[Music] in this video we're going to be looking at topic 16a and then redox equilibria is quite a large topic so we're going to break it down into certain parts and we're going to start by looking at the standard electrode potential this is a very complex topic that does take time to build up to the overall outcome so we're just going to take it step by step and this is the first video of the redox topic and this is for the ie2 course for nxl for chemistry so we're going to be looking at what we mean by oxidation and reduction why we use standard electropotentials and what this actually means how we measure it and how we relate that to the emf of a cell and then how that relates to oxidizing and reducing agents so a bit of a reminder where of what you should know by now you should know what oxidation and reduction are so if you remember you have your oil rig so oxidation is loss of electrons where we get our oxidation number increasing and reduction is the gain of electrons where we get our oxidation number decreasing and of course anything that is oxidized must have a reduction reaction that goes along with it and we call these our redox reactions what we are looking at now is a very specific type of redox reaction and it is between a metal and their solution so if we take a metal such as magnesium or copper and we place it into water or into a solution of itself some of the atoms will lose electrons and go into the solution as positive ions so we have our solid metal losing electrons to become positive ions and they lose n amount of electrons depending on what oxidation number they started with the electrons that are lost when this change happens will build up on the surface of the metal because remember of course electrons don't go into the solution so they stay built up on the metal and eventually that buildup of the negative charge attracts the positive ions back to it so our electron becomes negatively charged due to this buildup attracting the positive ions back and when this happens some of those positive ions will then regain their electrons and they will go back to the surface so we get the reverse reaction happening where our positive ions gain an n number of electrons and we go back to being the metal on the surface because of course we have a forward reaction and a reverse reaction happening eventually we will reach a dynamic equilibrium so remember dynamic equilibrium is when the rate of the forward equals the rate of the reverse reaction so we will reach this point and we get this equilibrium between the ions leaving the surface to go into solution and the ions leaving the solution to come back to the surface and we write it out like this and we have this reversible sign and notice i'm writing it in terms of the reduction reaction i always write this equilibrium in terms of the gain of electrons and that's going to become very important later on when we look at the data book and the information that is on that so once we build up this equilibrium we must be able to then use this in some way and this is where this is a starting point for a redox equilibrium so if we take the metals like magnesium and copper both of them will go undergo these reactions however their positions of equilibrium are going to be slightly different so for the magnesium it's actually going to be slightly more to the left so if i have magnesium gaining the two electrons to form magnesium metal i can also do a similar reaction with copper where it also gains two electrons and becomes the metal but these have different levels of their equilibrium position so for the magnesium the equilibrium lies slightly more to the left meaning this reaction is more likely to go from the magnesium metal to the ion this is because magnesium can release electrons much more easily than copper and you actually already know this because you know this in terms of reactivity series magnesium is above copper and the reactivity series meaning it prefers to become an ion so this position of equilibrium for the magnesium sits more to the left whereas for the copper your equilibrium sits more to the right this copper metal doesn't lose its electrons quite as easy we can then start to give a number to these difference in the positions of equilibrium because if we just talk about left and right it's very difficult to be able to compare different metals we want to give a numerical value so that we can look at the extent of an equilibrium and this is where we introduce these electrode potentials so this is essentially measuring with a voltmeter so and a voltage to show this difference and the position of equilibrium and it is measured essentially as how negative the metal is due to these electrons being lost compared to how positive the solution is due to the electrons being formed and this voltage this potential difference that is formed between the metal negativity and the positivity of the solution is known as the absolute potential difference now there is some problems with this unfortunately it is not possible to actually measure this this is not a value that we can physically and directly measure in order to do this the solution would need to be connected to the voltmeter using another metal we can't just put a crocodile clip into the solution and hope that it works we would need to use another metal as an electrode the problem is if we use another metal this will create its own potential difference so it will interfere with the one that we're trying to measure and you can't use the same metal because then you'll be measuring the potential difference between the two metals themselves and not the actual solution so it's not possible for us to measure this electro absolute potential difference what we do have the capability to do though is to create a reference electrode