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Lecture Comments (2)

1 answer

Last reply by: Professor Hovasapian
Tue Jul 23, 2013 8:10 PM

Post by Gift Nitchie on July 23, 2013

For H+ + e-, can you please explain why it's not H with no charge, since they negate each other? I mean, is that the same with 1/2 H2? Thank you!

Oxidation-Reduction Reactions

Lecture Slides are screen-captured images of important points in the lecture. Students can download and print out these lecture slide images to do practice problems as well as take notes while watching the lecture.

  • Intro 0:00
  • Oxidation-Reduction Reactions 1:32
    • Redox Reactions
    • Example 1: Mg + Al³⁺ → Mg²⁺ + Al
    • Reduction Potential Definition
    • Reduction Potential Example
    • Organic Example
    • Review: How To Find The Oxidation States For Carbon
  • Examples: Oxidation States For Carbon 27:45
    • Example 1: Oxidation States For Carbon
    • Example 2: Oxidation States For Carbon
    • Example 3: Oxidation States For Carbon
    • Example 4: Oxidation States For Carbon
    • Example 5: Oxidation States For Carbon
    • Example 6: Oxidation States For Carbon
    • Example 7: Oxidation States For Carbon
    • Example 8: Oxidation States For Carbon
    • Example 9: Oxidation States For Carbon
  • Oxidation-Reduction Reactions, cont'd 35:22
    • More On Reduction Potential
    • Lets' Start With ∆G = ∆G°' + RTlnQ
    • Example: Oxidation Reduction Reactions

Transcription: Oxidation-Reduction Reactions

Hello and welcome to, and well come back to Biochemistry.0000

Today, we are going to talk about a very, very important topic: oxidation-reduction reactions.0004

In previous lessons, we have talked about ATP; we have talked about phosphoryl transfer.0010

We saw how we can take an endergonic reaction, couple it with the energy available from the hydrolysis of ATP to actually make that reaction take place.0016

We also saw that there were other molecules that have very, very high -ΔGs, and we can couple that with the ATP reaction, reverse the ATP reaction and actually form ATP.0027

The truth is that most of the ATP that is made in the body happens through a very complex oxidation-reduction process, which is essentially what metabolism is.0042

I mean, all you are doing is, you are taking in food, these highly reduced compounds, and what the body does is it oxidizes these things.0051

It takes away the high-energy electrons, and it eventually gives them to oxygen using the energy that is available; that is what produces the ATP.0060

So, directly or indirectly, whether it is some reaction that is taking place or ultimately electrons moving from reduced compounds all the way to oxygen through the electron transport chain, directly or indirectly, oxidation-reduction is responsible for all of the work that is done by the body.0070

This is a profoundly important topic; let’s go ahead and get started.0090

OK, let’s see.0096

Now, of course, you have all seen oxidation-reduction reactions; you have all had general chemistry.0100

I am going to be going through it again, sort of, from the beginning; but there are certain things that I am going to not quite touch on in as much detail.0107

I would strongly urge you to take a look at the oxidation-reductions sections of my AP chemistry course for if certain fundamental things do not really make sense.0115

I go through a full, very complete, slow discussion with multiple examples during those lessons; but hopefully, what we do here will be enough for our purposes.0127

Redox reactions or oxidation-reduction reactions, they involve - hopefully, as you know - the loss and gain of electrons, the loss of electrons from one species and the gain of those electrons from another species, those electrons by another species; and that is it.0138

That is all that is really going on; you know, you do not want to let these sizes of the molecules scare you.0175

All that is happening is that one species is giving up electrons to another species.0181

It might pass through 5, 6, 7, 8 different intermediates, but it is still going to be something oxidized, something reduced.0186

Something loses electrons, something gains electrons.0195

During that process, hydrogens are involved; there is going to be, maybe, some rearrangements.0197

The molecules may not necessarily look the same afterward, but really, it is the electrons that we are concerned about because that is where the energy is.0202

