Welcome back to a new chemistry lesson today we are finally going to see the vaner bals forces and for those who missed the previous videos Obviously I recommend you go and watch and resume the videos regarding the hydrogen bond and in general all the intermolecular forces what are intermolecular forces also called intermolecular bonds are bonds that is, attractive forces different from the bonds that instead occur inside the molecules and therefore between the atoms these types of forces are much weaker than the forces between the atoms as forces between atoms we cite for example the covalent bond So these forces are much weaker than a covalent bond however they are forces that influence the macroscopic characteristics of matter that is the characteristics that we see of the substances that surround us As always I remind you that you can support the channel in many ways by subscribing to YouTube or patreon or by making a single donation for example via PayPal I leave you all the links in the description I also leave you on the screen the two references for Instagram and tiktok if you want Follow me there too And as always at the beginning of each video I thank all the people very much who already support my work Thank you because you allow me to continue working and with my work help so many students and now we can start let's do a small summary of the intermolecular bonds so bonds between molecules between these bonds we have the hydrogen bond the pole- pole bond the permanent pole-induced pole bond instantaneous pole-induced pole bond we have seen that these are the bonds that we can define as intermolecular we have already seen hydrogen bond in the previous video and let's see how the various bonds that we will see Today we can call them Vaner Wals forces By the way, under each force I have also put the name with which it can be found in some books such as Deb's Kisom force or London because we separate these forces from the hydrogen bond I have already said it in the previous video but the hydrogen bond is highly specific and above all it is much stronger than these other forces These are attractive or repulsive forces between weaker molecules but of the bonds both between the atoms but also of the hydrogen bonds Today we will see them one by one and so let's start from the first one The first force we are going to see is the bond defined by Po or of a pole, in reality the most correct name would be permanent dipole, permanent dipole, this is because it occurs between molecules that have permanent dipoles, that is, which are permanently definable as dipoles. If you have no idea what a dipole is, go and watch the video on the polar bond and I also recommend the one on the vsepr theory. This is because in these two videos I am going to define first with the polar bond. What is a permanently polar molecule and then with the VSP theory we are going to explain why some molecules are definable as polyi, that is, they have two poles, a more positive part and a more negative part. Here, for example, in the figure we see. What is a dipole, that is, a permanently polar molecule, eh, we define a polarity that does not change. This is because the molecule we see is hcl hcl has a part of the molecule that is made up of the hydrogen atom, the other part of the molecule is made up of the chlorine atom. Obviously, it is a diatomic molecule and therefore it only has two atoms, but the difference between these two atoms is important because the Chlorine is strongly electronegative, hydrogen is moderately electronegative, so we can say that it is slightly more electron positive than chlorine. This leads to an unbalanced molecule in terms of charges. Because we will have electrons that go a little more towards chlorine and a little less towards hydrogen. Remember, the electrons carry a negative charge. Consequently, the part of the molecule where the chlorine is will be affected by these electrons that move and will take on a slightly more electronegative charge than the molecule where the hydrogen is. It will be slightly more positive due to the fact that the electrons have left. Now, I must point out when we talk about these forces that when I say that the electrons move or leave an atom, this is not entirely true because the reality of things is that the electron spends more time circling around chlorine, for example, than around hydrogen. This is because, remember, when the electrons form a covalent bond, they are then slightly delocalized on the molecule, that is, we can say that they are not fixed but become electrons that rotate on molecular orbitals. they make slightly wider, slightly more particular rotations than the atomic orbitals. Obviously these molecular orbitals vary from molecule to molecule. We will see later in organic chemistry that there are very particular molecular orbitals. But now we are on the basic, we are on the simple and we see diatomic molecules. Now what happens is that every time I say that the electrons are more around chlorine than around hydrogen, it means that these electrons spend more time around chlorine, that is, they make many rotations around chlorine. And every now and then they make a little rotation around hydrogen. Many rotations around chlorine. And every now and then one around hydrogen. I am simplifying a lot, but this is a little bit, we can say the simplification of a real situation. We must also remember that when we define the electronic orbitals, we are talking about probability zones in which we can find these electrons, that is, everything I am telling you must always keep in mind that we can find the electrons more around chlorine. But it is a very large probability that we can find them around chlorine. There is also a small probability that those electrons could be found around hydrogen. And so far everything is normal, but if I told you that there is... even a minimal probability that the electrons of the hydrogen of the chlorine are found near another molecule and I'm not saying something totally wrong because even with very low probabilities we have a probability of finding the electrons even outside or far from the molecule that we are examining this is because when we go to define the position of the electrons we do not have absolute certainty everything This obviously depends on the theory of electronic configuration and quantum physics Among other things and therefore we are satisfied with saying that the electrons in this molecule revolve a little more time around the chlorine and a little less time around the hydrogen having defined all this preliminary information which in my opinion is essential to understand this part of the argument we can say that the molecule formed by