This is what's new in my life.

This kitten can be yours! His name is Jerry aka. Buster Bluth aka Catface .... He's incredibly sweet and I'm fostering him for the time being. He needs a permanent home!

This photo is sideways..

Attempting to claim the dog's couch...

What's up with all these phosphates, professor?

I think I talked free energy to death in the last two posts... Enough to be able to only need to mention ONE more time that it deals with which direction a reaction is going to be favorable in.

Reactions' abilities (or qualities?) of being favorable are really fucking important in terms of biological processes because for shit to happen in your body (like protein synthesis and shit) it has to be somehow favored. Otherwise...it won't happen.

Sometimes, things that happen in our bodies aren't...favorable...Remember once again terms like endergonic and exergonic. Endergonic relates to things that are bad for -ΔG...so they deal with processes that are +ΔG which means the reverse of that reaction is actually what's favored. And it wouldn't really be good times if say... the reaction of "Amino Acids ----> Proteins" was going in reverse ALL THE TIME if it were +ΔG for it to go forward, because then we'd just be spontaneously turning into amino acids and our proteins would be soup. That wouldn't be good. Nothing would get anywhere.

But sometimes, that's what's up, and reactions that get shit done are endergonic. So what can we do?
These reactions, since they're like, necessary for life and shit, are going to be coupled to reactions that ARE strongly favored in the forward direction.

Nice and BIG. So here we have two reactions, one which is ΔG+ (Thing ----> More Complicated thing) And one which is ΔG- (This ----> That)

If that's what's going on in the cell and the cell needs to do the first reaction, it can couple both the reactions together, adding up their ΔG's, and from there you get the cell converting This Thing to That Complicated Thing. That was totally unintentional when I was writing it on the board but it came out making sense-ish. And it's funny. To me.

In the body, the coupling of reactions is used to make things happen...bring things across membranes that membranes won't pass lazily, nerve impulse travellings, muscle contractions...

So what's up with all these phosphates, man?

For all this shit that's highly ΔG+, we need to make sure we have things that are ΔG- enough to make shit happen. For that, we have things called "transducers", which are basically things that convert stuff into slightly different stuff. These phosphates that we're going to talk about in a second are things that have all that highly negative free energy that's going to let other shit happen by coupling it to the stuff that happens with the phosphates. The phosphates undergo "hydrolytic release of phosphate groups" which basically means, in water, phosphate groups are released. The hydrolysis of these compounds with phosphates has highly negative free energy, enough to allow other things to happen.

Oh god, what is this chart?

So, THESE are THOSE COMPOUNDS that have phosphates that are released in hydrolysis, which provides us with enough negative free energy to couple shitty ΔG+ reactions to. And on the next bage, there's a second, almost as intimidating table, except smaller and with less colors..Oh god, can we do this one big word at a time?

Hydrolysis of _____ Gives us ____kj/mol ? Will that help?

PEP -61.9 kj/mol

1,3-BPG  -49.4 kj/mol

CP  -43.1 kj/mol

       ATP ... AMP (-45.6 kj/mol)

ATP ... ADP (-30.5 kj/mol)

PPi -19.2 kj/mol 

(And the less exergonic ones...)

AMP -13.8 kj/mol

G-1-P -9.2 kj/mol

So I guess we can call Pyrophosphate our cutoff point for...exergonicity?

ATP is kind of really important so lets look at successive hydrolysis' of ATP and then the things that come out of it, which are also hydrolyzed with phosphates falling off all over the place.

  So there's ATP. It's a Ribose ring (Note 2' OH group intact), with an Adenosine nitrogenous base, and three phosphate groups attached at the 5' end.

One thing we can pay attention to is which bonds are the phosphoanhydride bonds and which bond is the phosphate ester bond. Only one phosphate ester bond in ATP. Two phosphoanhydride bonds.

When you hydrolyze ATP, you get ADP (adenosine diphosphate) which is then hydrolyzed to AMP (adenosine monophosphate) which is then finally hydrolyzed to Adenosine by its sad lonely self.

With each round of hydrolysis, you get a Phosphate group released and some negative free energy being released as well. This is when a NEW ΔG term comes in (for this chart), which is ΔGs' (delta G standard-prime). This, also will be discussed, in a second. 

