Thermodynamics

Dr. Miller

8/29/2008

*This morning the course takes a slightly different turn from what you’ve had so far. We’re going to leave structural biology and basic biochemistry and go into metabolism.

Slide 1: Thermodynamics or Bioenergetics

  • In order to really fully understand metabolism, we have to give you a little background in thermodynamics or what is called more properly for this course, bioenergetics. We essentially in this part of this course treat energy. Energy is something that you can’t really see but you certainly can feel it and of course it is the necessary requirement to have any kind of movement or any kind of living organization that one might have even from the lowest single cell to the more complex or most complex multicellular organism that you can possibly imagine, including man.
  • We’re talking about energy here: it’s transformations, it’s conservation, and it’s utilization. Transformations comes in, in the sense that energy that we have is essentially derived from our sun and as far as biochemistry is concerned, the energy, the light energy, is used in photosynthesis to make food stuffs. So, light energy, or the energy of photons, is transformed into the energy of chemistry, chemical energy in the form of carbohydrates, lipids, a number of substances we have for our energy. And then that energy has to be conserved until we utilize it. When we utilize it we take some energy in, we conserve it, and then we utilize it ourselves. So it’s a whole system of energy, ultimately (as far as we are concerned as biochemical organisms and biochemical entities) our energy system starts at the sun. Everything is derived from the sun. About 1% of the sun’s outlet hits the earth every day. It’s an amazing situation that the sun is producing a great deal of energy and it supports our whole planet, but only 1% of the sun’s energy that is given off every day actually arrives at the earth. So that’s a little background in terms of energy generalizations. Now, let’s get a little specific.

Slide 2: The Basic Concepts of Thermodynamics

  • In thermodynamics we talk about a system and the system we’re most interested in is our own organs and our own system. Each individual is a system as far as energy and thermodynamics is concerned.
  • The surroundings are everything else; the other people around you, the desks, the atmosphere, the building. Those are the surroundings outside yourself.
  • If you have an isolated system it cannot exchange matter or energy. We are not isolated systems, we are open systems. We can exchange both matter and energy with our surroundings. We have a totally open system and the open system is diagrammed here cartoon-like. (slide 3)

Slide 3:

  • Isolated system
  • closed system- can share energy but not matter with it’s surroundings
  • open system- living organisms are open systems. We share both matter and energy with our surroundings.
  • You must realize by now that you are a warm body and you are radiating energy out into the atmosphere and into the universe all the time. You have a temperature of about 98 degrees Fahrenheit (37 degrees Centigrade) and you are radiating heat out all the time. That is a sign that you are metabolizing. You are using some energy but you are inefficient and you are actually losing some of that energy just like your car loses energy through the radiator when you travel. We do share energy, we take in matter, we secrete and excrete matter and so we are totally open systems. By virtue of being an open system, we are contributing to the entropy of the universe. Every living organism contributes to the disorder of the universe.

Slide 4: The 1st Law of Thermodynamics

  • The first law of thermodynamics says the total energy of the universe remains constant. It is neither be created or destroyed.
  • In other words, you can have some energy, you can utilize it, but in utilizing that energy, you are converting it to other forms of energy. So a perfect example of that an experiment, a chemical reaction which is done in what is called a bomb calorimeter.

Slide 5

  • Bomb Calorimeter: In other words, the reaction takes place inside this little chamber, stimulated or catalyzed by a necessary electrical system or electrical spark. The reaction takes place and the chemical energy involved in the reactants gives off some heat. The heat that is given off is measured by the fact that the water surrounding this calorimetry system is warmed up. You can measure the warmth derived from the reaction. So, you are changing energy; going from chemical energy to heat energy in that situation. Chemicals are foundations or fountains of energy. When they react they generally give off some other type of energy. It can be:
  • light energy or
  • heat energy
  • other forms of chemical energy
  • These are the major outcomes of a chemical reaction.

That you have to keep in mind in thermodynamics and there is another concept called Enthalpy.

Slide 6: Enthalpy (H)

  • Enthalpy is the heat content or internal energy of a system. At constant pressure and temperature ΔH is our symbol for this and the change in enthalpy is essential the change in internal energy of a system. For an example:
  • Take a bucket of gasoline verses a bucket of water. You know from personal experience that the gasoline is much more reactive and volatile than water. You can actually smell the gasoline as soon as it is poured out into the bucket because it is quite volatile. If you dropped a match in the gas it you would see there is a tremendous amount of energy there ready to be released. If you did the same with a bucket of water the flame would be gone, nothing would happen. So the internal energy of something like gasoline (which is composed of long chain hydrocarbon molecules) verses water (which is a rather inert substance) the gasoline is definitely higher energy than is the water. (So think of those contrasts. You can always think of contrasts to help you understand some of these concepts.)
  • Every substance has an internal energy, whether it is chemical energy, or energy for warmth, energy for explosiveness, whatever…it has an internal energy potential.