OK, oxidation- it is the loss of the electrons.0211

Reduction is the gain of electrons.0220

Let’s just take a look at an example here, so example 1.0230

We have an equation; let me go ahead and do this in blue.0236

We have magnesium plus, let’s say it comes in contact with some aluminum ion, and it goes to magnesium ion plus aluminum.0241

When we take a look at oxidation-reduction, we are accustomed to looking at oxidation state of the particular species to decide what is being oxidized and what is being reduced.0253

In this very simple example, magnesium has a 0 oxidation state, and it has gone to a +2 oxidation state, so it has been oxidized.0261

Well, aluminum+3 to a 0, it has been reduced.0269

So, magnesium, it has been oxidized; it has lost the electrons and aluminum.0276

It has gained those particular electrons.0285

Now, of course, this is not balanced; the charges do not balance here and that is fine.0289

We will go ahead and deal with the balancing in just a second here, but the idea is something loses, something gains.0293

Now, oxidation-reductions are broken down into 2 half reactions- the oxidation reaction and the reduction reaction.0300

Let’s go ahead and deal with the oxidation here.0307

Magnesium, it turns into magnesium2+ and 2 electrons, so we represent this this way.0312

It has lost 2 electrons to become magnesium2+, and the 2 electrons are on the right-hand side.0320

Now, the reduction process, we have the aluminium ion which is 3+ has been, when you add 3 electrons to it, aluminium becomes aluminium ion, becomes aluminium metal.0326

It has gained 3 electrons; now, we need to balance this, and in order to balance this, we need to cancel the electrons on both sides and then add the equations.0343

In this particular case, we have 2 electrons, 3 electrons.0354

We multiply this equation by 3, this by 2, to convert this to - let me go ahead and do this in red - we have 3Mg goes to 3Mg2+ + 6 electrons.0357

And over here, this one becomes 2Al + 6 electrons, goes to 2 A - I am sorry - Al3 + 2Al metal.0374

And now, the electrons cancel, and when you add this straight down, you get the final balanced equation.0385

You get 3Mg + 2Al3+ goes to 3Mg2+ + 2Al metal- that is it.0393

This is pretty much how all oxidation-reduction reactions are dealt with.0406

Individual half reactions- that is all that is happening there.0410

OK, now, pictorially, what is happening is the following.0415

Should I do this on this...yes, it is fine; I can go ahead and do it on this page.0419

So pictorially, what is happening is the following.0422

You have a magnesium; you have a magnesium, and you have a magnesium.0427

Now, this magnesium has, of course, 2 valence electrons on each one, and you also have an aluminum3+ and an aluminium3+.0432

Well, this does not have any electrons on it; what happens is the following.0442

This aluminium takes that electron, takes that electron, takes that electron; and this aluminium takes that electron, that electron and that electron.0447

It takes it, or magnesium gives it up just depending on your perspective because the aluminium ion has a greater affinity for the electrons than the magnesium does.0457

In a battle of the electrons, it is the aluminium ion that is going to win; and what you end up with is, of course, the magnesium ion, which is now, completely stripped off its electrons, and you have the 2 Als that now, have the 3 electrons.0469

6 electrons are lost from the 3 magnesiums; 6 electrons are gained by the 2 aluminiums- that is it.0470

This is a simple, direct transfer of electrons from one species to another.0477

OK, now, let’s examine this a little bit.0483

Now, when we wrote this equation, when we wrote this down, when we wrote this equation, the magnesium, the aluminium, we wrote it as a foregone conclusion.0487

In other words, we wrote it as if we knew what was going to happen.0534

Well, let’s examine this a little bit.0538

But, what if we had said the following?0543

Instead of writing the equation down, if we had just told you a solution of aluminium chloride is poured onto magnesium dust, how do we know what will happen?0557

How is it that the aluminium will actually pull electrons from the magnesium metal and turn into aluminium metal itself, turning the magnesium metal into magnesium ion?0579

How do we know this, OK, or what if we had said this?0589

OK, what if we had said a solution of magnesium chloride is poured onto aluminium dust?0600

