hello your twelves this is mr. Lim again and we are doing our seventh video on redox for standard reduction potential table it's gonna be a long video so let's start ok so today we're gonna be learning about the standard reduction potential table we are going to learn how to use it to predict the outcome of displacement reactions and work out in a galvanic cell which half cell is having oxidation and reduction occurring and then we're also going to describe the role of the hydrogen 1/2 so okay so what is the standard reduction potential table so previously in the displacement reactions video we used reactivity series which are helpful to predict which displacement reactions will occur but a better way to ring ranked things would be to including both metals and halogens would be to test each substance against all other substances right not just metals with metal ions or halogens with halide ions but to test metals against halogens or test metals against anything else in test halogens against anything else so you run a test them all against each other so that you can produce a 1-day table to predict all of your redox reactions ok so to do this we create galvanic cells which are used to test the oxidizing and reducing strings of substances in the oxidant and reductant forms ok and so if you watch the videos on down Vanek cells you'll remember that a galvanic cell is made up of two half cells and each half cell has an oxidant and the reductant form of a element or a substance right and then we compete them against each other alright so that's a look so each half cell is made of the oxygen and the reductant forms that's what we just said alright so that means for this one here might be a metal and a solution of metal ions or halogen in halide ions then they're connected to each other by the external circuit and the sole bridge at this point there will be two oxidants present and they all compete for electrons but not against each other not against the two oxygens but against the opposing reductant from the other cell so what does that mean there are two oxidants in this system okay there is the X plus as an oxidant and the y+ as an oxidant okay so these are the things that are pulling electrons now in metals we know that generally that they're talking about giving away electrons but we're just going to talk about in terms of reductants for now okay because it'll just help a little bit with understanding we're trying to talk about pulling electrons so these two things here are pulling electrons okay and what is it pulling it from this x+ is not pulling it from the X the anode it's metal it's pulling it from the metal why okay so it's trying to pull from the Y metal electrons from it okay and the y plus isn't pulling again from the Y cathode it's pulling it from the X metal okay so they're trying to pull electrons from each other this y plus is trying to pull it from the X okay and the other way around the X plus is trying to pull it from the Y okay so that's what's happening in these cells and then you can work out which one is actually winning via various of the information let's look at this diagram we can see that the electron flow is actually in this direction which means that the Y plus is pulling the electrons from the X okay so the X is giving up its electrons to the Y plus okay so information like that will tell us what's actually occurring okay so let's have a look they will pull out the electrons present until one of them wins that's what we were just talking about okay this will force the other reductants to give up their electrons and give them to the oxidant in the other half so you can tell who is one electrons from direct direction of the flow of electrons which we just had in the last slide or observations about changes in each half so from this you can work out which is the anode and which is the cathode so let's have a think if this half cell here this ss+ right let's just say that s plus is a blue-colored ion okay over time it becomes less blue becomes less blue what does that mean that means that the s plus ions are reducing and turning into s solid okay so that's what's happening because it's becoming less blue okay that's what it means so if this is occurring is this oxidation or reduction this is reduction okay and that means that it is gaining or losing electrons it is gaining electrons okay which means that this electrode s is where reduction occurs which is the cathode and that makes electrode are the anode and Eve we can even work out which way the electrons are flowing the electrons are flowing to the cathode okay the electron flow is that way okay so that's the kind of information that you'll be given you'll be given an electron flow you'll be given information about the observations and you have to work out what's going on from that okay so let's have a look at the next part there so the flow of electrons can do work so which means it has energy and the amount of energy it has is dependent on how much the oxidant or reductant of the oxidant wins by in the electron pulling competition so that means that the amount of energy can then be used to give further information about ranking the substance and can be used to insert new substances and to be existing lists without having to test it against every other substance what does that mean let's have a go through an example so we've just decided that our wins against the s plus and this produces 1.2 volts of electricity okay then we have another set of 1/2 cells okay another set of 1/2 cells this again is the s and the S plus and this again and we're going to put something different in here we're gonna put a Q in here a q-and-a q+ okay here's your external circuit here is your salt bridge okay and then I told you further information that Q also wins against s but only by 0.5 volts okay so in terms of which one is the strongest pulling which one do you think r plus Q plus or s plus is the strongest