[Music] hello and welcome to this video on topic two for the Edexcel specification bonding and structure my name is Chris Harris and I'm from Alawis tutors comm and the whole point of this video is basically just to go through topic two is a revision tool and so we're gonna give you an overview of the topic to make sure that you've got everything that you need now the slides that I'm using here and you can purchase them for a very reasonable price and if you just click on the link in the description box be able to access them there but they they are a really good value because you can use them for your revision you can use many smartphone you can cycle through take notes from them do whatever you want and then print them off you know so it's really good it's just not it's just a kind of to supplement the notes that you've already gotten revision material that you've got okay so like I say this is specifically geared toward Edexcel specification and it meets these specification points here that you can see on here so basically it matches topic two for EDX up okay so let's make a start and look at ionic bonding okay so ionic bonding basically is where you've got your obviously ions and these are oppositely charged ions you've got positives and negatives and the held together by electrostatic attractions okay so these are obviously the positive ions now ions are formed when an element loses an electron and another element gains an electron to get a full and shell of electrons now in this case you can see that sodium has got one electron the chlorines got seven and the sodium will give up one electron and underneath you've got a full shell underneath and chlorine will then have a full shell as well and then obviously they'll form ions and then the opposites will attract so like this then go and we put square brackets around them to show that their eye on so positive around the sodium negative on the chlorine and basically they attract it's an electrostatic attraction now we can obviously use the periodic table to work out the types of ions so group ones form one plus ions group twos form two pluses threes form three pluses now when you get to Group five there are five electrons an outer shell is easier for group fives to to gain three electrons and to lose five okay is energetically lower to gain the three so that's what it does and that's why they form three minus signs Group Six's form two minuses in group sevens form one minuses you also need to know your molecular high molecular ions as well so we've got hydroxides which Oh H minus nitrates are no.3 minus ammonium is NH four plus sulfates are so4 2 minus and carbonates are Co 3 2 minus so make sure you know the molecular answer is obviously going to use them quite a bit in your compounds okay and we can obviously work out the formula of ionic compounds and we can use something called a swap and drop method so the first thing we need to do is write down our two ions that's calcium + 2 + and no.3 - so I'm just picking these as an example we then swap the charges over so you can see we've got a - on the calcium now + 2 + on the nitrate then we drop the charges so we get rid of the minus and the plus and drop the numbers so they're lower there you can see you have put brackets around the no.3 now because we want both on both - lots of no.3 and put 2 outside and then the calcium and then we just merge them to form the formula and we simplify if we need to but this is already the simplest forms at CA no.3 - calcium nitrate look at another example CA 2 + + O 2 - again same thing swap the charges over so for - - + - + drop the charges CA 2 and O 2 combine and simplify so CA 2 O 2 can be simplified to CA o so we call this calcium oxide so you can see it's very neat method of just working out formulas of ionic compounds okay and alec compounds such as sodium chloride for example they have giant ionic structures these are big big structures here's an example of a sodium chloride structure and they are a regular structure so we've got these night nice neat rows of chloride ions and sodium ions it's a cubic shape and we've got this giant repeating pattern as well that just keeps on repeating repeating and repeating so they're big giant structures most ionic compounds they dissolve in water this is because of your water molecules are polar and they attract the positive ions and the negative ions and we'll see that later on as well and basically because they can attract a different part of the ionic compound they are soluble they do conduct electricity when they're molten as well the ions are free to move around and it's the same in solution as well when we dissolve it in solution the ions are free to move around so therefore and it can conduct electricity and they have really high melting points there's lots of strong electrostatic forces between oppositely charged ions so lots of energy is needed to overcome these forces got to make sure you including all these words that all underline there make sure that you're describing why sodium Clyde has a high melting point okay there were also quite brittle as well and basically when we hit them with a hammer when we strike them the layers will slide over each other and but what then happens is your positives well then that will be right next door to a positive and they'll repel and so sort of break apart really easy and this explains why an ionic compounds like sodium chloride are brittle and because of this repulsion when you knock them and you get the same ions and close by so two positives together and two negatives together okay so the size of the charge on the ion affects the strength of the ionic bond okay so the bigger the charge on the iron the stronger the electrostatic attraction between the ions and so more energy is required to overcome these forces and so they have high melting and boiling points okay so if we look at the TAS employed as an example so it's made of a k plus and Cl minus ions has a melting point of seven hundred and seventy degrees Celsius but if we look at calcium oxide which