bonding and electronegativity so remember we have two types of compounds that we've talked about and again compound being made up of more than one type of atom they can combine as molecules or they could combine as ionic compounds as depicted here we have our ionic compound on the left and remember it's the formula unit that makes up what we see for this compound here um it is a ratio the lowest whole number ratio is how we express it so again if it is sodium chloride there is one sodium ion and one chloride ion do not forget that these are charged though we do not include that in the actual formula it's an ionic compound it's made of ions and it is not that there is one sodium and one chloride ion there are one gazillion sodium ions for every one gazillion chloride ions is a one to one ratio so that's what we have for ionic compounds with molecules though they do form distinct units so for something like h2o you do actually have two hydrogens and one oxygen and that makes up one molecule of water and it goes around acting like water racks but with sodium chloride you don't have one sodium on one chloride you have one gazillion sodium ions and one gazillion chloride ions is the ratio the formula represents the formula unit luckily for molecules the formula actually is how the particles themselves end up looking now what we're going to be looking at today is how do you determine which of these two um the the compound that you're dealing with is going to behave like and so to do that we're going to be talking about this concept of electronegativity so electronegativity is the ability of an atom to attract electrons specifically in a bond not just an atom sitting by itself because that is electron affinity and that's not particularly useful for us but what is really useful is electronegativity so if you take that atom let's say x and you want to bind it to atom zed they are going to have different electronegativities in that one of them may want electrons more so they might end up forming a bond between them but one of those atoms may hog the electrons more than the other one so they can end up pulling on the electrons and we can determine this by calculating out what the electronegativity difference is between these two who wants the electrons more these values are on the chart here we'll talk about trends at the end um but this is essentially these numbers here are the electronegativity values for the atoms that we have here um so our periodic table um again if you want you can think of hydrogen as being over there but if you notice you can see a pattern here that we should be able to explain by the end of this but each atom has its own electronegativity value and that essentially represents it if these atoms were to bond or when they're in a bond this is how much they want or they're going to pull on or hog the electrons within that bond and it's the difference between their electronegativity and the thing they're bonding to is electronegativity that will determine what sort of compounds you end up forming so the type of bond that you're going to get between atoms and their and and then that will determine the type of compound you have it is all determined by the difference in electronegativity so we use this delta sign meaning the difference in electronegativity so we're going to look at two atoms at a time and we're going to look at the difference in their electronegativities and that will tell us about the bond that's going to be between them we get this difference by taking the higher electronegativity value whichever one has the higher one and we subtract from it the lower one because we want to find the difference that's how we do it mathematically and that will give us the delta e n and then we'll see what that means in a minute but let's try this with a loose dot diagram for the molecules so these are two non-metals so i'm going to assume it's going to be a molecule and i'm going to be drawing the electro the lewis dot diagram for them and then calculating the electronegativity difference so we have our nitrogen in the middle and you look at which column it's in it's in column 15 so we're gonna have five valence electrons around it and we have our bromine in column 17 so it's going to have seven valence electrons around it and in order to make them stable they're going to be sharing their pair of electrons which we don't want to draw like that instead we want to draw with a bond a line representing the two shared pair of electrons it's gonna be the same we have the formula is br3 so we know that there's going to be three of these ones and it's gonna be the same thing for each one of these a bit of a shortcut again circle replace that with the line that represents the shared pair and we're going to end up with the same thing happening on this side as well shared pair of electrons which we represent with a line as a bond so there's a lewis dot diagram for nbr3 now we've done that before what we want to do this time is we want to calculate the delta e n for this compound so we look at the electronegativity of nitrogen and the electronegativity of bromine and realize that depending on which periodic table you're looking at they may vary slightly so if you're not getting exact same ones as mine that's fine use whichever ones you happen to have um it won't be a terribly large difference but the the periodical i'm looking at gives me the electronegativity for nitrogen at 3.0 