Alright, we’re going to start this first Camtasia of Chapter 15, by talking about definitions of acids and bases. Acids and bases are
things that have been known about for hundreds, if not thousands, of years. Um,
there were lots of different compounds that people discovered over time that
they, ah, put in these different classifications. And originally acids
and bases were just defined based on some of the basic properties that they had. For instance,
acids were things that would taste sour, OK? Um, they didn’t know what it was that made it taste sour, but they knew that it tasted sour.
Um, they would react with certain metals, not all metals, but would react with SOME metals to make hydrogen gas,
which you have seen before. Alright. Um,
one of the things about, ah, acids, or bases, is that they would react with certain compounds and change their color. And one of the
first things that was really used as a test was litmus, it’s from a fungus, um,
but what it would do, an acid would turn litmus to a red color. OK, so
these were different things that they noticed that all these compounds had in common. Um, bases
would be things that would taste bitter, OK, they didn’t have a particular reaction with metals that they did, um, but they would
feel slippery or soapy, turns out that’s a reaction with your skin, um, but it would, feel slippery or soapy, um, when you touched it.
And they also would change the color of litmus, they would turn litmus, um, into a blue color. OK, so over the years,
there were MANY, MANY, different compounds, that when they were discovered or tested, um, it would taste sour. OK, and so when they,
well, they would add it to some of the metal, and …HUH, it might make a
little bit of hydrogen gas. But it would definitely change the color of the litmus.
If there’s something that they would, you know, feel that it was kind of soapy or slippery they’d taste it and it tastes bitter, they’d test against
litmus and it turned litmus blue, and so lots and lots of different compounds that they were able to separate in to
acids or bases based on these properties, but over the years as they started to
look at, you know, what was in these different compounds, there didn’t
seem to be anything really…that stood out. That made something acidic or basic. Um, the one other
thing about the acids and bases is that if you combine them, um, they would combine to make salt water. Uh,
so they would make a salt, and water if it was an aqueous solution, um,
and when you had the salt you really didn’t have your acidic or basic properties anymore; it was no longer
sour, it was no longer bitter, it didn’t really feel soapy any more, it didn’t react with metals. Um, and so
essentially, this was neutralization. OK? They essentially cancelled each other out and we got this salt water, OK? So, they were known about
for … MANY, MANY years, but they didn’t really know WHY they behaved that way, and there were lots and lots of different reasons
that were thrown out, um, some stuck around more than others. OK?
One of the first ones that really made...sense, in hind sight, um, was proposed
by a guy by the name of Arrhenius, um, this is the same guy who had the Arrhenius Equation that you learned for the
kinetics chapter. Um, and Arrhenius, what he was doing was he was working on his PhD thesis in the 1880’s.
And, (sigh) he was looking at acids and bases, and one of the things that … again, he was trying to
figure out what was going on. Um, and so to figure out what was going on, um, he actually was looking at…
conductivity of solutions. OK? Annnd, it’s been a while since we’ve done conductivity, so we are going to
look at a couple of videos real quick to remind ourselves of conductivity of solutions. Alriight.
So... "Well, we’ve gotta demonstrate that, I think is the best thing to do. So, what I have here,
is ah, just some water, I’ll put that there, and more water. Then I have this awesome testing machine, which is, ah, lethal in the wrong hands,
I’ll plug this in here, and there is a light bulb on the bottom as you can see, which will be
on the top in a minute, and there’s a couple of electrodes sticking out the top. But if something connects the electrodes the light bulb
lights, OK? So obviously a piece of metal conducts electricity very well, the electrons fly through it and the
light bulb lights up. Now, how about water? Is that an electrical conductor? Oh, …
some people,….no, it’s not distilled (inaudible) with that, aaaaah, some people say yes, some people
say no. Whadya think? Welll, not enough to light up the light bulb.
But if I was having a bath in there, and somebody threw this in I’d be dead.OK? There is enough
electrical current passing through there to zap a person but not enough to light up a 20-watt light bulb or whatever this is, so, not a whole lotta
current travels through. In a light bulb like this, you might have an ampere of current, to light it up nice and bright. In the bath tub it
only takes milliamps to go through your body to kill you. OK, so, distilled water or
just plain tap water, not much conductivity. Now, what is it that carries charge?
In the case of, of, this thing it’s electrons being transferred through the
spatula, and the electrons are mobile in the metal so they are what carry charge.
But what if we have something in solution-like some sodium chloride, common table salt, safety sealed.
