Winter_2021_Bis2A_Facciotti_Reading_12

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Learning Goals Associated with Winter_2021_Bis2A_Facciotti_Reading_12

Define a redox reaction and identify common biological redox reactions.
Given a redox reaction, identify the reducing agent, oxidizing agent, molecule that becomes oxidized, and the reduced species. Identify which species the electron(s) “starts” in, and to which species it “goes.”
Write a composite chemical equation when given two redox half-reactions.
Calculate the ΔE_₀’ for a given redox reaction using the equation ΔE_₀’ = E_₀’(oxidant) - E_₀’(reductant)
Predict whether a directional transfer of electrons between two chemical species is endergonic or exergonic by applying the concept of redox potential to provided data.
Qualitatively relate the difference in redox potentials with a corresponding delta of Gibbs enthalpy.
Define each variable and its role in the equation: ΔG₀’ = -nFΔE_₀’.
Convert between ΔG₀’ and ΔE_₀’ for a given redox reaction using the equation ΔG0’ = -nFΔE_₀’
Tell an energy story for a redox reaction that utilizes the electron carrier NAD⁺/NADH as the second substrate in the simple, generic reaction scheme: AH + NAD⁺ -> A⁺ + NADH.
Identify NAD⁺ from its molecular structure and identify the functional group involved in its function as an oxidizing or reducing agent.

Metabolism in General Biology

Cellular metabolism represents roughly 1/3 of the General Biology curriculum. While this may seem like a lot, we cover very little of what a classic course in metabolism covers, and a minuscule fraction of the metabolism that occurs on the planet. What we cover, however, is the very important foundational knowledge. You will learn about common chemical reactions that are associated with the transformation of life's molecular building blocks and about different core modes of energy transfer that you will encounter often in biology. The energy story and the design challenge rubrics introduced earlier will become increasingly important in these next few modules and beyond.

What have we learned? How will it relate to metabolism?

We have focused on the identification and chemical properties of common biological functional groups. As we dive into metabolism, this will help you be familiar with and sometimes even predict the chemical nature/reactivity of compounds you have never seen before.
We have practiced recognizing and classifying molecules into four major functional groups. This will help you as we discuss how to build and break down these molecules.
We have learned some basic thermodynamics. This gives us a common set of concepts with which to discuss whether a biochemical reaction or process is likely to occur, and if so in which direction and how fast. This will be critical as we consider some key reactions that take place in metabolism.
We have learned and practiced the energy story rubric. This will allow us to study new biochemical reactions and to discuss them with a common, consistent language and approach which also reinforces the lessons we learned about thermodynamics.

An overview of this section

We will introduce an important concept called reduction potential and you will be given the opportunity to use a redox tower. There is also a discussion on redox chemistry in your discussion manual. Use both resources.
We will introduce two major players in metabolism, ATP and NADH. We expect you to recognize their structures if shown on an exam.
We will cover the metabolic pathway glycolysis. Keep in mind that we want you to look at any reaction and tell us an energy story of that reaction. You should not try to memorize these pathways (though it will help to remember some big picture things - we will stress these). Often we will give you the pathway as a figure on the exams. Glycolysis produces 2 ATP via a process called substrate level phosphorylation, 2 NADH and 2 pyruvate compounds.
We will use the reactions of the TCA cycle to create multiple examples of energy stories. The TCA cycle will also produce more ATP, NADH and oxidize glucose into CO₂.
We will look at an alternative pathway to that of the TCA cycle, fermentation. Here, for the first time, we will see NADH used as a reactant in a metabolic reaction.
We will follow NADH to the end of its journey, as it donates its electrons to the electron transport chain (ETC). In this module, you will need to use a redox tower. The ETC produces a proton gradient. No ATP is directly generated in this process. However, the proton gradient is then used by the cell to run an enzyme called ATP synthase, which catalyzes the reaction ADP + Pi --> ATP. This method of ATP production (called oxidative phosphorylation) results in more ATP being produced than through substrate level phosphorylation.
And finally, we will go through the process of photosynthesis.

Reduction-Oxidation Reactions

In this class, most of the reduction/oxidation reactions (redox) that we discuss occur in metabolic pathways (connected sets of metabolic reactions). Here the cell breaks down the compounds it consumes into smaller parts and then reassembles them into larger macromolecules. Redox reactions also play critical roles in energy transfer, either from the environment or within the cell, in all known forms of life. For these reasons, it is important to develop at least an intuitive understanding and appreciation for redox reactions in biology.

Most students of biology will also study reduction and oxidation reactions in their chemistry courses; these kinds of reactions are important well beyond biology. Regardless of the order in which students are introduced to this concept (chemistry first or biology first), most will find the topic presented in very different ways in chemistry and biology. That can be confusing.

Chemists often introduce the concepts of oxidation and reduction from the technically more correct and inclusive standpoint of oxidation states. See this link for more information: <https://chem.libretexts.org/Bookshel...ation_Numbers)>. Fortunately, there’s no need to go into the details here (most of you will see that in chemistry at some point), just follow the argument for now. It might make things less confusing in both the long and short run. Anyhow, chemists will often ask students to apply a set of rules (see link above) to determine the oxidation states of individual atoms in a reaction. The chemistry formalism defines oxidation as an increase in oxidation state and reduction as a decrease in oxidation state.

All of this, of course, holds true in biology. However, biologists don’t typically think of redox reactions in this way. Why? We suspect it’s because most of the redox reactions encountered in biology involve a change in oxidation state that comes about because electrons are transferred between molecules. So, biologists typically define reduction as a gain of electrons and oxidation as a loss of electrons. The biological concept of redox is entirely consistent with the concept chemists use but it doesn’t account for redox reactions that can happen without the transfer of electrons. The biologist’s definition is therefore not as general as the chemist’s, but it works for most cases encountered in biology.