Well, now, it is the magnesium that is in ionized form, and it is being poured onto aluminium dust.0624

Now, it is the aluminium that actually has the electrons; it is the magnesium that is missing the electrons.0631

Are we able to say that magnesium is going to pull electrons from the aluminium and turn into magnesium itself, or what is going to happen?0635

How do we know this?0643

That is the whole point.0644

How do we know what will happen?0645

OK, here is how we know what will happen- something called reduction potential.0647

And you will see reduction potential listed in things called “a table of reduction potential for all kinds of species”.0658

OK, so what is a reduction potential?0664

OK, this is a quantitative - so it is a number - quantitative measure of the extent to which a given species wants to be reduced, wants to gain electrons.0668

It is a quantitative measure of the extent to which a given species wants to gain electrons- that is what a reduction potential is.0708

Now, when we measure things in science, we do not measure them absolutely.0718

We measure them relative to a standard.0724

So, when I say that a certain species has a reduction potential of, let’s say, 0.5V - and we will talk about the unit in just a minute - well, it is 0.5V relative to what?0728

OK, we need a standard reaction that we measure all of the other reactions relative to that.0739

If we have this one thing, and if we measure the reduction potentials of 50 things, it is all going to be relative to this one thing that we have chosen as our standard, as our 0 reduction potential.0749

Then, what we can do is, now that we have that list of reduction potentials, now, we can compare them among each other because they are all relative to a standard.0761

Because they are relative to a standard, they are relative to each other- that is what is going on.0770

We have chosen the following reaction as our standard reduction potential.0774

Let’s see; it is H+ + 1 electron goes to 1/2 H2.0780

You can multiply this by 2; it does not matter.0790

It does not change anything.0791

The symbol for reduction potential is this E, and this is the standard that is set at 0V.0794

OK, and a volt is a Joule per Coulomb.0801

OK, so relative to this reaction, all of the other species, the reduction potentials are measured relative to this.0811

OK, for example, in the case of our magnesium, so Mg2+ + 2 electrons goes to magnesium metal.0818

And in a table of reduction potentials, what you will see is this -2.37V.0832

OK, now, again, reduction potential, we had to choose either a reduction direction or an oxidation direction.0841

We chose the reduction so that we have a standard by which to measure everything.0850

What this means is that if I put these species, the magnesium, metal magnesium ion, hydrogen ion, hydrogen gas, if I put them in the vicinity of each other because 0 is higher, is more positive than -2.37, that means that this H+ will actually take electrons from the magnesium.0855

That means this has a higher reduction potential than the magnesium ion does.0880

So, in a competition for electrons, it is the hydrogen ion that will actually win the battle- that is what this means.0885

The numbers, the higher the reduction potential, that means it has a greater affinity for any electron that happens to be in the vicinity, and usually, the electron that happens to be in the vicinity is going to come from the other species if that species has electrons to give up- that is what is going on.0892

Now, in the case of aluminium, the aluminium is listed like this; again, they are all listed as terms of reductions.0909

There will always be some species on the left plus the electrons on the left plus any other species, but it is always the electrons on the left that are written as reduction potentials.0915

3 electrons goes to Al, and this one is listed as -1.66V- there you go.0926

Now, when you look at this, when you see magnesium ion, the reduction potential for magnesium ion is -2.37.0939

The reduction potential for aluminium ion is -1.66.0948

Well, the -1.66 is higher than the -2.37, which means that in a competition for electrons, the aluminium ion will win.0953

When you put aluminium ion, aluminium magnesium ion, magnesium together, aluminium will actually take electrons from magnesium, not the other way around.0963

That is how we know what will happen.0974

There we go; let’s write this out.0979

In a completion for electrons, the species with the more positive - and we say more positive because as you see, you can have 2 negative potentials, but relative to each other, one of them is more positive - with the more positive reduction potential, it will reduce, thereby reversing the other reaction - that is what is important - causing that species to oxidize.0982

