Hello my name is Chris Harris and I'm from Allo Chemistry and this video is for the CIE's that's the Cambridge internationals topic three 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 or kind of you know work through them different um i suppose 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 as part of the inorganic, sorry, the physical chemistry topic with other topics as well. But it'll be clear. If you go and have a look on there, you can see them. Great for revision and on the move, et cetera, and using alongside your revision books and notes and whatever you have there.
So hopefully these 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 nice and clear. So let's start with bonding first, ionic bonding first should I say. Now some of this stuff you might have seen from GCSE, 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. um seeing 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 concepts 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 We 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 in topic two, ions are oppositely charged. So ions are charged particles, but ionic bonding is where you've got oppositely charged ionic substances or ionic 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, and 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 let's look at the first one.
This is sodium and chlorine. So these are two atoms here, and we know they're atoms because this is just showing the valent electrons. So these electrons are in the outer shell. So obviously we know that sodium is in groups.
one so it has one electron in its outer shell and chlorine is in group seven so it has seven in the outer shell now you'll probably recall i think from topic two um that in order for these two and topic one is a little bit topic one as well in order for these to form ions then 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 in bonding later on but effectively the the positive and negatives they they effectively attract to each other so just as a reminder of the other ions that can be formed so obviously this is in relation to so the common ions in groups one two three five six and seven obviously the the kind of the ones right at the end don't form any ions there 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 uh charges they can form variable ions so they don't really form the same kind of pattern we see there so um 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 minus is nitrate ions ammonium is nh4 plus sulfate ions so42 minus carbonate ions 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 by 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 fine. So you go with whatever you feel comfortable with.
But I'm going to show you this method because it may help some people. So working out the formula of substances is quite important. Obviously, it's an integral part of chemistry. So in this case, we're going to show you this method here. So we're going to look at calcium and nitrate ions, for example.
So NO3-and Ca2+. Now, if we wanted to know what is the formula of calcium nitrate, this is the way in which we go about it. Or you could go about it, should I say.
So what we do is write out the charges first, write out the ions, and then we do a swap of the charges. So we basically move the 2 plus over to the NO3 and remove the minus bit from the NO3 over to the calcium bit, as you can see there. And then we drop the charges. So we basically drop them so they're kind of subscripts. 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 NO3, though, is that we have to put that in brackets to basically say, right, we've got an NO3 kind of unit and we have two NO3s. That's why we put the brackets around there.
We can't just put two next to the three. Otherwise, look, we've got looks like we've got 32 oxygens, which just doesn't happen. So yeah, very important to put your brackets around that as well. So we simplify to the 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 substance, which is calcium nitrate. So that's the formula for calcium nitrate. Let's look at another one.
Let's look at calcium and oxygen. Follow the same method. Swap the charges over. 2 minus and 2 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 the lowest common denominator, I suppose, in which case this is CaO, so that's calcium oxide. So that's the formula for calcium oxide there. So what we also need to be aware of is what is an actual ionic structure. And a common one, I suppose, is sodium chloride.
And there's many examples of this. But ionic substances generally form these giant structures. And giant meaning exactly what it is. They're big.
They're very large structures. Now, ions, obviously, the ionic bond itself is the attraction between oppositely charged particles. It's the electrostatic attraction.
And these two here are examples of ionic bonding. Now when we write them, we write them just as CaO, 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 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, 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-ions and the yellow spheres represent sodium ions. But you can see 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 sodium ion is trying and under the chloride iron 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 you know the kind of physical properties of it um kind of stack up So, um... Effectively they are soluble in water.
So 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've got a 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.
And we can see evidence of that because we see salt dissolves in water. And... However, they also do conduct electricity when they're molten or when they're dissolved in solution as well. So if we take salt water, for example, it will conduct electricity because you've got ions that are free to move around in solution.
And that's quite an important aspect of 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. 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 irons is there's an attraction between them but if it's surrounded by loads of these irons 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 um 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. In this case, it's probably more powerful than the thermal energy required to melt it.
