Transcript for:
Understanding Electrode Potentials and Cells

[Music] [Music] hello my name is Chris Harris and I'm from alibi chemistry and welcome to this video on AQA electrode potentials and electrochemical cells so this video is dedicated to AQA so if you are studying AQA a level chemistry then this is the video for you it's not like say some other resources where they may be quite generic and they'll which is nothing wrong and that of course on generic information but you might find that you look at it and think is this on my specification do I need to know this well this video solves that and actually goes through everything that you need to know for the AQA specification and in fact there is a full range of videos revision videos like this for year one and year two just have a look on my every chemistry YouTube channel there's a full range there there is also some pass paper walkthrough questions and some white board tutorials where you can use them to look at specific areas of chemistry which are a little bit more generalized so but there's a full range of things on there and they're all for free there's no charge for them whatsoever very comprehensive all I ask is that you just hit the subscribe button that would be absolutely fantastic just to show your support as long as I as long as people keep subscribing and keep watching then I'll just keep on making them as simple as that really so also these slides are available to purchase on my on my test shop so if you click on the link in the description box below you'll be able to get ahold of them there and they're great value for money you can use them on your tablet or your smart form when you're on the move you can also and I know some people who have printed them off and use them as revision notes then print off these to the individual slides and collected them all together so all the topics are here and they are dedicated to AQA so go and have a look there click on the link of the description box and you'll be able to get ahold of them there so like I say this is dedicated to AQA and obviously it meets the meets the specification points that you can see printed printed on there on the top there so we're mainly going to look at things like electrode potential electoral potentials and cells so we'll look at half cells at an EMF etc and then towards the end we're then going to look at and commercial applications so batteries and then electric and of electrochemical cells and we're going to look at fuel cells as well toward towards the the end of the video so this has a lot of electricity so if your physics minded then you're gonna absolutely love this put it that way so I like physics as well so so yes so we're gonna we're gonna start by looking it obviously the cell side first okay so let's look at half cells so half cell is one half of an electrochemical cell and they can be constructed of a metal dip to enter into its ions or a platinum electrode with two aqueous ions in okay now you might have done these if you go to school of college you might have seen these we do them independently I'm not sure if you have a not but certainly if you've if you look at school of college these types of practicals are pretty pretty good practicals so for example you have a half cell is basically just very simply just a bit of metal stuck in a solution of its ions dead simple dead straightforward so for example here we see we've got an iron electrode and the anion electrode is dipped into the in2 or in3 solution below and actually when we when the metal comes into contact with its ions there is a reaction that occurs believe it or not and the reaction is in equilibrium and we have an Fe 2 plus for example here we have Fe 2 plus ions so we have Fe 2 plus plus 2 electrons and will actually be in equilibrium with the iron so there is actually a reaction that's happening here at the at the end of this electrode so yeah so that's something you might not have known so and you can see here sometimes we have a situation where we might have two metal ions so it's alright if we've got a metal and an iron but some ions don't have a solid version so there is a solution to that literally there's a solution to that where we use a platinum electrode like this and we have a solution of example different iron so you might have the reaction between fe 2 plus and fe 3 plus so the equilibrium reaction of them two now obviously there is no metal version of this because all ions are dissolved in solutions of their aqueous so so there's no metal version so we still need an electrode so we need an electrode such as platinum platinum is very expensive but it's a it's got a good electrical conductivity so that's that's pretty good we're gonna make a cell and also it's inert the last thing we want is our electrode to be starting to interfere with the reaction in our solution because we just want it just to conduct electricity that's all we want it to do so thankfully platinum is perfect for that job and so an electrochemical cell is created by joining these two different half cells together and we get a full cell okay so let's have a look at what these and what these cells are and what they look like so it's made up of basically two half-cells joined together with a wire a voltmeter and we've got a salt bridge as well between the two so we're going to look at the what a electrochemical cell looks like so here's the picture here and you can see we've just taken two half-cells here we've used a zinc and a copper one okay and one side is actually undergoing a reduction process whereas the other side is actually undergoing an oxidation process so essentially we have a redox reaction it's not like the shower gel it's it's it's reduction and oxidation of course it is so yes so essentially we have a redox reaction happening here so let's have a look and label this diagram so we see we've got a voltmeter here and this is used to measure the voltage and I know physicists may be screaming at me it's not the voltage it's actually the potential difference I know so it's the potential difference then potential difference is measured in volts and it's between two half-cells and this is called the EMF or the e cell