Chemistry: Arrhenius/Bronsted-Lowry: Acid, Base

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We've already talked about the Arrhenius definition of an acid. An acid is something that, when dissolved in water, increases the concentration of protons above that that exist in neutral water. And, conversely, an Arrhenius base is something that increases the concentration of hydroxide in water above the concentration that exists in neutral water. So examples of acids are things like hydrochloric acid and acetic acid. Hydrochloric acid is stomach acid and acetic acid is the acid found in vinegar. Examples of bases, sodium hydroxide, which is drain cleaner, and ammonia, which is something that you mop the floors with. Now, both of those definitions, both the definition for the acid and base, are very phenomenological. In other words, you put something into the water, and then you just measure to see what the effect on the water is. And a more modern - but actually, not much more modern, dating from around 1923 - more modern definition of an acid and a base is that an acid is something that can donate a proton and a base is something that can accept a proton. And, if you think about it, that's much more on the atomic molecular level. And that concept of an acid base is due to two gentlemen, Brønsted and Lowry, who came up with the idea independently.
So let's look at what exactly this means. Here we have the reaction of nitric acid and water or, actually, taking nitric acid and putting it into water. And nitric acid, obviously, we call it an acid. What makes it an acid in the Brønsted-Lowry concept is that it has the ability to donate a proton. So, if it has the ability to donate a proton, something has to accept that proton. In other words, if we have the ability to donate money, something has to accept that money. And what accepts that money is the water. And so, water, in this case, is functioning as the base. And after the proton has been transferred, nitric acid becomes nitrate and the water becomes protonated and we call this species H[3]O+ hydronium cation. And we'll have more to talk about this later on.
So you can see that an acid reacts with the base to go to a new base and a new acid. And that's what happens when we transfer a proton in solution. Similarly, here's the reaction of ammonia, which is a base in the Arrhenius idea, but it's still a base in the Brønsted-Lowry idea. And it has the ability to accept a proton. Well, where does it get that proton? It gets it from water, and so it forms ammonium, where it has now taken that proton away from water. And then the water, having given up a proton, becomes hydroxide. So you can see that water can function as both an acid and a base, and we're going to talk about that later on. But on this page, what I want you to think about is just the idea that we're transferring this quantity, called the base - and analogously, you could talk about transferring money - and so, the idea is that something that can accept that money is a base and something that can donate the money is an acid.
Now, you'll also notice that these things appear in pairs. We have acid base, so nitric acid-nitrate and water hydronium base is ammonia-ammonium and water-hydroxide. And so, things appear in pairs in the Brønsted-Lowry idea. And we talk about them in terms of conjugate acid base pairs. So ammonia and ammonium are a conjugate acid-base pair. Another way to use that terminology is to say that ammonia is the conjugate base of ammonium. And a third way to say it is that ammonium is the conjugate acid of ammonia. So here we would say that chloride is the conjugate base of hydrochloric acid. And you notice that, in these conjugate acid-base pairs, what connects them is a proton. So ammonia to ammonium, there's one more proton. And hydrochloric acid to chloride, there's one less proton. And so these things form conjugate acid-base pairs.
Now, a concept in acids and bases that we haven't introduced yet is the concept of a strong acid or a strong base versus a weak acid or a weak base. And it's very important to understand that this has nothing to do with concentration. It has nothing to do with concentration. What it has to do with is the behavior of the substance when you put it into water. So a strong acid is something that's completely dissociated in aqueous solution. And examples, and this is pretty much a complete list of strong acids, HI, HBr, perchloric acid, hydrochloric acid, chloric acid, sulfuric acid and nitric acid. And I've underlined those compounds that you're very like to find in the lab, things like hydrochloric acid, sulfuric acid or nitric acid. Some of these others are a little less common. You can use them, but you're likelihood of encountering them is less. Now, when we say completely dissociated, what do we mean? We mean that, if you put, for instance, hydrochloric acid, into water, it completely comes apart into H+ and Cl minus. For all intents and purposes, it is completely dissociated, and that's what makes it a strong acid.
And, in contrast, things that are weak acids are only partially dissociated. So, for instance, here's acetic acid, which is the acid in vinegar, and the an equilibrium is established when you dissolve it in water, in which it becomes acetate and a proton, hydrofluoric acid becomes fluoride anion and a proton, but these are in equilibrium. So it isn't 100 percent over to the right, it's only partially dissociated. And I should point out that another way to write this - by the way, this is bisulfate and this is carbonic acid - another way to write this is to write it as HF plus water going to H[3]O+, plus fluoride. And what I've done is use the shorthand to write H+. But we know that, when I write H+, what I really mean is a solvated proton, H[3]O+. And I'll have some more to talk about that later on.
