Understanding Acids and Bases

This is the presentation for section 14.1. We're going to be taking a look at acids and bases and basically looking at the properties and then defining acids and bases. And so first of all, we need to kind of take a look at the properties of acids and bases. Acids have a tendency to taste sour. They turn litmus red. Acid turns red. And they definitely react with metals and they produce hydrogen gas. Bases have a tendency to taste bitter. They turn litmus blue. So to keep that straight, bases turn blue and they feel slippery. And most cleaning materials are basic and they kind of have that slippery feeling to it. And so we've got these properties of acids. And now we've got to kind of figure out, well, where are these properties coming from? And the first definition we're going to take a look at is the Arrhenius definition, which is historically the first one. And Arrhenius basically said acids produce hydrogen in water and bases produce hydroxide or OH-in water. What we found earlier in the, when we were looking at classifying reactions, we found that hydrogen from acids combines with OH from bases to produce water and we call that a neutralization reaction so we saw that a while a while back um in slot notes for quiz number four the problem with the Arrhenius definition is it's not all encompassing it doesn't show we we have compounds that will act like acids and bases And the ones that act like acids do donate hydrogen. But what we find is we have some bases out there that don't even have hydroxide in them. And so Bronsted and Lowry continued the definition, made it more encompassing, and said that instead of bases donating an OH, they accept the hydrogens from acids. And that comes from that idea of neutralization reactions. We can also find that hydrogen can be called a proton because if you go back to the slot notes quiz number two, we talked about atomic structure. And if you take a hydrogen atom that has one proton and one electron and you remove the electron, the only thing you have left is a proton. And that's what gives acids their strong and unique characteristics or properties. So acid-base reactions involve a proton exchange but we still just call it an H plus so classic hydrochloric acid reacts with water to produce the H3O plus we're gonna talk a little bit more about where that comes from in a moment and chlorine ions well if we have NH3 which is ammonia it'll turn litmus blue but it has no OH in it so if you'll notice it pulls a hydrogen off of the water to produce ammonium ions. That's at NH4 superscript plus. And if you'll notice, vicariously, it donates OH. So when we looked at the Arrhenius definition, it wasn't totally wrong. It's just we need to take a look at compounds that don't have OH directly in them. So if we take a closer look, the Bronsted acid is a proton donor, and now a Bronsted base is a proton acceptor. And so if you take a look, we've got this example of ammonia with water that was on the previous slide, but we're going to expand it out a little bit. Well, the NH3 acts like a base, and the H2O acts like an acid. That sounds kind of weird because you've always been told that water is neutral. Yeah, well, it's neutral because it has a hydrogen and a hydroxide. But if we shift the equilibrium, now we, you know why we're talking about Le Chatelier's in the previous section, previous set of slot notes. And if you'll notice, we have a double headed arrow. So now we have an acid and a base on the right hand side. So in our acid base reactions, we actually end up having two acids and two bases. So typically the base and the acid on the left hand side are left alone, but on the right hand side we add the word conjugate, one to tell us that it's on the right hand side of the equation and two that it's related to something on the reactant side. And so if you'll notice the NH3 and the NH4 goes together and the H2O and the OH go together. So as the OH, remember, going backwards, picks up a hydrogen, bases accept hydrogens. It turns into H2O liquid or the acid. So the species that forms when a proton is removed from an acid is now called the conjugate base of the acid. So if the acid is HB, the conjugate base is B minus. Sometimes we'll substitute. B means the conjugate base, but sometimes you'll see AHA instead of HB because A is actually the anion. Remember, negative ion. So you'll see an HA and an A minus. they're still representing the same thing. They're still representing the acid in its conjugate. There are some species that can do either. If you take a look, we got water. If we remove a hydrogen from water, it becomes OH. But H2O can pick up a hydrogen and it can become positive because the H2O is bent and it's polar and it's got this really negative N to it on the oxygen side of it. pulls in the hydrogen and so we get this H3O plus or hydronium ion. So let's take a look at some conjugate