okay and electrons these flow from a more reactive metal to a less reactive one okay that's very very important so they flow from a more reactive metal to a less one as long as you know the direction in which these electrons are traveling and we're going to look at that as well because a lot of cells a lot of these electrochemical cells as pride is predominantly to do with other metals in their solutions of course it is but also is to do with mainly to do with trapping the electrons so utilizing electrons that would normally transfer so between these two because if we were to mix these into beakers there would be reaction but what we're doing is were kind of tricking it what's in right yeah you are reacting but what we want to do is take the electrons and at the same time and actually make some use out of it so two for example and power electronics that I'm using now to record this so and the zinc 1/2 cell shows the loss of electrons as zinc loses electrons easier than copper and we're going to look at a little bit more this on the on the next slide looking at and you know how do we know ones more more likely to give up electrons and the other we're going to look at that later on basically it's the reduction oxidation so effectively the zinc and the zinc reaction is zinc forming zinc two-plus plus two electrons as you can see here so what we'd see our observation is actually our zinc electrode here would get thinner this is because our zinc is converting into its ions which are in solution as you can see there and I've seen it'll form two electrons as well okay so we're producing more zinc two plus ions in solution at the bottom there so let's have a look so electrons go round and they go round to the copper electrode okay so that's basically what's happening in the zinc side so what do you think might happen with copper then so let's have a look so our copper actually accepts them electrons that were produced by zinc on the other side and reduction happens here so effectively the copper two-plus ions in that solution there they react with em the electrons that came from the zinc and they form copper metal so see us and what we would see observation is that electrode will actually get thicker because the copper two-plus ions would accept the electrons to form the copper the copper that actually coats the bottom of that electrode so you can see here there's two bits to it there's always a reduction side and an oxidation site and depending on what we connect it with will depend on what type of reactions were going to get but the principles similar the similarities are that they all produce there's all the transfer of electrons and what we do is we use them electrons to do something with it that's all we're doing else essentially how a battery or cell works okay we look at batteries later on but batteries contain these cells okay so we have the salt bridge as well obviously our salt bridge is contains potassium nitrate a salt bridge is a funny thing because really you just use practically you just use some filter paper you convert that role the leg around filter paper you roll that into a strip you dip that into saturated potassium nitrate solution very very saturated solution and then you drape it between the two beakers as you can see on there so they're draped between there and there and really and the role of the salt bridge is just to complete the circuit you don't really need to know much about the salt bridge but it just allows the ions to flow through which balance out the charges in the two beakers so you don't need to know much about that other than the fact that the salt bridge must be dipped into the solution of each beaker and the salt bridge must not come into contact with the metals that the metals here so you just drape that across and just that allows the ions to flow through and balance out the charges okay so we're going to look at an electrode potential so remember what we said was a more about the more reactive and least reactive metals so what we're going to look at here is what we mean or how we can we identify if something's been reduced and something's being oxidized well actually what we what we can do is we can use something called electrode potentials which are enormous and what these do is these measure the electoral potential of each half cell in volts and it shows us the how easily the half cell gives up electrons which is oxidized ok so just remember that so you see from that previous slide we had the the two half cells which were zinc and copper half cells and you will have noticed that we have obviously the two half equations here as you can see there there's not two half equations so we've got zinc and copper and you can see they're both written in the reduced form so they're both plus electrons okay so that's the first thing so we always write them in that way and you'll notice that in the electrochemical series and we'll look at that later as well later on in the video but the electrochemical series is something that you will have a reference to so you will have a copy of this in the exam you may have a data book of use it within lessons or lectures so but you will always have the numbers you're not expected to remember the numbers thank God because there's a lots of information there's a lot of them there and so basically we always show in the electrochemical series we always show them in the reduced form so you'll say for example that each one of them species is gaining electrons that doesn't mean that that's actually happening in the cell this is just a theoretical a theoretical view okay so remember when we connect two half-cells together we always have one half cell undergoing reduction and we have one half cell undergoing oxidation okay so remember that so we need to work out which one is being reduced and which one's being oxidized and we essentially do that by looking at the electrode potential value and which is the e naught value and of so you'd find that in a datable so the key thing is we can see that from looking at this data and again this is data that you'd have access to we can see that the zinc two-plus zinc half cell so this one at the top here has a negative enopp