all right this is lecture thirty point two and we're continuing to talk about atomic properties and the first atomic property we're going to take into consideration and talk about in any detail is atomic size and atomic radii are the same thing is atomic size this follows two different trends and one of them is pretty easy to understand and the other takes a little bit more information to really understand what's going on so first and foremost atomic radii increases going down a group so that's going from top to bottom in the periodic table ie the arrow pointing down and the radii decrease going across a period and that's the one that's going to take a little more of information for you to really understand what's going on with that so atomic radius is measured as half the distance between adjacent nuclei in a molecule so we're basically going from the de nucleus to the nucleus so the radii is the distance between them as shown in the in the image here alright and so we're not going to be measuring any atomic radii per se we're going to just assume that we will be given measurements if we're added to ask to determine those radii but we want to talk about why atomic radii change as we are looking at different elements in the periodic table so to start with let's go with the easy one first the trend going down a group is that it increases and the reason it's increasing is well we're adding shells essentially if you think of the Bohr model and you're going from N equals 1 to N equals 2 N equals 3 and so forth you know that these shells are higher and higher energy and getting further and further away from the nucleus and so basically as a principal quantum number changes or the principal energy level increases in electrons move further and further nucleus obviously this is going to increase the radius of the element and so basically more and more electrons bigger and bigger atoms as we go down the periodic table so atomic radii increases so that's easy enough to understand and here's it just an example and as you can see as we are going down the N number is increasing the shells you know in two three four and so forth and as you see the radii increase as we go down the list and that makes sense we've got more electrons we've got a larger atom and so that's not anything you know real difficult to understand now the thing that's a little bit trickier to understand and is why atomic radii decrease right so we're talking about decreasing as we go across the period in a so these are the rows so we're going across a row in a periodic table and as we go from the alkali metals all the way over to the halogens the atomic radii decrease and that's a little bit harder to to to understand but it's it's it's it makes sense let's just say that so what we're going to be looking at is both looking at the atomic number and of course in a neutral atom atomic number tells you not only the proton count but also the electron count and where are we actually adding the additional electrons and that's where the the crooks of this comes in so as we move across a periodic table we notice that the atomic numbers increase right that's how its organized in each each element is one proton larger than the one before it and this means that the positive nuclear charge is increasing so you're getting more and more positive charge as you go across the period now we know that the periodic table is organized such that you know N equals one is the top row n equals two is the second in equals three four and so forth so this means that each row in the periodic table is filling a shell that means that the electrons that we're placing into these shells whereas when we're going down the periodic table they're going electrons are being placed in it successively larger in values going across the periodic table they're all going into the same end value so they're all going sort of into the same general area around the nucleus so the electrons are not going into higher levels they're going into the lower or the same levels now let me explain this this nomenclature or so to speak here as you know in equals one only has two electrons in it any you know in equals two has two and then six so 8 possible electrons and so what these numbers are here 2 & 8 well this would be the N equals 1 and N equals 2 sub you know the shells so this is telling us that these are full and the electrons now this is the N equals 3 the electrons there's one going in there and then 2 going in there and then 3 going in there and so as we go across the periodic table the core electrons are remaining exactly the same and we're only placing electrons in the outermost shell and and so what's happening as we go across the periodic table the the number of protons is going up right it's increasing but when we're putting in electrons so we're putting electrons in these shells we're not going out to put them in we're putting them all in at the same level so what's happening is the pull in other words the nuclear pull is increasing whereas all of the electrons that are around here are going in to the same level now the electrons that are in here these are what we call shielding electrons and so this is the the amount of shielding right here in the inner inner circle these are the shielding electrons and what they do is they shield the outermost electrons from the nuclear charge well as we go across a period in the periodic table the number of shielding electrons as you see this is not changing it's it's basically staying exactly the same so they're not getting any more shielding but the number of protons is going up so the the charge the effective nuclear charge is what we call it is increasing steadily going to the right so think of it like a magnet we're getting a stronger and stronger magnet the shielding between what the magnet is pulling on is is staying exactly the same so as we go across the periodic table the shielding becomes weaker and weaker because the effective nuclear charge is becoming stronger and stronger and that means things at the outermost that count on that shielding are being pulled closer and closer in as the magnetic charge so to speak increases and so as the atomic number increases across the row the additional electrons are all added just to that that same or even lower as you see here where we're showing this these are where we're actually filling in electrons in even lower fields this is remember they don't all feel sort of sequentially they they jump around when we get to the higher levels and so you're basically going to have an effective nuclear charge that increases and that means that these electrons are being pulled more strongly towards the nucleus and so effective nuclear charge this is one of the few equations that you have to know it's really more of a relationship but the effective nuclear charge which is symbolized as Z effective is just the nuclear charge this is the number of protons okay and you're from that you're going to subtract this is the core electrons okay so these are the shielding electrons and these would be the electrons from the previous in value shells so in this case this is of course N equals 1 and N equals 2 shells completely full right that's what these shielding of electrons are so if we were calculating this we would have 11 so you know we'd have 11 as our number of protons minus 10 so our effective nuclear charge would be plus 1 and as you go across of course the effective nuclear charge increases a lot so here's the sodium we were just looking at this is the two eight with the one electron in N equals 3 shell that's its electron configuration