in other words we create a standard and from this particular standard we measure the difference for each of our different metals and that then allows us to rank the metals in terms of how easily they lose their electrons in other words where their position of equilibrium 6 relative to the standard so a way to think about this is relative atomic mass we set carbon 12 as our standard and we measure the masses relative to to the carbon-12 we are doing a similar thing here except we're not using carbon-12 we're using this special reference electrode and everything is measured relative to this electrode and the electrode that we use is known as the standard hydrogen electrode you need to be able to draw this diagram you need to be able to identify each of the different parts of it and be able to talk about the conditions because these are very very important so we have hydrogen gas being used at 100 kilopascals and that is pumped in through into some sort of delivery ship we have a platinum wire that acts as an electrode and we use platinum because of course it is very unreactive we then have some platinum foil at the bottom which is covered in porous platinum and that is so that we have a large surface area for this reaction to take place and then we have hydrochloric acid which is our source of hydrogen ions at one mole per decimeter cubed and the whole thing is carried out at 298 kelvin so our key things here in terms of the standard conditions are the temperature the pressure and the concentration you are expected to be able to replicate this and then exam so when we've got this standard hydrogen electrode with the gas bubbling over into the platinum foil we're now building up an equilibrium and we're getting an equilibrium between the hydrogen ions that are in the solution which are of course aqueous gaining an electron to form the hydrogen gas so the hydrogen gas that is being pumped in is setting up this equilibrium on the surface of this porous platinum with the hydrogen ions and the solution and this equilibrium is very quickly established this is our reference electrode and this would give you a voltage or what we call an e naught value of zero volts so this is our standard reference that we measure everything else from we can then look at a specific electro potential for an ion and a metal system by connecting the standard hydrogen electrode on one side with a metal electrode and the metal solution on the other side and we have a circuit going around the outside to connect the two things together using a high resistance voltmeter that's very important which we'll come back to and we also need a salt bridge and that's generally a concentrated solution of potassium nitrate the two components so the standard hydrogen electrodes and the metal electrodes each of these are known as half cells and together they will combine to form what is known as an electrochemical cell and we will go into a lot more detail about electrochemical cells and how they work and what we use them for in the second video in this topic we want to focus at the moment about measuring these standard electrode potentials so our setup looks something like this we have our standard hydrogen electrode on the left hand side and this is the left hand side electrode that's how we would refer to it and on the right hand side our right electrode is the thing that we are measuring in this case magnesium and you can see that we have our salt bridge connecting the two substances and we always use any gases at 100 kilopascals our temperature is always 298 kelvin and the solution that is the ion is always at one mole per decimeter cubed these are our standard conditions now technically the salt bridge can contain any ionic salt but the ions shouldn't interfere with any of the components in the half cell so the reason that we use potassium nitrate is that all potassium salts and all nitrate salts are completely soluble so if they have any sort of reaction with the ions that are in solution they are not going to form any precipitates we don't want that we want there to be a salt solution flowing and a salt bridge is essentially there to connect the two parts of the electrochemical cell together we need a complete circuit in order to allow the charge to potentially flow or to be able to get this potential difference and it can be as simple as a piece of filter paper soaked in the electrolyte or it could be a glass tube stoppered with cotton wool that has got the electrolyte inside it the main thing is that it connects the two sides together the other part is the high resistance voltmeter now this is required to stop the flow of current we do not want the current or the movement of electrons to flow from one metal to the other that is not the point here what we're actually trying to measure remember is the difference in terms of the negativity because of that buildup of electrons to the positivity of the solution so if we allow the electrons to flow from one side to the other that's going to affect that measurement so we don't want the electrons to flow so we use a very high resistance voltmeter that stops this it allows us to measure a voltage but it means that we're measuring a voltage without current flowing which is an unusual concept for us to think about but this is essentially what we're trying to measure here so if we go back to what's actually happening here if we look at the magnesium we said that we're going to get this equilibrium existing so we're going to have the magnesium two plus plus the two electrons giving magnesium metal and we know for hydrogen we're going to have another equilibrium so this time i've balanced out the electrons here just so that we can be comparing what we need to understand is that we've got this difference and the extent of the