This is a biology reading for a biology class. We, therefore, approach redox from the “gain/loss of electrons” conceptualization that is commonly taught in biology classes. In our opinion, it’s easier to use (no long list of rules to memorize and apply), more intuitive, and works for almost all cases we care about in undergraduate biology. So, if you had chemistry already and this topic seems a little different in biology, remember that at its core it’s the same thing you learned about before. Biologists just adapted what you learned in chemistry to make more intuitive sense in biology. If you haven’t learned about redox yet don’t worry. If you can understand what we’re trying to do here when you cover this concept in chemistry class you’ll be a few steps ahead. You’ll just need to generalize your thinking a bit instead of seeing the topic for the first time.

Let's start with some generic reactions

Transferring electrons between two compounds results in one of these compounds loosing an electron and one compound gaining an electron. For example, look at the figure below. If we use the energy story rubric to look at the overall reaction, we can compare the before and after characteristics of the reactants and products. What happens to the matter (stuff) before and after the reaction? Compound A starts as neutral and becomes positively charged. Compound B starts as neutral and becomes negatively charged. Because electrons are negatively charged, we can explain this reaction with the movement of an electron from Compound A to B. That is consistent with the changes in charge. Compound A loses an electron (becoming positively charged), and we say that A has become oxidized. For biologists, oxidation is associated with the loss of electron(s). B gains the electron (becoming negatively charged), and we say that B has become reduced. Reduction is associated with the gain of electrons. We also know, since a reaction occurred (something happened), that energy must have been transferred and/or reorganized in this process and we'll consider this shortly.

Marys redox figure.png

Figure 1. Generic redox reaction with half-reactions

Attribution: Mary O. Aina

Knowledge Check Quiz

To reiterate: When an electron(s) is lost, or a molecule is oxidized, the electron(s) must then pass to another molecule. We say that the molecule gaining the electron becomes reduced. Together these paired electron gain-loss reactions are known as an oxidation-reduction reaction (also called a redox reaction).

This idea of paired half-reactions is critical to the biological concept of redox. Electrons don’t drop out of the universe for “free” to reduce a molecule or jump off a molecule into the ether. Donated electrons MUST come from a donor molecule and be transferred to some other acceptor molecule. For example in the figure above the electron the reduces molecule B in half-reaction 2 must come from a donor - it just doesn't appear from nowhere! Likewise, the electron that leaves A in half-reaction 1 above just "land" on another molecule - it doesn't just disappear from the universe.

Therefore, oxidation and reduction reactions must ALWAYS be paired. We’ll examine this idea in more detail below when we discuss the idea of “half-reactions”.

A tip to help you remember: The mnemonic LEO says GER (Lose Electrons = Oxidation and Gain Electrons = Reduction) can help you remember the biological definitions of oxidation and reduction.

Figure 2. A figure for the mnemonic "LEO the lion says GER." LEO: Loss of Electrons = Oxidation. GER: Gain of Electrons = Reduction

Attribution: Kamali Sripathi

• The vocabulary of redox can be confusing: Students studying redox chemistry can often become confused by the vocabulary used to describe the reactions. Terms like oxidation/oxidant and reduction/reductant look and sound very similar but mean distinctly different things. An electron donor is also sometimes called a reductant because it is the compound that causes the reduction (gain of electrons) of another compound (the oxidant). In other words, the reductant is donating it’s electrons to the oxidant which is gaining those electrons. Conversely, the electron acceptor is called the oxidant because it is the compound that is causing the oxidation (loss of electrons) of the other compound. Again, this simply means the oxidant is gaining electrons from the reductant who is donating those electrons. Confused yet?

Yet another way to think about definitions is to remember that describing a compound as reduced/oxidized is describing the state that the compound itself is in, whereas labeling a compound as a reductant/oxidant describes how the compound can act, to either reduce or oxidize another compound. Keep in mind that the term reductant is also synonymous with reducing agent and oxidant is also synonymous with oxidizing agent. The chemists who developed this vocabulary need to be brought up on charges of "willful thickheadedness" at science trial and then be forced to explain to the rest of us why they needed to be so deliberately obtuse.

The confusing language of redox: quick summary

1. A compound can be described as “reduced” - term used to describe the compound's state

2. A compound can be a “reductant” - term used to describe a compound's capability (it can reduce something else). The synonymous term "reducing agent" can be used to describe the same capability (the term "agent" refers to the thing that can "do something" - in this case reduce another molecule).

3. A compound can be an “oxidant” - term used to describe a compound's capability (it can oxidize something else). The synonymous term "oxidizing agent" can be used to describe the same capability (the term "agent" refers to the thing that can "do something" - in this case oxidize another molecule).

4. A compound can “become reduced” or "become oxidized"- term used to describe the transition to a new state

Since all of these terms are used in biology, in General Biology we expect you to become familiar with this terminology. Try to learn it and use it as soon as possible - we will use the terms frequently and will not have the time to define terms each time.

Knowledge Check Quiz

Knowledge Check Quiz

The Half Reaction

Here we introduce the concept of the half reaction. We can think each half reaction as a description of what happens to one of the two molecules (i.e. the donor and the acceptor) involved in a "full" redox reaction. A "full" redox reaction requires two half reactions. We illustrate this below. In the example below, half reaction #1 depicts the molecule AH becoming losing two electrons and a proton and in the process becoming A⁺. This reaction depicts the oxidation of AH. Half reaction #2 depicts the molecule B⁺ gaining two electrons and a proton to become BH. This reaction depicts the reduction of B⁺. Each of these two half reactions is conceptual and can't happen on their own. The electrons lost in half reaction #1 MUST go somewhere, they can't just disappear. Likewise, the electrons gained in half reaction #2 must come from something. They too just can't appear out of nowhere.

One can imagine that there might be different molecules that can serve as potential acceptors (the place for the electrons to go) for the electrons lost in half reaction #1. Likewise, there might be many potential reduced molecules that can serve as the electron donors (the source of electrons) for half reaction #2. In the example below, we show what happens (the reaction) when molecule AH is the donor of electrons for molecule B⁺. When we put the donor and acceptor half reactions together, we get a "full" redox reaction that can actually happen. In the figure below we call that reaction "Reaction #1". When this happens we call the two half reactions coupled.

