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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.
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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:
o light energy or
o heat energy
o 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.)
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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
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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 get A + B to go to C
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+ 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 + T ΔS 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
 There is one complicating factor, (and I’ll show you why that is a complicating factor) in
a sense that we need to have a standard free energy for biochemical reactions because
many biochemical reactions either utilize or give rise to a proton. In biochemistry, where
those reactions occur, the strength or concentration of the proton is nowhere near 1
Molar. If it were, the pH would be 0. You can’t live at a pH of zero, you could not
survive. Consequently, in biological systems we set up a standard now which is exactly
the same as the ordinary chemical standard except when you set up the standard, the
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proton concentration is given as 1×10 . In other words, a pH of 7. Everything else is the
same; 1 molar for everything else, except the proton. In this situation the actual free
energy given will be equal to the standard biochemical free energy plus the same entropic
term that we had here before.
Slide 15: Importance of this consideration
 Now, to give you an example of why that’s valuable. Well, first of all: Biological
systems operate by keeping a favorable ratio of reactants to products. And by that, I
think if you’re wise enough to perceive this: In order to make this factor here a minus,
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you have to have a high concentration of reactants [A] and [B] verses [C] and [D]. If you
do that, the log of that fraction less than 1 becomes a minus. If it is a minus then this
whole factor is a minus and that contributes to the free energy of standardization to give
you an increasingly minus figure here for the free energy.
So biological systems essentially get a big bang for the buck because they organize
themselves to make sure that the reactants are always favorably disposed. That is- you
have much more reactants than you do product and that drives reactions. That drives
reactions in the direction that you would like it to go. That is about as clear as I can make
that concept that favorable ratio of reactants to products. So any reaction may be made to
proceed in a given direction by the proper adjustments of that ratio of reactants to
products. You would like to have the reactants very high.
Slide 16: Free Energy Calculations
 An example of why you need to use a biochemical free energy system is: Take this
particular reaction which we talked about a long time ago.
 Acetic Acid ionizing to acetate anion and a proton. Now the equilibrium constant for that
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is 1.76 x 10 . The ΔG for that (if you remember the formula) would be minus the gas
constant times the absolute temperature times the log of this number here (which is
-10.94) This gives you 6.4 Kcal/mole.
 This reaction will never go because it has a positive ΔG. Why won’t it go? Because
under standard conditions the proton here is at 1 Molar, the pH is zero. So acetic acid is
not going to ionize, to give up a proton, against a gradient of 1 Molar protons already. It
is a weak acid.
 If we rearrange our system (and I won’t do the mathematics) you can set up a biological
ΔG°′ = ΔG° -9.5 kcal/mole because we’re taking in the negative, the log of a low
concentration of protons here to make our ΔG°′. Then, we have that minus the 9.5
kcal/mole gives us -3.1 kcal/mole. So, that tells you that under standard biological
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conditions that reaction will go. (when the concentration of protons is 1×10 )
 So all we’ve done here is we have a new standard of ΔG°′ composed of a new calculation
of standard conditions. The mathematics of that stuff is a little more than you need to
worry about but just keep in mind of why we have to have that thing for biology verses
standard chemistry.
Slide 17: Energy Transfer
 We have all of these notations. Keep track of these notations summarized on the board.
 Standard Chemical
 Standard Actual
 Standard Biochemical
 Actual Biochemical
 You’ll see those notations in all of your textbooks and will help you to understand what
you read.
 Energy acquired from nutrients and everything: you can’t use them immediately. For
instance, a mole of glucose would yield (going from glucose to carbon dioxide to water)
would give you a standard biochemical yield of -686 kcal/mole. A mole of glucose
would give you something like 686 thousand calories. At one step, if you eat that and
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liberate that type of energy right away, you would have the biggest fever that you could
ever imagine. You would be extremely warm and there would not be any way for you to
seek very much comfort.
So, we use the reactions that go from glucose to this and store the energy. There are two
classes of biomolecules which take up this energy. They are:
o Reduced coenzymes (NADH, FADH2)
o High-energy phosphate compounds
We are going to talk about them now.
Slide 18: High-Energy Biomolecules
 Biochemists consider a high-energy bond as one which yields a great deal of energy on
hydrolysis. That is why ATP is considered to be our favorite storage domain for energy,
because the anhydride-phosphate bonds that you find in ATP are extremely volatile from
the point of view that when they are hydrolyzed they yield the energy. That is why we
call them high-energy bonds.
 In chemistry a high-energy bond is a bond between two atoms which requires a great deal
of energy to break the bond.
 Biochemists look upon high-energy yield as a substance or energy that is released on
hydrolysis, on a reaction. As opposed to putting the energy into a substance and breaking
the bond. The bond is broken in a reaction and the reaction then gives off the energy in
biochemistry.
Slide 19
 We’re going to deal now with high-energy phosphate bonds, ATP is our major reservoir
(our battery so to speak). It has a side chain, adenosine triphosphate. (3 phosphate
groups). This phosphate group and this phosphate group (both bonds in red with yellow
highlighting) are joined together by an anhydride bond. 2 phosphoric acid molecules
have been utilized to make this and the release of a water molecule. This particular bond
(the phosphate to the methyline group) is an ester bond. This is a phosphate ester of
adenosine. This is a phosphoanhydride (red highlighted in yellow) and another
phosphoanhydride. These two bonds here are the ones that are the so-called high-energy
bonds. When they are hydrolyzed and water is added back and the phosphates leave,
high-energy, a great deal of energy is released. The amount of energy released, we will
talk about in a minute.
