[Music] hello and welcome to um this video this is the AQA group s the halogens video um this is for uh uh revision purposes so it's just a quick overview of the group seven topic um my name is Chris Harris and I'm from allery tutors.com um and basically like I say we are going to have a quick look at this this is very specific to AQA um the um PowerPoint that I'm going to use here you can use uh and you can actually purchase from um from by clicking the link in the description box for this video so you just click on that you'll be able to get a hold of it and you can use it for revision and you can use it for whatever you want print them off or whatever so um yeah so basically if you if you like what you see here then you can get them for yourself okay um also like I say these are tailored specifically to the AQA specification and so obviously the points are reflected there for for this topic okay so um we're going to start with the halogens obviously because this is this is all about group seven um and obviously the halogens are made up of non-metals they're on the right hand side of the periodic table um and the top of the hogen is Florine which is a p pale yellow gas and it has that electron configuration 1 S2 2 S2 2p5 makees sure you can give the electron configurations of these you should be able to do it for um the elements such as Florine and chlorine in particular uh chlorine gas is pale green um it's cl2 and there's the electric configuration for it bromine is a browny orange liquid and iodine is a gray solid um you see the electric configuration there is much much longer for iodine but it's just to kind of give you an idea okay now generally as we go down the group starting from Florine down to iodine the uh the boiling points increase as we go down the group it's because the um the molecules and the size of the atoms and the molecules here well these These are molecules uh and the relative masses of them um are getting larger so therefore we have larger Vander viles forces um and so this means that you need a higher amount of energy to um turn these into a gas now you can see some of these are obviously in different states the top two are gases bromines are liquid and iodines are solid so um that basically just explain means that the the Vander Val's forces are increasing in that room temperature um the as we go down the group they become more solid the electr negativity decreases as well so when we Bond these halogens to say like a carbon for example or another element uh in a calent bond um then the electro negativity decreases now the reason why is because the ability for one of these atoms to attract electrons towards itself weakens um the atoms get larger um and the distance between the positive nucleus and the bonding electrons increases um and basically we have more shielding so we have more shells between the nucleus and the alter electron so that means it's easier to um um the well the the electro negativity is a lot less um for that for that reason okay right so let's have a look at some displacement reactions now these are pretty important you do see a lot of these for halogens and it they pretty much explain the reactivity that we've described before uh and the fundamental rule is that a mo more reactive or the most reactive should say more reactive halogens will displace a less reactive halide ion um so basically like we said before the reactivity in hens they decrease as we go down group seven okay and for a reaction to occur um basically an electron is gained and we form that halion um atoms with a smaller radius they attract electrons better than the larger ones so that's quite important uh basically how hens are less oxidizing as we go down the group so remember oxidation is the loss of electrons so um so they become less oxidizing as we go down the group and we can show this by reacting halogens with the halide ions okay so halogens will displace a halide from a solution if the halide is lower in the periodic table so we're going to look at these two here okay we going look at all of these in fact and we're going to look at the the reactivity of the hallogen the reactivity of the hiide ion so remember as we go down the group the halogens the reactivity of them decreases okay so let's look at the um first one now because these are two chlorines chlorine water cl2 and chlorine uh KCl chloride sorry which is KCl potassium chloride and because they're just as reactive we get no reaction this one though we do chlorine is more reactive than a bromide uh ion so we do get a displacement reaction you can see here there's the chlorine there's the bromide ion these react and we get a displacement so we get chloride ions and bromine is made so we get an orange solution being formed here and that orange solution tells us that actually we're making a bromine solution here okay looking at the next one again chlorine is more reactive than iodine or than iodide should we say so we do get a displacement reaction now we get a brown solution I when it's in solution goes like a brownie color um so this brown solution is made of iodine chlorine again reacts with the iodide ions there it is there we get displacement chloride ions are made and iodine solution is made an iodine solution is brown remember iodine is a solid is is gray and but your iodine solution is brown okay so look at bromine water no reaction here bromine is less reactive than chloride ions so we get no reaction and no reaction for this one as well CU that's bromine with bromide so they're both just as likely to react but the final one