How organisms go about acquiring nutrients is called trophism and is variable between living things. There are two broad categories that organisms may fall within, those that use light energy and those that use chemical energy. The strategy an organism uses to complete this life process is one criteria used to classify them into different branches of the phylogenetic tree. tree of life. While the autotrophs make many of the previous reservoirs available, they are not significant agents of disease.
We will briefly discuss them here, but be aware that the remainder of the chapter will focus on heterotrophic strategies. Autotrophs are self feeders and may use either inorganic light or chemical energy to produce foods. Phototrophs include organisms like cyanobacteria that are significant producers of oxygen and sugars. However, it is also possible to use chemical energy to produce foods.
Remember there were successful living cells millions of years before the first photoautotrophs. The The exceptionally large Thio margarita is an example of a chemoautotroph. Their nickname is sulfur pearls, and they harness energy from their sulfur reservoirs. There are also lithoautotrophs that help to break down and weather rocks.
We will focus primarily on heterotrophic processes in this chapter, because these are the ones that interact with us. Either through positive mechanisms, like gut microbes freeing up metabolites for us to use, or negative mechanisms, disease. All of the processes we are about to describe are examples of chemoheterotrophy, meaning that energy-rich organic chemicals will be harvested from the environment.
But before we discuss the actual pathways, it is vital to understand the underlying rules of thermodynamics. Understanding the kinetics helps us to better appreciate why things have to happen. as they do in the metabolic pathways.
This will involve a brief overview of thermodynamics. Thermodynamics allow us to describe changes in energy. Energy is defined as the capacity to do work.
However, what is work? Work is further defined as the capacity to move or change something against an opposing force. Can you think of some examples? examples of opposing forces. Gravity, friction, and even concentration gradients are all examples of opposing forces.
For the biological sciences, we focus on the first two laws of thermodynamics, conservation entropy. First, when we talk about energy, we talk about it being transferred or transformed. Energy is neither created nor destroyed. It simply gets converted.
Second, Second, from a universal standpoint, energy moves from an ordered to a disordered state. This is best observed through concentration gradients. Think about dropping food coloring into water. Even without doing anything, the food coloring will diffuse from a concentrated drop to a more dilute state.
Now can you create a more ordered state? Can you get the dye back out of the water? Yes! But you will have to invest a large amount of energy to do so.
Free energy is the way that scientists combine these two laws in order to describe how energy works in a system. This was explained in the 1870s by the work of Willard Gibbs. For this reason, the way that we calculate free energy is to use the energy of the system.
changes in energy is called Gibbs free energy and is represented by the variable G. This can tell us how much energy is available for work. You will not have to actually calculate free energy but but you should recognize it.
Additionally, there are several ways to manipulate this equation to account for different scenarios. Since this is an introductory look at the topic, you will not be held accountable for these variants. That being said, you should become familiar with the variables in the basic delta G equals delta H minus T times delta S equation. As mentioned, delta G is the change in free energy and tells us the energy available to do work. This is important because organisms will need to accomplish work in order to maintain their internal order.
Delta H is the change in enthalpy in the system. Basically, this is a calculation of heat loss. Enthalpy literally means heat to warm with.
Heat is an easy path to disorder and is a natural consequence of the energy conservation law. T is simply a temperature constant. And then delta S is the change in entropy, or disorder. When we take this into account, we can determine whether a reaction is spontaneous or non-spontaneous.
Let's start with negative delta G. If the sign is negative, that means that energy has to be negative. has not had to be invested and there is still energy available for work.
These reactions are spontaneous. It can also be said that these reactions are favorable and that energy is released so they are exergonic. On the other hand, a positive delta G is going to be non-spontaneous.
There is not excess energy here and in fact energy has to be invested for this reaction to proceed. These reactions are said to be unfavorable. and they are endergonic because they require energy to be put into them to make the reaction occur.
Now, spontaneity is not a measure of how fast a reaction will occur. It only tells us if it can occur without the input of external energy. Let's draw out a way to think about this.