Remember talking about K and Q and equilibrium constants and all that crap in the last post? Well....with the whole ATP thing, and the ΔGs' value being -30.5 kj/mol, we're discussing a reaction where the equilibrium constant is greater than 10^5....which means that the equilibrium lies way far to the right, so far as to the point that we can consider this reaction irreversible. Now, remember how in BIG SCARY TABLE, some things were really really exergonic while some things were less exergonic? Now we can talk about why. In organic chemistry 1, the big "joke" and also the serious thing was that "The answer is always resonance". Well...let's talk about the resonance in this situation of figuring out why some things give us all this negative free energy when they are hydrolyzed and others...not so much.

Look at this thing that gets released, which is called orthophosphate ion (HPO4 2-)

Look how stable that sucker is. All these versions of thing thing are of equal energy, which means high entropy. Why does it have higher entropy? Remember, entropy likes randomness. When the Pi is bound up in an ester (like, before it gets released), it has fewer possible resonance forms, which means it's less resonance stablizied (okay that's kind of redundant), but that means that when Pi is released from underneath it's mother's wing, there's more entropy, which is favorable.

Another factor as to figuring out the reasons behind the difference in exergonicity of stuff that's hydrolyzed and why some stuff has more free energy than other stuff is that the stuff that's the products of hydrolysis can then ALSO be hydrolyzed. When the Pi gets released, especially if the other thing that's being release has a charge, that's good times for hydration land. IONS LIKE TO BE HYDRATED.

A third factor is backed up by us remembering that like repels like and opposites attract. Lets look at some of these phosphate-having compounds that get hydrolyzed and how there is repulsion between the like charges that they have inside of them.

It might actually be in my best interests that this is a bit blurry since, there's bad words in it...But this just shows how in these compounds that get hydrolyzed, hydrolysis is favorable since it lets all these negative charges stay the fuck away from each other. The charge-charge repulsion forces are strong with these ones. 

The next thing deals with keto-enol tautomerization...One of the things we learned about in Orgo 2, happening to be one of the questions I blew on the first Orgo2 exam and then later when Lenny yelled at us, it was really drilled into my head to understand what the fuck is going on. Because rule #1 is to LISTEN TO WHAT LENNY SAYS or else you're going to fail life and not be a doctor. So yeah...keto-enol, I think I can draw that shit in my sleep now.

In the little box on the left I drew the typical example of a keto-enol tautomerism, and now let's talk about what the fuck the deal is with phosphoenol pyruvate and why the fuck the ΔGs' for this thing is a whopping -61.9kj/mol...Well THEN. When PEP is hydrolyzed, what's up is that there's a keto-enol tautomerization that occurs in the product. Since the DIRECT product is in the enol pyruvate form, it quickly tautomerizes to the keto form of pyruvate, which is thermodynamically favored. The extra ΔGs' negative-ness basically swarms in from the goodness that comes out of the tautomerization. 

Another thing we see in these hydrolysis reactions that release Pi when the compounds are hydrolyzed is that an H+ is released. This is where I guess that second less scary chart comes in (didn't take a picture of it, but we're gonna talk about it now) where they bring in how pH factors into some of these compounds, not others. 

Here's what's good with ATP hydrolysis, and the ΔG that we can calculate from it. We usually deal with reactions that are around neutral pH so we need to talk about the chemical potentials (measure of how much that thing contributes to ΔG) of the water and the H+ in a different way than we did before. In the previous situations, with biochemical "standard state" stuff, the activity of water is not affected by any reactions that are going to use it up or produce it...another way of saying this is that in the biochemical standard state stuff, the "activity" of water is "unity". While standard state concentration is 1 M solutions, in our bodies, the concentration of H+ is 10^-7 which is a LOT less than 1M. So now we need to start looking at the H+'s contribution to ΔG as a "thing" that is found at 10^-7 M and not as something that's going to be a standard 1M concentration. 

So in this new state that we're talking about (living shit state), 1M is not going to be considered unity for H+. We're now going to call H+'s activity unity when the concentration of H+ is 10^-7 M. THIS is where this millionth new way of talking about ΔG (this time, we add two things to it, to make it ΔGs') comes in. And now we can plug some pseudostuff into that equation:

Now, we need more numbers, make stuff more real for us. You can tell that this stuff is real because it has NUMBERS and an ANSWER which is what "matters in life". Having answers. Which, we won't always have, in life.