So let’s put some salt in here. Stir it up a little bit, now this is dissolving, and my claim is that this is turning into ions, but, ah, the only way we could
know that there are ions in there would be to show that the solution has some electrical conductivity. And you can see that, of course it,
does. So, what’s happening here? It’s the ions now in the solution that are carrying the charge. Positive and negative charges; positive
sodium ions, negative chloride ions, those are the things that are carrying the charge from one electrode to the other and completing the
circuit. In ordinary water there are not sufficient number of ions – there are some ions but not a sufficient number to, ah, to
see that, ah, enough current would be, um, carried to light up this thing. Light up this bulb. Now, does it mean that when we dissolve
something it necessarily conducts electricity? Well, here’s some sugar. So I’ll throw some sugar in,
stir that around. You know that sugar is fairly soluble in water. I put about the same amount of sugar in there
that I had, ah, sodium chloride in the other one. So let’s see what happens. Zippo! So just
because something dissolves in water does not mean that it has separated into ions. The sodium chloride does,
the sugar doesn’t but yet it’s still soluble so there’s two different things going on in there. This part of the way
tells how you can really kill someone, you throw sodium chloride in the water first and then throw the toaster in bath tub.
One of my favorite shows is called, “Mythbusters,” I just, I laugh my head off when they do
things, and they, they demonstrated this, and they threw sodium chloride in
and threw their dummy in and yes, he got electrocuted more than if there was no sodium chloride in. Somebody have a question?
(student in background) Well, your body has salt on it, that’s true just not enough. If you put more in you get more conductivity and …
the guy’s even deader.” Alright. So that pretty much just demonstrates the idea of conductivity.
And as he said when we have ions present in solution, then, um, the
light bulb lights up. When there weren’t ions present in solution, like with the, um,
the, ah, sugar, then there was no conductivity. Now, he just
did salt and sugar, we’re gonna look at a couple of other solutions, I think, if it goes, there we go. There we go. Um, so
this has a variety of different things. Same type of idea, he has a probe
(TAP WATER), little bit of lighting up, you can actually see that one, you couldn’t see it on the other one.
(DISTILLED WATER, huh, what? SALT WATER, one teaspoon per cup.) Nice and bright.
(HYDROCHLORIC ACID), HCl here, nice and bright.
(SODIUM HYDROXIDE, just as concentrated) Nice and bright. (SUGAR WATER), here’s our sugar again.
(SWEET), Vinegar (VINEGAR, a weak acid), lights up, not quite as much as the HCl did before. (ETHANOL, a dissociative
that doesn’t associate – don’t try this at home, BARIUM SULFATE, it’s insoluble. Nooo conductivity here!) Alright,
dorky looking guy. Alright, um, so, the point there, though was that um, some solutions conduct, some don’t. And, what was
known about a lot of the acids at the time was that they had formulas of molecular compounds.
So we have things like, ah, HCl, where it’s known that hydrogen bonds
to chlorine. Um, and that we have things like acidic acid – CH3COOH, ok? So, a lot of the
acids were known to be, um, molecular compounds. OK?
And it was assumed that if you have molecular compounds, something like sugar, um, that when you put it in,
water, it is not going to split up into ions, ‘cause it is a molecule. If you have salt, something
like Sodium Chloride, and you put it into sugar, OK? Um, you’re going to get your individual ions. OK? Because it
is ionic. Now, this was the excepted reasoning. What Arrhenius said, though, was he said, “You know what? When I put acids in water, whether
it’s a strong acid or a weak acid, I get things lighting up.” And so, what he said was that even though acids are molecular compounds,
they can ionize. OK? Um, helps if I spell it right. They can ionize. They can create ions in solution. And so what that
means is that something has to come off, and what he figured out was,
well, all of them have some hydrogen floating around that is able to fall off.
And so Arrhenius said that acids are things which increase the concentration of H plus in solution, bases are
things which increase the concentration of hydroxides in solution. OK? Now, this was a very bold statement
back in the 1880’s, when he made it. Um, his PhD committee, um, that you present your research to, that say “yay or nay,”
um, did not agree with it, he almost failed. If he had we probably
wouldn’t have had our Arrhenius Equation. (laughs so inaudible word) named after him.
It was only through some political maneuvering that he was able to pass, um, and it turns out he was right. OK, acids are things that
increase the concentration of H plus, bases are things that increase the concentration of hydroxide, um, but this was a very radical
idea at the time, Ok? Now one of the things we talk about acids and bases, um, lot of times we are dealing with water.