pulling strongest at pulling electrons hopefully you said R plus is the strongest of pulling electrons so that means that if I was to write build a circuit or build a galvanic cell of Q versus R who do you wear which way do you think the electrons will flow when I have this system set up do you think that it will flow to the R or to the Q hopefully you've said that it will flow to the are and what's even more important is that we can predict how much it will win by so RBS are beat s by 1.2 volts Q beat s by 0.5 volts so Q is slightly stronger than s so ours not going to win by as much and hopefully you've worked out that it should win by something like point 7 volts okay which is the difference between these two values okay so and what's great is that you would be able to predict this value this value without even having to build this cell because we've tested out this two systems and in theory if these two systems are standard and done correctly then this value should be true okay so because we're just dealing with one substance winning against another substance we need to have one standard one so that we can say that all these other values are more or less than this one standard so here so what we need to do is we need to have a standard half cell to which all of us here can be compared by which which how who wins against this standard half cell and then by how much and then we can rank everyone as a better or worse reducing agent than oxidizing agent than the standard half cell and the standard half cell that is assigned is the hydrogen gas in hydrogen ion half cell okay which is effectively a gas and acid this is a sign desert value of zero volts and which is the unit of energy and then all other half cells are compared to it okay so this hydrogen hostel is not special in any way it's just given the value of zero which does not mean it's that it's weak and so remember we are talking about a different concept to reactivity alright because when we talk about who is won or who is lost we're not talking about reactivity because reactivity is kind of misleading when we deal with metals because the more reactive metals is the ones that will want to give away electrons whereas we've been talking about pulling electrons all this time okay so we need to give up that idea of reactivity and this I now just move on to this idea of this oxidizing strength okay so let's have a look so half cells with strong oxidizing agents can oxidize the hydrogen gas in the standard hydrogen half cell and I placed above it in the ranking system so for example we have I 2 here and here is the standard half cell in the middle ok so it has the ability to take electrons from the h2 and force it to give it up and then the I 2 will go to I minus and the h2 will go to H+ okay so it's a stronger oxidizing agent than H+ and therefore will oxidize the h2 to h plus and in turn reduce itself to I minus because it wants those electrons okay so that is what happens when it's above above the ranking system above that in the ranking system half cells with oxidizing agents which are weaker than the hydrogen half cell so let's just say Fe 2 plus yeah ok Fe 2 plus is a weaker oxidizing agent and so it's going to have to gain electrons right and lose electrons so what's gonna happen is that these two things are going to pull against each other and instead of the iron winning the iron our ion for hydrogen is going to win so it's not gonna go that way anymore it's instead going to go that way and it's going to force this one to go that way okay so half cells of oxidizing agents which are weaker than the hydrogen ions in the standard hydrogen half cell will have their reductants oxidized that's this one oxidized by the hydrogen ions that's this one here in the standard hydrogen 1/2 cell and those half cells are placed below it in the ranking system so if you win against the standard half so we get put above it if you lose against the standard half so you get put below it and by how much you lose or win by is this value over here right and that gives us an ability to then compare not only against the standard 1/2 cell but now also it gives us the ability to compare against each other all right so the energy is a combination of the two half-cells can produce and listen in the ranking system that's over here the half cells above it having a positive and the heart cells below and having a native okay so um so above and below whether they win or lose against the standard 1/2 cell all right so what that means is that we can apply that to any of the different substances here right so if you are below it so let's just say zinc and the zinc hostel is run against the cadmium and the cadmium half cell the cadmium two plus will be the winner because it's above it in the standard reduction potential table right and so therefore this will pull electrons pull electrons from the zinc forcing it to oxidize and this one can reduce okay so that's the concept of the standard reduction potential table now the most important idea is like all of that stuff is just all kind of theory about how the standard reduction potential was created all you really need to know is this downhill arrangement so things will occur if you have that and that alright so when you have a galvanic cell you you will always have both parts so in a galvanic cell you will have that and that in one half cell and this and this is another half so which reaction will occur it'll be be downhill arrangement downhill okay from that's the right downhill that's the one that will occur and for that one will go forward and that one will go backwards okay and that's how you predict whether reactions will go be spontaneous because the downhill reactions will be spontaneous if the opposite reaction that's this one and this one will not be spontaneous okay so we're looking for the downhill arrangement the two half equations will be forward reaction of the upper half so forward