Our next module is going to deal with acids and bases, which you have done acid-base chemistry in Chem Lab, or we have learned about acid-base chemistry reactions in different contexts. What we're going to look at it now is as an application of equilibrium. So there will luckily be some cases where we don't have to apply equilibrium concepts, but we are doing a big section on acids and bases now because so much of it relies on equilibrium principles. So I mean, let's define some things to get started and orient you to acids and bases in general, and then we'll bring in the equilibrium stuff. So there are a couple of different definitions that you can use for acids and bases.
For acids, if it yields hydronium ions, which is H3O plus aqueous. As these positive ions in an aqueous solution, this is the Arrhenius definition of an acid. If you say that it donates a hydrogen ion or donates H plus to another species in solution, this is using the Bronsted-Lowry definition. Now there are other, there's a Lewis definition that you will learn about and apply. No matter what, when we're talking about acids, I mean, Well, when we are talking about acids in Gen Chem in aqueous solution, we mean something that has a pH less than 7 in solution.
For bases, if it yields hydroxide ions, so yields OH- as a negative ion in aqueous solution, that's the Arrhenius definition. The Bronsted-Lowry definition is that it accepts a hydrogen ion. So it accepts H-plus from another species in solution. Again, there's a Lewis definition for bases, but what we mean is that in aqueous solution, it's a pH greater than seven.
Then we had less than seven, greater than seven. Equal to seven means that it's a neutral solution, which means it has H plus and OH minus in equal concentrations. Now for every acid-base chemistry reaction we do, the solvent is going to be water. So we are dealing completely with aqueous systems.
There will be different considerations to make when you're not in water. So keep that in the back of your mind when you learn new things about acids and bases as you go forward in chemistry. So let's look at some examples. So sodium hydroxide aqueous in solution turns into a sodium plus cation and an OH minus anion.
Something that you know already because sodium hydroxide is a soluble ionic compound. You know that it would split up into this. Sodium hydroxide is a base because this generates OH-. So we're using the Arrhenius definition. Hydrochloric acid, well, it's got acid in the name, but HCl aqueous reacting with water to make H3O plus aqueous and Cl minus.
Now, I mean, you could have written this with or without the water and been correct in what would happen to HCl in solution. But let's, I mean, look at a couple of things. The HCl generates H3O plus. So by the Arrhenius definition, HCl is an acid.
The other thing we could look at is that it donates a proton. Right, I mean it left the HCl, so it donated it. We will talk about where it goes, but donating that proton is the Bronsted-Lowry definition of an acid. Now, that proton was accepted by water. So we'll come back overall to look at water in these two reactions, but proton accepted by water, water was the base in this reaction.
So hydrochloric acid, water, the two reacted with each other. HCl acted like an acid, H2 acted like a base. If the reaction is NH3 aqueous plus water in equilibrium with NH4 plus aqueous and OH minus, again, a reaction that you would be able to tell me happens.
The NH3, it generates an OH- in solution. It's hard to tell directly. You can't see it from the starting molecule, but it generates an OH-.
It also accepts an H-plus from water, right? The H-plus that got connected to the NH3 came from water. So NH3 accepting a proton, that's a base under Bronsted-Lowry definitions. Water donating a proton, that's an acid under Bronsted-Lowry definitions. So we have water acting as an acid, water acting as a base, depending on the context of the reaction.
So water is amphoteric or amphiprotic. Now specifically it's an amphiprotic because it can donate or accept a proton. Amphoteric is a little more generic, but it is a substance that can behave as an acid or a base. It completely depends on the context. So we have seen this on just the previous slide, but some acid-base reactions exist as equilibria.
We'll talk about how to figure out. when this happens in a little bit. But if we're looking at this reaction, ammonia reacting with water in equilibrium with ammonium and hydroxide, in the forward reaction, so going to the right, NH3 was the base and H2O was the acid. In order for the reverse reaction to happen, NH4 plus would donate a proton to OH minus. That would make NH4 plus the acid and OH minus the base in the reverse reaction.
Now we connect these to each other. So the NH4 plus acid came from the NH3 base. So we call them a conjugate acid base pair. So NH3 When I added a proton to it, I got NH4+. NH3 acts like a base, so it's the base.
NH4 plus acts like an acid. It's the acid, but we call it the conjugate acid because it is connected to the NH3. H2O and OH- are also a conjugate pair.
H2O is the acid and OH- is the conjugate base in this situation. So let's identify these pairs in the reactions that are given here. So H2SO4 aqueous reacting with water to produce HSO4- and H3O+.
H2SO4 donates a proton to water. That makes H2SO4 the acid. And starting from just identifying the acid, I could go over and find when this acts like an acid, what am I left with? HSO4-. That is the conjugate base.
So an acid will have a conjugate base, a base will have a conjugate acid. I mean, let's explicitly mark them. Water accepts a proton from H2SO4. So water is acting like a base.
And on the other side of the reaction, its conjugate acid is H3O+. Of course, I mean, we said that water was amphiprotic. Let's identify it in these reactions.
So HCO3-, reacting with water in equilibrium with H2CO3 and OH-, HCO3- is accepting a proton from water. That makes HCO3- the base, and H2CO3 its conjugate acid. This makes H2O the acid, and OH- its conjugate base. Okay.