Understanding Redox Reactions in Chemistry

Okay, so we started talking about redox reactions on our last lecture. This is the lecture for Chem 115 for 2020, October 5th, Monday. At the end of the last lecture, we talked about what is a redox reaction. A redox reaction is a chemical reaction where electrons are transferred from one species to another. And this is one of the most important type of chemical reactions. We discussed briefly how some of the most important chemical reactions to humankind are redox reactions such as photosynthesis, your body breaking down the food. produce energy and carbon dioxide and other chemical reactions such as any kind of battery reactions. are all in the category of redox reaction. So what is a typical redox reaction, and what happens, how does the electrons change hands? Let's see it from a simple example here. Here's the chemical reaction between two free elements, metal sodium and chlorine gas, and making sodium chloride. Okay? And this is a... typical redox reaction. How? Well, any redox reaction can be divided into two halves. One half is one species. losing the electron and turning into something with a higher charge. In this case it's sodium breaking down into sodium ions and give out one electron. So this process is a process of losing electron, and this is called an oxidation. We say sodium in this process is oxidized. This is a process of a chlorine molecule getting electrons. Each chlorine atom would get one, and you ended up producing two chloride ions. This is a process of gaining electrons. This is called a reduction. And together, of course... You have redox. And now, the two parts, what we call the oxidation half reaction and the reduction half reaction, has to happen together. And when you think about it, it does make perfect sense. Mass is supposed to be conserved during a chemical reaction, and when you lose an electron, that electron has got to have somewhere to go. Somebody else needs to pick it up. Otherwise, it's not a full-blown chemical reaction there. And now... The two processes can be easily memorized through a short acronym called OIL-RIG. Oxidation involves loss of electron and reduction. Involve. Evolves. Gain. Of. Electron. Okay, and, now. During any redox reaction, at the end of the day, you can always break it into two hops. There's always going to be a species losing electron. There's always going to be a species gaining electron. Okay, and here's another very important concept of the process. And it sounds a little bit confusing initially, but when you think about it, it does make sense. Okay. The reagent goes through, the reactant goes through oxidation. We call it the reducing. agent and the reactant goes through reduction we call it oxidizing agent agent of course is referring to their identity of being a reactant. Why do we change words here? If something goes through oxidation, why do we call it a reducing agent? If something goes through reduction, why do we call it an oxidizing agent? Remember, oxidation and reduction are describing what happened to this particular reactant. And the reducing and the oxidizing refers to their role in this process, what they did to the other compound. reactant, okay, by going through oxidation, by losing a electron, okay, this reactant is forcing the electron onto another reactant, okay, that other reactant is being reduced, okay, so In our case of the chemical reaction, chlorine is, oh, sorry, that's not reduced. Chlorine is reduced, okay, and sodium is oxidized. Well, so if you think about it, the part that makes perfect sense, if chlorine is reduced, who reduced chlorine? Okay, gotta be sodium. The only other reactant in the reaction. And if sodium is oxidized, who oxidized chlorine? Okay, sodium is oxidized. Who oxidized the sodium? It gotta be the chlorine. So chlorine is the oxidizing agent. Oxygen is, uh, sodium is the reducing agent. It's like someone being robbed. Okay. Who's doing the robbing? Not him. somebody else did it to him. So if you're oxidized, something else oxidized you, and that something that oxidized you is the oxidizing agent. Someone is robbed. Whoever robbed him is the robbing criminal. And now, given a... Chemical reaction given a redox reaction. How do I know? Who is oxidized and who is reduced? Okay How do I tell or you gotta be able to figure out who's losing electron who's gaining electron if you can tell? this particular element, this particular species, lost electron, okay, then you know that it is oxidized, it's your reducing agent, and vice versa. Okay, so how do we determine whether electrons change hand? How many electrons change hand? And now, if it's for just a... ion, or in this situation, from a free element to a positive ion, from a free neutral element to a negative ion. It'll be easy to see, okay? But in the vast majority of the redox reaction that you're going to see, compounds get more complicated, and sometimes you also see molecular compound involved. And in those situations, How do I determine? I don't have no ions. I have molecules. I have covalent bonds. For example, the reaction of Carbon and oxygen gas producing carbon dioxide. Okay. It's something reacting with oxygen. You have to call this oxidation process a... oxidizing reaction of oxidizing process. But you ended up with a molecular compound which doesn't contain any ions. So how do I determine the lose and gain of electron? And to be able to do that, people designed a system that we call oxidation states or it's also called oxidation numbers okay these are numbers assigned to every element in a redox reaction to track electron transfer Okay, now, and we have a set of rules to assign oxidation numbers to elements, and once you get a redox reaction equation, If you're asked to tell who's being oxidized going through oxidation, who's being reduced going through reduction, or who's the oxidizing agent, who's the reducing agent, then you apply the rules on every compound, every atom. element and you get the numbers and from the numbers change if it's increasing okay it's losing electron if it's decreasing it's gaining electron then you have a clear view of who's going through what okay so here are the rules to assign oxidation state. rule number one, free elements. so what is a free element? this is one element by itself with no charge And for all of these guys, oxidation state