The Science of Everything Podcast - Episode 37: Oxidation and Reduction

Episode Date: November 22, 2012

An overview of oxidation, reduction, and redox reactions, including a discussion of the definitions of these basic concepts, and an explanation of how they relate to oxidation number and electronegativity. We then apply these concepts to several common examples of redox reactions, including combustion, rust, batteries, fuel cells, and respiration.

You're listening to The Science of Everything Podcast episode 37, oxidation and reduction. And I'm your host, James Fodor. So it's been a little while since the last episode. Sorry about that. I've been rather busy with my uni subjects, and I actually did five instead of the usual four subjects last semester, so that meant I had less time than usual. But fear not.

I'm on holidays now, which means more extra time, or a lot of extra time, for putting our podcast episodes. And today, we're going to be looking out oxidation and reduction, which is a very important topic in chemistry. Specifically in this episode, we'll be defining what we mean by oxidation and reduction and talking about how that works, what the concept means. And then basically we'll be applying that basic idea to a number of different applications, including combustion, which is a type of redox reaction, or oxidation reduction reaction, rust, batteries,

fuel cells and respiration, all of which are examples of or significantly involve the concept of oxidation reduction reactions and henceforth I'm going to refer to these as redox reactions which is just a word that's taken from reduction and oxidation redox it's just a much shorter way of referring to reduction oxidation reactions okay so let's get started oh I should have said before prerequisites for this episode or recommended pre-listening include episode 23 on chemical reactions and episode 15 on chemical bonding they're the two most important ones that you'll want to have listened to before this one. First we'll start with some basic definitions. What do we mean by reduction and oxidation? A redox reaction very simply is simply any chemical reaction in which atoms have their

oxidation state or their oxidation number changed, and Fear Not will define these terms momentarily, but a very large portion of all chemical reactions are redox reactions in one form or another. So one of the other big classes of chemical reactions are acid-base reactions, which we'll look at in the future episode. But between acid base and redox reactions that have come, we've already encompassed virtually all the chemical reactions that we're generally interested in. Okay, so, redox reactions are a very broad class of reactions, but what will we mean by oxidation and reduction? Any redox reaction involves sort of two, not exactly stages, but two aspects. Oxidation and reduction. In other words, in any redox reaction, some thing,

some atom or some molecule or some substance must be oxidized, and some, and some, some other substance must be reduced. These are sort of opposites of each other. They go together. You can't have one without the other in a given reaction. Of course, you can look at aspects of partial reactions they're referred to where you just look at the oxidation or just look at the reduction part, but overall they have to occur together as part of the same reaction.

So what is oxidation and reduction that I keep talking about? Oxidation is simply the loss of electrons, or alternatively it's an increase in the oxidation state or oxidation number by a molecule, atom or ion. So in this entire podcast, basically, we're thinking about this can happen to molecules, atoms, or ions. It doesn't really matter too much for our purposes. Probably we'll mostly be talking about atoms, because that's the easiest way of thinking about it, but it can be molecules as well.

So oxidation is the loss of electrons, reduction is the opposite. It's the gain of electrons. So you see why they have to occur together, because electrons sort of can't disappear or can't appear out of thin air. If electrons are lost from some atom or molecule, they have to then, be gained or reappear in some other atom or molecule. Now, I know what you're thinking. Couldn't the electrons just go and float away and be non-attached?

In other words, be free electrons, not attached to any particular atom or molecule. Well, yes, that is possible, but it's very rare. Essentially because electrons are negatively charged, and if there's been some material or molecules or atoms that have been losing electrons, they're going to be positively charged, and therefore the negatively charged electrons will be attracted to them, and therefore the electrons won't. remain unbound for very long. And if electrons are just floating around in the environment or in any environment, it's likely that before long they're going to come across something that's positively

charged or partially positively charged and they'll be bounded by that. So basically, it's just very rare that electrons are found unattached to atoms. And so in practice, loss of electrons by some atom means gain of electrons by another atom. So the way that I remember this is by using the mnemonic Leo says GER that's LEO says GER

in other words loss of electrons is oxidation and the gain of electrons is reduction Leo says GER as in Leo the line says gur

but there are other ones as well that's just the one that I learned in chemistry now remember I had two definitions here oxidation is the loss of electrons or an increase in the oxidation state

those might sound like they're completely different things but I'll explain what that means momentarily just think about oxidation as loss of electrons, reduction as the gain of electrons. The idea of reduction, by the way, is that when you gain electrons, your oxidation state decreases, and so it's reduced your oxidation state, hence reduction. The word oxidation comes from the fact that oxygen is just a really good oxidizer. In other words, oxygen is really good at grabbing electrons

from other atoms. Now, here's another source of confusion, which still confuses me sometimes, is the difference in the terms oxidation and oxidizing and oxidant, which are all very, sound, very similar, but refer to different things. So let's see if we can really clarify these concepts to avoid confusion. A reducing agent is itself always oxidized, whereas an oxidizing agent or an oxidizer is always reduced. So let's be clear about that. Remember, oxidation and reduction go together there, sort of two signs of the same coin, the yin and the yang, if you like. So a reducing agent is something, molecule or atom or whatever, that causes something else to be reduced. So a reducing agent, it's the agent of reduction.