pairs. HCl is a classic acid. Notice the hydrogen is hydrochloric acid and so its conjugate base would be the Cl. Notice how they're related to each other through the chlorines. And that means the H2O and the H3O are going to be related to each other. And so the H2O is going to act like a base when I add an acid to it. And the H3O is going to be its conjugate acid. If I take a look at the NH3, it's a base because if you look on the right hand side, it's picked up by hydrogen. So its conjugate acid is NH4+. That means the H2O is an acid and the OH is the conjugate base. Notice they always come in pairs. And in reality, if you can find one, you can probably find the other three. So what is not a characteristic property of acids is react with CO2 to form carbonate. No, they react with carbonates to produce CO2. That's your classic vinegar and bacon soda, but definitely the neutralized bases. Remember, acid turns red, so the litmus goes from blue to red. They'll produce hydrogen gas. So that H3O plus actually has a name, and you'll hear me mention it's called the hydronium ion. H2O does not like to exist by itself. If there's hydrogens around, it likes to pick up those spare hydrogens or protons and become H3O+. But as a lazy, efficient chemist, I'm probably going to just write H+, and make the assumption that everybody understands that it's technically an H3O+. So we've got the reaction of... H2CO3, H2O, HCO3-and H3O+. And we want to know what the Bronsted acids are. Well, remember, acids by definition donate hydrogen. And we're going to have one on the left-hand side and one on the right-hand side. So it looks like the H2CO3 is going to lose a hydrogen and the H3O is going to lose a hydrogen as they respectively go across that double-headed arrow. Now we've got HSO4 plus OH minus. Let's get a little bit more exotic. SO4 2 minus plus H2O. It says the conjugate acid base pairs. So if you'll notice, they look like each other. These ones with sulfur in them go together. So if I pull those two out, what I get left with is H2O and OH. And what I've done is I've done the acids in red. And the bases in blue notice the HSO4 loses a hydrogen to become SO4 and H2O loses a hydrogen to become OH. And so this says identify the conjugate base. Conjugate means it's on the right hand side. And remember the base is the one that's going to pick up or the one that's going to have the sulfur in it. And so we're going to end up with the SO4 2-. Remember bases gain hydrogen. Conjugate means look on the product or the right hand side. This says identify the base, which means look for the one, not the conjugate, but the base on the left-hand side. Remember, the bases are things that pick up hydrogens. And so that's going to be our CO3 is going to pick up a hydrogen to look like HCO3-. Hopefully you've realized that this really sounded a little bit complicated, but it's really fairly simple once you get this idea that acids lose hydrogens, bases gain hydrogens, they pair up. with similar that they look similar and the conjugates are on the right hand side so now we want the acid left hand side of SO4 okay acids lose hydrogens so it's our HSO4 so it says show by suitable net ionic equations that each of the following species can act as a Bronsted-Lowry acid. So what we're trying to show is how each of these can lose a hydrogen. Well, we've seen the H3O and we've seen the NH3. D is what's called an organic acid or a carboxylic acid. Notice how it loses its hydrogen off the end. Sometimes you'll see this written as COO-. And then if, and the reason you have F is you notice the HSO4 negative loses hydrogen, so it becomes 2 negative. Got to make sure our charges balance. This is now, how do they act like a base? Remember, bases pick up hydrogens. So that's, remember, we talked about the H3O or the hydronium ion for A. we've got the ammonium ion the reason i've had that a couple times is that that's a classic weak base type problem cn picks up a hydrogen becoming hydrogen cyanide that cn ion's really dangerous um it's a really strong poison and if you'll notice the h2po4 which is negative now picks up a hydrogen to become the h3 PO4. Notice there's now no charge. So the charges all balance each other out too. Number 16, it says what is the conjugate acid of the following? What is the conjugate base? I've done the acids in red and the bases in blue. The acid picks up the hydrogen. The base loses a hydrogen. Again, this whole section, acids lose hydrogens, bases pick up hydrogens. And then if it's the conjugate, we're looking on the right-hand side. And finally, we have things that can be amphiprotic that can either gain a hydrogen or lose a hydrogen. The OH gains a hydrogen to look like H2O or loses a hydrogen to become O2-. The S2-can't be amphiprotic because it doesn't have a hydrogen to donate. And then we have the HSO4. Okay, notice it can gain a hydrogen and lose a hydrogen.

