochemistry. And then finally I also like to point out that the book does an excellent job of giving you key concepts at the end to checklist through,
1:37
to make sure you understand when you've read the chapter,
1:41
the key equations, discussion questions to really make sure that you've understood conceptually the material covered as well as practical exercises beyond what's given in the textbook as an example, exercises and test, example test questions, there's a whole series of in a sense homework problems or exercise problems, and then also some really in depth kind of project-based problems as well. So this video is mainly designed just to kind of overview some of the material
2:13
in this chapter and by no means is a substitute for really getting in and reading the material yourself and going through the examples that they give and working problems yourself to really make sure you understand the concepts behind the first law of thermodynamics and its relationship in biochemical systems.
2:38
So what I'm doing is gonna provide a slide for kind of each one of these topic areas, each of these 13 specific subtopics that's covered in Chapter 1.
2:52
So the first is kind of some bookkeeping. And so the one thing I'll remind you is almost everything we're covering in this is the first law of thermodynamics. In fact the first several chapters, this is all you know something that has been covered in introductory chemistry and physics courses before. The main goal is to first kind of give the background information that we've had from some of those previous classes and then carry that on to really looking at biochemical systems
3:23
and making sure we really approach it from maybe a more rigorous mathematical perspective and molecular perspective as well this time around looking at this topic. So the first part of this is reminding us when we look at basic thermodynamics is how we get through some of the basic nomenclature and definitions that we use. So for example Figure 1.1 and 1.2 talk about
3:52
how we define you know the system, the part of the universe that we're interested in studying, and because we're typically, we're talking about chemistry, biochemistry here, you know it's pretty appropriate to kind of show it as a test tube 'cause it's usually often in a container that we're doing experiments on you know the part of the universe that we're interested in, and everything else being the surroundings. Now if we block off that surroundings and isolate it to have an isolated system,
4:23
this gets immediately at the first law of thermodynamics which is conservation of energy. So in a closed isolated system the change in the internal energy will be equal to zero. And again from freshman chemistry we'll often do the change as a delta, but now we're moving on to really calculus base where we're gonna look at the change as the derivative.
4:50
So the change in the internal energy for an isolated system being zero but the overall energy of the universe, the overall isolated system, can't create nor destroy energy in the process, but oftentimes we're allowing energy to exchange and so we can do that by sometimes not isolating the system so that it can exchange energy, but say for example we close the system so it can't exchange particles or number of moles. And so we would often say this
5:23
as the change in the number of moles is equal to zero, so that we can't add or subtract material, mass of any sort or any atoms, molecules, compounds, et cetera in and out of the system, or we can relax that, open up this and allow material into and out of the system. Or in other words you can change the number of moles in the system. You can do different types of chemical reactions and do chemical work, et cetera, in a system. And so this is just kind of reminding us
5:58
as well as later on defining, I like to think of this as defining the walls as well, like do you allow temperature diathermal versus adiabatic changes where you can't change the temperature in a system, for example in like a vacuum type container.
6:19
And so we often define thermodynamics in what variables are gonna be held constant versus not by defining rigorously the walls of our container that separate the system from the surrounding.
6:34
Now thermodynamics is all about energy. It's really looking at how energy changes in a system and we're typically used to often defining energy as splitting up into having multiple components. One of the most familiar is probably splitting up into either kinetic or potential energy, and this is you know most commonly done in introductory physics and physics courses. For macroscopic thermodynamics the split that makes the most intuitive sense is to split it up into different types of work,
7:11
mechanical work, pdV, chemical work, the chemical potential into the change in the number of moles, maybe electrical work, electrochemistry through changing voltages, right,
7:25
magnetic work, et cetera. So we have numerous types of work and then heat, and this is really at the heart of thermodynamics as implied by its name, is heat as a type of energy. And really there's a big distinction between these, work being a coherent type of energy while heat being a more random type of energy that we'll define later on at the molecular layer, at the molecular level. And really at this point the book goes through you know some intuition into how you can think about these
8:06
and for example if you lift, you do some type of pdV work on something, you would be coherently lifting that weight off, all the molecules in one direction, while heat we think of as kind of temperature changing and that's really how you know the internal motion of molecules which are random at fundamental levels are changing in a system.
