hello my name is chris harris and i'm from alley chemistry and welcome to this video on group 17 elements or group 7. this is topic 11 for the cie specifications so this is the um cambridge internationals um um syllabus so this video obviously will go through group 7 or group 17 chemistry and look at the trends and reactions of it so if you're studying the cie subject so then exam board should i say then this uh video is perfect for you um there's the full range of a-level videos on allergy chemistry youtube channel be a massive help if you can hit the subscribe button just to show your support that would be absolutely brilliant also these um these are just powerpoint slides that they're available to purchase which are ideal for your revision purposes and for your supplement and your other revision material that you've got if you click on the link in the description box below it will take you to the right place where you can where you can get a hold of them right so let's make a start then let's look at see what um group 17 has to offer so this is obviously part of the inorganic chemistry topic for year one obviously that is there is an inorganic for year two as well so this looks obviously at the year one stuff um or the s stuff and with regards to the power points that are available purchase i have kind of grouped them together in as an inorganic package anyway so you can have a look anyway so let's have a look at halogens so um halogens um make up group 7 or also known as group 17 i'll probably just call it group 7 from now on um they make up group seven of the periodic table so um the first one is fluorine that's right at the top of group seven in the periodic table and this is a pale yellow gas and it has the electric configuration of 1s2 2s2 2p5 and the next one down is chlorine which has a lovely green gas bromine is a brownie orange liquid at room temperature and pressure and iodine is a great solid so these are kind of main elements within the halogen group now as we go down the group the boiling points increase as we go down the group and this is because we have um an increase in van der waals forces you may also see it as london forces but it's exactly the same they just give it two different names so let's say vanderbilt so this is because this van der waals forces increases due to the increasing size of the relative mass of the atoms that increases the physical state goes from gas at the top of group seven which is obviously fluorine all the way down to solid at the bottom so as we as the obviously as the atom gets bigger the mass gets bigger the intermolecular forces increase and so therefore the physical state becomes increasingly more solid as we go down the group with regards to electronegativity obviously this decreases as we go down the group um so electronegativity is the ability for an atom to attract electrons towards itself in a covenant bond if you can remember that that's from the bonding topic that you will have seen earlier um and the atoms get larger the distance between the positive nucleus and the bonding electrons obviously increases and also there is more shielding as well so the most electronegative element in the periodic table is fluorine so that is incredibly electronegative so say for example when it's bonded to say hydrogen to form hf that is incredibly um electronegative and obviously iodine isn't so the increased shielding just weakens that ability to pull electrons towards itself okay so let's have a look actually some hydrogen halides i mentioned one there before which is hf um so hydrogen halides they are acidic they're definitely acidic um so hydrogen halides they're basically gases and they can dissolve in water to form acidic solutions obviously remember acids are aqueous okay so acids that you see for example that you use in the lab are dissolved in water so they're solutions so they react with water um in the air and they can form white misty fumes and they can dissociate which means basically to break apart you would have seen that before when they're dissolved in water so let's look at these equations here so we've got hcl which is a gas um hydrogen chloride gas and that can dissociate in um in contact with any moisture or water to form h plus ions and cl minus signs and remember it's these ions here that um that are effectively acidic in reality though um h plus what makes something acidic is hydroxonium ions which are h3 or plus that again you would have seen this in a previous topic when we looked at acids and bases so it's this bit that makes it acidic so basically hcl hydrogen chloride gas because there's water vapor in the atmosphere it reacts with the water vapor to form white mr fumes and that is really nasty stuff we definitely don't want to be breathing that in so obviously they produce acids when they're dissolved in water so hydrogen chloride forms hydrochloric acid hydrogen bromide forms hydrobromic acid and hydrogen iodide forms hydroiodic acid so you've got all these different types of acids here so fairly straightforward um hf which is um obviously hydrogen fluoride dissociates weakly so it doesn't really form a very strong acid um obviously hydrogen chloride is probably the more common one that you will have seen um so far and