is made of calcium 2 plus and OH two minus signs it has a melting point of two thousand five hundred seventy two if you look at the ion difference there calcium's obviously a bigger two plus and OH oxygen is no.2 - it's a bigger charge difference there so therefore they have a much much higher melting point compared to cently potassium chloride okay so the size of the Aisne which is the ionic radii that affects the strength ionic bond - so the smaller the eye the stronger the electrostatic attraction between the ions small iron these can pack together much more closely more energies needed to overcome these forces much much stronger forces and therefore and ions which are smaller have higher melting and boiling points sodium chloride for example is made up of na+ and cl- it has a melting point of 801 however if we look at potassium chloride now potassium is a bigger ion it has a melting point that is seven hundred and seventy degrees Celsius so it's lower because the potassium is a bigger iron so therefore and the attraction isn't it's great and generally the smaller the iron and the higher the charge the stronger the electrostatic attraction so we have high melting points the word we use to describe this as charge density we say they have a high charge density so a big charge and a small iron is going to pick a charge density than a small charge on a big iron so remember that word charge density okay right so let's look at some trends so the ionic radius increases as we go down a group okay so as you go down the group the number of electron shells increases so the ionic radius increases - so that's quite logical really so you might let that from topic one so the ionic radius and isoelectric ions decreases as the atomic number increases okay right so we need to have a look at some of these so isoelectric I and these are just different atoms of the same number of electrons okay so obviously they would form ions so let's have a look at these so we've got here's your I and then three - o - - f - na plus mg 2 plus al 3 plus so these are all isoelectric because they have the same number of electrons but look at the number of protons here okay it's increasing and then look at the ionic radius as well the ionic radius gets smaller okay and so we need to know why and the reason why it's because we've got these increased attractive force and because we have more protons look at the number protons increasing okay this is pulling these atoms in a little bit more and they have the same number of electrons okay so the outer electrons pulled in a little bit more okay remember the outer shell is the same for all of these ions in the in this table here because they're all isoelectric okay so you look at the number of protons here okay so we can look for evidence for charged particles and we can use this kind of nifty little bit of a kit here and we use electrolysis of copper to chromate player which is pair chromate six on wet filter paper and it shows evidence for charged particles so you see on the diagram here we've got wet wet Filip here is just mounted on a glass slide like a microscope slide we've connected the electrodes positive and negative to either side okay so when we put a drop of green copper tube chromate six it's placed on wet filter paper electricity is passed through it it starts to separate out you see there's the green copper to chromate six in the middle and then what happens when we switch the electricity on we get the positive copper bit the see YouTube Pro signs they move towards the negative cathode okay because opposites attract and what you'll see is a blue solution move so this is quite useful because it's colored because you can see the actual and ion separating out the yellow chromate six lines these ones here these have a negative charge cro4 - - these move towards the positive anode and you see the yellow solutions start to move towards this side and all this basically shows evidence for charged particles because remember ions are positive and negative charge and so if we pass and their electricity through them they can conduct electricity and also obviously they'll separate out according to the charges on the poles okay so let's have a look at color bombing so covalent bonding is the sharing of outer electrons okay in order in order for atoms to obtain a full shell so instead of and like ionic bonding when we had and after moving electrons from one place to another these ones are now sharing so there's no electrostatic attraction there's that word again between the shared electrons and the positive nucleus in the middle of the atom okay and we have single double and triple bonds and basically this is where we've got more electrons being shared so in the single bond but two electrons being shared in the double bond about four and in the triple bond about six okay and covalent bonds can be represented with lines as well as you can see here so this is a displayed formula and we can also have another type of bond we call this a dative covalent bond or a coordinate bond and this is where one atom donates both electrons to another atom or anion so I've got this example here of ammonia this is nh3 with a lone pair of electrons and hydrogen heat plus ion and doesn't actually have any electrons so normally this can't form any current bond because it doesn't have any electrons to share so if it's going to bond with something and the electrons have to both electrons have to come from another acid and this is what we call a dative covalent or coordinate bond so let's have a look they go H+ sits on top of there and we form this dative covalent bond both electrons are being donated from the nitrogen to the H+ ion and we can symbolize it using an arrow like down here so this is represents a coordinate bond remember the arrow shows the direction of where the electrons are moving from and where they're going to is going from the nitrogen to the hydrogen and overall we've got that positive charge also carbon monoxide okay it has a double covalent bond and a coordinate bond okay so it's