[Music] hello my name is chris harris and i'm from ally chemistry and this video is for the cias that's the cambridge internationals topic 3 chemical bonding topic so this topic is going to go through obviously different types of bonds so covalent ionic metallic and there's also some quite tough stuff in here to do with hybridization and orbitals and things like that so there's quite a bit of stuff in here so this is really designed to talk through um or kind of you know work through them different as well as bits of content in topic three now this is obviously a powerpoint slide and they are available to purchase through my test shop if you click on the link in the description box below you'll be able to get a hold of them there they're great for revision they're actually bundled up um as part of the inorganic uh sorry the the physical chemistry uh topic with other of the topics as well but it'll be clear if you go and have a look on there you can see them um great for revision and on the move etcetera and using alongside your revision books and notes and whatever you have there so hopefully they should be nice and straightforward nice and clear so again it's designed specifically for the cie topic so hopefully everything on here should be um nice and clear um so let's start with um bonding first so ionic bonding first to say um now some of this stuff you might have seen from gcse um some of the stuff you'll see particularly with the orbitals bit and energy levels is kind of derived from topics one and two so you'll find a lot of the topics in cie kind of merge together so i will be kind of reflecting back on them topics as well it's really important that if you haven't seen them videos or you're not really comfortable with some of them areas i can't emphasize enough that you should be comfortable and familiar with them at least to kind of get some of the concept here i'll try and bridge it as much as i can but it's just to kind of make you aware of it um so let's start with ionic bonding first um so obviously bonding is quite an important part of chemistry because it's how um atoms and ions join together and it's how you create molecules and obviously that's just fundamental to chemistry so one of them is obviously ionic bonding and ions as you've seen earlier um in topic two uh ions are um oppositely charged uh so ions are charged particles but ionic bonding is where you've got oppositely charged ionic um substances or ionic um entities should i say and these are held together by electrostatic attractions okay so electrostatic meaning obviously electro meaning kind of charged static mean they're not moving um an attraction is obviously where they kind of attract together that's really all it means they always give these kind of fancy names in chemistry so um let's look at the first one um this is sodium and chlorine so these are two atoms here and we're no other atoms because this is just showing the valence electrons so these electrons in the outer shell so obviously we know that sodium is in group one so it has one electron it's out of the shell and chlorine is in group seven so it has seven in the outer shell now you'll probably recall think from topic two um that in order for these two and topic ones a little bit in topic one as well um in order for these two and four mines then um the chlorine needs to accept an electron and sodium needs to lose an electron and effectively they form charged particles so there we are okay so sodium forms a plus charge and chlorine is minus and when we draw these diagrams these are called dot cross diagrams by the way and you'll see they've been covered and bonding later on um but effectively the positive and negatives they um they effectively attract to each other so just as a reminder um of the other ions that can be formed um so obviously this is in relation to so the common um ions in groups one two and three five six and seven obviously the the kind of the ones right at the end don't form any ions they're your noble gases so they already have a full shell of electrons and this group in the middle your transition elements um or these ones or d block elements should we say um so these ones can have variable or transition elements anyway these can have variable um charges they can form variable ions so they don't really form the same kind of pattern we see there so and we need to know some molecular ions as well and again you would have seen these in topic two but just to kind of bring some of them back in um hydroxides oh minus you need to be familiar with these um no3 minuses nitrate ions ammonium is nh4 plus sulfate ions so4 2 minus carbonate iron co3 2 minus you might see hydrogen carbonate as well which is h co3 minus so you might see various different molecular ions you do need to be familiar with these because you will see compounds and substances with these ions that exist okay so um we can work out the formula actually of a compound um by um using what i like to call a swap and drop method now you might have your own methods and you might be able to work this out without using this method it's just a way in which you can work something out and you might have your own which is you know which is fine so you go with whatever you feel comfortable with but i'm going to show you this method because it you know it may help some people so um working out the formula of substances is quite important obviously it's an integral part of chemistry so um in this case we're going to show you this method here so we're going to look at calcium and um nitrate ions for example so no3 minus and ca2 plus now if we wanted to know what is the formula of calcium nitrate this is the way in which we go about it so or you could go about it should i say so what we do is we write out the charges first right of the ions and then we do a swap of the charges so we basically move the two plus over to the no3 and remove the um the minus bit from the l3 over to the calcium bit as you can see there and then we drop the charges so we basically drop them to so they're kind of subscripts um now you'll notice that with the calcium we don't bother putting the one there we just kind of omit that because we know just one there and then the two is there notice what we do with the nl3 though is that we have to put that in brackets to basically say right we've got an n or three kind of unit and we have two n or threes that's why we put the brackets around there we can't just put two next to the three otherwise look we've got look as though we've got 32 oxygens which just doesn't happen and so yeah very important to put your brackets around that as well um so and we simplify to the whole whole number ratio if it's possible and if needed in this case it's already in its simplest form and so therefore we have formed our um um our substance which is calcium nitrate so that's the formula for calcium nitrate um let's look at another one let's look at calcium and oxygen follow the same method swap the charges over two minus and two plus drop the charges ca2 and o2 so take the little negatives and positives away and then in this case we do need to simplify it to its whole ratio whole number ratio which is going to be or the kind of lowest the the lowest common denominator i suppose um in which case this is cao so that's calcium oxide okay so that's the formula for calcium oxide there so um what we also need to be aware of is what is an actual ionic structure um and a common one i suppose is sodium chloride and there's many examples of this but ionic substances are generally from these giant um structures and giant meaning exactly what it is they're big they're very large structures um now ions obviously the ionic bond itself is the attraction between oppositely charged particles as the electrostatic attraction and these two here are examples of ionic bonding now when we write them we write them just as c a o and it looks as though it kind of floats around on its own just as a pair quite nicely in reality that just isn't the case these atoms want to be surrounded by oppositely charged particles and they will arrange themselves as long as it's you know feasible they'll arrange themselves to be next door to as many different oppositely charged particles as they possibly can and this is where you get these giant structures and and this is why it's important to kind of understand this so if we look at sodium chloride an example which is table salt effectively and you can see on this this is an example of a giant ionic structure you can see in this example it's cube shaped and the purple spheres represent cl minus ions and the yellow spheres represent sodium ions but you can see um from the structure you have this regular kind of pattern of different ions that stack together this is cubic in shape and you have this giant repeating pattern and if you notice if you look carefully at some of them you'll see that the saudi mine is trying and under the chloride as well is trying to be next door to as many opposites as it possibly can and this gives it this incredible robustness it's this really um the kind of physical properties of it um kind of stack up so um effectively they are all soluble in water um so if you say if you take salt for example this is just salt and you put that into water it will break down in water because each one of these has an ion so you got negative and a positive charge kind of stacked up here when you put it into water water is polar as you'll