and so as it goes across right you know neon this would be neon three S one would be the electron configuration of sodium and so when we look at this you notice that you know there's just the one electron after the noble gas the noble gas configuration is the core electrons so again the effective nuclear charge is plus one as we go across that period when we get to the the next element here this is your sulfur you'll notice that we've added more electrons to this shell in this electron configuration but the effective nuclear charge is only the core electrons and so the core 10 electrons have not changed and that means that the effective nuclear charge is plus 6 so no doubt a plus 6 charge is going to pull these outer most electrons these outer six electrons closer in because the core shielding electrons have not changed and so this is what causes the atomic radii as you go across periods in the periodic table to decrease and here's here's basically sort of a schematic of these trends and as you can see as you go across these things do shrink so to speak they increase going down and that makes sense because you're getting more and more shells and more and more electrons but going across because of this effective nuclear charge you see this reduction in the size of the elements radii okay so these are the two trends and you need to understand the trends and what causes those trends okay and you should probably be able to calculate an effective nuclear charge for for an element or elements all right now we've done the the neutral atom so what happens when we turn these atoms into ions and so we want to go through how the ionic radii will vary so positive ions we know that positive ions occur when they lose their outermost electrons they always pull from the outermost shell and we talked about that and you make your electron configuration you have to pull from the outermost shell aluminum's electron configuration is essentially represented here again this is in you know so this would be one s 2 and then 2s2 2p6 and then we have 3 s 2 3 P 1 basically okay and so this is the electron configuration for aluminum if I want to make a anion out of aluminum it likes to lose three electrons right because then it can be like the nearest noble gas so this would be neons electron configuration with just the 2 and the 8 fill and so these are you know how positive ions form and this is just a review so you should be able to do this and know what electrons in an electron configuration need to be lost to make a stable iron out of an element so we can go to get element aluminum 3 plus and so when we look at this we have positive ions and they're going to essentially remove the outermost shell electrons so you're you're basically losing a shell and if you lose a shell obviously it's going to shrink so positive ions get smaller right so they get smaller alright and their trend you know has this you know no more positive charge and and so you have a stronger elect effective nuclear charge because think about it you basically now have 13 protons for aluminum but you only have 10 electrons and so this is going to pull even those core electrons closer to the nucleus so positive ions are very much smaller okay they're they're smaller than their element and so here's here's a table and even the ionic electron configuration is the same as the Iranian it's progressively smaller as you go across so sodium is +1 notice that it has an atomic radius of 190 because they're fairly large on the positive side of things but the ionic radius is is much smaller and your ionic configuration is the same but again why is this happening so if our ionic okay so this is the electron configuration right for each of these elements they're attempting to match their configuration to the nearest noble gas which is neon and so they have the same exact same number of electrons but they have different numbers of protons and so what's happening is the proton is the pull towards the nucleus if you have the same number of electrons and you have a stronger pull as you go it's going to get smaller and smaller and that's exactly what it has you can tell what's happening and so the effective nuclear charge also means positive ions get smaller and smaller as you go to the right in the periodic table okay and so here are some comparisons and you look at this look at the change that takes place in terms of the ionic nucleus you know this is really a very very dramatic change in the in the radii because of that charge and that effective nuclear charge now you'll notice in these you'll see a trend starting over here in these negative ions and you can tell when you start adding electrons what happens of course you get a larger ionic radius and so if you look at the nonmetals here we got chlorine we're gonna go from 2 a 7 2 to 8 8 we've added electrons and of course that means the effective nuclear charge is decreasing okay so we've gone to a you know a full shell so to speak and so we are basically just making a bigger atom so if you add more electrons you make a bigger atom all right and so again you know if you go down the principal energy level everything gets bigger right we're going down the you know in the periodic table if you're going down size increases and that's because you're adding shells but going across periods in year you're looking at two different trends with respect to ion charge and size and think about it this way so when you're talking about positive ions right where do we switch from positive to negative ions well we do that when we go from metals to nonmetals right so as long as you're in the metals you're going to see the size decrease steadily as we go across the periodic table so if we go from left to right you're gonna see that the size decreases steadily and that's because the charges are increasing right they're becoming more positively charged and we just showed why that would happen you know the effective nuclear charge is getting greater and greater as we go from left to right but then as soon as we get to the negative ions then you're going to be seeing a difference right we're going to be seeing that the size goes up when we start going into the negative ions and so here's sort of a I guess a summary you can see in this this sort of table here so let's go across let's try this row so we see that potassium plus is 138 then calcium 2 plus this 99 Allium is 62 and again we still in positive ions this is t9 which is a drop from here but it's not as much of a job you see it starts and this is in the kind of the metalloid section around here you'll see it starting not to drop as much and then as we go up what you're going to see is that it's going to jump up but then it's going to start going down again and the reason for that is we add three electrons here you know we had two here and add one here so you're gonna go kind of to a low and then you're gonna have a so it's going to go down and then it's going to jump up and then it's gonna start going down again okay and the reason it's going down is because you only added one electron here you add it to here and you added three there so these are the trends that you see in the periodic table but the important part is you know why those trends exist you know it's if you need to go back over this or you need to listen to somebody else describe it you definitely want to read the book and get your mind around why these trends exist knowing the trends you know it goes up it goes down it's increasing decreasing that's great knowing why the trends exist is the most important thing alright so that's the end of this one