Periodic Trends Cont and IMFs

Added: 2026-05-09

[00:00:04]all right guys so we're gonna pick up um yesterday or excuse me on Friday um we introduced the periodic trend of electronegativity um and talked about how that relates to polarity we are going to start in and look at the other periodic trends that we need to be aware of um first one that we're going to look at is atomic radius or size um this is very similar to radius of a circle that you have done in math class um but in an atom there is not a nice defined outer edge of that electron cloud so how they Define atomic radius is it's one half the distance between the nuclei I have two atoms of the same element when they're joined so roughly halfway between the nuclei is going to be where the edge of that atom is but like I said it does not have a nice defined Edge the actual Trend you need to be aware of um is atomic size is going to increase from top to bottom so it's going to get bigger as you go down agreed and then it's going to decrease from left to right across a period so the biggest elements or biggest atoms are going to be in the bottom left-hand corner of the periodic table smallest ones are going to be in the top right hand corner of the periodic table and kind of why this increases and decreases you're going down a group the number of energy levels increases so you're going from if we were in the First Column you're going from putting electrons in 1s to 2s to 3s as we go down to each new row we're adding a new energy level and the electrons in that outermost energy level are further away from the nucleus and they're shielded by electrons closer to the nucleus from the pool of the protons so I'm trying to think of a good way to explain this um if someone were in the front of the classroom and they threw a tennis ball the people that are most protective from being hit by the tennis ball are the people that sit at the back because they have everyone in front of them kind of blocking or shielding them from that tennis ball being thrown at the front of the room um same kind of thing occurs whenever they have electrons in those inner energy levels they're blocking the pull of the nucleus from those outer electrons um so since those outer electrons are blocked from the pull of the nucleus they can spread out and the size of that outermost energy level increases whereas going across a group the number of energy levels doesn't change so if we look at like the second row we would add electrons in that 2s energy level and then into p those are these same energy levels so they're the same distance away from the nucleus um so there's no increase in shielding there's no additional inner levels of electrons to block the pull of the nucleus but going across the atomic number increases the number of protons increases which is going to be able to pull on that layer of electrons more kind of causing it to shrink in towards the nucleus which is going to cause the size to decrease bone across a period with um radius we also have to look at the radius of ions um so remember ions are charged particles that are produced when an atom gains or loses electrons I'm sorry cations are positively charged ions these have lost electrons to become more positively charged these are typically going to be our metals and what happens is since an electron is lost there are more protons than electrons the protons are able to pull on the electrons more tightly pulling them closer to the nucleus making the ion smaller than its original atom whereas an anion are negatively charged ions they gain electrons to have that negative charge these are typically going to be our non-metals and since electrons are gained there are more electrons than protons the protons cannot hold on to the electrons as tightly which in turn makes the ion larger than its original atom a way to remember these is I kind of use the analogy of weight so think of if you lose weight you get smaller cations lose electrons and get smaller if you gain weight you get bigger anions gain electrons and get bigger so that's one way that you can kind of remember the trend with those we also are going to look at ionization energy and this is the energy required to remove an electron from an atom so for it to completely lose that electron and it be transferred to another atom um we're going to typically deal with the first ionization energy which is the energy needed to remove one electron even if it's something that is going to lose more than one electron it only loses one at a time so the trend for that first ionization energy is that it decreases from top to bottom within a group so it gets smaller going down and it increases from left to right across a period this is the same trend from election activity that we did on Friday so the largest ionization energy is going to be in that top right hand corner of the periodic table the smallest is going to be the bottom left hand corner um if it is something that loses more than one electron there is a value for the second ionization energy so the energy in the interior made that second electron there is also a third ionization energy if it has to lose a third electron and generally as more electrons are being removed it requires more energy so if it's something that is going to lose three electrons the second ionization energy takes more energy to remove that second electron than the first one did that third ionization energy is going to take more energy than the second or the first so as they remove more electrons it gets harder and harder to remove them requiring more energy so the actual trend for ionization energy and I'm going to try to