assigned, simple, zero. Okay, so examples are the sodium that we just talked about, sodium metal, okay, or chlorine gas, okay. oxygen molecules, okay, and phosphorous molecules. This is a special carbon molecule that people developed in the and very hot in the nano material field nowadays. C60. Sixty carbon atoms bonded together in the shape of a soccer ball. And this is called the Buckyball. And all of these guys are free elements. So you can see, as long as you have one element, no other elements combined or bonded to it, then regardless of your subscripts whether you have one neutral atom or you have multiple atoms giving multiple of the same atoms giving you a molecule your oxidation your oxidation state is always assigned as zero all right number two mono atomic ions Okay, and of course this is cations, positive ones, and anions. Okay, their oxidation state assigned is equal to the charge on them. Okay, so by that I mean if you have Na+, oxidation state would be plus 1. If you have Cl-, oxidation state is negative 1. If you have ionic compound Fe2. O3. Then we know how to do the charge. Oxygen is in group 16. Always have the negative 2 charge. And that's 3 negative 2. That's a total of negative 6. And 2 ions. 2 ions are balancing it off. So each ion is plus 3. So ion would have a plus 3 oxidation number. And oxygen would have a... Negative 2 oxidation number. Fluorine. And rule number 3. Fluorine in any compound. Okay, so by that I'm saying fluorine bonded together with some other elements. So you can't say F2. F2 is a free element. It's not a compound. And oxidation. number of F2 for each fluorine is 0, okay? But fluorine in any other compound, when it's bonded with other atoms, okay, oxidation state is always negative 1, okay? Number four, oxygen in compound, okay? The vast majority of the compound, oxygen, would have an oxidation number of negative 2, with one exception, except in peroxides. What is a peroxide? This is hydrogen peroxide. This is the thing that... the household chemical that you could swap on a scraped knee to just kill germs. And in peroxides, oxygen would have the oxidation number of negative one. Okay, outside of peroxide, in every other compound that you could see in Chem 115, oxygen have the oxidation number of negative 2. And number 5 is for hydrogen. Hydrogen in compound. almost always have the oxidation state of plus one, okay? Just like in any acid, when you have hydrogen, and when hydrogen comes off, it assumes a plus one charge as a hydrogen ion, okay? And hydrogen in compound normally assumes a plus one oxidation state, okay? Except... when it acts as a non-metal in metal hydrides, such as NaH. That's an ionic compound. Sodium is always plus one, and that makes your hydrogen negative one because the charge needs to balance off. So in metal hydrides, when hydrogen acts as a nonmetal, hydrogen has the oxidation, and that makes the hydrogen the hydride ion. And then the monoatomic ion rule kicks in, and that gives the hydrogen negative one charge. All right, and so these are the five rules focusing on different elements. what kind of oxidation state they should have and rule number six is a calculation rule okay for elements that we that's not covered in this first five rules okay you could use the calculation rule to figure out a certain element of what oxidation number it has. And the rule goes as the sum of oxidation state of all Atoms in a molecule or a polyatomic ion should equal to the charge. on the group. Now, what do you mean by this? Well, if you see a molecular formula that contains elements that you don't know, or their oxidation number are not defined in are five earlier rules then you can apply the sixth rule and using some of the elements that you already know oxidation numbers such as a monoatomic ion or a fluoride or oxygen or hydrogen, then you can use these and the calculation rule to figure out the unknown element's oxidation number. A very simple example, think about carbon dioxide. Now we just said that carbon reacting with... oxygen-producing carbon dioxide is a redox reaction. Well, who's the oxidizing agent? Who's the reducing agent? I need to know how electrons change hands, and I need to know the oxidation number of these guys. But carbon is not covered in one of these rules because, well, carbon is... Carbon dioxide is a molecular compound. You don't have individual ions in there, and carbon certainly is not fluorine, oxygen, or hydrogen. Okay, so how should I get the oxidation number of carbon element? simple. Oxygen in compound, and this is not a peroxide, so oxygen has negative 2. And the total, so 1 carbon plus 2 oxygen, which is 2 times negative 2. The total of oxidation number of all atoms in a molecule should equal to the charge on the group. What's the charge on the group? The group is a molecule. And what do we know about molecules? Molecules are neutral. Neutral means zero charge on the group. So carbon plus 2 times negative 2, so negative 4, equals zero. What does that make carbon? Oh, okay. Carbon has a plus 4 oxidation number here. And what if I have a polyatomic ion like... Sulfate ion. Okay, I don't know what sulfur is, but I know I have four oxygen, and each oxygen should carry a negative two oxidation state, and that total should be equal to the charge on the group. The charge on the group is the charge on the polyatomic ion, negative two. So sulfur plus negative eight equal to negative two, and that makes sulfur plus six. Okay, what about... Okay. sulfide ion or in this case sulfur plus three times negative two should equal to the charge on the group is still negative two and that makes sulfur a plus four oxidation number okay and I'm using both of them point is an element could assume okay different oxidation numbers when it plays different role in a redox reaction. And now, so going back to that equation that we had earlier, okay? Carbon reacting with oxygen producing carbon dioxide. We said this is a redox reaction, so who's oxidized, who's reduced, who's the oxidizing agent, who's the reducing agent? And once you put the... oxidation number on them, it should be very clear. Carbon is a free element. Oxygen O2 is also a free element. So both have zero oxidation number. And carbon dioxide, we just figured carbon has plus four, oxygen has negative two. So let's see. Carbon went from zero to plus four. What does