Using this idea, we can theoretically couple and think about any two half reactions, one half reaction serving as the electron donor for the other half reaction that accepts the donated electrons. For instance, using the example above, we could consider coupling the reduction of B⁺ that happens in half reaction 2 with another half reaction describing the oxidation of the molecule NADH. In that case, the NADH would be the electron donor for B⁺. Likewise you could couple the oxidation of AH that happens in half reaction #1 with a half reaction describing the the reduction of hypothetical molecule Z⁺. You can mix-and-match half reactions together as you please provided one half is describing the oxidation of a compound (it's donating electrons) and the reduction of another compound (it's accepting the donated electrons).

A note on how we write full reactions versus half reactions: In the example above, when we write Reaction #1 as an equation, the 2 electrons and the H⁺ that are explicitly described in the underlying half reactions, are not explicitly included in the text of the full reaction. In the reaction above you must infer that an exchange of electrons happens. This can be observed by trying to balance charges between each reactant and it's corresponding product. Reactant AH becomes product A⁺. In this case, you can infer that some movement of electrons must have taken place. To balance the charges on this compound (make the sum of charges on each side of the equation equal) you need to add 2 electrons to the right side of the equation, one to account for the "+" charge on A⁺ and a second to go with the H⁺ that was also lost. The other reactant B⁺ is converted to BH. It must therefore gain 2 electrons to balance charges, one for B⁺ and a second for the additional H⁺ that was added. Together this information leads you to conclude that the most likely thing to have happened is that two electrons were exchanged between AH and B⁺.
This will also be the case for most redox reactions in biology. Fortunately, in most cases, either the context of the reaction, the presence of chemical groups often engaged in redox (e.g. metal ions), or the presence of commonly used electron carriers (e.g. NAD⁺/NADH, FAD⁺/FADH₂, ferredoxin, etc.) will alert you that the reaction is of class "redox". You will be expected to learn to recognize some of these common molecules.

Reduction Potential

By convention, we quantitatively characterize redox reactions using an measure called reduction potentials. The reduction potential attempts to quantitatively describe the “ability” of a compound or molecule to gain or lose electrons. The specific value of the reduction potential is determined experimentally, but for the purpose of this course we assume that the reader will accept that the values in provided tables are reasonably correct. We can anthropomorphize the reduction potential by saying that it is related to the strength with which a compound can “attract” or “pull” or “capture” electrons. Not surprisingly this is is related to but not identical to electronegativity.

What is this intrinsic property to attract electrons?

Different compounds, based on their structure and atomic composition have intrinsic and distinct attractions for electrons. This quality leads each molecule to have its own standard reduction potential or E₀’. The reduction potential is a relative quantity (relative to some “standard” reaction). If a test compound has a stronger "attraction" to electrons than the standard (if the two competed, the test compound would "take" electrons from the standard compound), we say that the test compound has a positive reduction potential. The magnitude of the difference in E₀’ between any two compounds (including the standard) is proportional to how much more or less the compounds "want" electrons. The relative strength of the reduction potential is measured and reported in units of Volts (V) (sometimes written as electron volts or eV) or milliVolts (mV). The reference compound in most redox towers is H₂.

Possible NB Discussion Point

Rephrase for yourself: How do you describe or think about the difference between the concept of electronegativity and red/ox potential?

Redox student misconception alert: The standard redox potential for a compound reports how strongly a substance wants to hold onto an electron in comparison to hydrogen. Since both redox potential and electronegativity are both discussed as measurements for how strongly something "wants" an electron, they are sometimes conflated or confused for one another. However, they are not. While the electronegativity of atoms in a molecule may influence its redox potential, it is not the only factor that does. You don't need to worry about how this works. For now, try to keep them as different and distinct ideas in your mind. The physical relationship between these two concepts is well beyond the scope of this general biology class.

The Redox Tower

All kinds of compounds can take part in redox reactions. Scientists have developed a graphical tool, the redox tower, to tabulate redox half reactions based on their E₀^' values. This tool can help predict the direction of electron flow between potential electron donors and acceptors and how much free energy change might be expected from a specific reaction. By convention, all half reactions in the table are written in the direction of reduction for each compound listed.

In the biology context, the electron tower usually ranks a variety of common compounds (their half reactions) from most negative E₀^'(compounds that readily get rid of electrons), to the most positive E₀^' (compounds most likely to accept electrons). The tower below lists the number of electrons that are transferred in each reaction. For example, the reduction of NAD⁺ to NADH involves two electrons, written in the table as 2e^-.

