hey guys welcome back to pop-mop cam sorry about the delay on some of the uploads in this video we're going to be starting unit 4 14 vip specification that structure and bonding and we're going to start it off with looking at covalent bonding so what we're going to do is we're going to define covalent bonding look at how electronegativity can affect covalent bonds and then we're going to contrast simple and giant covalent structures firstly let's think about bonding in general so there's kind of two major parts of bonding there is chemical and physical bonding so chemical bonding includes covalent bonds which we're doing in this video ionic and metallic bonding which will be coming later in the unit physical bonds are intermolecular forces sometimes called van der waals forces these split into three main categories london forces which are also called induced dipole instantaneous dipole and sometimes called dispersion forces we also have dubai forces which are also called induced to permanent dipole forces and lastly we have keysome forces which are also called dipole dipole forces or permanent dipole to permanent dipole forces and can also include hydrogen bonds for now though let's bring our attention back to covalent bonding so how do we define a covalent bond well a covalent bond is the attraction of two atoms to a shared pair of electrons we can conceptualize this as thinking of two nuclei with outer electrons and we know they want to complete their outer shell and so they may place these electrons for one to the better word between each other and what that does is that creates a mutual attraction between the nuclei and the electrons that are located in between the two nuclei and we can think about this as completing their outer electrons if we think about here we've got a 1s orbital and so that wants two electrons to complete it so we can think of that as each nuclei adding an electron to its outer shell and therefore each nuclei having a complete shell when they are paired together in this system hence why we describe them as shared electrons typically we see this kind of bonding between non-metals to help us predict these kind of bonds we can use the octet rule and what the octet rule states is the atoms usually form bonds based on the octet rule that is they want to have eight electrons in their outer shell we could also think about this as atoms forming as many bonds as they have gaps in that octet for example if they have seven electrons they have one gap and so can form a single bond something in group five would have five electrons in the outer shell and therefore would have three gaps and would have the ability to form three bonds now these bonds can organize themselves in different ways we can have single double triple and even dative bonds for example in cl2 we have two atoms with one gap so they form a single bond which is represented by a line between them which indicates that they have two shared electrons between them and so is a single covalent bond in co2 however we have oxygen which has six outer electrons and carbon which has four so carbon can form four bonds and oxygen conform two which leaves these two lines representing double bonds and that is four shared electrons between each of the carbon and oxygen atoms however oxygen can also form two single bonds as we said it has six electrons in its outer shell and so has two gaps so it forms two bonds but they could be in a single configuration or in a double bond configuration as i said before carbon can have four gaps and so make four bonds so here we have carbon bonded to another carbon with three lines representing a triple bond or six shared electrons and obviously one single bond with hydrogen to form a total of four overall bonds you'll notice that that is congruent with co2 which we did just above where the carbon has two double bonds either side still totaling four overall bonds now there are cases in which some atoms can contribute both of the overall electrons that are in the covalent bond between the two atoms this is called a dative or coordinate bond we can use the example of the nitrogen in ammonia remember nitrogen has five outer electrons and so in nh3 it forms three single bonds gives one electron to each of those and then has a lone pair all to itself and what this lone pair can do is it can find a h plus and it can actually donate both of the electrons to a bond between the nitrogen and the hydrogen and this forms nh4 plus which is the ammonium ion and we represent that with a arrow going towards the hydrogen atom from the nitrogen now it is important to mention that these two bonds are identical once they are formed they have the same bond enthalpy they have the same length they have the same strength so they are not different once they are formed it's only the formation that is different and then we put some square brackets around it with a plus to indicate that this is an ion and has a charge now the bonds get stronger as we go from single double to triple bonds so we get an increasing bond strength however we get the opposite when we look at the length we get a decreasing bond length as we go from single double to triple now just looking in a little bit more detail at coordinate or dative bonding as i mentioned of course the dative bond differs only in its formation the rest of the bond is the same we can also say to watch out donor species will always have a spare pair of electrons and acceptor species will always have an incomplete octet looking back to our formation of ammonium ion example we can see that the nh3 before bonding has a spare pair of electrons so it is the donor and the acceptor is h plus which has an incomplete 1s orbital so the nitrogen can donate both electrons and form that dative bond so like any good rule there are indeed exceptions to the octet rule now carbon nitrogen oxygen and fluorine always obey the rule and these make up a certain majority of compounds so that's good there are two types of exceptions firstly we'll look at reduced octet exceptions these are atoms that will happily be stable with less than eight now the first most obvious one is hydrogen because this will complete its outer shell with one s2 shell and so therefore two is actually a complete outer shell lithium will also be happy to just have a 1s2 filled outer shell beryllium will happily just have four electrons in the outer shell so actually half an octet just like in bef2 another common reduced octet violator is boron which has three electrons in the outer shell to begin with and will also be quite complacent with just three bonds or six electrons in the outer shell just like in bf3 or bh3 and boron is actually a really common one that's used in questions in exams so watch out for that one aluminium in the same group obviously as boron can also form these reduced six electron