1.12: Oxidation-Reduction Reactions

Metabolism

Metabolism refers to the full set of chemical reactions carried out by a biological unit. For example, human metabolism refers to all the chemical reactions that occur in people, while adipocyte metabolism refers to those reactions that occur in fat cells, and lysosomal metabolism refers to the reactions of those organelles.

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 (Electron Transfer)

Most of the reduction/oxidation reactions (redox) that we discuss here occur in metabolic pathways (connected sets of biochemical reactions) and can be described as electron transfers between two molecules. These electron transfers play two fundamental roles. First, they are involved in the building, breakdown and/or conversions of molecules into chemically different forms. Second, in all known forms of life these electron transfer reactions also play critical roles in energy transfer, either between a cell and its environment or within the cell itself. It is therefore important to develop, at the very least, an intuitive understanding and appreciation for redox/electron transfer reactions in biology.

Since redox reactions are generally important, most students of biology will also study this chemical process in chemistry courses. Unfortunately, biologists and chemists tend to approach teaching this common foundational topic in different ways (even if they're both teaching the same thing). That can be confusing for new students.

Chemists often introduce the concepts of oxidation and reduction using the concept of oxidation states. See this link for more information: <https://chem.libretexts.org/Bookshel...ation_Numbers)>. Chemists usually ask students to apply a set of rules (see link) to determine the oxidation states of individual atoms in the molecules involved in a chemical reaction. The chemistry formalism defines oxidation as an increase in oxidation state and reduction as a decrease in oxidation state.

However, biologists don’t typically teach the topic in this way. Why? It's probably because most of the redox reactions encountered in biology involve a change in oxidation state that comes about through a transfer of electron(s). Consequently biologists, typically define reduction as a gain of electrons and oxidation as a loss of electrons. We note that the biological electron-exchange view of redox reactions is entirely consistent with the more general definition associated strictly with changes in oxidation states. The electron-exchange model does not, however, explain redox reactions that do not involve a transfer of electrons, which sometimes occur in the context of a chemistry class. The biologist's view of redox chemistry has the advantage (in the context of biology) of being relatively easy to create a mental picture for. There are no lists of rules to remember or much inspection of molecular structure involved in developing at least a basic conceptual picture of the topic. We simply imagine an exchange between two parties - one molecule handing off one or more electrons to a partner who accepts them.

Since this is a biology reading for a biology class we approach redox from the “gain/loss of electrons” conceptualization. If you have already taken a chemistry class and this topic seems to be presented a little different in your biology course, remember that at its core, you are learning the same thing. Biologists just adapted what you learned in chemistry to make more intuitive sense in the context of biology. If you haven’t learned about redox, yet don’t worry. If you can understand what we are trying to do here, when you cover this concept in chemistry class you will be a few steps ahead. You will just need to work to generalize your thinking a little bit.

Let's start with some generic reactions

Transferring electrons between two compounds results in one of these compounds losing 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.

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Figure 1. Generic redox reaction with half-reactions Attribution: Mary O. Aina

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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 nor do they 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 must "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.

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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

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• 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

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 or the acceptor) involved in a "full" redox reaction. A "full" redox reaction requires two half reactions. In the example below, half reaction #1 depicts the molecule AH 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 neither can happen on its 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 not 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. In the figure below we call that reaction "Reaction #1". When this happens, we say that the two half reactions coupled. NOTE: In Chemistry class, because there are sometimes more than two reactions involved in a complete redox reaction, instead of the term "coupled reaction", you might encounter the more generic terms "simultaneous reactions" or "consecutive reactions". In all cases, these terms are meant to communicate that the reactions must happen together, either at the same time or in sequence.

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Figure 3. Generic redox reaction where compound AH is being oxidized by compound B⁺. Each half reaction represents a single species or compound to either lose or gain electrons (and a subsequent proton as shown in the figure above). In half reaction #1 AH loses a proton and 2 electrons: in the second half reaction, B⁺ gains 2 electrons and a proton. In this example HA is oxidized to A⁺ while B⁺ is reduced to BH.

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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 - in Chemistry class, these electrons might be referred to as "intermediates" because they do not appear in the overall redox reaction - 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 its 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.

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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?

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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 the same. 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.

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REQUEST

As you go through the Redox section of the reading please use NB to also comment on parts of the reading that are hard to understand, confusing, and places that you think are well explained and clear. If you have suggested edits please let us know. Use the following emoji to tag these comments:

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Table of Standard Reduction Potentials (aka. The Redox Tower)

All kinds of compounds can take part in redox reactions. Scientists have developed a graphical tool, a table of standard reduction potentials (aka. a 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 to change in 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 redox 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^-.

