okay so this video is uh going to run through the topic of intermolecular forces um it's a topic from unit one or chem one on the a specification so AQA um it comes as part of the bonding topic um there is a shapes molecules video this one's going to purely just look at intermolecular forces uh what they are how they work within molecules uh then it's going to end up looking at some uh some exam questions as to kind of how you can apply your understanding um in an exam situation so first off uh intermolecular forces so kind of really what does this turn intermolecular forces or inter mular Force actually mean well with anything if you can break the term down and you can work out what is actually what it means then sort of you're halfway there you've you've almost won the battle so interm forces the term inter um that actually is talking about between so between in the places you might see that so an inter city train is a train that travels between cities not within it's a between and that's very very important here term molecular is really kind of Fairly self-explanatory we're talking just really molecules with reference to a molecule so these are forces then that are acting between molecules not within the within Force um within molecule force uh within molecular Force I should say uh things like your calent bonding your Ion bonding and your uh your metallic bonding we are talking about the forces that are actually made between the molecules and I can't stress that enough really really can't so um what are these forces then well there's three as that you need to know um and it's important that you certainly do know what these forces actually are you know the name of them uh and I'm going to start off with what the name of names of them are uh before going into detail about each one and as I said then ending up with some exam questions sort of pulling together all of the ideas is so the forces are uh we have ferales forces uh you could call them V DW an abbreviation we have uh dipole dipole and we have hydrogen bonding so they are are three forces there are some different names for these forces you could stick the term permanent in front of our dipole dipole here uh and these could be classed as induced dipole dipole um but I'm going to stick with Vander viles dipo dipole just as a general term there is a separate one and hydrogen bonding there as as a different one um there again if you look through a mark scheme you'll see there are all sorts of other names as well U I'm not going to go through those other names these are the terms I'm going to use use U these are the terms I'm happy with and that I think you know they work fine if you can stick with these in the exam you're absolutely you're absolutely laughing uh the important bit here oh not the straightest line in the world the important thing here is that these forces are are not equal um particularly in the way they work within a molecule particular the strength of the forces the force strength is not equal we find that and I've done this in in order like this um on purpose starting here at Van deval's forces this is our weakest Force and this up here is our strongest force now not these are not a strong even at hydrogen bond here they're not as strong as our proper bonds these are inter lecular forces they are much weaker they have an effect they have a huge effect and stuff that you have done that you do all throughout um the a level and at gcsc and you know key3 and all the rest they have an effect they're very important but they are not as strong as the proper bonds if you like so ionic coent in the metallic bonds right so that's generally kind of a bit about interal forces and the three forces particularly this strength that we're talking here that's very very important I'm going to go through these now in order breaking them down into sort of more how they work and looking in a little bit more detail about them giving some examples and all the rest starting then with uh van deval's forces so van deval's forces let's see what color I want to use let's go with that one so van to valves so this is the weakest intermolecular force as you saw from the previous slid so these are the weakest that you're going to deal with in uh in the as uh unit one what actually are though well Vander Val's forces exist in all molecules okay that's that's the key thing they exist in all molecules are the weakest intermolecular force but they are there in all molecules whereas the other two aren't necessarily going to be there the way they work is that if you imagine you have a molecule and we'll imagine just say we've got uh let's just say we've got iodine we'll stick with iodine so the iodine molecule I2 that is made of an iodine atom bonded to another iodine atom via calent bonds now if we were to think about actually kind of what we find we are we've got here is that around these atoms within this molecule we have electrons so there are electrons within the molecule and these electrons are constantly moving they're not static they don't sit in these delightful little orbitals um just sitting there waiting for reactions to take place they are they are U they are moving they are active and they're going all over the place but the key thing is at any one point there are likely to be electrons more electrons in one place in the molecule than there are in the others so for example if I talk about sort of what um electron density I might find if I were to look around this molecule Maybe I would have something like this and what I'm implying here is that I have less electrons here and more electrons here so the electrons have moved within the molecule and I have