Reduction/Oxidation Reactions - Mary and Caidon
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In General Biology, most of the reduction/oxidation reactions (redox) that we discuss occur in metabolic pathways (connected sets of biochemical reactions). Here, the cell breaks down the compounds it consumes into smaller parts and then reassembles these and other molecules 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 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 think about or teach 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 trough a transfer of electron(s) between molecules. Biologists, therefore, 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.
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 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”.
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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 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.
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).
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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+.
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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 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.
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-.
oxidized form |
reduced form |
n (electrons) |
E0’ (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 E0’ 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; 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 |
P840GS (ox) |
PS840GS (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 |
PSIGS (ox) |
PSIGS (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 |
PSIIGS (ox) |
PSIIGS (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 redox 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, 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 following equation:
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
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.