This introductory physical chemistry course examines the connections between molecular properties and the behavior of macroscopic chemical systems.

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From the course by University of Minnesota

Statistical Molecular Thermodynamics

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This introductory physical chemistry course examines the connections between molecular properties and the behavior of macroscopic chemical systems.

From the lesson

Module 5

This module is the most extensive in the course, so you may want to set aside a little extra time this week to address all of the material. We will encounter the First Law of Thermodynamics and discuss the nature of internal energy, heat, and work. Especially, we will focus on internal energy as a state function and heat and work as path functions. We will examine how gases can do (or have done on them) pressure-volume (PV) work and how the nature of gas expansion (or compression) affects that work as well as possible heat transfer between the gas and its surroundings. We will examine the molecular level details of pressure that permit its derivation from the partition function. Finally, we will consider another state function, enthalpy, its associated constant pressure heat capacity, and their utilities in the context of making predictions of standard thermochemistries of reaction or phase change. Homework problems will provide you the opportunity to demonstrate mastery in the application of the above concepts.

- Dr. Christopher J. CramerDistinguished McKnight and University Teaching Professor of Chemistry and Chemical Physics

Chemistry

I want to return to something we've seen before, namely heat capacity, but now

Â within the context of considering constant pressure as well as constant

Â volume processes. And so heat capacity is a path function

Â not a state function. Let me remind you of the definition of

Â heat capacity, it's the amount of energy required to raise the temperature of a

Â substance by one degree. And it's different when it's done at

Â constant volume, compared to say, constant pressure, that's how we know

Â it's a path dependent function. So at constant volume, the energy added

Â is heat qv is equal to delta u. Indeed that's the definition we have of

Â heat capacity. It is change in u with respect to change

Â in temperature at constant volume. And if I write it in a infinite decimal

Â way, that would be partial derivative of u with respect to t at constant volume,

Â all right? So, at constant volume change and

Â internal energy is heat, and that's an expression then of heat capacity.

Â It's heat, relative to change in temperature.

Â At constant pressure, on the other hand, we will define a constant pressure heat

Â capacity as, the change in enthalpy with respect to the change in temperature.

Â Because at constant pressure the heat is equal to the enthalpy.

Â Alright, so the amount of heat Required to change the temperature by one degree,

Â that would make the denominator one degree.

Â That is the heat capacity. That's the definition.

Â So partial h, partial t at constant pressure.

Â So with those definitions in mind, let me give you a chance to think about what

Â their relationships might be and then we'll come back and look at them more

Â closely. Alright, so the question there was, which

Â of the two heat capacities, constant volume or constant pressure would be

Â larger. Lemme take a specific example, now that

Â you've had a chance to think about the answer, and hopefully answer correctly.

Â Then we're going to get to here. let me do an ideal gas, and so recall for

Â an ideal gas definition of enthalpy is h equals u plus pv.

Â And, so, for an ideal gas, pv is nrt. If I differentiate both sides with

Â respect to temperature, I'll get dh dt is equal to du dt plus nr.

Â Now, for an ideal gas, you and hence, h, depend only on temperature, not on

Â pressure or volume. so, I can write this more generally, this

Â exact differential is also equal to the partial of h, with respect to t.

Â It's just a differetn way to write it. There are no other things h depends on,

Â but I'll use the partial symbol. At constant pressure, equals partial u,

Â partial t, at constant volume plus nr. That is, cp equals cv plus nr.

Â So the heat capacity at constant pressure is always going to be greater than the

Â heat capacity at constant volume. How much great for an ideal gas?

Â By a factor of n, number of moles of the gas, times r.

Â And I'll just remind you if, if you haven't got a feel, is that a lot, a

Â little? Well remember that for a monatomic ideal

Â gas, the molar heat capacity at constant volume was three halves r.

Â So that would make the heat capacity, the molar constant pressure heat capacity,

Â five halves r. That's, that's a 67% change, right?

Â r is 67% of three halves r. and that's a nontrivial difference.

Â So it takes considerably more heat added to raise the temperature by one degree

Â when working at constant pressure, compared to when working at constant

Â volume. So and of course the reason by the way is

Â that you've got to expand the gas. You're doing pv work and you've gotta put

Â the heat in to do that. Well the reason this is interesting, this

Â heat capacity at constant pressure is, it allows us to potentially determine

Â enthalpy experimentally. Or more accurately, perhaps enthalpy

Â changes. Remember, thermodynamics is usually about

Â enthalpy changes. But, we'll get to establishing zeros and

Â being able to tabulate numbers in not too long.

Â But let's talk about determining enthalpy then.

Â The difference in enthalpy at two different temperatures Is going to be

Â determined by integrating cp over the temperature range.

Â We've already seen this for internal energy, integrating cv.

Â Now I'm going to take cp equals partial h partial t.

Â So I rewrite this as dh equals cpdt. So, if I want to know h2 minus ht1, I

Â integrate dh from t1 to t2, that's equivalent to integrating this from t1 to

Â t2, and here that is, integral t1 to t2 cpdt.

