okay so we're going to continue with this electrochemistry chapter and we're gonna start by looking at an electrolytic cell in an electrolytic cell a non spontaneous reaction is driven by an external power supply so now you're actually going up in terms of free energy right it's non it's not spontaneous and so you've got to put power energy in to roll the rock uphill right so a couple of things we can do with that first of all you can just like I said take a non spontaneous reaction something that wouldn't occur naturally so you sell in that case would be negative this could be the actual cell potential or the standard cell potential but whatever it is it's negative it's not spontaneous but we can force it to occur if I apply a greater potential in the opposite direction and that can be from a battery maybe I have a 9-volt battery I'm using or maybe I'm using a power supply that I plug into the wall but in some way I'm generating a potential and that potential is greater than this potential is and in the opposite direction and greater is gonna be a key word there if I put the exact same potentially opposite direction nothing's gonna happen I'm gonna have to put an over potential in order to get it to actually run backwards I could also do the same thing with if I have a spontaneous reaction I can force it to run in the non spontaneous direction if again I apply a greater potential in the opposite direction this would be an example of this would be recharging batteries right I can't recharge every battery because if I just have a strip and it's eroding away on one side and plating on the other if I reverse it the reactions will go backwards but it's not gonna plate on in the same shape that it eroded away and so what's gonna happen is my electrodes are gonna change shapes and eventually they're gonna break and fall apart and not make contact anymore so I can't just recharge any battery but we have batteries that we've designed specifically so the electrodes are not plating or eroding there's something on their surface that's coming off or going on and so the electrodes themselves are maintaining their same shape so it's a spontaneous a galvanic cell when you're discharging the battery when you're using it and then when you go to recharge it you're applying a greater potential the charger that's plugged into the wall in the opposite direction so it's making it run in the non-spontaneous direction and recharging the batteries so we can do that an example of this one would be like maybe electroplating maybe I'm gonna put gold plating or chrome plating on something I can do that electrochemically by using an electrolytic cell okay at this point I want to add a brief note here I said greater potential rate or potential and that's the over potential so let's say this is the the negative cell potential right this is what I need to overcome isn't my theoretical minimum amount of energy to have to put in to get it to run backwards well that's not gonna fly if I put that potential in in the opposite direction nothing's gonna happen right so I have to dial up the potential I'm I'm giving with my power supply and at some point the reaction will actually start occurring in the reverse direction or in the non-spontaneous direction so that's the potential I actually need this is a theoretical minimum that I could get I could have and that difference between those two that is the over potential all right so it's how much more I have to put in over the theoretical value and in order to actually get it to go in the reverse direction now kind of like a activation energy this could be small it could be a small barrier or this over tential could be huge it could be a huge barrier so usually the over potential is just kind of a pain for us because we got to put in more energy than theoretically we should to overcome the over potential but sometimes every once in a while it is kind of a nice thing every once in a while I have two different reactions maybe I have two different things in the same solution I want one to react I want the other one not to they may have the same potential and so it doesn't seem like I would be able to get one to react and not the other but if they have different over potentials I might be able to find a potential in between the two that actually causes one to react but not the other so every once in a while over potential can be used for our advantage but usually it's just zap stuff that's of energy okay so now let's see how this is going to what gonna happen what's gonna change if I run a galvanic cell in Reverse and make it now on electrolytic cell what's gonna change okay so if I take a galvanic cell and run in Reverse now it's gonna be an electrolytic cell because it's being powered by a power supply you can imagine that if everything is going spontaneously in one direction everything's gonna reverse and go in the opposite direction that's true of most things but not everything we got to be careful so all arrows will switch direction so if the electrons were going one way now they're gonna be any other way the cations are going one way now they'll be going the other way if I had an electrode that was plating now it's gonna be eroding right so all of those direction all those arrows will switch directions because where I had previously oxidation I now have reduction and vice versa the anode and cathode will also switch sides so almost everything switches but but but but there's a couple