Another important thing when we look at that big scary table of phosphate-containing compounds and how those compounds can be hydrolyzed and release a phosphate, is that some of them have something called phosphate transfer potential which is important. Phosphate transfer potential havingness means that some compounds can phosphorylate others. This potential is defined as -ΔGs'. So looking at bigscarytable again... 

See that blue thing on the right that is labeled TRANSFER POTENTIAL? And it even has a SCALE that tells you what's good? Well our good friend PEP is at the top of the chart..remember this is the thing that gets all its bragging rights from keto-enol tautomerization....So, PEP is at the top of the list, with the highest transfer potential, and that sad G-1-P is at the bottom of the list with a weaksauce transfer potential. Things that are above things can drive the phosphorylation of anything below them as long as you can couple stuff. Coupling is done by making the reaction occur on the surfaces of protein molecules, like enzymes. Let's look at these coupling reactions for this stuff: 

Hydrolysis of PEP is ΔGs' = -62 kj/mol

Phosphorylation of ADP    = +30. kj/mol

So when you couple, phosphorylation of ADP by PEP is going to be ΔGs' = -31.4 kj/mol

PEP has more phosphate transfer potential than ATP, SO THAT MEANS it can give ADP a phosphate group and it's going to be thermodynamically favored and YOU'RE GOING TO LIKE IT. Likewise, ATP can give glucose a phosphate group (I'm calling phosphorylation "giving") because the phosphate transfer potential of ATP is greater than that of G6P:

ATP Hydrolysis is ΔGs' -30.5 kj/mol

Glucose Phosphorylation is ΔGs' +13.8 kj/mol

So when you couple, the phosphorylation of glucose by ATP is ΔGs' = -16.7 kj/mol

Finally, the chapter 3 ends off by talking about ΔGs' for REDOX reactions and I'm going to make the last chapter 3 post discussing that shit.

What's in the real world? Sad dogs and really mean people.

Also, the glycolytic pathway is in the real world.

Glycolysis? What's that? Something Lenny can probably draw out in full in his sleep, without waking up.

But for US, it's what converts glucose that we eat into pyruvate. The free energy (shit, there's that word again) that gets released during glycolysis is used to form high-energy compounds ATP and NADH. I have a feeling I'm going to really, really like talking about this, or I'm going to dislike it so much that I'm going to crawl into bed and cry immediately after.

But let's see !

This section talks about the second step in the glycolysis pathways..isomerization of glucose-6-phosphate (G6P) to fructose-6-phosphate (F6P). I've conveniently drawn out the structural forms and forgot to label the drawing on the right as F6P but that's what that is. For anyone who doesn't know the word isomerization...isomers are compounds with the same molecular formulas but different structural appearances. They have the same crap in them but they look different. 

G6P <----> F6P is ΔGs = +1.7kJ/mol

Remember words like endergonic and exergonic? And remember how we keep forgetting what they mean? Endergonic processes are related to ΔG+ reactions while exergonic are related to ΔG- reactions.

This reaction, G6P <--> F6P is endergonic. When both products and reactants are at 1M concentrations both, since ΔGs is positive, we end up with the equilibrium lying to the Left, with "more reactants than products", so favoring the REVERSE reaction, being F6P < ---- > G6P

What kind of numbers are we talking about here? Just how MUCH is it unfavorable to make F6P? The next three photos are really slow math, going back to formula we've talked about and incorporating this concept INTO it.

Here, I'm just showing what we're going to be solving for...the equilibrium constant thing....which way things want to go.

Here, I just show that we're still talking about ΔGs stuff. I always get lost and forget where numbers come from so it's a good idea to keep pretty good track of that.

Now, let's remember that we're not at equilibrium cause if we were, we would be dead.

And such, the concentrations of crap inside of our cells isn't at equilibirum, either. 

A little while ago, I said that K and Q are "sort of" interchangeable. Well, here's a clarification:

Since ΔG isn't zero, that means it's either + or -, meaning the reaction is either being driven forward or in the reverse direction, which means....something has to be driving that reaction.

That puts things more into perspective.

So how can we maintain that Q < K, and form the products of the reaction as written favorably?

The system can consume the products as soon as they are formed, such that it makes the product concentration low compared to the reactant concentration, which will drive the system to make more of the product in turn. When you remove the product, it'll drive more of the reactant to make more of the product.

You can also maintain Q < K by keeping the concentrations of reactants high. Keeping the concentration of the reactants high will drive the reaction to make more of the products. 