position of equilibrium so when we get this equilibrium happening we've got some electrons being left on the electrode some of the positive ions being in solution so for example we're going to have say some of the hydrogens here and we will have some electrons here for the magnesium because this position of equilibrium sits quite far to the right what we want to basically see is do we have more electrons on the electrode and more positive ions in the solution so we are looking at to what extent do these equilibriums exist and you will find that one of the electrodes has more electrons on it either the metal you're measuring or the platinum because it has more positive ions formed in other words the position of equilibrium sits more to the left and when we get this we then get a difference a charge difference on the electrodes and it's that charge difference that is being measured by the voltmeter so this this is why we don't want the electrons to flow we want the electrons to build up on the electrodes so that we can look at what is the difference in the amount of charge between these two and that value that we get on this voltmeter is known as the standard electrode potential and this standard electrode potential has the symbol e naught which is the e with the small sign as the superscript and it will contain a sine it will either be positive or it will be negative depending on its position of equilibrium with respect to the standard hydrogen electrode so the more positive the value the more that position of the equilibrium sits more to the right meaning it sits more towards the metal the more negative the value the more it pushes towards the ions therefore we're going to have the position of equilibrium shifting left and it is all in reference to where the elect the hydrogen sits so we can write out these standard hydrogen standard electrode potentials with this relevant half cell so for example if we stick with our magnesium and our copper we write out the half cell equation and we give the value so we have the e naught value for magnesium being minus 2.37 volts telling you that this position of equilibrium sits much more to the left compared to the hydrogen because it is a much more negative number meaning that the build-up of electrons on the magnesium is much higher in comparison to the buildup of the electrons on the platinum electrode so this one has more electrons meaning it is the more negative electrode in comparison with copper copper is a positive value plus 0.34 that tells you that that position of equilibrium sits much more to the right compared to the hydrogen meaning this is the more positive electrode in other words this has a lot less electrons building up on the surface in comparison to the hydrogen and these are a couple more examples of some e naught values and you can actually find a full list of the e naught values and your data book you are not expected to memorize them they are all given to you and the question are in the data book depending on how they want you to use them and they the exam it will become very clear which ones they want you to use and we can then use these values to compare the position of equilibrium between the metal and the metal ion so we're looking at these positions of equilibrium in terms of the the number so we're starting to quantify it we are not simply talking about left and right anymore so if we take these standard electrode potentials as we've mentioned the hydrogen is zero meaning it sits right in the center because the magnesium is more negative that position of equilibrium sits much further to the left meaning the magnesium releases the electrons more easily than the hydrogen which is what we said just a few minutes ago that that negative sign tells us that we're sitting much more to the left so we have the magnesium ions being made more we're getting a build up of charge on that electrode the comparison is copper copper being positive it sits much more to the right so it releases the electrons less easily than the hydrogen we get a less buildup of electrons on that copper electrodes therefore we see that positive value which is essentially what this is saying so this key point down here that these values the more negative the value tells us that it releases electrons more negatively so we can summarize that and two bullet points now everything that we've been working up to here when we measure these metals against the standard hydrogen electrode we get this e naught value the more negative the e naught value the further the equilibrium lies to the left i.e the more readily the metal loses electrons to form ions the more positive that you not value the further the equilibrium lies to the right in other words the less readily the metal loses electrons so our examples each time have been magnesium prefers to exist as the magnesium ion whereas the copper prefers to exist as the copper metal because this was a negative value which was minus 2.37 volts whereas this was plus 0.34 volts now as we've previously mentioned the standard hydrogen electrodes when it is connected with the metal to give us the standard electrode potential it's measured when there are no electrons flowing through the external circuit that's why we use that high resistance voltmeter and actually what we call this is we give it a specific name and we call this the electromotive force or the emf of a cell and when we're comparing two metals to each other which we will talk about in another video we can be figuring out the emf or the overall voltage difference of the entire cell and it is given this symbol e naught cell and emf values can be positive or negative depending on the the reference potential it's measured against so if we have the hydrogen electrode if it is the positive then the emf will be