things that switch locations but don't switch sides and that is the source and sink of electrons and the corresponding positive and negative terminals for negative bills with source positives to the sink terminals they are now going to move to the power supply they're knocking previously in a galvanic cell the source and sink were the electrodes and the positive negative and positive terminals were the electrodes now the driving force is not the spontaneous reactions occurring on the electrodes the driving force is the power supply so it becomes the source and sink of electrons and it becomes the positive and negative terminals it turns out that those are going to be on the same side as they were before in our diagram they're just going to move up to the power supply instead of being written next to the electrodes they'll be look written next to the power supply and how we're going to show our power supply in our diagram is with this symbol right here this symbol can be written this way or flipped and written the other way depending on which way we want the electrons to flow but the short fat side is the source of electrons it is the negative terminal so electrons are flowing out of that side and the tall skinny side is the sink of electrons so it's the positive terminal and trongs are flowing him there so electrons are flowing out of here and electrons are flowing in to there okay so that's that's gonna be the driving force so it's the source and the negative terminal it's the sink in the positive terminal for some reason I always think of Don Quixote and Sancho Panza when I see this so I don't know if that helps at all the short back all right it's just going okay so we're gonna do the same cell we did as a galvanic cell except now we're gonna do it as an electrolytic cell and I'm gonna put my power supply in there in my circuit and we're gonna see what happens here now this is a little different than in a galvanic cells in a galvanic cell the first place we go is the standard table of standard reduction potentials to find out which direction it's going to spontaneously go but this isn't going to go in the spontaneous direction it's going to go in whatever direction the power supply makes it go which is usually the non spontaneous direction so instead of starting here we're gonna start here this is where we start so I know that in my power supply the electrons are gonna be flowing out this way and in this way right based on the symbols I drew it in the opposite direction as I did here right but I can draw whichever direction I want it and so whichever direction I'm gonna apply the potential in it so I'm gonna start up there I'm gonna start that this is my source of electrons and this is my sink of electrons and this is gonna be the negative terminal of my power supply and this is gonna be the positive terminal my power supply so instead of writing them down here by the electrodes I'm writing them up here so I know the directions of my electrons are flowing this is flowing this way this is flowing this way now that's gonna tell me what my reactions are occurring so this is gaining electrons ger it's gotta be reduction so this is gonna be the reduction half reaction and so reduction is going to be zinc ion is going to zinc metal so zinc two plus plus two electrons which I'm getting from the wire it's gonna become zinc metal this has a reduction potential of negative 0.76 volts all right now this one I'm losing electrons from here so this is going to be loss of electrons is going to be oxidation so the copper is going to be eroding like this so I have copper going to copper to ions plus two electrons which are going into the wire right and so this e nought ox for this one is going to be negative zero point three four volts so combining those two I end up with the same potential but now it's negative right negative 1.10 volts assuming these are both 1 molar solutions if they're not I just have to use the Nernst equation but let's say they are so my power supply would have to be greater than one point one volts in that direction depending on what the over potential is all right and then now let's look at our ions in the salt bridge so now my positive charges are going this way right and my negative charges are going this way so cations going that way and ions going that way all right now what else I have my source and my sink in my terminals I still need to do my hand out in cathode so those still are the electrodes because those are defined as the electrode where oxidations occurring in the electrode where reduction is occurring so this electrode is still the anode or is the anode now and this is the cathode because that's where reduction is occurring out this forcing sinker now the power supplies the terminal or now the power supplies but eroding and plating this is now eroding and this one is now pleading okay so if we bring back the one we had originally this is the galvanic and this is the electrolytic okay so notice that this direction switched this direction switched these directions switched these directions switched the anode and the cathode switched sides and the plating and eroding switched sides but notice that I before I had the source and the negative terminal on this side I still have the source of the negative terminal on this side except that they're no longer the electrode they're now the power supply and I had the sink and positive terminal on this side I still have the sink and