see in a moment and effectively it starts to pull this structure apart so when we can see evidence of that because we see salt dissolves in water however and they also do conduct electricity when they're molten or when they're dissolved in solution as well so if we take um salt water for example it will conduct electricity and because you've got ions that are free to move around in solution and that's quite an important aspect of you know something being able to conduct electricity same if it's molten as well so if we melt salt which i'll come back to in a moment um then the ions are free to move around as well and so therefore it will also conduct electricity so it's a it's quite a common feature but as i mentioned before this kind of really kind of robust structure has a high melting point there are lots of strong electrostatic forces between oppositely charged ions okay so in other words sodium is next door to loads of different chloride ions each each one of them ions is there's a an attraction between them but if it's surrounded by loads of these ions that's a lot of force to break down to actually um you know melt this so salts are actually really difficult to melt they've got incredibly high melting points and table salt for example if you put that in a frying pan and on a high heat it's unlikely to melt um because of the the strong electrostatic forces however if you put it in water it will break down quite readily so it just shows you the power of these forces i suppose between molecules um in this case is probably more powerful than the thermal energy required to melt it so it's quite uh quite strange quite um you know quite an unusual concept so and they are brittle law obviously you know we know salt you know if you take a clump of salt and then you know push it with your thumb it'll probably break um so it is that you know the layers do kind of slide around quite a lot um if you hit it and if you get two negatives together so chloride and a chloride obviously they repel and so therefore breaks so they are quite negative um and uh sorry they are quite brittle um so ionic compounds generally okay so um another way to kind of just prove for charged particles you might see this in in your practicals um you might not it's just a a way in which you can demonstrate ionic substances so in this case we're going to use a process called electrolysis which again you might have seen at gcse um and we're going to use copper in this case we're going to use copper ii chromate on wet filter paper the reason why we do this um is because um copper ii chromate is coloured um and it's green so what we do is we basically put a drop of green copper two chromate on wet filter paper and we pass electricity through it and effectively what we're doing is pulling the charges apart now you can apply this and you know that copper two again from topic two um the roman numerals tell us that copper is a two plus charge and obviously the chromate um is six um minus which is here so that you chromate um now when we switch the electricity on it presto your copper two ions move towards the negative side and copper ii is blue so you get this lovely blue color that appears on there then if you look on the other side your chromate ions are negative and so they move towards the anode the positive side of it and you get this yellow solution so this is just i mean this example here copper chromate is used deliberately because it forms collard that it forms coloured substances so obviously it's green to start with and you get blue and yellow either side other ionic compounds aren't as colorful so therefore you wouldn't see it it's invisible it's just a neat way in which you can demonstrate that you do have an ionic compound that's all okay all right so let's look at another type of bonding which is covalent bonding now covalent bonding um is a different type of bonding so ionic bonding is where you've got ions and effectively you have charged particles and it's an attraction of the charged particles and that's how it works covalent bonding um means it's going to break it down core means to share and valence means the outer electrons so we're sharing the outer electrons in this type of bonding there's loads of different examples here and and these are probably some of the more common ones you will see a few more as well but um as you can see there's quite a few on the end i think you'll probably get the idea when i talk through them um but like i say equivalent bonding is the sharing of outer electrons in order to obtain a full shell of electrons so this is a very similar concept to ionic bonding because they were giving up and accepting electrons to get a full shell this is the same concept they're just doing it in a different way and generally um corina bonding happens between two non-metals whereas ionic bonding happens between metals and non-metals generally and so that's as a you know to give you an idea of which one's covalent which ones aren't so you come in at once as you can see all these are non-metals trying to bond together now in this case there still is an electrostatic attraction just like what you had with ionic bonding except the the attraction is not between ions it's between the electrons that are being shared in the actual um in the actual substances i suppose a bit like this here so these are shared electrons in the middle um and the nucleus of the atom in the middle so it's this kind of attractive force between the nucleus and the shared electrons that effect we create the bond so we have single double and triple covalent bonds um and obviously you've got more electrons are being shared in the triple bonds than they are in the single bonds so your triple bonds are going to be stronger than your single bonds and you'll need to know a little bit more we'll come on to the hybridization orbitals bit later that gets a bit complicated there but um for this purpose it's just basically looking at what a bond is what covalent bond is and now covalent bonds these can be represented by lines as you can see in this diagram you might have seen these before and the line represents a single bond a double line represents a double bond and a triple line represents a triple bond so you've got another type of bond as well and they've got two names for this uh they're either called dative covalent or coordinate bonds um and this is um an unusual bond it's not as strong as a traditional covalent bond generally um but this is where you have one atom that donates both electrons to another atom or ion to form a bond okay so an example you can probably see on the diagram here um is ammonia which is nh3 which you can see here so you have your traditional covalent bonds between the nitrogen and the hydrogen atoms we have this kind of lone pair of electrons on the nitrogen and we have this hydrogen ion here that doesn't have any electrons of its own it's electrons being removed and that's why it's got a positive charge so if we remove that across there there we are this is a dative covalent bond okay and the dot cross diagram just illustrates that a little bit clearer because we know that these electrons have come from the same atom which cases is nitrogen this is an example of a um a coordinate bond or covalent and it's represented with an arrow like that so if you have to draw this out as a diagram as you can see there and you've got your nitrogen with your three bonds either side and you've got this arrow here which represents the electrons are moving from the nitrogen or the electrons are part of nitrogen atom and they're being donated to hydrogen in a shared way and we call that a coordinate bond okay so effect you form nh4 plus that's an ammonium ion which you will have seen in topic two okay so um co this is another one this is carbon monoxide um it has a double bond and a covalent bond uh sorry a double covalent bond and a date of covalent or coordinate bond and so carbon monoxide almost is that poisonous gas it's it's odorless you can't see it it's it's toxic obviously um this is um obviously the bonds here you kind of a mixture of the two so you have a double bond and a coordinate bond in the same and this is an example of it which is carbon monoxide okay so let's look at so another example um and so this is the one where you've got a mixture of different bonds here so this is um an example of um basically aluminium chloride so you've got alcl3 but the two molecules are kind of joining together and forming a coordinate bond between the two so um this one's aluminium chloride and actually the formula of aluminium chloride is al2 cl6 because this is the most stable form that exists in so it's not alcl3 but it's just really here just to show you a another example of that okay so let's have a look let's kind of dig a little bit deeper into kind of covalent bonding and what this kind of forces these electrostatic forces are so here we've got um a diagram here and we've got a nucleus in the middle as you can see here and we've got our electrons here we've got this covalent bond that we've just seen before now you remember um that bond um the bond enthalpies the strength of a bond is linked to how long that bond is and the length of the bond is dictated by the forces that are involved between the nucleus and the shared electrons as you can see here so basically the shorter the bond the higher the bond enthalpy okay so we need to know this relationship between the two so in covalent