outer shell now alcl3 is an example of this although it would actually exist in the dimer form not as i'm drawing it here just to illustrate that there would be three bonds to that central aluminium the other type of exception is of course an expanded octet these are atoms which readily form compounds that have more than eight electrons to the central atom main culprits here are the three p section of the periodic table that's from silicon phosphorus sulfur and chlorine and the reason for the ability of these atoms to form an expanded octet is because there is the low energy 3 d orbital which is very near the 3p orbital in energy and so extra electrons can kind of be shuffled up into that orbital and so you can have something that has more than eight electrons and still be stable so3 is an example of this in which we have sulfur in the center bonded to three oxygen atoms to which each it has a double bond giving us a total of 12 electrons seen by that central sulfur atom we'll actually deal with this in a little bit more detail when we look at formal charge now the period two elements cannot do this because they don't have that 3d orbital or even and there's no such thing as the 2d orbital just above for them to put those extra electrons in so what we end up with then is something a little bit more like the octet rule of thumb than a rule now this sharing can sometimes be completely equal between two elements however electronegativity that's the relative attraction of an atom of an element to a shared pair of electrons dictates how much the electrons lie towards one nuclei or the other you can look at these values at taper 8 in the data booklet for example if we have fluorine and fluorine the electron cloud is going to be shared equally over both nuclei as they both have the same value for electronegativity however if we look at hf and we look at table 8 we see that fluorine has a much higher value of electronegativity and so therefore the electrons lie much closer to the fluorine atom and this gives us an unbalanced electron cloud which lies towards the fluorine this is what we call a polar covalent bond and we represent this using the lowercase greek letter for delta which indicates that there is a partial negative or partial positive charge and because in this molecule that partial charge is unbalanced we end up with an overall dipole moment that means that this molecule in general has the electron shifting to one side giving us a dipole dipole just meaning two sides so we have a positive and a negative side so we have polar bonds caused by a large difference in electronegativity and non-polar covalent bonds which are caused when there is similar or the same electronegativity now you don't have to remember these values and you can also use table 29 in the data booklet to actually be able to calculate whether you're going to have a polar or non-polar covalent bond you can use the average electronegativity and the difference in the electronegativity to identify which type of bonding you would have based on this table for example if you had an average of 2.5 and a difference of one draw two lines across or use a ruler and you'll see that they cross over right in the middle of polar covalent so that bond would be polar covalent we can have a look at an example let's say co2 what type of bonding would we expect in co2 well we can use our two tables we look at the electronegativity of carbon and oxygen which is 2.6 and 3.4 respectively and then we take the average of those two values which is 2.6 plus 3.4 over 2 which equals 3.0 and we want the difference between them which is 3.4 take away 2.6 which is 0.8 we can then use the table 29 to find the type of bonding using those two values now there was also a kind of side note that in general differences between 0.5 and 1.8 are usually polar covalent which is what we see here from the graph the crossover point being just inside the polar covalent area looking at another example of ch4 we look for carbon and hydrogen that's 2.6 and 2.2 we take an average which is going to be 2.4 then we find the difference which is 2.6 minus 2.2 which is 0.4 and then we use those values on the chart just as we did before and now we can see that we have a non-polar covalent bond which we may have guessed by the very small difference in electronegativity so going on from what we mentioned on the last slide any difference between 0 and 0.5 is usually non-polar covalent bonds just as anything above 1.8 is usually an ionic bond which we'll come to later in the unit let's try some questions having the data book with you is going to be helpful here first question determine the nature of bonding you'd expect between carbon and oxygen pause the video give yourself a moment so once again finding the average and the distance by using table 8 2.6 and 3.4 gives us an average of 3 and a difference of 0.8 which is going to be a polar covalent bond okay have a go at doing the same but with lithium and oxygen this time pause the video pop em up so lithium having electronegativity of 1.0 and oxygen being 3.4 we take the average 3.4 plus 1 over 2 is 2.2 and the difference is 3.4 minus 1 which is 2.4 that is going to be an ionic bond looking beyond the individual bonds now we can look at how the organization of covalent structures can occur this happens in two groups we have simple covalent compounds and giant or macromolecular covalent structures simple can be broken down into polar and non-polar simple covalent compounds so let's have a look at the physical properties of these kind of structures firstly they do not conduct electricity now the solubility will depend on whether they are in an organic or inorganic solvent and their structure relating to that their boiling point is generally low because they're not extended molecules and the only thing that is holding them together are weak intermolecular forces which are easy to break we'll have a look at those more in detail later in the unit so they don't conduct electricity for the same reason there's no way to carry charge throughout the compound because they are separate molecules and that means that it's easy to separate the molecules so they're likely to be gases or liquids at room temperature in fact i2 which is iodine is one of the only simple covalent structures that is a solid at room temperature and that's mainly to do with the intermolecular forces again which we'll cover later in the unit and it sublimates which is also to do with these intermolecular forces that when they are overcome we go straight from a solid to a gas it's actually quite a beautiful purple gas so giant uh covalent structures such as diamond graphite silica etc these differ slightly in that they don't have individual discrete molecules instead we have these extended lattice structures which have covalent bonds that go