oxidized form	reduced form	n (electrons)	Eo´ (volts)
PS1* (ox)	PS1* (red)	-	-1.20
Acetate + CO₂	pyruvate	2	-0.7
ferredoxin (ox) version 1	ferredoxin (red) version 1	1	-0.7
succinate + CO₂+ 2H⁺	a-ketoglutarate + H₂O	2	-0.67
PSII* (ox)	PSII* (red)	-	-0.67
P840*(ox)	PS840*(red)	-	-0.67
acetate	acetaldehyde	2	-0.6
glycerate-3-P	glyceraldehyde-3-P + H₂O	2	-0.55
O₂	O₂^-	1	-0.45
ferredoxin (ox) version 2	ferredoxin (red) version 2	1	-0.43
CO₂	glucose	24	-0.43
CO₂	formate	2	-0.42
2H+	H₂	2	-0.42 (at [H⁺] = 10^-7; pH=7) Note: at [H⁺] = 1; pH=0 the Eo' for hydrogen is ZERO. You will see this in chemistry class.
α-ketoglutarate + CO₂+ 2H⁺	isocitrate	2	-0.38
acetoacetate	b-hydroxybutyrate	2	-0.35
Cystine	cysteine	2	-0.34
Pyruvate + CO₂	malate	2	-0.33
NAD⁺+ 2H⁺	NADH + H⁺	2	-0.32
NADP⁺+ 2H⁺	NADPH + H⁺	2	-0.32
Complex I FMN (enzyme bound)	FMNH₂	2	-0.3
Lipoic acid, (ox)	Lipoic acid, (red)	2	-0.29
1,3 bisphosphoglycerate + 2H⁺	glyceraldehyde-3-P + P_i	2	-0.29
Glutathione, (ox)	Glutathione, (red)	2	-0.23
FAD⁺ (free) + 2H⁺	FADH₂	2	-0.22
Acetaldehyde + 2H⁺	ethanol	2	-0.2
Pyruvate + 2H⁺	lactate	2	-0.19
Oxalacetate + 2H⁺	malate	2	-0.17
α-ketoglutarate + NH₄+	glutamate	2	-0.14
FAD⁺ + 2H⁺ (bound)	FADH₂(bound)	2	0.003-0.09
Methylene blue, (ox)	Methylene blue, (red)	2	0.01
Fumarate + 2H⁺	succinate	2	0.03
CoQ (Ubiquinone - UQ + H⁺)	UQH.	1	0.031
UQ + 2H⁺	UQH₂	2	0.06
Dehydroascorbic acid	ascorbic acid	2	0.06
Plastoquinone; (ox)	Plastoquinone; (red)	-	0.08
Ubiquinone; (ox)	Ubiquinone; (red)	2	0.1
Complex III Cytochrome b₂; Fe³⁺	Cytochrome b₂; Fe²⁺	1	0.12
Fe³⁺(pH = 7)	Fe²⁺(pH = 7)	1	0.20
Complex III Cytochrome c₁; Fe³⁺	Cytochrome c₁; Fe²⁺	1	0.22
Cytochrome c; Fe³⁺	Cytochrome c; Fe²⁺	1	0.25
Complex IV Cytochrome a; Fe³⁺	Cytochrome a; Fe²⁺	1	0.29
1/2 O₂ + H₂O	H₂O₂	2	0.3
P840^GS(ox)	PS840^GS(red)	-	0.33
Complex IV Cytochrome a₃; Fe³⁺	Cytochrome a₃; Fe²⁺	1	0.35
Ferricyanide	ferrocyanide	2	0.36
Cytochrome f; Fe³⁺	Cytochrome f; Fe²⁺	1	0.37
PSI^GS(ox)	PSI^GS(red)	.	0.37
Nitrate	nitrite	1	0.42
Fe³⁺(pH = 2)	Fe²⁺(pH = 2)	1	0.77
1/2 O₂ + 2H⁺	H₂O	2	0.816
PSII^GS(ox)	PSII^GS(red)	-	1.10
* Excited State, after absorbing a photon of light GS Ground State, state prior to absorbing a photon of light PS1: Oxygenic photosystem I P840: Bacterial reaction center containing bacteriochlorophyll (anoxygenic) PSII: Oxygenic photosystem II

Video on electron tower

For a short video on how to use the electron tower in redox problems click here or below. This video was made by Dr. Easlon for Bis2A students. (This is quite informative.)

**What is the relationship between ΔE₀^' and ΔG?**

How do we know if any given redox reaction (the specific combination of two half reactions) is energetically spontaneous or not (exergonic or endergonic)? Moreover, regardless of the direction of spontaneity, how can we determine what the quantitative change in free energy is for a specific redox reaction? The answer lies in the difference in the reduction potentials of the two compounds. The difference in the reduction potential for the reaction (∆E₀'), can be calculated by taking the difference between the E₀^' for the oxidant (the compound getting the electrons and causing the oxidation of the other compound) and the reductant (the compound losing the electrons). In our generic example below, AH is the reductant and B⁺ is the oxidant. Electrons are moving from AH to B⁺. Using the E₀^' of -0.32 for the reductant and +0.82 for the oxidant the total change in E₀^' or ∆E₀' is 1.14 eV.

Figure 4. Generic red/ox reaction with half reactions written with reduction potential (E₀^') of the two half reactions indicated.

∆E₀' between oxidant and reductant can tell us about the spontaneity of a proposed electron transfer. Intuitively, if electrons are proposed to move from a compound that "wants" electrons less to a compound that "wants" electrons more (i.e. a move from a compound with a lower E₀^' to a compound with a higher E₀^' , the reaction will be energetically spontaneous). If the electrons are proposed to move from a compound that "wants" electrons more to a compound that "wants" electrons less (i.e. a move from a compound with a higher E₀^' to a compound with a lower E₀^', the reaction will be energetically non-spontaneous). Because of the way biological/biochemical redox tables are ordered (small E₀^' on top and larger E₀^' on the bottom) transfers of electrons from donors higher on the table to acceptors lower on the table will be spontaneous.

It is also possible to quantify the amount of free energy change associated with a specific redox reaction. The relationship is given by the Nernst equation:

Figure 5. The Nernst equation relates free energy of a redox reaction to the difference in reduction potential between the reduced products of the reaction and oxidized reactant.
Attribution: Marc T. Facciotti

Where:

n is the number of moles of electrons transferred
F is the Faraday constant of 96.485 kJ/V. Sometimes it is given in units of kcal/V which is 23.062 kcal/V, which is the amount of energy (in kJ or kcal) released when one mole of electrons passes through a potential drop of 1 volt

Note that the signs of ∆E and ∆G are opposite one another. When ∆E is positive, ∆G will be negative. When ∆E is negative, ∆G will be positive.

Introduction to Mobile Energy Carriers

Section Summary

Energy moves and transfers within the cell in a variety of ways. One critical mechanism that nature has developed is the use of recyclable molecular energy carriers. While there are several major recyclable energy carriers, they all share some common functional features:

Properties of Key Cellular Molecular Energy Carriers

We think of the energy carriers as existing in "pools" of available carriers. One could, by analogy, consider these mobile energy carriers analogous to the delivery vehicles of parcel carriers—the company has a certain "pool" of available vehicles at any one time to pickup and make deliveries.
Each individual carrier in the pool can exist in one of multiple distinct states: it is either carrying a "load" of energy, a fractional load, or is "empty". The molecule can interconvert between "loaded" and empty and thus can be recycled. Again by analogy, the delivery vehicles can be either carrying packages or be empty and switch between these states.
The balance or ratio in the pool between "loaded" and "unloaded" carriers is important for cellular function, is regulated by the cell, and can often tell us something about the state of a cell. Likewise, a parcel carrier service keeps close tabs on how full or empty their delivery vehicles are—if they are too full, there may be insufficient "empty" trucks to pick up new packages; if they are too empty, business must not be going well or they shut it down. There is an appropriate balance for different situations.