molecules there are forces of attraction there we are and you can see them on there between the positive nuclei and the negative electrons being shared so there's your positive there's your attractive forces here and it's this that effectively just pulls these two atoms together and that's effective what the what the bond is um we then have um repulsive forces as well between the two um positive nuclei obviously the nucleus is the nuclei of the two atoms don't want to be near each other but we also have it between the electrons the electrons don't want to be near each other as well so um you've got these two forces at player here so you've got so much help to pull the atoms together but not too close because there's constituents in the element in the atom sorry they're also repellents that kind of push it away so there's a balance between the two forces and as a result we get something um well we get a bond length we get a distance between um you know well the distance in the bond actually you know so the the total length of that so the greater the electron density between the atoms okay so the stronger the attractive force um and so this means that the atoms are pulled in further towards each other and this leads to a shorter bond and a higher bond enthalpy so if we've got really strong attractive forces in other words the distance between the shared electrons this electron density here and the nucleus is quite short then clearly we're going to have a very strong force here which is going to pull them closer together and shorten that bond which makes the bond um stronger overall and we can look at this in terms of your single double and triple bonds um and high electron density is shared effectively between them and shorter bonds as we go down here and we get shorter bonds and a higher bond enthalpy so your triple bond is shorter than a single bond and a double bond is longer than a single bond etcetera etcetera so um and the reason why is because this density is shared between them and this again this will become a little bit more kind of apparent when we look at um orbitals right so we need to know um excuse me um we need to know some shapes of molecules as well and we need to know some of the rules associated with the shapes and molecules so let's have a look so the number of bond pairs and lone pairs of electrons um will dictate what the shape of the molecule is and you need to be familiar with the names of these okay so molecules they have a specific shape with specific angles and this is because bonds actually repel each other equally so remember when we're looking at bond repulsion of repulsion of electrons etcetera it's a similar type of concept so remember in a bond you have shared electrons and electrons in that shared that sharing kind of mechanism don't want to be near other um shared electrons they want to try and repel themselves um as much as they possibly can um and and that basically gives rise to a ship particularly if i say you've got a molecule here where you've got multiple different bonds here so each one of these has a shared pair of electrons but they're going to try and repel each other in a three-dimensional space okay so um a lone pair next to a bond pair these repel more than two bond pairs together okay and two lone pairs together repel even further so if you have a molecule that has a lone pair of electrons a bit like um the example we looked at ammonia before which is nh3 that had a lone pair of electrons on there and it had three bond pairs i hadn't three nh bonds and had the lone pair on the top so that lawn pair has a much bigger repulsive force than between two bonds that are next to each other you'll see some examples in a moment i'll try and explain that so there's your ammonia there it is actually there's your example so there's ammonia it could be ammonia so your nitrogen in the middle um and then you've obviously got um your um your bond pairs there as well okay okay and there's another one which is water so this could be oxygen with two hydrogens for example it has two lone pairs of electrons and that pulls that further apart now you'll notice that we've actually got some bond angles here now this is a i mean you'll see in a moment we call this tetrahedral so the tetrahedral shape which is um which has a bond angle of 109.5 degrees very important to not miss out at the 0.5 so basically that's the distance between each one of the bonds in a three-dimensional space and the wedge basically shows the bond coming towards you and the dotted line is the bond going away from you um this obviously with the lone pair squeezes that a little bit further so the bond angle is a bit tighter between these bonds here which is 107 because remember this repels this a little bit more and obviously the two lone pairs repels even further and it squashes these two bonds a little bit tighter together and we get 104.5 so basically you do need to be aware of these bond angles now you'll notice there's a bit of a pattern here and effectively the lone pairs change the shape and the bond angles as you can see here and it pushes them um closer now generally for every lone pair that we have we reduce the kind of bond angles the remaining angle between bonds by two and a half degrees um so that's generally you'll see some examples where it doesn't quite happen but in this example you can see if we take two and a half degrees off 109.5 we get 107 and then another two and a half degrees brings it down to 104.5 we'll go through some of the other examples later on okay so let's let's look at um look at some actual shapes and look at some specific examples so when you're trying to work out what shape your compound has or your substance has you've got to follow a specific method now again you might have a different method to what i'm going to use here but this is just a way in which you can do it so what i would do is i'll draw a dot cross diagram out to work out how many bond pairs we have and how many lone pairs we have so here i've got methane which is ch4 and you can see we've got four bond pairs and all lone pairs in this example and then with ions if we have ions and we just add electrons to the central atom for negative ions and remove them for positive ions so for example for nh4 plus so ammonium nitrogen would have four electrons um all involved in bonding and so therefore would be tetrahedral okay and you'll see you'll see some examples later so this in this example here we have four bond pairs as we said there we don't have any lone pairs so the total is four and the total tells you the shape in this case it is tetrahedral if we had lone pairs here you would need to replace the bonds for lone pairs and change the shape of the bond angle as we've just seen before so when we showed you them three different diagrams obviously a lone pair generally reduces the bond angle by two and a half degrees for every lone pair that you've got okay so here's another example here this is water and you can see we've got our lone pair there so this has got two bond pairs two lone pairs and that means we've got a total of four so we're starting from a tetrahedral position there we are okay so it's based on a tetrahedral position because we have a total of four that's the kind of start position um but we reduce the bond angle by five degrees because we've got two lots of lone pairs in there so what that means um is that we we basically see um a different structure we don't have something that's tetrahedral we have something that's called non-linear which sounds very exciting or bent we call it a bent molecule so basically any molecule which um it was based on tetrahedral and has two lone pairs then that's where it is so the the purpose of this really is to give us a start point so when we got two bond pairs two lone pairs this gives us a starting position and then we can alter according to the number of lone pairs okay so let's have a look at some specific shapes then let's have a look at some examples now you really do need to be familiar with these okay so hopefully this should make it a bit clearer so we're going to use the number of bond pairs and lone pairs as we've just seen before um to work out the shape of the molecule okay so let's have a look let's look at the first one so we've got um a molecule with two bond pairs and no lone pairs is called linear okay now all these examples here are not going to have any lone pairs of electrons so this just gives you a an idea so an example would be um beryllium chloride so becl2 so another example and hopefully you'll probably see the kind of format here we're going to go in order so this one's got three bond pairs no lone pairs all of these don't have any lone pairs on them example is bf3 this is what we call a trigonal planar molecule and the bond angle is 120 degrees so planet because it's flat so that's the furthest to where these atoms can be um you know within this molecule and four bond pairs and no lone pairs is called tetrahedral as you've seen before and that has a mole sorry a bond angle of 109.5 really important you remember 109.5 don't just put 109 down and that's called tetrahedral um five bond pairs and no lone pairs is called drag trigonal bipyramidal looks a bit odd um so effectively an example is a phosphorus pentachloride and which is pcl5 so just to kind of pause on this one a little bit just kind of show you the shape so imagine you've got these three here these three atoms here these are all in