and that of bromine at 2.8 so then obviously the difference between those would be 0.2 so what that means is that the electronegativity for this bond right here would be the difference would be 0.2 and it would also be i could do the math again if i want to but it's going to be the exact same thing this bond right here has a delta e n also of 0.2 and same math same answer delta en here equals 0.2 as well so this is going to tell me what the electronegativity difference is for each bond now you have to do each one separately you don't like multiply by three or anything like that you have to calculate it for a particular atom that is bonded to another particular atom and you do that by finding the difference in electronegativity of those two atoms and that will tell you what the difference electronegativity is for that bond so you're going to be getting answers somewhere between zero which again depending on which periodic table you use you might have got for the last question but it will start at zero if they have the exact same electronegativities then the difference between them has to be zero or it can be getting up into the three range so the delta e n i should say is going to go from zero to three the difference in electronegativity of any two atoms on the periodic table will fall somewhere in this range so you notice that if their difference electronegativity is low the amount that the bond is ionic is low so essentially what we're saying is that the higher the electronegativity differences the more those atoms are going to be like ions and that's saying basically the more charged they're going to be because if one wants the electrons more it's going to be pulling on those electrons more and the difference in electronegativity is going to be greater if they both want the electrons about the same amount then you'll end up with a low difference in electronegativity and the atoms themselves will not be much like ions at all if they have a high electronegativity difference if one doesn't really want the electrons and one really does there's going to be a high difference in electronegativity then the atoms themselves end up being quite ionic in nature they have a charge to them and therefore this is what we can use to figure out what type of compound we're going to have do the atoms within it share their electrons really well and therefore they don't have a big charge to them or do some of the atoms hog the electrons and that therefore would make them more like ions and give us an ionic compound realize that this is a pretty gradual change over the course of a difference in electronegativity there's a gradual change in the ionicness of the bonds but as humans we like to put things in in specific categories so what we do is okay well let's do a cut-off point somewhere in here where we can say okay if you're like getting over the 50 ionic-ness in your bonds we're going to give you one name and so if you're on this side of the line we're going to say you're ionic you're not showing your electrons well enough so we're going to say okay you're just your ions and if you're on this side of the line we're going to say okay well you are sharing your electrons quite well you've got a relatively small difference in electronegativity so on this side we're going to say okay these are molecules and so this is how we determine which type of compound is which we look at the nature of the bonds within them we calculate their difference in electronegativity if they have a large difference we treat them as if they're going to be ions if they have a small difference we treat them as if they're going to be molecules now realize this is not a easy not not a real categorization right there's going to be gray areas between the two so we have set the the line arbitrary line in the sand at 1.7 as the difference in electronegativity if it's over that we call it ionic if it's under that we call it molecular but if you're somewhere around there you might be kind of ionic and kind of molecular so there are some substances that are in the middle here and then we go leave this theoretical approach behind and we actually look at the evidence what do these molecules look like if their electronegativity difference is somewhere in the middle we start looking at their properties and say okay well is it acting like it's ionic or is it acting like it's molecular so this is just a simplification a way to identify molecules versus ionic compounds but it is always going to fall back on what does the evidence say what are the properties so here's that in a bit of a summary a little bit easier to look at if the electronegativity falls between 0 and 1.7 it is a molecule that we're dealing with molecule that says molecule if it's over 1.7 we're talking about an ionic compound and again this is over simplification this is to make our lives easier you can always look at the properties to say whether this is true conclusively or not realize that for the molecules if they are perfectly sharing their electrons the difference would be zero and so we refer to them as nonpolar covalent bonds so that the bond itself is exactly evenly shared so we can refer to that as a pure covalent bond or a nonpolar covalent bond if you're between 0 and 1.7 if you're somewhere in that range with your electronegativity difference they're still going to be classified as sharing but it's uneven sharing and when we say it's uneven sharing one's going