volume and the Cu two plus at CU half cell has a positive enopp volume okay so they both got different values here okay everything will sum so what what does that mean well I've got an acronym here now you might remember it in a different way you may be taught this in a slightly different way there's a few ways in which you can teach electrical potentials and and they'll probably get to the same point which is fine okay there's no right or wrong way of getting there and necessarily and it's about making sure that you understand what's going on so I'm going to introduce you to an acronym which I like to use because it helps to and keep everything neat and tidy and you'll see this when amused actually later on in the video I'll use it a lot so I'm gonna introduce it to you and it's not very creative but at least it'll help you to remember what's going on so and we remember the rule no problem ok it's not really an acronym but kind of so no problem so the most negative half cell will undergo oxidation the most positive half cell will undergo reduction okay so you can see there we've got negative half stolen to go with oxidation so this is the most negative so that's going to go and undergo oxidation and this one's the most positive so that one's going to undergo reduction okay so you can see most negative oxidation takes place so because that's the most negative and oxidation takes place remember oxidation is the loss of electrons so you remember the acronym from year one chemistry that I use which is oil-rig oxidation is loss of electrons reduction is gain of electrons so because this is oxidation we've got to show the equation losing electrons really simple all we do anything that's been oxidized or the oxidation process we flip that equation back to front so that's exactly what we've done there with the zinc we've taken the zinc forms ink two plus plus two electrons okay so we flipped out the other way around keep the other one the same because that's still shown reduction so then in this cell we have zinc giving up electrons and is it copper two plus accepting them and what we can do is combine these two half equations together to form a full lie on equation and so this is the product of it here so you've got zinc solid there plus copper two plus will form zinc two plus and copper solid okay so um remember that acronym no problem that's the if you forget anything in this video don't forget that acronym because it's going to really help you substantially later on when you're looking at more complicated examples okay right so these inordinately negative and positive values etc now these unit values will only exist if we connect it with something else and you might think well how do we know that zinc has got a - not point seven six and how do we know copper has got plus not point three four how do we know these figures well I hear you'll cry so I'm we have something called the she it's called a standard hydrogen electrode okay and basically the standard hydrogen electrode is used as a reference to measure standard electrode potential so all them figures there were calculated by connecting it with a sheet and then that's allowed to mammal to see that that's then allow us to much measure the the e naught value so electrode potentials of half cells they can't be measured on their own no surprise we can't just put a bit of metal in the solution and say it's generating electricity because it won't it's just going to react with there it's going to exist in equilibrium so what we can do is you can measure the inert and reference this to a standard hydrogen electrode so a she and with Enoch value which equals zero it's very important in science to to have a reference because it allows us to if we have a reference say in this case it's gonna be the stand hydrogen electrode it allows us to compare values globally so if it was somebody in say in America or Germany or France of Japan or anybody who's doing these types of reactions okay and these reactions will be these reactions will be standard them under standard conditions and so that allows us to compare the values and it's a universal language effectively so you know we measure everything against this value so this is the setup here okay so we've got here we've got a copper half cell and we want to measure the electrode potential of this copper half cell and we've got our platinum electrode you see on the left there so let's have a look at the set up so what goes in here is hydrogen goes in this has got to be a 298 Kelvin and 100 kiloPascals because that is just they are your standard conditions we must also have 1 mole per diem cubed of H+ ions in solution there okay so that's very important and we must also have 1 mole per diem cubed of copper ions which are obviously in the beaker there on the right so these are key things that you need to know you need to know the general setup so your standard hydrogen rod has a platinum electrode and it has a little glass tube around it and it's got hydrogen going in at 298 100 kilo Pascal's and we've got one mol per diem cubed of our hydrogen ions in that solution there as well as one more PM kujan or copper so these are the standards so we must stick to these standards and you must be able to remember them so they'll be expecting you to be able to recall the standard conditions for a standard hydrogen electrode so that's temperature 298 pressure at 100 kilo Pascal's and concentration of ions at 1 mol bgm cubed so the diagram on the Left this obviously shows a standard hydrogen electrode connected to a copper copper two-plus half cell now assume in the conditions above are met then the voltmeter should actually tell us the standard hydrogen electrode for copper and the copper copper 2 plus 1/2 sum so what we've got to do though we're going to be careful with this okay and there's another potential banana-skin with the exam so we've got and the movie said we need 1 mole per diem cubed of h+ ions in that beaker so remember we needed this here so we need 1 mole per diem cubed it's of h+ signs not of acid so don't get this confused and the reason why we shouldn't get this confused it's because um it's the UM they might say the right we're