Â Now, it's important to mention this is only true if the phase of the system

Â remains unchanged over that temperature range.

Â t1 to t2. We actually did an example not so long

Â ago of ice melting or water boiling. That's a phase change and it takes

Â additional heat which is enthalpy at constant pressure added to the system in

Â order to accomplish a phase change. So, at a phase change the heat capacity

Â becomes infinite, right? You are adding heat into the system

Â without changing the temperature. So the denominator if you like is, is

Â zero, and that's why it goes to infinity. but you can measure heat capacity over

Â the range of a pure phase, and then you'll eventually get to a phase change

Â and you'll need to measure that differently.

Â But it, it can be measured. And so, if you like, a way to think about

Â crossing a phase boundary would be if I want to know the enthalpy at a

Â temperature t and I'll start from zero. I'll have to assign some number to the

Â enthalpy at absolute zero, we can talk about how you might do that, but anyway

Â there is some number associated with that.

Â How would I do it? Well, I'll integrate from zero up to the

Â melting point. I'm going to assume that it's absolute

Â zero, it's probably a solid. I'll integrate the solid's heat capacity

Â up to the melting point, then, I will add the enthalpy of fusion.

Â And then, I'll integrate from that temperature, which is still the same

Â temperature, the temperature of fusion, up to the temperature of interest t, the

Â heat capacity of the liquid. And it'll have a distinct heat capacity

Â from that of the solid. So just what I said, solid from t equals

Â zero to t fusion, enthalpy of fusion, which is the enthalpy of the liquid,

Â minus the enthalpy of a solid at that fusion temperature, and then the liquid.

Â So let me show you what that looks like in practice.

Â Show you some experimental data. Benzene, so benzene and aromatic organic

Â molecule. found in oil and it has a melting point

Â of 278.7 kelvin and a boiling point of 353.2 kelvin.

Â And if you measure its heat capacity, temperature by temperature.

Â And what does that mean to measure the heat capacity?

Â It, it's pretty simple. You put a thermometer in the substance,

Â you add heat in some controlled fashion that you can quantify how much heat

Â you're putting in. And you measure how much you put in to

Â get it to go up one degree. Maybe that's how many meters of methane

Â did you burn in a Bunsen burner? Maybe it's how many calories of sugar did

Â you expend while you were turning a crank?

Â That would be a hard experiment to do, but in any case you can measure it.

Â And you can measure it degree by degree. So, it does vary, this is temperature on

Â this axis, this is the constant pressure heat capacity and what you see is, at

Â absolute zero it takes hardly any. And then, as you rise in temperature,

Â it's taking increasingly more energy added in order to raise the temperature.

Â And that's because that energy is flowing into more accessible modes.

Â Rotations, vibrations are beginning to pick up some of the energy, and they're

Â not being put into translation, which increases temperature.

Â That's really a gas explanation, and we're, we're sitting in a solid region,

Â but it's the same conceptual ideas. In any case, we rise, we rise, we rise,

Â we finally hit the melting point. Here the heat capacity would go infinite.

Â We can't measure it directly. We would have to do a, a measurement

Â where we just watch until all the solid melts and becomes liquid.

Â And we know how much heat we put in, so that's the heat of fusion.

Â And now we go measure the liquid and then, it boils and we measure the gas,

Â and this would be key experimental data. The enthalpy itself then, because these

Â are enthalpy changes with respect to temperature, to get the enthalpy is the

Â integral under these curves. And so that's what's shown over here.

Â As I raise the temperature, what is the additional enthalpy, that is the

Â additional heat that's been absorbed into the system, that's there.

Â I can extract it maybe to do interesting things, as a function of temperature.

Â So I don't have anything in there when I start, and I'm pouring it in, pouring it

Â in. It's going up, up, up as a solid, and

Â then the temperature doesn't change anymore but I pour in a whole bunch to

Â make it change phase. I pour in more to increase the enthalpy

Â of the liquid. Stops absorbing but does start turning

Â into a gas. Finally I've got all this enthalpy in the

Â gas. So I'm integrating that is for t above

Â the vaporization temperature. The enthalpy relative to zero would be

Â integral over the solid, plus the, the heat of fusion.

Â Integral over the liquid, plus the heat of vaporization.

Â And then finally integral over the temperature to whatever temperature I'm

Â at in the vapor. Alright, hopefully that made clear.

Â The, the experiments are relatively simple, and you would have a way to

Â record the enthalpy change relative to absolute zero.

Â Might not be the most convenient place to anchor your scale to.

Â It might be hard to do a measurement at absolute zero.

Â So, we're going to talk next about thermal chemistry in general, using these

Â enthalpies in order to. Understand reaction energetics make

Â predictions, not like what I illustrated for the thermite reaction in week one.

Â But we'll explore it in a literally more generally here and standardize it some

Â [SOUND].

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