Now, strong bases are those that are completely dissociated. Just like strong acids are completely dissociated, strong bases are completely dissociated. Or, alternatively, there are those that are completely converted to hydroxide in aqueous solution. So the kinds of things that you're like to encounter are things like sodium hydroxide, potassium hydroxide, calcium hydroxide, but also sodium hydride, which reacts with water to form sodium cations plus hydroxide plus giving off some hydrogen gas, and sodium oxide, which reacts with water to form sodium hydroxide. But, actually, [unintelligible] of sodium hydroxide. And, quite naturally, weak bases are those that are not completely converted to hydroxide in aqueous solution, so ammonia reacts with water to form ammonium plus hydroxide, but not completely. So there's an equilibrium established. We'll put some numbers to these equilibria later on. Similarly, fluoride anion reacts with water to form hydrofluoric acid or hydroxide. And this is not a complete reaction, so fluoride doesn't completely go to HF and hydroxide.
Now, another point that I'd like to make about conjugate acid-base pairs is sort of their relative strength in their respective categories. So the conjugate base of a strong acid is a very weak base. So hydrochloric acid, we said, was a strong acid. It's fully dissociated. And another way to think of this is that chloride doesn't have much of an affinity for the proton. In other words, if chloride is willing to give away its money in the form of a proton, then the other way you can say that is that chloride without the money doesn't have any great affinity to get that proton back, to get that money back. And so, chloride is a very generous anion and it's very willing to give up the money. So that's why HCl is a very strong acid, but chloride is the conjugate base of HCl and it doesn't have any great propensity to get that proton back. So it makes chloride a very weak base. And, similarly, the conjugate acid of a strong base is a very weak acid, so hydroxide is a strong base. Sodium hydroxide is a strong base. And so, hydroxide does have a great propensity to pickup that proton, and so that means water, it's conjugate acid, doesn't have any great propensity to give up the proton.
Well, that notion of this sort of inverse relationship between the strength of an acid and the weakness of its conjugate base is expressed in this table that I have for you here, where we've organized the acids and bases, with strong acids up here at the top, going down to weaker and weaker acids. And the idea is that, if we have a strong acid, then it's conjugate base is relatively weak. As we go to weaker and weaker acids here - as we go down the table, we're going to weaker and weaker acids - then the conjugate bases are getting stronger and stronger. All the way down at the bottom, we have those species that have negligible acidity, and the other way to say that is that the conjugate bases of those acids are very, very strong. So, in other words, these down here have a great propensity to pickup that proton, these bases down here. Now, it's important to understand that these bases in here, up to this point, these are still weak bases, but they're relatively stronger. So we use the notion of weak and strong both in an absolute sense and in a relative sense. So, for instance, these are the strong acids, stopping right here, but these are relatively strong. So, bisulfate, for instance, is a relative strong weak acid. And, going all the way down to getting weaker and weaker acids, again, we get stronger conjugate bases.
Now, the last point I would like to make about acids and bases here is that we can talk about concentration. So again, remember, strong and weak have nothing to do with concentration, but we can talk about concentration. And we use the terms dilute and concentrated. So, for instance, dilute is something where there's not a whole lot of solute in the solution, but it's a relative concept. And, similarly, concentrated would mean that there's a lot of solute. But it's, again, a relative concept. So I would characterize 10^-4 molar hydrochloric acid as being a dilute strong acid; strong, because HCl is fully dissociated, but dilute, because 10^-4 isn't a whole lot of hydrochloric acid, particularly compared to concentrated hydrochloric acid, which is probably close to ten molar. In contrast, one molar hydrofluoric acid I would characterize as a concentrated weak acid; concentrated, because there's a relatively large amount of solute, 10^4 times as much solute as the solution above, but it's weak, because hydrofluoric acid is weak. It's not fully dissociated when you put it into water. So again, these terms are relative, and so there's some ambiguity associated with the terms dilute and concentrated. And then also, strong and weak are both absolute terms and relative, as in relatively strong weak acid or relatively weak strong acid. As we work with these more in this tutorial, I think you'll become more comfortable with these ideas, but I wanted to introduce them here, just to make sure we had a common language for the rest of our discussion.
Acids and Bases
Acid-Base Concepts
Arrhenius/Bronsted-Lowry Definitions of Acids and Bases Page [2 of 2]