We can make the reactions that are thermodynamically unfavorable go by either keeping the Q < K situation going (by continuously removing products or continuously supplying reactants) or by coupling unfavorable reactions to favorable reactions. Types of favorable reactions to which unfavorable reactions are coupled are ion transport across membranes and hydrolysis of ATP. 

ΔGstandard, ΔGshmandard

Last time I made a bit of a mess of things trying to talk about them...Let's maybe not be messy tonight.
I need to finish talking about chapter three so I can move onto four and five. Some of the pictures have not been loading and I think that's just cause I've been copy and pasting them and maybe this site doesn't like that, so tonight I'll try to go through imgur.

When you've got a reaction and you are talking about which direction it's going to be favored to go in, you use terms like enthalpy (the energy of a reaction) and entropy (randomness of things that are happening, freedom of molecules, etc). 

The formula ΔG=ΔH-TΔS allow us to calculate the ΔG "Free Energy" of a system, which is in turn going to tell is if a reaction will be favored or if the reverse of a reaction is favored. Typically, reactions that are ΔG- are favored to go forward and if a reaction is ΔG+, the reverse of the reaction is favored. And all THAT depends on the interplay of entropy and enthalpy (TΔS and ΔH)

So when we take the first reaction in the book, fermenting glucose to ethanol, we see that ΔH is -82kj/mol.

Let's remember how messy we were last time. 

The reaction Glucose ---> Ethanol is ΔG - 218kj/mol

The enthalpy is strongly negative value, which is good for ΔG

The entropy happens to also be a negative value, which is good for ΔG

So ΔG gets to be negative and the reaction proceeds in that direction.

Thanks to Lenny, I can finally understand the term "kinetically stable" (metastable). It's used to describe something that is stable and not going to be torn apart by hydrolysis...at the moment. The nucleic acids chapter talks about how DNA is metastable. Eventually that shit will get hydrolyzed and we'll get broken down and things are going to eat us apart...but not yet. Kinetics deals with rates, and that's something we can't get ΔG tangled up in. Page 67, the textbook says: "The favorability of a process has nothing to do with its rate" and goes on to say that large negative free energy doesn't mean that it's going to happen fast....but it IS going to happen. 

ΔG's sign tells us if a process (or process's reverse reaction) is favored

ΔG's magnitude tells us how far the process is from being at equilibrium

ΔG also tells us how much useful work we can get out of something...it can sort of tell us if a process is going to happen, for what it's going to be used. Sort of. 

A different way to talk about this shit explains how the ΔG of a system depends on the concentrations of products/reactants you got going in a reaction.

That beast.

The book plays a game called not defining ΔGs (ΔG standard, too hard to find a way to put a degree sign as I keep just copy & pasting the Δ) until the next section, so we can play along.

This equation is supposed to give us the energy to do work from a reaction that happens. 

The next part is Chemical Potential. 

We start off by talking about the relationship between ΔG and what crap you have in your "flask". The chemical potential of a thing you have is the measure of what it contributes to the ΔG of the system. 

When they say "state function" they mean it depends on the initial and final states for what is going on in your reaction. 

Let's take this for example, something the book talks about

You've got two solutions of stuff that's separated by something through which stuff can pass.

Let's say this is what you have on Monday. On Monday, there's lots of stuff in area one, we call the concentration A1. There's less stuff in Area 2, we call the concentration of that stuff A2. Then you wait until Tuesday, for some of the stuff to move into the "bottom" region. The direction that the stuff is going to move in is going to help us evaluate ΔG

Alright. Now it's Tuesday, and our stuff has moved. Now we have about the same amount of stuff in both regions. 

When we're looking at the two days individually, we can compare their free energies with the formula from before, one formula per day, and then use the concept something being a state function and dependent on the final (tuesday) and initial (monday) states, apply all that to the formula. When the ΔGs values cancel out and you know how to use log/ln properties, we're left with:

ΔG = RT ln ([A1]/[A2])

This formula is pretty fucking important.

If the concentration of A2 is less than the concentration of A1, ΔG is going to be negative, meaning all the stuff that was transferred was favorable. 

If, however, the concentration of A2 is more than the concentration of A1, ΔG is going to be positive, meaning the movement of stuff is unfavorable into the "lower region" but the movement of stuff out of there IS favorable.