So, going back to the question, if I had a solution of aluminium chloride, which I know is aluminium ion, and if I poured that onto magnesium dust, well, the reduction potential, looking at that, it tells me that in a completion for electron, aluminium ion will take electrons from any species that has the electrons to give.1049

Well, since magnesium is in metallic form, and it has 2 valence electrons, in this particular case, the reduction of aluminium stays as written.1071

That stays like this; let me go to blue.1081

Al3+ + 3 electrons goes to aluminium.1086

OK, and we said that the E for that is -1.66.1091

And because now, we know that the magnesium, it is going to take from the magnesium, so magnesium is going to be oxidized.1097

We reverse that reaction; instead of writing it as magnesium + 2 electrons, goes magnesium ion + 2 electrons goes to magnesium metal, we switch that around, and we write it this way.1104

Magnesium goes to magnesium ion + 2 electrons, and in the process of switching it, we also reverse the potential, so now, it is +2.37.1117

That is how we know; we take a look at the reduction potential, the one that has the more positive reduction potential.1130

We leave that one alone, and the other reaction we flip.1137

Once we flip it, then we go through the process of actually balancing that reaction, and when you balance it and you add everything, then you just add the individual potentials to get the final potential of the complete oxidation-reduction reaction.1140

Once we balance this - right, we multiply this by 2, we multiply this by 3, we cancel the electrons - we ended up with the following equation.1155

2Al3+ + 3Mg goes to 2Al + 3Mg2+.1165

And the potential for that is equal to +0.71V, and I hope that you will confirm the arithmetic for me.1179

OK, this is how we know what will happen if we are not told explicitly what will happen, and oftentime a reaction will be written in such a way, but when you look at the reduction potentials, it is actually going in the reverse direction.1189

This reaction, if you put aluminum ion together with magnesium metal, it will spontaneously move in this direction.1203

That is what this positive total net, this is the reduction; this is the potential.1210

This is the electrical potential for this particular oxidation-reduction reaction.1219

This is the complete reaction; each half reaction has a reduction potential.1224

You decide which one is going to oxidize, which one is going to reduce; you balance the equation.1228

This is our final equation; because this is positive, this reaction will happen spontaneously.1232

Electrons will flow spontaneously without you doing anything from the magnesium to the aluminum ion- that is what is happening here.1238

OK, now, in the case of the magnesium chloride solution being poured onto aluminum dust, magnesium chloride solution and aluminum dust, well, in this particular case, it is the magnesium that is in ionized form, and it is the aluminum now, that actually has - it is in metallic form - it has its electrons.1247

Well, based on the reduction potential, magnesium ion has a -2.37 reduction potential.1277

Aluminum ion has a -1.66.1287

Magnesium ion is not strong enough to actually pull electrons off of the aluminum.1292

So, in this particular case, nothing will happen; absolutely, a reaction will not take place.1298

We can write it down, but just because we can write it down, it does not mean that it will happen.1304

The reduction potential tells us so; magnesium ion is not strong enough to actually pull electrons from the aluminum.1307

That is what is going on here, and it is this is reduction potential that tells us what will happen and what will not happen.1315

OK, here we go; let’s see.1322

Let’s go ahead and look at an organic example now; that is what is going to be important.1327

Let me go ahead and go back to black ink; let me actually start this on another page.1332

OK, well, that is fine; I guess I can go ahead and start here, and then we can continue.1345

Let’s look at an organic example.1349

OK, let’s start off with…well, that is fine.1360

Let’s go ahead and do RC, OH + Cu2+, goes to RC double bond, OOH + Cu2O.1365

OK, in this particular case, we have a couple of things going on.1381

Now, first of all, we have to decide which one is being oxidized and which one is being reduced.1386

Here, we have this aldehyde and we have copper ion.1392

Well, the oxidation state of copper is +2; and over here, oxygen is a -2.1396

There are 2 coppers, so that means, now, the individual copper is in oxidation state of +1.1402