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oxidized form	reduced form	n (electrons)	E^0’ (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 E^0’ 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 + Pi	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; Fe^3^+	Cytochrome b2; Fe^2^+	1	0.12
Fe^3^+(pH = 7)	Fe^2^+(pH = 7)	1	0.20
Complex III Cytochrome c1; Fe^3^+	Cytochrome c1; Fe^2^+	1	0.22
Cytochrome c; Fe^3^+	Cytochrome c; Fe^2^+	1	0.25
Complex IV Cytochrome a; Fe^3^+	Cytochrome a; Fe^2^+	1	0.29
1/2 O2 + H2O	H2O2	2	0.3
P840^GS(ox)	PS840^GS(red)	-	0.33
Complex IV Cytochrome a3; Fe^3^+	Cytochrome a3; Fe^2^+	1	0.35
Ferricyanide	ferrocyanide	2	0.36
Cytochrome f; Fe^3^+	Cytochrome f; Fe^2^+	1	0.37
PSI^GS(ox)	PSI^GS(red)	.	0.37
Nitrate	nitrite	1	0.42
Fe^3^+(pH = 2)	Fe^2^+(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

Table 1. Common redox tower used in Bis2A. By convention the tower half reactions are written with the oxidized form of the compound on the left and the reduced form on the right. Compounds that make good electron donors have highly negative reduction potentials. Compounds such as Glucose and Hydrogen gas are excellent electron donors. By contrast compounds that make excellent electron acceptors, such as Oxygen and Nitrite are excellent electron acceptors.

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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 ΔG⁰'?

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, 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.

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

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∆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 following equation:

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Figure 5. The relationship between 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

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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 ∆E⁰' and ∆G⁰' are opposite one another. When ∆E⁰' is positive, ∆G⁰' will be negative. When ∆E⁰' is negative, ∆G⁰' will be positive.

Alternative Views 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 encounter 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.

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Finding ΔE⁰'

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 of kiwi fruit in the form of 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 property "kiwi-potential". We can then propose 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 kiwi 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_{person 1} < Kp_{person 2}. 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_{person 2} - Kp_{person 1}.

Doing it with electrons instead of kiwi fruit: To find ∆E^0' 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 E^0' to one with a higher E^0'. 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 ΔE^0', you solve for ΔE^0' = E^0'_-final/end - E^0'_{-initial/start}. Alternatively, you can think of it as ΔE^0' = E^0'_-acceptor - E^0'_-donor.

When evaluating a redox reaction for E^0' you therefore need to:

1. 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).
2. 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.
3. Third, calculate ΔE^0' = E^0'_-acceptor - E^0'_-donor or if you prefer, ΔE^0' = 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 reaction electrons flow onto NAD⁺ from the donor H₂ creating the reduced product NADH and oxidized product H⁺. To calculate ∆E^0' 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 ∆E^0' requires 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 E^0' values for the start and end molecules shows us that NAD⁺/NADH has an E^0' of -0.30 while H⁺/H₂ has an E^0' of -0.42.

Therefore, ΔE^0' = (-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 (E^0' of H⁺/H₂ = -0.42) to a molecule that wants them more (E^0' 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 confusing. 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 (E^0'). 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 E^0' reported on the vertical axis.

Redox Tower: Type 2

Electron Acceptor	Electron Donor	E^0' (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 box + 1e^- →	Cytochrome bred	+ 0.035
Ubiquinoneox + 2e^- →	Ubiquinonered	+ 0.11
Fe^3^+ + 1e- → (pH 7)	Fe^2^+	+ 0.2
Cytochrome cox + 1e^- →	Cytochrome cred	+ 0.25
Cytochrome aox + 1e^- →	Cytochrome ared	+ 0.39
NO3^- + 2e^- →	NO2^-	+ 0.42
NO3^- + 5e^- →	1/2 N2	+ 0.74
Fe^3^+ + 1e^- → (pH 2)	Fe^2^+	+ 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 E^0' 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)	E^o´ (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
S^0	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 box	Cytochrome bred	1	0.035
Ubiquinone; (ox)	Ubiquinone; (red)	2	0.1
Fe^3^+(pH = 7)	Fe^2^+(pH = 7)	1	0.20
Cytochrome c; Fe^3^+	Cytochrome c; Fe^2^+	1	0.25
Cytochrome a	Cytochrome a	1	0.39
Nitrate	nitrite	2	0.42
Nitrate	1/2 N2	5	0.74
Fe^3^+(pH = 2)	Fe^2^+(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 E^0' 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 E^0' are listed starting at the top of the table and the compounds with the most negative E^0' are listed at the bottom. In bioloigal redox tables molecules with the largest E^0' are listed at the bottom while those with the smallest E^0' 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 E^0') to molecules that "want" electrons more (higher E^0').

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 Slightly 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. The oxidation state 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 states, 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:

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

etc.

For more on calculating oxidation states 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 states between the reactants and the products. Here is an example:

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

1. 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.
2. 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 FADH₂ lost 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 oxidation 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!!