all right now this video is going to be covering the general forces of attraction that are really observed throughout some of CC and kist unit one I actually forgot to record the audio for our first recording of the class so I'm just going to be going through the PowerPoint and offering explanations where I would think somebody may or may not ask a question all right so let me just begin now so when we talk about forces of attraction it generally deals with any type of interaction in a way between two molecules or any types of interactions within a molecule that keeps it together now those are generally what we describe as forces of you know attraction but generally here we tend to have two different categories that we talk about um when we describe these interactions and they are int molecular and Inter molecular forces of attraction now intra generally refers to within a molecule so forces that keep a molecule together while intermolecular forces are generally these forces between molecules that kind of determines their physical properties um in a sense right if you have a lot of um inter molec forces between two different molecules right you can increase its melting point um boiling point Etc all right so those type of stuff we're going to be looking at now intermolecular forces tend to be much weaker than intramolecular forces right so as it says here inter is between two molecules and intra is with within the a molecule itself so we see some examples here such as London dispersion dipole dipole ion dipole hydrogen bonding ion ion calent bonds so we're going to be talking about intermolicular forces actually as they are the weakest forms and then go into intr molecular forces now intermolecular forces generally talk about now the vaval forces London London dispersion forces are really they're not exactly the same but synonymous with vandol forces now this really occurs when two electronic species right two species that bear electrons right um have a difference in position of electrons and have a higher charge density on one side for example we have let's view that as a proton and here is the electron cloud if we have for example a helium ad example right we have two electrons all right in the 1s orbital here so what if these two electrons actually move to one side of the electron cloud this will make this side of the electron cloud slightly negative and this slide which is electron deficient slightly positive right this creates a pole so this molecule or this atom itself this species is said to be dipolar or polar in nature right as it has a more negative side and has a more positive side now if these interactions can happen right or this phenomenon can happen in a atom itself right or in a molecule itself right we can see that these dispersion forces right as electrons disperse right we can actually see that they can interact right with all um species because in this diagram here if if we have a species with an instantaneous dipole right it's arising from the movement and the dispersion of electrons right if we have this right we're going to have let's say more electrons on this side you're going to have the dipole here that's a dipole Arrow this side is more positive right if we have this more positive side facing another species that is just non-polar electrons properly um diffused everything is fine let me just put in the electron take an example right this positive charge on this molecule will attract the electrons closer opposite charges attract right so they're going to be moving to this side of the molecule and as you can see here in this example now right we now have a pool being established all right so this is how we induce a dipole we start out with a adom or species that has an instantaneous dipole due to some London dispersion forces generally um we're going to have an interaction of the LA and dispersion forces between the two um substances right so we have an instantaneous dipole in one and it tends to induce a dipole on another all right based off of different of charges now this begs the question our understanding of polarizability all right now polarizability we're going to be disc discussing it in more depth in module 3 when we reach module 3 but polarizability generally refers to the moleculees ability to become polar right or to displace electrons right while yeah that is general term so when something is polarizable it means it can easily really disperse electrons it can really become polar really easily electrons can move to one side of the molecule or be concentrated to one end of the species and we have the creation of poles right if something has good polarizing power it easily polarizes something right another species but we're going to be really speaking about that and how these forces um help with polarizability in module 3 when we talk about periodic trends dipole dipole forces this generally speaks to two substances that have dipoles already that allow some type of electrostatic interaction so for example here we're going to have this for an example I find it a little bit weird but it's an example nonetheless so we have these two species all right within the molecule itself we have these two dipoles right and how would we know that a dipolar rizes how do we know that a substance is polar right we really have to look at the fact that one species has to be more electr negative than the other all right so we generally look at stuff like that sorry so in this case we have iodine right in this example specifically right and we have chlorine all right so we're looking at this example specifically so we're not looking at actual values or anything like that all right but what electr negativity really speaks to is the tendency to attract electrons right so when something is more electr negative it tends to drag electrons closer towards it so in this case the chlorine atom here is more Electro negative so it will actually drag electrons in the bond if we look at it in the case of just looking at electrons between the two molecules right if chlorine is going to be more electronegative we're going to see that the electrons are closer to the chlorine than to the iodine so we're going to have more electron density surrounded around the chlorine and if there's more electron density we're going to have the chlorine being