to be differently charged than the other or what we'll talk about is partial charge than the other and therefore use the word polar to describe that if you're over 1.7 one of them so much charge than the other we actually treat them as if they are ions um and in most the cases they will be so take a minute pause here and try this out for yourself whatever periodic table have look up the electronegativities of the atoms in here realize some of these have more than one bond you have to calculate it for each bond if it's the same math you obviously will get the same answer so calculate the electronegativity difference for these bonds and then you can label the bonds are they above 1.7 or below 1.7 ionic or are they going to be covalent and if they're exactly equal then you know what the answer is going to be pretty quick and we can even call that pure covalent or non-polar covalent there's the answers so with hydrogen and chlorine and again your periodic table may differ a little bit but it should be somewhere around 0.96 is what i got um you may have less precision as well well with your electronegative difference but it should be somewhere around there and therefore this is between 0 and 1.7 so this is a covalent bond but it's not zero it's it's sort of on the in the middle area here and so it is polar covalent they're sharing but not perfectly equally dioxide again is going to be somewhere in the middle here so they are sharing they're covalent but not equally so it is called polar covalent bromine whatever the numbers doesn't matter because it minus itself or the difference between the same number is going to be zero so this is actually a pure covalent bond or another name for it would be a non-polar covalent bond water good one to know this is definitely going to be a polar covalent bond you'll notice this is pretty high this is getting close to the almost ionic part of it so this is definitely a polar covalent bond and that's a relatively big number on that polar covalent bond methane here a slight difference between the two so just over zero here so again they're sharing quite well not perfect but quite well so it's a it's a small polar covalent bond not a huge difference but a little bit of a difference and kcl 2.34 you're definitely over the 1.7 so we refer to this as an ionic compound and again if you're trying to visualize these right these first where's that five um these are all molecules right because they all are below the 1.7 kcl though however is ionic so it's ionic substance because it's over 1.7 and so if you're trying to imagine it it is not that there is one potassium and one chlorine that is completely wrong instead there is a ratio of potassium ions to chloride ions that is a one to one ratio in this giant repeating crystal lattice structure so again the electronegativity difference tell us what type of compound we're dealing with the greater the electronegativity difference the more unequal the sharing so once it's gone one past 1.7 we say okay it's so unequal we're not even going to call it sharing we'll say it's been transferred the electrons are actually moving from one to the other but realize again as we saw with that graph earlier it's a gradation between perfectly equal sharing and not equal share so with something like lithium chloride you have your chlorine is sort of hogging the electrons so this cloud here is sort of the electron around the the atoms themselves and so chlorine is so such a high electronegativity it's as if the electron has left lithium and is now around chlorine and so you see the larger cloud there and we did this when we did ionic compounds that would give chlorine a negative charge making it chloride and lithium a positive charge making lithium ion and we get our ionic compound so anything over 1.7 we will imagine that this is happening anything under 1.7 can either be perfectly equal sharing so again you have one atom here one atom here and they have their electrons around them and so if we think of your lewis dot diagram we did before you got a hydrogen with electron there hydrogen with electron there they're going to share their electrons and so we draw a bond um i mean in reality there's no line between the atoms right it is the shared area where the electrons are going to exist and if it is equal sharing we say it's a covalent bond or a pure covalent bond or sometimes called a non-polar covalent bond but perfectly equal sharing if like with hcl we have a chlorine and a hydrogen the difference between them isn't huge it's not over 1.7 but the chlorine still is hogging the electrons you can see that it is got the electron from hydrogen spending more time around the chlorine i won't draw an arrow actually because it hasn't actually taken the electron it is not ionic it is still covalent but since chlorine has a high electronegativity the difference is getting close to 1.7 it's getting up there it's quite a bit larger than zero it is a polar covalent bond that we have here it is a sharing of electrons but an unequal sharing that's what polar means is you've got these differences and we use this little symbol here this is a lower case delta and it is saying that chlorine here has a negative charge but it's not it's not an ionic charge like it has in lithium chloride this one here is it's not taking the electron away from the hydrogen it is just hogging the electron and so we put this lowercase delta sign to represent a partial