like a trigonal um planet arrangement very similar to that we've just tipped it on its side okay so we've got these three here and then up and down so in the polar kind of the poles of it we've got um the kind of two more atoms here the reason why it's called trigonal bipyramidal is the trigonal is the triangle bit in the middle here so you can imagine draw a triangle and by pyramidal is because if we join lines up imagine if we could draw a line from each one of these upwards you would form a pyramid on the top there and the same on the bottom as well so that's why we call it by pyramidal so you've got two pyramids and from a trigonal base you see this one has two bond angles 120 degrees obviously between these which is just the same as that basically and obviously the difference here is 90 degrees because that's like t-shaped like a t-shirt molecule okay um okay so the last one on here um so these are all molecules with no lone pairs um is uh six bond pairs and no lone pairs this would be something like um sf6 for example um and this is called octahedral um the bond angles are 90 degrees between all of them um now with octahedral the reason why it's octahedral is again if we draw a line going one two three four up to these we've got the square in the middle and then you've got the two poles either side and then it would form an eight-sided three-dimensional shape so you would have one two three four so there's your difference your different kind of faces imagine you have to yeah you have to be a little bit more visual here might be quite tricky some of you might be able to see this somebody might be quite tricky um but that's why it's octahedral don't get that confused with the fact that it's got six bond pairs for something that's octahedral octa is obviously eight but it's just the shape if you have to draw that into a 3d shape you have eight faces on it okay right you must remember them that's really important right so now going to look at some examples with lone pairs this time okay so um let's have a look at this one then so this is um three bond pairs and one lone pair of electrons um this is clusters pyramidal so nh3 is a classic example of this and remember you have a lone pair of electrons this was based on so you've got one two three four so you've got four in total so it's based on a tetrahedral but it's got the lone pair in there so we reduce the 109.5 by two and a half degrees to give us a bond angle of 107 okay so that's total bond and lone pairs is four so that's our starting point but the lone pair obviously squeezes that a little bit further okay something with two bond pairs and two lone pairs is um h2o and it's a bent or non-linear as is also known as again the extra pair of the extra lone pair of electrons it creates 104.5 degrees um another example so three bond pairs and two lone pairs um an example is clf3 okay and this is trigonal planar so we've seen that one already we've seen that one you know in the in the previous examples there's no difference here this one is slightly different in the fact that it has three bond pairs and two lone pairs so it's not just three bond pairs it does have these two extra electrons but this is where that rule breaks down remember when i said that for every lone pair of electrons that you have you reduce the bond angle by two and a half degrees this here is one of the exceptions so um in this case you still do have your three bond pairs as you can see on here so there's your one two three but the two lone pairs of electrons kind of sit top and bottom and what they do is these electrons here will squeeze these ones down but these electrons here will squeeze them back up again so effectively they cancel each other out and you end up with trigonal planar which is what you've seen before so the bond angle remains the same at 120 degrees okay right so and the last example here um is um four bond pairs and two lone pairs and this is square planar so for example um xenon tetrafluoride so xcf4 um it has two lone pairs again a similar principle to the trigonal planar um you have a square in the middle here each of 90 degrees and the lone pairs cancel each other out so these ones will squeeze these bonds down these ones will squeeze them upwards but they cancel out so you get this nice square shape okay so the bond angle like i say it remains unchanged like i said they kind of repel equally from both of them okay right so make sure you're aware of them bond angles make sure you kind of know what they mean you've got to remember i know there's a lot to remember here and a lot of the time it's going to be practice you must keep practicing okay so giant covalent structures um so these are um um effectively obviously we looked at covalent structures just before then some of them were what we call simple covalent so they were very just small molecules but you can get giant covalent structures a bit like what you had with giant ionic as well so giant covalent structures some classic examples include graphite and diamond um so if we start with graphite first so obviously graphite um is found in pencils so it's not lead that's found in pencils it's obviously graphite um each carbon is bonded three times and the fourth electron is what we call d localized so it's not actually attached to any it's not bonded to anything in particular you might have seen some of this before probably um so graphite has loads of strong covalent bonds between the carbon atoms so it has a really high melting point you know if you try and melt graphite say from a pencil um it doesn't melt very well whatsoever now they do have these delocalized electrons and they kind of sit between these layers um that of carbon atoms which would sit in rings now these layers are held together by weak forces it's them delocalized electrons which create that weak forces and that means actually the layers slide quite well over each other so very useful obviously for a pencil because when you put the graphite against paper what you want is shards of graphite falling off the pencil and going onto the paper that's the whole point of a pencil so it's just as well that these layers are quite slippery and you know it can kind of slide off quite readily um also these delocalized electrons between the layers they actually allow graphite to conduct electricity so they're a really good conductor of electricity and quite lightweight as well actually so you know you can see quite a lot of good uses for it um and the layers are quite far apart in comparison to a covalent bond and this means that graphite has a lower density um so it's not it's not um it's not a dense material by a long shot say when we look at diamond for example which is the next example um so it's quite um it's quite low density um and it is insoluble as well it doesn't dissolve um the the covalent bonds are far too strong for water to break them apart so it's not like ionic compounds where um generally ionic compounds are quite soluble um you know water can break the eye on the ion structure down um it's not the case with graphite okay so let's look at diamond then so diamonds is giant covalent as well um this time again it's made from the same material same atom which is carbon but the structure of it the way in which these carbon atoms are bonded to each other is different um in this case it bonds four times in a tetrahedral shape and obviously you've seen tetrahedral before just before that when we looked at the bond angles or the shapes of the molecules um they're tightly packed rigid arrangement allows heat to conduct well in diamonds so they're they're pretty useful they're used in um like circular saws and drill bits and bits out in the construction industry obviously when they're used to cut stone and brickwork they can get hot quite readily and what you don't want is the is the blades to start melting or kind of warping that could be quite dangerous so they usually put diamonds on them to toughen the material up so it means it can take the heat if it's been used to cut something um and unlike graphite diamond can be cut to make gemstones so obviously you've seen diamond jewelry for example really expensive metal a really expensive um mineral sorry not metal definitely not a metal by the way um carbon's definitely a non-metal um it's got a very high melting point loads of strong covalent bonds okay and it also makes it quite hard as well as i mentioned it's used in cutting materials um diamond it doesn't conduct electricity it doesn't have delocalized electrons like graphite does and they're all occupied in bonding and uh diamond is insoluble as well so obviously it doesn't dissolve um the covalent bonds just like what they are in graphite they're just far too far too strong to actually break apart and dissolve in water okay so another example is silicon dioxide this is sand effectively so sio2 it has a very similar structure to diamonds so if the uh you might see in the exam they might talk about silicon dioxide and that has a similar structure to diamond as well the same properties so um makes it makes it quite easy to try and remember what's going on right so still sticking with covalent bonds um this is where we're gonna look into a little bit more detail and i would say out of all of this topic this is gonna be the most trickiest and demanding part so i'm gonna try and talk through it and as clear as clearly as i possibly can um but if you're sitting there i think and this looks really