It causes reduction. It causes something else to gain electrons. So if you're causing something else to gain electrons, if you're being a reducing agent, in other words, then you have to be giving up electrons. In other words, the electrons have to be coming from somewhere. So you're causing something else to gain electrons, you're giving up electrons. Therefore, you are being oxidized. you're losing electrons.

Conversely, an oxidizing agent is something that causes something else to be oxidized. In other words, you're an oxidizing agent or an oxidant, you're causing something else to lose electrons. So what's happening to those electrons? You have to be gaining them. You're taking the electrons away, which means you yourself are being reduced. So, for example, when we say that oxygen is a good oxidant or a good oxidizing agent, what we're saying is that oxygen causes the loss of electrons by other atoms or molecules.

Where are the electrons going? Oxygen is taking them. So when we say oxygen is a good oxidant or a very good oxidant, it means that it's great at stealing electrons from other atoms or ions or molecules. Conversely, if we say that something is a great reducing agent, it means that it's great at giving up its electrons to other atoms. Okay, so, again, quick recap, because it's very easy to get confused here. Oxidation is the loss of electrons. If I'm a good, oxidizing agent, it means I'm good at causing other things to lose electrons, which means I'm

great at grabbing electrons myself. Reduction is the gain of electrons. If I'm a good reducing agent, it means I'm good at causing other things to gain electrons, which means I myself am losing electrons. Okay, so hopefully we've got a basic idea of what I mean by oxidation reduction and how those concepts relate to each other. And remember, a redox reaction is simply one in which you have oxidation and reduction occurring. So something's being oxidized, something's being reduced, that's just a redox reaction. But why does this happen and how can we predict when it happens? What's the significance? What's the, what's happening there at the sort of atomic or subatomic level? Well, to understand this, we have to introduce the concept of oxidation number

or oxidation state. Now, if you'll remember, when I originally defined oxidation and reduction, I defined them as loss and gain of electrons, respectively, but I also said, or an increase or decrease in the oxidation state or number. So oxidation state and number, there's a subtle difference there, but I'm just going to use the terms interchangeably, because that's good enough for us. So now I'm going to explain more precisely what I mean by this oxidation state. Okay, so in chemistry, the oxidation state is an indicator of the degree of oxidation of an atom in a chemical compound. That's the formal definition, but that might not seem very helpful because we just said that oxidation is the increase in the oxidation state, where the oxidation state is defined as

the degree of oxidization. That seems to be rather circular, and actually a lot of the definitions in science are circular. Look at physics, for example, for that. But we can sort of unpack this concept a bit more. More formally, what the oxidation state orxation number refers to is the hypothetical charge that an atom would have if all of the bonds in the molecule or ionic lattice or whatever that it is in, were 100% ionic. Now remember, there are different types of bonds, ionic, covalent, etc.

And in covalent bonds, the different atoms share electrons, whereas in ionic bonds the electrons are completely exchanged from one atom to the other. The basic idea of an oxidation state is to try and quantify how many electrons a particular atom in a given compound has.

The trouble is, of course, in many covalent compounds, electrons are shared between atoms, and so what we have to do is there are various rules for adjusting for this, and these are very complicated, and we need to go through them here, because the details are not that important. What we want to understand is that the concept of oxidation number basically tries to say which atom is most strongly holding on to a particular electron. So who's sort of mostly got the electron, and therefore we assign the electron to that atom, or to the other atom, that may be holding it more

stronger, or whatever. So it's sort of like taking all of the shared electrons and assigning them to whatever atom is most strongly holding onto them. And by that process, we give each atom in a compound or molecule an oxidation number. So, and this oxidation number is usually an integer, that is a whole number. It can be positive or negative or zero. Sometimes they can be fractions, but that kind of makes it even more confusing because basically we're talking about what the charge would be if all of the bonds were ionic. So that would tend to mean that you either have an electron or you don't. There can be various reasons why that doesn't quite work out, but for the moment, we'll just ignore the fractional numbers and assume that the numbers, that oxidation numbers can only

be integers. So if you had a positive oxidation number, that means that you've lost electrons, because effectively your charge has gone down. So that atom, with that positive oxidation number, because every atom has its own oxidation number in a given chemical compound or molecule, that atom has lost electrons. Conversely, if you have a negative, oxidation number, then you've gained electrons. The more negative you are, the more electrons you've gained. And if you've got a zero oxidation number, that means that you're, effectively, you're just the normal full-adishol charged, sorry, not full-adishol, but the normal uncharged atom. So a lone atom by itself, that is non-ionized, would have zero oxidation number. But when you put them in molecules or compounds of various sorts,

their oxidation number will change based on the interactions that they have with other atoms that are in that molecule. So the key point to understand is that the oxidation number or state of an atom depends upon the molecule or compound that it is in. So the same atom can have different oxidation states depending on whether it's bound with one element or another, or in one chemical structure or another. So it's dependent upon the bonding arrangements and not on the actual atom itself. Well, I mean, it does depend on the atom itself, but it also depends crucially on what other atoms are there. So basically the oxidation number is detainable. determined by the relative electronegativities of the constituent elements in the compound or molecule.