H ecl therefore Hydrochloric acid is a permanent dipole That is, it has two poles which always remain one negative and one partially positive and because of this the partial dipole the partially positive pole Excuse me is attracted by the partially negative pole of another molecule now this is very important to understand that is, the hydrogen of this molecule is partially attracted by the chlorine of the other molecule but is not detached from the first molecule then in reality in some particular situations This is an acid so it can certainly detach itself from another chlorine But we can say that in the situation we are seeing that we are examining, chlorine attracts the hydrogen of another molecule and therefore we have molecules formed by HCl which however attract each other the positive pole of a molecule with the negative pole of another molecule very simple pole-pole bond very similar to the hydrogen bond and only a little weaker we can say that the hydrogen bond is a stronger and more specific pole-pole bond simply this having defined the pole-pole bond and having understood the basics we can move on to the other forces of vandals such as for example the permanent pole- induced pole bond also called Deb force Here in this case as we go forward with these bonds we are talking about increasingly weaker bonds because for example this type of bond is established between a polar molecule a permanent dipole so for example also HCl that we saw before and a non-polar molecule and you say But How is it possible if the molecule is non-polar it cannot have two poles which will then attract the pole of another molecule. That is, to make it simpler, the positive pole that you see here in the figure. How can it attract the negative pole of a molecule which does not have one because it is non-polar? Let's see that a dipolar molecule can actually induce a polarization of a polar molecule. And in fact it is called a permanent pole- induced pole bond because the polar molecule is induced to polarize and obviously when it polarizes it can be attracted by the molecule which was already polar and therefore what is an induced dipole is formed which when this main molecule moves away will then be lost because then obviously the molecule will return to being polar but at that moment it is a dipole. It must also be specified that obviously the induction of a polar molecule to become a dipole therefore a polar molecule is a temporary induction, that is, we will now see forces which are gradually weaker but also temporally weaker therefore which last less because obviously the induction of a dipole exists as long as the dipolar molecule is close to the polar molecule Because the moment this moves away this molecule will return to being apolar And now I will try to explain it to you with a drawing We have a non-polar molecule obviously Forgive me Eh the drawing is a bit ugly I'm doing it with the graphic board but on PowerPoint it's not easy this non-polar molecule is approached or rather it is found close to a molecule instead that we can define as polar so a molecule that has a positive pole and a negative pole now what happens what happens is that the negative pole of this molecule means that because a molecule has a negative pole because it means that it has many electrons that rotate in that point and I represent the electrons in this way that is there are many electrons that rotate in this point now what happens is that the electrons carry a negative charge when this molecule approaches the non-polar molecule the polar molecule has electrons that rotate throughout the molecule in a fairly equal way it's not that they group together near an atom or they rotate closer to an atom what happens however is that when they are close to other negative electrons with negative Neo repel each other so the electrons that are very close to the negatively polarized area of the molecule when they go close to the polar molecule they go to move the electrons of the polar molecule to the other side of the molecule so let's say so they go They go to create a repulsion of those electrons and the molecule finds itself with the electrons moved to one side now what happens is that if the electrons of that molecule that was non-polar are moved to one side of the molecule and well the molecule at that point has become polar this because obviously by moving the electrons from one side on the other side there will be fewer electrons and therefore We will find a charge that we can define as partially positive on the other side a charge that we can define as partially negative And here we have induced a dipole that is we have polarized a non- polar molecule How with a polar molecule therefore a polar molecule can do this thing to a non-polar molecule And here Obviously having a dipole and in this case another dipole however induced we will have a negative area From this positive side, from this other side they will attract. As already said, however, this force is weak and extremely limited in time. We can however go and see an even weaker force and even more limited in time, that is the bond of instantaneous pole and induced pole, also called very commonly, you can find it in all the books of London forces because I prefer to call them with the names of instantaneous pole and induced pole because it is much easier to understand. In the books very often a huge summary of this topic is made and it is not very clear why these forces act and why they are divided in this way, calling them instead instantaneous dipole, induced dipole and so on all the names that I told you. Go and understand perfectly what happens between these molecules that I will now explain to you also regarding the last last force, among other things, therefore the London force occurs between apolar molecules. So here we no longer have permanent poles, we have totally apolar molecules that if we study them separately are apolar but when they are found in everyday reality and in a substance they create temporary dipoles now we're going to see how I always explain it to you with the blackboard I have to tell you however that these forces are precisely the reason why certain types of substances that we can define as non-polar molecules behave in a non-ideal way that is, we do calculations we study how ideal gases should work and then we notice that there is always some congruence that is, a force that we have calculated in one way comes out slightly different from what we have calculated this because maybe we have not calculated the London forces which are very weak forces but on a large scale they have a big enough influence to explain why obviously imagine that if there is a molecule that goes to create a