got both so just make sure that when you're drawing things specifically for carbon monoxide they do have that double bond and a coordinate bond and you've got em both types in that molecule so as long as you're aware of that that's the main thing okay so let's have a look at bond enthalpy okay so this is the strength of a bond so bond has to be linked to the length of a bond okay so how long it is the shorter the bond the higher the bond enthalpy okay so in committed molecules there are forces of attraction you can see them just form in there okay these are between the positive nuclei and the negative electrons being shed so the green lines here represent the attractive forces which are there for this attractive force in the nucleus and the shared electrons so but there are repulsive forces between the two positive nuclei so there that's the red line in between and between the electrons in the atoms so we've got electrons here that are not involved in bonding but they're repelling the electrons that are involved in bonding so we do have that repulsion okay so there's a balance between these two forces and the result is something what we call a bond length okay so we've got this balance between attractive forces and repulsive forces so the greater the electron density okay between the atoms the stronger the attractive force so this means that the atoms are pulled further towards each other and this leads to a shorter bond and a higher bond enthalpy okay so the shorter the bond the higher the bond enthalpy so we've got a big electron density here loads of electrons we've got this really strong attractive force pull these atoms closer together and strengthens the bond so if you have a look at the CC single bond this has a pretty low electron density in comparison to a double and even a triple which has a much higher electron density and so the bond length is going to be much shorter for a triple bond than compared to a single bond and therefore the bond enthalpy to break a triple bond requires a lot more energy than it does for a single bond okay so let's look at some shapes of molecules because I would do it on the topic of covalent bonding here so we need to know some rules for determining the shapes of molecules so we're going to use the number of bond pairs and lone pairs of electrons to work out the shape of the molecule so molecules have a specific shape with specific angles and the reason is is because the bonds repel each other equally okay so the bonds contain electrons so they will want to be as far away as possible okay because they contain these electrons which are repelling each other so here's an example here this is a example of a tetrahedral shape and it's pushing each other apart equally okay so the lone pair next to the bond pairs repel more than the two bond pairs together and two lone pairs repel even further so you can see here they've got a lone pair this is repelling these bonds much much more squashing these down into the bond angle reducers in this one we've got 104 point 5 because it's reduced even further because we've got two lone pairs and I really squashing these bonds close together and so lone pairs what they do is they change the shape and the bond angles so lone pairs push bonding pairs closer together as we say generally for every lone pair you reduce the remaining bond angles by two and a half degrees you can see that's what's happened here but sometimes we don't do that and there is some exceptions which we'll show you later on you just have to remember the bond angles for them okay so the rules so use the number bond pairs and lone pairs of electrons to work out the shape of a molecule so what we have to do is draw a dot cross and this is to work out how many bond pairs and lone pairs we have so you can see here that we've got four bond pairs here in the lone pairs okay you're going to be careful with irons as well so with irons all we do is we add electrons to the central atom for negative ions and remove them for positive and so for example NH 4 plus nitrogen would have four electrons so normally nitrogen is in group 5 so there are 5 electrons because it's got this positive charge we just tick on let's run away we're left with 4 and all involved in bonding and so would be tetrahedral in this case now this is just a model ok this isn't actually what happens so when we have an ion we don't actually take an electron away from central atom it's just a method in which we can work out the shape of a molecule so this is just a model like I say this is the method of working out ok so like to say going back into this one for bond pairs here no lone pairs the total is 4 so this tells you it's a tetrahedral so if you have lone pairs that you need to replace bonds for then the lone pairs and the lone pairs will actually change the angle and the shape of the molecule but we'll look at them later when we come on to that so look at this one here this is water into water we've got to bond pairs here two lone pairs now if we do the same sum as what we did here before we have to bond pairs two lone pairs the total is for this molecules based on a tetrahedral but it has two lone pairs so what there's 4 tons it's based on that so what we do is reduce the bond angle by two lots of two and a half which is five degrees okay so what we're going to do is going to look at some of the he's going to look at the ones with no lone pairs first then we're going to look at the ones with lone pairs now for these you can have to remember these names and you're gonna have to remember the bond angles involved as well ok so we're going to use the number peasant long pairs of electrons to work out the shape of the molecule so once you've got that you then use this table as a reference but you will have to remember them unfortunately and basically work out the shape of them so if we've got a situation where we've got to bond pairs no lone pairs we form a linear shape and you can see here's the diagram here the picture being formed the linear has 180 degrees if we've got three bond pairs