going to use this acid such as hydrochloric and that would be fine so hydrochloric acid 1 mol per diem cubed of hydrochloric acid produces 1 mol per diem cubed of H+ ions but if the examiners are talking about sulfuric acid sulfuric acid is a diprotic acid so it produces two molecules or two H+ ions per molecule of h2 so4 so actually we only need half mole per diem cubed of sulfuric acid because that will itself produce one mole of H+ ions so a must stress that the concentration is the concentration of h+ ions not the concentration of acid okay if it's a monoprotic acids such as HCL then that's fine we've got to be careful because they they will and they well they may ask you to they may use sulfuric acid just to see if you pay attention so just watch out for that I've seen it before so okay so we're going to look at the electrochemical series so remember we've been using some of their mean or values just to some of the examples before and I've mentioned the electrochemical series is a big basically it's a big table of series so it's basically a say it's a lift list of half cell reactions and their standard electrical potential so these have all been calculated by measuring them against that standard hydrogen electrode under standard conditions okay so I'm going to put up an example here and so this is an example of some of the half cell reactions and that we need to know there is a massive list of these it's not just these ones that you can see on there honestly I can't fit them all on the screen but in the exam they will give you the ones that that that you will you know that you'll need so don't worry about remembering obviously these these equations nor the numbers it's about using the numbers and to work this out so you can see and the table is also written in descending order and you you may actually get it the other way around as well so in this case we're starting up C+ at the top here it been cursor ends there you go positive at the top and negative at the bottom as you can see there and notice that all just like we said before they're all shown in the reduced form this is the standard way in which they display them so it's always something plus electrons okay it's always in the reduced form but remember depending on what you're bonding it with one of them will be oxidized and we flip it the other way around but obviously we have to look at that later okay so one of the things the last cue in the exam as well is about oxidizing agents and reducing agents and so remember just a reminder and oxidation is loss of electrons but an oxidizing agent gains electrons okay so reduction is the loss of electrons but a reducing agent gains electrons okay so we're going to look at them using them definitions or using that knowledge there we're going to then apply that into the electrochemical series and try and identify which ones are the most powerful reducing agents and which ones are the most powerful or sizing agents so let's have a look at the oxidizing agents first so as we go up this table in this case this is a table gun in descending order remember if it was in a different way the other way around then it would be going down but this one show in this particular example here so just be vigilant with that agents on the left-hand side of the equation are more easily reduced okay so for example this is like chlorine for example so they have an increasing tendency to gain electrons a more powerful oxidizing agent so remember this positive value is telling us that this reaction is is very likely to go negative values tells us that is that just really isn't going to work so a positive value means it really is going to go so chlorine is going to be more than happy to accept these two electrons to form CL minus so because it's it's a positive value this is telling us that it is a more powerful oxidizing agent so you can see here that the most powerful oxidizing agent is chlorine in this example and the weakest oxidizing agent is magnesium 2 plus because that's at the bottom of this table here but so is on the left-hand side ok so we're always comment on the left-hand side not the right-hand side because that's to do with reducing agents which we're going to look at now so as we go down this table this particular table we have a stronger reducing agents that's been produced so you can see agents on the right-hand side of the equation are more easily oxidized so it's these ones here on the right-hand side and they have an increase in tendency to lose electrons in other words they are more powerful reducing agents so your most powerful producing agent here is going to be magnesium which is the one on the right here this is the most powerful reducing agent the weakest reducing agents can be CL - it's many of the most powerful ones that they're going to ask you about so make sure it's the most negative and the one on the right of the equation it's going to be the most powerful reducing agents the most powerful oxidizing agents is going to be the most positive and the one on the left the species on the left hand side of the arrow okay so make sure you understand that there's a lot of information there but you know as long as you can visualize that graph or that chart with the hours then you should be fine okay so what we're gonna do obviously we've got the electoral potentials we know what they do we know what reduction and oxidation means and we know we've got a whole list of these standard electrode potentials or these enoch values so now what we're going to do is we're going to calculate the standard cell potential okay so the standard cell potential e naught or the electoral potentials these can be used to work out the standard cell potential which is e naught cell okay so the electoral potentials are for the half cells that we've seen obviously and when we connect them together we form a cell we can actually work over the north of that value as well and we use a very simple formula here probably one of the most most simplest ones inor of the cell is e naught reduced minus seen or oxidized so the way in which I remember the which way round it is is I I normally basically call it redox so read and then ops okay so reduction