[Music] okay so this uh lecture covers so we start chapter nine and it covers this concept of ionic and covalent bonding so basically it examines how atoms are attached to each other or or attracted to each other to make compounds so basically the two main concepts it covers is ionic bonding and covalent bonding but there's other concepts covered in this chapter as well that allow us to look at the ionic and bonding and covalent bonding in more detail so electronegativity partial charges formal charges resonance core electrons valence electrons exceptions to the octet rule lewis dot structures and look at very general bond strength and bond length so all of these concepts in this chapter you should understand okay so we talked a little bit about core electrons and valency electrons and a little bit about ionic bonding in previous chapter uh so just to refresh your memory right quick so how are valence electrons and core electrons defined so if we take hydrogen lithium and sodium for example this is their electron configurations and in red um is the core electrons so that would be the noble gas core so two electrons is helium or for sodium uh two plus two plus six so that would be ten electrons so that would be neon core so that would be the noble gas core and and so we could write this in shorthand notation right like like this using neon and then 3s1 or helium 2s1 and so the electrons beyond the core then would be the valence electrons so hydrogen lithium sodium have each one valence electron right and so we can write the lewis dot structure using that valence electron so each dot represents the number of valence electrons the the atom has okay so and of course what does one so hydrogen lithium sodium they're all group one metals they all have one valence electron uh so how about oxygen sulfur selenium so if we did the same thing so an oxygen the noble gas core is helium two electrons then if you go down a period in the periodic table sulfur is below it so if you go down to the next noble gas so 10 electrons so that would be neon so here we have helium here neon and then selenium uh if we go down up here go down well let's just gonna do something different here so that would be argon right so that would be its noble gas core and so if we write this in sort 10 notation as i've done on the right so oxygen helium and then 6 electrons beyond the core so those 6 electrons would be the valence electrons and same for sulfur six electrons beyond the noble gas core so six valence electrons for selenium it's a little bit different uh because now we have d and d orbitals and if you go down another shell then f orbitals thrown in so now we have if you remember what's called a pseudonoble gas core uh where we take all of the shells that are full so the 1s the one first shell is full the second shell is full the third shell is 3s 3p and 3d and so that becomes what's called the pseudonoble gas core so if you were adding adding up the valence electrons for selenium then again it's just the s2 and the p4 so those six electrons uh would be the valence electrons so all of these are in group six um so that's the easiest way to determine valence electrons right just look at what group they're in and whatever group they're in that's the number of valence electrons so that works for the main group elements so six valence electrons so we could write a lewis dot structure right the atom with six dots around it to represent the six valence electrons okay so what is significant about valence electrons is that those are the electrons that are used to make bonds and those are the electrons that participate in chemical reactions so the core electrons for all practical purposes don't do anything besides occupy space obviously they'll influence the properties of the atom but they're not going to be participating in bonding in chemical reactions so it's the valence electrons that do that so those are the important electrons so for carbon what is carbon's valence electrons all right well that's the core helium and so it would have four valence electrons and of course the easiest way is just to look at carbon uh in the periodic table it's in group number four so it's got four valence electrons and for aluminum uh the core would be neon and so what's beyond the core then would be those three electrons so aluminum which would have three valency electrons and beryllium right its core is helium and so its valence electrons would be those two and beryllium is in group two so two valence electrons aluminum's in group three so three valence electrons carbons in group four four valence electrons okay so now let's look at the valence electrons in context of how they can make bonds so if you took a atom of sodium and an atom chlorine makes them well what charge does sodium like to be if you remember from way back in chapter two it likes to be plus one charged so it likes to get rid of one electron which electron would be that one right so the core electrons nothing is happening to it's the valence electron right it has one valence electron and those are that's those that's the electron that matters so sodium loves to get rid of that electron it has a very low ionization energy it can get rid of that electron and give it to something else uh what about chlorine what does chlorine like to do it likes to be negative