So we talked about electro-negativities in the episode on chemical bonding. For reference, that was episode 15, if you want to go check that out again. But just a quick review, electro-negativity is the chemical property of a given element. So all atoms of the same elements have the same, all atoms of the same element have the same electro-negativity. There may be slight differences with different isotopes coming to think of it, but I doubt that that would be a large effect. so for the moment we'll say that it just depends on the element.

Electron negativity describes a tendency of an atom to attract electrons towards itself, or in other words, the tendency to hold on to electrons. So, electronegativity measures the pulling power that an atom has for electrons. So clearly, if you have an atom that has a high electronegativity, that means it's going to be pulling electrons towards itself away from other atoms in a given compound or molecule. What does that mean? Well, remember, an oxidizing agent,

is something that causes other atoms to lose electrons. So a reducing agent is something that's gaining electrons itself. Well, why is that happening? Because, generally, oxidizing agents or strong oxidants, have high electro-negativities. And this is really the key concept that we need to understand. The relationship between oxidation or oxidants and electro-negativity. So oxygen has a high electro-negativity, and therefore it is a good oxidant. Both of these two statements, good oxidant and high electronegativity, effectively mean the same thing. Again, for our sort of high level of abstraction purpose of analysis,

they both mean that you're good at pulling electrons. So low electronegativity and poor oxidizer, or in other words, you're a reducing agent, means that you have poor pulling power of electrons, so people tend to pull electrons away from you. Electronigativity is just sort of the intrinsic ability of an atom or element to attract electrons. depend upon what other elements is bonded to or the compound that it's in. However, whether or not you actually sort of win the fight, so to speak, that is whether you actually get the electron or not depends upon the electron negativity of other atoms that are in that compound or molecule. And this is where oxidation states come in, because oxidation states essentially says who has won

the fight, who has actually pulled the electron or got the electron, or at least got more of the electron because remember we sort of average it out based on who has it, who has the electron more often, because remember, in covalent bonds, it's not a, it's not a, it's not a zero-one thing. It can be degrees of having the electron. Electron negativity says, what is your pulling power of electrons? Oxidation state says how many electrons do you actually have at the moment, where low, negative oxidation states are indicated that you've got a lot of electrons, whereas high electronegativity indicate that you have strong pulling power for electrons, so don't get confused there, that inversion is effectively because electrons are negatively charged. So the oxidation state

represents a charge, so if you have lots of electrons, you're a big negative. Electron negativity just represents pulling power and sort of so that's, there's no negatives there. So more pulling power just means a higher value of electronegativity. Oxygen is such a great oxidant because it has one of the highest electronegativities on the periodic table. So you pair up oxygen with almost any other element and it will win in a sense and pull the electrons from the other So that means that oxygen in that case is acting as oxidant, an oxidizing agent, causing the loss of electrons in other elements by gaining electrons itself. However, if you pair oxygen with fluorine, which is the only other element to have a higher electronegativity than oxygen, then fluorine will win out and it'll grab the electrons because it has a higher attractive force for electrons is higher. and therefore oxygen will actually be oxidized, whereas normally it's reduced, normally it gains electrons,

but in this case it would be oxidized because it would lose electrons to the fluorine. So whether a given element or a given atom acts as an oxidizing agent or reducing agent depends upon what other elements it's paired with in the given molecule or compound. Okay, so, to do a bit of a recap, redox reactions are defined as chemical reaction, in which one or more atoms have their oxidation state, or I guess two or more atoms have their oxidation state changed. And the oxidation state is essentially defined as sort of some kind of weighted average of how many electrons each atom has, averaging out from the effect of sharing of electrons. So it's not the same as the charge, the actual ionic charge on each atom, but it's

related to it. So when I started the podcast, or at the very start of the podcast, I defined redox reactions as a transfer of electrons. Now you see why that's not quite correct, because it doesn't actually have to be a physical transfer of electrons. It can simply be a change in which atom has relatively more pull on the electrons than another one. Okay, so that's what a redox reaction is. And we have, remember, two sides to any reidox reaction, oxidation, which is the loss of electrons, and reduction, which is the gain of electrons. Oxygen is a good oxidant, which means that it causes a lot of electrons in other atoms, meaning that it itself is reduced, that it gains electrons, and it's very easy to get these confused. So just think about who is doing the

gaining and who is doing the losing, and remember the mnemonic Leo says, go. Loss of electrons is oxidation, gain of electrons is reduction. Okay, so now that we've got the basic concepts down, or hopefully, more or less down, we'll try and consolidate them and further our understanding by looking at a number of cases of redox reactions or examples of where they occur and how we can use the concept of redox to understand a wide variety of chemical reactions. in different situations. So first of all, we'll start with one of the most famous examples of a redox reaction, which is combustion. So combustion is, well, I mean, the more common word for combustion is just burning or fire. Specifically, though, combustion is a redox reaction in which oxidation is