So it's quite strange, quite an unusual concept. So they are brittle law. Obviously, we know salt.
If you take a clump of salt and then push it with your thumb, it will probably break. So it is... you know the layers do kind of slide around quite a lot um if you hit it um and if you get two negatives together so a chloride and a chloride obviously they repel and so therefore it breaks so they are quite negative um and 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 your practicals you might not it's just a way of 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 and we're going to use copper in this case we're going to use copper 2 chromate on wet filter paper the reason why we do this is because copper 2 chromate is colored and it's green so what we do is we basically put a drop of of green copper to 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 is six minus which is here so that's your chromate um now when we switch the electricity on hey presto your copper two irons move towards the negative side and copper 2 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 2 chromate is used deliberately because it forms colored it forms colored 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 to kind of break it down core means to share and valent means the outer electrons so we're sharing the outer electrons. in this type of bonding.
There's loads of different examples here and these are probably some of the more common ones. You will see a few more as well but as you can see there's quite a few on here and I think you'll probably get the idea when I talk through them. But like I say current 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. Generally covalent bonding happens between two non-metals whereas ionic bonding happens between metals and non-metals generally. So that's to give you an idea of which ones covalent and which ones aren't. So your covalent bonds 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 and the nucleus of the atom in the middle so it's this kind of attractive force between the nucleus and these shared electrons that effectively create the bond so we have single double and triple covalent bonds 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 a bit later. That gets a bit complicated there, but for this purpose, it's just basically looking at what a bond is, what covalent bond is.
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. They're either called dative covalent or coordinate bonds. And this is an unusual bond. It's not as strong as a traditional covalent bond, generally. But this is where you have one atom that donates both electrons to another atom or ion to form a bond.
So an example you can probably see on the diagram here is ammonia, which is NH3, which you can see here. So you have your traditional covalent bonds between the nitrogen and the hydrogen atoms. You 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. Its electrons have been removed and that's why it's got a positive charge. So if we then move that across there, there we are.
This is a dative covalent bond. 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 in this case is nitrogen. This is an example of a coordinate bond or dative 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 a part of nitrogen atom and they're being donated to hydrogen in a shared way and we call that a coordinate bond okay so effectively you form nh4 plus and that's a ammonium ion which you will have seen in topic two okay So, CO, this is another one, this is carbon monoxide. It has a double bond and a covalent bond. Sorry, a double covalent bond and a dative covalent or coordinate bond. So carbon monoxide, which is that poisonous gas, it's odourless, you can't see it, it's toxic obviously.
This is, obviously the bonds here, you can't have a mixture of the two, so you can 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 another example and so this is the one where you've got a mixture of different bonds here. So this is an example of 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 this one's aluminium chloride. and actually the formula of aluminium chloride is Al2Cl6 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 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 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 that bond, the bond enthalpy, the strength of a bond, is linked to how long that bond lasts. 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 effectively 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 play here so you've got some which helps to pull the atoms together but not too close because there's constituents in the element in the atom sorry that also repel it that kind of push it away so there's a balance between the two forces and as a result we get something well we get a bond length we get a distance between you know in 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, 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 stronger overall.
And we can look at this in terms of your single, double and triple bonds. And high electron density is shared effectively between them. And shorter bonds, as we go down here, 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. etc etc 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 and we need to know some shapes of molecules as well and we need to know some of the rules associated with the shapes of molecules so let's have a look so the number bond pairs and lone pairs of electrons 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 etc it's a similar type of concept so remember in a bond you have shared electrons and electrons that shared that sharing kind of mechanism don't want to be near other shared electrons they want to try and repel themselves as much as they possibly can and that basically gives rise to a shape particularly if a 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 A lone pair next to a bond pair, these repel more than two bond pairs together. And two lone pairs together repel even further.