So one way to think about delta G's is to kind of picture it like a mountain. So maybe this is Mount Everest. And we're going to set a point as zero. So I've got a mountain.
I've got a number line. This is zero. And if I got dropped off on the middle of this mountain, And I just decided to stay right here at zero. This is Mount Everest and I'm halfway up. I can't stay at zero.
And if delta G equals zero, then for a cell, this is not going to work. This is actually no work being done and the cell is dead. So you have to do something. And...
If you think about it, just as a way to remember how Delta G's work, if I got dropped off unprepared halfway up Mount Everest, what would be my natural inclination? Well, my natural inclination would be to get off this mountain. I didn't plan for this.
I didn't prepare for this. And my natural inclination is to go down. If this is a number line and I'm dropping below zero, that's negative.
So we definitely want to do away with this idea that negative is bad. Negative is a sign on my delta G. And in this case, these are my spontaneous reactions. These are favorable. And they're going to be those reactions that are Exergonic, energy out, and we use the term exergonic for most cell functions.
In chemistry you may hear reactions referred to as exothermic. Those are also energy out reactions, but in a cell we don't really want to be able to feel the heat coming off of the reaction. That wouldn't be healthy for the cell. And so in cellular terms we use exergonic. Now I do also have the option to try to make the peak and to go up Mount Everest here.
But that is not going to be my spontaneous reaction. That's going to be a positive delta G. And those reactions are... Categorized as non-spontaneous, they are going to be unfavorable. And if I'm going to take this trek up Mount Everest, it's going to be endergonic.
I'm going to have to put some energy in to achieve this goal. It's possible. but I'm going to have to exert energy to get that done. Now as a cell, a cell is going to need both of these types of reactions to work. And how a cell is going to accomplish this is by coupling these reactions.
So in metabolism, we see lots and lots of coupled reactions. And what that means practically is that we have to see two things. One is there has to be a shared intermediate. And a shared intermediate just means that the product of one reaction is going to be a reactant in the next.
And it doesn't matter which product. just one of them has to show up over there. My other rule for a coupled reaction is that the overall delta G, the overall delta G should be negative, should be energy out.
Again, negative doesn't mean bad. It just tells me whether this reaction is spontaneous or not. And so if we look at an example, if I had, and I'm just going to use some letter designations here in place of actual reactants and products.
If I had something A that I was going to break down into B plus C, and the delta G of this was maybe... Negative 8.6, so this would be favorable and spontaneous, and I'm going to couple this to another reaction, maybe where I combine B and D to form E, and the delta G on this is positive 8.4. Overall, together, this is going to be negative 0.2, and the size of this doesn't matter.
Most reactions in the cell actually run really close to zero. So again, the magnitude of this doesn't tell me anything about how likely it is to happen or how fast it is to happen. It just says this is spontaneous.
And most cells are going to run a lot of reactions really close to zero. They might have a few reactions that are really far away. And they can actually run a couple of reactions that are. at zero as long as the overall delta G for the cell is doing work. So our cells do need to be accomplishing work.
Again, remember that spontaneity is not a measure of how fast something will occur. Glucose has a delta G around negative 686 kilocals per joule. And we know that living organisms can harvest energy from glucose. But if I poured a pile of sugar on the bench top, it's going to stay there a long time and not just spontaneously break down at a very fast rate.
In order to have reactions occur in a biologically relevant time frame, we are going to need to speed things up a bit. That's where catalysts come into play. Catalysts work to speed up reactions.
Be very clear, they do not affect the energetics or spontaneity of a reaction. They only affect the rate. This is achieved by lowering the activation energy, which is energy required to reach transition states. What happens at transition states?
In anabolic reactions, bond formation occurs. While in catabolic reactions, bond breaking occurs. In a variety of reactions, a functional group such as phosphate may be transferred between molecules. Catalyst is a very broad term and includes anything that can speed up a reaction while it itself remains unchanged. In biology, our favorite catalysts are the protein enzymes.