8:32
So now we wanna get to that we're gonna split up you know,
8:38
that the energy is conserved in a system but you can think of it in two different kind of components, in work and in heat components, and the sum of those two will add up to the overall energy of a system. And now it gets into some practical of measuring some of these quantities, and what we really mean by positive versus negative work. And the way I like to think about this and the book does an excellent job of going through you know defining how the energy is related
9:11
to the amount of work and heat in a system. And the way I like to think about this is we're thinking about this from you know from the standpoint of the system itself. So when work you know is being done by the surroundings onto the system so that at the end you know it has more, you know it has gained work or it has gained heat into the system, you know that's what we think of as therefore gaining energy and being a positive amount of work, a positive amount of heat with respect to the system,
9:49
and then vice versa if it loses, if the system loses internal energy because it's losing heat or work, or giving it off, doing work onto the surroundings, giving work or heat off to the surroundings, then that would come as the opposite sign negative. So, and the reason I really mention this is because this is really one of the fundamental things that often in practically working problems that often gets messed most, or missed, which is just the sign of work or heat in these,
10:22
and therefore when you start adding and subtracting them, if they have the wrong sign that can propagate a large amount of error in some type of calculation. And then since we're talking about work specifically here after we've kinda defined it in general, it introduces what I would say is the first type of chemical process where we look at work as far as its mechanical work on for example a gas as being minus the external pressure change in the volume. And again we're gonna write this
10:58
as the derivative in volume instead, while often in freshman you'll see it as dV like this. So we initially start looking at this pressure volume work and it explains why it comes as the negative because how the volume changes with pressure is inversely proportional.
11:23
And it goes on to talk about measurements of heat
11:27
and it gives, so and we typically do this through temperature. We typically measure heat through temperature. Work can be complicated, whether we measuring some you know height which is related to some volume of a system in pdV work, or whether we're in chemical work we're measuring how the number of moles of something change, electrical work, it can vary but almost inevitably heat is almost always measured through temperature and this is called experimentally calorimetry,
12:01
and the book goes through both constant volume calorimeters and constant pressure calorimeters, and I have several other videos on this topic as well because it's so important to thermodynamics especially in biochemistry. You know calorimetry is often our gateway into measuring thermodynamics in biological systems. And so this is reminding us of topics that have been covered before which is how the heat capacity, the capacity of a substance to store heat
12:35
and in a sense resist temperature change, and so, and you know it reminds us kind of at a molecular level as you make things, more molecules with more degrees of freedom, say more atoms associated with them in a molecule, the heat capacity goes up because you're gaining additional vibrational degrees of freedom that you can store stuff in, et cetera. Also you know some of the unique properties of water, specifically liquid water, having such a high heat capacity
13:11
even though you haven't changed overall you know the molecule in this case from ice to liquid water to vapor, but you do change the heat capacity when you go into the liquid state. And you know because of hydrogen bonding and again kind of approaching it from a molecular standpoint. And then reminding us that it's this derivative. So it's how the internal energy changes with temperature
13:40
and this is at constant volume you know which is you know defined as the heat capacity at constant volume. And you know, and then it goes on to now look at you know this internal energy directly. So we, it kind of discusses ways of measuring work and heat
14:04
and then now let's go on to really make sure we can put this all together to look at the first law, the conservation of energy that the change in energy of a system, which we call the internal energy and we usually represent by a capital U,
14:23
goes as the change in work and the change in heat. And so, and I've written it in this way. Oftentimes you'll also see it kind of written sometimes like this to indicate that these are path dependent changes while this is a state function and is path independent variable. So the change in internal energy is the change in work plus the change in heat of a system.
14:54
And we just looked at the internal energy as it changes with temperature. You can also look at it with volume as well. So you can look at all the partials of how different energies change with respect to different variables, holding other ones constant.
15:14
Now, so the internal energy of a system if we look at what we we're looking at with work here, so work being minus pdV, so the internal energy minus pdV,
15:28
and then heat plus the change in heat of a system, right,
15:35
what we find is is that we often want to look you know, we want to be able to look at the internal energy and have it being proportional to you know its change in heat. So we could do that by holding something at constant volume and that's what you know it's talked about here, the constant volume heat capacity, but it turns out especially in biochemical systems but also in a lot of chemical systems, what we typically work at practically is constant pressure.
16:06
And so what we do is change to a type of energy where it's at constant pressure that the work is changed.