one charged so it likes to gain one electron right where would that electron go it would go there into the 3p orbital because it's not full right and so again for chlorine um that's yeah that's chlorine it's core and so it's got seven valence electrons so three s and three p so seven valence electrons but it likes to gain more one more electron into the three p orbital so if you mix sodium and chlorine that's what's going to happen sodium is going to give up its 3s electron and that electron is going to go into chlorine's 3p electron and then sodium would become 1s2 2s22p6 which is in a plus and what else what else has this ice this electron configuration that would also equal neon 10 electrons so na plus and neon have the same number of electrons so they are said to be if you remember the word they are isoelectronic with each other and then if chlorine gains one electron then it becomes that 1s2 2s2 2p6 3s2 3p6 and so that would be cl minus what also has this electron configuration would be argon so cl minus and argon are isoelectronic so if you mix sodium and you mix chlorine that's all that happens is sodium gives up an electron to chlorine but it does so very violently right it's a very favorable reaction so it gives off a lot of energy okay so let's illustrate this reaction uh using lewis dot structures so sodium has one valence electron chlorine has seven valence electrons so if you mix the two together basically what happens is sodium gives up that electron and chlorine picks it up and now you have n a plus with minus that's lost one electron and co has now eight valence electrons so it's now co minus so it picks up one electron so it becomes negatively charged and so this is an example of an ionic compound so on an ionic compound you always have a cation and you always have an anion every single ionic compound has a cation and an anion and so this would be an example what's called an ionic bond so sodium and chlorine are not physically attached to each other right there's nothing linking them together but they are attracted to each other right why are they attracted to each other because something is positive and something is negative and opposite charges attract so that's what an ionic bond is is cation and anion that are attracted to each other but not attached to each other okay so let's do the same analysis so what if we had magnesium and fluorine so this is magnesium's electron configuration right those are the core electrons and so this would be two valence electrons right so what charges magnesium like to be if you remember magnesium likes to be plus two charged so how does it become plus two because it likes to get rid of those three acid electrons so it wants to get rid of two electrons what does fluorine like to do fluorine likes to be negative one right it likes to gain one electron to finish the 2p off and to have 10 electrons like neon so magnesium wants to get rid of two electrons but fluorine only wants to gain one electron so how do you make that work well you need two fluorines all right so one electron from magnesium can go here and one electron from magnesium can go to the other fluorine atom and so you would make mgf2 right so if you mix mg and f you need two fluorines to make mgf2 would be the compound that it makes so magnesium is now going to be 1s2 2s2 p6 so that would be mg2 plus which is isoelectronic with neon and then fluorine is going to become 1s2 2s2 2p6 also f minus which is also isoelectronic right with neon but again you need two fluorines for every one magnesium so if we wanted to illustrate this using lewis dot structure so we have magnesium with two valence electrons and we need two fluorine atoms that have seven valence electrons so one of these electrons goes to that fluorine one of these goes to that foreign and you would make an mg two plus cation and you would make two f minus anions right and they would not be physically attached to each other but they would be attracted to each other so even though it's wrote like mgf2 you have to recognize it's really an ionic compound and it really looks like that instead mg2 plus cation and two f minus anions not physically linked to each other but they are attracted to each other okay so how about calcium and nitrogen well what charge does calcium like to be it likes it's a group two metal so it likes to be plus two right and nitrogen if you remember likes to be three negative so the formula for the compound it would make would be ca3n2 calcium nitride so in order for this reaction to work so calcium right that's its core electrons argon and so it only has two valence electrons right um so calcium has right two valence electrons um but you're going to need three of these because you need six electrons because nitrogen each nitrogen wants to gain two three electrons so we're going to need three of these and then we're going to end up taking these six electrons right and so we're going to need then two nitrogen atoms uh because what is nitrogen wants to do it wants to gain three electrons there to become 2p6 and then it looks like neon so that would be n3 minus which is nitride which is isoelectronic with neon and then calcium of course is going to lose two electrons each each calcium [Music] will and so then that would be isoelectronic with argon okay so these six electrons