slightly slightly negative right and the iodine being slightly positive right because the chlorine attracts electrons from the iodine to become slightly negative and since the electrons and the bonds and everything is closer towards the chlorine then I will actually become a little bit electron deficient we won't really feel the negative charge that much and become slightly positive all right and we didn't know slightly negative slightly positive by Sigma as we can see here so in this molecule we're going to have slightly positive slightly negative charge negative attracts positive there and positive attracts negative here all right so the negative chlorine will attract the positive iodine and the negative chlorine here would attract this positive iodine here forming the interactions now bonding right and interactions inside molecules intra molecular reactions are generally just displayed by unbroken lines while intermolecular interaction is displayed by these broken lines so anywhere you see that you know that it's an intermolecular reaction um not reaction but interaction right so we have the negative species here Electro negative species here attracting an electropositive species here so that's how that would work permanent dipole dipole forces are really established through you know actual um permanent species right so it's really a kind of example of a type of dipole right other than the instantaneous dipole induced dipole species as we saw from this slide right one is nonpolar right and one it has a dipole so it causes the nonpolar one to become polar in nature but with permanent dipoles as we saw from this example here right both are have dipoles both are permanently dipolar right for example here we have hydrogen right which tends to be really electropositive right and Florine which is the most Electro negative element on the periodic table right so between these two now we can clearly see that there's a dipole that we can draw all right where the electrons are closer to Florine because of its electro negativity all right when we have molecules that have permanent dipoles right they will also interact in this way as the slightly negative side of the molecule will interact with the slightly positive side of the molecule induc dipole we looked at that already from the helium example but just putting it in this case again right we have chlorine as an example here so it means that one tends to be um there there's a possibility that one of the substances tends to be you know polar in nature and the other tends to be nonpolar right and creat a pole on the non-polar species all right so that generally tends to happen but in this case here right both tend to be non-polar right but due to the interaction of another we're going to have a pole a pole being created all right so chlorine chlorine bonds by itself is fine right but let's say that another chlorine molecule becomes closer right there is going to be some repulsive forces between the molecule right and it's up to really chance but the electron density on one molecule will actually move away right and make a pole all right and then induced dipoles can arise from that so it's a chance that one chlorine molecule may push electron density of another chlorine atom on another chlorine molecule onto another right so it's a tendency that this atom may push the electrons on this closer to this chlorine and establish this dipole here all right in that case now we're going to have these London dispersion forces right so remember that when we're talking about these things there usually tend to be different form of London dispersion forces right dipole dipole permanent dipole dipole or induced dipole dipole right interactions hydrogen bonding now is another form of kind of permanent dipole dipole interactions or just dipole dipole interactions itself right but these are specific these are a specific type right when we look at hydrogen bonding we're looking at an interaction between hydrogen on one molecule to a extremely electr negative atom on another molecule all right that must be bonded to hydrogen so the conditions established is that when we look at things like halides which are group 17 and chales which are um group 16 these tend to have high electro negativity electro negativity increases across the period from left to right so we tend to have hydrogen on one end of the spectrum and we tend to have chides like oxygen and sulfur and halides like Florine chlorine bromine iodine right these have high Electro negativities really really high all right hydrogen's electro negativity is generally like 2.20 but you can look up the other electr negativities across the periodic table and see how some of them are 3 something for maybe for Florine not really remembering it right now but they definitely have higher electri negativities than hydrogen right and this creates a pole right as we were discussing before once a species as high electro negativity is going to attract the electrons towards it and make a pole on the molecule right as you can see here hydrogen bonding in water so water tends to have a high m a high boiling point because cause of hydrogen bonds and stuff like that right so even though it's called a hydrogen bond it is an inter molecular interaction and not an intra molecular interaction right it's an intermolecular reaction caused by intramolecular um characteristics right so in this case hydrogen is bonded to un chalal or a group 16 element Like Oxygen and water right so oxygen is way more Electro negative electronegative than hydrogen so I'm drawing these ARS to see the direction of electron flow within the molecule and now you'll notice that electrons are more well I need to put in my lone pairs because oxygen does have lone pairs so I want you to see the sheer difference in number of electrons right so oxygens are are negative right slightly negative and if you drag electrons from one species to another the one that has less electron density will now become slightly positive all right so just drawing all the symbols here for you to see it so in a water molecule we're going to have a dipole being created dipole here slightly positive and here slightly negative all right and in in hydrogen bonds right we tend to have interactions between hydrogen specifically and the other electr negative