And again, if this idea of assigning oxidation state is still a little hazy for you, by all means, please take a look at the sections in the AP chemistry course under oxidation reduction where I go very, very carefully through this process.1408

It looks like the copper is reduced, which means the carbon is oxidized.1422

Here, we have an oxidation state of +1 on this carbon, and here we have an oxidation state of +3 on this carbon.1427

Now, of course, this is not balanced; but we are not concerned with the balancing, right now.1435

We are just concerned with recognizing what is oxidized and what is reduced.1440

So, carbon is oxidized; copper is reduced.1443

OK, now, let’s go ahead and take a moment to review how we actually find the oxidation states for carbon.1448

I will start that over there; let’s review...oops, there we go.1458

Let’s review how to find the oxidations states for carbon.1469

OK, we start off; for a given carbon, start with an oxidation state of 0.1486

Now, there are different ways to actually represent the oxidation state of carbon.1504

You can go from 0 to 8; you can go from -4 to +4.1508

Again, it is just depends on what you are looking at, in terms of how many electrons that particular carbon owns.1515

That is what oxidation state is; it is a measure.1522

It is a statement, a numerical measure of how many electrons it actually has, that it, actually - you can say - that it owns.1525

The range itself, the specific numbers do not really matter.1533

What matter is what is going on chemically.1537

I would like 0 because I think it is a great way to start.1540

You can sort of see from 0 the gain of electrons, the loss of electrons, so I have chosen that as my standard.1544

I am going to go up and down from there as opposed to, say, 0-8.1552

It is just depends on what particular book you are using, what your teachers teaching you; but hopefully, this will make sense.1556

So, no. 1, OK, for every bond - when you are looking at a given molecule for a given carbon - for every bond to another carbon, there is no change in oxidation state.1564

No. 2, for every bond to hydrogen - and this is actually very, very easy - you add a -1 or subtract 1, depending on how you want to.1589

I think of it as adding a -1 because you are adding an electron.1606

So, when you add a hydrogen, if you see a hydrogen bond into a carbon atom, that means that the oxidation state add a -1 to it.1609

In other words, the carbon is more electronegative than the hydrogens, so the carbon owns that electron.1619

It is carrying that extra negative charge- that is what that means.1623

And, of course, the third one, for every bond to oxygen - or in parentheses, I will put or a sulphur chlorine, basically anything that is more electronegative than carbon etc., but mostly it is going to be oxygen - you add a +1.1627

What that means is that it has lost ownership of that electron; the oxygen or the sulphur or the chlorine has taken the electron away from the carbon.1652

Let’s just do a bunch of examples, go through the range of oxidation states for carbon.1659

I think this is really, really important, so examples.1662

We are going to run through the entire range here.1668

OK, so let’s start off with methane H, H, H, H.1671

We start with an oxidation state of 0, and then we said for every bond to hydrogen, we add a -1.1677

So, we have -1, -1, -1, -1; our oxidation state of methane is -41684

It is a very, very, highly reduced form of carbon; in fact, it is the most reduced form of carbon.1693

It actually owns not only its 4 electrons that it brought to the table, it has also taken away the electrons from a hydrogen.1699

It is carrying a -4 charge; this is why I chose 0.1705

I think, because it tells you how many electrons it has actually taken from other species as supposed to the total.1710

OK, well, let’s take a look at another one; let’s take a look at H3, C, C, C, C, C.1717

This is going to be ethane; we said bond it to another carbon.1728

No change, it is 0; and then 3 hydrogens here, so -1, -1, -1.1733

Now, it is -3; OK, now, the oxidation state on this carbon is a -3.1738

OK, in other words, it owns its 4 electrons and the 3 that came from hydrogen carbon-carbon bond.1747

It is equally shared, so it does not really own anything.1754

OK, let’s go ahead and do this one.1758

Let’s do H2, C, C, H, H.1764

OK, one bonded carbon, another bonded carbon, so this is going to be 0 + 0 and then 2 hydrogens - 1, -1.1770