hydrogen halides react with ammonia as well ammonia gas and they make these white fumes of ammonium halides so for example ammonia and hcl sus um obviously hydrogen chloride gas will form ammonium chloride which gives you that white mr fumes that's been formed so again this is a neutralization reaction when you react in ammonia with an acid okay so let's have a look at the stability of some of these hydrogen halides um so obviously hydrogen halide stability um when heated decreases as we go down group seven so this kind of i know in the previous topic topic 10 you looked at group two you'll find there's some parallels for this in terms of the kind of things that the example want you to know which is about trends and stability and you know color and things like this so it's the same with um group two so you'll find it's quite a it's kind of bunched in a similar format so in this case we're looking at hydrogen fluoride and hydrogen chloride and these are quite stable when they're heated so they don't split into h plus and f minus or cl minus ions okay so they're relatively step but don't really break up that much hydrogen bromide will split up partially and hydrogen iodide will split up more easily i mean because that bond the bond between them is is quite um is relatively weak remember that's what we're talking about more about the shielding of these atoms um and obviously as we go down group seven the atom gets you know a good bit larger as we go down and obviously the bonding electrons are much further away from the nucleus and a shielded more so remember when you looked at bonding i think was topic three then we look at um you know what is it actually that holds the atoms together it's the attraction between the shared pair of electrons that form the bond and the nucleus in the middle and obviously if the nucleus is shielded by various shells in between that's going to make that bond a little bit weaker and so therefore it's going to break up much more readily and you have so you can see there obviously the data here to kind of explain that okay so let's have a look at some other reactions let's look at halide ions with um silver nitrate so we can test for halides using silver nitrate and then we can obviously confirm that by doing a further test by adding ammonia solution now you might have done this in the lab again um whereby um you're testing for halide ions or the presence of them and you get these funny kind of colours that appear here so we can test for chloride ions bromide ions and iodides and basically it's a very simple reaction it just involves adding dilute nitric acid so that's hno3 then a bit of silver nitrate solution um and then basically we get a precipitate that forms and the precipitate the color of that precipitate will help you to identify the halide ion so you can see a picture here this is obviously a diagram of what that test may look like very kind of boring colors aren't they're not very good um anyway chloride ions um when if we have a chloride line present in the solution and we're testing for the iron we add our dilute nitric acid and some silver nitrate and if it contains a chloride ion then we should get a white precipitate form and that white precipitate the silver chloride effectively the silver in the solution here so this ag plus here reacts with the chloride ions that are exist that exist in the solution and that forms agcl and that's a solid that's a precipitate that's been formed so now for bromide ions you get a cream precipitate forming so again silver reaction with br minus that will form silver bromide so that's your bromide ions and then finally you've got iodide ions now again these react with the silver ions and it forms silver iodide instead now this is a yellow precipitate that's actually being formed here now obviously we add nitric acid to this as well because you notice there's nitric acid involved and what this will do is react with any excuse me um it'll react with any anions other than the halides so for example any carbonates that might be knocking around in there and this could give a false result so the purpose of the nitric acid is just to try and react with any other ions so we know that if we are getting a positive result we know it's because of a halide ion rather than because of any other anion that's knocking around okay so one of the things you'll probably spot a problem here um one of the things is that that's really difficult to spot and we can just about say that well you might be able to i'm not sure if you can see it clearly you might just be able to see the variation in color if you had them side by side but if you just looked at them in isolation it'd be really difficult to tell the difference between a white precipitate a cream one and a yellow one they're very very similar so what we can do is kind of to prove this we can add another chemical which is ammonia and ammonia can really help to conclusively decide whether it's chloride bromide or iodides that exist so if we add ammonia to each one of these tests if we add it and it contains chloride iron then that white precipitate will dissolve in dilute ammonia first now if we add dilute ammonia to the bromide or the iodide they won't dissolve so that's a classic sign that it's