move this down a little bit that way we can still have our Trend there so going down a group we said that ionization energy decreases the atoms are larger going down a group making those electrons further away from the nucleus and since those outer electrons are shielded from the nucleus it's easier to remove those electrons so the outermost electrons when they're further away are easier to pick off um so this is why the energy into your movement electron decreases down a group so you could think if we're in a horror movie kind of the rule is you never like go anywhere by yourselves if we're in a big group and there's a serial killer the one that's the people that are easiest to pick off are the ones on the very outside of the group um whereas the ones kind of in the center of the group are going to be harder to pick off um and then kind of going across a period we said that atoms decrease in size going across the period so remember there was no shielding because we're in the same energy level however going across from left to right there are more protons that were pulling harder on the electrons in that energy level so since it's smaller and those electrons are closer to the nucleus it is harder to remove those electrons so that's why the energy needed to remove electrons increases going across the period and like I said this trend was the same as the trend for electronegativity however remember electronegativity did not include noble gases ionization energy includes noble gases so that's kind of a difference between the two so there is a way to remember the trends um and we call this the Roboto dance so you're using your arms um to kind of make these wedge shapes so I'm going to draw it in that top right hand corner I already have the right arm drawn and let me switch colors and I'll draw the left arm representation so the overall trend is with your left arm straight up and down if you look from the width of your pinky down to the width of your elbow your arm increases in size going down so that represents that going down a group ionic and atomic radii increases so your left arm is representing the radius trend and then you're going to turn it horizontal across your chest and notice from left to right it goes from your elbow to the tip of your pinky which is decreasing in size from left to right so where your elbow is is that largest radius and then your right arm is all the other trends this one you're going to kind of replace your left arm so it's going to start going across your chest horizontal and from left to right it goes from your pinky to your elbow so it increases in size going across and then you're gonna drop it down where your fingertips are pointing towards the floor and you have your elbow on top going down your pinky so it decreases in size going down so the largest ionization energy and electronegativity is that top right corner this is just kind of a way to help you remember the trends if you're not so good at memorizing those initially all right and we're going to switch gears and we are going to look at intermolecular forces these are forces between two different molecules that we're going to see um so molecules can attract one another in a variety of ways now these intermolecular forces are intermolecular attractions are weaker than actual bonds so these are weaker than ionic or covalent bonds um these a lot of times determine the state of matter that we're in so it'll determine if a molecular compound is a gas liquid or solid at a given temperature and it all depends on the strength of those horses and there are three different forces that we have to be aware of two of them are considered Van Der waals forces and then the third one are hydrogen bonds so we're going to look at these individually so our van der waals forces um these are split up into two different forces we have our dipole interactions and dispersion forces these are the two weakest attractions that there are um these were developed by a Dutch chemist um Johannes van der waals and the dipole interactions remember we talked about with differences in electronegativity we end up with that polar molecule that has the positive end and the negative end of a dipole so dipole interactions are where polar molecules are attracted to one another so the positive end of one molecule is going to be attracted to the negative end of another molecule so these occur between oppositely charged poles so if we have something like HF chlorine has a lot of electrons it's the most electronegative so that end is going to be slightly negative whereas the hydrogen is going to be slightly positive how it interacts with the hydrogen fluoride next to it is that positive and negative end are going to be attract attracted to one another so that dotted line is our dipole interaction just where those two oppositely charged poles are attracted to one another now the dispersion forces are London dispersion forces these we are going to see in everything but this is typically what we're going to look at in non-polar molecules and this force is caused by the motion of electrons because electrons are constantly moving and what happens is they will unequally distribute at some point in time while they're moving around which causes this dispersion force so what happens is when one atom's electrons move it can influence the electrons of other atoms around it to move in the same direction and this is due to electrons being attracted to the nucleus of another atom so I'm going to try to draw this this one like I said is the trickiest one to kind of wrap your head around so here I'm going to use the nucleus As an X so here we have