that mean? When, so for carbon oxidation state increase. Increase means losing electron. Because electrons have negative charge, losing negative charge means gaining positive charge, increase. And also losing electrons means you're going through oxidation. So carbon is oxidized. And of course, in this case, then oxygen is going from zero to negative two. So oxygen is reduced. Oxidation, I'm going to state, going down. And that makes your carbon the reducing agent. And that makes your compound O2 the oxidizing agent. Okay, now even if your chemical formula becomes more complicated and harder to see directly, okay, you can still use this the same way to figure out the different rows that a compound or a polyatomic ion plate in. A chemical reaction, so if you look at SO4 2 negative reacting with NO2 negative producing NO3 negative and SO3 2 negative. Okay, this is a... redox reaction, and you can see sulfate is changing into sulfite, and nitrite is changing into nitrate, okay? One, it looks like one oxygen, okay, actually changed hand, okay? What is the electron loss and gain during this reaction? Once you label everybody, that should be pretty clear. Okay. Sulfur, we just talked about it, is plus 6. Oxygen is negative 2. Sulfur here is plus 4. If you go back to the last page, you would see that. And oxygen, of course, in compound is always negative 2. Nitrogen. What kind of oxidation? state nitrogen has in this nitride ion. Okay, we'll think about it. Nitrogen plus two oxygen, okay, each oxygen is negative two, two negative... 2 oxygen is negative 4, and nitrogen, after added 2 negative 4, make the whole thing negative 1. And that makes your nitrogen here plus 3. And similarly, you could find out that your nitrogen here has plus 4, plus 5. Okay, so sulfur's oxidation number is decreasing. Okay, and by looking at this, you see going from plus 6 to plus 4, you can actually very easily determine... between the two sulfur polyatomic ions, two electrons were gained or loose. plus six to plus four your oxidation state is going down okay going down means gaining negative charge okay so you gained two electron okay and similarly from plus three to plus five okay you're having a higher positive charge and that means losing or lost two electrons. And once you determine that, whoever is losing electrons, so NO2 negative is losing electron, so it's going through oxidation, oxidized, and of course it's the reducing agent. Okay, and your SO2 negative is being reduced. And of course, it's your oxidizing agent. Okay, so being able to look at a chemical formula in a redox reaction, then assign oxidation state on all the elements is a very important step if you want to determine what every chemical compound's role, what every reactant's role in a given redox reaction. So... Couple more examples about assigning oxidation number if you see N2. Okay, what's its oxidation number? Oxidation state, zero. Free element. If you see NH3, ammonia. What's the oxidation state? Well, hydrogen in compound is plus one. This is not a metal hydride, obviously. Nitrogen is not a metal. And this whole thing is a neutral molecule, and you have three hydrogen, each have a plus one oxidation state, and that makes your nitrogen a negative three oxidation state. What about HNO3? Well, we just did this earlier, nitrate. So hydrogen is plus 1, oxygen is negative 2. You have 3 oxygen that make it negative 6. Hydrogen is only plus 1, so you would need the nitrogen to be plus 5. Then you can have a neutral acid molecule. Okay, so... N2O3, oxygen is negative 2, and nitrogen is plus 3. Or you could see things that's even more complicated. K2Cr2O7, let's say. First of all, what kind of compound is this? It's an ionic compound. Okay, obviously, because you have metals and nonmetals. No. You know oxygen in compound has negative 2, and you go, well, if I have 2 unknown and I only know the total equal to 0, I still can't make a judgment here. But think for a second. What kind of... wrote, potassium is playing here. Potassium is your positive, obviously, your positive ion in the compound. Okay? And potassium is from group one. And group one in... in chemical compounds, in ionic compounds, always show plus one charge. And monoatomic ion, potassium ion, has a plus one charge. And if you know that, well, plus one times two, and two times what chromium has, and plus negative two times seven, and that should equal to the total charge on the group, the total... total charge on the group, this is a neutral compound, total charge on the group, obviously, is zero, and that makes your chromium a plus six oxidation number, okay, and now, If you're wondering, well, yeah, I see you following the rules and put oxidation numbers on different elements in a compound. Now, I can get it in an ionic compound that Positive ions are losing electron and negative ions are gaining electron. We already talked about that before. Okay. They obviously have charges. But what about these molecular compounds? Yeah, you put an oxidant. number on it, but what does it mean? For example, carbon dioxide. You said carbon is plus 4, oxygen is negative 2. But didn't you also tell us that molecular compounds are sticking together by covalent bond, which is sharing electrons? Yes, we did. Okay, so what happened? If they're sharing electrons, why do you say one of them is positive and actually positive a lot, positive four, which means lost four electrons. And then you also have oxygen, okay, and negative two, which means... it's gained electron but they're actually sharing right? This is how sharing happens okay this is the electron structure of carbon dioxide Yes, it is sharing electrons with oxygen. And on the other hand, it also shares electrons with oxygen. See, they share. What is this negative 4 and negative 2 plus 4 stuff? One thing I want you to understand, you don't have to memorize this. Okay, I'm not going to ask you to explain why these things are assigned positive or negative oxidation numbers. As long as you can assign them, we're all good. But the explanation here is that sharing can be different. Sharing electrons can be un... Equal. What do you mean? Well, different elements, different atoms, even in a molecule, okay, their ability of attracting electrons are different. That should be obvious because every single one of them have a different nucleus that contains different