chloride we then with the rest of the other two here the chlorides and iodides if we add concentrated ammonia to these then this will dissolve the bromardine the silver bromide will dissolve the iodide one is insoluble so silver iodide won't dissolve in either dilute or concentrated ammonia so if you add dilute and then concentrated ammonia to this and it doesn't dissolve then you know you've got iodide ions basically so this is a really important kind of additional little test to to try okay um you'll probably notice that we've emitted fluoride ions um quite simply fluoride lines don't actually form a precipitate um they do react they do form silver fluoride but it is soluble so you won't actually see any precipitate at all okay so um basically it's kind of going in order of solubility isn't it so you've got your chloride iron silver chloride dissolves readily with dilute ammonia silver bromide with concentrated ammonia that doesn't dissolve at all um and obviously agf is obviously above fluoride is above chloride ions in the periodic table so this is definitely soluble you need to add any ammonia to get that to go okay so let's have a look at some displacement reactions and it wouldn't be halogens without looking at some of these to be honest so um if we add hexane which you'll see later on in um in the um which topic it is now so i think it's topic 13 i think yes it is topic 13 which is the introduction to organic chemistry so yes organic chemistry then you'll come across um hydrocarbons and you might have seen some of these anyway gcse but anyway we can add hexane to any one of these reactions and it'll be easier to observe these color changes so when undertaking halogen displacement reactions we can add an organic solvent like hexane to see the color change really easily okay so the reactions we just seen there were um the reactions between there were displacement reactions or swapping around so the halogen present will dissolve readily in the organic solvent which forms a layer above the aqueous layer and a coloured band will appear so for example if we have our aqueous years you've got your aqueous layer here okay and then we've got our organic layer now organic solvent is hex is hexane which is on the top layer here so let's have a look at some displacement reactions okay so these are quite common reactions and obviously we can use that hexane to help identify some of these color changes here and it makes it a little bit more pronounced so the idea is or the concept is more reactive halogens will displace less reactive halide ions okay and obviously using hexane adding hex into this will allow us to see these color changes that we're going to see on here a little bit clearer so the reactivity of halogens um decreases um as we go down group seven so as we as we've seen before um so for a reaction to occur an electron is gained so atoms with a smaller atomic radius and attract electrons much more readily than larger ones so if you think about chlorine oh sorry fluorine let's let's say fluorine right at the top of that periodic table throne is a very small atom not much shielding and so remember group sevens will want to accept an electron to try and get a full outer shell which is a little bit more stable so fluorine is going to be much more readily it was going to be more able to accept that electron because it's got less shielding so the attractive forces are stronger but as we go down the group and we get towards say so we go chlorine then bromine and iodine at the bottom ions much a much bigger atom there's a lot more shielding and so therefore its ability to to take an electron is not as powerful so that's where that comes from so halogens are less oxidizing as we go down the group and we can show this by reacting halogens with halide ions okay and we'll look at a bit more of oxidizing oxidizing reactions as well but a halogen will displace halide from solution if the halide is lower in the periodic table okay so we're talking about reactivity here so remember the most reactive is at the top of group seven the least reactive elements are at the bottom so let's have a look at some examples so we've got chlorine and so chlorine water so cl2 reacting with potassium chloride now both are just as reactive as each other so there is no reaction um so basically um in terms of the aqueous solution that we've seen before remember we had your hexane and the aqueous solution the aqueous solution will be colorless and the organic solution will be colorless as well okay so if we take chlorine and react with potassium bromide now here we will get a change because chlorine is more reactive than bromine it is higher up in the periodic table so we will get a displacement reaction so um the aqueous layer will be yellow and the organic layer will be orange and the the kind of reaction i showed you before so chlorine will react with the bromide ions from potassium bromide and that'll form chloride ions cl minus and br2 so effectively the bromine is being oxidized okay and the chlorine is being reduced okay so remember that's gone from minus one to zero and that's gone from zero to minus one so remember this redox reaction so reduction and oxidation so that's mentioned in a previous topic okay so let's have a