and no lone pairs we form a trigonal planar structure now trigonal planar has this shape and the bond angle is 120 degrees it's a flat molecule if we have four bond pays and no lone pairs we form a tetrahedral shape now the tetrahedral shape obviously has four different bits coming on you can see it being constructed there the bond angles 109.5 for all tetrahedral shapes if we've got a molecule with five bond pairs and no lone pairs then we form a trigonal by pyramidal shape and you can see on here that we have the trigonal bipyramidal shape being formed and the bond angles are has to 124 that be 90 degrees for the two axes now these ones are three-dimensional and the tetrahedral and trigonal bi-pyramidal you got the wedge showing the acid coming towards you dotted lines showing them going away and then we've got this solid line which shows us in the plane of vision you can see here we've got 120 the trigonal bit in the middle baby it's just this turned on its side and then we've got the two poles top and bottom okay if we've got a situation where we have six bond pairs and no lone pairs then the shape would be octahedral and you can see it's been constructed right there now and the octahedral shape has a square bit in the middle so it's like a flat square you can see here there it is and we have ninety degrees between each of these so it's all 90 degrees and it's called octahedral okay so we've got this shape here too coming towards us to going away and to in the air pole there okay it's called octahedral by the way because even it's got six bonds and if you draw a line from each of these and peripheral atoms here on the side up to the one the axis at the top it have a four-sided pyramid on the top if you do the same in the bottom it'll be four sliders as well obviously the number of sides and total would be eight that's why we call it octahedral okay so let's look at the ones with no long pairs right these get a little bit tricky here okay so you need to really remember these so for three bond pairs and one lone pair okay we've all a trigonal pyramidal shape there it is there look so this is an example of ammonia look at the bond angle 107 degrees okay because we've reduced that by two and a half degrees for the lone pair that's on there if you have a situation we've got two bond pairs and two lone pairs we get a fence molecule and again two lone pairs here they repel the bond angle and we get 104 point five degrees remember we're reducing it for every lone pair that we add we reduce it by two and a half degrees so we get one hundred and four point five gets a little bit confusing at this bit okay this is where that that kind of electron will breaks down that we mentioned so bond pairs three lone pairs two so what we get is a distorted t-shaped and we get something like this here again it would have been 90 degrees at this point here but reduce again we were just the bond angle by two and a half degrees but just look how we've drawn it okay it's all in the plane but we've just distorted this bit here it looks like a letter T but you reduced it by two and a half degrees from 90 remember is 19 that one so it's two and a half degrees so that one's going to eighty-seven point five degrees and if you've got four bond pairs and one lone pair and we actually get a seesaw as well so that's all look there it is okay now again if you look at the bond angles here okay this is where it kind of breaks down you've got to be really on the ball here to know these bond angles for these two one lone pair and four bond pairs and so again you draw your dot cross and if you've got that scenario this is the type you're looking for so it's called the seesaw maybe because if I flip that on its side it looks a little bit like a seesaw okay five bond pairs and one lone pair right what we get is we get a square pyramidal shape so for example I f5 is an example and here it is here and now you can see where our lone pair on here and while this is going to do it's going to push these four up and it's going to just close this angle down a little bit so it's going to be eighty one point nine degrees for this angle 90 degrees remains for the square bit in the middle and four for bond pairs and two lone pairs we get this square planar shape now you can see they've got too low lone pairs of electrons here now this is the one where the bond actually remains unchanged because we've got the lone pair pushing these ones down and this lone pair pushing them ones back up again so actually has no effect on the bond angle so this one remains the same at 90 degrees because they repel equally from opposite sides okay so let's look at some soon as we're still on the covalent side LS that some giant covalent so examples of giant common structures include graphite and diamond okay so that's our look at graphite graphite is a big structure and it contains lots of carbons each carbons bonded three times and the fourth electron that is involved in carbon is actually delocalized and it forms the layers now we've got lots of strong covalent bonds between the carbon atoms means graphite as a really high melting point however what we have let you say we got these layers and they slide over each other really easily you have these weak forces between the layers and that allows it to slide which makes it ideal for using a pencil we have these delocalized electrons as well these are between the layers this allows graphite to and usually as a nonmetal air conduct electricity because the electrons can carry a charge okay the layers are really far apart from each other as well compared to and the covalent bond length and this means that Graphite's got a low density and so and that's pretty useful because it makes it lightweight so you can carry these things around like a say in pencils graphite is insoluble it doesn't dissolve the covalent bonds are far too strong and too break for water so thankfully put a pen to the water and it won't actually dissolve so that's pretty useful just in case you wanted to do that right let's look at diamonds