very rapid and is accompanied by heat and usually the emission of light. Generally, combustion involves oxygen, although, I mean, it depends how you define it. If you define combustion such that it requires oxygen then by definition it does, but there are combustion reactions which, I mean, are very similar to what we would call fire or combustion, which don't involve oxygen, but most of them do. So if you burn logs, if you burn oil, the, you know, fueling, the spurn jet engines, anything like that, all of that's combustion. It's all fundamentally the same reaction, although the particular fuel might be somewhat different. So to explain a bit more detail what's happening there, when any substance that contains carbon is burned completely, then the carbon forms carbon dioxide.

When a substance that contains hydrogen is burned completely, the hydrogen forms into water as a result of, both of these are remembers as a result of reaction with air. So if you get carbon and combine it with oxygen, you get carbon dioxide, which kind of makes sense. And if you get hydrogen and combine it with oxygen, you get a compound, sorry, a molecule that has hydrogen and oxygen, which is called water. So those two reactions kind of make sense. Now, remember hydrocarbons, this comes back to the episode we did on biochemistry basics. Hydrocarbons are just molecules, often very large molecules, or at least large-ish molecules, that contain just carbon and oxygen.

And we can sort of extend that to, they contain mostly carbon and oxygen and some other bits and pieces of other functional groups. But anyway, for the moment, just imagine long backbones of carbon with, I'm sorry, did I say oxygen? I meant to say hydrogen, carbon and hydrogen, hence hydrocarbons. So think of long backbones of carbon with hydrogens around the edges. Hydrocarbons are what comprise oil, coal, natural gas, all of those things that we think about as fuel sources. So we've said that when carbon burns it forms carbon dioxide, when hydrogen burns it forms water, well, when you burn a hydrocarbon, in other words, both at once or both bonded together, you get both of them.

You get CO2, carbon dioxide, and H2O, water. So effectively what you're doing is you're taking some combination of C and H, carbon and hydrogen, and combining it with oxygen, the O2 molecules of the atmosphere, and what you get is carbon dioxide, CO2, and H2, water. If you write this reaction down and balance it, you'll see it all works out. So we've got the same elements on both sides of the chemical reaction. So that sort of stylized model that I just gave you is what happens when there's complete combustion. In other words, all of the carbon, all of the hydrocarbon is burned, and we have exactly the right amount of oxygen we need for that, and therefore the only byproducts are carbon dioxide and water.

That's called complete combustion, and it's very rare because in practice you don't have the exact right amounts of oxygen, and it's unlikely that all of your hydrocarbon source will be burned, and there's also likely to be impurities and other factors which affect the reaction. So in practice, you usually get partial combustion, or some complete combustion, and a lot of partial combustion. In particular, if there's insufficient oxygen to completely burn the carbon or the hydrocarbon, then what happens is some of the carbon

is converted into carbon monoxide and various other forms of pure carbon, for example, soot and ash. So the smoke plumes that you see coming out of particularly something like a coal power power plantar, for example, well, I mean, there's a lot of stuff that's in there, and nowadays there's various scrubbing technologies which are used to extract many of the pollutants from there,

but sort of traditionally, at a simple level, a lot of what that stuff is, is sort of an ash, which is basically just carbon, just pure carbon. Now, again, there are going to be impurities there, but you can just think about it as pure carbon that's in particular form suspended in the air. The carbon dioxide is basically invisible, so you're not going to see that. What you actually see, the black stuff, that's sort of ash.

Carbon monoxide is also invisible, however, it's a very toxic gas, which is why it can be problematic. And if you think about why it's formed when there's why it and also pure carbon are formed when there's insufficient oxygen. If you just think about it, CO2 requires one, two oxygen atoms for every carbon atom. So you need a fair amount of oxygen for that to happen. If you don't have enough oxygen, then some of the carbon atoms will only get to bind with one oxygen atom, in which case they film carbon monoxide, or maybe they won't get any oxygen atoms, in which case they just remain pure carbon.

It'll become pure carbon. So it kind of makes sense that that would happen if you have incomprehendium. complete combustion. Also, any combustion in atmospheric air, which is 78% nitrogen, will create various forms of nitrogen oxides, that is nitrogen bonded with oxygen, because in addition to the carbon interacting with the oxygen, you've also got the nitrogen in the air, which interacts with oxygen in the air, catalyzed by the high temperatures of the reaction. But again, that's an additional complexity that we don't want to worry too much about when we're just talking about the combustion in an abstract sense.

So I've been talking a lot about combustion, but I haven't really explained how it's a redox reaction. So where is the reduction in oxidation occurring? Well, you might have already guessed because this reaction involves oxygen. And oxygen is a great oxidant. That means it causes other atoms to lose electrons, or in other words, it gains electrons itself. Because remember, oxygen has a very high electronegativity, so it's grabbing all those electrons. In particular, in this case, it's grabbing them from carbon and hydrogen, which have much lower electronegativities than oxygen does.