So, when we have an H plus in water, OK, we have a bare proton, floating around in water, and water has all
those lone pair electrons on the oxygen, OK? And so what is going to happen is that, in water, they are
going to make, um, a coordinate covalent bond, and they are going to make
H3O plus, which is our hydronium ion, OK? So technically, when we are dealing
with acids in aqueous solution-which is where we see them most of the time-um, what we get is not
an increase in the concentration of H plus, what we get is an increase in the concentration of hydronium ion. OK? Um,
but, ah, just some terminology to be aware of, OK, because sometimes we’ll write our acids as H plus, sometimes we’ll write them as
H3O plus. Sometimes we will call them a proton, because H plus is just a bare proton floating around, or a hydrogen
ion, not as often but we could, sometimes we’ll call it hydronium.
So the terms are used interchangeably, sometimes you’ll see H plus people call hydronium,
sometimes you’ll see H3O plus people call it protons, they use the terms and, um, symbols interchangeably. OK. But that is
our acid; we increase the H plus in solution, our base increase is hydroxide, according to Arrhenius. OK? This is a
really, really great definition of acid-base. It’s what we usually
use to start out when we’re teaching it, um, but there are a couple of problems.
OK? First of all, you have to be in water. If you’re not in water, it doesn’t really work very well. OK. The second one is you have to
have both H plus and hydroxide, when you’re not in water you’ve often don’t have hydroxide, OK?
Um, and so there are some things that we … um, describe as acids and bases which um, …may or may not work, OK, using Arrhenius’s
definition. So to illustrate that, quickie little video here, um, what we have is a flask that
has ammonium, and then, ah, a Q-tip that has some HCl, OK? Now,
HCl we know is an acid, ammonia we know is a base, OK? Um, we kind of put the two right by each other and we are getting ammonium
chloride, which is a salt, so we’re even doing a neutralization. Um, but if we think about the reaction between, um,
HCl gas and NH3 gas, what we wind up getting is NH4Cl solid.
No H plus floating around in solution to increase its concentration, definitely no
hydroxide floating around in solution to increase its concentration. So even though we have something
that we know is an acid and something that we know is a base, and they’re coming together to make salt, a neutralization like we said acids and
bases do, um, according to the Arrhenius definition, um, we can’t classify this as an acid-base reaction. OK? So there had to be
other ways of describing it. Um, the next best way of describing it was proposed in 1923 by a pair of guys, one named
Bronsted, the other named Lowry; um, they were not working together, they published their definitions
within a couple months of each other, and so this has just become the Bronsted-Lowry definition. And so what these two guys did,
was they said, “Well, you know what? We don’t want to use both H plus AND hydroxide, um, so we’re just going to focus
on the H plus.” OK? And so what they did was they looked at the reaction, if I have an acid and a base, OK, they come together
and make water. And so, they said, “OK, we’re gonna say that our H pluses are acid, what is the base doing
in relation to the H plus?” And they figured out that what it’s doing is it is picking up an H plus, OK? And so they said
that an acid is something that donates a proton and base is therefore something that accepts the proton. OK. Um,
all of the previous definitions of acid and base that Arrhenius gave still worked.
OK, so if you think about it, ah, if we have HCl, OK, and we put it in water,
OK, HCl we said is an acid. OK? It’s going to make hydronium ion and chloride ion, OK? And so HCl still
functions as an acid. It is donating a proton, what it’s donating it to here in this case is water, and so water is going to function as a base.
Um, something like ammonia, OK? Um, when we put it in water, we get ammonium ion and hydroxide, OK?
It’s the hydroxide here that, um, allowed Arrhenius to call it a
base-ammonia base. Um, in this case though, for the Bronsted-Lowry, what is
happening is it is picking up a proton from the water. OK? That’s what gives us our NH4 plus and what we have left over is
the hydroxide, OK, so things that were acids according to the Arrhenius
definition are still acids, things that were bases according to the Arrhenius definition,
um, are still bases, the one thing that is different now though, is that I can do something like my gas phase reaction, OK? I can have
HCl and NH3, OK? And they can still function as an acid and a base. So the HCl donates, so it is an acid, um, the NH3 accepts, so it is a
base, OK? And I make my NH4 plus ion and my chloride ion, which come together
to make the ammonium chloride salt. OK? So. It was a much more, ah,
generic definition. Um, it applied to things that didn’t have hydroxides floating around in solution, and it also
got rid of, um, the need for water. So we have gas phase
things now, this definition works really, really good in organic solvents where we don’t have any
water present. And so, it’s a much more generic definition of acid-base than the Arrhenius was. OK? We have
couple of results from this, though. Um, one of the things is we have some species that can act as an acid or base. Here with the water,
um, when the water was combined with the HCl it functioned as a base. When the
water was combined with the ammonia it functioned as an acid. OK?
That is what is known as an amphoteric species. OK? It is something that
can act as an acid or a base, depending on what is put in it. OK?