reduction oxidation redox okay so it's always reduction first minus oxidized okay so remember your half cell equations with the most negative a naught value is being oxidized remember that no problem negative oxidized positive reduction okay or reduced so if you have two positives or two negatives then it is the most negative that is oxidized okay because that may happen so that's the simple rules remember that no problem negative oxidized positive is reduced so here's your first example so we're going to use that electrochemical series that would see before and we're going to calculate the e nought of the cell when a CL 2 CL minus and a zinc two-plus and zinc half cells are connected together so we're going to look at our data so we need a CL 2 in the CL minus here's our first one so that's plus 1.36 and we've got a zinc 2 plus and zinc half cell which is going to be down here so that's going to give us minus not 0.76 so what we're going to do the first thing is we need to identify which is being oxidized okay so in this case the one that's been oxidized is our zinc 2 Zink half-cell that is the most negatives remember- no problem negative oxidized so that's the one which is being oxidized so what do we do with the oxidized one well we'll put that on the right-hand side of the equation so we're going to put all the figures in and we get that okay so we get plus two point one two volts so it's one point three six minus - not 0.76 okay let's have a look at another example so we're going to use the data in the electrochemical series again and this time we're going to work it out the e naught the cell when the chlorine chloride electrode half cell and actually reacts with a copper and copper two-plus electrode okay so when we connect these so again we need to Aqua we need to work out what's been oxidized so we've got our chlorine here which is one point three six plus one point three six and our copper which is plus noir point three four so these are both positive values so it's the one which is the most negative has been oxidized which in this case will be this one so there it is so the copper copper two plus one is the most negative and this one's been oxidized so we put that on the right-hand side of the equation here and we subtract them away from each other and we get plus one point zero two volts from this cell here okay okay so that's fine so we know how to work at the you know at the cell it's fairly straightforward just remember that acronym no problem now what we need to be able to do is be able to draw our cell so instead of drawing two beakers with a salt bridge in between and the electrons etc and it's there's a better way of drawing it which is much quicker and much neater and we call this a cell notation okay and cell notations are basically used to simplify how we draw the setup of a cell okay it's a universal way of doing it and it allows us to draw it really created quickly and you'll see this well but there's a few rules that we need to know so as a standards they are represented like this and the most negative half cell goes to the left hand side of the double line okay so you can see here that we've got our M standard cell set up here so effectively the solid lines in a cell notation show a physical state change so that might be for example a solid electrode in contact with eight quiz ions there's a physical state change between the two the double line choses a salt bridge so this is the bit in the middle and so let's look at a specific example of the zinc copper salt that we'd seen before and so using that zinc and copper cell we can set it up like this and you can see that we have our zinc on the left there it is and we've got our copper on the right so zinc was the most negative so it sits to the left of that double line okay now you can see we've got reduced and oxidized form so what we're looking at is you have two elements to the half cell you have the metal which is zinc and we have a zinc two-plus which is the ion this is more oxidized compared to that because that's got an oxidation state of zero because it's an element's this has got an oxidation state of plus two so that's more oxidized so that sits closest to that salt bridge and likewise on the other side as well okay so there we are so that was an oxidation state of plus two zinc has zero so that one sits closest to the salt bridge you've got to be careful though because we have ions okay so we have ions in solution so obviously all of these and is looking at solid with the ions okay but if you have two ions in solution that's going to be difficult because we don't have a physical state change between them for example Fe 2 plus Fe 3 plus we can't put a solid line so luckily we have a method of doing of tackling that as well and that method is actually using a comma so we use a comma not a solid line because they're in the same state so let's have a look at an example so you can see here we're going to use Fe 2 plus Fe 3 plus and we use magnesium 2 plus and magnesium cell so you can see here that we've got to have these obviously solid and aqueous so that's a physical state change so you put a solid line down there but between Fe three-plus and Fe 2 plus there is the same state change here so this is a quiz and a quiz so there's no change there between these so put a comma between them because they're in the same state clearly there's a solid line here we must on a solid electrode and in this case we're using a platinum electrode because it's Platinum solid okay so remember and that's what we need to use for these types of cells because we've got to eight cosines and no solid electrode here so we use platinum okay right so now we know these cells we know how to draw them out and we've seen the electrochemical series we know what's oxidized and what's reduced and we know how to identify that and we know how to form these cells we're going to put all this together and use it to predict if a reaction is going to go now this unfortunately will not predict a lottery numbers because that would be brilliant if it did and but it will it does the next best thing which is to predict reactions so we're going to predict the feasibility of reaction