again we form our um Brown solution again it's the same reason uh bromine is more reactive than iodide ions so we get a displacement reaction we form 2 BR minus and we get iodine produced as well and obviously this is in solution so it makes it turn brown uh and the last one um it's no reaction here again chloride is um iodine sorry is less reactive than chlorides we get no reaction iodine is also less reactive than bromides and then obviously iodine and iodide are just as equally reactive so they won't displace either so yeah make sure you know your half equations and you know the colors and you know the reasons why it's all about this displacement okay right bleach okay so it wouldn't be halogens without talk about bleach really let's be honest so um this is made by using a disproportionation reaction okay transy see that one quick um it's basically um a reaction where we're getting oxidation an element is simultaneously ox oxidized and reduced as we'll see in a minute I'll show you the reaction but if we mix chlorine and sodium hydroxide together okay will form sodium chlorate one solution um and this is known as bleach now the reaction like I say as you can see here is the sodium hydroxide plus chlorine literally that's what it is sodium chlorate one is this stuff here NAC okay so that's that one plus you form your salt and you form water as well okay so really important this is your bleach that's bleach there okay so let's look at this disproportionation thing again chlorine is an oxidation state of zero it's an element Chlorine here is plus one because Oxygen's minus two sodium is plus one so if the whole thing to be neutral chlorine must be+ one and chlorine here has an oxidation state of minus one because sodium is+ one so chlorine has been simultaneously reduced and oxidized like I say um and we call this a disproportion reaction uh and we can use bleach uh bleach is pretty useful obviously uh we can use it to treat water uh we can bleach paper and fabric so we can make um uh paper whiter or Fabrics whiter uh and obviously we can use it as a cleaning uh re agent obviously you put bleach Us in bathrooms clean toilets Etc so um yeah it's got a pretty important use okay another important use is obviously water sterilization now without this uh we would have had um uh outbreaks of chalora and actually in some countries today where they don't have um efficient water sterilization techniques they do get unfortunately do get spreads of Cher which is a waterborne disease um but um the whole point of adding chlorine to water is to kill the bacteria in the water um including diseases um including microbes sorry that cause chora so yeah adding water to chlorine will produce chlorate ions CL minus a little bit like the bleach that we've seen just just before these kill bacteria and they're useful in drinking Waters and swimming pools so let's have a look at these again now this is slightly slightly different again we got water and chlorine no sodium hydroxide we're using water instead and what we form is H+ ions CL minus ions and cl minus now you might also see this as HCL and hclo um so these H pluses can be attached to this but because they're soluble and it dissociates um pretty strongly and then we normally write them is separate ions but again if we look at our oxidation state chlorine has an oxidation state zero obviously minus one here and plus one here so again this is a disproportionation reaction the chlorine at the start has been simultaneously reduced and oxidized okay now this is quite important especially in things like swimming pools because sunlight can decompose chlorinated water too um and the problem is is that we get no CL o minus that is made and remember it's the clo minus that actually kills the bacteria so if we look at these two here let's have a look we've got water and we've got chlorine now if we use if we have sunlight that's present actually what we produce is 4 H+ CL minus and O2 obviously this can pose a problem in swimming pools because if you get sunlight um like bleaching into the um into the side of the building maybe is where the swimming pool is um the swimmming the the the sunlight can actually destroy the bleach that's used in swim pools the CL minus uh and what we form is CL minus NO2 now obviously because we don't have any clo minus because the sunlight is broken it up then um obviously we don't have any active ingredient to kill the bacteria so this is why the um chemical reagents or the chemicals that we add to swimming pools have to be replaced quite regularly because sunlight breaks them down into these compounds here which are obviously different to that bleach up there okay drink water okay so advantages and disadvantages obviously it's in the UK and we do chlorinate water um but there are like with anything adding any chemical to anything there's advantages doing it and disadvantages and you need to know what they are so the advantages they destroy microorganisms that cause disease so that's pretty good thing they're longlasting so they reduce bacteria buildup further down the supply um which is pretty useful and reduces the growth of algae that discolors water and give it a bad smell and taste last thing you want come out your TAP is brownish water um that doesn't smell too good it would put you off pretty much so yeah it basically the the chlorine helps to reduce that but the disadvantages chlorine is toxic um we can't hide that um and it can irritate the respiratory system so people who add this chemical um into the water have to take obviously precautions