Acid turns red. And they definitely react with metals and they produce hydrogen gas. Bases have a tendency to taste bitter. They turn litmus blue. So to keep that straight, bases turn blue and they feel slippery.

And most cleaning materials are basic and they kind of have that slippery feeling to it. And so we've got these properties of acids. And now we've got to kind of figure out, well, where are these properties coming from? And the first definition we're going to take a look at is the Arrhenius definition, which is historically the first one.

And Arrhenius basically said acids produce hydrogen in water and bases produce hydroxide or OH-in water. What we found earlier in the, when we were looking at classifying reactions, we found that hydrogen from acids combines with OH from bases to produce water and we call that a neutralization reaction so we saw that a while a while back um in slot notes for quiz number four the problem with the Arrhenius definition is it's not all encompassing it doesn't show we we have compounds that will act like acids and bases And the ones that act like acids do donate hydrogen. But what we find is we have some bases out there that don't even have hydroxide in them.

And so Bronsted and Lowry continued the definition, made it more encompassing, and said that instead of bases donating an OH, they accept the hydrogens from acids. And that comes from that idea of neutralization reactions. We can also find that hydrogen can be called a proton because if you go back to the slot notes quiz number two, we talked about atomic structure.

And if you take a hydrogen atom that has one proton and one electron and you remove the electron, the only thing you have left is a proton. And that's what gives acids their strong and unique characteristics or properties. So acid-base reactions involve a proton exchange but we still just call it an H plus so classic hydrochloric acid reacts with water to produce the H3O plus we're gonna talk a little bit more about where that comes from in a moment and chlorine ions well if we have NH3 which is ammonia it'll turn litmus blue but it has no OH in it so if you'll notice it pulls a hydrogen off of the water to produce ammonium ions.

That's at NH4 superscript plus. And if you'll notice, vicariously, it donates OH. So when we looked at the Arrhenius definition, it wasn't totally wrong. It's just we need to take a look at compounds that don't have OH directly in them.

So if we take a closer look, the Bronsted acid is a proton donor, and now a Bronsted base is a proton acceptor. And so if you take a look, we've got this example of ammonia with water that was on the previous slide, but we're going to expand it out a little bit. Well, the NH3 acts like a base, and the H2O acts like an acid.

That sounds kind of weird because you've always been told that water is neutral. Yeah, well, it's neutral because it has a hydrogen and a hydroxide. But if we shift the equilibrium, now we, you know why we're talking about Le Chatelier's in the previous section, previous set of slot notes.

And if you'll notice, we have a double headed arrow. So now we have an acid and a base on the right hand side. So in our acid base reactions, we actually end up having two acids and two bases.

So typically the base and the acid on the left hand side are left alone, but on the right hand side we add the word conjugate, one to tell us that it's on the right hand side of the equation and two that it's related to something on the reactant side. And so if you'll notice the NH3 and the NH4 goes together and the H2O and the OH go together. So as the OH, remember, going backwards, picks up a hydrogen, bases accept hydrogens. It turns into H2O liquid or the acid.

So the species that forms when a proton is removed from an acid is now called the conjugate base of the acid. So if the acid is HB, the conjugate base is B minus. Sometimes we'll substitute.

B means the conjugate base, but sometimes you'll see AHA instead of HB because A is actually the anion. Remember, negative ion. So you'll see an HA and an A minus.

they're still representing the same thing. They're still representing the acid in its conjugate. There are some species that can do either. If you take a look, we got water.

If we remove a hydrogen from water, it becomes OH. But H2O can pick up a hydrogen and it can become positive because the H2O is bent and it's polar and it's got this really negative N to it on the oxygen side of it. pulls in the hydrogen and so we get this H3O plus or hydronium ion. So let's take a look at some conjugate pairs. HCl is a classic acid.

Notice the hydrogen is hydrochloric acid and so its conjugate base would be the Cl. Notice how they're related to each other through the chlorines. And that means the H2O and the H3O are going to be related to each other. And so the H2O is going to act like a base when I add an acid to it.