Slide 20: Phosphoric Acid Anhydrides
 ADP and ATP have these phosphoric anhydrides and they are largely negative free
energy change due to hydrolysis because of the electrostatic repulsion, stabilization of the
products by ionization and resonance factors and entropy factors which we will talk about
in this particular slide here (slide 21)
 For instance: if you hydrolyze the 3rd phosphate group from ATP, you liberate an
inorganic phosphate molecule, you liberate a proton. You now have three different
molecules as opposed to one original. That has created a little disorder. You have
relieved the charge-charge repulsion here and that is chemically very favorable. The
products that have been made, these substances organize water much less readily than
does the ATP molecule itself. Consequently, you have relieved charges, you’ve got more
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particles than you had before, and you have also release water molecules. So, you’ve
created a lot of disorder and you’ve also gone from molecule of high-energy to a
molecule of lower energy.
You can have a second hydrolysis. The 2nd phosphor anhydride can be cleaved to give
more energy for the same reason that was given for the 1st hydrolysis.
ATP has a double whammy with respect to the amount of energy that can be given and
we’re going to see on this next slide
Slide 22
(don’t worry so much about this)
 This is enzymology here that the standard biochemical energy from ATP hydrolysis is 30 kJ/mole (or -7.3 kcal/mole)
 However, the actual biochemical energy that you can get is larger than the standard. It is
52 kJ/mole (and-12.4 kcal/mole).
 Why do you get more energy from the system than the standard system would say?
o In cellular conditions, the reactant here (ATP) is far in excess of the products;
ADP + inorganic phosphate. So, the reaction has been set up in an actual system
where the reactants and products are not totally equal but there is much more
reactant material here than there is products. So, you will get more energy from
the reaction in that situation than you will from standard conditions. Standard
conditions are productive enough of energy. (30.5 kJ/mole or 7.3 kcal) But the
actual energy delivered by a molecule, 1 hydrolysis of a molecule of ATP, is -52
kJ/mol (12.4 kcal/mol). Precisely because of this entropic factor here. The ratio
of products to reactants. Products being here: [ADP][P] and reactants being in the
denominator.
o ATP is the reactant! Right? Right. Reactants: ATP. Products: ADP and
inorganic phosphate.
 There are highly reactive compounds in biology. ATP is not THE biggest reactant. It is
the one quantitatively that is the most important but it is not the most high-powered one.
 There are substances like phosphoenolpyruvate (which we will talk about next Tuesday
when we talk about glycolysis) there is 1,3-Bisphosphoglycerate (which we will also talk
about next week in glycolysis) and there is phosphocreatine (which Baggot will talk to
you about in terms of protein synthesis and protein metabolism. All of these compounds
have high-energy phosphate bonds (which are sufficiently high in energy) to actually
phosphorylate adenosine monophosphate. So you can have adenosine monophosphate be
phosphorylated by phosphoenolpyruvate or phosphocreatine or 1,3-Bisphosphoglycerate.
The molecules of higher energy phosphates will actually donate their phosphate groups to
the lower responsive element like adenine. ATP will actually phosphorylate substances
like glucose and glycerall. So, these are low phosphate energy compounds and this is
going from a rather high energy to a very low energy phosphate compound.
Slide 23: Phosphate Energy Levels
 These are low energy phosphate bonds and these are high energy phosphate bonds. This
type of phosphate bond is formed intracellularly in preparation for glycolysis (as we will
see next week)
Slide 24-25: Phosphoric-Carboxylic Anhydrides
 There are other carboxylic anhydrides and these are substances which you will run into
from time to time.
 This is acetic acid forming an anhydride with phosphoric acid so this acetyl phosphate is
also a high-energy bond. (-43.3 kJ/mol under standard conditions)
 And this is 1,3-Bisphosphoglycerate which has a high-energy phosphate
 This particular molecule here phosphorylates ADP during glycolysis to make ATP. It is
one of the ways in which you get ATP out of a primitive metabolic process called
glycolysis (which is an anaerobic process) We can make a little ATP in a primitive way
and that is done largely through this compound here.
Slide 26-27: Enol Phosphates
 Phosphoenolpyruvate (We will see more of this when we talk about it in glycolysis) But,
you can have phosphoglycerate. This is glyceric acid (glycerate) is phosphorylated at
carbon atom number 2. There is an enolase enzyme in glycolysis which takes water away
from this molecule and gives you this enolic form here. So, this is phosphoenolpyruvate
(PEP) and PEP is strongly enough exergonic when it is hydrolyzed to actually
phosphorylate ADP. ADP goes to ATP and then you have pyruvic acid left.
Slide 28
 The reason this goes so well is shown on this slide. You have phosphoenolpyruvate
giving up phosphate and the phosphate is lost, the pyruvate-enol compound is left behind
and it can then isomerizes to pyruvate. This isomerization reaction is really part of the
whole phenomenon. The loss of the phosphate is worth minus 28.6 kJ per mole. The
tautamerization is worth minus 33.6 kJ/mole. This a very robust reaction.
The rest of the slides are not really pertinent but I want you to keep in mind (as we go
forward with biochemical metabolism) that the larger the negative number, the more
robust the reaction is. In other words, the reaction, if it has a minus 112 kJ/mol is a much
stronger more robust reaction than one that has a minus 2 or minus 3 or minus 4 or minus
5 kJ/mol. The vigor with which the reaction occurs is indicated by the absolute number
here on the delta G.