positive charges, okay? And when the sharing is not equal, What you see for a carbon dioxide molecule is not like what we draw here anymore. What's the reality? The reality is they are sharing four electrons, but those four electrons are going to be so much closer to the oxygen compared to the carbon. And in this case, it really feels like... Okay, carbon originally had four electrons to himself. Okay, out of the eight electrons shared here, four of them belong to carbon. Okay, but oxygen is sharing two pairs with him and really kind of taking the electrons away from carbon. and the other oxygen took another two electrons away from carbon, and that gave carbon plus four oxidation number. It feels like carbon lost four electrons. Now, did it really? I mean, think about it as this. Let's say you and your friend, okay, ponied up some money, and you bought an Xbox. Okay, now each of you paid the same amount, half of the money, but then it turned out that where is this Xbox? It was in your friend's apartment. Well, yeah, on the title half of that Xbox is mine. But when I'm sitting in my apartment, I don't see Xbox. And that feeling should be kind of similar to that you lose the money of half of an Xbox. And for your friend, because it's in his apartment, he's playing it every day. It feels like he gained something. He got in on a really good deal. He paid half the money, but he has that Xbox so much closer to him than to you. So, negative two, negative two. All right? So, one thing that you can do to practice on your... Ability to assign oxidation number, okay, to figure out oxidizing agent and reducing agent is really going through your textbook where the oxidation redox chapter is. Find any chemical... chemical reaction there, write it down, then start labeling, okay? And then try to determine, okay, who's losing electron, who's gaining electron, who's oxidized, who's reduced, who's the oxidizing agent, who's the reducing agent. With some practice, this can come eventually pretty easy, okay? Alright, so this is how we initially handle a redox reaction and figure out what role. each reactant plate in that redox reaction, okay? And the other aspect of redox reaction that we're going to put a significant amount of time on, okay, is balancing redox reaction equations, okay? And why is that important? Well, at this point, you already know how important a balanced chemical equation is. Remember, the whole stoichiometry thing was based on what? Balanced chemical equation. Because balanced chemical equation gave me coefficients, and those coefficients are model ratios, and we can use them to do calculations. But also, when we were... Balancing chemical equations before either the equations we balanced are very simple, okay? Or, if you remember, one of our rules is do not break down your polyatomic ions. Does that apply in a redox reaction? Well, look at this thing we just did there. Originally you have a polyatomic ion of SO4 2 negative. In the end it turned into a polyatomic ion of SO3 2 negative. Do you still have any SO4 around in the product side? No. So the part of balancing equation to focus to not breaking down the polyatomic ions. doesn't really apply to redox reaction okay and this makes the balancing process of redox okay a little bit more complicated than normal okay but does that mean okay we're gonna still trial and error in balancing redox equations and you know Try on things until we find a solution, until we find a set number to put in. Actually, no. Okay. Why? Because we have a systematic way of balancing redox reaction. We actually have quite a few different methods of balancing redox equation. Okay. And the one that you're going to see in my class. The one that I choose to explain in detail, okay, is one that has a very detailed preset procedures, okay? And if you can follow those procedures, you can always get your balancing. Regardless of how complicated your redox reaction is, you can always get the balancing done. Okay, so take a look at this next example. This is a bromine molecule reacting with a uranium 4 plus ion and producing bromide ion and UO2 plus polyatomic ions. How do I balance this redox equation? Now, the first look tells you something that you really can't balance this as it is. Why? On the left side in your reactant, you see bromine. You see uranium. On the right side, you see bromine. You see uranium. But you also see oxygen. I don't even have oxygen on the other hand, on the other side. How am I supposed to balance this? So look at the method that I'm going to introduce here. We're going to call it a half reaction method. Okay, I want to be very clear about this. Half reaction method is not the only method on balancing redox equations. Okay, those of you who's been through. Chem 110, okay, know that in Chem 110, okay, we use a different way, a different method to balance redox equations, okay, at least a lot of the Chem 110 classes that I know, okay, use a different method to balance redox equation, okay, but why do you choose, and if you're really good at any other methods of balancing redox equations, trust me, I have no intention to make you change, okay? Now, but why don't you use other methods? Why are you focusing on this half reaction method? Okay, a couple reasons. Number one is that although this method, the half reaction method, is a little longer than the other method that I know, okay, but this one is very, very organized, okay. There's no, there's always, okay, preset individual steps that you can follow, okay, and if you follow those steps, there's no way that you get balancing wrong. So half reaction method for balancing redox equation is more of a foolproof type of method okay now how do I do this well it's called half reaction method so you got to get your half reactions out so the first step you always go okay is identify half reactions and this is actually very easy to do okay And you can go my bromine is turned into bromide and my U4 plus is turning into UO2 plus ion. Okay, I don't know who's gaining electron, who's losing electron. I'm going to figure it out later on. The first step of identifying half reaction is simply who is turning into what. Oh, okay, it's about time. We'll pick up from here on our Wednesday lecture. We'll finish this balancing process. Right now, it's only on first step, breaking a redox reaction into two half reactions.