three goes there and three electrons go there and you make two you make two n3 minus anions and three ca2 plus cations so we would need so nitrogen has five valence electrons right because this is its core and these are valence of course nitrogen is in group five so that's another way again to know that it has how many valence electrons it has so we need two nitrogen atoms and we need three calcium atoms and then what has to happen is one two three electrons has to go to this nitrogen and then one two three electrons has to go to the other nitrogen so you make three ca2 plus and you make two in three minuses or you make so that would be ca ca3n2 but you would have to recognize it's really three calcium two plus cations and two n3 minus anions okay so what about carbon and fluorine so what does carbon want to do well carbon's in group four right so that's core so it's got four valence electrons so to look like a noble gas if you remember carbon would like to be either c4 plus or c4 minus so either likes to gain four electrons here so then it would be c4 minus which is isoelectronic with neon or it likes to lose those four valence electrons to become 1s2 and then it would be c4 plus which is isoelectronic with helium and fluorine of course just wants to gain one electron here to become f minus which f minus which is isoelectronic with neon so the formula for the compound that it's going to make so you need uh four fluorines because each one can only take one electron but carbon has four electrons to either gain or give away so it's going to make a molecule cf4 right but the problem is is that carbon does not really want to give up four electrons because when you give one away electron away the next one becomes even harder to give away right because now you have so carbon's got six protons hanging on to six electrons it gives one electron away now it has six protons hanging onto five electrons if it gives another one away now it has six protons hanging onto four electrons so every time you take a electron away the next one becomes harder to take away and it doesn't really want to gain four electrons either to become c4 minus because every time you add an electron then you increase the electron electron repulsion so what carbon would rather do is it would rather share its four valence electrons so this is illustrated here so we have carbon in the middle with its four dots it's four valence electrons and then fluorine with its seven valence electrons right so seven valence electrons for fluorine this is its core helium so what they like to do is they like to share electrons so carbon shares one electron with one fluorine and it does that four times so if it shares electrons that's what's called a covalent bond and so another way you may simply see it depicted as like is like so so the two dots being two shared electrons so that's a covalent bond so what does it mean to atoms to share electrons that means the electron on carbon can orbit carbon's nucleus or it can orbit fluorine's nucleus and conversely the electron on fluorine the one that's shared with carbon that one can be orbiting fluorines in nucleus or it can be orbiting carbon's nucleus so shared electrons means those electrons are free to basically orbit around either of the two nuclei that's in the bond so you may see it represented like this you may see it represented like that or you simply may see it represented like this where the line so each line equals two shared two shared electrons or a covalent bond okay and then h4 and then has six electrons that it's not using for anything so those are what are called non-bonded electrons or you see them called lone pair electrons so those so each fluorine then has three lone pairs or six non-bonded electrons so those those six non-bonded electrons are basically not doing anything they're just hanging out but they are valence electrons so they could do something it just depends on if you gave the molecule something to react with then those electrons could be utilized for something but just with a molecule of cf4 they're not doing anything those those non-bonded electrons are just hanging out okay so how do you know if a molecule is going to be ionic or if it's going to be covalent uh well it's if it's a metal and a non-metal on opposite sides of the periodic table then those that's typically going to form an ionic bond um if it's two non-metals attached to each other then that's going to be a covalent bond so a metal and a non-metal typically an ionic bond two non-metals attached to each other then it's typically going to be a covalent bond where they're going to share electrons so let's put that into practice so let's take these four molecules and predict are they going to be ionic or covalent okay so pcl3 so phosphorus is here in the periodic table and then if it's attached to chlorine so if you remember where your metalloids are so metalloids are here right um so those elements in in the right boron silicon germanium um your metalloids everything to the right of that if you remember are non-metals so these are the nonmetals so phosphorus and chlorine are both nonmetals so this is going to be a covalent bond between phosphorus and chlorine uh nitrogen oh let's see so nitrogen is here oops not that so nitrogen is here and hydrogen is here so on the left of the periodic