difficult it's because it is okay that's not too frightening it's just to kind of think well actually okay you should be finding this quite tricky hopefully if you understand what's going on here it makes it a little bit easier hopefully okay so here we go right so and we all have to learn somewhere so i'd learn this as well you know um you know when i was uh many years ago when i was at school or even university or anything like that so you know we all have to pick this up so hopefully it should be nice and straightforward fingers crossed so anyway when we look to covalent bonds we've seen single bonds and double bonds and triple bonds um and we've seen in topic two where we looked at orbitals and remember we had different types of orbitals you had s orbitals you had p orbitals d orbitals and f orbitals you had all these different orbitals and if you remember from topic two the s orbitals were spherical so they were round um and p orbitals were um well they look like um eights number of eights figures of eight and you had three different types of p orbital you had p x p y and p z so you've got three different types there and you only have one type of s orbital okay so if you can recall that from topic two if you're not too sure on that i wouldn't urge you at this point you could just go back and look at that because i'm going to be kind of talking about this in a bit more detail building on that knowledge now bonding is when these orbitals overlap okay so they start to merge together and as chemists do they give them fancy names and they use the greek alphabet quite a lot to do this so you have what you call sigma bonds and pi bonds and these are bonds that are involved in covalent bonding okay so they overlap and they form this covalent bond as we've seen here before so like i said from topic to i urge you to have a look at that first if you're not sure what i'm talking about here okay i will try and bridge it the best way i can but obviously it would help if you know this bit so you've got 2s and 2p are quite close in energy okay now on the diagrams in topic 2 you'll see we drew energy diagrams and we showed the s orbitals and the p orbitals and we drew kind of them in order of energy now the s and the p are quite close in energy and that gives this kind of unique feature in terms of what they can do for bonding so given the right amount of energy this allows electrons to move from the 2s to empty 2p orbitals quite easily so it's a bit like i'm using analogies but like a double decker bus okay so you've got people sitting on the bottom layer and people sitting on the top layer now it doesn't really take that much effort to go from downstairs to upstairs in a bus and if you wanted to go upstairs you could do but in theory if you really didn't want to expend that much energy you'd just sit down stairs wouldn't you so it's kind of the same see orbitals in a similar way you've got the s orbital which is like the downstairs bit of the bus and the p orbitals which are upstairs and there's some empty chairs upstairs and we can move an electron or a person from downstairs to upstairs in a free seat if if we wanted to okay so and see in this way so this is like energy i suppose the energy difference so here what we're going to do in this example is we're going to be using um we're going to use bonding with carbon atoms and for bonding we must well for bonding to occur covalent bonding to occur we must have singly occupied orbitals okay that's really important so in other words if we use that bus analogy again in order for somebody in order for a bond to occur in an atom there must be a free seat upstairs okay with nobody sitting in it at all so it's the same with this so in order for something to happen you must have a singly occupied orbital okay so um actually if i use a different that's probably not a good example let's say if you want to kind of this might be a bit weird i don't mean it to be weird at all but um if you have um say a bond between two passengers saying you want them passages to maybe have a conversation with each other you've got to have a spare seat haven't you so one person might sit on the seat and they might have a spare one you know on the you know next to you nobody sit next to you let's say the bus stops at the next stop it picks some people up that person then comes onto the bus and sits next to you now for that person to have a conversation with you they need to realistically be sitting next to you don't they so um they need to have a free space next to you to do that otherwise they'll be shouting across the other side of the bus um now in reality does that happen i don't know most people probably don't they just keep themselves themselves don't they but in atoms it's exactly the same so in other words for two atoms to kind of bond together they've got to have an i've got of a singly occupied orbital you can't have it where you know it's fully occupied okay and it must be singly occupied i hope you get that so far so um so there are two ways in which ch bonds can be formed and here i'm just going to kind of talk through um the two main ways but there's obviously one way which it it does actually kind of form and that's what i want to talk about here so um we can either overlap the two p orbitals with the s orbital in hydrogen so we're going to form a ch bond so remember in hydrogen you only have a 1s orbital because it only has one electron in there so it's singly occupied and with your carbon um you have six um electrons i've seen in in total with carbon um so you have some in the p orbital as well so you can overlap the two p orbitals with the s orbital and hydrogen and there we are you formed a bond or you could be quite clever and you can migrate an electron from the 2s orbital that's a bit like downstairs in the bus um into one of the empty 2p orbitals and we form a brand new orbital called a hybrid sp3 orbital wow okay right let me just kind of pause there and just explain what this means so hybrid um you might have seen cars which are hybrid cars okay so cars are either um petrol petrol or diesel so the fossil fuel run or they're electric and then you can have a car that kind of sits in between the two and has a bit of petrol and a bit of electric in it and we call that a hybrid so basically it's kind of it's a merger of two different types of technology put into one car it's the same with atoms so effectively you've got your s orbitals you've got your p orbitals and you can kind of have somewhere in between which we call an sp orbital okay so they're just kind of it's like petrol it's a petrol electric orbital okay so um so this is called an sp sp orbital now s p comes from the s orbital and there's a p orbital in there as well um now the three bit tells you how many orbitals are involved so we have three p orbitals involved and one s orbital and this is where we've got an s one s orbital and three p orbitals um kind of forming a brand new orbital so it's neither petrol nor electric it's kind of a mixture of the two and this is exactly what we can do with orbitals so it's not an s orbital it's not a p orbital it's an sp3 orbital and they're all it's a completely different type of orbital okay but it's it's formed from the mixture of the other two so hopefully you understand that because that kind of makes it a bit clearer hopefully i told you this is quite tricky didn't i so if we just go back and again i urge you to look at topic two if you don't know this but a carbon atom has um in its outer shell so it's valence shell obviously it's got a one s orbital it's got two electrons in the one s but we're just really looking at the outer ones because these are what involved in bonding so it's got a 2s2 and it's got an electron in the px orbital and an electron in the py orbital okay so this is what this shows here in case you get the 2s2 2px and 2py so this s orbital is full this one here and these ones i've got one electron in them each there is a pz but that has no electrons in at all okay so we can go one way and say right so we can form two ch bonds and what we can do is these are singly occupied remember that one's full so that one can't get involved okay because that one's full but these two can and these can overlap and effectively form two ch bonds fine okay and that expels some energy obviously releases energy when it forms a bond uh and it's minus 824 kilojoules per ball okay so that's one way so that's option one option two is we can effectively this is like downstairs in the bus one of these electrons can get up go upstairs and occupy an unoccupied seat which is a two pz seat yep quite fancy um and it can sit upstairs and effectively you now have four seats where there's one electron or four kind of orbitals should i say stop using seats four orbitals which have one electron in each of them so you've got four singly occupied ones there it is okay so now this is our setup you have a 2s1 or 2s should we just say um 2px 2py and 2pz now this takes energy okay so if you had to get up and go upstairs you have to burn some calories not a lot like but you know you have to burn some calories to do that it's the same with electrons as well so that's going to take some energy to do that that's plus 404 kilojoules per mole and there we are and that's the setup we have there okay so now remember for bonding to occur we must have singly occupied orbitals this arrangement here we have four singly occupied orbitals now that means we