positive terminal on this side but they're now the power supply and not the electrode because this is what's driving in here so it's the source and sink and terminals in the spontaneous direction this the reactions occurring here spontaneously is what was driving it so those electrodes were the source and sink and positive and negative terminals all right so that's an electrolytic cell let's do another example where we're gonna split some water electrolytically okay so the electrolysis of water this is our second example of an electrolytic cell and what I'm gonna do is I don't know look at these two reactions right here that I have my pen caps pointing towards and normally this reaction would go in the forward direction in the reduction direction and if they were paired up this reaction would go in the reverse direction right that's the spontaneous direction so water and hydrogen is oxygen hydrogen ions would react to form water hydroxide and hydrogen gas would react to form water and that kind of makes sense right because if you have oxygen and hydrogen how would those react with each other well they would combust and make water H+ and OH H - those would neutralize and make waters to that that the spontaneous direction totally makes sense but I'm gonna force it to run in the opposite direction and split water into oxygen and hydrogen so I'm going to take this reaction and instead of running it forwards I'm gonna run it backwards so that's what I've done right here and now that's an oxidation so I've got to change the potential from positive one point three volts one point to three volts to negative one point to three volts okay and this one I'm gonna run in the four direction even though us a non-spontaneous direction for those two and I'm gonna put that there and its potential will just be the same as it says on the chart right the reduction potential I do have to multiply this top reaction through by two because I have four electrons lost so I'm gonna need to gain four electrons as well so I'm gonna multiply that through by two notice when I did switch the directions I did change the sign but when I multiplied through by two I did not multiply this through by two and the reason for that is again potential is an intensive property so I'm not if I double the amount I don't double the potential okay so now when I bring those down what I'm getting is six waters are gonna split into two hydrogen's and an oxygen and then I'm also gonna get four protons and four hydroxides out of that so it's gonna also release acid and base now this makes a lot of sense so when I actually do this I'm gonna get twice as many hydrogen bubbles as I'm getting oxygen bubbles which makes sense because water has twice as much hydrogen in it that it has oxygen so I'm gonna get more hydrogen gas out of it than I am AUSA Djinn gas but interestingly I'm gonna get the same amount of protons there H+ ions that I do hydroxide ions so these would just neutralize each other and go back to making more water again but I would get hydrogen gas and oxygen gas so let's look at the demo set up for this okay so this is my setup here I have a petri dish and in the petri dish I have an electrolyte solution I have a sodium sulfate solution and I also have some universal indicator in there it's starting out neutral so the universal indicator starts out green I also have two platinum electrodes in there Platinum is expensive so these are very thin platinum foils but I have those as my inert electrode so it's neither gonna plate nor erode on the surface it's just going to react on the surface and then I have that hooked up to a power supply which I just use a 9-volt battery so if I look at here my potentials my cell potential is negative two point zero six volts that's a negative it's non-spontaneous I'm gonna need more than two volts one six zero six volts in the other direction to make it split water and again I just use okay you know an excess I choose a nine volt battery and that's gonna make water split like crazy so that's a good thing I am also careful about what I chose as my salt bridge here and this direction is I'm forcing in the oxidation direction I chose sulfate because that's actually higher up on the list and it's even less likely to get oxidized to go in the oxidation direction so water will split before sulfate reacts so that's a good thing I don't want my salt bridge to be what's reacting on the electrodes I want my reaction to be happening on the electrodes and the sodium ion one is lower where is it right there lower than this one so this is going forward so this is gonna prefer to go forward over this one's even lower on the list all right so I thoughts why I chose those particular cations and anions all right now if I hook up my power supply in this case I can hook it up either way and it doesn't matter cuz it's just gonna you the reactions are gonna happen on this side and the other reaction on this side or if I hook it up the other way they'll switch not a big deal but if I do have it hooked up this way and this is just the 9-volt battery right this would be the negative terminal of the battery this would be the positive terminal of the battery that means that this is the source of electrons this is the sink of electrons and electrons are flowing in this direction and in this direction okay so here on this side we're going to have lots of electrons so this is going to be oxidation so that is going to be my anode on this