and diamonds are the giant coil and slightly different structure though this time each carbon atom is bonded four times and like graphite weights just the three platon pats rigid arrangement allows the heat to conduct well in diamonds and so that's pretty useful but unlike graphite diamonds can be cut to make gemstones so you make it into jewelry a really high melting point just like graphite loads of strong covalent bonds really hard to to break these because you need so which end you to overcome these strong covalent bonds diamond doesn't conduct electricity well and it doesn't have any delocalized electrons but it's an insulator really don't have any delocalized electrons so unlike graphite and again just like graphite diamonds and soluble commenter bonds are far too strong to break so you can put a diamond in water so if you've got a diamond ring and you wash it up look like it won't dissolve in the water thankfully this can be a very expensive and material to pay for and silicon dioxide is another example that she has the same structure as diamonds same arrangement of atoms and obviously the properties maps as well so yeah so as long as you aware that the structure facility and dioxide is the same as for diamonds that's the main thing okay so a little example of a giant karrine structure is graphene okay no graphene is a bit unusual very useful property actually it's one layer of graphite so we just take one layer from that one awesome thick is made of loads of exact little carbon rings as graphene is really because it's really thin it's only one cell thick it's lightweight and it's transparent so it makes it ideal for electronic uses and this because it has delocalized free moving electrons between this sheet here excellent conductor of electricity and they can carry a charge so the same delocalized electrons strengthen the covalent bonds okay and this gives graphene a really high strength property so you can twist it and bend it it's really really difficult to to break this structure and the uses of it well aircraft gels is obviously quite useful because of its strength and its light weight super computers because in conducts electricity to high-speed computing very little very low resistance electrical resistance and this stuff compared to silicon and it's using smartphone screens for that so and you'll find maybe in the next 10 20 years you're going to get these materials very commonly aware very commonly around you'll probably get these really thin like sheet like materials and you already seeing them in some applications now for example Google specs and you get these really thin screens that are mounted onto lenses so you'll see a lot of this as well and particularly the rise of graphene so what's this space to suppose depends when you're watching this I suppose okay metallic bonding is an unsightly bombing so these obviously occur in metals only so they have giant metallic lattice structures what they have is a positive net lines okay so these are in the middle these are formed from them to the metal atoms they donate electrons into a sea of delocalized electrons that you can see here so there's got a few of them floating around there is an electrostatic attraction between these positive net lines and the negative delocalized electrons so there they are there there's your delocalized electrons and they're attracted to the positive metal ions and the more electrons an atom can donate to the delocalized system the higher the melting point so for example magnesium as a higher melting point than sodium because magnesium is in Group two can donate two electrons and to the delocalized system where sodium can only donate one per atom so so magnesium has a higher melting point they're really good thermal conductors because they have the delocalized electrons they have it in transfer this kinetic energy then when we heat these up but giving the electrons energy they start to vibrate and then knock into neighboring atoms and they pass on this kinetic energy but me some good conductors of heat obviously they're good electrical conductors as well because they've got these free moving electrons they can carry this current and they can move it through the electron or the metallic structure thumbs are no surprises battle's a good conductors of electricity they have high melting points as well because they've got these lots of strong electrostatic attractions between the delocalized electrons and the positive metal lines and they are insoluble and because of the strong electrostatic attraction between the pasta methylenes and the delocalized electrons now you can obviously see a theme here there's a lot going on about the strong electrostatic attractions and metals are malleable and they're ductile as well as the ion layers can slide over each other when we hit them with a hammer and they can still retain this attraction between the delocalized electrons and the positive meth lines so because this can still be retained despite distorting the kind of ion structure and this gives metals this property where you can hit them and hammer them into shape and draw them into into wires as well okay electronegativity so electronegativity is the ability for an atom to attract electrons towards itself in a covalent bond a really really important definition you must must remember so the further up and right you go in the periodic table excluding the noble gases the more electronegative the element is so fluorine is the most electronegative element in this series so the power leaning scale helps us to quantify how electronegative an element is and you can see here that fluorine is the most electronegative it has a value of 4 whereas on this scale here hydrogen is the least electronegative with a value of 2.2 and basically the bigger the difference in electronegativity and the more ionic a compound will be and if there's no difference then we purely covalent so for example CL 2 two chlorines because they both got an electronegative electronegativity value of 3 the difference is 0 so that would be an example