However, enzymes aren't the only type of catalyst. Heat can also be a catalyst. And in Chapter 4, we'll see...
catalytic RNA in the ribosome. We'll also see that enzymes can work in different locations. The majority of enzymes are classified as endoenzymes and will be the main focus of this chapter.
These enzymes function inside the cell. Just keep in mind that later we will also explore exoenzymes. These are secreted by the cell and perform their function outside the cell.
This might be to break down an exceptionally large cell. nutrients such as starch, which we will see in lab later this semester. When the cell is undergoing growth, exoenzymes are also used to help build new cell walls.
Enzymes lower activation energy by binding and positioning substrates. This is often likened to a lock and key, however this description first coined by Fisher in the 1890s doesn't accurately describe the kinetics of the enzyme sub-enzymes. substrate interaction. Following research regarding enzyme reaction rates by Michaelis and Mitten in the early 1900s, Koshland arrived at the induced fit model in the 1950s.
This model describes the enzyme active site as able to fit several substrates, but only the correct substrate will induce bond interactions. These bond interactions should be weak and easily reversible. What are some examples of the Koshland model?
of weak bonds. The most common are hydrogen bonds. However, we may also see ionic bonds, which are easily disrupted by water, or van der Waals interactions.
Remember, the definition of a catalyst is that it is unchanged itself by the reaction, so we shouldn't see any permanent types of bonds here. Ultimately, catalysts can increase the rate of reaction by 10 to the 8th, up to 10 to the 20th of the normal reaction rate. It should also be noted that many enzymes do not work alone, and many require non-protein helpers to function at optimal capacity.
It is very common for enzymes to bind to metallic ions, such as iron, to work more efficiently. In fact, there are whole classes of membrane-bound enzymes that use iron that are called siderophores. And as we will see in the next presentation, there are non-protein coenzymes such as NADH that are essential for transferring energy in the form of electrons and bonds through the metabolic pathways. Let's finish up this presentation by graphing the impact of catalysts on a reaction. So let's talk about what the overall transition state and activation energy would look like for a reaction.
What do we mean when we say an enzyme lowers my activation activation energy and when we say an enzyme does not affect the energetics. So we're going to draw a graph here and what we're looking at is an energy hill. Most reactions can be described in relation to an energy hill.
So maybe this is from earlier, maybe this is A going to B plus C. Now, right at the tip, right here at the peak of the hill, this is my transition state. That's the moment when A breaks apart and I get B plus C. So right there at the tip, that's my transition state.
How tall that hill is... That's my energy of activation. That's that energy barrier that I have to overcome. And what I'll see is that for each side of this hill I have the delta G of the forward of this reaction and I have what it would take to back this reaction up to make it run in reverse. That's the delta G of the reverse of this reaction.
Now, there is an overall delta G to this. So I have the individual forward and reverse. But sitting here underneath the hill, I have the delta G of the total reaction.
And that's going to be the delta G of that forward minus the delta G of the reverse. So For this hill, what is my overall delta G? Well, I don't really need any numbers to be able to see what the overall delta G is here, because this forward is a little short arrow, and this reverse is a big tall arrow. So this is going to be a little number minus a big number. So overall, my delta G should, for this reaction, be negative.
If it makes you feel better to put in some numbers, let's say I have this is 100, this is much bigger, it's 200, and overall negative 100. What the enzyme is going to do, what a catalyst is going to do, is going to lower that energy of activation. So my enzyme is going to serve as a catalyst and it's going to lower this energy of activation. This does not change the energetics of this reaction.
Sometimes this really confuses people because we're lowering this hill and we have this forward and reverse, but trust me the delta G of the reaction is not affected. It just makes this hill shorter. It's the studs of Mount Everest. Maybe I have Mount Scott down by Lawton.
And so that hill is not nearly as tall. It's going to be a lot easier to get over. But the energy doesn't change because however much I lower this forward, I also equally lowered the reverse. And so my overall delta G of this reaction, the energy is unchanged.