look at the next one so this is again chlorine reacting with potassium iodide obviously chlorine is more reactive than iodide or iodine the aqueous layer will be brown the organic layer will be purple in that kind of test that showed you before um sometimes you can just and again the example that might just look at the aqueous layer the organic layer might not be there you might not be um required to give that um that color change but the aqueous one you definitely do so you should really know these um chlorine reacts with iodide ions to form cl minus and i2 now iodine um obviously is a brownie color you form this brown solution here that's been formed and again iodide is being oxidized to iodine and chlorine has been reduced to chloride ions so there's a displacement reaction okay so let's have a look at the next one addition of bromine water to kcl um cl minus or cl is more reactive than bromine so you get no displacements obviously bromine and bromine they are just as reactive as each other so there's no reaction there but with the final one there will be a reaction because bromine is more reactive than iodide ions so therefore we do get displacement reaction and the bromine which we started with here will react with the iodide ions to form br minus and i2 so you get iodine that's been formed here and obviously you get that brown solution as well and finally the last ones iodine reacting with kcl obviously iodine is less reactive than kcl so we get no reaction it's also less reactive than bromine and obviously the final one iodine and i iodide sorry and iodine are equally as reactive so you get no reaction there so this is what we call see a displacement reaction so this is where we get the obviously the exchange of of um halogens or the kind of displacement of halogens depending on their reactivity you might have seen this gcse possibly so this is just adding a little bit extra to it this is your kind of equations for it okay so there we are okay so if we look at some uh halogen ions let's have a look a little bit more so most halogens have different oxidation states okay so they don't just have one so let's have a look at some examples here so let's say we've got uh the ion which is cl minus or br minus for example so these will have um a an oxidation state of minus one as you can see on there um there we are okay um chlorine um so cl or br on its own as an oxidation state of zero remember if it's an element it'll be zero chlorate ion clo minus is plus one okay because oxygen um in this case is uh is is minus two but it's obviously gonna over have that negative uh that negative charge that's plus one uh bromate is plus one to bromine and or minus again a lot of these and you'll see bromate three iodide five and i date um i did seven um these you need to be able to know these and if you're not sure you need to look back at topic six which looks at oxidation states so you can work these out using the rules that's in topic six so if you're looking at this thing and i don't really know what all this means go back to topic six have a look at the how these are worked out and look at the rules and then come back to these and then you'll see them so there's nothing really new here this is just looking at in context of halogen ions and fluorine only forms oxidation states of zero obviously if it's an element or minus one if it's f for example if it's in a compound um and they can't form any positive oxidation states at all and flowing obviously apart from zero okay so just make sure you're aware of that obviously get these different types of reagents here okay so um let's have a look at some of these funny reactions now i kind of mentioned some of these i think in a previous topic as well but very briefly so this is where we kind of go through it a little bit more detail so this is called again this was in topic six i think um disproportionation reactions so this is in relation in particular to halogens so halogens react with cold alkalis in disproportionation reactions so they've got to be cold alkalis so here we are so we've got x2 and x will represent a halogen you know could be any halogen react that with sodium hydroxide and we form your um your sodium um with your chlorate so might be a chlorate for example it might be a cla bromate or anything like that you form your sodium halide here and you form water as well okay now a disproportionation there's your kind of ionic equation as well so remember the ionic equations you've obviously got hydroxides you remove the spectator ions which in this case will be sodium and you just concentrate on the ions which matter so a disproportionation reaction is where we have a simultaneous reduction and oxidation process of the same element so in this case we're going to look at halogens so the halogen here starts off with an oxidation state of zero because it's an element um and then here uh the halogen here has obviously been oxidized so it's gone from um zero to plus one here because it's bonded with an oxygen so it must be plus one um but it's also been reduced here because it's gone to minus one at this point here so this has obviously been simultaneously oxidized and reduced so let's have a look at an example here so bromine reacting with sodium hydroxide