in this video we're gonna look at trends for the periodic table of elements for dimensions like ionization energy atomic and ionic radii electron affinity and electronegativity and to do so we're going to start with a very fundamental idea in chemistry or physics and that's Coulomb's law and for our point of view we can view Coulomb's law as saying is that the magnitude of the force between two charged particles is going to be proportional that just means proportional right there is going to be proportional to the charge on the first particle times the charge on the second particle divided by the distance between those two particles squared when we're thinking about it in context of the periodic table of elements and various atoms you can view q1 as the effective positive charge from the protons in the nucleus of an atom you can view q2 as the charge of an electron now any given electron has is going to have the same negative charge but as we try to understand trends in the periodic table of elements it's really the outermost shell electrons the valence electrons that are most interesting because those are the ones that describe the reactivity and so when we think about the distance between the two charges were mainly going to be thinking about the distance between the nucleus and those outermost valence electrons now we can view this effective charge so I'll call it Z effective as being equal to the difference between the charge in the nucleus so you can just view this as the atomic number atomic number or the number of protons that a given element or an atom of that element has and the difference meaning that and what is often known as s or how much shielding there is now there is complicated models for that but for an introductory chemistry class this is often approximated by the number of core electrons remember we really want to think about what's going on with the valence electrons and so if you imagine a nucleus here let me do that orange color that has protons in it and so you have core electrons these are the core electrons in the first shell and then you have some core electrons in the second shell and let's say the valence electrons are in the third shell so let's say these are some valence electrons here and they're blurred around there in these orbitals those valence electrons which have a negative charge are going to be attracted to the positive charge in the nucleus but they're also going to be repulsed by all these core electrons that are in between them and so that's why an approximation of the effective charge that these valence electrons might experience is going to be the charge of the nucleus - and this is an approximation the number of core electrons that you have so if we use that roughly as a way to think about Z effective what do you think are going to be the trends in the periodic table of elements so I want to be the effective charge for the Group 1 elements over here well hydrogen has no core electrons and it has atomic number of 1 so 1 - 0 it's going to have an effective charge of roughly 1 lithium atomic number of 3 - two core electrons that are in 1s so you're once again you're going to have 3 minus 2 effective charge of 1 so roughly speaking all of these Group 1 elements have an effective charge of 1 what if you were to go to the halogens what are the what's the effective charge there well if you look at fluorine atomic number of 9 has two core electrons in the first shell so as an effective charge of seven chlorine actually has an effective charge of 7 for the same reason atomic number of 17 but 10 core electrons if you go even further to the right to the noble gases you see that helium is going to have an effective charge of 2 atomic number of 2 minus and 0 core electrons but then when you get to neon you have an atomic number of 10 and then - only two core electrons and you'll see as you go down these noble gases other than helium they have an effective charge of 8 and so the general trend is your effective charge is low at the left effective charge low for group 1 and then when you go to the right of the periodic table you have a Z effective Z effective is going to be hot so within a given period or within a given row on the periodic table of elements your outer electron surveillance electrons are in the same shell but the effective charge is increasing as you go from left to right so this q1 right over here is going to be increasing so what is that going to do to the radius of the atom well Coulomb's law will say that the magnitude of the attractive force between those opposite charges is going to be stronger and so even though you're adding electrons as you go from left to right within a row within a period the atoms in general are actually going to get smaller so let me write it this way so as you go from left to right generally speaking radius decreases now what's the trend within a column well one way to think about it is as you go down a column as you go down a group you're filling shells that are further out and so you would expect radius to increase as you go down a column or as you go down a group or you could say radius decreases as you go up a group so radius decreases so overall what's the trend in the periodic table of elements well radius is going to decrease as you go up and to the right and so you could draw an arrow something like this and it is indeed the case that by most measures helium is considered to be the smallest atom a neutral helium atom and francium is considered to be the largest atom so can we use this to think about other trends in the periodic table of elements what about for example ionization energy just as a reminder the first ionization energy is the minimum energy required to remove that first electron from a neutral version of that element and since it's the minimum energy it's going to be one of those outermost electrons it's going to be one of the valence electrons and so what's going to drive that well you can imagine the ionization energy is going to be high in cases where the Coulomb forces are high and what are the situations where the Coulomb forces are high well this is going to be a situation where you have a high effective charge and where you have a low radius lower radius makes the Coulomb forces high an effective charge makes the Coulomb forces I so where is that true so you have the lowest radiuses at the top right and you have the highest effective charge at the right so you would expect the highest ionization energies to occur in the top right so high high ionization ionization energy and that actually makes intuitive sense these noble gases are very stable they don't want to release an electron so it's gonna take a lot of energy to take one of those electrons away fluorine or chlorine they're so close to completing a shell the last thing they want to do is lose an electron so once again it takes a lot of energy to take that to take that first electron away on the other hand if you go to something like francium it has one valence electron and that valence electron is pretty far from the nucleus and there is a low effective charge despite all the protons because there's so much shielding from all those core electrons so it's not surprising that it doesn't take a ton of energy to remove that first electron from francium now another trend that we can think about which is in some ways the opposite is electron affinity ionization energy is talking about the energy it takes to remove an electron electron affinity thinks about how much energy is released if we add an electron to a neutral version of a given element so high electron affinity elements these are the ones that really want electrons so they should have a high Coulomb force between their nucleus and the outermost electrons and so that means they should have a high effective Z and that also means that they should have a low R so one way to think about it you're going to have a similar trend with the one difference that the noble gases don't like gaining or losing electrons but we do know that the fluorines are the chlorines of the world can become more stable if they gain an electron they can actually release energy so you actually have high electron affinities for the top right especially the halogens and you low electron affinities at the bottom left now there's one little quirk in chemistry conventions people will generally say that fluorine and chlorine and the things in the top right that aren't noble gases have a high electron affinity and it is the case that energy is released when you add an electron to a neutral version of them it just happens to be that the convention and this can get a little confusing is that when you release energy you have a negative electron affinity but generally speaking people are real when they say a high electron affinity this is going to release more energy when it's able to grab an electron now a notion that is related to electron affinity is electronegativity and the difference between the two can sometimes be a little bit confusing electronegativity is all about when an atom shares a pair of electrons with another atom how likely it is it to attract that pair to itself versus for the pair to be attracted away from it to the other one and so you could imagine it correlates very strongly with electron affinity things that release energy when they are able to be ionized to grab an electron if they form a bond and they're sharing a pair of electrons they are more likely to hog those electrons electron affinity is easier to measure you can actually see when this elements in a gaseous state if you add electron how much energy is released it's normally measured in kilojoules per mole of the atom in question while electronegativity isn't as clear-cut on how to measure it but it can be a useful concept in future videos as we think about different atoms sharing pairs of electrons and where do the electrons spend most of their time so I'll leave you there we started with Coulomb forces and we were able to intuit a whole bunch of trends just thinking about the Coulomb's law and the periodic table of elements