can form four bonds instead of two now these four bonds when you form a bond energy is released and the energy um to form these bonds is four times four one two okay and this expels one minus 1 648 kilojoules per mole and the four bonds they form this new hybrid so these are not these are not sp px pypz these are effectively scrapped and we form four brand new orbitals called sp3 orbitals and you have four of these which is a hybrid so it's effectively having say um a petrol uh so we have an electric car here and a petrol and petrol and petrol effect what we've done is mashed them all together and formed four hybrids okay instead so it's got a mixture of all of them so they're brand new orbitals so let's look at the energy here so forming two ch bonds because that's an option um releases minus two 824 kilojoules per mole now forming four ch bonds by overlapping all of the orbitals the new hybrid orbitals we need to put energy in which is not good okay um which is plus four or four but we release 1648 kilojoules per mole of energy when we form four new bonds so the total energy change is minus 1 244 now it doesn't take einstein to work it out that actually um you get a much bigger energy release with doing it this way in other words moving an electron into a different orbital first and then forming four bonds then you do by just overlapping existing bonds here now remember atoms are incredibly lazy chemistry is like the lazy signs i like to call it um molecules want to be in a position where they're in the lowest energy form possible and a sign where they give out a large amount of energy like this overall is a good sign as far as molecule is concerned because it will sit in a much lower energy kind of status than if it was here and obviously because this is expelling 1244 kilojoules per mole of energy as opposed to 824 using this one clearly this is going to be the kind of root or this is the kind of correct version of how carbon bonds with hydrogen to form the ch bond and that is why carbon bonds four times with hydrogen so it normally exists as ch4 rather than just ch2 because of this set up here okay so really really important and this is because we have this ability to form an sp3 hybrid and just to kind of point on that as well we do form sp2 hybrids which i'll show you in a moment um and effectively that is where you have an s orbital um an s orbital um merging with two p orbitals so that's sp2 but you also have sp orbitals where you've got one s orbital um and one p orbital um merging to form an sp hybrid orbital so that's all that means so it's basically just a mixture of this bit here just says right we've got an s and we've got three p orbitals overlapping um and this is a brand new hybrid okay this is a brand new kind of orbital this is not an s it's not a p it's an sp3 it's kind of somewhere in between the two i hope you understand that it's quite tricky okay so let's have a look let's kind of carry on and look at a little bit more a little bit more detail so that was a single bond a single ch bond you do need to be aware of um sigma and um pi bonds now pi bonds are found in double bonds and triple bonds for that matter as well so sigma bonds is where we have two orbitals that overlap okay really important so in the ch bond example that we've seen previously ch4 has four um sigma bonds um some with s and p characteristics hence why we say sp3 hybrids so the bond is effectively then four sp3 hybrids kind of overlapping forming that single sigma bond so another type of hybrid is where the s orbital is hybridized with two p orbitals instead of the three that we've seen in the previous slide so we get something called an sp2 hybrid model okay so effectively one of the p orbitals is not going to be involved in forming this hybrid okay so um we found um we can find this example in alkenes and benzene as well now benzene you're going to see a lot more in year two so don't worry too much about it now i'm just going to show you as an example though um alkenes you do need to know about that okay um so here um we have um an s so what we can have is we can have three sp2 orbitals and a p orbital that's separate so remember go back to that model from topic two you have an s orbital okay which has which can hold two electrons maximum and then you have your three p orbitals that kind of sits just above it in energy so what we're doing here is using the s orbital still but we're kind of merging that with two only two p orbitals okay and that forms what we call an sp2 orbital so we have a spare p orbital that's left okay so what we have is we have these three sp2 orbitals um and the sp2 orbitals these three sit at 120 degrees apart from each other and they form this planar structure okay and have the remaining there's the planar structure there now the remaining p orbital will sit at 90 degrees so imagine this kind of sticking out at the top here and then below like that so you've kind of got this kind of looks like a windmill like a wind turbine and you've got this bit sticking out like that up and down now this one sticking up and down is your traditional p orbital that hasn't been hybridized um and the remaining s and the other two p orbitals have merged to form this new hybrid which is an sp2 and these three orbitals kind of sit in a nice kind of triangle shape like this okay so in alkenes a pair of sp2 orbitals so these two there's one there's an example you might have another molecule somewhere over here so another imagine you've got two of these side beside these two orbitals this one underneath the molecule can merge together to form a sigma bond and now what that does it helps to pull the molecules a little bit closer together and then that allows a creation of a pi bond okay so let's have a look at the diagram so here we have um an s orbital okay so your s is to kind of imagine these is it's kind of quite difficult to show but imagine this and another molecule over here they kind of overlap the orbitals overlap and we form this sigma bond here what you do have though remember you've got this kind of pi bond that's kind of sort of pi bond the p orbital that's kind of on the top here kind of 90 degrees to it that then starts to overlap with a neighboring molecule so you've got this kind of pipe the p bond so the p orbital um here and you've got another molecule near it and effectively these kind of vertical ones here they are line up top and bottom okay so this is not a hybrid version and what these do is these overlap to form this okay so this is your pi bond okay and so there's your sigma there this might kind of make it a bit easier so you've got your sigma so that's your carbon there that's your sigma bond these overlap and you've got these two p orbitals that kind of on the top and bottom on each of them and they kind of merge together they kind of mix and they form this pi bond here and that is a double bond so it looks a bit like a hot dog and a bun to an extent so you hot dogs like the sigma and the bun bits you got two bits of it which is the pie bit there okay so that is effect what a pie orbital is i hope you understand that bit okay so just go back it is quite tricky if you're not sure go back and have revisit and have a look so just to kind of look at um make it a little bit more clearer again i'm throwing a lot of diagrams here just to kind of show you what these look like hopefully it'll just reinforce it this is an example of six carbon atoms bonded together again you'll see a lot more of this benzene um you'll see more this in year two if you're not doing year two you don't need to be too concerned over it but it's really just to try and kind of show you what's happening so you've got your six carbon atoms here they're all kind of joined up in a ring and these red circles here hydrogen okay so you've got your sp2 hybrids kind of bonding here here and here so they are sp2s okay so that's your trigonal bit there you're kind of playing a bit and then you've got this p orbital that wasn't involved in the hybrid kind of top and bottom okay and you've got that with each of your carbon atoms along there now in this orbital here you've got an electron that's kind of whizzing around in this figure of at the other electrons involved in bonding now what happens with with this effect these are just single bonds as these electrons can overlap and it doesn't look like it in the diagram because it's for diagrammatic purposes to show any kind of what it looks like but these are close enough to overlap with each other top and bottom like that and effectively within benzene they kind of merge together and form this kind of um like a donut shape i suppose um and in this case and benzene's case is quite unique and again you'll don't need to be too concerned about this in year one um but they will merge together and form this and what what it's trying to demonstrate here is that orbitals overlap to form bonds and this is effectively a bit like a like a pi bond okay okay so let's kind of move away from that bit okay so that's the trickiest bit so breathe okay so um let's look at some of the areas as well which is electronegativity so electronegativity happens in covalent bonds as well so there's a lot of stuff in covalent bonds here um so electronegativity is the ability for an atom to attract electrons towards itself in a covalent bond so this happens in covalent bonds only okay so the further up and right you go in the periodic table excluding the noble