using these inor and electrode potentials and it's basically going to tell us if a reaction is likely to proceed under standard conditions so we'll bring back our electrochemical series again so there it is in the bottom left and we're going to look at this specific example here so we're going to use the data and the electrochemical series to predict whether magnesium will react with copper two-plus ions in solution under standard conditions so the first thing we're going to remember remember the no problem the no problem acronym that we used so half cell with the most negative e naught value has been oxidized so the first thing we need to do here is identify what's being oxidized so in this case the mg 2 plus mg 1/2 cell has the most negative e naught value so is oxidized okay so can you remember what we do with oxidized reactions well what we do is we need to reverse it we need to flip it round the other way and then once we've flipped it around the other way we write both equations that we need side by side so you can see here that our magnesium equation has been flipped so we've got mg mg 2 plus + 2 electrons that's been oxidized and this one remains the same because that's just been reduced so write them side by side like that and then what we need to do is combine these two equations and the equation that we form here is actually the feasible reaction so we can see here the one which is feasible is magnesium reacting with copper so what we do is we compare that reaction that we've just developed there and compare it with the question and we can see that actually yes magnesium will react with copper two-plus ions because it tells us here it says that magnesium will be out with these copper two-plus ions and this is a feasible reaction so it's a match so it does work but if you've seen any other videos I like to make sure it's definitely correct and it's a good technique to have in the exam because it gives it up a bit of confidence if you if you do these little checks to make sure that it is correct so one way of making sure this is right is you just chuck your numbers into that inor reaction remember that equation sorry so remember that equation that we'd seen that's the one on the top there so we put the figures in reduced minus oxidized and we get a positive value and basically if it's a positive inor cell value that means the reaction is feasible okay any inor cell that's positive it means it's a feasible reaction under standard conditions okay right let's take a look at another example here and so this one is looking at rust so there might ask you to justify why something happens so same principle will just confirm it so basically the questions told you iron nails become rusty wedding contact with air and moisture oh that's obvious because you look around and you can see that iron rusts we can see it around us and we know it exists what the examiners are after is a justification so we need to prove that that is the case using the e naught values there you know to justify the a shear this is definitely right we're given some credibility to this to this statement so the first thing is identify what we need to actually use so we're using our water and oxygen equation here because it's air and moisture so we're using this half cell and obviously Fe 2 plus 2 Fe and we're using this half cell as well because this is iron going to iron 2 plus which is rust okay so iron oxide is is the rust on on the nail as you can see on there so the first thing we need to identify what has been oxidized so what's been oxidized here is the fe 2 plus Fe 1/2 equation is the most negative inor that with the 2 so therefore that one is being oxidized so can you remember what we need to do now well we need to flip that equation around and write the two equations side by side so we reverse it and we write the equations there now this is where it's a little bit different because you'll notice that we have two electrons on the top equation and we have four electrons on the bottom now we can't combine these until the number of electrons is equal so for all those who do maths this is just a similar sukwon to try and get one of the factors here the same and cancel them out so what we're gonna do is we've got four E and two E we're gonna cross that one out and cross that one out well we're gonna multiply this top row by two first there we are multiply that by two first to get to Fe 2 Fe 2 plus and for electrons and then we cancel out and combine so when we combine them equations we get something like this now this is our feasible reaction okay so this is the reaction that's feasible now what we need to do is you use this equation that we've just come up with there and compare it with what we've been asked for in the question and it's Astor's for to justify this and so basically the statement that we can say on this is that iron does indeed react with oxygen and water and it will form iron two-plus which is your iron oxide that'll react obviously to form iron oxide which is you've lost and an alkaline solution as well so it's a china s-- and of course we can improve this using our Enochs cell equation so if we put the numbers in inor reduced - II not oxidized put all your numbers in and if we get a positive value which in this case we do it's not a massive positive value but it's still positive so under standard conditions this reaction will go because we've gotten a positive enon cell value okay right so once we've got all this information we can now use these cells that we've been looking at and scale it up a little bit bigger and look at something called batteries which of course you know what a battery is so and batteries are just electrochemical cells and that are joined up and batteries come in two main forms and these are rechargeable and non-rechargeable batteries okay so non-rechargeable batteries tend to be cheaper than rechargeables so but your rechargeables are reversible and they can last longer so cheaper in the long run so a classic example of a rechargeable battery is your mobile phone that's exactly the same so your phone when you plug in your phone it recharges the battery up and if you if you have a nun with non rechargeable