when they adding this in liquid chlorine causes severe chemical burns to the Skin So if there was high levels of chlorine in drinking water that could obviously cause a problems but there's systems and safeguards from doing that but I suppose the biggest one here is the fact that um there is this uh belief by some people that the um are some scientists that the chlorine could react with organic compounds that present in the water and they make chloro alanes and these have been linked um with causing cancer um but obviously the risk of not chlorinating the water um and could lead to what like like we said before Cher epidemic and Cher is lethal um as well so um the basically what we're seeing here is the the advantages outweigh the risks or the benefits outwear the risks in this case the levels of chlorine that are added to water are very very very small and scientists believe that the risk of cancer um from this is obviously very low incredibly low because of the low levels that are added anyway they only add just enough just to make sure that the bacteria is killed in the um in the uh water okay so um right let's look at these hals now these are quite difficult actually these things so um we obviously looked at halogens so far we looked at a little bit the Halal ions with the um with the displacement reactions but halons already have that electron already now these lose that extra electron they've got and so this makes them really good reducing agents okay so remember reducing agents are oxidized themselves okay an oxidation is the loss of electrons so halide ions lose electrons because they've got that extra one and so they make some good reducing agents okay so basically as we go down the group the ionic radius increases and I've just drawn that diagram there basically the distance between the nucleus and the outer electrons becomes larger there's more shielding use all these words in your exam and the attractive Force gets weaker okay so it's not as strong and the outer electron is lost a lot more rly and this is the reason why I minus is more powerful reducing agent than F minus I minus loses that um that electron much more readily than F minus okay now there's two tests you need to know okay um to prove this to prove this trend so one's with sulfuric acid which we'll look at next and the other one is with silver nitrate solution um and you've got to know the results of both of them and I think the most difficult one is the one with sulc acid because you get a lot of equations here so let's just talk it through make sure you know what's going on okay um so some halide ions can reduce concentrated sulfuric acid so that must be concentrated that's very important so what I'm going to show you here is basically um a little kind of it looks a bit like a train map I like to call it um I think it's probably easier maybe um seeing these things in this way rather than in other ways that I've seen in in some um in some notes and sheets and books Etc I think this probably is an easier way of doing it U but you can you can use it however you want basically what I've done is I've written my reduction products of sulfuric acid on the top row the only one which isn't a reduction product is this one sodium hydrogen sulfate and but these ones are reduction products and I'll come on to these in a minute and then what I've done is I've looked at sulfur okay because sulfur is the key element here and I've written the oxidation state of sulfur underneath so we've got plus 6 here plus 4 0 and minus two and what I've done is I've written our helide ions down here these are the helide ions that we're going to be using going to react it with concentrated sulfuric acid so we're going to react to Chloride iion with concentrated sulfuric acid and see what products would produce we're going to react to bromide iion with the same stuff see what products would produce when we react to nine see what products would produce okay so let's start with the first one okay here's our like underground train like train line thing um right so we're going to look at a okay now the good thing is that chlorides bromides and iodides they all produce this compound here sodium hydrogen sulfate and we need to know the reagents of these as well the reactions of these so this is point a now point a this is not a Redux reaction okay this is just a standard chemical reaction between a salt and an acid so here we are this is where sodium chloride sulfuric acid react sodium chloride all we get is one of the hydrogens is given up um to form HCL you get white Misty fumes here uh and you form sodium hydrogen sulfate um you also get the same with sodium broide and sodium iodide as well however as you will see that these reagents here these sodium hydrogen sulfates will then go on to react further so actually you have very little of these anyway um in the overall reaction but certainly with the chloride the only product you get with sodium chloride is that and that's it you can see the arrow stops there no more products produced okay so now we're going to bring our attention onto bromide okay so bromide um is let's have a look okay so these are going to be looking at reaction B so bromides can actually reduce further the sulfur further they can reduce it to produce sulfur dioxide and iodides can do the same as well okay so they could produce sulfur dioxide so let's have a look at their reactions so this is point B okay you can see um in addition to a as well Okay so we've got obviously all the air reactions as well um but with sodium bromide because that's what we're