16:18
So we define this, we define the enthalpy as U plus pV
16:25
and I give a separate video kind of showing the Legendre transformers, et cetera, of how this changes from being internal energy which is its dependent variables are volume,
16:38
to its now variable being pressure dependent instead of volume dependence. So now if you're holding things at constant pressure, then the change in enthalpy will be related to the change in heat in a system. And so before we get on to defining heat from a molecular level and defining entropy et cetera, this is often a common way to look at thermochemistry. And so we're introducing this again here just like you've seen it in introductory chemistry courses.
17:12
And constant pressure is much more common under biological conditions to work at.
17:19
So it's often the enthalpy from this standpoint on that we'll be looking at things, not the internal energy directly, but this enthalpy is an energy, right, and it's, so in fact most of thermodynamics that we look at, we're just looking at changing the type of energy we're looking at by changing what dependent variables are in play. And so the enthalpy is U plus pV, and this is typically a positive quantity
17:51
so it's always a little higher than the internal energy and it's sloped just like the slope of this is related to the heat capacity. So how the enthalpy changes with temperature at constant pressure is the heat capacity at constant pressure. So, and that's the slope at any specific temperature here.
18:19
So, and this even shows of how you can look at this heat capacity, and this is for example of the unfolding of ubiquitin, it gives us an example and it shows several, I show a differential scanning calorimeter figure here, but it also shows bomb calorimetry or constant-volume calorimeters as well, and I have separate videos that discuss calorimetry specifically. And this is typically our experimental gateway to measuring these energies through measurements of changes in temperature of a system.
18:56
So now that we have these enthalpies, we wanna look at it to get some intuitive feel kind of from a molecular standpoint of either doing physical changes or chemical changes. And so starting with kind of physical processes, things that we're very familiar with for example ice going to liquid water, liquid water boiling to a vapor, et cetera, or vice versa, right, or even you know sublimation, ice going directly to the vapor phase,
19:29
and looking at that, those energies through looking at their enthalpies, because we're typically doing this at constant pressure.
19:41
So the relevant energy is enthalpy that we wanna look at in this for how the overall energy, 'cause we're, 'cause that will directly relate the energy to how much heat not doing any direct mechanical work
20:00
in this system.
20:02
And so those start with kind of intuitive processes and then we move on to kind of drilling it down more to the molecular level. And again this is kind of reminding us of things we've seen in introductory chemistry which is we can always because enthalpy just like internal energy is a state function, meaning it's path independent, and from a mathematical standpoint it means that you know these are exact differentials, and so we have separate math tutorials and videos
20:34
to kind of cover that topic more specifically,
20:38
we can always break this down because of Hess' law which is basically just a statement that you can take any path you want, it's just, the enthalpy's just dependent on its initial and final conditions, so we can always take this all the way back to the elements, and so one thing you can always do is just break every bond in the system and then reform them. You can break every bond in the reactants and reform them all into the products you want and look at that overall delta H, the reaction.
21:12
And so it's often very useful as this table provides that you have some general numbers of what these bond enthalpies are.
21:24
So taking that, another important process is looking at combustion or the reaction with oxygen of different materials to form carbon dioxide and water, and you can relate this to the formation enthalpies from you know basically taking it all the way back you know to the elements and how they would form different compounds. And this is, so we give you know some of the combustion enthalpies and therefore what some of the specific enthalpy densities are, and it puts this
22:02
in relationship to common fuels you would use, both what we would call common fuels for you know gasoline, you know things we would use to, in kind of a mechanical sense to fuel our vehicles et cetera, as well as more a biological sense. So you know fats and stuff, or fatty acids that we would use to fuel our bodies, and so it discusses it kind of from both standpoints and how much specific energy you have in some of these things and what their heat of combustion are,
22:33
how much heat they would release upon combustion, and you can also look at this in terms of formation as well. So when we get to biological systems we often get to combining multiple reactions enthalpies. And this is again reminding us that when we're looking at a specific variable, and you know some things are path dependent and some things are path independent. And so the energies, whether it be the internal energy or the enthalpy is path independent. We can take any path we want.
23:11
However some things are, like the amount of heat or the amount of work, those are path dependent variables. And so it's often, we're often looking at these energies and breaking it into paths that are easy for us to calculate and they're often multiple steps. And so we often using Hess' law to get to the enthalpies
23:36
of some you know more a combination of reactions to get to where we wanna be.