And this is one of the most important type of chemical reactions. We discussed briefly how some of the most important chemical reactions to humankind are redox reactions such as photosynthesis, your body breaking down the food. produce energy and carbon dioxide and other chemical reactions such as any kind of battery reactions. are all in the category of redox reaction. So what is a typical redox reaction, and what happens, how does the electrons change hands?

Let's see it from a simple example here. Here's the chemical reaction between two free elements, metal sodium and chlorine gas, and making sodium chloride. Okay? And this is a...

typical redox reaction. How? Well, any redox reaction can be divided into two halves.

One half is one species. losing the electron and turning into something with a higher charge. In this case it's sodium breaking down into sodium ions and give out one electron.

So this process is a process of losing electron, and this is called an oxidation. We say sodium in this process is oxidized. This is a process of a chlorine molecule getting electrons. Each chlorine atom would get one, and you ended up producing two chloride ions. This is a process of gaining electrons.

This is called a reduction. And together, of course... You have redox.

And now, the two parts, what we call the oxidation half reaction and the reduction half reaction, has to happen together. And when you think about it, it does make perfect sense. Mass is supposed to be conserved during a chemical reaction, and when you lose an electron, that electron has got to have somewhere to go.

Somebody else needs to pick it up. Otherwise, it's not a full-blown chemical reaction there. And now... The two processes can be easily memorized through a short acronym called OIL-RIG.

Oxidation involves loss of electron and reduction. Involve. Evolves. Gain.

Of. Electron. Okay, and, now.

During any redox reaction, at the end of the day, you can always break it into two hops. There's always going to be a species losing electron. There's always going to be a species gaining electron. Okay, and here's another very important concept of the process. And it sounds a little bit confusing initially, but when you think about it, it does make sense.

Okay. The reagent goes through, the reactant goes through oxidation. We call it the reducing. agent and the reactant goes through reduction we call it oxidizing agent agent of course is referring to their identity of being a reactant. Why do we change words here?

If something goes through oxidation, why do we call it a reducing agent? If something goes through reduction, why do we call it an oxidizing agent? Remember, oxidation and reduction are describing what happened to this particular reactant.

And the reducing and the oxidizing refers to their role in this process, what they did to the other compound. reactant, okay, by going through oxidation, by losing a electron, okay, this reactant is forcing the electron onto another reactant, okay, that other reactant is being reduced, okay, so In our case of the chemical reaction, chlorine is, oh, sorry, that's not reduced. Chlorine is reduced, okay, and sodium is oxidized.

Well, so if you think about it, the part that makes perfect sense, if chlorine is reduced, who reduced chlorine? Okay, gotta be sodium. The only other reactant in the reaction.

And if sodium is oxidized, who oxidized chlorine? Okay, sodium is oxidized. Who oxidized the sodium?

It gotta be the chlorine. So chlorine is the oxidizing agent. Oxygen is, uh, sodium is the reducing agent. It's like someone being robbed.

Okay. Who's doing the robbing? Not him. somebody else did it to him. So if you're oxidized, something else oxidized you, and that something that oxidized you is the oxidizing agent.

Someone is robbed. Whoever robbed him is the robbing criminal. And now, given a...

Chemical reaction given a redox reaction. How do I know? Who is oxidized and who is reduced? Okay How do I tell or you gotta be able to figure out who's losing electron who's gaining electron if you can tell?

this particular element, this particular species, lost electron, okay, then you know that it is oxidized, it's your reducing agent, and vice versa. Okay, so how do we determine whether electrons change hand? How many electrons change hand?

And now, if it's for just a... ion, or in this situation, from a free element to a positive ion, from a free neutral element to a negative ion. It'll be easy to see, okay?

But in the vast majority of the redox reaction that you're going to see, compounds get more complicated, and sometimes you also see molecular compound involved. And in those situations, How do I determine? I don't have no ions. I have molecules.

I have covalent bonds. For example, the reaction of Carbon and oxygen gas producing carbon dioxide. Okay. It's something reacting with oxygen.

You have to call this oxidation process a... oxidizing reaction of oxidizing process. But you ended up with a molecular compound which doesn't contain any ions.

So how do I determine the lose and gain of electron? And to be able to do that, people designed a system that we call oxidation states or it's also called oxidation numbers okay these are numbers assigned to every element in a redox reaction to track electron transfer Okay, now, and we have a set of rules to assign oxidation numbers to elements, and once you get a redox reaction equation, If you're asked to tell who's being oxidized going through oxidation, who's being reduced going through reduction, or who's the oxidizing agent, who's the reducing agent, then you apply the rules on every compound, every atom. element and you get the numbers and from the numbers change if it's increasing okay it's losing electron if it's decreasing it's gaining electron then you have a clear view of who's going through what okay so here are the rules to assign oxidation state. rule number one, free elements. so what is a free element?

this is one element by itself with no charge And for all of these guys, oxidation state assigned, simple, zero. Okay, so examples are the sodium that we just talked about, sodium metal, okay, or chlorine gas, okay. oxygen molecules, okay, and phosphorous molecules. This is a special carbon molecule that people developed in the and very hot in the nano material field nowadays. C60.

Sixty carbon atoms bonded together in the shape of a soccer ball. And this is called the Buckyball. And all of these guys are free elements.