Slide 7: The 2nd Law of Thermodynamics

Now, when we talk also about thermodynamics, we have the second law of thermodynamics.

  • The second law of thermodynamics is that the disorder in the universe constantly increases
  • There is no way in which you can stop disorder. The only way you can stop disorder is to put some energy into something and to fix it up. If you don’t take care of your automobile, it is going to deteriorate, then you have to put some energy in it to restore it. If you don’t take care of your room, your apartment, your home, and you let it sit with no repairs for years, it will deteriorate.
  • The entire universe is expanding very rapidly. We are becoming less and less concentrated with respect to the mass of material in the universe. That is a creation of disorder; going from a concentrated solution to a dilute solution. This is very easy. It is the natural way of things.
  • You will never find a dilute solution becoming concentrated all of a sudden by itself (you would have to put energy into that). So a concentrated solution is a higher level of energy than a dilute solution. If you put a concentrated solution inside a dialysis bag and put that in pure water, the ions are going to leave through the dialysis bag and go out into the area. It is going to simply dissociate the concentration and make it lower, lower concentration.
  • So, randomness is designated as entropy and the symbol for entropy is S. You can have a change in entropy. If it is a positive change, that is a positive disorder. More disorder is a good thing. That is something that is natural and occurs all the time. It is a process that contributes to the energy you get out of a reaction. If you can create some disorder in a reaction, you will get more energy out of it than if some of your explosive power or some of your chemical energy has to be devoted to ordering things. Then, that cuts down on your energy output, cuts down on the energy that you get from the reaction.
  • The ideal thing is to create disorder. If you have a minus ΔS, that is-entropy is being decreased, that from the point of view of energy output is a bad thing because that means in order to order something, you’ve had to use some of your energy to create this order situation. That is generally not a good thing for deriving chemical energy. So, entropy is S.

Slide 8: The 3rd Law of Thermodynamics

  • The third law of thermodynamics is that the entropy of any crystalline, perfectly ordered substance must approach zero as the temperature approaches absolute zero (0° K).
  • In other words; the only time that you can have a perfectly ordered system with no entropy, no movement on the part of the molecules, is when you have a perfect crystal at absolute temperature of zero.
  • Third law kind of a way to express what disorder is and how you would get rid of disorder. But, outside of those conditions, you always have a tendency for disorder to arise. You can see that in your own self. You can see it in your parents and your grandparents. A human being or any living organism for a while can sustain a certain amount of order because of the energy we put into ourselves. Ultimately, at the age of around 80-90-100-120, you no longer can keep up with it and the result is essentially cessation of life.
  • Disorder is the order of things even for living organisms. We can delay/inhibit disorder but you can’t overcome it completely.

Slide 9: Gibbs Free Energy (G)

  • In biochemistry, reactions occur most often, virtually always, at constant pressure and temperature.
  • Temperature and pressure deviations are very minimal.
  • Gibbs suggested we use this particular formula for energy in living systems.
  • ΔG, the free energy of the system, is equal to the change in enthalpy (the change in the internal energy of the system) minus the absolute temperature times the disorder equivalent.
  • What all this means is: for a system to deliver us energy, you must have a negative ΔG (free energy). You must have free energy liberated for work. The system must lose it’s energy. The way it can lose energy is it can lose internal energy and it can also create disorder and have ΔS be positive. If the system loses energy, that makes this (ΔG) a minus, and if it creates disorder, (if S is a positive), this(ΔG) remains a minus and that is a good outcome for the total free energy that is liberated from the reaction
  • On the other hand: you can have a reaction where you liberate heat or some type of energy from the reactants and products, but you create order and this (ΔS) is a negative. Negative times a negative gives you a positive. So, you will lose some of the free energy you might get from a reaction if you create order. If disorder is minus, you are creating order.

This is elementary algebra. There is nothing mysterious about this kind of thing.

  • The free energy available for many reactions is the difference between change in heat content or energy lost or gained by randomization processes at a given temperature. Randomization processes are measured by how much randomization you create. If you create a lot of randomization, you increase entropy, you make this term a minus and that helps your total free energy.