element right so this type of interaction is the hydrogen bond right so hydrogen B Bonds in in themselves are weak right but why do we tend to describe hydrogen bonds of having you know such strength right when we look at hydrogen bonding within substances it leads to a higher heat of latent heat of vaporization latent heat of fusion um looking at higher melting well higher boiling points actually all right because these Jil in liquids in a sense all right usually but that's because of the scaling of it one or two hydrogen bonds may be weak intermolecular forces but hydrogen bonds really stack so if you look at this this water molecule example here we only have two hydrogen bonds on this example right but as you tend to increase the number of water molecules let's say we have one entire mole of water molecules we're going to have a lot of hydrogen bonds and the amount of hydrogen bonding increases the number of forces of attraction between the molecules and make it much harder make you it makes you require or make it so that you require way much more energy right in order to break these bonds right so that's really important there okay so that's how it really um relates to higher boiling points Etc right it also scales if the molecule is capable of having more hydrogen bonds per molecule right so looking at something like ammonia ammonia has three three hydrogen atoms right and one nitrogen atom so it has actually four areas in which it can participate in hydrogen bonding right so if you look at a water molecule let's look at this specifically it has hydrogen yet it has two lone pairs all right of electrons so you have four areas where you can have hydrogen bonding all right same thing with nitrogen you can have hydrogen bonding to the nitrogen there to one of the hydrogens and two other hydrogens right can be attracted to nitrogens so looking at the fact that one molecule can make four hydrogen bonds so if you scale it to like an entire liter right you're going to have a lot of interactions right so that's how it really scales so that's another this is another example of a molecule in which hydrogen bonding can take place as nitrogen is more electronegative than hydrogen all right so that's generally it for our general overview of the intermolecular forces but for our intr molecular forces now we have to inspect the forces within now for ionic bonding right we're going to be looking at this in the sense of well in the sense that it's same stereotypical type of bonding that we see in CC now in ionic bonding we look at the fact that it is between a metal and a non-metal right right a non-metal tends to accept an electron right to be to follow the octet rule to become um to have the same configuration of a noble gas noble gas configuration is important right while the metal tends to give off electrons to reach noble gas configuration right so in this case ionic bonding is really the bonding between a metal and a non-metal right so for example in this case we have a sod atom all right so we're looking at the attractive forces between these two right as the sodium gives one of its electron to the chlorine right and we technically denote these substances like this in a molecule right it's it is a neutral salt or sometimes we can denote them like this to show the charges between them the attractive forces a positive cation and a negative anion now if we denote it like this this shows the true um binary nature of this type of salt right but ionic compounds tend to form these type of lazes all right so the attractive forces of ions are isotropic this means that they attract in all directions so if we have a sodium atom sodium ion sorry let me draw a little bit closer right it can attract a chlorine chloride atom chloride ion sorry about that in any direction right it's like if you have a magnet sitting somewhere right so if you have let me just draw a small magnet n s a small magnet sitting somewhere the magnet can attract metal from generally any Direction all right if you're looking at like that it can attract from any direction so it's the same thing with ions all right they don't attract on One Direction only all right so in order to maintain the stability of these ionic compounds the positive and um calion tends to be surrounded with negative annion to St to stabilize it right and the negative annion tends to be surrounded by positive cations to make sure it's stable as well so it forms this intricate repeated ltis structure this threedimensional lce structure as we can see here I'm going to be pointing out this sodium atom right here sodium ion rather the sodium ion right here right is surrounded by 1 2 3 four chloride ions right and if I single out this chloride ion here it's surounded by one two 3 four sodium ions so it's a stabilized structure in which all the charges right um tend to really cancel each other out to create a neutral salt all right and they're tightly bounded in this repeated Arrangement this threedimensional latis structure all right so that is the structure of ionic compounds now the properties of ionic compounds now ionic compounds are crystalline in a sense right now these type of compounds are synonymous with the interactions of cations and anion right and this make this tends to make ionic Crystal lates really really strong right CU we know that the interactions between different charges are strong right so they tend to have high melting points right so ionic compounds tend to have high melting points it's hard to really break the ionic bond they require a large amount of energy to Break um the bonds in the crystal latice structures right so that's important ionic compounds tend to be bural as well even though it takes a lot of energy to really make the molten to really melt to them to break these bonds right it is really easy to break them apart so So within crystals and stuff like that you can actually tend to break them into powder salt if you have large grains of salt you can Crush salt right table salt sodium chloride now this is because right even though within the structured lattice there are strong attractions between the ions right if you implement a large enough amount of mechanical force we disrupt the structure of the lattice that you can see here we disrupt disrupt the structure if you disrupt the structure