a greater weight of electrons on this side than this side and what that does is that creates a an induced dipole so I can say this side because there are less electrons it's therefore I guess Delta positive and this side would therefore be Delta negative as a result now the key thing is these electrons are moving so no sooner as this become apparent then it might flip and it might go to the other side where we've got lots here and then a lesser number here and again that switches so this now becomes Delta negative and this now becomes Delta positive the key thing is in a group when you've got lots of molecules together this random switching is happening all the time so one molecule becomes Delta negative there and Delta positive there which is attracted therefore to this next molecule it's Delta negative and Delta positive and that's constantly switching around there are constant attractions occurring constantly switching so where there was an attraction then there's not an attraction where there was there's not and all the rest but the overall effect is that this uh these temporary dipoles cause basically an overall an overall attraction and that's really quite important that's really quite key so we have an overall R attraction you don't to worry too much about sort of the theory behind this but it kind of helps to explain it and gives you a bit of a bit of a grounding as to sort of what's happening so the overall switching U of where the electrons are within the molecule means that they may therefore be attracted to other molecules due to the uh opposing opposing charges so Delta positive would be attracted this Delta negative here which would be attracted to another molecules Delta positive and all the rest going through this whole huge molecule and they all switch and they go for it again the important thing is this force uh this Vander Val's Force it its strength changes depending on the size of the molecule ultimately the number of electrons that are present so bigger molecule equals more electrons and this leads to Stronger vaves forces and the opposite is therefore true a smaller molecule has less electrons therefore has a weaker Vander vales forces and and you can see this particularly if we look at um a graph of something like the noble gases in particular so these guys here we start with helium which is the smallest going up to radium sorry raidon which is the the biggest um and as you can see um the size increases we get more electrons we get bigger Atomic radi and all the rest and the boiling point and melting points here each than these lines being representing the boiling and melting point they increase because the forces get bigger and all really we're talking about with melting and boiling points is we are separating the forces between the molecules and that's very very important so melting and boiling point is really the the energy required to separate um the molecules okay melting we're talking about sort of I guess a partial separation uh boiling we're looking at a total separation where we are we obviously tur into a vapor into a gas U where the forces between them are are relatively relatively low but this fits exactly what I've just said bigger molecule and this case obviously these are atoms but again more electrons therefore stronger Vander vales forces going from small relatively small to relatively big there across or down the group but uh going from Helium there through to radon so that's really quite important this is one way that you could see this you could certainly be asked an exam to describe this trend and explain this trend um and all you would need to do is say that as as we go down the group from Helium through to radon this size of the atom increases therefore there's more electrons therefore there are stronger Vander vales forces therefore more energy is required to overcome the attractions overcome the forces between the atoms so that's really our Vander Val forces okay that's our Vander vales forces key thing there and whilst it's important that you have understand the Vander valves is that actually within all of these the idea of the melting um and I'm going to put melting SL boiling points it's really really we're looking at the energy required to separate molecules or in terms of obviously what we've done there we we could also say atoms but it's more like hence that being inter molecular forces it's more likely you're going to be dealing with molecules rather than atoms um last bit on Van Devar before I move on to the dipole dipole stuff um and this is a case particularly when we look at things like uh fats are a really good example so straight chain saturated fats versus um unsaturated fats so saturated fats and unsaturated fats U this term saturated and unsaturated applies to the bonding that's present in the chains of the fats I'm not going to go into detail about what the fats look like really key thing is saturated they are all single bonds and we end up with sort of chains that look kind of like this and if we were to draw the next one it would sort of be like this and the next one like this when we look at unsaturated fats we tend to find the chains have sort of these kinks in them uh and as a result they tend to sort of be like this or or this also applies to branched chains I guess so here I'm looking at straight chains um and again Vander Val's forces are coming into into play here if this is one section of my saturated fat or my Straight chain molecule and this is one section of my Branch chain or my in this case the unsaturated fat where there's a kink in it note that when we have our straight chains they can lie very closely together and