And so oxygen is being reduced because it's gaining electrons, and carbon and hydrogen are being oxidized because they're losing electrons. The electrons are moving away from the carbon and hydrogen and moving towards the oxygen. Now, of course, what we're actually forming are molecules that contain carbon oxygen and hydrogen and oxygen, namely our carbon dioxide of water again. But remember, the actual definition of a redox reaction is when the oxidation number or state changes. And so even though carbon dioxide and water are both

covalent molecules where the electrons are technically shared, remember that according to the oxidation number, we allocate the electrons to whichever atom is most strongly pulling them, and that, of course, is oxygen. And so the oxidation state of oxygen is reduced because it's pulling on those electrons getting up a big negative charge, so you get a big negative number, and that of the carbons and hydrogens is being increased because they're losing electrons that they previously had. You can see why combustion is a redox reaction. There's one thing that I need to mention, which is the fact that if oxygen is pulling on electrons more strongly than carbon and hydrogen, if that meant many other elements as well,

then why doesn't oxygen just eat up all of the electrons and everything else lose out? In other words, why aren't fires occurring all the time? Why is it difficult potentially to start? Well, in many cases, it's difficult to start a fire. It's difficult to get this reaction going. Well, the reason is because remember from one of our previous chemistry episodes, the concept of a catalyst, that in order to reach a lower overall energy state, you often need to, for a while, move into a high energy state.

Remember, we can consider the energy states of different molecules or atoms as being a sort of a graph which has hills and valleys, and we can imagine starting at a high level, and we have to initially move across a hump, which is the increase in potential energy before we can get down to the valley on the other side, which is the reduction, the lower level of potential energy. So H2O and CO2 have lower levels of potential energy

than things like hydrocarbons and di-oxygen molecules. That's just two oxygens bonded together, which is what we have in the air. However, in order to break up these initial molecules, especially the di-oxygen molecules, you need to have a decent amount of energy to overcome this initial energy bump. catalysts to remember are molecules or substances that reduce this amount of initial energy that you need to get the reaction started.

But for this to happen, you need energy and that comes in the form of basically high temperatures or, in other words, fire. So that's why it's difficult to start to fire, because you've got to get the temperatures high enough. In order to overcome this initial amount of energy that is keeping the dihydrogen molecules and also the hydrocarbon molecules together before you can move down into that valley where of lower potential energy when the oxygen is able to actually pull the electrons away from the carbon and hydrogen, therefore actually formed the carbon dioxide and the water molecules. Some other interesting points on fire. The visible part of fire, the part that we actually see is called the flame.

I think you probably knew that. And what it actually is is essentially just hot gas or possibly a plasma. There's some dispute about that and it might vary from case to case. Remember a plasma is just a gas that's been ionites basically, where there's so much energy that the atoms have lost their outer electrons. So basically what you're observing when you see a flame is a hot gas version of whatever is being burned. And one of the reasons it's visible for a couple of reasons. One is because it's so high temperature, so it glows and radiates energy incandescently, according to the black body spectrum.

I don't think we've talked about that, but I'll talk about that in a future physics episode. So that's one reason it glows, which basically is so hot and hot things glow. And a second reason is because of photons being excited as part of the chemical reaction. So that's why different substances have different colours when they're burned, because the energy gaps between the outer and inner electron shells are different for different atoms, and therefore the amount of energy that is emitted and absorbed when electrons jump between those shells is different for different atoms, and therefore the energy of the photon that is emitted when the electrons jump between these shells is different for different atoms.

Because remember, the photon has to carry away the exact right amount of energy that is lost or gained when the electron jumps between these different shells. So the difference in energy levels between the different shells of different atoms is why we see different frequency photons being emitted from different elements when they are burned and different frequency photons carry different amounts of energy,

but to us we also see them with different colours. That's a little bit of an aside though. So don't worry too much about that if you didn't fully understand everything I said there because a lot of it does relate back to previous physics and chemistry episodes that we've looked at. Also, one other thing that I wanted to point out is, which we'll probably take up in more detail in a future episode, which I'd like to do about movie physics, or lack thereof, really, is how in movies, well, basically everything explodes, especially cars when they crash into things or planes when they crash into things,

or just people sometimes apparently. Everything explodes when they hit things or when you shoot them or whatever. Now, what is an explosion? It's basically just a very rapid occurrence of combustion, where you don't have a flame in the same way that it sort of burns slowly. It just all happens at once, basically. So it's just a big, really fast redox reaction, releasing a lot of energy. But as we just found out,

combustion involves, is a reaction that involves oxygen atoms in the air, essentially taking away the electrons from basically high. hydrocarbons or similar molecules. So for this to happen, obviously, there needs to be enough oxygen atoms around, or in other words there has to be enough air. And you need quite a bit of air, because remember, air is a gas, whereas hydrocarbons that we burn are generally either liquid or a solid, and solid and liquids are much, much denser than gases.