force with another molecule and creates a London force this force is quite negligible because it is very weak as I told you but imagine that there are 10 molecules that create this London force well you understand that all the London forces of 10 molecules are already a little more relevant than the force that occurs between two molecules now imagine that in a glass there are billions and billions of molecules So in a glass in a container with billions of billions of molecules, billions of billions of London forces are established and at that point it is not so negligible they are forces that in everyday reality we will see in the next slide have their relevance but now let's go and see how you can create an instantaneous aneous temporal dipole that then goes to polarize another polar molecule creating an induced dipole let's take a molecule let's take the diatomic For example it could be oxygen it could be any other type of molecule in this molecule the electrons rotate in a fairly uniform way around the whole molecule We have however said that the electrons rotate around the molecule and have a probability of being here or of being here the same thing obviously can happen to another molecule of the same substance because we perhaps have a container with many molecules of this substance and also in this substance the electrons could be found around both atoms that make up the molecule in a fairly indiscriminate way in a fairly equal way they could be found in this point as in this point as in this point Let's say that in any case the molecule is non-polar so it does not have a difference in charge therefore the electrons are not all on one side or all on the other Now however I mark in red the specific case and that is since the electrons rotate it can always be there imagine a bit like the traffic of cars there are maybe 100 cars on a road that contains them and generally the traffic we can say that it is quite regular however at the moment in which all the cars i.e. all the electrons are on one side and it can always happen let's say that the electrons are all on this side for some reason only one remains on this side This one here the others are all on this side of the molecule you understand well that obviously an instantaneous phenomenon So here I underline it again instantaneous but at that instant Effectively all the electrons are on that side and then this side of the molecule is weakly polarized the molecule takes a slightly negative side this however leads as we saw before being close to another molecule of the same type It causes the electrons of this molecule to move away from this side going to accumulate also in this molecule On this side and you can see clearly that in this case we have created an induced dipole obviously on the other side where there are no electrons we can symbolize it with more and we have created what I now draw for you in a dotted way which is precisely the London force And that is this electromagnetic traction due to the fact that an instantaneous dipole is created which is totally normal in the everyday life of a molecule and then consequently an induced dipole between the molecules that are close to it Obviously the molecule which is an induced dipole can go on to induce another nearby molecule and so here the forces in London take on a non-negligible importance As we said before and precisely on the basis of this non-negligible relevance we see a real life case in which the van der Wals forces come to act in a rather important way here in the slide In fact You see a Geo What happens with the Jechs although the van der Wals forces seem to be confined to a purely academic, i.e. scholastic, field and of little relevance in reality they have a notable importance in the reality that surrounds us For example in the Jackys which are these little lizards that climb the walls you find them very often in the summer stuck to the wall often perhaps of a little house countryside and these lizards have a spectacular adhesion to the wall because it is carried by the paws because the paws under the palm have very very thin lamellae. Here these lamellae are very many and very very thin and the adhesion of these lamellae to the wall is due precisely to the Vaner Wals forces that are established between the lamellae of the geo's paws and the surface of the wall in most cases they are rather weak forces as we have just said but not on the micro and nanoscale there yes they become significant and the smaller the lamellae are the greater the adhesion to the wall will be and precisely to symbolize this I project to you again precisely what are the bond energies of the various bonds that we have seen up to now. This is the final lesson of the entire part regarding bonds both intra and intermolecular and we see that intramolecular bonds such as the covalent bond Ionic and metallic have forces therefore very high bond energies while instead when we go on the Vaner Wals forces and hydrogen bond we have much lower forces but obviously as we have already seen not negligible okay Having said that, I'll stop here with this lesson which is already quite long and in my opinion also quite complex. If you have any questions, answer them below in the comments and I always encourage my entire community. So all of you to answer each other, even someone who already knows can answer someone who asks the question obviously try to inform yourself and answer correctly this because I always try to supervise everything but first of all I can sometimes make some mistakes or forget something in the video but then above all I also have extreme difficulty responding to all the comments There are so many of you now and I am very happy about this but I have a bit of difficulty responding to everyone so I would like all of you my entire community to collaborate to answer each other and also be able to create something that is useful to all those who will then watch the video This is why I always say that even when you notice some error in the video try to go to the comments because or in the description Because I probably I could have already corrected it or some user could have already corrected it I maybe make a small note and bring that comment to everyone's view because I put it among the most important and then So you see it immediately in the comments so before commenting to correct Look if someone before you has already done it and to whom maybe I have already responded because at least this way you can trivially avoid another comment but certainly every comment is welcome and surely any comment that helps the community is always welcome. Having said that, I say goodbye and I 'll see you in the next video and I wish you good study.