charge so that sort of kind of looks like a curvy looking d or an s or some of that with a negative that means it is a partially charged it's not an ion yet it's a partially charged atom it is hogging the electrons so they are still if we're drawing our lewis structure they would still have a line between them representing the bond but they are partially charged hydrogen has its electrons being pulled away from it and so it is partially positively charged so see if you can try this out where you're actually drawing the lewis structure then showing the bonds and then writing in those partial charges that we saw in the last one so take a pause here and see if you can do it for carbon and fluorine all right should do the math first so we know what sort of bond we're drawing notice that this is less than 1.7 therefore it is not ionic it is covalent and it is not zero so it is actually polar covalent it's cutting on the high side of polar ground so quite polar covalent um so we have our polar covalent here we have these partial charges and whichever one has the larger number it's the one that's going to hog the large electronegativity number it's the one that's going to hog the electrons so here fluorine has the electronegativity somewhere around 4. it's going to be hogging the electrons away from the carbon so we go to draw this again carbon is in group number 14 has four valence electrons fluorine is in group number 17 it has seven valence electrons so we start off with that they're going to want to share a pair of electrons that'll make fluorine happy and of course we don't draw the circle in the dots we draw instead a single line representing that shared pair of electrons that makes the fluorine happy carbon is not however so as you'd imagine hopefully you can start seeing these in your head without having to draw every aspect of it um you're gonna need another flooring to share a pair there with carbon but again we replace that and draw a line so the same thing is gonna happen with another fluorine and another fluorine so as you'd imagine it takes four fluorines to bond with carbon and we end up with our formula cf4 so our carbon tetrafluoride this is our molecule and i know it's a bond i know i know it's a covalent bond i should be drawing these lines in here because the electronegativity difference is less than 1.7 so i draw lines now i can include the partial charges i can say okay well yes they're sharing this pair of electrons but fluorine is hogging it so fluorine is going to end up with a partially negative charge that's that lowercase delta that's not very good one and this carbon is going to have a partially positive charge to it now that is for every single bond each one of the bonds and again feel free to do this calculation four times if you like it's going to give you the same answer each time for the electronegativity difference of each of these bonds is 1.43 so each of these is a polar covalent bond and each of them have these partial charges so this fluorine would have a partially negative charge this fluorine would have a partial negative charge let's get this one there we go that's better um that fluorine has a partial negative charge as well and then the carbon would have a partially positive charge for each of those bonds i'm just going to draw it once because it applies to all four of those bonds so the trend for electronegativity as we saw with the when we actually saw the numbers of them um is that it increases up and to the right and so we see with fluorine the electronegativity is really high now be careful with the noble gases right remember the definition of electronegativity this is the ability to track electrons within a bond and those noble gases do not like to form bonds that's not entirely true you get some bonds of the larger ones but essentially most of your periodic table will probably give you no electronegativity value at least for the first three noble gases and you're probably not going to do any bonding with the other ones so you can ignore them for electronegativity essentially but other than those noble gases it increases up into the right making fluorine the most electronegative element on the periodic table and then again decreasing as you move down into the left is going to be less and less so again you should be able to explain why this trend is which is a whole lot easier than trying to memorize the trends if you understand that again it's all about the effect of nuclear charge it's all about the attraction of those electrons to the protons and for the same reason we see the trend with ionization energy and atomic radii it would also cause the exact same consequences for electronegativity if they have a stronger pull on electrons when you put them in a bond they're going to pull on their partner's electrons strongly as well so as you move from left to right across a period for the same reason you're getting more protons pulling on the same electron shell you get a higher effective nuclear charge the electronegativity will increase as you move up a group again you're getting fewer inner shells you're going to have more traction for those other electrons that are that are there the outer electrons for the nucleus and therefore you're going to have a stronger attraction as well and again higher effective nuclear charge that also will attract the electrons in other atoms as well so therefore increasing the electronegativity as well