So if you have a molecule that has a lone pair of electrons, a bit like 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, so it had three NH bonds, and had the lone pair on the top. So that lone 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 could be ammonia So your nitrogen in the middle and then you've obviously got Your your bond pairs there as well. Okay. Okay, and there's another one which is water So this could be oxygen with two hydrogen's 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, you'll see in a moment, we call this tetrahedral. So there's a tetrahedral shape, which has a bond angle of 109.5 degrees. Very important to not miss out 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.
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 it even further. 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 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. 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 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 would 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 no lone pairs in this example. And then with ions if we have ions 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. 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. But we reduce the bond angle by five degrees because we've got two lots of lone pairs in there. So what that means is that we basically see a different structure.
We don't have something that's tetrahedral. We have something that's called nonlinear, which sounds very exciting, or bent. and 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've 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, to work out the shape of the molecule. Okay, so let's have a look. Let's have a look at the first one.
So we've got 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 an idea. So an example would be 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 planar because it's flat so that's the furthest away these atoms can be um you know within this molecule um four bond pairs and no lone pairs is called tetrahedral as we've seen before and that has a molecule sorry a bond angle of 109.5 really important you remember the 109.5 don't just put 109 down and that's called tetrahedral um five bond pairs and no lone pairs is called drink trigonal bipyramidal looks a bit odd um so effectively an example is a phosphorus pentachloride which is pcl5 so just to kind of pause on this one a little bit just to 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 planar arrangement very similar to that we've just tipped it on its side okay so you'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 bipyramidal 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 bipyramidal 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-shaped 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 sf6 for example And this is called octahedral. The bond angles are 90 degrees between all of them. Now, with octahedral, the reason why it's octahedral is, again, if we draw a line going 1, 2, 3, 4 up to these, we've got the square in the middle, and then you've got the two poles either side, then it would form an eight-sided, three-dimensional shape.
So you would have 1, 2, 3, 4. So there's your difference. your different kind of faces imagine 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 had to draw that into a 3d shape you'd 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 let's have a look at this one then so this is three bond pairs and one lone pair of electrons this is classed as 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 our 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 long pair of electrons creates 104.5 degrees. Another example so three bond pairs and two loan pairs 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 loan pairs so it's not just three bond pairs it's 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 lawn 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 is four bond pairs and two lone pairs, and this is square planar. So, for example, xenon tetrafluoride, so XEF4, it has two lone pairs, again, similar principle to the trigonal planar. 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, remains unchanged.
Like I say, 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 them. I know there's a lot to remember here. A lot of the time it's going to be practice. You must keep practicing.
Okay, so giant covalent structures. So these are effectively, obviously we looked at covalent structures just before that. And some of them are what we call simple covalents.
So they were very just small molecules. But you can get giant covalent structures. It's a bit like what you had with giant ionic as well.
So giant covalent structures, some classic examples include graphite and diamond. So if we start with graphite first. So obviously graphite is found in pencils.
So it's not lead that's found in pencils, it's obviously graphite. Each carbon is bonded three times and the fourth electron is what we call delocalized. 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. So graphite has loads of strong covalent bonds between the carbon atoms so it has a really high melting point.
If you try and melt graphite from a pencil it doesn't melt very well whatsoever. Now they do have these delocalised electrons and they kind of sit between these layers of carbon atoms which sit in rings. Now these layers are held together by weak forces. It's thin. 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 also these delocalized electrons between the layers you 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. And the layers are quite far apart in comparison to a covalent bond.
And this means that graphite has a low density. So it's not a dense material by a long shot. Say when we look at diamond, for example, which is the next example. So it's quite low density.
And it is insoluble as well. It doesn't dissolve. The covalent bonds are far too strong for water to break them apart.
So it's not like ionic compounds where generally ionic compounds are quite soluble. You know, water can break the ion structure down. It's not the case with graphite.
Okay, so let's look at diamond then. So diamond is giant covalent as well. 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.