If the two concentrations are equal, ΔG is zero meaning the system is at equilibrium (which is actually what's going on Tuesday)

If some stuff can pass through a "membrane", the direction that is it going to want to go is going to be from a place with LOTS of stuff to a place with NOT SO MUCH stuff. It's going to go from a place where the chemical potential of stuff is high to a place where chemical potential of stuff is low. This chemical potential is the driving force behind this. It's going to move the system toward equilibrium, where the concentrations of stuff on both sides of the "membrane" are equal and there will be no more driving force since ΔG will be zero. 

In situations where stuff happily goes the reverse way, from LOW concentration to HIGH concentration, there will be a free energy price to pay and we can pay it by "coupling" this movement of stuff to some type of reaction that is thermodynamically favorable. The book says: "Two or more reactions are coupled when one reaction cannot occur without the other reaction also occurring"

We can wiggle around that last ΔG formula

ΔG = RT ln ([A1]/[A2])

And change it so that it reflects a reaction that is occurring with products and reactants and shit.

The terms Q and K become sort of interchangeable and it is basically the equilibrium constant for the reaction...the concentrations of Products/concentrations of reactants such that, for a reaction AT equilibrium, where ΔG = 0....

ΔG = ΔGs + RT ln (Procuts/Reactants)eq = 0

0 = ΔGs + RT ln (Products/Reactants)eq (and we can move ΔGs to the other side to make..)

-ΔGs = RT ln K

The book calls ΔGs a "reference value for free energy change" ... you compare the free energy changes in reactions under equilibrium circumstances. 

All of THIS and especially reaction coupling is what's going to get me into the next post. Using real worlds. Not Monday and Tuesday worlds.

Rewinding in the brain map

There are things that sometimes click in my brain, literally, like someone turned on a lightswitch. And this mostly happens in Lenny's sessions, or when I'm reviewing things after Lenny's sessions. And it feels like my mind is rewinding or maybe even relapsing back to previous parts of my life where I learned things, and realize that I'm relearning something that I should have already learned.

Except that I feel like high school was just five-hour-long memorization that didn't mean anything. I've been able to tell the difference between "I've learned this before, but I don't remember what it is, so now I need to relearn it" and "I've learned this before, this is that thing" and it's kind of shocking how much more I've been experiencing the former. I don't know if it's the educational system or if it's me, I'm hoping it's just me now learning how to really learn things so that they mean REAL THINGS to me, so I can quantify or illustrate or just be able to teach someone else what it is that's up. This might make no sense, but it's kind of both empowering and demotivating at the same time. I haven't added a meaning to it yet, so it's just currently this thing sitting on my table. On every table I sit at or near.

Enthalpy? Entropy? Enthalpy shmenthalpy.

So last time we talked about constant volume and constant pressure scenarios for carrying out a reaction.
When we want to see how a reaction occurs outside the body (but resembling what happens inside the body), we used constant pressure systems. The difference is going to lie in the heat that comes out of doing the reaction in the scenarios.

The term enthalpy is used to talk about the change in heat that occurs in a constant-pressure reaction, and we measure it by the internal energy and PV (pressure * vol).

We talk about state functions, and how all this shit exists as a function of states...I just look at it by seeing what's up in the initial state vs. what's good in the final state of what you're doing in these reactions, whatever they may be. That's where all our Δ's come in, because we're looking at initial conditions vs. final conditions.

That ΔH value we get is what we got after we did all the math. When we measured the heat of reaction at constant pressure, we're actually measuring the change in enthalpy.

The book then goes on to start confusing me a little bit, since up to here I can sort of follow it. We then talk about how in most cases the difference between ΔU (internal energy change) and ΔH (change in enthalpy, or, change in heat that occurs in constant-pressure system) is practically negligible and to think of ΔH (enthalpy) as a measure of the energy change of a process. So enthalpy is change in energy? Google...help me out.

en·thal·py

ˈenˌTHalpē,enˈTHalpē/

noun

PHYSICS

1. 1.

   a thermodynamic quantity equivalent to the total heat content of a system. It is equal to the internal energy of the system plus the product of pressure and volume.

   
   

Okay.

Now .... entropy



We use entropy to answer questions like that. Under a certain set of circumstances, some processes might want to go one way, and when you change the circumstances, they'll want to do something else. Like melting of water when it's warm and reforming of ice when it's cold.

But some things cannot go back and forth, they're irreversible. Once you introduce a certain set of circumstances to a certain thing, a reaction will take place, and you'll never be able to go back to your starting reactants.