So the best example of that is our water, OK? Other things are polyprotic acids, that have lost
at least one but not all, of their protons, OK? So for example, if I have
bicarbonate, HCO3 minus, OK? It lost one, but it hasn’t lost both, OK? So what it can do
is it can pick up that proton, so it can accept a proton to make H2CO3. If it accepts a proton um, then it is acting
as a base. OK? But it also has this H plus here, that it can get rid of – so it can donate a proton and act as an acid to make CO3 with a negative
2 charge. OK. So. Amphoteric species. Water is one of the great examples, but also these polyprotic acids that
have lost AT LEAST ONE, so it can get it back and act as a base, BUT NOT ALL, so it can lose more and act as an acid
um, of the protons. OK? That is one result. Another result of, um, our, ah, Bronsted-Lowry definition, is that
we have are called conjugate acid-base pairs. OK? Um, one of the
things that maybe you noticed when we were looking at the definitions, um, I have an acid and a base; OK? I have an acid and a base, alright?
Um, every time that we have a reaction, we always have a reaction between an acid and an a base, ‘cause something has to donate,
something has to accept. OK? The other thing, though, about our acid-base reactions we figured out in the last chapter, a lot of
our reactions are technically reversible, OK? Um, sometimes the reverse reaction isn’t very favorable, but our reactions
are reversible. OK? And so, um,
we start out with an acid-base, and if we’re gonna reverse our reaction, then the proton is going
BACK, and so our products are also an acid and a base. Alright? So, if we
look at one of these reactions, OK, if we look at, say the, um, ammonia reacting with the water, OK? Um, we said that we
made ammonium ion and we made hydroxide, OK? Um, actually I’m gonna do this on the next page. Alright, so, um, our
ammonia and our water going to make ammonium ion and hydroxide. So we said, in the forward reaction, what happened was that
the water donated a proton to the, um, ammonia, and so the water donates it acts as an acid; the ammonia accepts
so it acts as a base. OK? However, if I’m going to do the reverse reaction, OK, I have to
give that proton back, so now NH4 plus is donating a proton, so IT acts as an acid;
the hydroxide accepts, so it acts as a base, alright? So, on the
reactant side we have an acid and a base, on the products side we have an acid and a base.
OK? If I pair up my base on my reactive side with my acid on the product side, um, I have NH3 and NH4
plus, OK? One is a base, one is an acid, the only difference between the two is an H plus. OK? If I look at the other pair, I have a water
that’s an acid, I have a hydroxide that is a base. OK? Again, one is a reactant, one is a product, one is an acid, one is a base, um, and again if I
look at the two, the only difference is that there’s just that one H plus that’s going back and forth, OK? And so these are my conjugate
acid-base pairs that I get from the Bronsted- Lowry definition: one of them is the NH3
-NH4 plus, the other pair is the water and the hydroxide. And so one of the things that we get from the Bronsted-Lowry definition
is we get these conjugate acid-base pairs in all of our reactions. OK? So, as example, gonna put a couple of these in here,
what I want you to do is identify what are the conjugate acid-base pairs in these two reactions. So which one is the acid,
which one is the base? (Sound effect) oops, if I write it right…HCO3 minus…which one is the acid, which one is the base, on
each side, and what is the acid-base pairs? OK? So, take
a couple minutes here, and do that, and when you’re done, start it back up and we’ll go over it. (Sigh) Alright.
So, the first one: carbonate and water make bicarbonate and hydroxide, OK?
It helps to pick one of the compounds, um, and then look and see how it changes on the other side.
So if I start out with CO3 here, OK, I have CO3; the other side the thing that has carbon
is HCO3. So what I do, this starts out without hydrogen, it gets hydrogen, so that means we must be accepting,
‘cause it picked something up, so that means it acts as a base and my water acts as an acid because it is giving it up. OK? On the other
side, um, what I have, ah, here I have the hydrogen, I need to get rid of it to
go back to the other side. So I need to give this one up. So I’m donating, and so this is my acid;
hydroxide, maybe obviously, is going to be my base, OK? Um, so one pair
is carbonate and bicarbonate, the other pair is again water and hydroxide, OK? Alright, second reaction.
I have acetate and nitrous acid. Um, going to make acetic
acid and nitrite. OK? So again, I’m gonna look and see which things are similar; I have C2H3O2 HC2H3O2, so that is going to be one of my pairs.
Don't know which one is the acid or the base yet, but they only
differ by a proton, so I know that that is one of my pairs, and then the other one that
has nitrogen is going to be my other pair. OK, so those are my pairs, which one is the acid, which one is the base? This one starts out without a
hydrogen, winds up with a hydrogen, so this one is accepting so it’s my base. My nitrous acid is my acid, it is donating.