side so on my anode I have this reaction right here so oxidation lots of electrons and so water is going to be touching the Platinum electrode it's going to be coming off of that after it gives up electrons to the wire it's going to be coming off as oxygen gas bubbles so I'll have some bubbles of oxygen gas and I'll also have some hydrogen ions so the universal indicator on this side it's gonna start off green but on this side where I'm making oxygen it's gonna start turning red because you eye is red and an acidic solution and then the side is a little bit more acidic on the other side I'm gonna have the cathode okay on the cathode side I'm going to have the reduction reaction so remember to multiply that through by two to make the number of electrons match and gaining electrons thoughts reduction on the cathode side I'm going to be producing twice as much hydrogen gas as I am oxygen gas so I have a lot more bubbles on this side I'm off but I'm producing an equal amount of base as I produced acid so on this side you eye is gonna go from green to purple I didn't have a purple pen so I just worth having blue but it's gonna go from green to purple so this side turns purple this side turns red and the interesting thing is when I do this demo is it looks like there's a lot more purple then there is red the red kind of hangs out right around the electrode the purple kind of spreads out through the petri dish much more so than the red does so it looks like there's more purple but what's really happening is that you just have twice as many bubbles forming so there's more convection from the rising bubbles on this side and so that spreads out you know agitates the solution more and spreads out the purple the hydroxide further and so it looks like this side is totally beating this side but then the cool thing is at the end when I disconnect the 9-volt battery I swirl it and it goes if first it gives all kinds of weird ugly colors but then it eventually goes back to green because I created a stoichiometric amount of hydrogen ions and hydroxide ions so the exactly new to realize each other and I go back to being green in terms of UI right another cool thing about this is notice I didn't have to do this in two different half cells I didn't have to do it in two different containers with a salt bridge in between I could do it in the same container and the whole solution is the salt bridge because it's not a spontaneous process so I don't have to get be worried about getting kicked out of the middleman position because I'm not I'm not taking advantage of this it's I'm actually putting energy in to drive it so I don't have to worry about you know getting cut out as the middleman as I would in a galvanic cell because it wouldn't happen by itself anyway it's a non-spontaneous I'm driving it with my power supply so also when we electroplate something we can do it in one solution we don't have to split it into two separate half cells because again we're not trying to prevent a spontaneous reaction from happening we're forcing a non spontaneous okay so let's go on okay so another electrolytic example electroplating this is something you would have done in lab if you'd have been able to this semester but we were going to pleat copper onto a brass key and we're gonna do that electrolytically so what we wanted to happen is we wanted copper to plate onto the key so that is a reduction going from the ion to the metal this would be plating on to the key right and then since that's a reduction that would be the cathode on the other side we want we're gonna put a strip of copper in there and we're gonna have that erode so we're gonna go from copper metal to copper ions and the solution was already gonna have copper sulfate in it as our electrolyte so copper would be plating on this copper ions would be plating into copper on this side and more copper ions would be coming into solution replacing those copper ions that were plating out from the copper strip so this copper strip would be the anode because it's oxidizing and it would be eroding you'll be getting thinner and thinner and then we're gonna measure the difference in mass of the key and the and the and the copper strip to see what would happen now when we were hooking that up to the power supply we had be careful which way we hook it up doesn't want the copper to plate here if you hook it up backwards it's gonna play here so let's make sure we've done that right so this would be the source of electrons or the negative terminal and so electrons would be flowing in to the key which is good because we need them as a reactant and then here this would be the sink of electrons and this would be the positive terminal up here the power supply and that's good because we're losing electrons here so we want those viewing out and into the wire okay so that would be all set up the other thing we're gonna do in that is a little stoichiometry we wanted to calculate based on the current and the time how much copper should have plated onto the key and how much copper should have eroded from the anode from the piece of copper so let's set up those calculations and I'm gonna give you a couple definitions first okay we already know that a volt is a Joule per Coulomb right so how much energy per unit of charge and that's a measure of potential now current current is measured in amps and that's coulombs per second so how much charge is flowing in a second right if I multiply those two together