Uhmm...this may be a bit morbid...but that's my cat. She used to be a cat. Then she died, and we sent her body to the crematory agency, and got back her ashes (and Dana helped me make a clay paw).

When you introduced the conditions of FIRE to CAT, you got ashes, and you can never have CAT back. How's that for irreversibly sad?

On the other hand, we have reversible processes, which will want to be near a state of equilibrium, being, the lowest possible energy state that a system can be in. Lower energy states of systems are favored over higher energy states, which is why systems want to be in lower energy states, therefore, they want to be at equilibrium. Irreversible shit is set up far from equilibrium, but it wants to be at equilibrium, so that's where it drives. I guess my dead cat is now at equilibrium. 

In thermodynamics, favorable and spontaneous mean the same thing...it's just what the system wants to happen. But that doesn't mean that the favorable thing is going to happen quickly. It's just GOING to happen, we're not talking about rates of reactions, just what's happening IN them.

When you put water in the freezer, it's going to freeze. Just not right away. It's about what's favorable for the water in the freezer. Not how fast the water freezes.

Thermodynamics just has conversations with itself as to what is favored in a given situation.



Let's get a little philosophical and a little landmarky. I once asked my co-workers son what his favorite color was and why. He said: "blue is my favorite color because blue is my favorite color" and it reminded me of the difference between deciding and choosing as discussed in the landmark forum. But biochemistry isn't about choosing chocolate because it's chocolate, the favorable thing isn't the favorable thing because it's the favorable thing....there's more to it than that.

We can try to say that the favorable thing is the favorable thing because that's the lowest energy state of "thing". But minimal energy isn't going to explain everything. When ice melts, energy is absorbed.

So there has to be something else involved. The book talks about layering water on top of a solution of sucrose (a dissaccharide made of glucose & fructose). The stuff in the sugary solution is going to want to spread out throughout the newly added aqueous parts. There's not going to be changes in energy, no changes in heat, and no changes in work being done in the situation. The sugars are not going to decide to gather together and become a different layer. This is where this entropy crap comes in. Things are going to have a natural tendency to be random. And that's what entropy is, it's a "measure" of randomness. How much disorder there is in a system.

In terms of comparing energy states of systems (like let's say a bunch of legos), the entropy of an ordered state (like having them stacked up tightly to make a complicated wall) is lower than the entropy of a disordered state (having all the legos randomly dispersed throughout your living room).

Let's just, for now, say that shit wants to be random. That's where it's happiest. That's where you get high entropy.



We now can get into the second law of thermodynamics, stating that in an isolated system (where you put nothing in and nothing out), the entropy (amount of randomness) will increase to some...maximum.

Things aren't going to get more orderly on their own....not without...help



So...we don't really deal with isolated systems here. We deal with open systems. The world is a closed system but all the shit IN it is open systems. I put nutrients in the form of dog food and carrots into my cockroaches and they shit all that out as feces. Open systems :)

And in open systems, you need to account for both energy and entropy changes in the crap that happens to things living in the world, if you want to study it properly. That's how we ...somehow..get to Gibbs Free Energy, which is that thing that I kept saying that I don't believe in.



The thing Lenny talked about some, was that you cannot possibly directly measure this value. It's just a thing that's dependent on other crap that you can SORT OF measure, and then it gives you the free energy thing... ΔG...which is what we're going to use to explain if something is "favorable" or "unfavorable"

Free energy refers to the amount of ΔH that you've got that is available to be used to do work.
So what makes a process favorable?

Lower energy things, where ΔH is negative

Increases in entropy, where ΔS is positive

That's how you get favorable processes. Both of those situations will make ΔG negative which is what describes a favorable process. 

This is a sentence from the book: "...for a favorable process in a nonisolated system, at constant temperature and pressure, is that ΔG be negative..."

THAT MEANS that if you have +ΔG, the process is NOT favorable, and the reverse of that process is favorable.

You then get words like exergonic and endergonic....

Exergonic processes are ΔG- 

Endergonic processes are ΔG+

From the formula ΔG = ΔH - TΔS,

if ΔH and TΔS balance each other out, ΔG will be zero, meanings the processes is not favored in either direction, meaning the system is at equilibrium. The process is going to be reversible, and it can be pushed in going either direction by some external push.

In this next part, we're going to talk about how entropy and enthalpy play together when discussing liquid water/ice.

