Monday, May 07, 2007

Oxidation-Reduction Reactions

 
Biochemical reactions are just a complicated form of organic chemistry. Living organisms have evolved enzymes that make these reactions go faster but the underlying chemistry is unchanged.

Cells are constantly having to deal with the problem of shuffling electrons and channeling them to the right place. You might be familiar with the classic fuel metabolism example of glycolysis where the breakdown of glucose to CO2 releases electrons. When a reaction results in the loss of electrons it's called an oxidation reaction. When electrons are gained it's a reduction reaction. Oxidations and reductions always go together since electrons are passed from one molecule to another.


Loss of Electrons is Oxidation (LEO)


Gain of Electrons is Reduction (GER)


LEO says GER



Oxidation Is Loss of electrons (OIL)


Reduction Is Gain of electrons (RIG)


OIL RIG


During glycolysis, the electrons that are released have to be deposited in some sort of electron sink and expelled as waste. If a cell couldn't get rid of its electrons it would build up a huge negative charge.

Oxygen is the electron sink in mammalian fuel metabolism. A molecule of oxygen takes up electrons and combines with protons to make water. The easiest way to see this is to draw the molecules as Lewis structures showing the valence electrons (Pushing Electrons).

There are sixteen electron on each side of this equation for the reduction of molecular oxygen. Remember, reduction is a gain of electrons. This is a half-reaction, there is no corresponding oxidation that provides the electrons so this isn't a valid oxidation-reduction reaction. It just shows the reduction part.

None of the reactions of glycolysis result in the direct reduction of molecular oxygen. In all cases, the release of electrons when glucose is broken down to CO2 is coupled to temporary electron storage in various coenzymes. We have already encountered several of these electron storage molecules such as ubiquinone, FMN & FAD, and NADPH.

We discussed a simple electron transport chain where electrons were passed from pyruvate to NAD+ in the pyruvate dehydrogenase reaction. This is a classic oxidation-reduction reaction.

How do we know which direction electron are going to flow? For example, if ubiquinone is reduced to ubiquinol by acquiring two electrons then where do the electrons come from? Can NADH pass electrons to ubiquinone or does ubiquinol pass its two electrons to NAD+? And where does FAD+ fit? Can it receive electrons from NADH?

The answer is related to the reduction potential of the various electron carriers. In order to understand reduction potentials we need to learn a little inorganic chemistry.

The reduction potential of a reducing agent is a measure of its thermodynamic reactivity. Reduction potential can be measured in electrochemical cells. An example of a simple inorganic oxidation-reduction reaction is the transfer of a pair of electrons from a zinc atom (Zn) to a copper ion (Cu2+) as shown below. When a pair of electrons is removed from zinc it leaves a zinc ion that's deficient in two negative charges (Zn2+). These electrons can be taken up by a copper ion (Cu2+) resulting in a copper atom (Cu) with no charge.

This reaction can be carried out in two separate solutions that divide the overall reaction into two half-reactions. At the zinc electrode, two electrons are given up by each zinc atom that reacts (the reducing agent). The electrons flow through a wire to the copper electrode, where they reduce Cu2+ (the oxidizing agent) to metallic copper. A salt bridge, consisting of a tube with a porous partition filled with electrolyte, preserves electrical neutrality by providing an aqueous path for the flow of nonreactive counterions between the two solutions. The flow of ions and the flow of electons are separated in an electrochemical cell. Electron flow (i.e., electric energy) can be measured using a voltmeter.
The direction of the current through the circuit in the figure indicates that Zn is more easily oxidized than Cu (i.e., Zn is a stronger reducing agent than Cu). The reading on the voltmeter represents a potential difference—the difference between the reduction potential of the reaction on the left and that on the right. The measured potential difference is the electromotive force.

It is useful to have a reference standard for measurements of reduction potentials just as in measurements of Gibbs free energy changes. For reduction potentials, the reference is not simply a set of reaction conditions but a reference half-reaction to which all other half-reactions can be compared. The reference half-reaction is the reduction of H+ to hydrogen gas (H2). The reduction potential of this half-reaction under standard conditions (E̊) is arbitrarily set at 0.0 V. The standard reduction potential of any other half-reaction is measured with an oxidation-reduction coupled reaction in which the reference half-cell contains a solution of 1M H+ and 1 atm H2(gaseous), and the sample half-cell contains 1 M each of the oxidized and reduced species of the substance whose reduction potential is to be determined. Under standard conditions for biological measurements, the hydrogen ion concentration in the sample half-cell is (pH 7.0). The voltmeter across the oxidation-reduction couple measures the electromotive force, or the difference in the reduction potential, between the reference and sample half-reactions. Since the standard reduction potential of the reference half-reaction is 0.0 V, the measured potential is that of the sample half-reaction.
The table below gives the standard reduction potentials at pH 7.0 (E̊́) of some important biological half-reactions. Electrons flow spontaneously from the more readily oxidized substance (the one with the more negative reduction potential) to the more readily reduced substance (the one with the more positive reduction potential). Therefore, more negative potentials are assigned to reaction systems that have a greater tendency to donate electrons (i.e., systems that tend to oxidize most easily).