table but if you remember hydrogen is not a nonmetal hydrogen as a metal so sometimes you see hydrogen wrote here on the periodic table as well so both nitrogen and hydrogen are non-metals so this is going to be a covalent compound as well potassium and bromine so let's see potassium is here in the periodic table on the left and bromine is here so that's a metal on a non-metal on opposite sides of the periodic table so this is going to be ionic and c4h4 so carbon is here in the periodic table and hydrogen is a non so they're both nonmetals so this is going to be covalent compound okay so let's uh let's draw a lewis structure for each of these so this is the method that i use for coming up with the lewis dot structure uh so we'll go through this as we draw these molecules so we're going to count the number of total valence electrons first so phosphorus has five it's group five in the periodic table right so one two three sorry three a four a five a right four five six seven and eight so phosphorus has five valence electrons and chlorine has seven and there's three of those so there's 26 electrons so when you draw the structure of this molecule it better have 26 electrons in it bonded and non-bonded electrons or it's wrong so what we're going to do then is we're going to put the atom so this is one strategy put the atom that you only have one of in the middle and attach everything else to it with a single bond so let's put phosphorus in the middle attach the three chlorines to it with single bonds and then what we're going to do is now number three we're going to add enough non-bonded electrons to give each atom eight eight total electrons bonded and non-bonded electrons so we're going to basically satisfy what's called the octet rule so phosphorus has each bond is two electrons right so it's got three bonds so that's six electrons so it needs two more to make eight chlorine has one bond so that's two electrons so it needs three lone pairs so now that has eight and this one would have three lone pairs and that would not have three lone pairs so what is the basis of the octet rule if we go back to looking at it like this so like phosphorus or fluorine for example if it picks up one electron uh so now right now it has seven electrons in its valence shell if it picks up one more electron now it has eight and now it has a now it has a noble gas electron configuration so that's the basis for for this same with nitrogen if nitrogen picks up three electrons then now it has eight in its valence shell right so you're trying to satisfy the octet rule basically you're trying to give the atom an uh electron configuration of an of of a noble gas okay so um now what we want to do in number four so when all of the atoms are happy that means they all have an octet um or except for hydrogen hydrogen of course has one electron that only wants one more to look like helium so hydrogen only wants two but the other elements want eight so if when they have eight they're happy they have an octet they look like a noble gas so we want to count the number of electrons once everything is happy so once we have enough bonds and dots around everything to make eight and see if we have the right number of electrons so we have let's see so one let's not do that so one two three uh four five six seven eight nine ten lone pairs times two that would be twenty and then bonds we have two four six more electrons so that would be 26 electrons total which matches what we're supposed to have so that's the structure of pcl3 everything is happy we got the right number of electrons okay so on this rule number five so what if we didn't have the right number of electrons what if we had too many electrons if we had two too many electrons then that would mean we need a double bond or a ring in the structure instead so we would erase four non-bonded electrons and we would add a double bond or we would add a ring and in the next lectures there's examples of that or if we had if we counted our electrons say instead of 26 we had 30 if we had four too many electrons if you have four too many electrons and that means you either need two double bonds or you need one triple bond or you need one ring and one double bond and again there's examples of that in some of the other lectures okay so let's take this example nh3 so nitrogen has five valence electrons hydrogen has one and there's three so there's only 80 electrons to build the molecule so then we go to number two we take the atom that there's only one of which is nitrogen in this case and we put it in the middle and then we attach the other atoms with a single bond so now we've done number two and now number three we add enough non-bonded electrons to give everything eight right except hydrogen which only wants two so hydrogen's happy it's got two electrons so nitrogen needs a lone pair now it's got eight so nitrogen is happy hydrogen's happy because hydrogen's got two nitrogen's got eight count our electrons there's three bonds that's six electrons one lone pair that's two more so that's eighty electrons which is what we're supposed to have so that's the structure of the molecule okay how about kbr4 so this is supposed to be ionic right so if you remember potassium wants to give up one electron it has one valence electron bromine has seven valence electrons it's in