So you can see, as long as you have one element, no other elements combined or bonded to it, then regardless of your subscripts whether you have one neutral atom or you have multiple atoms giving multiple of the same atoms giving you a molecule your oxidation your oxidation state is always assigned as zero all right number two mono atomic ions Okay, and of course this is cations, positive ones, and anions. Okay, their oxidation state assigned is equal to the charge on them. Okay, so by that I mean if you have Na+, oxidation state would be plus 1. If you have Cl-, oxidation state is negative 1. If you have ionic compound Fe2.

O3. Then we know how to do the charge. Oxygen is in group 16. Always have the negative 2 charge.

And that's 3 negative 2. That's a total of negative 6. And 2 ions. 2 ions are balancing it off. So each ion is plus 3. So ion would have a plus 3 oxidation number.

And oxygen would have a... Negative 2 oxidation number. Fluorine. And rule number 3. Fluorine in any compound. Okay, so by that I'm saying fluorine bonded together with some other elements.

So you can't say F2. F2 is a free element. It's not a compound.

And oxidation. number of F2 for each fluorine is 0, okay? But fluorine in any other compound, when it's bonded with other atoms, okay, oxidation state is always negative 1, okay? Number four, oxygen in compound, okay?

The vast majority of the compound, oxygen, would have an oxidation number of negative 2, with one exception, except in peroxides. What is a peroxide? This is hydrogen peroxide. This is the thing that... the household chemical that you could swap on a scraped knee to just kill germs.

And in peroxides, oxygen would have the oxidation number of negative one. Okay, outside of peroxide, in every other compound that you could see in Chem 115, oxygen have the oxidation number of negative 2. And number 5 is for hydrogen. Hydrogen in compound.

almost always have the oxidation state of plus one, okay? Just like in any acid, when you have hydrogen, and when hydrogen comes off, it assumes a plus one charge as a hydrogen ion, okay? And hydrogen in compound normally assumes a plus one oxidation state, okay?

Except... when it acts as a non-metal in metal hydrides, such as NaH. That's an ionic compound.

Sodium is always plus one, and that makes your hydrogen negative one because the charge needs to balance off. So in metal hydrides, when hydrogen acts as a nonmetal, hydrogen has the oxidation, and that makes the hydrogen the hydride ion. And then the monoatomic ion rule kicks in, and that gives the hydrogen negative one charge.

All right, and so these are the five rules focusing on different elements. what kind of oxidation state they should have and rule number six is a calculation rule okay for elements that we that's not covered in this first five rules okay you could use the calculation rule to figure out a certain element of what oxidation number it has. And the rule goes as the sum of oxidation state of all Atoms in a molecule or a polyatomic ion should equal to the charge. on the group. Now, what do you mean by this?

Well, if you see a molecular formula that contains elements that you don't know, or their oxidation number are not defined in are five earlier rules then you can apply the sixth rule and using some of the elements that you already know oxidation numbers such as a monoatomic ion or a fluoride or oxygen or hydrogen, then you can use these and the calculation rule to figure out the unknown element's oxidation number. A very simple example, think about carbon dioxide. Now we just said that carbon reacting with...

oxygen-producing carbon dioxide is a redox reaction. Well, who's the oxidizing agent? Who's the reducing agent? I need to know how electrons change hands, and I need to know the oxidation number of these guys.

But carbon is not covered in one of these rules because, well, carbon is... Carbon dioxide is a molecular compound. You don't have individual ions in there, and carbon certainly is not fluorine, oxygen, or hydrogen. Okay, so how should I get the oxidation number of carbon element?

simple. Oxygen in compound, and this is not a peroxide, so oxygen has negative 2. And the total, so 1 carbon plus 2 oxygen, which is 2 times negative 2. The total of oxidation number of all atoms in a molecule should equal to the charge on the group. What's the charge on the group?

The group is a molecule. And what do we know about molecules? Molecules are neutral. Neutral means zero charge on the group.

So carbon plus 2 times negative 2, so negative 4, equals zero. What does that make carbon? Oh, okay. Carbon has a plus 4 oxidation number here. And what if I have a polyatomic ion like...

Sulfate ion. Okay, I don't know what sulfur is, but I know I have four oxygen, and each oxygen should carry a negative two oxidation state, and that total should be equal to the charge on the group. The charge on the group is the charge on the polyatomic ion, negative two.

So sulfur plus negative eight equal to negative two, and that makes sulfur plus six. Okay, what about... Okay.

sulfide ion or in this case sulfur plus three times negative two should equal to the charge on the group is still negative two and that makes sulfur a plus four oxidation number okay and I'm using both of them point is an element could assume okay different oxidation numbers when it plays different role in a redox reaction. And now, so going back to that equation that we had earlier, okay? Carbon reacting with oxygen producing carbon dioxide.

We said this is a redox reaction, so who's oxidized, who's reduced, who's the oxidizing agent, who's the reducing agent? And once you put the... oxidation number on them, it should be very clear. Carbon is a free element. Oxygen O2 is also a free element.

So both have zero oxidation number. And carbon dioxide, we just figured carbon has plus four, oxygen has negative two. So let's see.

Carbon went from zero to plus four. What does that mean? When, so for carbon oxidation state increase.

Increase means losing electron. Because electrons have negative charge, losing negative charge means gaining positive charge, increase. And also losing electrons means you're going through oxidation.