of a purely covalent compound whereas something like H F because you've got this big difference here to point to and for this we'll have more ionic character and then just purely covalent so even though it would be a covalent bond there would be some ionic character there because the difference in electronegativity so yes you can have covalent and ionic character in the same thing effectively that's what I'm trying to say right ok polar bonds so covalent bonds can become pala it's very important this is only available ok and if there is a difference in so the bigger the difference in electronegativity the more polar the bond will be so let's have a look at that example that we looked at and will we use HF before with HCl ok chlorines more electronegative so what that's going to do is pull these electrons towards itself in this covalent bond the electrons are going to be closer towards the chlorine than the hydrogen we have this Delta negative on the top seeing that this is obviously electronegative Delta positive on the hydrogen so this is basically to show the polarity and you can see in this example this is CL 2 the atoms are bonded with the same or similar electronegativity value so these both have a value of 3 appalling value of 3 so these are not polar ok and the shared electrons you can see set bang in the middle and hydrocarbons I've also crusted nonpolar as well so even though there might be a slight difference in electronegativity between carbon and hydrogen they are classed as nonpolar and look out for molecules like this you might have uneven distribution of electrons so these are polar so water is a classic example of that carbon dioxide and is a molecule it appears to be polar and looks as though it's polar but because it's a symmetrical molecule in terms of the distribution of electrons and this isn't actually an polar this is a non-polar molecule that the overall polarity is 0 because we have this symmetry so carbon dioxide is an example of that okay so intermolecular forces so we need to know about the and the 3 types I'm going to look at London first so London or also known as instantaneous dipole induced dipole you can see why we just call it London these forces exist between atoms and molecules now the London forces are the weakest ones ok now any molecule with electrons can form a dipole when they move near another Actimel molecule so here we're going to use CL 2 as an example so this occurs as electrons in the molecule or an atom they can move from one end to the other and you can see here that we've got a distortion of electron so we've got more electrons at this end of the molecule and we do at this end because we've got more on this side we have this Delta negative because more electrons are on that side than on this side now this happens when this molecule goes close to another one it might be another chlorine molecule for example and obviously the electron clouds and both molecules will repel each other and so we form this temporary dipole here this instantaneous dipole induced dipole okay so they basically do this but when the molecules move away that interaction is destroyed so you can see here that we've got two got a few chlorine atoms here and now obviously the electrons in here will repel the electrons in the neighboring molecule so that would push them to that end of the molecule and what we have is this Delta negative and positive attracting to each other this week attraction is called a London force and so this is only exists when these two molecules nearby when this molecule leaves or moves somewhere else then the electrons in here will move back to its original position and the whole thing will start again okay an example of a London forces in iodine now these can hold these crystal structures together which remember a Deans is gray solid it comes in crystals so iodine looks like this you see this nice regularly arranged structure of i2 molecules now all we have between these molecules are weak London forces and these hold the i2 molecules together however what we've got between the atoms of iodine we've got strong covalent bonds and these hold the 2id and atoms together they've got to know the difference between them a force happens between molecules hence the word intermolecular and a bond happens between atoms okay and the bond is much stronger than a force so when we actually heat this up I didn't actually sublimes and then what we're doing is were weaken in the London forces not breaking the covalent bonds okay so the bigger the molecule or the atom and then the more London forces we have because we've got larger electron clouds remember if the electron clouds that get involved with London forces so when we boil liquid then it could be a liquid such as and a fuel for example and although that would ignite everything that's not maybe maybe that's not a good example but yeah what would you engines were breaking the weak London forces not the corona buttons okay is really important when you're talking about boiling points of all simple molecular molecules like this which what about forces not bonds okay we've got up enough energy to overcome these forces and that basically determines their boiling points or melting points as a solid so longer straight chain hydrocarbons these have more London forces so more energy is needed to overcome these forces this means that the boiling point increases so if we start off with something like say like butane and compare that with some like decane which is got 10 carbons in it that's going to a higher boiling point than butan because of the fact that it's got and more London forces branched hydrocarbons over these branching here they can't pack closely together and so this weakens the London force interaction between these molecules and it lowers their boiling points so hydrocarbons of branching have lower boiling points than their counterparts with the same molecular mass ok let's look at the next strongest intermolecular forces it's called a permanent dipole dipole ok these exist we've got these permanent