therefore due to this the vand deval's forces are stronger because the chains are lying very close together when we look at the the branched ones they are not as close to each other and because they're not lying as close to each other here what I find is that my vander's forces are actually weaker as a result so it's not just dependent upon the size of the molecule it's also dependent upon the branching how closely they can lay next to each other in this case branch in causes them to be weaker because that really the forces here less surface area I guess for them to be in contact with and therefore an overall lesser um lesser uh less attraction between the between the chains and and the molecules and the proof really is that you have seen this um saturated fats things like butter for example or lad or coconut oil or something like that all tend to be solid at room temperature okay or pretty solid at room temperature and that's CU they have a high um High proportion of saturated fats within them and because of this uh these straight change which lie close to each other they therefore requires more energy and room temperature doesn't have the energy to actually to overcome those forces and as a result the butter and the L tend to be solid unsaturated fats we're talking our olive oils or General sort of veg oils uh and in this case they tend to be liquid they contain more unsaturated fats the Vander Val forces are weaker therefore the energy provided at room temperature is enough to overcome the forces and therefore we end up with liquids if we cooled these down we would find them solidifying but key thing is at room temperature the comparison between these two is an application and is uh it's a way for us to actually see vander's forces in action particularly the straight Chain versus the the branch chain so that's vavar forces okay that's that's it all it all basically done it's electrons moving around causing an overall attraction and that's that's basically it um going on to the next one we'll have a look at permanent uh dipoles so permanent dipoles what are these well the vavas I mentioned the the term the idea of temporary or induced dipoles due to those uh those movement of electrons well here we have well not temporary it's it's permanent obviously it's it's in the name uh it's permanent dipole dipole so we're looking at situations where that Delta negative Delta positive is always present and the classic one here is something like a hydrogen halide so I'll go for hydrogen uh chloride right there there's my hydrogen chlorine and what I get there is if we look at that in a bit closer then it change color I really like that color um hydrogen is here and my chlorine is here due to the electro negativity difference between these I know hopefully that the electrons are going to lie closer to the chlorine in this double in sorry in this single Bond here my single calent Bond because chlorine is more electronegative than hydrogen therefore they lie closer to here now that's not going to change that's the general sort of pattern there we're looking at that being how it is generally in the bond in terms of the effect that has on the molecule this is always therefore a Delta positive end this is a Delta negative end and you'll come across this idea again in um some of the mechanisms in unit 2 but in terms of the actual intermolecular forces here Delta positive end Delta negative end because of where the electrons lie now if I have to stick another uh molecule here I could put a I'll try and get back to that greeny color another hydrogen chloride there uh put another hydrogen chloride right there again now what we're getting is when we look at the actual uh the individual molecules when we add more in we can see where the attraction causes there's a Delta positive Delta positive Delta positive end this is our Delta negative end and we have attraction we have attraction here between those molecules and so it's a very similar idea to the Vander vales but here we're talking about permanent dipoles um occurring really there uh which is which is obviously quite useful it allows for uh Stronger attraction than we had with um the the vaval so we have stronger forces of attraction but works in a similar way it's an attraction between the Delta positive and Delta negative areas okay therefore this this electrostatic attraction it can be overcome again melting boiling points dependent upon how much energy is required to overcome these uh the attractions between these these different ends here um in terms of diagrams in an exam if you're expected to draw stuff or here's your sort of diagram you could put a little sort of uh the idea of an attraction between these two guys here um but that's basically uh our sort of s the simple uh simple description really simple explanation of this dipole dipole situation okay final one um is hydrogen bonding and hydrogen bonding isn't necessarily any more difficult really there's a little bit more to sort of remember um but it's the general idea is pretty is pretty similar um to the last two we certainly we talking about attractions between opposites and all the rest now hydrogen bonding occurs and it's no it's got this name bonding it is still an intermolecular force it is the strongest of the three um and it crops up all over the place it crops up in biology it's between the base pairs in in the DNA Helix um which we know can be broken because we can separate it and we can carry out DNA replication all the rest but hydrogen bonding has quite a key effect um and comes up certainly in A2 as well as as uh and the questions are all very very roughly the same and it basically revolves