In this course, we will examine two major types of molecular recyclable energy carriers: (1) the adenine nucleotides: nicotinamide adenine dinucleotide (NAD⁺), a close relative, nicotinamide adenine dinucleotide phosphate (NADP⁺), and flavin adenine dinucleotide (FAD₂⁺) and (2) nucleotide mono-, di-, and triphosphates, with particular attention paid to adenosine triphosphate (ATP). These two types of molecules participate in a variety of energy transfer reactions. We primarily associate adenine nucleotides with red/ox chemistry, and nucleotide triphosphates with energy transfers linked to the hydrolysis or condensation of inorganic phosphates.

Red/ox chemistry and electron carriers

The oxidation of, or removal of an electron from, a molecule (whether accompanied with the removal of an accompanying proton or not) results in a change of free energy for that molecule—matter, internal energy, and entropy have all changed. Likewise, the reduction of a molecule also changes its free energy. The magnitude of change in free energy and its direction (positive or negative) for a red/ox reaction dictates the spontaneity of the reaction and how much energy it transfers. In biological systems, where a great deal of energy transfer happens via red/ox reactions, it is important to understand how these reactions are mediated and to consider ideas or hypotheses for why these reactions are mediated most times by a small family of electron carriers.

Note: BURNING BEAR

Relate the burning of (the full oxidation of the sugar in) a gummy bear with the last paragraph above. What does that demonstration have to do with our upcoming discussion on red/ox carriers? There is some mention above already—can you find it?

Note: DESIGN CHALLENGE

The problem alluded to in the previous discussion question is a great place to bring in the design challenge rubric. If you recall, the first step of the rubric asks that you define a problem or question. Here, let's imagine that there is a problem to define for which the mobile electron carriers below helped Nature solve.

***Remember, evolution DOES NOT forward-engineer solutions to problems, but in retrospect, we can use our imagination and logic to infer that what we see preserved by natural selection provided a selective advantage, because the natural innovation "solved" a problem that limited success.***

Design challenge for red/ox carriers

What was a problem(s) that the evolution of mobile electron carriers helped solve?
The next step of the design challenge asks you to identify criteria for successful solutions. What are criteria for success in the problem you've identified?
Step 3 in the design challenge asks you to identify solutions. Well, here Nature has identified some for us—we consider three in the reading below. It looks like Nature is happy to have multiple solutions to the problem.
The penultimate step of the design challenge rubric asks you to test the proposed solutions against the criteria for success. This should make you think/discuss why there are multiple different electron carriers. Are there different criteria for success? Are they each solving slightly different problems? What do you think? Be on the lookout as we go through metabolism for clues.

NAD⁺/H and FADH/H₂

In living systems, a small class of compounds function as electron shuttles: they bind and carry electrons between compounds in different metabolic pathways. The principal electron carriers we will consider derive from the B vitamin group and nucleotides. These compounds can both become reduced (that is, they accept electrons) or oxidized (they lose electrons) depending on the reduction potential of a potential electron donor or acceptor that they might transfer electrons to and from. Nicotinamide adenine dinucleotide (NAD⁺) (we show the structure below) derives from vitamin B₃, niacin. NAD⁺ is the oxidized form of the molecule; NADH is the reduced form of the molecule after it has accepted two electrons and a proton (which together are the equivalent of a hydrogen atom with an extra electron).

We are expecting you to memorize the two forms of NAD⁺/NADH, know which form is oxidized and which is reduced, and be able to recognize either form on the spot in a chemical reaction.

NAD⁺ can accept electrons from an organic molecule according to the general equation:

Here is some vocabulary review: when electrons are added to a compound, we say the compound has been reduced. A compound that reduces (donates electrons to) another is called a reducing agent. In the above equation, RH is a reducing agent, and NAD⁺ becomes reduced to NADH. When electrons leave a compound, it becomes oxidized. A compound that oxidizes another is called an oxidizing agent. In the above equation, NAD⁺ is an oxidizing agent, and RH is oxidized to R. Put another way, the reducing agent gets oxidized and the oxidizing agent gets reduced.

You need to get this down! We will (a) test specifically on your ability to do so (as "easy" questions), and (b) we will use the terms with the expectation that you know what they mean and can relate them to biochemical reactions correctly (in class and on tests).

You will also encounter a second variation of NAD⁺, NADP⁺. It is structurally very similar to NAD⁺, but it contains an extra phosphate group and plays an important role in anabolic reactions, such as photosynthesis. Another nucleotide-based electron carrier that you will also encounter in this course and beyond, flavin adenine dinucleotide (FAD⁺), derives from vitamin B₂, also called riboflavin. Its reduced form is FADH₂. Learn to recognize these molecules as electron carriers.

Figure 1. The oxidized form of the electron carrier (NAD⁺) is shown on the left, and the reduced form (NADH) is shown on the right. The nitrogenous base in NADH has one more hydrogen ion and two more electrons than in NAD⁺.

The cell uses NAD+ to "pull" electrons off of compounds and to "carry" them to other locations within the cell; thus we call it an electron carrier. Many metabolic processes we will discuss in this class involve NAD(P)⁺/H compounds. For example, in its oxidized form, NAD⁺ is used as a reactant in glycolysis and the TCA cycle, whereas in its reduced form (NADH), it is a reactant in fermentation reactions and electron transport chains (ETC). We will discuss each of these processes in later modules.