Here's how we're going to do this. We're going from ice -- > liquid and we have two different y-axes. 

In ice, there's a certain number of bonds between molecules. When we melt ice, you're breaking some of those bonds. The difference in enthalpy between ice and liquid is similar to the energy you put into the system to break the bonds. So, the enthalpy change for ice --> liquid is positive:

Next, we want to consider what is happening with entropy in this system. There is a change in entropy because ice is more ordered than liquid. Water molecules get to play around doing their happy thing, while in ice, they are structured in an orderly fashion. So, when we go from ice to liquid, the entropy change is going to be positive because we're increasing randomness in the system. 

This gives us some freedom to talk about gibbs free energy and the rest of the shit in its equation.

When we talk about ΔG for ice to water, when the temperature is low (and you have ice), ΔH (enthalpy) is going to dominate, meanings the ΔH term is going to be greater than the TΔS term, meaning ΔG is going to be positive....THAT means that since ΔG is POSITIVE, ice becoming water is NOT favored at low temperatures.

However, when we have higher temperatures, above the melting point of ice, the ice cube is going to melt.

ΔG is going to be NEGATIVE because higher temperatures, the TΔS value is going to be higher than the ΔH value when the T is large enough. So, at higher temperatures, the process of ice becoming water is favorable.

And up here, we can also see that at the melting point, 273 K, ΔH and TΔS are equivalent, and ΔG is going to be zero. This means the system is in equilibrium and any changes are reversible. 

The melting point of something is the temperature where ΔH and TΔS lines intersect, which is the temperature where freezing is in balance with melting. 

So when you're looking at reactions and you want to figure out what the favorable direction is, you're going to be looking at both entropy and enthalpy, and temperature is definitely also a factor. 

Here's two important figures from the book:

This one has there processes. In the first one, both enthalpy and entropy work together and make the reaction favorable. In the second, the reaction is enthalpy favored (since ΔH is negative) but entropy opposes the reaction because entropy is a positive value...which you can say is bad for ΔG. However, ΔG still ends up being negative. Because enthalpy won, the reactionis enthalpy driven. 

In the third process, you have an entropy driven situation. The enthalpy is a positive value, which is bad for ΔG, since it'll try to make ΔG positive, but we have enough entropy in the TΔS value (-140) to still have the reaction be favored and have ΔG be negative. That's why it's entropy driven. Because entropy won.

This is the second important figure. It's basically a summary. But you can't just fucking memorize this shit because it's going to make no sense if you don't understand it.

At low temperatures and high temperatures, the reaction is ONLY going to be favored if ΔG is negative. Having a positive ΔG makes the reaction not favored.

The rest of chapter three, I'm going to talk about in the next post.

Let's talk about Energy...and other crap

While my dad spends the day in bed watching old Tom Hanks movies, I'm going to sit here until I understand biochemistry.



We start off this chapter talking about energy, heat, and work in terms of a system and all the shit that's around it...the surroundings. A system can be isolated, closed, or open.

An isolated system can't exchange anything with it's surroundings...not energy and not matter. I equate this to forming a bacterial lawn on an agar plate and letting the thing grow, but not taking anything away (waste products) and not adding anything (like extra nutrients) while the bacterial colony undergoes whatever processes it undergoes.

A closed system can exchange energy but not matter with the surroundings...That's basically our planet.



We don't really exchange much with outer space short of sending out the occasional astronaut or satellite

A system can also be open, exchanging both energy and matter with the surroundings.



A system's internal energy is basically all the energy that it exchanges by means of physical processes and chemical reactions. This includes whatever it's atoms/molecules do, the energy stored in bonds (that can then be broken/reformed), and the energy of non-covalent interactions (the crap that I talked about earlier....charge-charge, charge-dipole, etc). Those are all associated with a certain amount of energy.

A system's internal energy depends on the thermodynamic state.



The thermodynamic state is defined by asking questions like...how much and what kind of crap is in your system? You also use 2/3 variables to define it, including temperature, pressure, and volume.

Only isolated systems cannot exchange energy with the surroundings. Exchanging energy means changing it's internal energy, thereby giving us ΔU



When a system is closed (like earth), exchange of energy and thereby change in internal energy occurs by either transfer of heat or work done. The work includes either the system doing work on the surroundings or the surroundings doing work on the system. When you exert a force against a resistance, you are doing work.