dipoles here as you can see in HCl in other words these these their polarity doesn't just exist when it's near another molecule this one is permanent so the exist in molecules the polarity for example like HCl what we have is a weak electrostatic force between the molecules a little bit like London forces again Delta negative on one side the molecule Delta positive except these are permanent so the Delta negative part one molecule is attracted to the Delta positive on the other and then make London forces dipole-dipole interactions involves molecules with a permanent dipole and so these forces are stronger than London forces it is important to note though that these molecules such as HCl for example that has these permanent dipole dipoles they also have London forces as well so they have both forces it's just the strongest intermolecular force is a permanent dipole dipole and we can test polar molecules by just running them against a and basically we take the polar liquid could be water for example in this case and we run it through a burette image so we get a steady stream of liquid running through the US if we put a charged rod next to that steady stream of polar liquid the liquid should bend towards the rod so in this case what's happening here is we're positive rods and the negative oxygens there Delta negative option on the water here is attracted to the positive rod and so it'll bend towards it so we've got this kind of attraction and that's the test for polar molecules or if you've got a polar bond okay and the strongest type of intermolecular force is hydrogen bonding okay and they have or we get these when we have very electronegative elements okay so we're looking at the likes of nitrogen oxygen and fluorine these are the most electronegative elements and obviously cuz it's hydrogen bonding we've got to have hydrogen involved in there so what you're looking for is molecules of nitrogen oxygen or fluorine and it must have hydrogen there for it to be able to hydrogen bond so an example is water water can hydrogen bond with each other and what we can do is we can show hydrogen bonding it's basically an interaction between the lone pair of electrons on an oxygen and a hydrogen on another molecule and when you're drawing these make sure you draw all your partial charges your lone pairs as well on the oxygen you must show that the interaction and the interaction to a hydrogen hydrogen bonding always occurs between nitrogen oxygen or fluorine and hydrogen so it's important to note as well just like with the permanent dipole dye for that any molecules that have hydrogen bonding also have the two other weaker ones before it so water for example has hydrogen bonding but it also has permanent dipole dipole and London forces so it has all three it's just its strongest intermolecular force its hydrogen bonding so just be aware of that okay so we're going to look at some examples here of hydrogen bonding but this is in terms of ice now ice is a bit peculiar because when you cool water down obviously turns into ice but it actually expands it takes up more room there normally with most materials the vast majority of tails when you cool them down they contract they get smaller but ice doesn't okay and this is why ice is actually a regular structure and so it's got loads of water Gill's arranged regularly we have the hydrogen bonding obviously as we've seen before between the delta the lone pair on the oxygen and the hydrogen but what this does is this as she pushes the water molecules slightly further apart than if they were liquid and this obviously makes them less dense and hence the reason why water expands when it's frozen and obviously you probably would have seen this as well with face to put ice water it obviously floats so alcohols though are not as volatile as alkanes with similar masses and this is due to the hydrogen bonding in alcohol so alcohols can hydrogen bond a bit like water can and but if we obviously take a similar mass alkane compared to the alcohol it's not as volatile because of these stronger forces remember alkanes only have London forces their strongest intermolecular forces much weaker and makes them a lot more volatile okay so let's look at some data here you see the boiling points of hydrogen halide you can see here we've got to see different hydrogen here right here this 1 haytch F is got the muck that's got the highest boiling point out a lot of them it has hydrogen bonding loads of energy needed to overcome these stronger forces between the molecules and the rest of them don't have hydrogen bonding and so if we go down here we've got a mixture between permanent dipole dipole and London forces and because these a lot weaker than hydrogen bonding they obviously have a lower boiling point but you'll notice it does kind of gradually increase from h CL to H I and this is because we have as we go along here we've got bigger here light iron and that bigger Hill iodine comes with larger London forces so that means we've got an increased mass molecular mass of the molecule and we've got a higher intermolecular force between the molecules okay solubility so remember we talked about last time about these ionic compounds being soluble so polar substances can dissolve in polar solvents okay so for a substance to dissolve the solvent bonds must break in case we look at the solvent bonds breaking the substance bonds must break as well and the new bonds formed between the solvent and the substance okay so we've got a lot of things here so basically breaking the solvents breaking the substance and then forming new bonds between the two them so polar solvents these are molecules have a polarity some like water can hydrogen bonds that's called an eighth Chris solvents and some like propanone can have a permanent dipole dipole interactions and London forces these are called non aqueous solvents these are like organic forms so let's have a look most ionic compounds okay these can dissolve in polar solvents like water and basically what's happening is the Delta positive on the H