the coulombs cancel and I get joules per second which is power which is watts but you dude you don't need to know that but you need to know that what these two are in order to make your units work out okay so in this example we ran at first a half an hour and we set it to a certain current so we had a current system a constant current system so we could measure the current that was flowing and then the voltage didn't really matter as long as it was greater great enough to make the the non spontaneous process happen all right so now what if I wanted to figure out a theoretical amount that should have played it and should have eroded well I can do that and what I'm gonna do is I'm gonna start with one of these two numbers I don't want to start with this because that's actually a conversion factor coulombs and seconds so I want to save that as a conversion factor what I want to start with is my time I want to start with the time because if I played for a longer more is gonna play I played for shorter less is gonna play even if the current stays the same and so let's start there so I'm gonna start with 30 minutes we'll put like a decimal in there all right I'm gonna have to go four minutes two seconds so I can use this so that's not a problem so one minute is 60 seconds okay and then I can now use this as a conversion factor so that point 3o goes with the coulombs and it's per second so zero point three zero coulombs for every second and then now what am I gonna do I'm in coulombs what else do I know that has coulombs in it and maybe moles something I needed to mole somehow oh look at that Faraday's constant ninety six thousand four eighty five point three four coulombs per mole electrons why'd you get two moles of electrons from there so let's see I want the coulombs on the bottom so ninety six thousand four eighty five point three four coulombs is one mole of electrons that's so bad not so bad alright now I have mm C two moles of electrons for every mole of copper one mole of Cu and now the last thing I need is the molar mass of copper so I'm just gonna go to the periodic table and find that what a terrible chemist I am actually had to go get a periodic table my should have that tattooed on me somewhere alright so anyway this is the the the atomic mass of copper and so now I'm gonna go from moles to grams and now I have grams of copper all right so now see I have two sig figs there that's infinite two sig figs I got a lot of sig figs they're infinite four sig figs okay so it looks like two sig figs is it so my answer is 0.18 grams of copper and that's the amount that should play on the side theoretically and that's the amount that should erode from that side theoretically should be the same amount all right let's go on to our last topics okay our next topic is metal ores and the corrosion of free metals so if we love looking for metals in nature most metals are not found as they're free metal or the elemental form in nature and there's a reason for that so we have an oxidizing atmosphere here on earth we didn't always the for photosynthesis was actually reducing atmosphere but now we've got plenty of oxygen and if I look at where oxygen is on the table of reduction potentials it's way up here near the top most metals are below that so oxygen is spontaneously going to get reduced to water and whatever metal is going to get oxidized up to the cation right that's the spontaneous direction so once we had oxygen available in the atmosphere our metals were in the form of cations and it's really interesting we can actually tell what those distances first happened because there's a band of iron oxide in the sediment layer and so when photosynthesis occurred and oxygen start being produced all the iron precipitated out of the ocean as an oxide and then form this layer so we can tell the sediment when that happened which is really interesting but there are a few metals that are above this reaction though for example Gold is above that reaction and so Gold wants to be reduced more than oxygen does so gold will stay in its elemental form in nature even with the presence of oxygen and that's why we call it one of the noble metals so you do find gold in its elemental form but most metals we find in their cations and those cations are going to be combined and ionic compounds that we call metal ores and these ores are usually oxides or sulphides as the anion the metals the cation the oxide of the sulfide is the anion so aluminum oxide or iron 3 oxide or titanium dioxide right so we have all these these are some common ores that we that we mined for so then what we have to do is you have to take that ore and we have to isolate the components that we want and then we have to reduce it down to the elemental form so we mine the ore is out of the ground we process them and then we reduce the cation to its elemental form the metal form and we call that reduction smelting and I like that word is just a silly smelting it's not melting it's some melting this is actually a huge user of energy and cause of pollution a lot of sulphuric oxide sulfur oxides go in the atmosphere as a result of this process okay but now we've got the metal and we want that we want the metal because it's got the properties we want right metal metals are malleable they're ductile they conduct electricity they have all these great properties that we want could we use the ore to build a skyscraper no because ionic compounds are brittle the skyscraper