Energy story for a red/ox reaction

***As a rule of thumb, when we see NAD+/H as a reactant or product, we know we are looking at a red/ox reaction.***

When NADH is a product and NAD⁺ is a reactant, we know that NAD⁺ has become reduced (forming NADH); therefore, the other reactant must have been the electron donor and become oxidized. The reverse is also true. If NADH has become NAD⁺, then the other reactant must have gained the electron from NADH and become reduced.

Figure 2. This reaction shows the conversion of pyruvate to lactic acid coupled with the conversion of NADH to NAD⁺. Source: https://en.wikibooks.org/wiki/Structural_Biochemistry/Enzyme/sequential_reactions

In the figure above, we see pyruvate becoming lactic acid, coupled with the conversion of NADH into NAD⁺. LDH catalyses this reaction. Using our "rule of thumb" above, we categorize this reaction as a red/ox reaction. NADH is the reduced form of the electron carrier, and NADH is converted into NAD⁺. This half of the reaction results in the oxidation of the electron carrier. Pyruvate is converted into lactic acid in this reaction. Both sugars are negatively charged, so it would be difficult to see which compound is more reduced using the charges of the compounds. However, we know that pyruvate has become reduced to form lactic acid, because this conversion is coupled to the oxidation of NADH into NAD⁺. But how can we tell that lactic acid is more reduced than pyruvate? The answer is to look at the carbon-hydrogen bonds in both compounds. As electrons transfer, they are often accompanied by a hydrogen atom. Pyruvate has a total of three C-H bonds, while lactic acid has four C-H bonds. When we compare these two compounds in the before and after states, we see that lactic acid has one more C-H bond; therefore, lactic acid is more reduced than pyruvate. This holds true for multiple compounds. For example, in the figure below, you should be able to rank the compounds from most to least reduced using the C-H bonds as your guide.

Figure 3. Above are a series of compounds than can be ranked or reorganized from most to least reduced. Compare the number of C-H bonds in each compound. Carbon dioxide has no C-H bonds and is the most oxidized form of carbon we will discuss in this class. Answer: the most reduced is methane (compound 3), then methanol (4), formaldehyde (1), carboxylic acid (2), and finally carbon dioxide (5).

Figure 4. This reaction shows the conversion of G3P, NAD⁺, and P_iinto NADH and 1,3-BPG. This reaction is catalyzed by glyceraldehyde-3-phosphate dehydrogenase.

Energy story for the reaction catalyzed by glyceraldehyde-3-phosphate dehydrogenase:

Lets make an energy story for the reaction above.

First, let us characterize the reactants and products. The reactants are glyceraldehyde-3-phosphate (a carbon compound), P_i (inorganic phosphate), and NAD⁺. These three reactants enter a chemical reaction to produce two products, NADH and 1,3-bisphosphoglycerate. If you look closely, you can see that the 1,3-BPG contains two phosphates. This is important since a chemical reaction should lose no mass between its beginning and its end. There are two phosphates in the reactants, so there must be two phosphates in the products (conservation of mass!). You can double check the book keeping of mass for all other atoms. It should also tabulate correctly. An enzyme called glyceraldehyde-3-phosphate dehydrogenasethat catalyzes this reaction. The standard free energy change of this reaction is ~6.3 kJ/mol, so under standard conditions, we can say that the free energy of the products is higher than that of the reactants and that this reaction is not spontaneous under standard conditions.

What can we say about this reaction when it is catalyzed by glyceraldehyde-3-phosphate dehydrogenase?

This is a red/ox reaction. We know that because we have produced a reduced electron carrier (NADH) as a product and NAD⁺ is a reactant. Where did the electron come from to make NADH? The electron must have come from the other reactant (the carbon compound).

BONUS SECTION

Alternative View of Some Common Confusing Issues in Basic Redox Chemistry for Biology

This reading tries to break down some of the more challenging topics that students run into when studying redox chemistry in General Biology. This reading is not a substitute for your main reading but rather a complement to it that revisits some of the same topics through a different lens.

Finding ΔE

Students often struggle with finding the ∆E for a given redox reaction. One of the main barriers to developing this skill seems to be associated with developing a picture of the redox reaction itself. From the context of most biological redox reactions it is useful to imagine/picture a redox reaction as a simple exchange of electrons between two molecules, an electron donor and an electron acceptor that accepts electrons from the donor.

An analogy with kiwi fruit: To help build this mental picture we offer an analogy. Two people are standing next to one another. At the start, one person is holding a kiwi fruit in their hand and the second person's hands are empty. In this reaction, person 1 gives the kiwi to person 2. At the end of the reaction person 2 is holding a kiwi and person 1 is not. We can write this exchange as we might a chemical reaction:

person 1(kiwi) + person 2() <-> person 1() + person 2(kiwi).

start/initial state <-> final/end state

If we read this "reaction" from left to right, person 1 is a kiwi donor and person 2 is a kiwi acceptor. We can extend this analogy a little by proposing that person 1 and person 2 have different desire and ability to grab and hold kiwi fruit - we'll call that kiwi-potential. We can then propose to set up a situation where person 1 and person 2 compete for a kiwi. Let's propose that person 2 has a higher "kiwi-potential" than person 1 - that is, person 2 has a stronger desire and ability to grab and hold wiki than person 1.

If we set up a competition where person 1 starts with the kiwi and person 2 competes for it, we should expect that after some time the kiwi will be exchanged to person 2 and stay there most often. At the end of the reaction the kiwi will be with person 2. Due to the difference in "kiwi-potential" between person 1 and person 2, we can say that the spontaneous direction of kiwi flow is from person 1 to person 2. If we ever observed the kiwi flow from person 2 to person 1 we could probably conclude that person 1 required some extra help/energy to make that happen - flow from person 2 to person 1 would be non-spontaneous.

Let's call the "kiwi-potential", Kp. In our analogy, Kp_person₁ < Kp_person₂. We can calculate ∆Kp, the difference in Kp between the two people, and that will tell us something about how likely we can expect to see kiwi exchange hands between these two people. The bigger the difference in Kp the more likely the kiwi will move from the person who has a lower Kp to the person who has the higher Kp.