In this case, it bonds four times in a tetrahedral shape. And obviously, you've seen tetrahedral before, just before when we looked at the bond angles or the shapes of the molecules. They're tightly packed.
Rigid arrangement allows heat to conduct well in diamonds. So they're pretty useful. They're used in... 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 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 and unlike graphite diamond can be cut to make gemstones so obviously diamond jewelry for example really expensive metal a really expensive mineral sorry not metal definitely not a metal by the way um carbon is 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 Diamond it doesn't conduct electricity it doesn't have delocalized electrons like graphite does and they're all occupied in bonding and diamond is insoluble as well so obviously it doesn't dissolve 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 diamond. So if you might see in the exam, they might talk about silicon dioxide. That has a similar structure to diamond as well, the same properties.
So 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 going to look into a little bit more detail and i would say out of all of this topic this is going to be the most trickiest and demanding part so i'm going to try and talk through it and it's clear as clearly as i possibly can um but if you're sitting there thinking this looks really difficult it's because it is okay that's not to frighten you 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 we all have to learn somewhere so I'd learn this as well you know when you know when I was there 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 looked at covalent bonds we've seen single bonds and double bonds and triple bonds 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. And p orbitals were, well, they looked like eights, a number of eights, figures of eight.
And you had three different types of p orbital. You had px, py and pz. 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 would urge you at this point to 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 say from topic two 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 two, 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 going to use an analogy, it's a bit 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 downstairs wouldn't you so it's kind of the same c orbitals in a similar way you 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 we wanted to.
So, and see it 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, we're going to use bonding with carbon atoms. And for bonding, we must, well, for bonding to occur, for 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 say and you want them passengers 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's sitting 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 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 now in reality does that happen i don't know most people probably don't they just keep themselves to 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 a singly occupied orbital. You can't have it where it's fully occupied. And it must be singly occupied. I hope you get that so far. So there are two ways in which a CH bond can be formed.
And here I'm just going to kind of talk through the two main ways. But there's obviously one way in which it does actually kind of form. And that's what I want to talk about here. So... So 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 one s orbital because it only has one electron in there. So it's singly occupied. And with your carbon, you have six electrons in total carbon. So you have some in the p orbital as well.
So you can overlap the two p orbitals with the s orbital in hydrogen. And there we are. You've 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, 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, you might have seen cars which are hybrid cars. Okay, so cars are either petrol, petrol or diesel, so they 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 sp 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 you 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 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 you 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 in its outer shell, so its valence shell, obviously it's got a 1s orbital, it's got two electrons in the 1s, but we're just really looking at the outer ones because these are what are 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. So this is what this shows here.
So you've got 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 you 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 2pz seat.
Yeah, quite fancy. 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 you 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 to form these bonds is four times four one two. Okay.
And this expels minus 1,648 kilojoules per mole. And the four bonds, they form this new hybrid. So these are not SP, PX, PY, PZ. 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, a petrol. Sorry, we have an electric car here and a petrol and petrol and petrol. In fact, 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, 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 404 but we release 1648 kilojoules per mole of energy when we form four new bonds so the total energy change is minus 1244 now it doesn't take einstein to work it out that actually um you get a much bigger energy release um with doing it this way in other words moving an electron into a different orbital first and then forming four bonds than you do by just overlapping existing bonds here. Now remember atoms are incredibly lazy, chemistry is like the lazy science I like to call it. 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 the 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 1,244 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...
with... hydrogen so it normally exists as CH4 rather than just CH2 because of this setup 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 and effectively that is where you have an S orbital an S orbital merging with two P orbitals so that's SP2 but you also have SP orbitals where you've got one S orbital and one P orbital 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 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 it in a little bit more detail. So that was a single bond, a single CH bond.
You do need to be aware of sigma and 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 sigma bonds. Some with S and P characteristics, hence why we say SP3 hybrids. So the bond is effectively them 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. So 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. Alkenes you do need to know about that. So here we have 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 sit 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. And the sp2 orbitals, these three sit at 120 degrees apart from each other. And they form this planar structure.