around the idea of a molecule containing oxygen nitrogen or Florine and these must be bound to um bonded to hydrogen so as the name implies hydrogen bonding we are dealing with some hydrogen present here somewhere the example that you'll or the one of the classic examples is in water where if I was to draw a water molecule here here's my hydrogen here's my oxygen and here's my other hydrogen H2O water now note the water molecule itself has calent bonds within the molecule but we're looking about between water molecules so I'll stick another another water molecule down here uh and then I'll go for um a another water molecule let's put one here really the greatest positioning but there we go so there's my water molecules done now the key thing is on these water molecules I have um some lone Pairs and it's those lone pairs that are very very important each oxygen has two of them so there's my lone pairs two lone pairs on each one of my oxygen atoms hydrogens have no lone pairs that's fine but what's important if I deal with one water molecule to start with is that oxygen is electr negative okay well that's not how you spell Electro negative let's try that again oxygen is electr negative much better um hydrogen isn't and therefore we have a difference we have a polar bond between the two here where we have the electrons in here much closer to the oxygen creating a Delta positive region and a Delta negative region and what we actually find in in sort of in reality here is that so this is our Delta positive Delta positive Delta positive Delta positive this is tiring Delta positive Delta positive the is the lone pairs of really that create this sort of Delta negative region um and and what really are quite important in this um in this uh process of hydrogen bonding so we have water molecules here we have a Delta positive region uh in my hydrogen and I have this Delta negative region here and I particularly have these lone pairs what we get is we get an attraction between Delta positive hydrogen and the lone pairs of other molecules and that is a high hydrogen bond right there it's the attraction between the hydrogen and the lone pair on the oxygen on the nitrogen or on the Florine that's the important thing all three of these have lone pairs on them and with the bonding with the hydrogen present as well when we throw in other molecules we will have this attraction between the hydrogen and the oxygen this comes up in an exam in exam situations and we'll show show you later on and you'll see the um the kind of things that you are you are shown um the kind of things that you see in terms of the question that well they will expect you to draw something like this it's very important that you include partial charges it's very important that you include the lone pairs as well um regardless of which example we're talking of it's probably going to be water to be honest that they're going to use in an exam but they could use any molecule it's about applying your understanding to it um so that really is is is hydron bonding if we were to have a different example when you might say what about if it was something like uh we use Florine so hydrogen fluoride same concept here Florine very electr negative and here we would have the electrons much closer to the Florine end and so we end up with uh lone pairs on the Florine again lone pairs but that attraction occurring between the two between the hydrogen and the florine's lone pairs there and we could keep going we could have a molecule we could have a big web setup but it's the same concept again we could do it with ammonia we could do it with all sorts of different molecules which there's no point doing it again again it's the same concept lone pairs with hydrogens attraction between them boom you've got yourself a couple of easy marks there in terms of application for this where might you see what where might the questions go well sometimes there are questions relating to uh water um and hydrogen bonding particularly in relation to Ice uh but if you've ever seen an iceberg icebergs float okay icebergs tend to not really sink very well they tend to sit on above water uh when H in winter when a Pond freezes the the top of the pond freezes water essentially the ice is is less dense it doesn't sink it's less dense than the water which is very important because in terms of life and things the The Ice floating to the top means that life underneath actually ends up the water underne is a little bit warmer as a result kind of traps some of that heat in and actually life can still thri Thrive under under that icy icy layer but the reason for that is that it's all down to really the hydrogen bonding so when we have ice we tend to find in a in a solid form so in ice more hydrogen bonding or or more hydrogen bonds and what it happens these these hydrogen bonds the water tend to have a very set length and we end up forming a really regular lattice um of of water molecules in in ice and as a result um it actually the the volume of the water increases so the water expands when when it freezes because this regular Arrangement becomes set up where the water molecules are actually now further apart than they were within the liquid form and because of that they become less dense but it expands and the one way to see this if you get a uh a bottle of water fill it right to the top screw the lid on put it in the freezer maybe put it in a plastic bag as well uh once it's frozen you'll find often it will it will burst somehow the top of the water will will come off and all the rest because actually it's expanded um and and basically split the bottle or four for the lid off and it's because of