By definition, to calculate ∆Kp we obtain the solution to ∆Kp = Kp_final/end - Kp_{initial/start}. Since the kiwi is with person 2 at the end of the reaction and it starts with person 1 at the beginning of the reaction we would calculate ∆Kp = Kp_person₂ - Kp_person₁.

Doing it with electrons instead of kiwi fruit: To find ∆E for a redox reaction we can translate this analogy to the molecular space. Instead of people we have two molecules. Instead of a kiwi, we have electrons. Different molecules have different inherent abilities to grab and hold electrons and this can be measured by the value E. If two molecules exchange one or more electrons we can imagine that electrons will flow spontaneously from a molecule with lower E₀ to one with a higher E₀. We can write a familiar reaction with those substitutions.

molecule 1(electron) + molecule 2() <-> person 1() + molecule 2(electron).

start/initial state <-> final/end state

To find ΔE₀, you solve for ΔE₀ = E_0-final/end - E_{0-initial/start}. Alternatively, you can think of it as ΔE₀ = E_0-acceptor - E_0-donor.

When evaluating a redox reaction for ∆E you therefore need to:

First, find which of the reactants is the electron donor. The donor can also be associated with the initial state because it is the molecule that initially (before the start of the reaction) has the electron(s) to donate. This will always be one of the reactants and will be the molecule that gets oxidized (i.e. the molecule that loses electrons).
Second, find which of the reactants is the electron acceptor. This will also always be a reactant and will be the molecule that becomes reduced by the reaction (i.e. gains electrons). This molecule can be also associated with the final state since, in its reduced form, it is the molecule that has the electrons at the end of the reaction.
Third, calculate ΔE₀ = E_0-acceptor - E_0-donor or if you prefer, ΔE₀ = E_0-final/end - E_{0-initial/start}.

In the example above, we can examine the reactants and determine that NAD⁺ is the oxidized form of the electron carrier - it can, therefore, not be the donor. This means that H₂ must act as the electron donor in this reactant. During the reactantion electrons flow onto NAD⁺from the donor H₂creating the reduced product NADH and oxidized product H⁺. To calculate ∆E₀ we say that at the start of the reaction the exchanged electrons are on the donor H₂. We say that at the end of the reaction the electrons are found on NADH. Calculating ∆E₀requires us to evaluate the difference:

E_0-acceptor - E_0-donor

or equivalently,

E_0-final/end - E_{0-initial/start}.

Using a redox table to find E₀ values for the start and end molecules shows us that NAD⁺/NADH has an E₀ of -0.30 while H⁺/H₂ has an E₀ of -0.42.

Therefore, ΔE₀ = (-0.30) - (-0.42) = 0.12 V.

We can see intuitively that this reaction is spontaneous: electrons are flowing from a molecule that "wants" electrons less (E₀ of H⁺/H₂ = -0.42) to a molecule that wants them more (E₀ of NAD⁺/NADH = -0.30).

Reading Different Looking Redox Towers

Novice students of redox chemistry will all undoubtedly run across different ways of representing a redox tower. These different representations may look different but contain the same information. Without explanation, however, reading these tables - when they look different - can be confusion. We will compare and contrast different common forms of redox towers.

Redox Tower: Type 1

Figure 1. A Generic redox tower with oxidized/reduced couple listed with its reduction potential (E₀^') .

Attribution: Caidon Iwuagwu

In this type of redox tower, the oxidized and reduced forms of a molecule are separated by a slash. There is a line drawn from each half-reaction to its redox potential E₀ reported on the vertical axis.

Redox Tower: Type 2

Electron Acceptor	Electron Donor	E₀ (eV)
CO₂ + 24e^- →	glucose	- 0.43
2H⁺ + 2e^- →	H₂	- 0.42
CO₂ + 6e^- →	methanol	- 0.38
NAD⁺ + 2e^- →	NADH	- 0.32
CO₂ + 8e^- →	acetate	- 0.28
S₀ + 2e^- →	H₂S	- 0.28
SO₄^2- + 8e^- →	H₂S	- 0.22
Pyruvate + 2e^- →	lactate	- 0.19
S₄O₆^2- + 2e^- →	S₂O₃^2-	+ 0.024
Fumarate + 2e^- →	succinate	+ 0.03
Cytochrome b_ox + 1e^- →	Cytochrome b_red	+ 0.035
Ubiquinone_ox + 2e^- →	Ubiquinone_red	+ 0.11
Fe³⁺ + 1e- → (pH 7)	Fe²⁺	+ 0.2
Cytochrome c_ox + 1e^- →	Cytochrome c_red	+ 0.25
Cytochrome a_ox + 1e^- →	Cytochrome a_red	+ 0.39
NO₃^- + 2e^- →	NO₂^-	+ 0.42
NO₃^- + 5e^- →	1/2 N₂	+ 0.74
Fe³⁺ + 1e^- → (pH 2)	Fe²⁺	+ 0.77
1/2 O₂ + 2e^- →	H₂O	+ 0.82

In this type of redox tower, each row consists of a half-reaction. The oxidized form of a molecule is shown in the first column, the reduced form of the molecule is shown in the second column. Finally, the E₀ value of the molecule is listed in the third column from the left. The number of electrons transferred to reduce the oxidized form of the molecule is shown in column 1. While the format of the table looks different from Type 1 tower, both contain the exact same information.