So carbon is oxidized. And of course, in this case, then oxygen is going from zero to negative two. So oxygen is reduced.

Oxidation, I'm going to state, going down. And that makes your carbon the reducing agent. And that makes your compound O2 the oxidizing agent. Okay, now even if your chemical formula becomes more complicated and harder to see directly, okay, you can still use this the same way to figure out the different rows that a compound or a polyatomic ion plate in.

A chemical reaction, so if you look at SO4 2 negative reacting with NO2 negative producing NO3 negative and SO3 2 negative. Okay, this is a... redox reaction, and you can see sulfate is changing into sulfite, and nitrite is changing into nitrate, okay? One, it looks like one oxygen, okay, actually changed hand, okay?

What is the electron loss and gain during this reaction? Once you label everybody, that should be pretty clear. Okay. Sulfur, we just talked about it, is plus 6. Oxygen is negative 2. Sulfur here is plus 4. If you go back to the last page, you would see that.

And oxygen, of course, in compound is always negative 2. Nitrogen. What kind of oxidation? state nitrogen has in this nitride ion.

Okay, we'll think about it. Nitrogen plus two oxygen, okay, each oxygen is negative two, two negative... 2 oxygen is negative 4, and nitrogen, after added 2 negative 4, make the whole thing negative 1. And that makes your nitrogen here plus 3. And similarly, you could find out that your nitrogen here has plus 4, plus 5. Okay, so sulfur's oxidation number is decreasing. Okay, and by looking at this, you see going from plus 6 to plus 4, you can actually very easily determine...

between the two sulfur polyatomic ions, two electrons were gained or loose. plus six to plus four your oxidation state is going down okay going down means gaining negative charge okay so you gained two electron okay and similarly from plus three to plus five okay you're having a higher positive charge and that means losing or lost two electrons. And once you determine that, whoever is losing electrons, so NO2 negative is losing electron, so it's going through oxidation, oxidized, and of course it's the reducing agent. Okay, and your SO2 negative is being reduced.

And of course, it's your oxidizing agent. Okay, so being able to look at a chemical formula in a redox reaction, then assign oxidation state on all the elements is a very important step if you want to determine what every chemical compound's role, what every reactant's role in a given redox reaction. So... Couple more examples about assigning oxidation number if you see N2.

Okay, what's its oxidation number? Oxidation state, zero. Free element. If you see NH3, ammonia. What's the oxidation state?

Well, hydrogen in compound is plus one. This is not a metal hydride, obviously. Nitrogen is not a metal.

And this whole thing is a neutral molecule, and you have three hydrogen, each have a plus one oxidation state, and that makes your nitrogen a negative three oxidation state. What about HNO3? Well, we just did this earlier, nitrate.

So hydrogen is plus 1, oxygen is negative 2. You have 3 oxygen that make it negative 6. Hydrogen is only plus 1, so you would need the nitrogen to be plus 5. Then you can have a neutral acid molecule. Okay, so... N2O3, oxygen is negative 2, and nitrogen is plus 3. Or you could see things that's even more complicated.

K2Cr2O7, let's say. First of all, what kind of compound is this? It's an ionic compound.

Okay, obviously, because you have metals and nonmetals. No. You know oxygen in compound has negative 2, and you go, well, if I have 2 unknown and I only know the total equal to 0, I still can't make a judgment here.

But think for a second. What kind of... wrote, potassium is playing here.

Potassium is your positive, obviously, your positive ion in the compound. Okay? And potassium is from group one. And group one in... in chemical compounds, in ionic compounds, always show plus one charge.

And monoatomic ion, potassium ion, has a plus one charge. And if you know that, well, plus one times two, and two times what chromium has, and plus negative two times seven, and that should equal to the total charge on the group, the total... total charge on the group, this is a neutral compound, total charge on the group, obviously, is zero, and that makes your chromium a plus six oxidation number, okay, and now, If you're wondering, well, yeah, I see you following the rules and put oxidation numbers on different elements in a compound.

Now, I can get it in an ionic compound that Positive ions are losing electron and negative ions are gaining electron. We already talked about that before. Okay. They obviously have charges. But what about these molecular compounds?

Yeah, you put an oxidant. number on it, but what does it mean? For example, carbon dioxide.

You said carbon is plus 4, oxygen is negative 2. But didn't you also tell us that molecular compounds are sticking together by covalent bond, which is sharing electrons? Yes, we did. Okay, so what happened?

If they're sharing electrons, why do you say one of them is positive and actually positive a lot, positive four, which means lost four electrons. And then you also have oxygen, okay, and negative two, which means... it's gained electron but they're actually sharing right?

This is how sharing happens okay this is the electron structure of carbon dioxide Yes, it is sharing electrons with oxygen. And on the other hand, it also shares electrons with oxygen. See, they share.

What is this negative 4 and negative 2 plus 4 stuff? One thing I want you to understand, you don't have to memorize this. Okay, I'm not going to ask you to explain why these things are assigned positive or negative oxidation numbers.

As long as you can assign them, we're all good. But the explanation here is that sharing can be different. Sharing electrons can be un... Equal. What do you mean?