So in order to get enough oxygen atoms to burn a given amount of carbon, or hydrocarbon, you need a heck of a lot of air. And so, in practice, it's not so easy to get things like cars, or even. even petrol tankers to explode. If you drop a lit match into a barrel of oil, it'll just go out. Crash cars very rarely explode, certainly straight away. Even a crashed fuel tanker probably wouldn't explode, certainly straight away. It would likely catch fire if the tank was breached, but explosions are much less probable. So, explosions in movies are way overdone, which you probably knew already anyway. Okay, that's enough on combustion. First example of a redox reaction.

Second example we're going to look at is rust, which is another well-known phenolon. Rust is a general term that describes iron oxides. These are chemical compounds composed of iron and oxygen, you may have guessed. Iron oxides are very common in the earth's crust. Things like iron ore is an iron oxide. And in fact, most of the things we think of as rocks, or at least the minerals that go into rocks, so like silicate, for example, basically silicon oxide or iron oxide and various other metal oxides, are basically just metals that have really.

reacted with oxygen to form metal oxides. And rust is a particular subset of that, although some people use the word rust a bit more generally to refer to metal oxides, but correctly it just refers to iron oxides. Now, iron oxides occur effectively because, again, the oxygen is being reduced. It's pulling electrons away from the iron, and therefore it is reducing its oxidation number and increasing the oxidation number of the iron. And this happens essentially because it allows both atoms to reach a lower energy state. For this type of reaction, though, there isn't really much of an initial barrier of energy that needs to be overcome.

Generally, the reactions just take a little bit of time to occur. We all know that rust is quite brittle and sort of just crumbling in fragments under the touch. So what tends to happen is that the outmost layer of a given mass of iron rusts first, and then it crumbles away and falls or blows off, and then the layer just underneath that oxidizes, and then it keeps going until the whole thing's crumbled away and is completely rusted through. You can avoid this by, well, there are many means of trying to protect metal structure, iron structures from rusting. One of them is to cover them with a layer of another metal which doesn't oxidize so easily. Another is what's called a sacrificial anode, which effectively puts a bunch

of another metal which oxidizes more easily than the iron, and so it's called sacrificial because that lump of metal, it's often like buried in the ground or something, is pretty referentially rusted away and therefore protects the main iron substance. Of course, that needs to be replaced periodically because the sacrificial anode is slowly being corroded away. Many other different methods, too. Water tends to accelerate the rate of rust. So again, just to understand exactly how this is a redox reaction, iron atoms lose electrons, so they're oxidized, to the oxygens in the air, which are reduced, forming iron oxide, which is a combination of the iron ions, that are left behind dissolved in the water,

and the essentially oxygen or hydroxide ions that are being created and potentially dissolved in water as well. So these deposit on the surface of the metal and combine with the iron ions and what we get is essentially iron oxide. Rust or more generally, the oxidation of metals is essentially the main reason, or one of the big reasons why you generally can't find metals in their native form that is in the pure form on the earth unless they've been processed by humans. Because basically, if you just had a lump of iron or most other metals, it would over time be oxidized by oxygen in the atmosphere and turn into rust. or the equivalent of rust. Effectively, it would turn into a metal oxide, which is generally what we refer to as an ore. It's just the metal oxide, and that needs to be purified,

or the reaction essentially reversed in order to get the metal separated from the oxygen. And that's why alls need to be treated. It's also why it's not that easy, because the process of rust is favorable in that it releases energy. It allows the metal. and the oxygen to enter a lower energy state. So that means in order to undo it, you're going to have to introduce energy. So it costs energy to do.

You have to generally heat it up to high temperatures and so on. So it's not that easy to separate metals out from their oxides. One other interesting fact that this concept of oxidation and also electro-negativity helps us understand is why some metals, like gold being the most prominent example, don't rust and don't really react at all. They just sort of sit around for thousands of years without tarnishing or reacting or rusting or anything

And one of the big reasons for this, particularly in the case of gold, is because gold has a very high electro-negativity. Now, it's not as high as, say, oxygen and other non-metals like that, but for a metal, gold has a very high electro-negativity, so it's relatively resistant to having its electrons stolen, in a sense, by oxygen in the atmosphere, or more correctly, it's much less likely that gold is going to undergo oxidation because its electronegativity is much higher

and therefore it's able to hold onto its electrons in relation to oxygen more readily. Okay, so that's our discussion of rust. Our second example of a redox reaction. Now we're going to take our third example of a redox reaction which are batteries and also fuel cells that kind of go together because they're very similar. We'll mostly talk about batteries.

