is about a covalent bond formation so um now we're gonna switch uh ionic compounds electrons are transferred and covalent bonds electrons are shared okay um and then one of the the common uh kind of uh like abbreviations that i use to remind myself what is covalent is looking at where the atoms are on the periodic table and typically atoms that are involved in covalent bonding are either going to be non-metal plus a non-metal or maybe a metalloid plus a non-metal um interestingly though uh it that's just a that's just a generalization okay you really have to do calculations um with ionization energy and electron affinity and you have to look at z effective and you have to look at the size of the atoms and then you have to look at a couple other things and then from that you can say where on this spectrum between ionic and molecular or ionic and bonding and covalent bonding do these atoms fall okay so it's a it's not just a generalization it's a pretty good one though so if i'm looking at the upper right hand periodic side of the periodic table or if i'm looking at hydrogen which happens to be in the s block right because it's one valence electron even though it's not a metal um i would know like oh i will form a covalent bond hydrogen can either form a covalent bond with itself or other atoms or it can form an ionic bond so you can do both like sodium hydride okay or hydrochloric sodium hydride it's a good example all right so let's take a look at this diagram you're going to see this diagram a lot and on the left hand side we have energy okay it's going from low energy to high energy um starting here if the atoms are too far apart then the attractions are weak and no bonding occurs now what are the attractions it's simply a negative electron is attracted to a positive proton on the other atom so the electron on this atom is attracted to the proton on another atom okay so and vice versa right so i have a proton on the atom on the left i have an electron on the atom on the right and they have to be attracted to each other okay so the electrons are feeling the protons in the nuclei they're also repelling each other so there is an optimal separation where the energy is at a minimum and um what's interesting is this bond length here is typically an average bond length and you have to really look at is this hydrogen h2 in the gas state is it in the crystalline state how is this bond length determined typically it's from nuclei to nuclei okay so i would say it's basically from nucleus to nucleus okay and often we give it the the letter r for radius okay um so if there was a radius um i don't know why we do that but we do it's the distance between the two uh nuclei um and then at this place at this point the electron and proton attraction on opposite atoms is optimum the electrons aren't repelling each other and neither are the protons so strong repulsions occur so you're going to get this shooting up of let me actually get this shooting up when you get these two atoms too close to each other because the protons and electrons are all repelling each other they do not want to be that close all right so it's very difficult to smash these two atoms together and the inner nuclear distance the distance between the nuclei is often again abbreviated as r so what's happening well atoms are going to bond because it results in a more stable electron configuration this is the theory from lewis and we see this energetically that is true atoms bond when they it have an energy minimum so basically nature is lazy it wants to sit it wants to minimize energy always minimize energy okay energy's expensive so minimize it so that's why atoms bond particularly in if you're looking at a covalent bond and um let's look at how atoms can bond equally unequally okay so atoms can bond oh i'm sorry i'm going to go through all the properties before i get into electronegativity yes okay so now the distance or that optimum bond length is going to be average for many similar bonds for many compounds so um here's what you guys need to know do not google the answers to my quiz questions um because you will find uh some like something on brainly or something on whatever and they're looking at a different averaging of bond length than i am explaining to you guys right now where maybe they're not in the gas state or maybe they're looking at just the crystalline state or maybe they're looking at you know an atom in a vacuum or a molecule in a vacuum okay so what i want you to know is um that the values that i have here from a textbook uh are going to vary depending on the circumstances but the general trends and this is what i want you to know is based on what's the size of the atom okay and how many electrons are shared between them so in general more electrons equal shorter covalent bond so more electrons shorter bond um you have to be comparing bonds between like atoms um so i can look at carbon and carbon or carbon and nitrogen and what we see is a trend that remember each line represents two electrons so a triple bond is going to be six electrons looks like triple equals double i'll use a colon there instead of equals so three lines is six electrons and we do see that they're shorter right so um another thing to look at is that bond length is going to decrease from left to right across the period and that's because um the size of atoms decreases from left to right across the periodic table and we talked about why that is you have an increase in effective nuclear charge and so carbon is larger than nitrogen is larger than oxygen or nitrogen is smaller than carbon oxygen is smaller than nitrogen so of course there will be smaller bonds now that's just one explanation there are other explanations as to why bond length is decreasing okay um other things that are going on is that you begin to have partial separate uh partial charges and then if you have charge separation that causes bonds to shrink so generally bond length is decreasing from left to right across the period and then finally it increases down the column because remember that as you go down the column you're increasing in n right so fluorine is in period two and then chlorine is three bromine is four so period two three and four so n is increasing and so these are getting just larger atoms right larger atom larger bond length between the two of them so that's another way to think of it um that summarizes that and then one of the takeaways from this that um i haven't really talked about actually is that as bonds get longer they're also weaker so what's happening is that you have less attractive forces between longer bonds so less attractive forces means you have a weaker bond okay so longer bonds means weaker bonds longer is weaker less attractive forces means longer bonds means weaker so when we look at bond energy we're going to see that the energy which is basically how much energy is is between these two bonds and how much it take to break it we look at that trend we'll see that um the weaker bonds are going to be the longer ones and the stronger bonds are going to be the shorter ones so again we're looking at a summary of what we looked at in the end of chapter three so bond energies more electrons to atom share the stronger the covalent bond so now i'm actually looking at instead of the the inner nuclear distance okay now i'm looking at the actual values and there's something that's very curious about this that we're going to see uh when you look at valence bond theory b b t valence bond theory is going to explain what and so will molecular orbital theory okay it's going to be explaining the fact that you have here 347 kilojoules for a single bond but only 611 for double and 837 for a triple hmm you just added two more electrons between these carbons why is this not 690 or something like that right why is this only 611 why didn't i double the um amount of energy required to break a bond or break the both bonds in in carbon and we see the same pattern for carbon and nitrogen something else is different not only is it not double it's like 305 versus 6 15. so what's going on there so then we see the difference in electronegativity matters okay so um we'll get to some of the reasons around this but uh right away you should be saying hmm more electrons shorter bond stronger bond and then if you look at uh bonding between bromine and the other halogens you'll find that again i'm comparing similar types of bonds so i'm comparing just the halogens here and bromine is the consistent atom so brf is a stronger bond than brcl is a stronger bond than brbr bonds are going to get weaker down the column atoms are larger longer bonds get stronger across the period atoms are smaller uh so good time to brush up on your periodic properties if you um don't have that memorized yet or understood yet and hopefully you have a periodic table where you've drawn all this stuff out like okay smaller atom right upper right helium's the smallest atom on the periodic table francium's the largest so go ahead and make sure you have that available and you should have a periodic table out too all the time we're talking about this