It equals an oxidation state of -2.1780

OK, let’s try another molecule here; let’s try this H, H, OH.1784

This is ethanol; OK, bonded to carbon, that is a -0 - 1.1792

Hydrogen is -1; oxygen is +1, so you get a oxidation state of -1.1799

OK, let’s try CH3, CH3, CH3.1812

This time, we have a carbon surrounded by 4 carbons.1821

This is just 0 + 0 + 0 + 0, so the oxidation state is 0.1825

This is a nice, basic, normal, everyday carbon; it owns its 4 electrons, and that is it.1831

Nothing is lost; nothing is gained.1837

Everything is good.1839

OK, now, let’s take a look at an aldehyde, so C, H, CH3.1841

OK, this is acetaldehyde.1851

We have the carbon, which is 0; we have the 2 oxygens, so +1 +1 - not the 2 oxygens, the 2 bonds - the 2 oxygens, each bond, and then -1, we get a +1.1854

In the example that we were looking at for the oxidation-reduction, that was some sort of an aldehyde; it had an oxidation state of +11872

OK, let’s take a look at a ketone here, so C.1881

This is CH3; this is CH3.1887

We have carbon, carbon, so it is going to be 0 + 0, and then, +1 +1 for the 2 bonds to oxygen.1892

Now, it has an oxidation state of +2.1899

What that means is it brought 4 electrons to the table; it has lost 2 of those electrons to oxygen.1903

Oxygen has come; it is bound with it, and it has pulled electrons away from it.1909

I mean yes, it is still sharing; it is involved with the bond, but the electrons are with oxygen mostly, not with carbon.1913

That is what that +2 means; it has lost 2 electrons.1919

This is why I think it is best to start with a 0 because from 0, you can tell how much you have gained, how much you have lost.1923

OK, let’s take a look at a carboxylic acid here, so acetic acid.1930

This is CH3, so carbon is a 0; we have +1 and +1 for the 2 bonds to oxygen.1936

And then, we have, oh another +1 for another bond to oxygen; so this is +3.1944

In the example that we were just looking at, the oxidation by copper ion, it went from a +1 state, an aldehyde to a carboxylic acid, which was a +3 state, so it lost 2 electrons.1949

So, +3 means that it brought 4 electrons to the table; it has given up 3 of them to oxygen.1963

It has been oxidized, and let’s take a look at the final.1970

Now, of course, there are multiple variations on this, but now, we have +1, +1, +1, +1.1975

This is a +4 oxidation state; carbon brings 4 electrons to the table.1985

Oxygen has taken away 2 of them; the other oxygen has taken away 2 of them- that is it.1989

Carbon is completely oxidized at this point; when you exhale, you exhale carbon dioxide.1994

It is a waste product; all of the electrons from carbon, those high-energy electrons, have been taken away, used for other purposes.2000

Now, this carbon is completely spent; there is nothing else for it to do.2007

The range, as you see, is -4 to +4; it is those 9 oxidation states including 0.2012

Those are the oxidation states of carbon- that is it.2021

That is all that is going on here; there is certain number of electrons.2025

The body takes those high-energy electrons and uses them for other purposes- that is oxidation-reduction.2029

OK, a couple of things to notice.2035

You know what, I think I am going to write it on this page.2042

Notice how as carbon becomes more oxidized, in other words, as its oxidation state rises from negative to positive, it is losing hydrogens.2045

In biological systems, oxidation will often mean losing hydrogens.2073

Yes, you are losing electrons with those hydrogens, but the electrons usually come in the form of hydrogen, so this is going to be a very, very big deal.2080

When we think about oxidation, we think about the loss of hydrogens.2089

That is not the only way oxidation happens, but to a large extent in physiological systems, that is how it happens.2094

We think of reduction as the gain of hydrogens, gain of hydrogen atom, because the hydrogen atom brings, it has an H and an electron, an H and an electron or maybe a hydride.2100

A hydride has an H + 2 electrons.2111

So, often, oxidation reduction in biological systems will take place like this; we often look at the gain and loss of hydrogens.2114