A battery is simply a device. that converts chemical energy into electrical energy via redox reactions. So, the best way of understanding fundamentally what's going on in a battery is to imagine two beakers of solutions sitting next to each other with electrodes of different metals stuck into each of the solutions. And these electrodes, which is essentially thinking of them as rods of metal, are connected by a wire. In order to complete the circuit that we're sort of conceptually building here,

we need to put what is called an electrolyte bridge or a salt bridge that connects the two solutions to each other directly, because remember to have a circuit, we need to have a sort of a circular path for the current to flow. So we've got one half of that, which is the wire connecting the electrodes. The second half is the salt bridge or the electrolytic bridge. There's many ways of doing that, but sort of the common laboratory, simple high school laboratory ways just to use a piece of paper that's been soaked in water, basically. or soaked in one of the electrolytic solutions. All we need is for the ions to be able to travel along this bridge from one solution to the other,

so it doesn't terribly matter what form it takes in, but we need to prevent the two solutions from mixing directly. So we've got two beakers sitting next to each other with electrodes stuck in them, a salt bridge connecting the solutions, and wire connecting the two electrodes. So that's going to be our battery, or an electrochemical cell, as we call it. Real batteries that we actually use are obviously not constructed this way. But conceptually they're constructed in the same way. So physically they're put together differently.

The electrodes are separated in different ways, and the electrolytic solutions are separating different ways, and salt bridges are done differently. But fundamentally, what's going on is the same thing. So we use this simplistic electrochemical cell to understand what's happening. So we've said that a battery is a device that converts chemical energy into electrical energy using a redox reaction. Now let's think for a moment about what electrical energy is.

We haven't done an episode on this yet, but that's coming up. So stay tuned. But electrical energy is essentially, for our purposes here, we'll think of it as the flow of electrons or the flow of electric charge from one place to another. That's what we use for electric energy, or that's what generates an electric current. Well, I mean, that's what an electric current is. It's just a flow of charge, of electrical charge, particularly electrons. Okay, so to get electrical energy, all we need is to get electrons to move around.

But think about what redox reaction is. It's when electrons move from one atom to another. So if we can harness the energy from these electrons moving around, We've essentially created, or not created, but transformed chemical energy into electrical energy, and we can tap that energy by essentially tapping into the energy released by electrons as they move from one atom to the other. So that's where batteries fundamentally get the energy from. Remember that, for example, in the case of a combustion reaction, the reactants have a relatively high potential energy, the hydrocarbons and the dioxygen molecules,

the products have a relatively low potential energy, that is the carbon dioxide and the water. That means that energy is lost in the process of the reaction, or energy is emitted in the process of the reaction. In a battery, we can harness that release in energy in the form of electrical current, basically, which we can use to do work and other useful stuff, like, for example, record podcasts. Okay, so that's where the energy is coming from, but let's look into a little bit more detail about how the process occurs and how it is a redox reaction. Now, I've referred to the two big cases containing electrolytic solutions. I haven't defined that term yet, so I'll do that now.

Remember, a solution is just one substance dissolved inside another, and refer back to our episode on solutions and mixtures, if you're a bit hazy about that. But the distinction here that we're drawing is between an electrolytic solution and a non-electrolytic solution. An electrolytic solution is simply one in which the dissolved molecules are electrically charged, that is their ions. In a non-electrolytic solution, and sugar in water is an example,

the molecules are neutral. And so they can't conduct electricity because remember, electricity is the flow of ions or charged particles. So electro-olytic solutions are necessary for the battery to work, so otherwise we're not going to get a flow of charge. Precisely what electrolytic solution it is doesn't matter too much. I mean, it will for practical

applications, because some will be cheaper any other than others and so on, but conceptually it doesn't really matter. It doesn't need to be an electro-elytic solution. Now, I also mentioned the two electrodes, one of each is stuck inside each of the electrolytic solutions. What's crucial is that these electrodes be made of metals that have differing electronegativities, sufficiently differing so we can get an interesting response. So now that we've got our setup, let's explain what's happening. Remember,

we've got two electrodes and two beakers. We'll call the first electrode the zinc electrode and the second electrode, the copper electrode, just to simplify things a bit. But again, it's not restricted to zinc and copper, we're just using this as an example. And it's also important that in the electrolytic solution, At least for our example, again, to simplify things, we'll assume that one electrolytic solution has zinc ions in it, and the other one has mostly zinc ions and mostly copper ions, and those are in their respective beakers. So in other words, one beaker has zinc ions in the solution, and some other things as well, but that's not crucial. And it also has zinc electrode, and the other one has copper electrode and copper ions. And both of those ions are positively charged.