is attracted to the negative ions in the ionic compound and the Delta negative on the oxygen is attracted to the positive ions and the structure starts to break down and so this is how powerful water is really so you can see here there's an example of it's you've got the positive ion here and the waters align themselves where you've got the Delta negative oxygen aligning towards the positive ions pretty clever isn't it and then you've got the Delta positive hydrogens here surrounding the negative and so what the done is effectively they've surrounded it and pulled it apart so we call this hydration we've hydrated the salt so for this to happen the new bonds formed must be the same strength or greater than those broken otherwise there's no point in doing it it's got to be and it's got to be energetically favorable to do this if not the substance is very unlikely to dissolve so aluminium oxide for example doesn't dissolve the ionic bonding is too strong even though even though it's obviously an ionic compound and most ionic compounds do dissolve in water because the ionic bonds between the aluminium the oxygen is so strong water can't break it down okay so let's look at some other examples here so some non ionic substances can dissolve too so alcohols these dissolve in polar solvents as they can hydrogen bond with water molecules so we know obviously and alcohols are covalently bonded but because of this lone pair and the oxygen and the hydrogen are more so they can hydrogen bonds so we've got that interaction so that allows them to dissolve but this hydrocarbon part is nonpolar so this bit of the alcohol this dissolve in water and basically the bigger this pit is the less soluble the our college so if we've got an alcohol with ten carbons on it it's not really going to dissolve very well because it's got such a huge chunk that can't dissolve so some polar molecules don't dissolve in water and here alkenes don't dissolve and their dipoles are not very strong so they'll really dissolve very well and water and water form stronger hydrogen bonds between each other than with the hero alkanes so hero alkanes are actually insoluble so the water is more than happy just interacting with itself rather than bothering with the halo alkane with a weak dipole okay so they can dissolve in solvents that interact by a permanent dipole dipole interactions though so we can do that so for example you might have propanil okay nonpolar substances can dissolve best in non-polar solvents so these are molecules nonpolar solvents basically are molecules that don't have a polarity so for example these are things like alkanes they're typically boring molecules okay so like butane all these have is London forces there's no polarity in there so alkanes dissolve best in nonpolar solvents as they can form London forces between the molecules so nonpolar molecules these tend to dissolve well in water as water forms and stronger hydrogen bonds and between each other than interacting with the non-polar molecule so again the waters more interested interacting with its own molecules than with the nonpolar on okay so let's summarize some of this bonding now you can see here that we've tried to summarize all of the bonds that we've seen here so giant covalent macromolecular graphite diamond silicon dioxide remember these ones graphene's and on usual state they're solid they don't conduct electricity when they're solid or liquid Graphite's the only exception and graphene of course and so because they don't have really difficult to melt a little conducts they're not soluble in water remember them covalent bonds are too strong and their melting points are high lots of energy needed to break them strong covalent bonds simple molecular there's two types of covalent here ID ammonia water little molecules like that normally liquids or gases they don't conduct electricity or when they're solid or liquid soluble in water it depends on the polarity really and obviously we just looks at solubility just before they're melting boiling points are low because got weak forces we're not breaking bonds remember when we talked about these we're just weakening forces so when you sort of melt melting and boiling point which on that weak forces giant ionic sodium chloride calcium oxide magnesium bromide these are examples of giant ionics and normally solid at room temperature don't conduct electricity in the solid because the ions are not free to move around when they're as liquid they can or even dissolve the water they can't because they've got free ions and these allow electrical conduction soluble in water yes they are again we've seen that just before and the melting boiling points they're high lots of energy needed to break them strong electrostatic attractions between the oppositely charged ions metallic tone bonding metallic magnesium sodium copper these are all examples of metallic bonding usual state they're solids obviously these are metals and they conduct electricity sort of yes because they got delocalized electrons and they do it as a liquid as well they're not soluble the water bonds are far too strong to break melting and boiling points normally high because they have strong electrostatic attractions and so it's very difficult to break them so this polarity bit here just a reminder that polar molecules dissolve well in polar solvents but like water but your nonpolar is don't they they dissolve in nonpolar solvents so like hydrocarbons for example and that's it that's just a very quick overview of the topic to bonding and structure please support this channel if you and if you want updates on the videos that are put on and new videos and replacement videos then the best thing to do is to subscribe and so we just click on that central button the circle button in the middle and that should get you subscribe to the channel also just remind if you want to purchase these powerpoints for revision they're great value if you just click on the link in the description box and it will take to space rate and get hold them but that's it bye-bye