Okay. And now 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. 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, 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-by-side these two orbitals this one and a neighboring molecule can merge together to form a Sigma bond and not 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 an s orbital Okay, so your S's, imagine these as, it's 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, sorry, 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 neighbouring molecule. So you've got this kind of pi, the p bond, so the p orbital here.
and you've got another molecule near it and effectively these kind of vertical ones here they align 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 the 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 your hot dogs like the sigma and the bun bits you got two bits of it which is the pi bit there okay so that is in fact what a pi 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 a revisit and have a look you So just to kind of look at and 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. You'll see more of 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.
These red circles here, they're hydrogen. So you've got your sp2 hybrids kind of bonding here, here, 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 you carbon atoms along there now in this orbital here You've got an electron that's kind of whizzing around in this figure of eight here the other electrons involved in bonding now What happens with with? This effectively these are just single bonds is these electrons can overlap. I know it doesn't look like it in the diagram because it's for diagrammatic purposes just showing you 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, like a donut shape, I suppose. And in this case, in benzene's case, it's quite unique. And again, you don't need to be too concerned about this in year one. But they will merge together and form this.
And what it's trying to demonstrate here is that orbitals overlap to form bonds. And this is effectively a bit like a pi bond. Okay, so let's kind of move away from that bit. Okay, so that's the trickiest bit. So breathe.
Okay, so let's look at some of the bare 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.
So, electronegativity is the ability for an atom to attract electrons towards itself in a covalent bond. So, this happens in covalent bonds only. 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, the more electronegative an element is. So, in this case, fluorine is the most electronegative element. in the periodic table and we can use this scale which was called a powering scale um quite a 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 if 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 the difference in electronegativity is so big. But if we look at something, say, such as carbon and hydrogen, which 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.
And obviously anything which has no difference in electronegativity will be purely covalent. So there'll be no ionic characteristics already. And in reality, molecules have some covalent characteristics and some ionic characteristics.
Some are more purely ionic than others though. So 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 so covalent bonds can become polar if the atoms attached to it spit it out 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 has been kind of moved over towards chlorine and delta negative for chlorine here so the electrons have been moved across to one side and effect 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 that within the molecule itself So, like I say, to show that, we put these little delta positives and delta negatives next to it.
But atoms with the same or similar electronegativity values are not polar, so they're non-polar, and the electrons kind of sit bang in the middle, so they're nicely shared. And hydrocarbons are classed as non-polar, so things like, you know, methane, butane, ethane, etc. You'll see that later in the introduction to... as 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 um 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, as you can see here, are being pulled towards oxygen, which is more electronegative than hydrogen.
So you get two sets of polarities. You've got one here and one here. And it's this property that helps water dissolve or helps to break up ionic compounds because of this charge here.
It kind of muscles its way into the giant ionic structures and breaks them up. And that's why you get ionic compounds that are generally soluble. You've got other ones as well. You've got to be a little bit careful with these. So you've got carbon dioxide on the face of it.
It might look polar because you've got an oxygen and a 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 and then what we'll do is look at some metallic bonding and then that's would 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 waals which is dutch i believe i believe he's dutch um so van der waals um 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'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 forces. Okay, so look at the first one, which is van der Waals.
Now, van der Waals are the weakest type of force that exists between molecules. So you need to know the criteria for these as well. How do we know which force exists where? So any molecule or atom with electrons can form a dipole. As we've kind of seen in the earlier example, some of them dipoles can be permanent if they exist.
Some of them can be induced if they come near another molecule. So van der Waals, if they've got electrons in there, basically it can be induced. So this occurs, so for example chlorine. So chlorine's got no kind of.... 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 you 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 I say, when you've got two molecules or atoms nearby, and when they move away, when they kind of drift apart, the interaction is destroyed.