will form sodium bromate which is naobr sodium sodium bromide and water so this is a an example of a disproportionation but it's basically where the halogen here is simultaneously oxidized and reduced so it's come from that same halogen there okay right so halogens they can react in hot alkalis as well in disproportionation reactions so let's have a look so again x represents your halogen um and that loves to react with sodium hydroxide as you can see there and it'll form um your your halide compounds as you can see there's our ionic equation and again there's our reactions there so we've got zero obviously there that goes to plus five in this case so remember these reactions here and that's minus one at this point so you can see it's been simultaneously oxidized and reduced so obviously forming a different substance within hot alkalis so a classic example um is forming sodium chlorate five solution which is this here um and obviously this shows the overall reaction using chlorine instead okay so this is where we're now going to look at heal iodines separately so we've looked at halogens merely and we've looked at the reactions of halogens and using sodium hydroxide we've looked at displacement reactions um as well so this is where we're now just going to focus on halide ions and say right well how can these kind of be used so here iodines um so these um lose an electron in reactions okay so obviously they have the extra electron already but they're losing electron reactions so they are effectively reducing agents again these terms are in topic six so if you're unsure of this please go back to topic six and have a look at that so reducing agents remember oxidation is the loss of electrons reduction is the gain so reducing agents um lose electrons okay so reducing agents um anything that loses an electron is classed as a reducing agent so and halides do that quite well hail iodine should say so as we go down the group the ionic radius increases so obviously we're adding an extra shell as we go down the group so the radius will increase at that point so the distance between the nucleus and the outer shell becomes obviously greater so it increases as we go down there's more shielding between the nucleus and the outer shell but also the distance between the nucleus and the outer shell is also increased and obviously this force means it's weaker so what that means is as we go down the group the electrons are lost much more readily and this is why um iodines are iodide sorry i minor so much more powerful reducing agents than fluoride ions so iodide ions can lose that electron much more readily than fluoride ions can okay so this is going to be important when we're looking at reactions and we're looking at reducing a substance so it might be reducing a chemical and we'll see that in a moment with sulfuric acid um where we reduce using halide ions okay so there's two tests to prove this as i mentioned before um so we've got um the reaction with sulfuric acid um so we can see we can have a look at that um in a moment and then the reaction with silver nitrate solution as well and in fact we've seen the silver nitrate solution reaction already um and i'm sure we've seen that with um the the white the cream and the yellow precipitate before we know kind of how to test for that and then adding ammonia to see which one dissolves the next slide we're going to look at here is going to be a similar reaction so it's a way in which we can prove this but using sulfuric acid now this reaction is going to be quite um this is going to be quite tricky this bit so hope you're following some of the equations look really technical as well they're not really um effectively you need to know how to form a half equation of your own and again that's going back to topic six as a big emphasis for topic six for this link the link between topic six and this um but going back there will show you how to write these half equations and just ensure that you're writing them correctly and if you don't know how to do that please go back and look at this because i'm going to flash up some half equations here and you're going to think oh my god do i have to know all of them equations the answer is yes but you can kind of form them yourself you can kind of work them out work by yourself okay so this is where we're going to have um i suppose reacting halide ions with sulfuric acid and some halide ions they can reduce sulfuric acid now this is a very nasty um reaction there's some really horrible products being produced here as you probably wouldn't do this in a lab it's probably unlikely so it's gonna be quite difficult for you to kind of visualize this but we'll try so i'm gonna do this as a bit of a a bit of a diagram i suppose to try and illustrate this because i think there's a lot of reactions here and i think if we kind of spell out as a diagram it makes it a bit easier so here what you can see on the left here i've got our hyladines here what chloride bromide and iodides all the way down the bottom here and along the top here i've got some reduction products now we're starting with sulfuric acid which is h2so4 and sulfuric acid um some of these are some of the products of reactions that we can that we can achieve now this isn't a reduction