which is crumble and shear and that would be terrible so the metal though has the useful properties but what's gonna happen as that metal is exposed to the atmosphere well I forced it this way to my great effort it's just gonna go back that way because that's the spontaneous direction so that's what corrosion is corrosion is just the spontaneous oxidizing of the metal back into its cationic form okay so we've invested a lot of time and money in an energy and caused a lot of pollution reducing these metals we want to keep them reduced for our purposes so we spend a lot of time trying to prevent the corrosion of metals we prevent iron from rusting we prevent copper from getting that patina you know so on and so forth one way we can do that is really clever and it's called cathodic protection or also called the use of a sacrificial anode all right so what you do is you have an object that's metal and you want to protect it see it's a dam or a bridge or a or something you don't want it to corrode so what I'm gonna do is I'm going to physically attach a metal that has a lower reduction potential so something lower on the table then the metal I'm trying to protect now so for example let's say I have iron and I want the iron ship and I wanted to I don't want it to rust through I could take a metal lower than that and attach it to it I don't have to directly attach it to it as long as it's attached with a conductor that's good enough but so for example I could take zinc that's lower than iron the zinc would get oxidized and the iron would stay reduced or I could take magnesium the magnesium look at oxidized and the iron would stay reduced so what's gonna happen is because the metal that's lower on the table has a lower reduction potential it has a higher oxidation potential so it's gonna be preferentially oxidized so that makes it be anode so the metal that were sacrificing is the anode the sacrificial anode the other metal that it's attached to is going to be the cathode and that's going to be protected because the the anode is going to get oxidized instead and the cathode is going to remain in the metallic form right so this is a very good idea so here we have a few slides and in this first slide we have an iron pipe and we're gonna it's underground it's buried in the ground it's gonna be you know prone to rusting so as we go through this the reduction reactions that's this one the oxygen being reduced down the water and the oxidation reactions are gonna be these ones here iron going to iron to and then iron to going up to iron three okay and then we have one more reaction in there where iron three and water react to form rust which is iron three oxide monohydrate and you get a little bit of acid out of that to you which then accelerates the corrosion alright so what would be happening is that the iron is both the and the cathode in this situation when you have corrosion right in one part the oxygen is getting reduced so that's the cathode in another part the iron is oxidizing so that's the anode and this whole thing is connected by the ion which is conductor and the damp soil is the salt bridge right you ever noticed that when you drive somewhere where they salt the roads your car rusts faster or if you live by the coast I say Hawaii your car was faster because you have a salt bridge that salt spray on your vehicle in this case the damp soil is gonna act as the salt bridge all right now what if we attach a piece of magnesium well now the reduction reaction is going to be the same and it's still going to be occurring on the iron pipe so the iron pipe is still the cathode but now what's happening instead of the iron pipe oxidizing and also being made I know the magnesium is gonna oxidize instead and that's gonna be the anode and we can see here some examples of sacrificial magnesium anodes ready to go ready to be attached at something and you have the wire coming off to connect it to whatever you're trying to protect and here we have an example of a piece of magnesium a sacrificial anode that was dug up out of the ground and you can see how it's pitted and eroded away as it was oxidized to protect whatever it was attached to so eventually the sacrificial anode will completely be oxidized and you'll have to replace it every so often and here are some other magnesium sacrificial anodes in various forms some are going to be connected to offshore oil rigs a little be welded onto the side others or bracelets and they will be connected around pipelines but they all work the same way you can even have say a bag of magnesium pellets and have a wire in there and connect water to say underground propane tank now when we when we use a we usually use magnesium when we're on terrestrial applications because the wet ground is not a great conductor it's not a great salt bridge so we want something as far down as as realistically possible from what we're trying to protect in those situations because the ground doesn't conduct very well and you know we do want to protect the object now when you're using a ship in salt water salt water is a fantastic salt bridge and so you can use something that's not quite as far down the list that's a little cheaper so you can use oftentimes they'll use zinc for that application instead of magnesium it's closer to hyoeun but it works well in that application they're also now trying to use alloys of aluminium as sacrificial anodes on ships as well okay another way we can prevent the