group seven and so potassium is going to give up that electron so it's going to become k plus and br minus and it's going to be an ionic compound right so you don't want to do this because that would be totally wrong because what does a line mean a line means a covalent bond it means it's sharing electrons so this is not covalent this is ionic and you don't want to do this because if you draw them too close to each other to me it looks like you're trying to tell me they're sharing electrons and all right they're not sharing electrons so you want to draw them further far enough apart that i can clearly tell that you're not trying to share electrons between them you have a cation and you have an anion that are not physically attached to each other but they are attracted to each other okay so that would be the lewis structure for kbr and then lastly um so c4h10 so carbon's got four valence electrons and there's four carbons hydrogen's got one valence electron and there's ten of them so that would be sixteen plus 10 so that would be 26 electrons to build this molecule right okay so this one's a little bit harder than the others because we can't apply number two we don't have an atom that we just have one of that we can stick in the middle so this one's a little bit harder so how do you solve a problem like this well the way that i solve most problems is trial and error you just start trying stuff until you can find something that works so the easiest thing to do here is well let's just attach the atoms that can make more than one bond right hydrogen can't make more than one bond so if you ever put two bonds on hydrogen you know what's going to be wrong all the time because that would mean there's four electrons on hydrogen and you can't put four electrons on hydrogen you can only put two right because hydrogen is one that's one and it only wants one more electron to become one s2 so hydrogen can have two bonds so if you put two bonds on hydrogen in this class or if you go into organic chemistry and you put two bonds on hydrogen it will be wrong one hundred percent of the time right hydrogen likes hydrogen likes one bond so a carbon contain more than one bond so let's string the carbons along and then let's take the hydrogens and attach them with one bond so carbon wants four bonds right so now it's going to have eight electrons around it right carbon has four valence electrons so it can share those four to make four bonds so the two end carbons want three hydrogens the two metal ones only want two hydrogens so that would be one two three four five six seven eight nine ten hydrogens so that's the right number of carbons that's the right number of hydrogens is it the right number of electrons well how many bonds are there um there's one well let's not do that let's not make the picture too messy but there's one two three four five six seven eight nine 10 11 12 13 bonds so that's 26 k electrons so carbon's happy it's got each carbon's got an octet hydrogen's happy each hydrogen has two electrons one bond around it 26 electrons is what we're supposed to have so that's the structure of the molecule now you could have also come up with a different structure you could have come up with you could have attached the four carbons like that and then if we attach four hog uh i'll do that then if we attach each hydrogen each carbon to give four bonds now the one in the middle only wants one hydrogen because it's got three bonds to carbon so if we do that that is also c4h10o each carbon's got four bonds each hydrogen's got one bond that's 26 electrons so that's a perfectly valid structure as well so in this class if i gave you c4h10 and ask you to draw a lewis structure i would accept either one because either one is perfectly fine but i wouldn't expect you to necessarily have to come up with both of them if you go into organic chemistry i would expect you to come up with both of them right but for this class if you can just come up with one lewis structure that that works i'm happy with that okay and then lastly in this class uh this lecture just wanted to point out that some compounds uh can have both covalent and ionic bonds in the same molecule so you can have a mixture of covalent covalent and ionic bonds so an example would be sodium acetate so if we had that and we came up with a lewis structure so if we come up with this lewis structure then so if we counted our valence electrons right two carbons four electrons each three hydrogens bring one electron each um oxygen brings six electrons each and sodium has one valence electron so it should be 24 electrons in that molecule so so this is a good lewis structure um right each carbon's got four bonds so a double bond we'll talk about a double bond in the next lecture but this has covalent and ionic bonds right all every single line so line line line right line line both of those lines those are all covalent bonds but this would be an ionic bond right so oxygen is a non-metal sodium is a metal so that's going to be an ionic bond between oxygen and sodium but between oxygen and carbon carbon and carbon and carbon and hydrogen those are all covalent bonds right so you would not want to do this because if you draw a line between oxygen and sodium that means covalent bond and it's not covalent covalent it is ionic okay so we will stop there