Slide 10: Parameters for ΔG

The parameters that we’re talking about are

  • ΔG must be minus if the reaction is exergonic and will proceed as written. So I have a reaction written here A + B ↔ C + D and the reaction has given up some free energy. If ΔG is minus, then the internal energy of the products, C & D, must be lower than the internal energy of the reactants, A & B. If that is not the case, then you are creating something because you can never have a spontaneous reaction which gives you higher energy products than you have in the reactants. You would have to put energy into it and ΔG would be positive. So, you can’t get something for nothing. There is no free lunch! If you’re going to have a spontaneous reaction which gives you some free energy, the products are going to have to have a lower energy than the reactants. We would like to have ΔG be minus.
  • If ΔG is positive, the reaction is endergonic and will proceed spontaneously only in the reverse direction. That is- if we have to put energy in the system to getA + B to go to C + D, then the products are a higher energy level and the reverse will be spontaneous. This is a very simple deduction from almost intuitive reasoning.
  • Now, ΔG is measured as calories per mole. We’ll use calories here even though calories is not international (Joules is international measure), we’re eventually going to have to talk about calories because this is the way we measure the energy input of our food stuffs (in this country anyway)
  • ΔG is calorie per mole
  • 1 scientific calorie equals roughly 4.2 Joules
  • ΔH is calorie per mole
  • ΔS is calorie per degree mole
  • If you use those units you will come out and see that ΔH + TΔS is equal to calories per mole. That gives you the formula that we need to work with.

Slide 11

  • Temperature has some effect on these systems, obviously. This is the denaturation plot of chymotrypsinogen. This is an enzyme that we talked about before. All this does is show you that you can have free energy liberated from the denaturation of trypsin. Trypsin becomes denatured giving off free energy as the temperature rises. In other words, the higher the temperature, the more rapidly chymotrypsinogen will actually denature. At the lower temperatures, chymotrypsinogen remains native. It takes the energy to denature it. The ΔG is essentially negative at higher temperatures. So temperature has an effect on these things but in biology we’re dealing largely with a constant temperature.

Slide 12: Standard Free Energy Change

Consider this reaction

  • A + B ↔ C + D
  • The reactants and products are initially at 1.0 M. In other words, you set up this reaction with all of the reactants A & B and C & D at 1.0 M concentration. You do the reaction at 25°C (298° K) and at 1 atmosphere. If you do that, the system proceeds to equilibrium.
  • The ΔG, which is ΔG° here (the standard ΔG), is equal to -RT ln Keq
  • All that means is, if you set up the reaction and it goes so products A + B go to C + D, then the equilibrium constant will be greater than 1. C × D, these products will be more prominent than A × B and this fraction ([C][D]/[A][B]), then, is greater than 1. The natural log will be a positive and the whole equation becomes a minus. The standard ΔG (ΔG°) is a minus.
  • The gas constant is 1.98 calories per degree mole.
  • All this is: A way of describing a chemical reaction which is a standard way. But anybody that does a reaction, (whether they do it in Sumatra, in Australia, South America, Canada, Russia, wherever it’s done) scientists will say that reaction gives off this much energy. Because it’s all done under controlled conditions. 1 molar, 1 molar, 1 molar, 1 molar concentration, atmospheric pressure, temperature 25°Centigrade. So this is a standard measure of free energy. That is how that particular situation is derived and why it is derived.
  • So whenever you see the ΔG° here, that is the standard chemical free energy given off by a given reaction.

Slide 13: Actual Free Energy Change

  • Now, you have to modify that because no reactions ever really take place at the standard conditions, this would be something that’s kind of unusual. So, the actual free energy change given off by any reaction is just ΔG and that is equal to ΔG° plus RT ln Keq. (RT times natural log of the equilibrium constant)
  • And this factor right here; is the factor which contains [A] and [B] and [C] and [D] in the actual concentrations, that is the concentrations that are present in your particular reaction. (NOT the standard conditions where everything is 1 Molar) This can be a condition such that [C]=1/2 Molar, [D]=3/4 mole per liter, [A] and [B] are 1/10 mole per liter each. So, that is the actual situation that’s an entropic term which modifies the standard and gives you the actual free energy that is released.
  • So, the ΔG here is the standard free energy change plus an entropic term for the actual reactant conditions. The basic phenomenon here is that you are going from a standard condition to an actual condition and you can then calculate the free energy that would then come from a situation where there is the actual concentration present. This is an important phenomenon because ΔG = -RT ln K`eq
  • When the reaction that you set up goes to equilibrium, then the equilibrium factor here is equal to the equilibrium factor here and then ΔG = 0. In other words, when the reaction has proceeded to equilibrium and is going in this direction and this direction (forward and reverse) in equal fashion at equal rates, then you no longer are getting any energy from that reaction. So, at equilibrium in a reaction there is a no free energy given off.
  • The moral of that story is that in biological systems the body sets up reactions so that there is never any equilibrium obtained. If you are in equilibrium, you are dead, essentially. There cannot be an equilibrium system because you cannot produce free energy and if you’re not producing free energy you can’t breathe, you can’t metabolize, you simply…you’re gone. =(
  • The moral of that story is that when standard free energy is equaled by your reaction here, the minus and the plus are the same and the free energy is equal to zero.

Slide 14: Biochemical Free Energy Changes