CH240 Chapter 1 Video 2 Bond Formation and Lewis Structures

Added: 2026-05-10

[00:00:01]chapter one bond formation and Lewis dot structures with Lewis dot Lewis understood that the atoms transfer or share electrons in such a way as to attain a filled shell of a full sale of electrons in other words the S orbital will always have total of two electrons when it's full the p orbital will always have six electrons and when you put them together you will get eight electrons and what that does and has allowed these atoms to become isoelectric which means they have the same electronic configuration as helium neon or argon so hydrogen is gonna have two electrons so it is technically than 1s - but it's gonna have a negative charge in this case except Rick has a gained electron is either transferred shares that electron allows that to happen so what they're trying to do except for hydrogen is fulfill the octet rule hydrogen is an exception the filled shell has eight valence electrons two from the S and six from the P for a total of eight and this will always hold true for the second row of the periodic table always always always so hydrogen helium doesn't have the eight but the second row does there are some exceptions of that which we'll get to here elements in the third row are higher can have what is known as the expanded octet rule which means there's the d-block working for them so solver we could have sulfur tetrafluoride and we'll eventually get into ocular structure so here is sulfur hexafluoride so sulfur had a total of let's meet leptons they're two four six eight ten electrons now around it instead of eight we can also have sulfur hexafluoride in this case it's gonna have twelve electrons that is the expanded octet rule there are times when aluminum may have three but it can because of this d-block except of a pair of electrons in others there's bonds here so that's still expanded and we will this happens in organic chemistry quite a bit with aluminum especially later on in the semester so with bonding two types of bonds and we have these in organic chemistry also and they're extremely easy to see when it happens we can have ionic bonds where there is a transfer of electrons so lithium has one electron chloride has seven based upon its placement on the periodic table so the electron from the lithium goes to the chloride and now we have lit if I write that a little bit straighter with a plus charge and chloride now has eight electrons around it and it has now a negative charge and then normally we would write them lithium with a plus charge and chloride with a minus charge for our ionic bond and we remember ionic bond is usually a metal bonded to a nonmetal and that will hold fairly true in organic chemistry also and these normally form nice crystal lattices covalent bonds is where they share of electrons so we have a hydrogen atom and a hydrogen atom forming h2 so s overlaps with a s and the two overlap here and this overlap is where they share the electrons and this is an example of a covalent bond and most covalent bonds are non metal with nonmetal and carbon is non metal hydrogen so we're gonna try to draw Lewis dot structures for this so with Lewis dot structures we're interested in the valence electrons only so we have methane which is ch4 it has four valence electrons so according to the Hans rule we have to singly occupy each space so I just pork four dots around it and for a hydrogen it has one and the electrons can move but I'm just this one is made simple intentional so we have put four hydrogen on here and it's pretty obvious where we're gonna stick the hydrogen we have four single electrons and there we have it sometimes we just draw each line represents two electrons and this is the Lewis dot structure for methane and there are eight carbon eight electrons around this carbon for the octet rule so we have two four six eight it's gonna be important because some of you'll try to draw ten some of you'll try to draw six or any number but to in between there and there done that so let's look at methane or excuse me methanol it is an alcohol so our carbons gonna have four valence electrons around it our hydrogen's gonna have one now oxygen has six one two three four but before around it now we can go back and pair them up so oxygen looks like it has two bonding sites so now we have to try to build this molecule and I'm trying to give it away there's gonna be three hydrogen now I can't put another hydrogen on there because then there's no more free electrons so I need to put the oxygen so I'll put it in blue now if we look at this carbon carbon right now has eight electrons around it we still have one more hydrogen to go so we can put it right here so what we have is and I'm gonna straighten it out it's not linear but I'm gonna draw it like this and this is methanol it has a pair of non bonding electrons sometimes we just call these a lone pair and you want to make a habit of always showing the lone pair because if you remember right electrons oh we eventually gonna figure out that reactions always occur around electrons so those lone pairs are important fact this would be called electron rich so if I have Corr chloro ethane I'm gonna put my two carbons together and we know carbon always has to have or bonds so now that carbon has for this carbon now has three so it's gonna bond with a chloride and this chloride is gonna have three sets of lone electrons so these are our lone pair so there are compounds that have multiple bonds so we can have a double bond or a triple bond in a double bond there's four sets of electrons and a triple bond there's six sets of electrons so an example of a double bond we're just going to go ch2 oh we have carbon and I'm just going to put my hydrogen here and then there's gonna be two bonds for my family to hide so I have a C double bond o with an H and an H coming off the carbons with two lone pair eventually we're gonna learn that oxygen is very slightly electronegative so it's gonna have a slight negative charge and that means this carbons gonna have a slight positive charge and be called polar that's for another video just to remember that whoops we have triple bond this is exciting it has a triple bond but if you notice the carbons in both of these cases still have a total of eight electrons around it 1 2 3 4 5 6 7 8 1 2 3 4 s 2 4 6 8 and this one also has 8 so it will always follow the octet rule there is a general rule to help when you're drawing Lewis dot structures and it's for patterns and these are for uncharged molecules because when they become charged like a hydronium ion eventually it's gonna check the role will change so carbon will always have four valence electrons nitrogen which means it's can always have a total electrons nitrogen will always have an electron pair one electron pair it has three bonding electrons and one electron pair oxygen will always have two electron pairs and two bonds hydrogen will always have just one bond the halides by chlorine bromine fluorine and iodine they will always have one bond and six pairs of electrons sulfur will always have two electron pairs and two bonds well I wonder why because they're in the same family on the periodic table why are the halides the same because they're in the same family in the periodic table so that will always hold true for uncharged species we will do more Lewis dot structures