binary acids oxoacids and polyprotic acid is going to be the topics of this lesson my name is chad and welcome to chad's prep where my goal is to take the stress out of learning science now in addition to high school and college science prep we also do mcat dat and oet prep as well i'll be sure to leave links in the description below for where you can find those courses now this lesson is part of my new general chemistry playlist i'm releasing several lessons a week throughout the school year so if you want to be notified every time i post a new lesson subscribe to the channel click the bell notification all right so we're going to start with binary acids we're going to talk about trends in acid strength for these binary acids and in this case you first got to understand what is a binary acid well a binary acid if you think of that word binary the prefix bi bias on the front there which means two a binary acid is typically one that's going to be composed of exactly two elements hydrogen and one other now it turns out there are some exceptions to that like hcn is technically a binary acid so but it's not one we're actually going to even discuss so for you in this lesson you're probably again the only binary acids you're going to see are going to have two elements hydrogen and one other so in this case the most common ones are going to be the hydrohalic acids hf hcl hbr and hi so however water technically could be classed as a binary acid and so when i say it has exactly two elements i don't mean that necessarily only has two atoms so you're going to have you know some variable number of hydrogens and then one other element so cool so these are our binary acids and there's it turns out an explicit trend for their strength and that first part of that trend is they're going to increase as you go down a group and it turns out that's going to be related to size and so size it turns out is the most important factor for determining the strength of one of these binary acids and it turns out it's related to the fact of how long the bond is so bond length is defined as nucleus to nucleus distance here so if i want nucleus of hydrogen to nucleus of some other atom so it turns out the longer that bond is the weaker that bond is and a weaker bond is going to dissociate more and so we talked about the difference between strong acids and weak acids in the last lesson and a strong acid dissociates one hundred percent whereas a weak acid is only going to associate partially often less than five percent and it turns out in terms of relative strengths the more an acid associates with the stronger we say that acid is and so it turns out longer bonds are going to dissociate more resulting in a stronger acid and so that's why they increase in strength as you go down a group so it's all about increases in size weaker bonds and greater dissociation however it's also going to increase from left to right as you go across a period and this one's going to be a little bit strange but it's not about size so it turns out when you're in the same period you're going to have some you know changes in size but not like when you go down a group and stuff like this because you don't have such large differences in size well then the second most important criterion becomes the most important one for this comparison that's electronegativity cool and as you get a more electronegative atom as part of your binary acid it turns into a stronger acid and again you you know both of these trends you probably know these if you've already from the last lesson memorized your strong acids if we look at you know going down a group hf is not a strong acid but hcl hbr and hr hi are strong acids so if there's trend it must be increasing strength as you go down well same thing here you probably know that hf now is a weak acid water you'd probably consider as neutral and then nh3 we learned was a weak base and so going across for our binary acids here we go from a weak base to a neutral compound to a weak acid we're definitely increasing in acidity as we go to the right and so in this case it turns out it's not about the size of the bond or the length of the bond or anything like that it turns out it's actually the stability of the conjugate base so if you look at hf when he dissociates you get his conjugate base right here so you find out that one of the best ways we have for comparing acid strength is actually looking at the stability of the conjugate base instead and so in this case the conjugate bases has a negative charge and it turns out a negative charge is more stable on a more electronegative atom we saw this when we drew lewis structures and we had to compare like some certain resonant structures on occasion and stuff so uh and in this case then you know any one of these three when they dissociate is going to have a negative formal charge either on fluorine or on oxygen or on nitrogen and if i had my choice i'd prefer to have that negative charge on a flooring it's more stable when it's there well it turns out a more stable base is a weaker conjugate base and it turns out there's always this inverse relationship the stronger an acid the weaker its conjugate base the stronger a base the weaker its conjugate acid and so here if fluoride here is more electronegative than either nitrogen or oxygen then this fluoride conjugate base is going to be more stable than either of the other two and a more stable base is a weaker base and a weaker base comes from a stronger conjugate acid so we'll find out when we talk about the oxoacid trend a little bit we're going to allude back to this kind of a thing we'll relate you know stronger acids have weaker and more stable conjugate bases so there's also this you know always this kind of inverse relationship between stability and reactivity so something that is more stable is less reactive so increase in stability is always going to correspond to a lowering of say that's right increase in stability is going to be lower reactivity and so if you're stable the way you are well then you don't want to change and if you don't want to change well then you don't want to react to become something new it's kind of the way to think about this so but for binary acids you're definitely going to have to know this trend but now based on what we just talked about with the stronger and acid the weaker the conjugate base well now they can ask you about a trend for the conjugate basis two in a question so if i gave you four conjugate bases here and i gave you i minus br minus cl minus and f minus and said which one is the strongest conjugate base well again the trend we learned is for acidity and that's the really the trend you not need to memorize so if you're looking at the conjugate bases well again the stronger the acid the weaker the conjugate base so if i want the strongest conjugate base well then i want to choose the one that's the conjugate base of the weakest acid and so for these guys with their conjugate bases the weakest acid is hf and so hf being the weakest of these four