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- Water is amphoteric, which means it can act as an acid or as a base. And so let's say that this water molecule functions as a Bronsted-Lowry acid, so it's gonna be a proton donor, and this water molecule functions as a Bronsted-Lowry base, so it's going to be a proton accepter. So a lone pair of electrons on the Oxygen take this proton and leave these electrons behind. So we're gonna form a Hydronium. So we make H3O+, so let me go ahead and draw out H3O+ here. So lone pair of electrons on this Oxygen, plus one formal charge, and let's show these electrons here in red. So these electrons in red are going to take this proton to form this bond here. So we make Hydronium, and in the process, these electrons over here in green come off onto this Oxygen. So let's go ahead and draw out what we would form. So we have that Oxygen, right, and then we had two lone pairs of electrons on this Oxygen, and we picked up another lone pair, so the ones in green right here. It's going to give this Oxygen a negative one formal charge. And this, of course, is the Hydroxide anion. So this is the autoionization of water and we can write an equilibrium expression for this reaction. So we would write Ka, but for this reaction, it's special, so we write Kw instead of Ka and Kw is called the autoionization constant. So this is the autoionization constant or you might hear different terms for this. You might hear ion product constant. So ion product constant. So whatever term you want to use or whatever term your textbook uses. And remember, when we're writing an equilibrium expression, you're going to put your products over your reactants, your concentration. So we go over here and we look at our products and we see H3O+. So we do the concentration of Hydronium ions times the concentration of our Hydroxide anions and then we don't worry about our reactants because we have pure water over here. So we're done writing our equilibrium expression. Alright, the concentration of Hydroxide ion... Actually, Hydronium is what I underlined there. So the concentration of Hydronium ion and pure water at 25 degrees Celcius has been determined experimentally to be 1.0 times 10 to the negative seven more. So this is the concentration of Hydronium ions and the concentration of Hydroxide, again determined experimentally, is also 1.0 times 10 to the negative seven more. So the concentration of Hydroxide at 25 degrees Celcius of pure water turns out to be 1.0 times 10 to the negative seven. So we can calculate the value for the autoionization constant. We can calculate Kw. Kw would just be equal to... This would be 1.0 times 10 to the negative 14. So we can go up here and write it. All this is equal to 1.0 times 10 to the negative 14. And with an equilibrium constant much less than one, you can think about the fact that the equilibrium lies far to the left. So you're not gonna have... That's why you have such low concentration of ions. So this is the autoionization of water. Now let's think about the concentration of Hydronium compared to the concentration of Hydroxide. For this example, they're the exact same. So the concentrations are the same. The concentration of Hydronium is equal to the concentration of Hydroxide, and so let me go ahead and write that. And so the concentration of Hydronium ions is equal to the concentration of Hydroxide. And when that happens, we say we're dealing with something that's neutral. So water is neutral, pure H2O is neutral. And we use that as a comparison, so if we had a solution where we had a concentration of Hydronium ions that was greater than the concentration of Hydroxide anions, that's not a neutral solution. We call this an acidic solution. So this is an acidic solution. And if we think about the opposite, if we have a greater concentration of Hydroxide anions. So a greater concentration of Hydroxide anions than Hydronium ions here, then that's a basic solution. So that would refer to a basic solution, here. So let's go ahead and do a problem. Let's say that... Let's say we had some lemon juice and the concentration of Hydronium ions of our lemon juice has been measured experimentally to be 2.2 times 10 to the negative third more. And let's say we're expected to find the concentration of Hydroxide anions, and also to classify our solution. So are we dealing with a neutral solution, an acidic solution, or a basic solution? So we can find the concentration of Hydroxide anions by using the equation that we have up here. So let me go back up here and let's look at this equation that we came up with. So the concentration of Hydronium ions times the concentration of Hydroxide anions is equal to Kw. So let's go ahead and take our equation and let's plug in our numbers. Let me go ahead and rewrite that down here. So we have our concentration of Hydronium ions times concentration of Hydroxide ions is equal to 1.0 times 10 to the negative 14. So we can plug this into here, so now we have 2.2 times 10 to the negative third or we could just make the concentration of Hydroxide anions x is equal to 1.0 times 10 to the negative 14. So we have a simple calculation, so we'll just get out our calculator here and take 1.0 times 10 to the negative 14 and then we just divide that by 2.2 times 10 to the negative third. And that will give us the concentration of Hydroxide ions. So x is equal to 4.5 times 10 to the negative 12. So x is equal to the concentration of Hydroxide anions which is equal to 4.5 times 10 to the negative 12 more. And so let's compare our concentrations, here. I'm gonna go in and write more here. So let's compare the concentration of Hydronium to Hydroxide. The concentration of Hydronium ions was 2.2 times 10 to the negative third and that number is larger than this number, right? 2.2 times 10 to the negative third is larger than 4.5 times 10 to the negative 12, so the concentration of Hydronium ions is greater than the concentration of Hydroxide. So let me go ahead and use red for this. The concentration of Hydronium is greater than the concentration of Hydroxide, and so we're dealing with an acidic solution. So lemon juice is acidic.