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Bis2A 5.1 REDOX Chemistry and the
∗
REDOX Tower
The BIS2A Team
This work is produced by OpenStax-CNX and licensed under the
Creative Commons Attribution License 4.0†
Abstract
This module will discuss REDOX chemistry basics including the denitions of the terms and several
examples of chemical reactions where REDOX is occurring.
Section Summary
An oxidation-reduction
(redox) reaction is a type of chemical reaction that involves a transfer of electrons
between two compounds or atoms. For example, the transfer of an electron from sodium (Na) to chloride (Cl)
resulting in a positively charged Na and negatively charged Cl is a Redox reaction. An oxidation reaction
strips an electron from an atom in a compound, and the addition of this electron to another compound is
a reduction reaction. Because oxidation and reduction usually occur together, these pairs of reactions are
called oxidation reduction reactions, or
redox reactions.
Redox reactions are common and vital to some
of the basic functions of life, including photosynthesis, respiration, combustion, and corrosion or rusting.
Every red/ox reaction can be thought of as 2 half reactions, in one reaction a compound looses electrons and
in the second reaction a dierent compound gains electrons. The amount of energy transfered in a redox
'
reaction is associated with the dierence in each half reactions' reduction potential, E0 . The electron tower
is a tool that ranks dierent common half reactions (and therefore various compounds) based on how likely
they are to donate or accept electrons.
The lower, more negative, the electrochemical potential for each
half reaction, the higher it sits in the electron tower. Reduced compounds can donate electrons to oxidized
compounds that are below it on the electron tower.
Oxidized compounds can accept electrons from any
compound that are above it in the electron tower. The use of the electron tower will be more evident as we
discuss electron transport chains in a few modules.
note: Sometime a redox tower will list compounds in order of decreasing redox potentials (high
values on top and low values on the bottom). Does this change the redox potential of a compound
compared to a table that lists compounds in increasing order as described above?
∗ Version
1.1: Jan 18, 2016 3:20 am -0600
† http://creativecommons.org/licenses/by/4.0/
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Figure 1: The rusting of iron is an electrochemical process that begins with the transfer of electrons
from iron to oxygen. The iron is the reducing agent (gives up electrons) while the oxygen is the oxidizing
agent (gains electrons). Source: http://tetonmotors.blogspot.com/2013/08/why-do-cars-rust.html
1 Reduction-Oxidation Reactions
In this class we are going to focus on REDOX reactions that are biologically associated. The majority of
the reactions we discuss occur in the context of metabolic pathways (connected sets of metabolic reactions)
where compounds may be consumed by the cell, broken down into smaller parts and then reassembled into
larger macromolecules.
Lets start with some Generic Reactions
Transferring electrons between two compounds results in one of these compounds loosing an electron, and
one of the compounds gaining an electron.
For example, look at the gure below.
If we use the energy
story rubric here to look at the overall reaction we can compare the before and after characteristics of the
reactants and products. What happens to the matter (stu ) 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 follow the movement of electrons from compound
A to B by looking at the change in charge. A looses an electron (becoming positively charged), and in so
doing we say that A has become oxidized.
Oxidation
is loss of electron(s). B gains the electron (becoming
negatively charged), and we say that B has become reduced.
Reduction
is gain of electrons. We also
know, since something happened that energy must have been either transfered and/or reorganized in this
process and we'll consider this shortly.
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Figure 2: A generic redox reaction. The full reaction is A +B goes to A + B . The two half reactions
are shown in the blue box. A is oxidized by the reaction and B is reduced by the reaction.
+
-
oxidized, the electron(s) must then passed to another
reduced. The oxidation and reduction reactions
are always paired in what is known as an oxidation-reduction reaction (also called a red/ox reaction).
Remember the Denitions:
When an electron(s) is lost, or a molecule is
molecule. The molecule gaining the electron is said to be
oxidation
=
loss of electrons
reduction
=
gain of electrons
2 The Half Reaction
To formalize our common understanding of red/ox reactions, we introduce the concept of the half reaction.
Two half reactions are required to make the full red/ox reaction.
Each half reaction can be thought of
as a description of what happens to one of the two molecules involved in the red/ox reaction.
This is
+
illustrated below. In this example compound AH is being oxidized by compound B ; electrons are moving
+
from AH to B
+
to generate A
and BH. Each reaction can be thought of as two half reactions: Where AH is
+
being oxidized and a second reaction where B
is being reduced to BH. These two reactions are considered
coupled, a term that indicates that these two reactions occur together, at the same time.
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Figure 3: Generic red/ox 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 gure 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.
+
+
Exercise 1
+
+
(Solution on p. 9.)
In reaction #1, AH becomes:
a. oxidized
b. reduced
c. degraded
note: If you consider a generic redox reaction and reect back on the thermodynamic lectures what
factor will determine whether a redox reaction will "go" in a particular direction spontaneously and
what might determine its rate?
3 Reduction Potential
By convention we analyze and describe red/ox reactions with respect to
reduction potentials that is, with
respect to the ability of a compound to gain electrons. This value is determined experimentally but for the
purpose of this course we assume that the reader will accept that the reported values are reasonably correct.
