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conjugate acids and bases

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Common aqueous acids

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1. Group 1 and 3 Read pgs. 452-454 (Define Bronsted model of an acid and base)
2.   Group 2 and 4:  Read pgs. 457-459 (indicating the acid, base, conjugate acid and conjugate base)When chemical equations, formula or space-filling models are given.

Expert leader from each group of the same topic will meet together to discuss their findings

3. Group 5: will prepare an assessment for the class

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- [Voiceover] You've probably heard the term acid used in your everyday life. But what we want to do in this video is get a more formal definition of an acid. And particular, we'll focus on the one that is most typically used. Although we'll see future videos that there's other fairly common definitions of acids used as well beyond the one that we're going to see here. But the one that we're going to focus on is the Bronsted-Lowry definition. The Bronsted-Lowry definition of acids and bases. And this is a picture of Bronsted. This is a picture of Lowry. And they came up with this acid-base definition in the 1920s. So, we're going to do the Bronsted-Lowry, Bronsted-Lowry definition, definition of acids and bases. So, according to them, according to them, an acid, an acid is a proton, proton, or instead of writing proton we could actually write hydrogen ion donor. So why is a proton and a hydrogen ion the same thing? Well, in the most common isotope of hydrogen, we would, in it's nucleus, we would find just a proton and no neutron. And if it's neutral, you would have an electron buzzing around, jumping around in its orbital. So, you would have it's electron jumping around in its orbital. But if you were to ionize it, you're getting rid of its electron. So, if you're getting rid of it's electron, so, if you're getting rid of this, all you're going to be left with is a proton. So that's why a proton, an H plus, is usually referring to the exact same, is referring to the exact same thing. So, that's what an acid is. So what would a base be? Well, you could imagine by this definition A base, a base would be a proton, would be a proton, or you could say a hydrogen ion acceptor, acceptor. So let's make this a little bit more tangible with some examples. So one of the stronger acids we know is hydrochloric acid. Let me, let me draw. So, it's a hydrogen having a, having a covalent bond. Having a covalent bond with chlorine. With chlorine, with chlorine right over there. And if we want to, let's draw actually chlorine's lone pairs. So outside of the electron that is contributing to this pair in the covalent bond. It also has, it also has three other lone pairs. It also has three other lone pairs, just like that. So, if you were to take hydrochloric acid, place it in an aqueous solution, so it's in an aqueous solution right over here. And actually an aqueous solution, you'll see this written like that. That just means it's in a solution of water. So you could write like this, you could write hey, hydrochloric acid in an aqueous solution if you want to make it a little bit more explicit. You could say hey, look, this is going to be around some water molecules in its liquid form. Aqueous solution just means it's dissolved in liquid water. So, some water molecules in their liquid form. So, this is a water molecule. Whoops, water molecule. Right over here. So, an oxygen bonded to two hydrogens. And sometimes you'll see it written like this, that it's in its liquid, it's in its liquid form. Well, what do you think is going to happen? Well, I already said that this is a strong acid right over here. So this is going to really want to donate protons. It's really going to want to donate this hydrogen, but not let the hydrogen keep its electrons. So what's likely to happen here? Well, the both of these electrons in this pair are going to be grabbed by this chlorine. And then this hydrogen ion, because its electron was grabbed, well this could be nabbed by some water molecule passing by. Remember, in a real solution, it's not like they know what to do. They're just all bumping past each other. And based on how badly they want to do things, these reactions happen. And so you can imagine this lone pair right over here, well maybe it's able to form a covalent bond with this hydrogen. And so what's going to happen? What's going to happen? And I'll draw it with just an arrow because this reaction favorably goes, very strongly goes to the right, because this is such a strong acid. Well, then you're going to be left with, you're gonna be left with, the chlorine is now going to have its three lone pairs that it had before. And then it also grabbed these two electrons right over here. It also grabbed those two electrons right over there, so it gained an extra electron. It now has a negative charge. It is now the chloride anion. So it has a negative charge. And what about this water molecule? Well this water molecule, you have your oxygen, you have your hydrogens, you have your hydrogens, but now you don't just have two hydrogens, you grabbed this hydrogen right over here. And maybe I'll do this hydrogen in a slightly different color so that you could keep track of it. You have this hydrogen right over there. And this lone pair, this lone pair you can view it as now forming this covalent bond. You had your other two covalent bonds to the other two hydrogens. And then you still have this lone pair right over here. You still have that lone pair sitting right over there. And what just happened? Well, this water molecule just gained a proton. This hydrogen did not come with an electron. So if you just gain a proton, you are now, if you were neutral before, you are now going to have a positive charge. So what just happened? You put hydrochloric acid in a water solution, in an aqueous solution, this thing has donated a proton to a water molecule. And so, what is the acid and what is the base here? Well, when we look at the reaction this way, we see that this is the acid, the hydrochloric acid, it's literally called hydrochloric acid. And here, water is acting as a base. Water is acting as a base. And as you could see, water can actually act as an acid or a base. So, water is acting as a base. Now you might be saying, okay, this reaction goes strongly to the right, hey, but like you know, I could imagine in certain circumstances where chloride might accept a proton because it has this negative charge. And you would be right. This reaction goes strongly to the right, but once an acid has donated its proton, the thing that is left over, this is called a conjugate base. And I'll do the same color. So, this is the conjugate base of hydrochloric acid. The chloride anion. Conjugate, conjugate base of hydrochloric acid. And this right over here is the conjugate acid because you could imagine this hydronium ion, this could, under the right circumstances, donate protons to other things. Donate a hydrogen without donating electron to other things. And so this is actually the conjugate acid of H2O. Conjugate acid of water, of a water molecule. And as we'll see, water can act as an acid or a base. But this this gives you a kind of a baseline of at least the Bronsted-Lowry definition of acids and bases. And actually, one other thing I want to add. In some books here, so over here I said, hey, put this in an aqueous solution you're gonna form some hydronium, sometimes you'll see it written like this. And I'll just write it a little bit, a little bit, sometimes you'll see it like this. So you have your hydrochloric acid, and I won't draw the details this time, in an aqueous solution. So it's in a solution of water. And they'll just draw the reaction going like this, where they say hey, you're gonna be left with, you're gonna be left with some hydrogen ions, these protons. And you're going to be left with, and actually we could say it's gonna be in a aqueous solution, aqueous solution. And you're gonna be left with some chloride anions. Some chloride anions and it's in an aqueous solution. Now this isn't incorrect, but it's important to realize what they're talking about when they're talking about these hydrogen ions right over here. We know that if you have the hydrogen ions in an aqueous solution they don't just hang out by themselves. They get grabbed by a water molecule and they form hydronium. So, it's much more, I guess, it's much more close to the actual of what's happening, is if you actually talk about hydronium forming. As opposed to just the protons. 'Cause these protons in an aqueous solution, in a water solution, they're gonna be grabbed by a water molecule to form hydronium. And that's why I did it the way, this way up here.

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Bronsted-Lowry definition of acids and bases

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Voiceover] You've probably heard the term acid used in your everyday life. But what we want to do in this video is get a more formal definition of an acid. And particular, we'll focus on the one that is most typically used.

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Use the simulation to define a weak acid/base and a strong acid/base

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Conjugate acid and bases

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Give some real life application on (some reactions ) that are used for our own benefit

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Use the label "Proton" to determine whether you are acting like an acid or a base ( Bronsted Theory)