on and on forming a kind of infinite covalent lattice in which every single bond is a covalent bond being very strong that means that all of these bonds require lots of energy to break them which gives us a different set of properties they still don't conduct electricity because there's no way for them to pass electrons or charge through the molecules except in the case of graphite which we'll look at they have very high melting and boiling points again because if we want to melt or boil these we have to break every single carbon bond and they also are very hard again apart from graphite because of the very strong covalent bonds in that lattice arrangement so the next covalent structure we're going to have a look at is silica and silica actually resembles the diamond structure in many ways because it has this tetrahedron like shape except now we have both carbon and oxygen atoms composing the overall lattice structure but still all bonded by these individual very strong covalent bonds but now we have oxygen atoms in between each of the silicon atoms so this is kind of one repeating unit if you like and then we would have the same kind of tetrahedral continuation further along the lattice structure as we had with diamond but with the oxygen and silicon repeating now with the empirical formula sio2 sometimes you can look at this overall repeating structure that i've drawn here and think well that looks like it's a one-to-one relationship between silicon and oxygen however if we look a little more closely what we see if we think of the black as the silicon and the orange as the oxygen every silicon is bonded to four other oxygen and every oxygen is bonded to two other silicon and hence we have the ratio sio2 a two to one relationship silicon dioxide of course being the major component of sand and uh many rocks and this has similar properties to diamond i.e a very high melting point and boiling point and they're very strong and that again comes from this infinite tetrahedral lattice that is formed from the silicon and oxygen bonds now because there's no mobile electrons or ions there is no electrical conductivity through silica just like in diamond now as i mentioned diamond i've already alluded to the fact that there are different forms that carbon can take and these forms are called allotropes so an allotrope is where there's two or more different forms which an element can exist so for carbon this is diamond graphite charcoal buckminster fullerenes we also have things like graphene and carbon nanotubes but basically they're all carbon and the only difference between them is the way in which the atoms are bonded together and the structure that they form this means we have an incredibly high melting point in fact carbon doesn't have a boiling point because it sublimates at over 4 000 degrees centigrade that means going straight from solid to gas it's very strong and very hard indeed the hardest substance that humans know of and there's no free electrons and so it is not an electronic conductor now graphite is another carbon allotrope so still made completely of carbon however the structure is very different and this structure leads to a couple of significantly different properties than we would expect compared to that of diamond we basically form these six carbon rings which are bonded in an extended lattice as i've shown here and instead of having four bonds between each carbon we have three bonds between each carbon and then those layers are on top of each other so that extra electron that carbon has instead of forming that fourth bond kind of gets delocalized between the two layers so it still has a very high melting point and sublimates just like that of diamond however this delocalized layer of electrons causes it to be soft and that's why graphite is used in the lead of pencils because it is easy to smear effectively that layer of graphite onto the page which is what we associate with a pencil mark that is because in between these two layers we only have weak intermolecular forces we don't have those strong covalent bonds that we see along those three other bonds between the carbon atoms we have these interactions and those interactions are much easier to break and require much less energy allowing the layers to easily slide over each other the other thing that these delocalized electrons allow is they allow the conduction of an electrical current through the substance hence why graphite is a good electrical conductor now if we take just one of these layers of graphite we end up with graphene which quite simply is just one layer of graphite but changing some properties interestingly raising the melting point even further to sublimation above 5000 degrees it's also the strongest material that's ever been tested is transparent and it conducts electricity extremely well and has in many applications been touted as a potential super material for some applications and carbon doesn't stop there indeed fullerenes buckminster fullerenes that were first theorized before they were discovered are these kind of football shaped where we have five rings and six carbon rings all joined together to form this large sphere have a reasonably high melting point of 600 degrees but these are individual discrete molecules so they don't exist in the same way as that long extended lattice structure of the other allotropes of carbon and this makes them quite soft because these individual molecules are only held together by weak intermolecular forces and so these individual ball shaped molecules can slide over each other quite easily for the same reason it does not conduct electricity because there's no free electrons however it does have a really important property which that it can kind of be opened and formed into these carbon nanotubes which again also have been touted as an important potential future material before we finish up then let's get our whiteboards ready and try some questions firstly explain using the structure of the allotropes why graphite can conduct electricity but diamond cannot pause the video to have a go at that pop em up so hopefully remembering that graphite has three moving electrons between its layers which allow it to carry charge next question why does silica have a higher melting point than iodine pause the video to have a go at that pop em up so of course silica having a giant covalent structure and iodine only having a molecular or simple covalent structure and so therefore requiring much less energy to overcome those intermolecular forces okay guys no practical to go with this and no questions just yet as we're going to do lewis structures in the next video and then we'll try and apply our knowledge to some problems thanks again for joining me guys don't forget to like subscribe hit the bell icon you know the drill as always practice makes slightly better