That's just how, because metals tend to lose electrons when they form ions. Now, each of these beakers with the electrode and the electrolytic solution is called a half-cell, because the whole thing is called an electrolytic cell. It's also referred to as a galvanic cell or a voltaic cell, you may have heard of referred to. The whole thing is a cell, so each beaker with its electrode is referred to as a half-cell. In each half cell, half of the redoxiraction occurs. In other words, in one, oxidation occurs, that's our zinc beaker, and the other one, reduction occurs, that is our cut.

copper electrode. What happens is copper has a higher electro-negativity than zinc, so it has more pulling power for electrons and zinc. So what happens is that the zinc atoms in the metal atus are oxidized, that is, they lose electrons, and the remaining positively charged zinc are disassociates and goes into the electro-electric solution. The electrons then move along the wire that connects the two electrodes to the copper beaker. And there, what happens is that the copper ions in solution, so not the ones that are actually on the electrode, but the ones that are in the solution,

the positively charged copper ions, they are reduced. That is, they gain the electrons. They pick up the electrons that have been released by the zinc. And now, if a copper ion picks up two electrons, it will be neutral, and so it will then precipitate out of the solution and onto, generally onto the...

copper electrode. And this then continues. Now, this is the key reason we need the salt bridge, because if that was all that happened, what we'd have is a buildup of charge in the, with electrons being continually deposited in the copper electrode and taken away from the zinc electrode, the copper electrode would become negatively charged and the zinc weren't positively charged, and that would, the potential difference thereby established would offset the redox effect that we're talking about and the current flow would stop. So in order for that not to occur, what has to happen is that the, in effect, a circuit needs to be completed, and we need to have a flow of charge back into the zinc electrode. And that occurs

via the salt bridge that we were talking about. That allows ions to travel from one side to the other, thereby balancing out the flow of charge that has occurred as a result of the electrons moving. And these charges that are flowing across the salt bridge could be positive or negative, it doesn't matter because, you know, they'll just flow in the correct direction and offset the flow of electrons. what's happening here is that the zinc electrode is being corroded away, although Corro is not quite the right word, because it's not a

reaction with oxygen, but it's being effectively corroded away as the zinc atoms are giving up their electrons and disassociating into the electrolytic solution as zinc ions. The electrons are going over into the copper electrode combining with the copper ions in the electrolytic solution

and depositing on the surface of the copper electrode. So the copper electrode is growing. There's actually physically more metal there than was before, and the zinc electrode is shrinking. And that's happening because the copper is, the copper atoms, sorry, the copper ions that are in the solution are being reduced. They're picking up the electrons from the zinc. Oxidation occurs at the zinc and reduction occurs at the copper. And remember the entire reason that this occurs in the first place is because electrons are seeking out their lowest energy state and high electronegativity, copper is able to, in a sense offer or provide

the electrons a lower energy state, which they can settle into. And so this entire process releases energy, which we can tap in the form of the electric current flowing from one half cell to the other. And that's effectively how a battery works. Now, you can use many different, as I said before, forms of metals for electrodes and forms of electrolytic solution. There are ways that you can have voltaic cells that are not, the non-actually liquid cells, that you use solid as the electrolytes, and that's what a lot of batteries are, how a lot of batteries are constructed nowadays, because you don't have the risk of leakage and so on. But fundamentally, it's the same principle that's happening. It's a redox reaction that allows chemical energy to be converted into a

electrical energy. I'm going to skip fuel cells because they work fairly similar to a battery except rather than using metal electrodes which are corroding and depositing and salt bridges and so on. It basically just uses oxygen and hydrogen which reacted with each other to produce water and we know that's essentially just a form of a combustion reaction. It's just basically burning hydrogen in oxygen to produce carbon dioxide of water. We've already discussed that as being an example of a combustion reaction. Again if we set that up in the right way we can get electrons to move through a wire and therefore we get electric current. So fundamentally, that's all the fuel cell is.

Just a way of combustion occurring in a way that we can tap the energy out of the electrons as they're moving. A final little bit that I want to do is a word on respiration, which is going to be our fourth and final example of a redox reaction. Many important biological processes involve redox reaction. So cellular respiration is an example of this. In this case, glucose, which is not exactly a hydrocarbon because it has oxygen, molecules in there as well, but it's got, it's sort of like a hydrocarbon, because it's got a lot of carbon, hydrogen, and also some oxygen in there. So glucose is oxidized into CO2 and water,

again, in the presence of oxygen. So essentially, respiration is quite similar to a combustion reaction. I think depending on definitions, it is an example of a combustion reaction. And the photosynthesis that occurs in plants is, not exactly the reverse of the redox reaction in cell respiration, but conceptually, essentially is. It's just undoing that reaction in order to produce glucose out of carbon dioxide and water. And of course, that requires the input of energy which comes from sunlight. So just to be clear about what's going on here. The respiration is a bit like a combustion reaction in that, in this case, glucose, in the place of, say, a hydrocarbon is being oxidized and is losing electrons to carbon dioxide and water. So again, we can understand a wide variety

of chemical phenomena from batteries, respiration to rust and combustion as examples of the broad category of oxidation and reduction or redox reactions, which involve the change in the number of electrons held by different atoms. Okay, so that's all we have for today's episode. Hope you enjoyed it. Please visit the Facebook page we have. If you just do a search for The Science of Everything podcast on Facebook, you should find it and join up there where I'll be posting updates and news about upcoming episodes and also additional material, for example, some diagrams that will be helpful in understanding some of the things I talk about in the episodes. Also, if you can jump on iTunes and give me a favourable review, that would be helpful in spreading the word about the podcast and increasing the listener base.

So again, thank you for listening, and I'll talk to you next time.