corrosion of metals is by putting a coating on it so there are plenty of manmade coatings we use paints we use plastics all kinds of things paints on our bridges and plastics on the inside of our cans of food all kinds of things like that also galvanization is a form of coating that we use can organization is putting zinc on the surface of steel and so what that does is it's actually a physical barrier it's a coating just like painter plastic would be that prevents the steel from getting oxidized but the cool thing about using zinc is that even if you scratch it through the zinc and expose the metal the zinc is below iron on the table and so even if scratch it'll still act as a sacrificial anode so that's kind of a double bonus of galvanization some metals actually have self forming coatings though like if you've ever noticed a piece of aluminum never looks rusty or patinaed or tarnished so what's up with that is aluminum just not gonna get oxidized is it not reactive aluminum is way down here look how far below the oxygen reaction it is aluminum super reactive in fact almost instantaneously it forms an oxide layer on the surface magnesium does the same thing the difference though is that these oxide layers happen to adhere very well to the surface and they are impermeable to oxygen and water so it's a layer that's so thin is there but it's so thin that you can't really see it but it protects the layers underneath from getting oxidized because it's impermeable to water and oxygen so that's why you'll see magnesium rims or you know aluminum well that looks shiny and doesn't think we have to be protected because it's protecting itself other coatings that form though are not impermeable so like for example rust on iron the iron 3 oxide monohydrate it is flaky and it is porous and so it doesn't protect the metal underneath and corrosion continues all the way through but these to form their own coatings that protect them which is interesting when we do the lab where we make our own table of reduction potentials you have people have trouble with these too because it's really hard to make good contact with the metal with the probe tip because of that oxide layer being impermeable okay so coatings are another way we can do it and then finally we have alloys he's so an alloy if you recall is a homogeneous mixture of a metal and other elements the other elements can be metals or nonmetals but the key is that the mixture has to retain metallic properties it still has to behave like a metal and we have various types of alloys we have a substitutional alloys we have interstitial alloys but that's that's another topic but the interesting thing about this is that they are mixtures they are mixtures meaning I don't have to have fixed proportions I can mix them in any ratio I want but the difference between an alloy and a normal mixture in a normal mixture each component retains its own properties so if I mix sand and sugar together this still sand and the sugar is still sugar they retain their individual properties but alloys are different because an alloy a metal has a sea of delocalized valence electrons so the electrons are shared by all of the atoms and so when I start putting impurities in there that changes the sea of electrons and it changes the property of the alloys so alloys have different properties than their constituent elements do including the reduction potential so an alloy will be in a different place on the table compared to this oxygen reaction then a pure metal would be so an example of an alloy would be stainless steel so steel is already an alloy it's iron is the main element but we put some carbon in there and the carbon is an interstitial alloy it fits in between the iron atoms because carbon is pretty small and that makes it harder so it's not as easy to shift its shape and if you put too much carbon in there it makes it brittle so but we put some carbon and that's what makes it steel and then we can make it stainless steel by also adding some chromium and some nickel and some other elements and so we can change the properties by doing that and that and that's pretty that's a pretty cool thing to do one one of my favorite alloys I like to tell this story is is the alloy that they used for the Sacagawea dollar that golden colored alloy and they had a lot of trouble finding one that would work the trouble was that they used to have susan b anthony dollars but they looked like a quarter and so people would accidentally spend a dollar when they minutes to spend a quarter and so they wanted to make a new coin that was a dollar coin that was visually distinct so I had a golden color he would not ever confuse it with a quarter but the problem was vending machines because vending machines identify the coin not by the size or shape but by its conductance of electricity and so if you're gonna make a new coin it's got to have the same conductivity as the Susan being a penny dollar did because otherwise you'd have to replace the Tronics um you know millions of vending machines and so they through trial-and-error found an alloy eventually that had the same conductivity as the other coin but had a different color and that's kind of a just picking what you're gonna mix in in what quantities all right so that's another way we can protect metals is by using alloys of metals instead of the pure metal and that brings us to the end of this chapter so next we're going to look at coordinate covalent complexes