23:43
And so, and this is just reminding us of that any reaction can be written as its formation enthalpies of the products minus the reactants times it, so the sum of all of the product formation enthalpies in its standard state, taking into account the stoichiometric coefficients, the number of moles of each of these, and the same for the reactants, summing 'em all over, and then that giving the overall enthalpy of the reaction. And I really like that you know thinking of it
24:22
from these enthalpy or energy standpoints, you know that we can think of it as that we can take everything up to the elements and then form the reactants and form the products, and that this you know delta between the reactants and products you know comes out
24:44
of taking it up to any other path or any other state we want. And that some of these, this one showing things that are all exothermic, because a lot of what we're looking at especially in biological systems, the molecules are exothermic
25:01
as far as their formation enthalpies, but you can have endothermic where it requires heat as well. And again keeping in mind that these are all kind of under you know constant pressure which I often write as just the change in pressure is zero.
25:19
Now one of the things I love about the chapter in the book is getting to you know a point to of looking at computation as a method to really look at enthalpies, and specifically formation enthalpies, giving two nice examples, one kind of classic fuel type or hydrocarbon based methylcyclohexane
25:48
and another one biological looking it at glycine in several different states. And this one is looking at kind of its zwitterion properties under different solvation dynamics, for example in water.
26:04
And this one just talking about its, you know whether the methyl is in an axial or an equatorial plane and how that can have a change in energy between that and that can show up as you know a change in its formation enthalpy, so that when you subtracted one formation enthalpy from the other it would give you the change in enthalpy or energy between these two different molecular structures,
26:38
or molecular forms of the same molecule. And I've done this just to show practically in mole calc, and it's at molecalc.org, while this is not the highest level of theory, it uses an ab initio of AM1 which is a semiempirical type, it still is a good, it's something anyone can use on the web and you can get some initial experience with looking at some kind of computational chemistry and seeing these type of differences. And this is the ones we calculated, I did with mole calc,
27:15
work out similar to what is discussed in the book. And I'll potentially maker a separate video just going through this 'cause I think this is a very important point that computational chemistry, especially with kind of smaller molecules, is often very insightful from a molecular standpoint as well as just an accuracy of getting a feel for some of these enthalpies and energy changes between different molecules, different states of molecules, et cetera.
27:50
And then finally looking at oftentimes while we're often doing this under constant pressure so that we're looking at enthalpy, we'll often do this either at a nonstandard state temperature, something away from room temperature, or something where we're changing the temperature. As we said oftentimes measuring the change in temperature experimentally is calorimetry, we do quite often. And so how that energy at constant pressure, so the enthalpy changes as a function of temperature,
28:28
you know goes as its heat capacity, and the heat capacity can have a temperature dependence. This is showing it being temperature independent to first approximation, but by having different, a delta Cp or different heat capacities between reactants and products, that can cause a different delta H at various temperatures. This one's showing that the enthalpy of the reactants is going up faster, has a higher temperature dependence than the enthalpy of the products, and you can see therefore
29:01
as you get to higher and higher temperatures that delta, the enthalpy of that reaction, keeps getting larger and larger. It can go the opposite way. It can go where the products rises faster than the reactants, et cetera. So, but this gives a nice visualization of this and you know again we often use Hess' law where we'll do the reaction at 298 but we will be able to take you know the reactants
29:33
through their heat capacity and change in temperature to 298, do the reaction, and then take those products up to whatever temperature, and then that overall path will allow us to calculate the delta H at some nonstandard temperature, okay, where you'll then, you need to know the heat capacities of all the relevant reactants and products. And when we're doing this enthalpy, you need to know the heat capacities at constant pressure for the system. And then the book does an excellent job
30:06
of going through looking at glutamine glutamate and the enzyme that causes the change here,
30:18
and keeping in mind that we rarely do biochemistry at room temperature, at 298 Kelvin or 25 degrees Celsius, we often do it at psychological temperatures. We think of our own physiology being around 37 degrees Celsius, right, so it's often at these a little higher temperatures. Now do we expect that to change these reactions a little or a lot? Oftentimes the heat capacity is in the order of joules per K mole and we're looking at reaction changes in the kilojoules,
30:56
so it often doesn't change it a lot, sometimes only a few percent over, but it matters how far in temperature, et cetera, but to get really accurate changes in energy, or enthalpy in this case, we often need to account for this.
31:15
So and then like I said at the beginning, the textbook then goes through where you can, after you've gone through this material, you can go through the checklist of key concepts, key equations which I show here,
31:29
the discussion questions, exercises and projects to really make sure you have a fundamental and conceptual grasp of the material covered.
31:39
Thank you.
Here to help you learn
Sign in to keep a private conversation with your course tutor.
Ask me anything about the material…
Sign in to chat