negative and vice versa this is something we'll come back to later on when we're talking about electrochemical cells in more detail the key thing here is that when we have the standard electro potential being measured we have no electrons flowing and this is called emf now anybody that does physics may talk about emf in a slightly different way in physics we have a slightly different definition so in chemistry this is the definition that we use so now we can take all of this information and we can put it together and we can now define what we mean by our standard electrode potential so this is the standard electro potential of a half cell as the emf of a cell because remember we have no current flowing so the emf of a cell containing a half cell connected to the standard hydrogen electrode and the conditions of 298 kelvin 100 kilopascals and solution concentrations of one mole per decimeter cubed apply so this is an explanation of that diagram that we were previously looking at we can get some more slightly complicated redox systems and we just need to be aware of how these are formed if we want to measure the standard electrical potential of a gaseous non-metal such as chlorine we obviously can't use a metal because it has to be the equilibrium between the atom and the iron so this one for the chlorine would be set up very similar to the hydrogen the only difference is that we would use chlorine gas and our solution would have to have chlorine chloride ions in it so we might use sodium chloride for example so we just need to make sure that we are aware that we can have these standard electro potentials of substances that are not gases we can also look at when we have the both substances in the equilibrium and solution so for example bromine to bromide those are both liquids fe3 plus to fe2 plus is another one and when we do this again it's similar to the hydrogen electrode we just don't have the gas pumping in but we use a platinum wire and we use that platinum electrode to allow this um potential to be measured and we just put both of the solutions and both of the ions into solution each at one mole per decimeter cubed and we use the platinum with the porous platinum as our surface in order to carry this out so now that we have these standard electrode potentials of lots of different substances metals non-metals whichever it is that we want to look at we can now start to order them and we build up this thing called the electrochemical series and you'll notice it's very similar to the reactivity series that you've previously met back in gcse the only difference is that now we're actually quantifying it so we have this series being built up in order of their standard electrode potentials you have the most negative values at the top and you have the most positive values at the bottom and you can see that every single half cell reaction is written as the reduction reaction and this is the table that you can find in your data book and this is the one that you will use throughout your exam we can then from this electrochemical series identify oxidizing and reducing agents everything that is on the right hand side of the half cell reaction is a reducing agent because they can lose electrons if you remember our reducing agent is oxidized itself because it's going to reduce something else so it loses its own electrons to be able to give those to something else for it to undergo reduction so everything on the right hand side everything is being oxidized that could go to lose electrons as a reducing agent and these values actually tell us something about how powerful these reducing our oxidizing agents are the most powerful reducing agent in the table is the lithium because it has the most negative e naught value meaning its equilibrium sits far to the left i.e it is more likely to have its possession of equilibrium sit all the way to here it is very happy to lose those electrons in comparison the fluoride ion is less likely to lose its electrons so we have something like this that position of equilibrium does not set very far to the the left it doesn't want to lose those electrons so that's a very very positive value so that's the least powerful reducing agent when we look at the other side of the half cell equations the left-hand side are the oxidizing agents because they are the substances that are reduced itself so it is reduced by gaining electrons because it takes the electrons from something that is losing it and the most powerful oxidizing agent in the table this time is the fluiding so again if we think about that equation like that because this is such a positive value it means that position of equilibrium sits so far to the right hand side in comparison to the lithium ion where that equilibrium does not sit to the right hand side it doesn't want to become lithium metal it prefers to remain as lithium ions so that is the least powerful oxidizing agents so when we then build up this information from the electrochemical series we can start to give this more information so we have our oxidized form on the left hand side of the reversible reaction sign the reduced form on the right hand side and we can then identify most powerful oxidizing and reducing agents and you just need to remember top right bottom left so the top right hand side substances are the most powerful reducing agents the bottom left-hand side of the electrochemical series are the most powerful oxidizing agents and this is things that you actually already know you've already discussed back in year 12 the oxidizing power of bromine chlorine and fluorine when we talked about group seven so you've actually met this before all we're doing is just backing up using these values that we get in our electrochemical series so again now we can make a summary when we figure out