this idea of hydrogen bonding there are more hydrogen bonds in the solid form of water ice um and those hydrogen bonds lead to a really regular Arrangement uh which looks something like this where we have the water molecules sell this really regular Arrangement lots of space between them therefore lower density than the liquid form of uh of water uh and so it floats uh and we get this expanded volume those that really there actually one more what do one more one more diagram which is quite a good one which uh highlights a couple of the points um this graph shows uh it shows you a couple of um sort of trends that you can see this first one is the trend Down group seven here so Florine to chlorine to bromine to iodine the next one is the trend Down group uh six for what it was then Down group six going from oxygen sulfur selenium um to tum uh what we've got here is we've got each of those elements in the respective group seven here or group six here bonded to a hydrogen now remember what I said hydrogen bonding involved either oxygen nitrogen or Florine and straight away we can see how much an effect it has here's our boiling points both of these exhibit hydrogen bonding they are so much higher than the other examples it's so so much higher when we drop down to no hydrogen bonding it goes right down okay big difference between that you know 100 sort of Celsius degree difference there between boiling point booms down then notice that it increases well this goes back now to the vaner vales so in all of these cases Vander vales is increasing due to the size of the molecule increasing so we've got a couple of things we can see here bam straight away hydrogen bonding really high boiling point in comparison dropping down no hydrogen bonding but increasing due to the vaner vir and this is quite a nice example here where it's application and description really of of graphs explanation of this graph due to what you hopefully now know about hydrogen bonding and the Vander valves as well right what I'm going to do now is I'm going to hopefully apply some of those ideas to the uh to some exam questions uh and show you the kind of things that come up and actually the repetition that's often seen okay so going to look at some questions now um I actually just went through and did all these um there's five papers I'm going to go through and so five five questions I actually went through and did all these I had the video on pause so I'll do them again and hopefully the video is actually recording this time um but I'm sure we'll see in a minute so for the second time although you haven't seen it for the second time um this question here was it's uh from June 2010 just picked five random papers really looked through you almost guaranteed in every single Paper to get a question that um relates to inter molecular forces so this one here uh very not much on intermolecular forces but this one says about this molecule C2 F2 so suggest the strongest type of intermolecular force between C2 F2 molecules so my advice would be draw that molecule out you don't want to miss anything you don't want to get stuff wrong um it takes a couple of seconds to draw it out and at least then you can look at it and you can really visualize what's happening so my molecule what have I got have I got Vander vales well yes because I've got electrons so tick for Vander vales we know that's our weakest though have I got dipole dipole forces and have I got H bonding that's the key questions now so have I got hydrogen bond well to have I got dipole dipole well yes I have and the reason is because these guys are very electron negative chlorine and Florine and obviously the chlorine other Florine therefore electrons are very close to these halogens and as a result we find that there is a large difference in electro negativity between carbon and between the chlorine and between carbon and the the Florine therefore we do have dipole dipole we have these charges as such Delta negative Delta negative Delta negative do we have hydrogen bonding well we have some electronegative stuff particularly the Florine here which is part of oxy nitrogen Florine but we have no hydrogen so clearly not strongest one therefore is going to be the dipole dipole forces excellent one mark we are doing good okay next one um this one is from Jan 12 as you can see at the bottom there um and yeah this one here starts here so 1D set the strongest High interor force between hydren fluoride molecules well that was one of the examples that I gave straight away hydrogen bonding it has got Florine in it it has got hydrogen therefore boom hydrogen bonding nice and strong here we go draw a diagram to show how two molecules of hydren flid are attracted to other by the type in force you stated in part di include all partial charges and all lone pairs electrons in your diagram now I actually did a bit of a bad did a boo boo in in in the previous part of the video where I didn't draw enough loan pairs so please go back and correct that or uses as your standard loan version but you must include all loan pairs I did not um then I also didn't claim to include them all either so I'm not entirely wrong but also it was a bit bad of me but anyway let's go on with this question so hydrogen fluoride calent bond between the hydrogen the Florine three lone pairs on my Florine not on my hydren obviously draw my other molecule because it does specify two molecules don't start drawing 6 7 89 unneeded uh you have a couple of ways of drawing lone pairs I quite like the double dot one do not just do dots like that CU they're very easy to miss do them as a bit of metor looking couple of dots or stick that little