OK, now, let’s go ahead and return, and talk a little bit more about reduction potential, say a little bit more, so back to reduction potential.2121

Now, this E with a little 0 on top is the standard reduction potential, and standard means that when we ran these experiments to actually get these numbers that we got like for the aluminium and the magnesium - well, hold on a sec, let me just go ahead and finish writing this, it is a standard reaction potential - standard, it means that when we ran these experiments, the concentrations were all 1M.2141

The temperature was 25°C, and if there was a gas involved like hydrogen gas in the reference cell, the gas was pressurized at 1 atmosphere, so standard conditions.2168

Again, a standard, we need a reference.2178

Now, but concentrations in cells are not standard.2181

So, let us introduce an equation, which gives us a way of finding the actual reduction potential of either individual species, the half reactions, of either half reactions or the potential of fully balanced reactions, where we have actually put the half reactions together to find both the oxidation and reduction, our final reaction, our net reaction, our fully balanced reaction.2200

OK, so you remember we actually did this with the ΔG.2276

We said that we have a ΔG standard, but then if the concentrations are different, it is going to change the ΔG.2280

We saw that under standard conditions, ATP -30.5kJ/mol, but under physiological conditions, it could be -50, -55, even up to -60, 65kJ/mol, depending on the concentrations.2286

Well, it is the same thing, and in fact, we are going to start with the same equation, the ΔG; but now, we are going to be talking about electrical potential because it is electrons that are moving,2300

Let’s start with that equation; let’s start with ΔG = ΔG standard + RT ln Q.2311

That is the equation we are going to start with; now, the floor of electrons does work.2327

Anytime there is some sort of oxidation-reduction process taking place, work is being done.2331

There is energy that is often released in the spontaneous process like that.2335

Now, there is a relationship: the ΔG = NF times that of the reaction.2340

So, when we have a balanced oxidation-reduction reaction that has a certain electrical potential - that is this E thing, E of the reaction - well, there is also a ΔG associated with that reaction.2349

This is the relationship between those 2; here, N is the number of electrons that are transferred, and F is the Faraday constant, and it happens to be 96,485C/mol of electrons.2359

A coulomb is a unit of charge, so in 1mol of electrons, there is 96,485 units of charge that those electrons bring.2390

OK, now, let’s go ahead sell ΔG = ΔG standard + RT ln Q.2402

ΔG also equals this; we are going to put this in where we see ΔG, so here is what you get.2412

You get minus - oh this is a negative here, sorry about that - this is -nFE = -nFE standard + RT ln Q.2419

When I divide it by -nF, I am left with the following: E = E standard - RT/ nF ln Q.2434

This equation and this equation, these are the 2 important equations concerning oxidation-reduction chemistry.2446

This is called the Nernst equation.2455

If I have a particular oxidation-reduction reaction or a half reaction and if I know its potential, I can go ahead and calculate the free energy change of that reaction.2461

If I have the standard reduction potential, which is listed in tables, and if I have concentrations that are different or temperatures that are different, I can actually calculate the new potential in order to calculate the ΔG.2471

So, these two equations are what is important; this is really what this lesson is about- these 2 equations.2486

OK, so let’s go ahead and examine these equations by means of an example; that is going to be the best thing to do.2494

Now, let’s see, example.2503

That is the best way to make sense of anything, is to do examples.2510

In the table of reduction potentials, we have the following.2514

We have...let’s see, should I write it out or should I…yes, that is fine.2537

I will go ahead and draw the structure: H3, C, C, H.2543

This is acetaldehyde + 2H+ + 2 electrons goes to…you know what, actually I am going to make this a little bit easier.2548

I am not going to draw out the structures; I am just going to go ahead and write out the words.2558

I think that is better; we have acetaldehyde + 2H+ + 2 electrons goes to ethanol, and its standard reduction potential happens to be -0.197V.2562

OK, well, we also have the following; we have something else in the table, and it says NAD+ + H+ + 2 electrons goes to NADH.2584