gases so they've got a big black line through them and the more electronegative an element is so in this case fluorine is the most electronegative element in the periodic table and you can use this scale which is called a powling scale um quite useful scale quite straightforward and it basically tells us how how to quantify how electronegative an element is so the bigger the number the more electronegative it is so fluorine is the most electronegative element in the periodic table and so therefore has a electronegativity value of four and you can see some of these elements like oxygen for example is not quite as electronegative chlorine is less electronegative still nitrogen carbon hydrogen etc so these are the values that we give to them now essentially the bigger the difference in the electronegativity value the more ionic a compound will be okay so we've got basically if you've got sodium which is say at one end here and chlorine which is at the other clearly one's very electronegative and the other one isn't now we know that sodium chloride is ionic because you know the difference in electronegativity is so big but if we look at something serious such as carbon and hydrogen which is the c is 2.2 and 2.6 the difference in electronegativity between them two is not as great so they're more likely to be covalent um and obviously anything which has no difference in electronegativity will be purely covalent so there'll be no cover there'll be no ionic characteristics already and in reality molecules have a some covalent characteristics and some ionic characteristics some are more purely ionic than than others though so there's in reality there is a mixture of some of them there but we can kind of categorize broadly speaking molecules which are generally seen as covalent and molecules which are seen as ionic okay so let's look at some of these in a little bit more detail let's look at some polar bonds um so covalent bonds can become polar if the atoms attached to it spit it out uh have a difference in electronegativity so for example the bigger the difference in electronegativity the more polar a bond will be as we've seen before so let's look at an example here this is hcl now cl as you've seen before is is more electronegative than hydrogen and what it does is remember when we looked at the sharing of electrons between these atoms is it pulls these electrons towards itself because it's a lot more electronegative and you get this polarization this polar bond and we put these little symbols here this is a small delta so delta positive for hydrogen because effectively the electrons it was sharing is being kind of moved over towards chlorine and delta negative for chlorine here so the electrons have been moved across to one side and in fact we have a little little dipole we have a little positive charge here and a little negative charge here now this is not ionic this is still covalent because the electrons are still shared but it has a mini kind of polarity within the within the molecule itself so like i say to show that we put these um little delta positives and delta negatives next to it um but atoms with the same or similar electronegativity values are not polar so they're non-polar um and the electrons kind of sit bang in the middle so they're nicely shared um and hydrocarb hydrocarbons are clusters non-polar so things like you know methane butane ethane etc you'll you'll see that um later in the introduction to um a.s organic chemistry topic later on um but um yeah so hydrocarbons are basically um non-polar um you've also got some as well such as molecules which probably have more than one atom in there so for example water and so if you've got an uneven distribution of charge this leads to polar molecules as well so water is a classic example of a molecule that is polar so the um the electrons you can see here are being pulled towards oxygen which is more electronegative than hydrogen so you get two um sets of polarities you've got one here and one here okay and it's this property that helps water um dissolve or helps um to break up ionic compounds because of this charge here it kind of muscles its way in into the kind of giant ionic structures and breaks them up and that's why you get ionic compounds that are generally soluble um you've got other ones as well you've got to be a little bit careful with these um so um you've got carbon dioxide on the face of it might look polar because you've got an oxygen in the carbon and there is an electronegativity difference between the two however what you have here is symmetry so you've got one end pulling electrons to one end and the other one's pulling it equally but in the opposite direction on the opposite side so molecules like for example carbon dioxide they're classed as non-polar i know we have polar bonds we have polarity here but when we look at the molecule as a whole there is symmetry there so basically you've got two ends of this molecule which have got the electrons in the middle bit is kind of left exposed in the middle with very little electrons so this is is what we call a non-polar molecule so just be really cautious with that okay so let's look at some intermolecular forces um and then what we'll do is look at some metallic bonding and then that'll be it okay so it's quite quite a bit in this topic isn't it so um right so we talked a lot about bonds um before so bonds are these um attractions between atoms they're quite strong okay now intermolecular as the name suggests is inter means between and molecular is molecules so this is forces between molecules so all the other bits we've been looking at before these are um bonds between atoms or ions so it's a very different thing don't get these forces mixed up with bonds okay now there's three types of intermolecular force okay one of them is van der vaals which is dutch i believe i believe he's dutch so van der vaals or also known as induced dipole dipole um you've also got um permanent dipoles and you've got hydrogen bonding as well so you've got three types of intermolecular force now all of these forces are weaker than bonds okay so bonds are very strong and they've got a they're really difficult to break down forces are weaker than bonds okay so kind of categorize them into two containers one is bonds and one is um forces okay so look at the first one which is van der vaals now van der vaals are the weakest type of force that exists between molecules so um you need to know the criteria for these as well how do we know which which force exists where so any molecule or atom with electrons can form a dipole okay as we've kind of seen in the earlier example some of them dipoles can be permanent if they're you know if they exist some of them um you know can be induced if they come near another molecule so van der vaals if they've got electrons in there basically um it it can be induced so this occurs so for example chlorine so chlorine's got no kind of um there's no dipole in chlorine you know both elements are as electronegative as each other however we can induce a dipole so we can bring about a dipole and this is a temporary dipole and this can happen when you've got chlorine molecule floating around in a container and it comes into contact with say another chlorine molecule for example and these have electrons in them now when the electrons the electrons in this chlorine comes into contact with electrons in another chlorine molecule we get the electrons in the bond moving away because they kind of repel and for that moment in time we have this polarity so you have this kind of temporary kind of positive delta positive on one side and temporary delta negative on the other side and this is because you might have another chlorine molecule here that's kind of nudging the electrons in the molecule to one side and this is why we call it a temporary dipole now this only exists like you say when you've got two molecules or atoms nearby and when they move away when they kind of drift apart they the interaction is destroyed okay and as long as that interaction is there um you might have neighboring molecules that will have this polarity as well as temporary polarity and you have opposite subtraction on these two ends of the molecule here and so this is a very weak force because it only exists if this molecule is nearby but this is a van der waals this is a very weak force between molecules um and this is obviously a delta negative and delta positive and this is a van der waals force so this happens in in any molecule that has electrons and and obviously can form this temporary dipole okay so the next type um so if we go kind of a little bit higher in strength here these are called permanent dipole-dipole forces okay now the word permanent kind of gives you a bit of a hint here doesn't it so the van der vaals were dipoles but they were temporary so they only existed when another molecule kind of came near another one so two molecules coming together permanent dipoles basically means that the molecule itself has a has a dipole irrespective whether it's near another molecule or not i it's a permanent one so a classic example here is hcl as you can see obviously in this diagram here so um when you have hcl you have this permanent dipole here delta negative delta positive chlorine is the most electronegative element and you have these electrostatic forces between the delta negative chlorine and the delta positive hydrogen between them so the delta negative part