acids would have the strongest conjugate base and so you find out that the trend in basicity for the conjugate bases is exactly the opposite trend we saw for their corresponding conjugate acids all right so now we'll move on to these oxoacids also called oxyacids an oxoacid has three elements well they're acids they have hydrogen they're oxo acids so they also have oxygen and then they have one other element that i'm just going to call the heteroatom or the hetero element just the other one all right so there's two parts to this trend for oxoacids and it turns out more oxygens more acidic and so if we take a look here we've got hclo4 on top hclo3 right below it and then we've got here clo4 minus the conjugate base and clo3 minus the conjugate base so it turns out when we look at these oxoacids so we're always going to kind of look at the strength of the conjugate base and so it turns out with these oxoacids we're going to find out that the length of the bond never comes into play like it did with the binary acids because in an oxoacid the acidic hydrogen is always bonded to an oxygen it's never bonded to that heteroatom that other element in this case chlorine for both of these examples so it turns out then we're going to be reliant on looking at the stability of the conjugate base and a more stable conjugate base again will be a weaker conjugate base and a weaker conjugate base comes from a stronger conjugate acid all right so if we take a look at these conjugate bases we've got the perchlorate ion clo4 minus and the chlorate ion clo3 minus and the difference is that extra oxygen what difference does that make well all the difference in the world it turns out so it turns out if i was going to ask you to draw the proper lewis structure here i'd include some more lone pairs and i actually included those on the study guide so all the auctions got some lone pairs here and stuff like that so i didn't want to get bogged down with the details but the big thing i want to show here is that this lovely structure actually has three additional resonance structures i could make any three out of the four octanes have a double bond and then just one have a single bond with the negative charge on that oxygen and so as a result all four of these oxygens get a chance to share that negative charge well negative charges is an indication that you have too many electrons and it turns out for stability purposes you don't want to have too many electrons you don't want to have too few electrons you want to be as close to neutral as possible most of the time and so in this case negative charge means you have too many electrons that kind of takes you away from stability but at least you're sharing that negative charge with four oxygens total so that each one only really is bearing a fourth of that negative charge whereas with clo3 minus you're only going to have three resonant structures with a negative charge only being shared between three oxygens so that each of them has negative one-third of a charge if you will and so typically the more atoms you can share that negative charge with in this case the more oxygens the more stable it's going to be and so clo4 is a more stable conjugate base that makes it a weaker conjugate base and a weaker conjugate base comes from a stronger conjugate acid so the first part of the trend more oxygen is more acidic for these oxy or oxoacids and it really comes down to resonance stabilization of that conjugate base all right so the second part of this trend here is the more electronegative heteroatom the more acidic so i wrote them in this order because usually the greater number of oxygens is the most important thing but what if both your acids have exactly the same number of oxygens like hclo4 and hbr04 well then what do you do well that's when you go here and more electronegative heteroatom more acidic and big thing here is that for the heteroatom its size does not matter like it did with a binary acid it's purely about its electronegativity no if ands or buts doesn't matter if you're comparing things in the same group or in the same period it is all about electronegativity nothing about size all right so we see what's going on here in both cases here we're going to end up with four resonant structures you're gonna have that negative charge shared by four oxygens no difference there so rule number one doesn't help us we move on to rule number two well then it goes how electronegative is that heteroatom in this case that central atom chlorine versus bromine whichever one is more electronegative is going to give you a stronger acid well you might recall again fluorine is the most electronegative element on the periodic table and so chlorine is more electronegative than bromine and as a result if i had to pick the stronger acid i'd definitely be picking hclo4 now you probably already knew this though because hclo4 is on the strong acids list and hbro4 is not on the strong acids list so if one of them stronger than the other it's probably the strong one that's stronger than the weak one all right and how do we explain this well if you look at this electronegativity gives you the ability to pull electrons towards you through the bonds we learned this with like polar structures and things of this sort and so what's happening is chlorine is actually pulling electron density away from those oxygens more than bromine is and by pulling electrons towards them chlorine's taking on a little bit of that negative charge and making all four of these oxygens a little less negative themselves and so it spreads out the negative charge even more we said the more atoms that are sharing that negative charge and the less of that burden all of them are bearing and stuff like this so uh super important in this case and so it turns out chlorine's ability to pull electrons towards him though again has nothing to do with his size everything to do with how electronegative he is and that's why hclo4 ends up being a stronger acid than hbr04 because his conjugate base is more stable now the reason we give to this is what we call the inductive effect or simply induction so if you're pulling electrons through the bonds due to your electronegativity that's induction or the inductive effect here so rule number one was all about resonance stabilization of the conjugate base so rule number two is about inductive stabilization of the conjugate base instead now one thing to note again is i just want to emphasize this really important is that this had nothing to do with size you might have been like wait a minute chad you just told me a minute ago that hcl versus hbr you told me that hbr was stronger and it was due to its size you're right i did and that's when we were talking about binary acids but now we're talking about oxoacids not binary acids and for an oxoacid size is irrelevant and again why did we say that hbr was stronger well it was due to its size was relevant because that was the size of the bond to the hydrogen and hpr had that longer bond and a longer bond was a weaker