And the H3O is going to be its conjugate acid. If I take a look at the NH3, it's a base because if you look on the right hand side, it's picked up by hydrogen. So its conjugate acid is NH4+. That means the H2O is an acid and the OH is the conjugate base.

Notice they always come in pairs. And in reality, if you can find one, you can probably find the other three. So what is not a characteristic property of acids is react with CO2 to form carbonate.

No, they react with carbonates to produce CO2. That's your classic vinegar and bacon soda, but definitely the neutralized bases. Remember, acid turns red, so the litmus goes from blue to red. They'll produce hydrogen gas.

So that H3O plus actually has a name, and you'll hear me mention it's called the hydronium ion. H2O does not like to exist by itself. If there's hydrogens around, it likes to pick up those spare hydrogens or protons and become H3O+.

But as a lazy, efficient chemist, I'm probably going to just write H+, and make the assumption that everybody understands that it's technically an H3O+. So we've got the reaction of... H2CO3, H2O, HCO3-and H3O+. And we want to know what the Bronsted acids are.

Well, remember, acids by definition donate hydrogen. And we're going to have one on the left-hand side and one on the right-hand side. So it looks like the H2CO3 is going to lose a hydrogen and the H3O is going to lose a hydrogen as they respectively go across that double-headed arrow. Now we've got HSO4 plus OH minus. Let's get a little bit more exotic.

SO4 2 minus plus H2O. It says the conjugate acid base pairs. So if you'll notice, they look like each other.

These ones with sulfur in them go together. So if I pull those two out, what I get left with is H2O and OH. And what I've done is I've done the acids in red. And the bases in blue notice the HSO4 loses a hydrogen to become SO4 and H2O loses a hydrogen to become OH. And so this says identify the conjugate base.

Conjugate means it's on the right hand side. And remember the base is the one that's going to pick up or the one that's going to have the sulfur in it. And so we're going to end up with the SO4 2-.

Remember bases gain hydrogen. Conjugate means look on the product or the right hand side. This says identify the base, which means look for the one, not the conjugate, but the base on the left-hand side. Remember, the bases are things that pick up hydrogens.

And so that's going to be our CO3 is going to pick up a hydrogen to look like HCO3-. Hopefully you've realized that this really sounded a little bit complicated, but it's really fairly simple once you get this idea that acids lose hydrogens, bases gain hydrogens, they pair up. with similar that they look similar and the conjugates are on the right hand side so now we want the acid left hand side of SO4 okay acids lose hydrogens so it's our HSO4 so it says show by suitable net ionic equations that each of the following species can act as a Bronsted-Lowry acid.

So what we're trying to show is how each of these can lose a hydrogen. Well, we've seen the H3O and we've seen the NH3. D is what's called an organic acid or a carboxylic acid.

Notice how it loses its hydrogen off the end. Sometimes you'll see this written as COO-. And then if, and the reason you have F is you notice the HSO4 negative loses hydrogen, so it becomes 2 negative. Got to make sure our charges balance. This is now, how do they act like a base?

Remember, bases pick up hydrogens. So that's, remember, we talked about the H3O or the hydronium ion for A. we've got the ammonium ion the reason i've had that a couple times is that that's a classic weak base type problem cn picks up a hydrogen becoming hydrogen cyanide that cn ion's really dangerous um it's a really strong poison and if you'll notice the h2po4 which is negative now picks up a hydrogen to become the h3 PO4. Notice there's now no charge.

So the charges all balance each other out too. Number 16, it says what is the conjugate acid of the following? What is the conjugate base? I've done the acids in red and the bases in blue. The acid picks up the hydrogen.

The base loses a hydrogen. Again, this whole section, acids lose hydrogens, bases pick up hydrogens. And then if it's the conjugate, we're looking on the right-hand side.

And finally, we have things that can be amphiprotic that can either gain a hydrogen or lose a hydrogen. The OH gains a hydrogen to look like H2O or loses a hydrogen to become O2-. The S2-can't be amphiprotic because it doesn't have a hydrogen to donate. And then we have the HSO4. Okay, notice it can gain a hydrogen and lose a hydrogen.

Transcript for:Understanding Acids and Bases

Transcript for:
Understanding Acids and Bases