We can anthropomorphize the reduction potential by saying that it is related to the strength with which a
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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?
Electronegativity, the tendency of an atom or molecule to pull electrons to itself.
Dierent compounds,
based on their structure and composition have intrinsic and distinct attractions for electrons. This quality is
termed
reduction potential or E
0
'
and is a relative quantity (relative by comparison 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 whose magnitude is proportional to how much more it "wants" electrons than
the standard compound. The relative strength of the compound in comparison to the standard 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 .
note: Rephrase for yourself: What is the dierence between the concept of electronegativity and
redox potential?
4 The Redox Tower
All kinds of compounds can participate in red/ox reactions. A tool has been developed to rate red/ox half
reactions based on their E0
'
values and to help us predict the direction of electron ow between potential
electron donors and acceptors. Whether a particular compound can act as an electron donor (reductant) or
electron acceptor (oxidant) depends critically on what other compound it is interacting with. The electron
'
tower 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 organizes these half reactions based on the ability of compounds to accept electrons, with the most
electronegative at the bottom of the tower and the least electronegative values at the top. In addition each
half reaction is written by convention with the oxidized form on the left/followed by the reduced form on
the right of the slash.
+
For example the half reaction for the reduction of NAD
+
-
NAD /NADH + 2e .
An electron tower is shown below.
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to NADH is written:
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Figure 4: Common Red/ox 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 excellent electron donors are found at the top of the tower. Compounds such as Glucose and
Hydrogen gas are excellent electron donors. Notice, that they are found on the right hand side of the
red/ox pair half reactions. At the other end of the tower lies compounds that make excellent terminal
electron acceptors, such as Oxygen and Nitrite, these compounds are found on the left side of the red/ox
pair and have a positive E value.
0
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Video on electron tower
1
For a short video on how to use the electron tower in red/ox problems click here
. This video was made by
Dr. Easlon for Bis2A students. (This is quite informative.)
Exercise 2: Reading a Redox tower
(Solution on p. 9.)
The right and left sides of the chemical reactions in the redox tower are separated by a "/".
The form of the compound on the left of the slash is____________, and the form of the
compound on the right of the slash is______________.
a. oxidized, reduced
b. reduced, oxidized
c. oxidized at the top of the tower, reduced at the bottom of the tower
d. reduced at the top of the tower, oxidized at the bottom of the tower
What is the relationship between ∆E and ∆G?
0
'
The question now becomes: how do we know if any given redox reaction is energetically spontaneous or not
(exergonic or endergonic) and regardless of direction what the free energy dierence is? The answer lies in
the dierence of the reduction potentials of the two compounds. The dierence in the reduction potential
for the reaction or
E
0
'
for the reaction, is the dierence between the E0
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losing the electrons). In our generic example below, AH is the reductant and B
+
are moving from AH to B .
'
change in E0 or
∆E0 '
Using the E0
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oxidant (the compound
reductant (the compound
for the
getting the electrons and causing the oxidation of the other compound) and the
+
is the oxidant. Electrons
of -0.32 for the reductant and 0.82 for the oxidant the total
is 1.14 eV.
Figure 5: Generic Red/Ox reaction with half reactions written with reduction potential (E ) of the
two half reactions indicated.
0
1
http://youtu.be/HxbgKaTrxH8
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The change in
∆E 0 '
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correlates to changes in
is proportional to a large negative
∆G.
Gibbs free energy, ∆G. In general a large positive ∆E
The reactions are exergonic and spontaneous.
For a reaction to be exergonic the reaction needs to have a negative change in free energy or
will correspond to a positive
∆E0 ' .
0
'
-∆G, this
In other words, when electrons ow "downhill" in a redox reaction from
a compound with a higher (more positive) reduction potential to a second compound with a lower (less
positive) reduction potential, they release free energy (review module 4.0 and 4.1).
∆E0 ' ,
The greater the voltage,
between the two components, the greater the energy available when electron
ow occurs. It is, in fact, possible to quantify the amount of free energy available. The relationship is given
by the Nernst equation:
Figure 6: The Nernst equation relates free energy of a redox reaction to the dierence in reduction
potential between the reduced products of the reaction and oxidized reactant.
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:
∆G and ∆E have an inverse relationship: When ∆G
∆G is negative ∆E is positive. For a review see Red/Ox
What you should notice is that
is positive,
∆E
is negative and when
discussion in the Bis2A Discussion Manual.
Exercise 3
(Solution on p. 9.)
Using the table suggest which of the following could be used as an electron acceptor for an electron
transfer from Ubiquinonered . The subscript "red" indicates that ubiquinone is in a reduced state it has an electron to give to the "right" acceptor. If you would have seen the subscript "ox" this
would have indicated that ubiquinone was oxidized and thus not available to donate electrons, but
rather accept them.
a. FAD
b. NADH
c. cytrochrome a
d. cytochrome c
ox
red
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Solutions to Exercises in this Module
Solution to Exercise (p. 4)
a
Solution to Exercise (p. 7)
a
Solution to Exercise (p. 8)
C
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