So you've probably been thinking that it's been a while since you have done a resonance structure or thought about periodic trends. Well, I have great news for you. It's time for us to do this with as-is and bases. Okay, so you're probably not thinking any of that, but now is the time when we connect multiple topics and models together and use them to explain what we see.
in terms of how acids and bases behave in solution. Now this is specifically going to address acids, but you can apply most of these things to bases as well. So acid strength and molecular structure, there are three contributing factors. One is bond length, one is the difference in electronegativities, and the other is resonance stabilization of the conjugate.
Now we're going to do acid. So I have resonance stabilization of a conjugate base. Well, if we're going to have resonance anything, we need resonance structures, which means we need Lewis structures. So we will be bringing that back. First, let's look at the acid halides.
And these are things like HF, HCl, HBr, Hi, and the length of the halogen. hydrogen bond. These are given in units of angstroms where one angstrom is 10 times 10 to the minus meters. The bond length in HF is 1.01 angstroms.
HL is 1.36, HBAR is 1.51, and HI is 1.7. So what you're seeing is the halogen is getting bigger, the bond is getting longer, and a longer bond is a weaker bond. Meaning if it's a longer bond, it's easier to break.
So if it's easier to break, then you have more H plus in solution, which we define as being a stronger acid. Our strong versus weak definitions are just about percent dissociation or percent ionization. So what this means is that HI would be the strongest in this series. Then HBR, then HCL, then HF would be the weakest in this series.
Now, okay, that's fine, but we have kind of already categorized some of these as strong or weak. So I just, I want to get this on here, right? Like we are only calling hydrofluoric acid a weak acid.
We're calling all of the other ones strong. This is just because that is the behavior we observe in water. If you had these things in a different solvent, they would behave differently. You might get a differentiation between strong and weak at a different spot. We just happen to have three that fall into the strong category because we're in water.
If we ranked them though, right, HI would be stronger than HBR. It's just that under these conditions, all three of them can be strong. Now, don't confuse.
Strength with reactivity or a dangerousness, right? Hydrochloric acid is a strong acid, but it's something that you use in lab, even at higher concentrations. Hydrochloric acid has to be stored in a plastic container because it will dissolve glass.
Hydrochloric acid will leach the calcium out of your bones, right? So It may be a weak acid, but that is only telling us about how much it dissociates in solution. And it doesn't tell us about what terrible things it does once it gets there. Okay, that is bond length, which is really a proxy for bond strength.
The next is difference in electronegativity. A larger difference in electronegativity means that you have a more polar bond. This is something we've done before. A more polar bond means a higher partial positive on the hydrogen, which makes the hydrogen more likely to dissociate.
So a higher partial positive on the hydrogen makes it easier for the hydrogen to attach to water and make H3O+. The comparison here, we're going to go across the second row of the periodic table. NH3, H2O, HF, so nitrogen, oxygen, and fluorine with hydrogens attached.
Differences in electronegativity, so nitrogen and hydrogen, the difference is 0.8. Oxygen and hydrogen has a difference of 1.2, and hydrogen and fluorine have a difference of 1.8. So higher difference in electronegativity, HF would be a stronger acid than H2O, which would be a stronger acid than NH3. Now, because we're in water, because we have some reference points at this point in the module, you...
probably could have written this down because HF is acidic, H2O is neutral, and NH3 is basic. But still, if we're ranking in terms of acidity, well, how good is NH3 at being an acid? It's terrible. It's so terrible that it acts like a base.
We also see some electronegativity differences coming up in the oxoacids. Now, these are like... Some of the oxoacids that we have looked at, we've also looked at the ones that have more oxygens.
In something like an oxoacid, what you will have is a halogen connected to an oxygen, and the oxygen is connected to a hydrogen. The oxygen-hydrogen bond is the acidic bond. So I've got this still generically drawn up here in equilibrium with its conjugate base, which is the XO.
minus and the H plus. So the oxygen hydrogen bond is the one that's breaking, but the trend is related to the identity and electronegativity of the halogen in the oxoacid. So what we find is that as the electronegativity increases, the acidity increases. So HClO, electronegativity of chlorine is three.
Ka is 2.9 times 10 to the minus eighth. HBRO, electronegativity of bromine is 2.8. Ka is 2.3 times 10 to the minus ninth.
So less electronegative halogen, smaller Ka, less dissociated, weaker acid. HIO, electronegativity of 2.5 for iodine and a Ka of 2.3 times 10 to the minus 11th. So as the halogen electronegativity increases, the acidity increases. Now, when we look at this, it's really easy to initially say like, oh, well, the halogen is pulling on the electrons in the halogen oxygen bond.