When we take the energy that can be exchanged in a closed system, separate it into heat and work, you get positive and negative values of both.

When you have +q, the system is absorbing heat from the surroundings. -q means the system is releasing heat into the surroundings.

Likewise, when you have +w, work is being done by the system on the surroundings and when we have -w, work is being done by the surroundings, on the system.

Bringing this into more of a real thing perspective, when we eat crap, we're ingesting glucose. That gets oxidized to carbon dioxide and water, which has an energy change associated with that particular process. The energy released, we can use in the form of heat or work. It's kind of neat.

All of this ties into the First Law of Thermodynamics.
In a closed system (like the earth), internal changes in energy occur via heat or work that is exchanged with the surroundings.



So we get that formula, for overall change in internal energy, dependent on heat and work done.

If we consider a system that absorbs heat at the same rate that it does work on it's surroundings (where both q and w will be positive, and equal to each other), you'll have no overall change in internal energy because everything evens out.

Any change that could occur in internal energy is going to depend on the initial and final states of the system and is going to be "path independent" meaning if there are two ways to accomplish something, that is irrelevant and the only thing we care about is the change in energy. But the amounts of heat and work individually that are released/absorbed by the system are going to depend on what happens between the initial and final states....that shit is path dependent. This way, we get into this picture that I'm not going to draw, there is no way in hell, so I'll just take a photo of the book:

The book takes these two experiments that are basically doing the same thing...oxidizing one mol of palmitic acid. The reaction is basically

1 mol acid + 23 mol Oxygen --> 16 mol CO2 + 16 mol H20

The book talks about running this reaction in two ways, and then figuring out which of the ways is more alike to what happens when the reaction occurs in our body.



First, the reaction is run at constant volume, which is definitely not our body state, because if that were the case, we would explode.

This is done using a bomb calorimeter. So you've got some palmitic acid with some oxygen and you ignite the sucker in a sealed up box that's inside of a water bath, which has a thermometer to gauge the temperature difference in the water.

When you ignite the insides, you measure the heat that goes from the system (closed up bomb container) to the surroundings (the water with the thermometer that's measuring this). Since this shit is constant volume, no work is being done to/by the surroundings...so in our

Δ U = q-w

Δ U = q - 0

Δ U = q

The total change in internal energy that's going on here is just the heat that goes from the bomb thing to the water bath around it....At constant volume, the internal energy change that is going on is the heat that is released from the system to the surroundings.

Heat is measured in joules and the value of ΔU here is going to be -9941.4 kj/mol of palmitic acid.
The value is negative because the reaction is releasing energy that's stored in bonds. As heat in this situation went from the bomb thing (system) to the water bath (surroundings), the energy in the system decreased.

If you do this same reaction at constant pressure instead of constant volume, you get something different going on.



In this case, when you ignite the bomb thing, instead of being constrained by volume, there's this piston thing that move up and down. The bomb thing (I'm just going to call it bomb thing) is free to expand and contract because of the piston thing when you ignite the sucker. So you ignite it, the reaction happens, the piston moves up, then down, and finally we have it contract in proportion to how many moles of gas we LOST. Remember, we started with 23 mol of oxygen and we get only 16 mol of CO2 in our products. Since we now have less gas volume, this means that work was done by the surroundings on the system.



So in this constant pressure system, we can use some formulas and shit to get to a good way to calculate the work that is being done.

Work is done when volume is changed against pressure, and using PV=nRT, where R is the gas constant, T the absolute temp (K), and n is the number of moles changing, we get all that crap above.



And here is the actual calculation of the value of work, in J/mol of crap, which you can then convert to kj/mol of crap in the next photo...



Uhhhh so this is kind of important. You're figuring out some stuff here which differentiates the two scenarios.

You're kind of moving the formula ΔU = q-w ....since you already have the ΔU (from the constant volume experiment) and you have the w which we calculated from the previous part, we can get the q value by adding the change in internal energy to the work that was done in this constant pressure thing.

What this means is that under constant-pressure systems, we get some more heat released to the surroundings than in constant volume. In the constant pressure version, the surroundings are able to do work on the system, and you can measure this work as the extra heat that gets released from the system.

This is where all the path independence/dependence crap comes from. Heat and work that are varying in systems depend on path but the internal energy doesn't.

In the next post, we'll talk about the places where I get most confused...enthalpy and entropy. Entropy not so much, but I still get sad with enthalpy because I feel like I understand it wrong.