that's that one there and the reverse reaction of the lower half so that's that one there okay so it means it can be a predictive tool for halogen displacement reactions instead of the reactivity series also a predictive tool for the metal one there as well alright so let's go through a couple of examples on how this is used okay so consider the electrochemical cell bill shown below there is their determine the half equations occurring at each electrode okay so this is a asking you to predict what's going to happen whether oxidation or reduction is going to occur so first of all we say okay well this is a galvanic so there are not everything no this is not a standard galvanic cell but there are certain things here that will oxidize or reduce so let's go find them so let's have a look in 1/2 so we have graphite which is a nert which means it's not going to react and we had this iron 3 nitrate okay so this is an Fe 3 plus and if you look for Fe 3 plus it is here in the standard reduction potential table okay so that's the only thing in that half cell in the other half cell we have two things we have the iron and the iron but this is not I am three this is iron two plus okay so we have the iron and the iron two plus okay so those are the things that we have in this electrochemical cell now remember we want the downhill arrangement but downhill arrangement that means this will occur so therefore the forward reaction will occur at the upper half equation and the reverse will occur in the lower half equation so what are the two half equations going to be it's going to be Fe 3 plus turning to Fe 2 plus okay and then in the other half so it's going to be Fe joint fe2 plus okay that's what's going to occur at each of those half cells look at another question another type of electrochemical cell utilizes the following standard half cell reactions there okay complete the diagram below to show the construction of the operation for so ensure they fully label bla bla bla bla okay so first of all a couple of things you'll notice that both of these half cells or both of these equations are gaining electrons are therefore they're reducing their reduction equations okay so just like all of these are actually reduction equations okay so what you've got to recognize is that one of them will go forward one of them will go backwards okay so that they're not going to go exactly but waiver they show we have to work out which way they'll go so let's find these two half equations this one here is this one here okay and that one okay so if you didn't have this Daniel reduction particular well you would also be able to look at these here okay where these weren't in the standard reduction potential table you'd be able to look at those there and you should say okay well then this one must be higher because that one has a positive value and the other one has a negative value okay so that's another way to be able to tell which one would be on top and which won't be on point okay so we're looking for the downhill arrangement so that one and that one will react that will go the upper one will go forward the room bottom one will go reverse okay so you'd have to work out what's going on now the only pick thing information in this picture is that there's co2 coming out of this little into this little pipe here and it's coming out there and it's going onto this thing here now generally when you see that kind of ladder thing that's a platinum electrode electrode which means it's in an electrode which means it's just not gonna do much okay so this is a galvanic cell which means that it's a standard Hafsa which means that this co2 is being bubbled into its co- form so that must be co- of some sort okay so this is the red half cell here okay the blue half so that means that this must be C r3 plus and this must be C R because that one's the other half so we should be able to label all the other things what's this this is the soap bridge and what this is the external circuit okay so we know that the upper reaction here the red reaction here will go forward right forward if that's going forward that is gaining an electron which means that this is the real gaining electron means the reduction equation okay if this is the reduction equation therefore this one must be the cathode I'll label that some way cathode okay okay and this therefore bin the other side must be the anode well that's going to okay all right so we had the anode and cathode and including their respective polarities we'll deal with that later on okay yep so the electrolytes used okay so we've got that the Co 3 plus and the CL minus and the direction of movement of cations and anions in the salt bridge and the direction of movement of electrons so let's go through the electrons first and then we'll go through the cat night ions for the anions all right so the anode is where oxidation occurs and that's where we lose electrons and so therefore electrons are flowing this way through the system so therefore you draw a little arrow that way okay for the electrons and then we have to think okay if the electrons are going that way that means that this cell is becoming more negative more negative if that's becoming more negative it needs to have some ions to balance it out that means the positive ions are going to go that way and the negative ions are going to go that way okay why are the negative ions clean that way because this is becoming more positive because it's losing electrons and so therefore it must balance that out with some more negative things there okay so that's that we have all of the things just with this polarities the anode is always negative for galvanic cells and the cathode is always positive for galvanic cells so there we are there all right so that's that if you have any questions which I'm sure you will come and ask but remember just this downhill arrangement is how we predict what we reaction will actually occur okay that's it