you're going to ca like charge ions to interact with each other and repel each other so what ends up happening is one piece of the lattice ends up falling off or repelling itself so that's really how we have the brutal structure of ionic compounds another thing is our conductivity really really important ionic compounds can't really conduct electricity right can't allow electricity or electrons to pass through them in it ionic structure right it's l structure that is because everything is really fixed it doesn't allow electrons to pass through but let's say that you melt it you increase the temperature you make molten sodium chloride as an example right or you solve it you put it into a solution right you put it into a solvent right that allows for the ions to be freed up right that allows for the ions to be freed up and allows charges to move move throughout the solution right so positive and negative ions are fixed within this the L structure so they can't conduct anything but within solution positive and negative ions can move and allow electrons to pass through the substance all right so that's really important so that's generally for the ionic bonding all right for coent bonding really a equalent bond is really formed through the equal sharing of electrons this is generally what we understand through um CC right but we're going to look at the fact that it doesn't have to be equal sharing right we can have a substance donating both of its electrons to contribute to one Bond while the other is electron deficient and we're going to really be speaking about that real quickly all right so before we get into that we have to understand Atomic orbitals and just a recap of atomic orbitals Atomic orbitals are probable electron locations right or described more simply as places in which we can find electrons right so we know that shells don't really exist but we use them for explanations in chemistry but generally we see that electrons are clouds of density that surround the proton itself right so for example in this case we have some examples of you know some orbitals here we can I'm just going to see in the chat here right see if you can name what type of orbital this is I just put in the square right but these other orbitals here all right see if you guys can name all of them all right let me just drop some axes onto them right so one two three all right so let's see if we can name these they're kind of different all right let me just see if I can just make that a little bit like that one 2 3 all right so let me call this Zed X Zed y and x z YX Zed YX well so see if you can name them drop the names of these orbitals down in the description so this is a b c and d in the comment section just let me know what the name of these orbitals are all right now hybridization let's do a quick recap of hybridization now hybridization really speaks about the redistribution of energy or electrons right from orbitals of lower energies or different energies right to give orbitals of equivalent energies so so that we form hybrid species or hybrid orbitals that's really what it is so it speaks about orbital mixing we the example of carbon we see that carbon tends to have two unpaired electrons right so how can it really create four bonds we learned already that carbon can create four bonds but how so we have to speak about the fact that ground state carbon tends to have two electrons that are free right but once we excite the species right we have hybridization occurring so what tends to happen is that when we excite the species right we're going to have one of the electrons right from the 2s orbital here being promoted to a higher energy level right here in the 2p Z in an nty 2p orbital right and that allows for them to have similar energies denoted by the changing color here so now we not only have two electrons that were free for bonding but now we have four electrons that are free for bonding all right now if we have four orbitals right four electrons free for bonding there we're going to have what we call sp3 hybridization if we have three orbitals free for bonding we call it SP2 hybridization and if we have two freefor bonding we call it SP hybridization right how does this bonding occur right we're going to be really looking at it but SP hybridization looks a bit like this right so in Sp hybridization we have one s orbital and one p orbital merging to create one two hybrid orbitals and they're hybrid because they're not quite as orbitals and they're not quite P orbitals they're mixture between the two right so they tend to look like this all right so it's SP so it has a little bit of s characters a little bit of P character so we call it an SP hybrid orbital now in Sp hybridization we have two or hybrid orbitals in SP2 hybridization we see that 1 s and 2 p orbitals one and 2 p orbitals create three so note that the number of orbitals is always higher than the number so we're going to have three orbitals for spe2 hybridization and the information here on side is cool so if you want to just pause and watch it and well then read it then it's fine all right it's really important information and for sp3 hybridization we're going to have one s orbital but three p orbitals and that gives us four right SP hybridized orbital another way you can look at it is that s is always one as orbital and P3 speaks to three different people orbitals right and together they make four hybrid orbitals all right now the single calent Bond now now that we're done with that the single calent bond is really caused by the overlapping of these hybrid orbitals that's why we must have a species hybridizing before it can Bond a species cannot really Bond like this covalently if it's not hybridized so we have the hybridization and then we have the overlapping of SP orbitals along the inter nuclear axis right so along the same axis as the atoms right as the nuclei right to create a sigma Bond so these two must overlap in this way to create what we call a sigma Bond all right so that's a sigma Bond characteristics generally there are strong bonds generally right um the electron density in a sigma bond is concent along the nucle internuclear axis as I said before it allows for free rotation we're going to describe what that means in unit 2 actually and sigma bonds can still ex can exist in