It's important to note the direction of all these reactions is written in the form of a reduction or gain of electrons. That's not important when it comes to determining the direction of electron flow. For example, note that the reduction of acetyl-CoA to pyruvate is at the top of the list (E̊́= -0.48 V). This is the reaction catalyzed by pyruvate dehydrogenase. Electrons released by the oxidation of pyruvate will flow to any half reaction that has a higher (less negative) standard reduction potential. In this case the electrons end up in NADH (E̊́ = -0.32 V).

The reduction of oxygen is way down at the bottom of the list. That's why it's an effective electron sink for gettng rid of electrons.

Now we'd like to know something about the thermodynamics of these electron transport reactions so we can find out how much energy is available to do useful work. This will lead us to the Nobel Laureate for April 25th.

[©Laurence A. Moran. Some of the text is from Principles of Biochemistry 4th ed. ©Pearson/Prentice Hall]

3 comments :

  1. When I think about some difficult problem in electronics, I always take heart that chemists have it so much harder with thermodynamics of chemical potentials and the complexities of solid-liquid charge interfaces. ;-)

    That said, I feel the need to nitpick some of the EM part of the article.

    "Electron flow (i.e., electric energy)"

    Minor nitpick: Electron flow is one form of free electric current (moving charges). Energy is the amount of work a system can do on another system. Free currents can flow indefinitely without doing current (sic) work, in vacuum or in uncoupled superconductors, consisting potential energy instead. So while the text isn't exactly wrong, it isn't exactly the best picture either, IMHO.

    ("Free current" is to distinguish it from displacement current, without charges. Since I hate having nitpicks nitpicked. :-)

    "Electron flow ... can be measured using a voltmeter."

    Major nitpick: Um, no. Voltmeters measure voltages across electric potentials, where potentials have the potential (sic) to do work.

    There will not be any free current until (a group of) charges have the possibility to move in the potential, for example by bridging it with a wire where electrons can flow. That current is measured by an ampere meter.

    Minor nitpick: Since voltmeters should measure potential without affecting the measured system, they are high resistive to conduct virtually no free current. So one could quibble about if the first picture is misleading or not. For someone trying to make a battery do work, it could be. :-)

    "The direction of the current through the circuit in the figure"

    Minor nitpick: By convention, current flows from a positive source to a negative sink. Current direction is always opposite electron flow.

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  2. ""Electron flow ... can be measured using a voltmeter."

    Major nitpick: Um, no."

    In retrospect, I think I was too nitpicky; must have been lack of coffee. ;-) Replace that with a minor nitpick: Voltmeters doesn't measure current directly. Of course one can rig loads to measure potentials in nodes inside the active circuit such that one can deduce the currents.

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  3. I really appreciate these blogs. It really helps when someone "talks" a subject rather than the typical textbook style that just lays out what "is" and never discusses the "what isn't" or "why isn't it..." or "what could be but isn't because..." issues that help develop perspective.

    I'm still confused about oxidation-reduction and electrochemistry however because I can't bring it back to basic enough physics. It always seems something is glossed over or doesn't quite add up.

    Of course free energy is a very deep concept to me because the simple equilibrium equations cover up the microscopic details of entropy, although equilibrium thinking actually has helped me become more at peace with entropy concepts, free energy, and the directionality of reactions. One problem I have is: I tend to return to the "Feinman's Ratchet" question - if we can get energy from ocean waves, why can't we design a microscopic "ratchet" that winds up whenever random collisions due to thermal motion occur, thus converting thermal energy to potential energy?... that kind of thing.

    But back to electrochemistry. It always involves oxidation-reduction terminology, and I have never really seen any really good discussions that spend a length of time on what is NOT an oxidation-reaction... they just start spewing all these oxidation-reduction reactions. I want to see more comparisons of oxidation-reduction with non oxidation-reduction reactions.

    Now, here's my naive way of looking at it. First, you need something called "electronegativity" to get started, which is that for some mysterious reason (not explained too well), some neutral atoms like electrons so much that, given the chance, they will pull an electron (or electrons) right off another neutral atom that isn't so desperate to have electrons around. How do we quantify that in energies?

    Now we have basically two charged atoms, which seems very unnatural. However, the atoms apparently make a good deal since they can form a "crystal" which apparently gets them both to a lower energy state. All this is due to electronegativity I guess, otherwise each atom would be happy to keep its electrons to itself. So electronegativity is really a powerful force of some kind.

    Now a crystal isn't the only way to get to a lower energy state -- If these atoms happen to be in a medium with polar molecules like water, they will often actually prefer to surround themselves with the toady little polar molecules rather than stay in a crystal. Why and to what extent is not so easy to explain, though we can wave our hands and talk about thermodynamics and entropy. I'd like to see it broken down to bond (or disassociation) energies and statistical probabilities to try to understand it, and see why certain crystals dissolve better than others, etc.

    Now before moving to oxidation-reduction and electrochemistry, let's remember that electronegativity is not the only way for atoms to get to lower energy states. Atoms with similar electronegativities can still form covalent bonds and do something similar, but electrons haven't wildly changed loyalties in that case. So what is their role in electrochemistry if any?

    This is where I run out of steam and need to see clearly what forces are really responsible for moving electrons around in chemical reactions, and really see clearly the competition for electrons in electrochemical reactions, and the energies of these competitions. It's easy to talk about a hydrogen cell with a lot of H+ floating around, but there is also some negative ion floating around that no doubt competes with the electrode to supply electrons to the H+. So what makes the electrode win or lose? I'd like to see it broken down to some forces or energies or electronegativities or something.

    Good blog though.

    -SB

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