Redox Tower: Type 3

oxidized form	reduced form	n (electrons)	Eo´ (volts)
CO₂	glucose	24	-0.43
2H⁺	H₂	2	-0.42 (at [H⁺] = 10^-7; pH=7) Note: at [H⁺] = 1; pH=0 the Eo' for hydrogen is ZERO. You will see this in chemistry class.
CO₂	methanol	6	-0.38
NAD⁺+ 2H⁺	NADH + H⁺	2	-0.32
CO₂	acetate	8	-0.28
S⁰	H₂S	2	-0.28
SO₄^2-	H₂S	8	-0.22
Pyruvate + 2H⁺	lactate	2	-0.19
S₄O₆^2-	S₄O₆^2-	2	0.024
Fumarate	succinate	2	0.03
Cytochrome b_ox	Cytochrome b_red	1	0.035
Ubiquinone; (ox)	Ubiquinone; (red)	2	0.1
Fe³⁺(pH = 7)	Fe²⁺(pH = 7)	1	0.20
Cytochrome c; Fe³⁺	Cytochrome c; Fe²⁺	1	0.25
Cytochrome a	Cytochrome a	1	0.39
Nitrate	nitrite	2	0.42
Nitrate	1/2 N₂	5	0.74
Fe³⁺(pH = 2)	Fe²⁺(pH = 2)	1	0.77
1/2 O₂ + 2H⁺	H₂O	2	0.816

In this redox tower, the oxidized form of a molecule is in the leftmost column, its reduced form is in the second column from the left, the number of electrons transferred is in the third column from the left, and the E₀ is in the far right column.

Again, all of these towers contain the exact same information and are used in an identical manner.

Special note: If you have studied redox chemistry in a formal chemistry course, you might notice two key differences between the towers you use in a biology setting and those used by chemists.

1. In chemistry, the redox towers are flipped relative to those in biology: In chemistry, the molecules with the most positive E₀ are listed starting at the top of the table and the compounds with the most negative E₀ are listed at the bottom. In bioloigal redox tables molecules with the largest E₀ are listed at the bottom while those with the smallest E₀ are listed starting at the top. The biology orientation has the advantage of making it easy to picture electrons spontaneously falling down the table from molecules that "want" the electrons less (lower E₀) to molecules that "want" electrons more (higher E₀).

2. In chemistry, the redox potential for hydrogen (H⁺/H₂)is listed as 0. This is because (a) redox potentials for chemistry are measured under a set of non-biologically relevant standard conditions and (b) hydrogen is being used as the common standard redox potential against which all other redox potentials are measured. In biology, the redox potential for hydrogen (H⁺/H₂)is listed as -0.42. This difference between the chemistry and biology tables comes about because the redox potential for (H⁺/H₂) in biology is measured at a physiological pH of 7.0.

Familiarize yourself with how to read and interpret all three types of redox towers!

Chemistry and Biology Teach Redox Differently

For students who have been taught redox chemistry in a formal chemistry course, biological lessons in redox can sometimes seem like they're talking about something completely different. Not surprisingly, chemists tend to teach the most proper and universally applicable approach to evaluating redox reaction. This approach consists of using a set of rules to formally evaluate whether atoms in a molecule have undergone a change in oxidation state. Meanwhile, biologists tend to approach the discussion of redox reactions by thinking about electron transfers between molecules. It turns out that the approach biologists take is not as rigorous as how chemists approach redox reactions and can sometimes not identify bona fide redox reactions that wouldn't be missed using the chemist's approach. However, since the vast majority of biological redox reactions do involve a transfer of electron (and therefore change in redox states).

Let's look at a specific example to see the differences in approach.

The Chemistry Approach (oxidation numbers):

To evaluate/solve redox reactions in CHEMISTRY, we use the concept of oxidation states/numbers (we will just be saying oxidation number here). The oxidation number of an element refers to how the electrons are shared between atoms in a chemical compound and they tell us about the movement of the electrons in the redox reaction. There are specific rules to assigning oxidation numbers, we will not be going over all of them since they will not be applicable to the redox reactions you will see in General Biology, but here are a few:

A single element has an oxidation number of 0
Fluorine ALWAYS has an oxidation number of -1
Hydrogen has an oxidation number of +1 with nonmetals and -1 with metals.

etc.

For more on calculating oxidation numbers see: <https://chem.libretexts.org/Bookshel...ation_Numbers)>

So to find which elements are reduced/oxidized when given a redox reaction, you must track the change of the oxidation numbers between the reactants and the products. Here is an example:

In the unbalanced reaction NO₃^-+ FADH₂⟶ NO₂^-+ FAD⁺

Using the rules, we observe that in NO₃^-, the oxidation number of Nitrogen is +5. In NO₂^-, the oxidation number of Nitrogen is +3. So because +5 ⟶ +3, N is reduced in this reaction.
We could conduct a similar calculation for key atoms on FAD⁺ and FADH₂ to discover that FADH₂ is oxidized in the reaction.

The Biology/Biochemistry Approach (electron flow):

To evaluate/solve redox reactions in biology/biochemistry, we typically do not use or assign oxidation numbers to evaluate redox reactions. Rather we follow the exchange of electrons between molecules. Fortunately, biology tends to reuse a limited number of electron carriers and redox towers tell us which form of a compound is reduced or oxidized. As mentioned above, the basic approach to redox in biology is to define oxidation as: the loss of electrons. Reduction is defined as: the gain of electrons. Here is an example:

NO₃^-+ FADH₂⟶ NO₂^-+ FAD⁺

Here we examine the reactants and immediately spot the common electron carrier FADH₂, the reduced form of the electron carrier. In the products we observe the oxidized form of the electron carrier FAD⁺. We conclude that FADH₂lost electrons (became oxidized) in the reaction. Since the electrons had to go somewhere they were likely accepted by NO₃^-which then became reduced to NO₂^-. In this case the biologist's model arrives at the same conclusion as the chemist's approach through a more intuitive approach that doesn't require memorizing numerous rules and how to apply them.

In our General Biology class, we take the biology/biochemistry approach to redox. You will not need to know how to calculate redox states in this course.

DISCLAIMER: DO NOT WORRY IF YOU HAVE NOT TAKEN CHEMISTRY YET !! WE WILL NOT BE USING THE CHEMISTRY APPROACH WHEN IT COMES TO REDOX REACTIONS IN OUR CLASS. THE PURPOSE OF THIS IS JUST TO DISTINGUISH AND HOPEFULLY CLARIFY THE TWO APPROACHES FOR STUDENTS THAT MAY HAVE ALREADY TAKEN A CHEMISTRY COURSE!!