two kind of nicely balanced nuclei where all the electrons are evenly distributed in that electron cloud well what happens is one of those the electrons are moving around and they all decide to kind of cluster on that right side what that does is the atom next to it all of its electrons that were nice and evenly distributed they don't want to be close to the electrons of that other atom so they will all migrate away from those electrons that are there now this will then redistribute because the electrons are constantly moving around and it will go back to normal so this is just where the electrons since they're randomly moving around at any given point in time they're not evenly distributed for a split second which causes a dispersion force between these two molecules where it's kind of repelling the electrons from one atom and making the electrons and the other atom move around but like I said this is going to be a split second and then it's going to go back to normal so this is something that is constantly changing a way to think about this is kind of in our school building we generally know you know if it is seven o'clock in the morning if every student in the building is considered an electron most of our electrons are going to be kind of near the commons and cafeteria area or even out in the parking lot at that point in time well 715 Bell Rings everyone should be in classes so there's going to be students on every single hallway in the building well then class time or class change occurs there's still electrons in every room of the building whenever third period though we have a lunch we might have a whole hallway that's empty and then a cluster of people in the cafeteria so the electrons are the individuals in the building are not evenly distributed anymore but once a lunch is over those people go back to their classroom and then another group of people go to lunch then we have a different empty hallway so this is something that is going to be constantly changing throughout the day or throughout time all right and our next course that we look at are hydrogen bond and the name of these is kind of misleading because this is still an attractive force it is not an actual Bond like an ionic or covalent bond would be so a hydrogen bond is an attractive force in which a hydrogen is covalently bonded to a very electronegatively charged atom also weekly bonds are attracted to an unshared pair of electrons on another electronegative atom um and this you're going to see those very electronegatively charged atoms to remember which ones work I say hydrogen is fun f-o-n so hydrogen bonds are going to occur whenever hydrogen is attached to either a fluorine an oxygen or a nitrogen and that's because there's such a large difference in electronegativities that it causes this strong attraction this is what happens in water so I'll actually draw water so you can see what this looks like but you actually saw this in biology um whatever you were looking at like DNA what holds the DNA bases together are hydrogen bonds and also with like proteins being folded is do those hydrogen bonds and hydrogen bonds are the strongest of these intermolecular forces so remember we said these are weaker than actual ionic or covalent bonds but hydrogen bonding is stronger than dipole or London dispersion so what this looks like is when we have water excuse me my pen turned off when we have water now a hydrogen bond is still just a very strong dipole so we have kind of a negative end of water and the positive end of water well here we have that hydrogen oxygen bond that forms so what happens is the water that is next door that hydrogen is going to be attracted to that unshared pair of electrons on the oxygen so this attraction is our hydrogen bond so it's just that hydrogen because it's positively charged because all the electrons kind of migrated towards the oxygen is attracted to those electrons on the oxygen next door all right and to do just a quick recap of those intermolecular forces I remember we said the intermolecular forces are weaker than actual bonds so ionic and covalent bonds are stronger than any of these three forces we said hydrogen bonding is stronger than dipole forces which is going to be stronger than dispersion forces dispersion forces are the weakest of all of them and it's because they are so short-lived like we said it's going to unequally distribute electrons for a split second and then go back to normal so this slide attraction comes from electrons not being evenly distributed around the nucleus and then it goes back to normal those dipole forces we said are stronger than dispersion forces because they're constant if something is polar it is always polar so it's always going to have an attraction to other polar molecules so this attraction is from the slight different selection activity making part of the molecule slightly positive and the other part slightly negative so remember a positive and negative in from two different molecules can attract one another and interact and then hydrogen bonding is the strongest of the three and this is also constant because it is like I said a strong dipole is due to a larger difference in electronegativity between the hydrogen and its oxygen nitrogen fluorine it is attached to this comes from them being so far apart on the periodic table and by the heightened attraction due to small size of atoms you're not really going to have to identify these in molecules at this level that's more of what happens at the AP level but you should know which one is stronger than the other and you should be able to identify them based off of their definition only