battery that's just like a standard like say like a standard battery that you would purchase you know from the shop like an AAA battery for example and use it once and then you dispose of it so lithium-ion batteries like I say are a good example of you rechargeables and they're commonly used in wireless power tools tablets mobile phones and lectric cars so you can see a lot of different uses for these batteries now the lithium-ion batteries for example them use in mobile phones they have the following components they have electrode a which is lithium cobalt oxide so that's what's in a lithium ion battery and electrode B is graphite so we use graphite which is just carbon and the electrolyte which is the bit in between we'll look at the electrolyte when we look at fuel cells the electrolyte is the solution that contains that's within the battery that straddles between these two electrodes is a lithium salt dissolved in an organic solvent okay so remember this is the electrolyte is the part of the battery that acts as the conductor pathway for the ions to move from one electrode to the other so it's a bit like a sea of ions effectively okay now to identify which is the negative electrode we need to establish which is producing the electrons ie oxidation okay so what the example will expect you to do is to use this information here to answer questions on it using the same principles that we used before which is no problem negative oxidation positive reduction okay so to establish the overall reaction in a lithium ion battery we need to know the half equations at each electrode so you can see here here's our half equations and I've see these would be these to be given to you so we've got a positive in value and a negative not value for each of the different reactions that's happening at each of the electrodes so lithium the Li plus Ally has the most negative inor value so oxidation occurs here so remember we flip the most negative equation around and this shows the electrons are actually being produced so this is your negative electrode okay electrons being produced is the negative electrode so the negative electrode we have this reaction so we've just flipped that half equation around this one here because that was the most negative and obviously the positive electrode we keep that reaction keep that half equation the same we look at our electrons there's one electron in each so that's fine so we balance them cancel out the electrons and write the overall equation but this is the overall equation for the discharge so this is using the battery so for example and using your mobile phone without it plugged in so the reaction that occurs when you're using that phone is lithium reacting with your cobalt oxide or we cover oxide which is here and this will form your lithium salt okay which is this this product here so we can work over the enol to the cell for this by using the equation that we've seen before so we do produced minus oxidized put the numbers in and we get it Enoch the cell of plus three point six volts so that's quite a significant potential difference there so that's why these batteries are these this setup is really useful because you need a lot of power to power you know phones and laptops and tablets etc okay so rechargeable batteries and work by simply plugging them in to supply a current so that's a flow of electrons and so these this current forces electrons to flow in the opposite way and as we do what we do is we reverse this reaction so if it asks you to write an equation for and the recharging of the battery all we do is we flip that equation the other way around because we're just forcing the lithium lithium salt to reproduce your lithium and cobble oxide obviously with the help of electrons and to do that okay so that's really what you need to know about batteries now the the kind of final part is to do with fuel cells so fuel cells are light batteries they need a continuous supply of fuel okay rather than batteries which have a ready store of the chemicals so you don't need to keep supplying and fuel to a battery it already has the energy of the chemicals required to do that but a fuel cell needs a continuous supply so if you look here we're going to look at a specific example for this one this is an alkaline hydrogen oxygen fuel cell as an example of eating get acid ones as well but for this example we're going to look at an alkaline one so you can see we've got a picture of the fuel cell on the left and basically what we're going to do for the next few slides is just look at what these numbers mean what's actually happening at these because you are expected to know that and and then once we've done that we then need to we're then need to summarize it and look at the reactions overall so that's what we're gonna do first so we'll look at number one and so out point number one we've got a hydrogen feed here so the hydrogen feed as you can see and reacts with the O h minus signs that actually come from this side of the reaction which we'll look at in a moment and basically the hydrogen reacts with the hydroxide ions to form water and for electrons okay so then the flow of electrons this is endpoint to the electrons produced in that part travel through the Platinum electrode and through the wire so remember Platinum is good because it's inert it's good conductor of electricity so the electrons flow through here and this flows through to the third section which is the component so this is could be for example a car empty power or a car it could be empty power power in a house if you had a quite a few of these fuel cells so anything like that this is where the use comes out of it okay so now I'm moving on to step four which is this bit and this is basically oxygen is fed in here so this reacts and with water and the four electrons that have been produced from this side and to make the RH minus signs that were needed to react over here so the overall reaction is oxygen coming in reacting with to lots of water and four electrons and that forms for lots of h- ions and that are produced here and love to see them demo which minor signs are then used to react with the hydrogen that's coming in on this side so