going to look at mainly um we get the same reaction with sodium iodide as well what we do is we write down our half equations first okay um the bromide is being converted into bromine I've written a half equation there so the bromide ions are being oxidized BR minus going to br2 um obviously we balance it with two electrons okay this is being oxidized um I must stress as well if you don't know how to work out your half equations you really do need to know um how to do this I have put a video on that um describes how to balance half equations you've got to know that as well I'm just going to go go through it as an overview here assuming that you know how to do it but um this is the bromine and then crucially this is this bit here we're star with sulfuric acid and we're going to produce sulfur dioxide because that's one of our products what we do is we balance with water first then protons then electrons okay so we just Balan this half equation this is the sulfur being reduced okay so this is a Redux reaction because we got oxidation reduction cancel out all electrons and combine the equations and we should get an overall an equation which is this here so you see the sulfuric acid has been reduced to sulfur dioxide we should see in this reaction an orange Vapor of bromine being produced and that tells us that actually this reaction has occurred okay um so bromine tells us that we have produced sulfur produced bromine and we've also produced sulfur dioxide as well which is a choking gas okay let's look go further so bromides are not as powerful reducing agents so they can't reduce that any further but iodides can okay so they keep going so iodides can reduce sulfur even further to produce sulfur which is this sulfur solid again we're going to do the same thing look balance it out 6 IUS I2 so you're going to put I minus go into I2 uh and then we um basically try to balance that out we do the same here start with sulfuric acid forming sulfur and then we balance with water protons then electrons when we balance this we um effectively worked out that we have to multiply the top one by three so we can balance out the electrons the six electrons remember we've got to try and balance these out cancel them out rewrite the equation the overall equation here is obviously I minus going to I2 and sulfuric acid going to sulfur so producing the solid sulfur and iodine iodide ions are so powerful at reducing they reduce the sulfur even further and they produce H2S okay I should point out with this one as well you produce sulfur solid which is yellow so this one the final step is Step d um again same thing look we're just going from I minus to I2 just balancing out sulfuric acid this time we're forming H2S hydrogen sulfide we call it which is this one here um and the overall ionic equation obviously just like before again we' got to multiply this top one by four to make sure we get eight electrons on both sides um classic one here rotten egg smell H2S is a nerve agent it's really toxic shouldn't be breathing this in and it smells of rotten egg it's got a horrible smell it's also known as sewer gas um because it's found commonly in sewers so you probably get an idea of what H what that what they may smell like okay so that's your reaction with hels with sulfuric acid make sure you can balance your half equations really Redux is crucial here being able to balance them okay the other one which is probably a little bit safer to use let's be honest is reacting it with silver nitrate okay so we can use these halides add uh silver nitrate and we can confirm further with ammonia which will look at later on okay so this is how we do it we test for chlorides bromides and iodides all we do is we add nitric acid first dilute nitric acid then we add silver nitrate solution ag3 the color of the precipitate will help you identify the helide ion okay so I'll come on to the nitric acid a little bit later on as to why we add that but you can see here we've got the test tubes we've got different color test tubes and now this one here is the white test tube white precipitate is formed this is chloride ions Okay so we add silver nitrate and if we get a white precipitate that's a chloridine this is the reaction that happens the half equation ag+ plus CL minus forms agcl the cream one here is a sign of bromide ions cream precipitate formed the cream precipitate is silver bromide that's the actual precipitate and here's the ionic equation ag+ BR minus forms a agbr silver bromide which is a solid and the last one the yellow one is iodines now these form a yellow precipitate um and is the half equation for that as well uh the ionic equation sorry of that ag+ plus I minus forms AGI silver iodide is your yellow precipitate okay the nitric acid bit that we mentioned here this is important that we add this because what this does it reacts with any other anions um that could um and any other halides for example for like we're looking at carbonates um that could precipitate out as well um and when they react with the silver now this could give a false result because obviously we could get um silver and silver um silver ions reacting with any like say like carbonates and we could get precipitates forming and we might think oh we've got a um a chloride there but actually we don't because it's just reacted with an impurity so we just basically add the nitric acid just to mop up these kind of Rogue carbonates that are floating around in the water naturally um we can also test this further as well um and because these are really difficult to see I mean obviously