So oxygen is going to pull on the hydrogen. And that's how you get the partial positive. That's not, it's not quite what happens because all of these halogens are less electronegative. than the oxygen. So in every case, if you were to draw the bond dipole for the halogen-oxygen bond, and I'm just going to do this on the conjugate bases, so ClO minus, BrO minus, and Io minus, the bond dipole in the chlorine-oxygen bond is smaller than the bromine oxygen, which is smaller than the iodine oxygen.
The difference in electronegativity between the halogen and oxygen is smallest with chlorine and biggest with iodine. So how does this go together? That is the question. So chlorine is the most electronegative.
What you would say is that it is helping stabilize the conjugate base by sharing the load of the negative charge. So when that hydrogen leaves, and oxygen has a negative charge on it. In the case of chlorine and oxygen, chlorine is like not pushing electrons onto the oxygen.
Oxygen is not trying to pull as many electrons in that chlorine-oxygen bond because the difference in electronegativity is smaller. In the I-O minus, right, oxygen has the negative charge because we took the proton off. But then iodine is also like, yeah, here, you can have my electrons.
And then oxygen is just great. Oxygen is carrying the full load of this negative charge. In ClO minus, oxygen is not trying to also carry a ton of electron density from the chlorine through the chlorine-oxygen bond. So you may not like the word sharing a load, but chlorine is not actively making this negative charge.
bearing worse compared to iodine, which is relying on oxygen to just carry everything. Now, this works for HClO4, HClO3, HClO2, right? We have a whole series of oxoacids, which we have looked at with Lewis structures and resonance structures before. So, right, this is one of the things that's happening.
We're just looking for conjugate base. Happy means stronger acid. Now, the next piece of what we're going to do also works within this series of oxoacids.
We're going to look just at the conjugate base and at resonance stabilization within these structures. So we're going to do NO3- versus NO2-, because we're going to look at HNO3 versus HNO2. So nitric acid is a strong acid. And nitrous acid is a weak acid. The only difference is the one oxygen in the anion.
So when you want to compare the strength of an acid, unless it's one of those binary acids that we had at the beginning, you kind of have to look at the conjugate base and see what is it doing in solution. So we're going to compare. the conjugate bases NO3- NO2- and the way we are going to compare them is with Lewis structures and resonance. So for each of these conjugate bases, draw the Lewis structure, the best Lewis structure, and any equivalent resonance structures, and then let's see how they look. For NO3-, so 24 electrons, the best Lewis structure that you can come up with has nitrogen in the middle with a positive charge, a double bond to one oxygen, and single bonds to the other two oxygens.
Single bonded oxygens have negatives. So this, it doesn't look great, but because you can't expand the octet for nitrogen, this is really the best that we can do. This means that my equivalent resonance structures, I have two others because that double bond. between nitrogen and oxygen could have been with any of the oxygens. Now, okay, we had two negative charges and one positive, but the negative charge is shared over three oxygens.
So each oxygen doesn't have to carry the full brunt of that negative charge in any structure. If we look at NO2 minus, So this has 18 electrons, so my best Lewis structure would have a lone pair of nitrogen, one double bonded oxygen, one single bonded oxygen, and because that double bond could have been on either side, we do have one equivalent resonance structure. So what we would say about this is that the NO2- is less stable because there is less resonance stabilization. And so the amount of resonance stabilization just depends on the number of resonance structures that exist that are equivalent. So NO2 minus, less stable, less resonance stabilization, fewer resonance structures.
NO3 minus, more resonance structures, more stable conjugate base. A more stable conjugate base means a stronger acid. Essentially, and if you go back and listen to anything at the beginning, this should connect to some of our intros to definitions about acids and bases. If the conjugate base is really happy being on its own, meaning it's stabilized, then the conjugate base can stay on its own.
It is better able to stay on its own, meaning that the H plus is not connected. making it a stronger acid. So the increase in resonance stabilization or the presence of more resonance stabilization is why NO3 can stay on its own, making HNO3 a strong acid compared to HNO2.
So something you should do, try comparing the conjugate bases for perchloric acid, chloric acid, and other acids. Chloris acid and hypochloris acid. Make sure I got all my names right. This is a series of oxo acids.
The HClO4 and HClO3 are both strong acids. HClO2 is weak. HClO is weak.
but don't forget you have to compare the conjugate bases so CLO4 minus CLO3 minus and so on.