the fifth one is the negative electrode so this here is the negative electrode this is the cathode and so electrons flow to the negative electrons that come from here to the negative electrode and this is also made from platinum now number six this is the electrolyte so this is the green section here on this diagram so this is made from potassium hydroxide solution which is this bit here it carries the Oh H minus ions here from the cathode from this bit to the anode to this side at seven this is your positive electrode this is your anode this bit here and so electrons flow from this and it moves across to the to the cathode which is this side and then eight water is eliminated here and so obviously this is the product of step one some of the water then emits out here so the only emission with this type of fuel is water now this is very different to say a fossil fuel where you produce carbon dioxide and water and Maeby's carbon particulates carbon monoxide nitrates and sulfates and all sorts this fuel is much much cleaner the only product here is water that's produced that's the only by-product the last step the ninth step is the movement of the O H minus ions so the O H minus signs that were produced here at Step four are actually carried through the electrolyte here and towards the towards the anode at this point so that's all like nine symbolizes so you can see that we also have some ion exchange membranes as well and the ion exchange membranes actually sit on the electrodes and they sit on the side where the electrolyte is and these membranes are live with the rh- ions to pass through and but not the hydrogen oxygen gas the last thing we want to do is through our hydrogen oxygen gas to pass through and meet because the electrons will be transferred here and we won't actually get the benefit through the wire so that's why we must keep them separate so to summarize the first step is the hydrogen feet so the hydrogen is actually fed in through here reacts with yo h- ions from solution from step four which is here they come from the solution step nine so the overall reaction is 2h 2 4 or h- producing 4 h2o and 4 electrons and obviously the oxygen then is fed on this side so this reacts with the 4 electrons that come from and the hydrogen feed and this producer the forage oh it's minus ions which are then fed through to feed on this side so it's one big circle as you can see now we can combine these two equations and cancel out any species which are the same and such as your waters and electrons etc and we form this overall reaction which is showing what's happening so effectively it's - lots of hydrogen reacting with your oxygen which is here and this little form of C two molecules of water and which is produced as a as a byproduct ok so let's look at the pros and cons of using these it sounds like a really good idea using these doesn't it because there's no harmful emissions it's just water looks as though we're producing we're using oxygen as well you know which we can get and you know quite readily so let's look at the the pros and cons of using these fuel cells to generate electricity so the advantages and is that fuel cells are much more efficient than an internal combustion engine and some cars are now actually hydrogen powered I believe I think California I think new hydrogen power don't use much of it in the UK who have mainly gone down the line of electric cars rather than hydrogen powered and more energy is converted into kinetic energy and so therefore combustion engines waste a lot of energy as thermal energy so these hydrogen abso these fuel cells are much more efficient excuse me right so many electrical like so many electrical vehicles are battery powers however like batteries and fuel cells don't need to be recharged you literally just fill up you have a tank of hydrogen in your boot and you fill up with hydrogen and you just need a ready apply that hydrogen and obviously oxygen which can come from the air as well and also the only waste product is water as I've said no carbon dioxide is emitted so unlike what you get in a traditional and combustion engine so these are a lot better for the environment so the disadvantages so the disadvantages are that hydrogen is actually really highly flammable and it must be stored and transported correctly it's not the type of gas that you want near a naked flame so and we've got to handle it really carefully and and as a result it's expensive to store it so there we are so it's expensive to store and transport it effectively your transport and round a very light gas so that has to be to make it cost-efficient it's normally pressurized it's a really high pressure so you can transport more of it around and you know purple lorry so so yeah so storage it's got to be in pressurized containers and also energy is required to make the hydrogen and the oxygen in the first place and quite often unfortunately fossil fuels are used to use to pass water through an electrolysis process and this also uses fossil fuels so there's an indirect pollutant here because the actual manufacture of the hydrogen in the first place can use fossil fuels but of course if the making of the of the hydrogen is used by green sources then of course is quite a green way to go but as it stands you know we're still using a lot of new fossil fuels in terms of electrolysis process okay and that's it so that's the end of the video on electrode potentials and electrochemical cells so you can see there's quite a bit of information in there but the standard acronym like I say anything if you're going to forget anything don't forget this which is and you know no problem negative oxidized positive produced ok so just remember that and like I say there's a full series of AQA videos on Alawis chemistry YouTube channel and they're all for free as well as whiteboard tutorials and exam work throughs and constable 3 all I ask is you just hit the subscribe button and that would be absolutely fantastic I'd really appreciate that and and also you can purchase these like I said right at the start the video just click on the link below and you'll be able to get a hold of that great value for money but that's it bye bye