balloon thing on here as you doing shapes and molecules that would be completely adequate as well either way I've got three loone pairs it asks for partial charges here so partial charges well I've got a Delta positive region and I've got a Delta negative region another Delta positive another Delta negative it does say all partial charge so don't just do it on one molecule that's ridiculous do it on all of them the bond itself is formed between the Florine and the hydrogen so it's the lone pair of the Florine and the hydrogen that we're getting or the lone pair on the oxy um oxygen or nitrogen whatever but in this one obviously it's got to be the lone pair on the Florine between the hydrogen and you're three marks for this mean three marks for drawing this much it's ridiculous the Mark is for all wait bother about that it's all partial charges is one mark all loan pairs is the other so all four partial charges all six loan Pairs and the correct placing of the bond three marks Bish bash BOS um I like this question this is a comparison question here it's quite a clever one I've lost the pen there is um it gives you the bowling points of Florine and hydrogen Florine it asks you it gives you what they are and blah blah BL blah and then it says explain in terms of bonding why the boiling point of Florine is very low well this is a clever one because some people here will start talking about why hydrogen fluoride has such a high one respectively or or um comparably now that's not needed because it's not asking that it's only asking you about Florine so why is Florine so low well Florine F2 only has vanav Val's forces nothing else there vanales are weak therefore not much energy required to break them and that really is probably three marks worth almost there but certainly this one here is is definit it's one if not two marks um in this question though we are only awarded two marks so it's going to be one mark for that one one mark that one two marks overall lovely job here so that's not a bad one at all if you want to look up the mark scheme of things the Jan 12 paper U that was okay let's go on to our next paper our next question then another very similar kind of question State the strongest type of interor force holding the water molecules together in the Ice Crystal well we know it's water so straight away we going H bonding we're not scared to write that straight away tick and Mark next one state the strongest type again in methane well methane if we draw it out contains no halogens contains no real difference in electro negativity between the atoms in the bonds therefore no dipole dipole it's got hydrogens but it's got no oxygen nitrogen or Florine no hydrogen bonding then or dipole dipole so it's only vaner vales easy Mark though that really really easy Mark so done another one there another two marks that was the Jan 11 paper if you want to check out mark schemes and things next one keeping them going this one is the June 11 paper uh and this is one here asking for an explanation as to why iodine has a high melting point than Florine well break it down think about the the hallogen group florine's at the top followed by chlorine followed by bromine followed by iodine down here we have a size increase number of electrons also increases and this links into exactly what I said in the previous part of the video more electrons bigger the size stronger the force so say that iodine is bigger than Florine therefore more electrons and we could go iodine here the I2 it doesn't matter either way the iodine atoms are bigger therefore the overall molecule is also bigger so iodine is bigger than Florine therefore more electrons therefore requires iodine to be specific iodine requires uh more energy to overcome clear here more energy uh to overcome van Val's forces it's trying to be comparable there you really are trying to compare iodine to Florine two marks again the idea of size and then the idea of more energy required final one here and I like this one it's a slightly different way of ouring this was June 12 um this time telling you there is no hydrogen bonding between phosphine molecules phosphine if we go up and look at this you can see phosphine here is this ph3 so there's my phosphine so ph3 why is there no hydrogen bonding well I have hydrogen but I got no oxygen nitrogen fluid but that's not enough to say that is not an answer that is not an explanation there is no oxy nitr or Florine what is it about oxy nitrogen Florine that allow for hydrogen bonding that's the key thing so here between my phosphorus and my hydrogen obviously I've got three of those bonds there is not enough of an electro negativity difference between these two to allow for a dipole a permanent dipole to be formed whereas if this was nitrogen there would be enough and that's the key thing if we have this exact same format here NH3 ammonia there is hydrogen bonding because this is high enough in terms of electro negativity to create a big difference here ultimately leading us to a d dipole um ultimately leading us to sorry I said dipole I didn't mean dipole ultim leading to a big enough difference to uh to give us the uh the partial charges to allow for the attraction to occur so basically the answer is phosphorus not very Electro negative therefore not a big enough difference in electro negativity between phosphorus and hydrogen one Mark there so few questions there that should hopefully be the topic of inter molcular forces um hopefully that's made some sense to you if you have any problems please do let me know and I'll do my best to try and answer any any queries um again hopefully that's been some help uh and thank you for watching