bond well with an oxoacid again that hydrogen is never bonded to the chlorine of the bromine for an oxoacid hydrogen is always bonded to oxygen and so that means the length of that bond is roughly the same in either case and size doesn't come to play a role the size of the chlorine or the bromine so hydrogens not monitor them and so the only quality for chlorine and bromine for an oxyacid that is relevant is its electronegativity its ability to carry out inductive stabilization here its size doesn't affect the length of the bond of the hydrogen at all cool so i just want to make sure you see this difference so if you're comparing binary acids they get stronger as you go down a group but if you're comparing oxo acids like we were here they're actually getting stronger up the group one is based on size for binary acids the other is based on electronegativity only for our oxoacids hopefully you can see that difference let's take a look at polyprotic acids all right polyprotic acids so a polyprotic acid is an acid that can act as a proton donor more than once so we've got phosphoric acid here and it turns out phosphoric acid has three acidic protons definitely a polyprotic acid and then sulfuric acid is the only one of our strong acids that is a polyprotic acid all the rest we'd call monoprotic so these polyprotic he's got two acidic hydrogens here all right so the big trend with polyprotic acids it turns out as they donate h's it gets harder for them to donate successive h's so if you notice here we're going to donate an h and that's going to result in a species that is now negatively charged and so it's more attracted to the h plus ions that are still attached to it which makes it harder to give them away and now he's got a negative two charge and so the one he's got left it's even he's even more attracted to it and it's even harder to give him away and so their proton donation they get less and less acidic as you go down the chain in this poly what we call a polyprotic acid series so we'll find out their equilibrium constants to describe acid strengths so and the more they dissociate to form products the greater these equilibrium constants are going to be and so what we're going to find out that is if these are getting less and less acidic as we go along their equilibrium constants are getting smaller and smaller and smaller so each successive h plus ion donation proton donation is going to have its own equilibrium constant and they get smaller as we go and we'll see that in a couple lessons all right so now we've got it with a strong acid as well and in this case what you should know about sulfuric acid only one of the h's is strong the other one's weak and so only the first donation goes to completion effectively so donation of that second proton again it gets weaker and weaker as you go so the first one went strong we shouldn't be surprised if the second one is donated to a smaller extent in this case only as a weak acid so we get 100 dissociation for the first part and probably less than 5 for the second part all right so some things you should realize here so if we look up at the weak acid so with our weak acid again because he's weak he probably is going to dissociate less than five percent let's say it's one percent or something like this well the 99 is going to stay looking like h3po4 and one percent is going to turn into each of these and so you know we're going to have this equilibrium established between all these species but most of it's still just going to be in the form of h3po4 it's going to be the most abundant species in your solution so but after that you'd figure well you're probably going to have equal amounts of h plus and h2po4 minus and you might think that except you don't so if if none of the rest of this existed then that would be true it's a one-to-one ratio however because you get a little bit further dissociation so of the h2po4 minus one it's going to form a little more h plus so this concentration is actually going to go up but it actually uses up the h2po4 minus so his concentration is going to go down and so as a result these aren't going to end up being equal in the end the h plus concentration is going to be a little bit greater than that h2 po4 minus but again it's going to get less acidic though so if one percent of the h3po4 dissociated to give us a tiny amount of this one percent essentially well again only maybe say 0.01 percent of this is now going to dissociate and then maybe 0.001 percent of this would dissociate it gets much less acidic going along the way here and it's usually quite a bit less i might even be understating it so next on the chain then would be the hpo42 minus you're gonna have a lot less a hem then of that h2po4 minus and then even less at the end of the chain here is going to be the phosphate which you'll have the smallest amount of so and these kind of relative concentrations are something you should kind of realize now in the case of h2so4 we might think oh okay i know how this works then so i'm going to have the greatest amount of h2so4 and then i'm going to have h plus so and we would think that initially these are equal but some of him is going to go away and it's going to create some more h plus so actually the h plus is greater than the hso4 minus and then finally at the end of the chain is the concentration of so42 minus except we have a problem here we started with a strong polyprotic acid not a weak one like we did here for weak acid it associates less than five percent we said well let me how about one percent so 99 stays looking like this but h2so4 is a strong acid how much does he dissociate 100 which means how much h2so4 is left in your solution well effectively zero now the truth is it's not exactly zero but it's pretty darn close to zero and so way down at the other end that's where the h2so4 is now so if you're asked to kind of you know show some relative you know concentrations of species in the solution you should definitely take that into account in the case of h2so4 being a strong polyprotic acid so but this is kind of what to be aware of and notice in the middle of these polyprotic acid trends that's where you find your amphoteric compound so h2po4 minus is amphoteric he can act as an acid going to the right or it can act as a base going back to the left he can donate a proton or he can accept a proton he's amphoteric so is the hpo42 minus and so is the hso4 minus they're all amphoteric species and that's what you often find in the middle of any of these polyprotic acid series cool get your calculators handy because in the next lesson we're going to have a little fun with the ph scale so if you found this lesson helpful so hit that thumbs up button best thing you can do to support the channel and let youtube know that other students need to see this lesson as well if you are looking for practice problems on acids bases or anything else in general chemistry check out my general chemistry master course i'll leave links in the description free trials available happy studying