hi in this video we're looking at trends in the periodic table and a trend is just a pattern so if we look in the periodic table from top to bottom within a group we might see a pattern for a certain property similarly if we look left to right within a period in the periodic table we might also see a trend in a certain property and so we're gonna look at three properties the first property is something called atomic radius the second is ionization energy and the third is electronegativity so let's start with just what is atomic radius atomic radius is the size of the atom it's actually physically the radius of the atom itself the distance from the nucleus to the furthest out electron in the electron cloud and so if I just look at Group one here from hydrogen to cesium and I look at how many electron shells are occupied by electrons in hydrogen I really only see that there's one electron shell because there's only one electron so we only need one shell to house that one electron if I look at lithium though lithium has three total electrons that means two of those electrons can fit in the first shell but the third will have to go into the second shell so lithium has two shells and as you kind of go from top to bottom in this group we see that we just have an increasing number of electron shells because the electron amount is going up and we need more room to house those electrons the way that we house those electrons in an increasing amount is by just bringing in more electron shells to hold them and what you can see is that as we go top to bottom there is a bigger radius because we have more and more electron shells the way to remember this just quickly is to think of a snowman if you look at these bottom three here you see a snowman that was always my trick for remembering this and I'm sure it's plenty of other people's too it can be yours but the reason that that's happening is because there are more and more electron shells as we go from one period to the next down a group let me lay out the atomic radius for all of the elements on the periodic table what we see is that top to bottom trend of the atomic radius increasing but what you'll also see is that as we go from left to right within a period the atomic radius actually decreases now may seem strange to you because if we're adding electrons as we go top to bottom and that's causing our radius to get bigger shouldn't our radius get bigger as we go from left to right in a period because in that situation we're also adding electrons the explanation for this one is a little more complicated so let's just take a look at the elements in period two lithium to neon lithium has three protons and three electrons as I slide over from lithium to neon I'm adding one more proton but I'm also adding one more electron and as that happens the attraction of the electrons to the nucleus becomes greater and greater and that causes the entire electron cloud to just kind of scoot you in closer to the nucleus so you can see of this grouping here neon is actually the smallest atom and lithium lithium lithium is the biggest atom in period two so now we have our full trend here top to bottom in a in a group the atomic radius is going to increase think of the snowman effect there the radius actually decreases as we go from left to right within a period so the overall trend if you think of it this way this may help you helium is the smallest atom and francium or a caesium that would be on this slide here but francium is the largest atom and so knowing this trend you can actually get a lot of the information for the other two properties we're looking at ionization energy as the next property we'll investigate ionization energy has a fancy name but it actually is not all that fancy when you know what it is ionization energy is the amount of energy needed to remove an electron from an atom you can think of this almost like it's the cost of taking an electron from an atom a low ionization energy means it's really easy to take an electron from that atom a high ionization energy means it's more difficult now here's the overall idea the smaller and atom is the harder it is to remove an electron from it if I just look at an extreme case where I have a small atom the valence electrons are really close to the nucleus in this scenario but if I have a larger atom the valence electrons are much farther away from the nucleus and therefore they're much more willing to kind of be donated to somebody else another effect that's going on here is the fact that we have all of these rings that contain electrons already and so this is called shielding shielding is where the interior electrons kind of mitigate the attractive force of the nucleus the outside valence electrons so said differently the signal is kind of being blocked by the middle electrons and so the valence electrons are not as loyal to the atom and so it's much easier to steal from a bigger atom as compared to a smaller one now if we know the trend in atomic radius in the periodic table and we know that the bigger the atom is the easier it is to steal from it and therefore the lower lianas ation and energy is required then we can figure out what the trend in ionization energy would be just based on knowing the size of the atoms across the periodic table the ionization energy is going to increase from left to right because if you just compare the size of lithium for example to the size of neon this is much more difficult so higher energy over here because the size is smaller those electrons are being held really tightly to that nucleus lithium not as much as neon so that means lower energy over this way and remember the lower the energy requirement the easier it is to take an electron from that particular element if we know that from left to right across a period and it's really just based off of the size then we can also figure out what the pattern is from top to bottom and a group for ionization energy it's at the ionization energy energy decreases so again the bigger the atom the easier it is to steal from it and that means the lower the ionization energy will be okay so the last piece is going to be electronegativity electronegativity is going to come back when we talk about bonds and bond types so this is something that is certainly important to know it's going to help you when we when we talk about bonds but in an electronegativity what we're talking about is how much an atom wants an electron so it's a measure of an atoms action two electrons and it's based on a scale from zero to four where four is the element that really really wants electrons and zero is an element that really really does not want electrons so let's look back at the periodic table we fill in all the atomic radii for each of the elements here I'd actually just like to look at kind of the two extreme scenarios here let's let's compare lithium to fluorine fluorine has seven valence electrons if fluorine can just gain one more valence electron it'd be a lot like neon and that would fill its valence shell it would satisfy the octet rule and fluorine would absolutely love to gain an electron there four fluorines electronegativity is going to be really high in fact fluorine has the highest electronegativity of all the elements on the periodic table now contrast that against lithium they're in the same period but lithium is on the other side of the periodic table the left side lithium has one valence electron and does not want to gain any more valence electrons in fact lithium would actually prefer to lose that one valence electron to become a lot like helium and so the electronegativity for lithium and for all the elements on sort of the left side of the periodic table is going to be far lowers let's take a look at it just a special periodic table that only shows the electronegativity values if I'm yellow I really don't want electrons in this periodic table and if I'm red that means I really do again let's look at fluorine three point nine eight compared to lithium's point nine eight there it is and so if I just kind of look from left to right within a period I'm gonna see that the electronegativity increases if I go top to bottom in a group the electronegativity actually decreases and this again unsurprisingly has a lot to do with the radius of the atom if I just kind of look kind of down this way here all of these halogens would like to gain one electron to be like their noble gas neighbors it is much easier for fluorine to attract an electron that isn't originally its own because fluorine is so small its nucleus is located very close to its valence shell as you go from top to bottom the nucleus gets further and further tucked in the center the valence electrons get kind of further and further out so the electronegativity actually is a little bit lower as we go top to bottom not because they don't want that one electron but because they're pull on neighboring electrons is a little lower because they're just bigger atoms and so that's what's going on top to bottom in a group on the periodic table for electronegativity so that's it those are the three trends in the periodic table it's important that you know what the trends are for sure but I believe it is way more important that you know why these trends exist and if you think about it long enough it really all boils down to the size of the atoms in many cases it's about the distance between the nucleus and the valence electrons thank you [Music]