product and but these three are and you'll see they'll kind of come they'll come back into their own in a minute okay so here's your oxidation state of sulfur in each of these products here so you've got plus six for sulfur here plus four zero and minus two so you can see these numbers the oxidation numbers of these products are decreasing so this one's obviously the lowest and that one's the highest okay and here's arrows right so what we're going to do i'm going to try and break this down as simply as i possibly can so we're going to start by saying right if i had a beaker or test tube and i had some chloride ions in there some bromide ions and some iodides and into each three of these i'm going to add some sulfuric acid into them okay let's have a look and see what types of products could be made from each of them and so we're going to start with obviously point a first which is at the top here so rather test tube with chloride ions added some sulfuric acid what i would say is sodium hydrogen sulfate being produced and hcl so i'll see some white misty fumes being produced with the chloride ions and also see the same reaction as well with bromide ions and iodides so i'll get all of them okay so remember as we go down the group this becomes um this is a much more powerful reducing agent because the atom or the the ion sorry is bigger than chloride iron so this is more likely to be this is this is a more powerful reducing agent rather than chloride so effectively with sodium chloride uh sorry with sulfuric acid and and chloride ions we don't actually get any reduction happening whatsoever so we do get a reaction but this is not a reduction reaction because sulfur here is plus six and it's plus six here so there's no reduction but it's just a reaction okay so let's have a look at step b then so now we're going to focus on br minus and i minus because clearly these do reduce these are powerful enough to reduce sulfuric acid so in addition to a um with sodium bromide and obviously it's the same with sodium iodide as well you've got your two half equations so effectively the br minus is being oxidized to br2 and your sulfur in sulfuric acid which is here h2so4 is being reduced to so2 so it's gone from plus six to plus four okay um now these are the half equations that i said that you need to look at how you kind of construct these from topic six so please have a look there i'll not go through it too much detail here because there is another video that goes through that and obviously if we cancel out these electrons we get an overall ionic equation of what's happening here so sulfuric acid um reacting with br minus will form br2 so we get bromine vapor being produced sulfur dioxide and water now sulfur dioxide is acidic so that's not very good and this isn't really good for you either so really not very nasty products being put nasty products being produced here okay so let's go a little bit further and let's look at iodide so iodide's more powerful reducing agents so it can actually reduce even further so it can go even further so at point c so in addition to a and b then with sodium iodide we actually get iodide ions being oxidized to form i2 okay and then the sulfuric acid gets reduced to form sulfur okay so you get sulfur that's been formed here so that's more reduced than sulfur dioxide so that's further reduced so the overall equation is this here now what you notice is you get iodine you'll see iodine being formed you might see some vapor like a purple vapor but you'll get a yellow solid being produced and this is sulfur and that stinks of rotten eggs so it doesn't have a very nice smell whatsoever but it doesn't stop there iodide ions are really are that powerful as a reducing agent they reduce sulfur even further and this is point d here and they actually form um another compound which is called um hydrogen sulfide h2s and this is a rotten egg smell and it's toxic it's a nerve agent um it's also known as sewer gas so you really this isn't very pleasant whatsoever as you can imagine but the reaction again sulfuric acid reacting with iodide ions will form iodine and h2s as well okay so basically this is just showing you the power of reducing agents so iodide is the most powerful reducing agent and it reduces sulfuric acid to so2 s and h2s here all three of these plus all of them will produce sodium sodium hydrogen sulfate as well this is not a reduction product but this is just a product of reaction okay so hopefully that makes sense okay because this can be quite tricky but again i urge you to look back at topic six if you're not comfortable with how we form these reactions here okay so let's have a look at bleach okay so after all that smelliness and rotten egg smell let's get something that's a little bit cleaner shall we um now we've seen some reactions already with disproportionation reactions where we react them with sodium hydroxide and we get your disproportionation reactions formed bleach is actually a product of disproportionation so if we mix chlorine and sodium hydroxide we will form sodium chlorate one solution and that's known as bleach quite simply so here's the reaction so you've got sodium hydroxide reacting with chlorine gas that'll form your sodium chlorate solution which