single double and triple bonds which we will briefly look at and sigma bonds exhibit cylindrical symmetry along the bond axis which we'll have a look at in unit two so we have different types of fix Sigma bonds if two s orbitals one and two s orbitals overlap we call it an SS overlap for example within the hydrogen atoms all right if we have an S and A P overlapping we have SP overlapping for example within the hydrogen carbon bond in methane all right and if we have two P orbitals we call it PP overlapping for example between two carbon atoms in Ethan ethine right or ethine so ethine is example on the slide right so we have this here these different types of overlapping Pi bonding now it tends to occur between two unhybridized orbitals so let's say that we have okay we're going to be using this example right so we have these this PCI here all right so it's going to have its single Bond right denoted by the overlap of hybrid orbitals right so this is the S the sigma Bond the single Bond here right the same thing as that right so we're going to have a s Sigma bonds here here here and here right but notice now that there is a free unhybridized p orbital on both of the carbons what will happen is that there's going to be an interaction between two of these orbitals they're going to overlap and allow electron systems to you know occur right so why Within this type of bonding here it allows electrons to travel throughout that's why we have the sharing or bonding happening right we're going to have electrons traveling right throughout the systems here and it will create what we call a pi Bond denoted by the symbol pi right so a pi bond is formed when two unhybridized P orbitals overlap all right side to side right so they tend to be parallel right so as you can see here from this diagram they tend to be parallel they run parallel right and they are perpendicular this is a perpendicular line to the internuclear axis so we have the internuclear axis here and it's perpendicular to it all right so looking at that so they tend to overlap side to side like that right so we have the characteristics of P bonds right the limit rotation we're going to talk about that when we talk about um isomer and stuff like that um within geometrical isomer in unit 2 right they have high electron density right so high electron density is concentrated around them below and above the nuclear inter nuclear axis because we have one electron density here and one electron density here right due to the overlap of both of them right so they're generally weaker than Sigma bonds due to their side to side overlap so they're not head on they're not head-to-head overlap they do side to side so they tend to be a little bit weaker in p Bond electron density is distributed over a larger area and that's one of the reasons why it's kind of weaker and Pi bonds are commonly found within double bonds and triple bonds so let's have a look at this I actually forgot to edit this let me just edit it real quick I'm going to go for an example of an electron description of double bonds right an electronic well not really configuration but description I want to see a graphical description of the double and triple bond real quick there was a good example that I had on the slide within class h okay so this is it all right so let me just add it to the slide here now within this example now I want to show you guys the look of it right now within this type of structure right we're going to have a sigma Bond here we see the head head overlap all right and then we have unhybridized P orbitals for example these red orbitals here would then overlap to get a pi Bond right so when you have one Sigma Bond and two and one Pi Bond actually one Pi Bond you're going to end up with a double bond Right double bond but if you have one Sigma Bond and two Pi bonds we're going to end up with a triple Bond and that's because there are three different areas of electron density around the atom right well generally all right three different bonding areas right so we have three different types of overlaps right one being a sigma overlap one overlap from one Pi Bond and the other overlap from the other Pi Bond giving three bonds all right so now that we're over with that right that's the general bonding within covalent compounds right we tend to have coordinate bonding right or dative bonding right so this happens when one electron Rich species donates both of its electrons to form a bond with an electron um deficient species so an important understanding of this is get is given by ammonia ammonia and Boron trifluoride all right so so in this case Boron Tri fluide right has all of its electrons used up within its General species it's actually fine this way it defies the octet rule in which it's actually satisfied without all it on its villain shell but it can participate in coordinate bonding right so this is what we would call electron deficient because it does not have eight electrons on its valence shell yet it is fine all right but with nitrogen here on the ammonia molecule right we have a lone pair of electrons we say that this species can be somewhat electron R right it's not it doesn't have an expanded octet which we will look at in veence cell theory Jen electron per repulsion Theory but in this case now nitrogen gives both of its electrons to Boron so yes they are sharing the electrons but it's not from two individual species in this case it's solely from nitrogen so we tend to have a bond like this forming and I see this on a lot of paper ones right the arrow is pointing towards the Boron to show that this is coordinate bonding to show that hey the nitrogen is the one donating both of the electrons to this Bond all right so that's what's happening there now for the bond length and bond energy I put this information here a precursor to understanding what we're going to go through in energetics but it's important to just pause and go through the information here all right we also put a list of lengths and bond energies now for metallic bonding now we already generally understand what metallic bonding is right so we have the concept of an a metallic lattice structure right of made of cation surrounded by a sea of mobile mobile delocalized electrons right so the example that um the syllabus wants us to get is copper right so