we can see them side by side here we've got a white screen behind here but it's really difficult to tell in isolation so we can do a further test we add ammonia to each one of these Solutions so basically we add um dilute ammonia to each of these and the chloride one should dissolve the precipitate should dissolve with dilute ammonia uh then obviously we know that's got to be chloride so we put that to one side we should have these two left so then if we add concentrated ammonia to these two that's left the concentrated ammonia your silver Brom should dissolve in concentrated ammonia but your silver iodide shouldn't it's insoluble so bromide you get a crepitate dissolves in concentrated ammonia and IDE ions the elate is insoluble so it doesn't dissolve so really important make sure you know these little reactions and you know the colors and what the testing for and obviously the further test with ammonia as well okay um just to find little bit really because obviously we're looking at ions here and I thought it just be good to kind of look at some of these uh um other tests so for example we can test for group two ions and we can use flame tests uh basically and we take the solid sample you put on a nyome wire uh and then you put it into the bunson flame We can spray it on as well um but uh it's a problem if your if your compounds are insoluble obviously we can't dissolve them in solution so we can't spray them but the idea is that like I say you put it in Nyon wire dip it into hydrochloric acid to just clean the wire dip it into the sample place the loop into a blue onon flame CU it must be blue and observe the color the colors that you need to know are calcium is dark red strontium is red color just normal red color and barium is a green color okay so as long as you know these these colors here that's quite important okay uh more test RS uh ammonium compounds and hydroxides basically we can use litmus paper for these so testing for ammonium compounds what we do is we add sodium hydroxide to our ammonium compound or suspected ammonium compound gently heat it if ammonium is present you'll produce ammonia gas and all we do is we use a bit of damp litmus paper red litmus and it'll turn blue if ammonia gas is being produced here's the overall reaction look ammonium ions plus hydroxide ions will produce ammonia plus water okay and testing for hydroxides um basically These Are alkaline and they'll turn red litus blue which is pretty straightforward but it's really difficult because a lot of things turn red litmus blue it's not necessarily a sign of a hydroxide so you got to do further tests to effectively discount the other ones and just to confirm that is a hydroxide okay more tests for ions so carbonates and sulfates using Hydrochloric acid and berium chloride so this is a little bit of mixture of group two as well um so test for carbonates uh dead easy all we do is we're going to add an acid to a carbonate so hydrochloric acid if we think it's a carbonate if we add it to a um a carbonate we should get carbon dioxide gas given off um and then obviously carbon dioxide when the gas given off we can test for carbon dioxide bubble it through lime water and it will turn cloudy okay test of the sulfates um we've seen this already in group two um we just to kind of recap this um so basically if we need to find out if something contains sulfite ions all we do is we add hydrochloric acid it just removes any carbonates in there again we don't want to see them because they could give a false result um and then what we do is we add barium chloride um to the solution and if there is sulfates present we should see a white precipitate forms and this white precipitate is barium sulfate um and obviously as you know as you may know barium sulfate is insoluble so that's why we see this precipitate here's the ionic reaction here barium 2+ plus sulfate ions forms barium sulfate there's your solid look at your state symbols acris acris solid that means precipitate is formed okay and um basically the fin final thing is just make sure you test these things in a particular order just be methodical with your practical technique um and this just prevents any um misconceptions or false positives so the first thing really you should do is test for carbonates first that'll basically tell you if you have got any carbonates in there obviously add your acid um if there is carbon prod if there is carbon dioxide produced brilliant you've just identified a carbonate if there isn't then you need to test for sulfates so basically take your sample add baring sulfate to it um do you get a white precipitate formed if you do fantastic You've Got A sulfate if you don't you go on to the final test and this is basically testing for your halides you add your silver nitrate solution remember um and your nitric acid to mop up any Rogue carbonate ions that may be floating around talides white cream yellow and then add ammonia dilute will dissolve your chlorides concentrated will dissolve your bromide salt and iodides don't dissolve so then test for your ID and basically doing it this way means you can be certain that um or you try and narrow down what compound you've got and that's it so um that's group seven hope that was um pretty useful it's just a very quick overview really of the of of the topic um like I say if you want to purchase these uh PowerPoints that could be really useful for your revision then if you just click on the link in the description box below the video and you'll be able to get them um get them there all right thanks very much bye-bye