Okay, and as long as that interaction is there You might have neighboring molecules that will have this polarity as well. There's temporary polarity and you have Opposites attracted 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 And this is your Delta negative and Delta positive and this is a van der Waals force So this happens in any molecule that has electrons and obviously can form this temporary dipole.
Okay, so the next type, 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 Waals 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 on 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, 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 waals they don't have one or the other these ones will have both okay so permanent dipole dipole and van der waals 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 rubber duster on it to basically create a charge uh on the duster and if you have a trickle of water just the trickle coming out of your tap and move the kind of um the kind of ruler the plastic ruler to the um to the water stream you'll see the water will kind of bend slightly um and that's because water is polarized and you've got a rod which has got positive charges on it you What will happen is the negative bit of the water molecule will be attracted towards the positive rod that you put near it. You'll see this trickle. There we are.
Try that one. The last type of intermolecular force is hydrogen bonding. Hydrogen bonding is obviously, it's got the word bond in there.
Don't get that confused with ionic, covalent, metallic bonding. Hydrogen bonding is still a force. It's a weak intermolecular force.
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 electronegative elements in the periodic table and 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 they will be hydrogen bonding obviously we represent hydrogen bonding by drawing a dotted line 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, 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 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 metallic bonding then. So metallic bonding is obviously a type of bonding. 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 donates one per atom. So sodium is quite soft. You can cut it with a knife.
It's not very hard at all. There's magnesium. You might have seen it. It's like magnesium ribbon.
You put it in a bunsen flame and it gets really bright white light. Metals, they're really good thermal conductors as well. So obviously they're good for pans and 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.
Metals are obviously no surprise. They're generally good electrical conductors. Again, because of that delocalized electron system that they have there.
Very similar to ions, so giant ionic compounds where they have delocalised ions. They don't have delocalised electrons, but they have delocalised ions. Then delocalised electrons will also do the job as well in terms of conduction of electricity.
Hence why graphite will conduct electricity, because it has the delocalised electrons there between the layers. Metals, they have high melting points because of these strong electrostatic attractions. And obviously solid metals are insoluble. The bond between the positive metal ion and the delocalised sea of electrons is far too strong to break, and so therefore water can't actually break into that. And metals are also malleable, and they're ductile, as the ion layers, obviously these can slide around.
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 it 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.
This is just summarizing everything that we've seen here because there's an awful lot here. Now remember, this is looking at bonding. There are them intermolecular forces as well, but deliberately the forces are not on here. 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 obviously they're they're normally solid they don't conduct electricity apart from graphite they don't conduct electricity as a liquid either there's no ions moving there they're not soluble in water and they've really high melting points you Simple molecular, so these are your simple covalent structures like water and ammonia for example and iodine is another one. So these are generally how liquid and gas at room temperature and pressure, iodine is a solid, it's a bit of an exception.
They don't conduct electricity as a solid nor as a liquid. They're not really massively soluble but it depends on the polarity of the molecule as I mentioned before. and they have low melting and boiling points.
They're not giant structures like graphite or diamond. Ionic compounds obviously are giant ionic. They are giant by nature so they are definitely solid at room temperature.
They don't conduct electricity when they're solid but they do when they're a liquid because the ions are free to move around. They are soluble in water because of the ions. Water can get in between the ion 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. As we've just seen there before, they are solid.
They do conduct electricity because they have them delocalized sea of electrons. 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... Just... kind of chuck one last bit in there and your polar molecules dissolve well in polar solvents like water so um 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 uh cooking oil for example cooking oil is um is a hydrocarbon if you put cooking oil in water it won't mix um because they're emissible 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 it actually happens in covalent bonds. Make sure you understand about your electronegativities and your three different types of intermolecular forces, which are van der Waals.
your permanent dipole dipoles and your hydrogen bonding and that is it like i say if you these are available to purchase from the test shop have a good look around if you wish the link is in the description box below but that's it bye bye