Now, do not worry about what NAD+ and NADH are.2598

Right now, we are actually going to be discussing that in the next lesson; for right now, it is just a species.2601

It is the oxidized version of the species in the vicinity of a couple of electrons.2607

When it becomes reduced, it turns into this thing called NADH, and it has a reduction potential relative to the reference hydrogen electrode, and its reduction potential is the following.2612

These are -0.320.2626

Now, the question is “what happens when these species are brought in to contact with each other?”.2633

“When I have some acetaldehyde, some ethanol, some NAD+ and some NADH, what is going to happen?”.2637

Well, take a look at the reduction potentials; this is -0.197, and this is a -0.320.2644

This one is actually more positive than this one, so this will end up staying as written.2651

The acetaldehyde will reduce to ethanol; this one will end up having to reverse.2657

It is going to end up being the NADH; that is going to end up turning into NAD+.2662

Acetaldehyde will reduce, NADH will end up oxidizing- that is what is going to happen.2668

Well, the E of the acetaldehyde is greater than the E standard of the NAD+, so the NAD+ reaction gets reversed always- that is what we do.2679

Under spontaneous conditions, these numbers tell us that is what happens - gets reversed - so what we end up with is the following.2701

Let’s see; we end up with acetaldehyde + QH+ + 2 electrons goes to ethanol, and its standard is -0.197; and the other one gets reversed.2711

So, we write it as NADH goes to NAD+ + H+ + 2 electrons.2733

And because we reversed it, we actually end up writing it as now, a +0.320.2745

We go ahead and we cancel electrons; in this case, there is 2 here and 2 here.2753

We do not have to multiply it by anything to balance it.2758

We cancel this H+ with one of those H+s, and what we are left with is the net reaction.2760

Acetaldehyde reacts with this thing NADH under some slightly acidic conditions, and it turns into ethanol; and it releases NAD+.2766

The net for this, just add this and this; and what you end up with is +0.123- that is all we have done.2789

Aldehyde and ethanol, if you put them together with NAD+ and NADH spontaneously, what is going to happen is the aldehyde will react with the NADH.2802

The aldehyde will reduce the ethanol; the NADH will become oxidized to NAD+; and this is the measure of the extent to how fast it is going to go.2813

Now that I have this, I can actually calculate the free energy change for this based on the equation that I have got.2824

Now that I have this, my free energy change for this reaction is -nFE of the reaction.2831

It equals minus, well, N is the number of electrons that is transferred.2843

We have 2 electrons that are transferred; let me actually write here: 2 moles of electrons are transferred.2848

We have 96,485C/mol of electrons.2857

And, of course, we have the potential, which is +0.123V, which is a Joule per Coulomb - and I just wanted you to see that the units cancel - and when you multiply all this out, you get a ΔG for this reaction, is equal to -23,735J.2868

There you go; that is it.2883

That is all that is going on here.2886

We have some species; we have a table of reduction potential.2888

We have some species; its reduction potential is this.2892

We have another species; its reduction potential is this.2894

Under conditions when these things are brought together, what is going to happen spontaneously - well, spontaneously because this is larger than this - it is going to reverse the other one to induce, it is going to flip it around.2898

This will stay a reduction; this will become an oxidation.2909

We go ahead and write it; we cancel electrons.2913

We get the final balanced equation, so this is the reaction that is going to take place.2916

This reaction will take place spontaneously; we do not have to do anything.2921

It does not mean the reverse reaction will not take place because the reverse reaction does take place.2926

Enzymatically, it can happen; but spontaneously, this is what will happen; and this positive potential tells us that that will happen.2932

This reaction as written, because of this equation, actually gives us the free energy change.2959

Notice, -23.7kJ- that is a very, very highly exergonic reaction.2965

This reaction wants to go forward; that is all that is happening here.2973

OK, thank you for joining us here at

We will continue our discussion of oxidation-reduction chemistry in the next lesson; take care, bye-bye.2980