is there sodium chloride and water is being produced so this has an oxidation state of zero okay because it's an element this has an oxidation state of plus one and this has an oxidation state of minus one so you can see here that chlorine has been simultaneously oxidized here and reduced here and this is an example of a disproportionation reaction now obviously chlorine uh sorry bleach sodium chlorate and one which is this is obviously really useful for treating water so obviously tap water to keep it free of um diseases and a bleaching paper and fabric so obviously white paper has been bleached um using this and obviously cleaning agents obviously domestic bleach um that you can use or industrial bleach if you want to um you know to clean toilets and um and sinks and showers and things like that so obviously you know what bleach is so this is how you make it basically okay so in line with that we need to kind of have a little bit of understanding of water sterilization as well now obviously bleach is used to sterilize water as seen before and it's used to obviously kill bacteria so if we add water to chlorine it'll produce chlorate ions which are cl minus clo minus ions and these will kill the bacteria and obviously they're useful in drinking water and swimming pools as well obviously to a higher concentration of course so here you can see you've got um chlorine has actually been simultaneously oxidized and reduced here so um you've got this obviously got an oxidation state of zero so that should be point there and that's minus one and that's plus one so this is basically showing you what the reaction is when we add chlorine to water and we form these products here so it's that under very dilute solutions that can be added into tap water mainly fluoride based as well you know we add fluoride based compounds to be honest um but also mainly in swimming pools so this is used a lot in swimming pools as you would expect um and then chloric one acid so that's clo minus ionizers okay to um to make the chlorate ions which is the hypochlorite ion so this hclo that's produced here this reacts with more water to form cll minus signs and h3o plus you've got to be careful with this this remember is what makes it acidic so um you've got to be really careful with what's been produced here and because obviously if you go into pool water it is slightly acidic but because there's such a vast amount of water in there it heavily dilutes it down but obviously um chlorine in swimming pools can irritate the eye because of um obviously irritation from these chemicals here but because the sword dilutes um it doesn't actually you know cause too many problems which is thankful but there's enough in there to kill off you know the vast majority of bacteria which is what you want because you don't know you spin around with the peter's bacteria idea so um looking at the i suppose the kind of final use of halogen based compounds and one of them is pvc polyvinyl chloride and it's used everywhere around the home so pvc is just a polymer and you'll see this later on i've seen year two which contains chlorine it has um tough properties which make it ideal around the home okay so obviously let's have a look at some of these ideas here so um you can see you've got a halogen here there's your chlorine um and it's in it you'll see a little bit more this to be honest and and the polymerization topic later on in year one and in year two and but this is just to give you an idea this is a halogen attached to a um hydrocarbon as you can see here now it's what we call a repeat unit so basically this repeats multiple times in a chain to form a polymer and this polymer is called pvc so it's used in electrical coverings because it's it's really resistant to electricity just as well because you don't want to touch one of them switches and it's gonna um you know it's going to electrocute you it's the last thing you want um but also it's hard wearing um and it's obviously ideal to make in window frames as well so halogens are kind of mixed with other substances to form these products um but also it's got good heat stability as well which makes it ideal for outdoor use like guttering plumbing pipes etc so these might carry hot water or they may be exposed to the high heat like we've had over the past um i suppose over the past month we've had some extreme heat haven't we in in the uk um well everywhere to be honest obviously i live in the uk so that's um that's why i mentioned that but you need these plastics to try and be resistant to that so um pvc is pretty good for that as well so that's it so that's everything to do with group seven or group seventeen um like i say there is a full range of ci a level videos on allergy chemistry youtube channel please subscribe to the channel and show your support that'll be fantastic um these um powerpoints as well are available for purchase from the test shop if you click on the link in the description box you'll be able to get a hold of them there and like i said i've bundled the inorganic topics together for year one and year two but you can buy the the kind of the um uh bundles for year one and year two chemistry as well so you'd have to buy them kind of separately but yeah so i hope that's helpful bye bye