these standard electrode potentials when they're connected to the hydrogen electrode the more negative the e naught value the more the equilibrium position lies to the left the more readily the species loses its electrons and therefore the more powerful the reducing agent for example our lithium where we have the li plus this is minus 3.03 volts meaning that position of equilibrium sits so far to the left causing that species on the right to lose its electrons meaning it's an extremely powerful reducing agent if we look at the other side when we're looking at the more positive value so our fluidine for example that position of equilibrium lies so much further to the right hand side meaning that species on the left is going to lose its electrons very easily therefore it is the most powerful oxidizing agent so so far what we've done is how do we measure standard electric potentials what do they mean what is their hydrogen electrode and then what do these values tell us in terms of oxidizing and reducing agents in the next video we will talk about how we can use these values to determine feasibility of a particular reaction and how they link into entropy and the equilibrium constant but as normal let's finish off with some past paper questions so these are from the june 2018 past papers remember these are from the old specification um rather than the new specification but they get very similar questions and the new specification so standard electrode potentials are used to predict the thermodynamic feasibility of chemical reactions that's what we will talk about later in the second video of this particular topic standard electrical potentials can be determined by using the standard hydrogen electrode as a reference standard label the diagram indicating the essential conditions that have to be used so we've got our platinum electrode and we have two labels so if we start with the label on the right hand side we can see that that's where the gas is coming in so we need hydrogen gas and we must specify that this is at 100 kilopascals they will also allow you to see one atmosphere but it's better to put the kilopascals because that is the proper term that we would be using we should be talking about pressures and kilopascals or the temperature and kelvin so the temperature isn't given we're not asked about that because that's not where the arrow is pointed of course we know that that's 298 kelvin on the other side that the arrow is pointing towards the solution so we need to have a source of hydrogen ions so we have our h plus aqueous and we could even specify that this is hcl or h2so4 or whichever acid but we can keep it nice and simple and just say we need a source of hydrogen ions and this must be at one mole per decimeter cubed you get one mark for each of the two labels so you have to have the substance ie the h plus or the h2 plus the condition that goes along with it the platinum electrode that is in this standard hydrogen electrode is coated and finely divided platinum which is known as platinum black we want to suggest why platinum is used as the electrode and then secondly why is coated in platinum black this platinum black is essentially the other words for the porous platinum that we saw in our diagrams so this is two marks and we're being asked two different questions so first of all why is platinum used as the electrode well platinum is chemically inert in other words the platinum is not going to react with whatever it is that you're going to react with whichever is that you have set up so the platinum is chemically inert it's also a good thing that it conducts electricity but that's not quite enough or we could also say actually that platinum acts as a catalyst for this hydrogen reaction happening where it breaks down into the the ions but the main thing is that it is inert and why do we coat it in this finely divided platinum black well we do that because that increases the surface area and by increasing that surface area it allows more active sites and allows this equilibrium or this catalysis to be much more efficient but we can just keep it nice and simple again it increases the surface area part three we want to state the value assigned to the electrode potential for the standard hydrogen electrode well this would be the easiest question in the entire paper it would just be zero volts that is always what we set the standard hydrogen electrode as and then part four explain why a reference electrode is needed to measure the electrode potential of chemical systems well that is because it is not possible to measure the absolute potential difference and remember that absolute potential difference means it is between the metal and its solution it is not possible to measure that voltage being built up between the metal and the solution because we can only use one single electrode we cannot have anything else and if we don't have that reference it means our circuit is not complete we're not going to get any build up so anything like that would give you the the mark but the main thing is that it's not possible to measure that potential difference and you can see there's the mark scheme for those different questions those are sort of the start of the the redox topic the very basic part you'll probably find that these questions come from a question that could be sort of 15 marks where it could link in redox equilibria potentially with transition metals is these two things tend to go hand in hand that's everything for the first part of the redox equilibrium the standard electrode potentials as i mentioned we will be talking about thermodynamic feasibility and electrochemical cells and the second part of this topic if you've got any questions please do feel free to leave a comment below and we hope to see you back on the channel