I just use this picture because the color is synonymous with copper all right so because of this structure we have different different um classifications and characteristics coming out and this the first one is conductivity we know that metals can conduct electricity electricity speaks to the movement of electrons all right now we have to know specifically that even though the cations are within a rigid structure right a rigid lattice within Metals the electrons are delocalized meaning that no one cation right has an electron meaning that electrons move around the cations and no one cation is bounded to electron or no one electron is bounded to a cation so electrons are easily allowed to flow throughout the metallic lattice right and the flow of electrons gives us our electricity now if we apply a potential difference or a voltage right or set up a circuit right and allow current to run through this allow the metallic structure allows for the electrons right that are delocalized free moving allowing them to move around catons to complete a circuit this allows the metals to be conductive right conductivity is the movement of charge malleability and ductility now malleability speaks the fact that metals can be molded they can be shaped right they can be turned into shapes of trophies of of nails of hammers of chairs of cars they can be molded ductility speaks the fact that they can be dragged and into sheets and wires and stuff like that now metals can do this even though they're in a Lati format we understand that ions if you go back to ions ions form laes laes as well but they brittle they are easily broken repulsive forces um between the ions when they meet up within like charge ions when they meet up and they just break apart but with Metals now the free C of electron is not really used to explain the ductility and malleability right but it's really the surrounding protons or surrounding um cations I must say in this case right the surrounding cations now when a metal is moved right so when a metal is generally moved or or changed or bent or dragged or molded right the cations don't build up in energy or charge and repel each other like in anionic well well not like in ionic structures I must say right it doesn't do that right what tends to happen here is that when the cations move generally we're going to have electrons that actually restabilize and keep the metal structure intact right so there's a new formation of cations that allow the metallic structure to stay intact so what we're saying is that we may have a structure like this right that cation specifically let me not make it like that so cations like this surrounded by electrons right if we move it now if we change the structure and put it something like this right these C are too close right they're really close in the ionic slates what they'll do is repel and just break apart but in metals we have a stabilizing C of electrons right so it keeps a metal the metal structure intact so the cations may be rearranged but the Sea of electrons creat a new formation that keeps the metal intact right so this layer of ions in an the layer of ions generally in an electron C moves along one space with respect to the layer below it so the crystal structure does not fracture but it is deformed right so it can change shape it can go through deformation but it is not brutal enough to just break right so that's really the explanation there the last thing on this entire well the last two things on this entire powerpo is luster and melting points now we generally understand melting points and boiling points of metals right because of the strong metallic structure right but luster now this comes from our understanding of how light interact interacts with um electronic species right so if we look at the fact that electronic not meaning you know like computerized but substances that contain electrons right because there are lot of of um electrons free moving electrons there are free SE of electrons within Metals right upon interaction with a photon we have photo excitation of electrons now these electrons will move to higher energy levels right let's display using this photo excitation these electrons will move to higher energy levels really excited right and then when they drop back down the energy that they absorb to move to higher energ levels are redistributed as a specific wavelength of light right this light being reflected by the metal itself is what gives it this this shiny demeanor right this shiny look all right so that explains the luster of metals most metals have a lot of electrons and tend to reflect way more light right so they tend to be more shinier than other things like non-metals and ionic crystals and stuff like that so they tend to be way more shiny because of the reflection of all all of this lights all right melting points and boiling points strong metallic bonds are the result of more delocalized electrons so the more delocalized electrons there is the more there's effective charge distribution in the ltis of cations right so we tend to have really really really strong um connections between the electrons and cations and this makes um the metallic strong the metallic bond really strong so in effect making the size of the C Small all right the size of charge of the cation smaller it reduces the high positive charge and creates a more of a neutral tone within the lattice right so the more electrons there are to really um reduce the high charge density of cations the stronger metallic bond and the more more energy required to really break metallic bonds all right so that's why Metals T have high melting points or high boiling points for example metals like gallium that melt on body temperature or Mercury that is liquid at room temperature they may be liquids but they still do have high boiling points so it's something to really note all right so that is generally our description of the fundamental inter molecular and intra molecular forces all right so that's generally it so if you guys have watched until now just please um let me know right what are the interesting things that you have gained from this video or just let me know if there are any questions or concerns you'd like to share just drop them in the comment section and I'll be able to really address them thanks for watching