and one molecule is attracted to the delta positive on the other one as you can see on there and unlike van der waals forces dipole-dipole interactions or your permanent ones are stronger okay because they're permanent there's no um they don't exist because they're near each other so really important though molecules such as hcl for example they do have permanent dipole-dipole forces but they also do have van der vaals they don't have one or the other these ones will have both okay so permanent dipole dipole and van der vaals really really important that you know that so um a classic example of something for example with a permanent dipole is um water it also has hydrogen bonding which i'll come into in a moment but in this example um we can you can try this at home really if you take a ruler um and rub a duster on it to basically create a charge uh on the duster and if you have a trickle of water just a trickle coming out of your tap and move the kind of um the the kind of ruler the plastic ruler to the um to the water stream you'll see the water will kind of bend slightly and that's because water is is polarized and you've got a rod which has got positive charges on it and what will happen is the um the negative bit of the water molecule will be attracted towards the positive rod that you put near it and so you'll see this trickle so there we are so try that one um and the last type of intermolecular force is hydrogen bonding now hydrogen bonding is obviously um it's got the word bond in there don't get that confused with ionic covalent metallic bonding okay hydrogen bonding is still a force it's a weak intermolecular force um now hydrogen bonding is the strongest out of the three forces as you can see on there and water is an example of a molecule that will hydrogen bond and this is why so hydrogen bonding occurs when you have an interaction between hydrogen as the name suggests and three of the most electro negative elements in the periodic table in this case it's nitrogen oxygen and fluorine so if you have any of that combination if you have a combination we've got a hydrogen and another molecule has a nitrogen oxygen or fluorine on there then you will have hydrogen bonding between the molecules and so here's an example so there we are so you've got oxygen on one water molecule with hydrogen on another um they will be um hydrogen bonding obviously we represent hydrogen bonding by drawing a dotted line um and it goes between the delta negatives and the delta positives and it must be between the lone pairs okay really really important lone pair of electrons on the oxygen so and just like we mentioned before um hydrogen bonding any molecules which have hydrogen bonding will also have van der waals forces and they'll also have permanent dipole-dipole interactions as well so water has all three intermolecular forces between the molecules okay right so again don't get them confused these are intermolecular forces they're very different to um covalent ionic and the last example we're going to come on to here is metallic bonding they're kind of two different pots okay so let's look at um metallic bonding then um so metallic bonding is um obviously a type of bonding um it happens between metals as you would probably expect they have giant metallic structures so what you need to be aware of is what that structure looks like and what the properties are for metallic bonding so positive metal ions are formed as metals donate electrons and they form this sea of delocalized electrons that kind of float around in the structure okay so if you had to look at metal that's what it would look like now you have electrostatic attractions there's that word again between your positive metal ions and your delocalized electrons which are negative so you've got this attractive force between it and the more electrons an atom can donate to the delocalized system the higher the melting point okay so obviously you have stronger attractive forces between the positive metals and the negative electrons if you've got more of them so magnesium just to give you an example so magnesium has a higher melting point than sodium because magnesium can donate two electrons which is group two whereas sodium only generates one per atom so sodium is quite soft you can cut it with a knife it's not not very hard at all there's magnesium you might have seen it it's like magnesium ribbon you put it in a bunsen flare and it goes really bright white light um metals they're really good thermal conductors as well so obviously good for pans and you know cooking things and that's because the delocalized electrons can transfer the kinetic energy remember if you do physics you'll know that if you transfer heat through conduction it's the electrons that kind of bump into each other and pass that energy along and metals are obviously no surprise they're generally good electrical conductors again because of that delocalized electron system that they have there um very similar to ions the giant ionic compounds where they have um delocalized ions they don't have delocalized electrons but they have delocalized ions and then um a d log as the electrons will also do the job as well in terms of conduction of electricity hence why graphite will conduct electricity because it has the delocalized electrons there between the layers and metals they have high melting points because of these strong electrostatic attractions and obviously solid metals are insoluble and the bond between the positive metal ion and the d local ic of electrons is far too strong to break and so therefore water can't actually break into that um and metals are also malleable and they're ductile as the ion layers obviously these can slide around and you can see they're in a kind of nice neat layer these can kind of slide around so if we take a hammer to this on this side and hammer it on that side then these layers can kind of slide over each other but the delocalized electrons kind of shift to kind of merge and try and keep it all kind of held together nicely as well so obviously you can hammer metals into shape as well which is quite a useful property okay and so just kind of moving on to the final slide i suppose um this is just summarizing everything that we've seen here because there's an awful lot here um now remember this is looking at bonding there are them intermolecular forces as well but deliberately the forces are not on here because i don't want to kind of confuse you so looking at the bond types then so you've got four main types of bonding you've got giant covalent which are known as macromoleculars so these are graphite diamonds and silicon dioxide them obviously they're they're normally solid they don't conduct electricity apart from graphite um they don't conduct electricity as a liquid either there's no ions moving there they're not soluble in water and they're really high melting points simple moleculars so these are your simple covalent structures like water and ammonia for example and iodine is another one so these are generally our liquid and gas at room temperature and pressure um iodine's are solid it's a bit of an exception they do conduct they don't conduct electricity as a solid nor is a liquid they're not really massively soluble but it depends on the polarity of the molecule as mentioned before and they have low melting and boiling points they're not giant structures like graphite or diamond and ionic compounds obviously giant ionic they are giant by nature um so they are definitely solid at room temperature they don't conduct electricity when they're solid but they do when they're liquid because the ions are free to move around they are soluble in water because of the ions water can get in between the iron structure and break it up and they have high melting and boiling points because the strong electrostatic forces between these oppositely charged ions and the final one is obviously metallic bonding um as we've just seen there before they are solid and they do conduct electricity because they have them delocalized c of electrons um they conduct electricity as a liquid as well for the same reason they aren't soluble in water and they have really high melting and boiling points because of their strong electrostatic forces so and that oops there we are and your polar molecules just kind of took one last bit in there and your polar molecules dissolve well in polar solvents like water so and that's really important as well so for example um you know ammonia will dissolve well in water and hydrocarbons don't so if you take a classic example is um cooking oil for example cooking oil is um is a hydrocarbon if you put coconut oil in water it won't mix um because they're immiscible um and that's because um hydrocarbons are not polar and generally for things to dissolve you need them you need polar and polar to dissolve okay and that is it so that is chemical bonding just as a summary make sure you know you've got your covalence your ionics and metallic bonding make sure you understand about hybridization and why what actually happens in covalent bonds make sure you understand about your electronegativities and your three different types of intermolecular forces which are van der vaals your permanent dipole dipoles and your hydrogen bonding and um that is it like i say if you these are available to purchase from the test shop um have a good look around if you wish the link is in the description box below but um that's it bye bye