Well, different elements, different atoms, even in a molecule, okay, their ability of attracting electrons are different. That should be obvious because every single one of them have a different nucleus that contains different positive charges, okay? And when the sharing is not equal, What you see for a carbon dioxide molecule is not like what we draw here anymore.

What's the reality? The reality is they are sharing four electrons, but those four electrons are going to be so much closer to the oxygen compared to the carbon. And in this case, it really feels like...

Okay, carbon originally had four electrons to himself. Okay, out of the eight electrons shared here, four of them belong to carbon. Okay, but oxygen is sharing two pairs with him and really kind of taking the electrons away from carbon.

and the other oxygen took another two electrons away from carbon, and that gave carbon plus four oxidation number. It feels like carbon lost four electrons. Now, did it really? I mean, think about it as this.

Let's say you and your friend, okay, ponied up some money, and you bought an Xbox. Okay, now each of you paid the same amount, half of the money, but then it turned out that where is this Xbox? It was in your friend's apartment. Well, yeah, on the title half of that Xbox is mine.

But when I'm sitting in my apartment, I don't see Xbox. And that feeling should be kind of similar to that you lose the money of half of an Xbox. And for your friend, because it's in his apartment, he's playing it every day.

It feels like he gained something. He got in on a really good deal. He paid half the money, but he has that Xbox so much closer to him than to you.

So, negative two, negative two. All right? So, one thing that you can do to practice on your...

Ability to assign oxidation number, okay, to figure out oxidizing agent and reducing agent is really going through your textbook where the oxidation redox chapter is. Find any chemical... chemical reaction there, write it down, then start labeling, okay? And then try to determine, okay, who's losing electron, who's gaining electron, who's oxidized, who's reduced, who's the oxidizing agent, who's the reducing agent. With some practice, this can come eventually pretty easy, okay?

Alright, so this is how we initially handle a redox reaction and figure out what role. each reactant plate in that redox reaction, okay? And the other aspect of redox reaction that we're going to put a significant amount of time on, okay, is balancing redox reaction equations, okay? And why is that important?

Well, at this point, you already know how important a balanced chemical equation is. Remember, the whole stoichiometry thing was based on what? Balanced chemical equation. Because balanced chemical equation gave me coefficients, and those coefficients are model ratios, and we can use them to do calculations. But also, when we were...

Balancing chemical equations before either the equations we balanced are very simple, okay? Or, if you remember, one of our rules is do not break down your polyatomic ions. Does that apply in a redox reaction?

Well, look at this thing we just did there. Originally you have a polyatomic ion of SO4 2 negative. In the end it turned into a polyatomic ion of SO3 2 negative.

Do you still have any SO4 around in the product side? No. So the part of balancing equation to focus to not breaking down the polyatomic ions. doesn't really apply to redox reaction okay and this makes the balancing process of redox okay a little bit more complicated than normal okay but does that mean okay we're gonna still trial and error in balancing redox equations and you know Try on things until we find a solution, until we find a set number to put in.

Actually, no. Okay. Why? Because we have a systematic way of balancing redox reaction. We actually have quite a few different methods of balancing redox equation.

Okay. And the one that you're going to see in my class. The one that I choose to explain in detail, okay, is one that has a very detailed preset procedures, okay? And if you can follow those procedures, you can always get your balancing. Regardless of how complicated your redox reaction is, you can always get the balancing done.

Okay, so take a look at this next example. This is a bromine molecule reacting with a uranium 4 plus ion and producing bromide ion and UO2 plus polyatomic ions. How do I balance this redox equation? Now, the first look tells you something that you really can't balance this as it is. Why?

On the left side in your reactant, you see bromine. You see uranium. On the right side, you see bromine.

You see uranium. But you also see oxygen. I don't even have oxygen on the other hand, on the other side.

How am I supposed to balance this? So look at the method that I'm going to introduce here. We're going to call it a half reaction method.

Okay, I want to be very clear about this. Half reaction method is not the only method on balancing redox equations. Okay, those of you who's been through. Chem 110, okay, know that in Chem 110, okay, we use a different way, a different method to balance redox equations, okay, at least a lot of the Chem 110 classes that I know, okay, use a different method to balance redox equation, okay, but why do you choose, and if you're really good at any other methods of balancing redox equations, trust me, I have no intention to make you change, okay?

Now, but why don't you use other methods? Why are you focusing on this half reaction method? Okay, a couple reasons. Number one is that although this method, the half reaction method, is a little longer than the other method that I know, okay, but this one is very, very organized, okay.

There's no, there's always, okay, preset individual steps that you can follow, okay, and if you follow those steps, there's no way that you get balancing wrong. So half reaction method for balancing redox equation is more of a foolproof type of method okay now how do I do this well it's called half reaction method so you got to get your half reactions out so the first step you always go okay is identify half reactions and this is actually very easy to do okay And you can go my bromine is turned into bromide and my U4 plus is turning into UO2 plus ion. Okay, I don't know who's gaining electron, who's losing electron.

I'm going to figure it out later on. The first step of identifying half reaction is simply who is turning into what. Oh, okay, it's about time. We'll pick up from here on our Wednesday lecture.

We'll finish this balancing process. Right now, it's only on first step, breaking a redox reaction into two half reactions.

Transcript for:Understanding Redox Reactions in Chemistry

Transcript for:
Understanding Redox Reactions in Chemistry