On the reverse side, um, that needs to lose this hydrogen, so this is
the acid, the nitrite is going to be my base because it’s going to pick up that hydrogen.
OK? So this is my Bronsted-Lowry definition of an acid and a base. OK? And again, more generic, it covers a LOT more situations, OK.
I have one more definition of an acid and a base, OK, and that is the Lewis acid and base definition. OK?
Now, this one is going to be the most generic. OK? And it rises in situations where we need to get rid of the hydrogen as well.
Alright, (sigh) an example of this. Um, if I have a metal oxide, OK, sodium oxide;
metal oxides are known to be basic. OK. Um, non-metal oxides, like SO3 here, are acidic, OK.
where we get things like our acid rain, we have nitrous oxides and sulfur oxides, um, these are acidic oxides, OK.
So I have something that I know is a basic oxide, if I put it in water I get a base. Something that I know is an acidic oxide, if I
put it in water I get an acid. If I take these two and add them together, what I get is, I get
sodium sulfate, which is neutral. OK? So again, take a base and an acid and I’ve combined them together to get a neutral salt.
This is your classic acid-base reaction, but I don’t have any hydrogens. OK? Um,
so I can’t use the Arrhenius Equation, I can’t use the Bronsted-Lowry definition, how am I going to define this as an acid or a base? OK.
This is where Lewis came in. OK? Lewis. You have learned about Lewis when we talked about Lewis structures back in 1061.
Gilbert N. Lewis is the dude who came up with Lewis structures, and so, when we’re doing Lewis structures
we were looking at valence electrons, and where those valence electrons were, OK? And
so this is what he did with the acid-base reaction, was he looked at electrons. OK? So, again,
looking at our generic acid, OK, um, and our generic base, OK. Here we’re going to actually draw the Lewis structure of a hydroxide, OK.
Um, but what is going on in terms of our electrons when we are doing an acid-base reaction? What Lewis figured out was that the electrons
were going from the hydroxide to the H plus to create this coordinate covalent bond.
OK, and so what he said was that an acid is something that accepts electrons, OK, a base is therefore going to be something that donates
electrons. OK. So an acid either has to have or can make – AH – empty orbitals in order to pick up those electrons.
A base has to have lone pair electrons in order to donate them. OK? And so this now becomes our
most GENERIC definition, OK? An acid is something that can accept electrons, a base is something that can donate electrons.
So with our Sodium Oxide and our Sulfur Trioxide, OK. What we have, if we get rid of the sodium because that’s just floating around,
is we have oxide with a negative two charge, um, that is our base, OK. And we’re combining it with Sulfur Trioxide. OK.
Right now the Sulfur Trioxide does not have, um, really an empty orbital, but what it can do is,
it can pick up those pair of electrons and when it does that, um, it can rearrange – OK, it can dump ‘em out, sulfur can have an expanded
octet. Um, and what we get is we get the sulfate ion, OK? Um,
so. This is…the sulfur functions as our acid
because it can accept the lone pair electrons, um, the oxide is our base. OK? Again,
this is a real GENERIC definition, the most generic definition that we have acids and bases.
Um, we don’t use it most of the time, a lot of times we do have our H pluses floating around, and it’s easier to think about
the H pluses and where they’re going. Um, the Lewis definition is often used, um,
with complex ions. OK. Complex ions are metal ions that have things stuck on ‘em, like, water,
ammonia, cyanide, things like that, OK? Um, so, for example, ah, if we had silver ion combining with NH3
we can make, um, AgNH3 two plus, OK. Or I could take something like Boron Hydroxide and
add water to it, and in doing that get Boron with four Hydroxides and H plus. Alright?
So, um, in both of these cases, using the Lewis definition, what is going to be our acid? What is going to be our base?
Alright? So, what is going to be our acid, what is going to be our base? We need to think about electrons. OK. Silver ion? Probably doesn’t have
a whole lot of electrons, OK? Ammonia, though, NH3, there are lone pair electrons on that nitrogen,
OK, so those lone pair electrons can attack, so… our NH3 DONATES electrons, so it is a base;
silver ACCEPTS electrons, and so it is going to be an acid. OK? With my boron hydroxide, if you think about it, boron we said before,
likes to have three things stuck to it and it has this empty p orbital on it. My
water has two lone pair electrons that can go fill that empty p orbital, to create the fourth
bond. So, my boron ACCEPTS lone pair electrons so it is going to be an acid, the water DONATES so it is going to be a base. OK,
and after it donates one of the hydrogens just falls off and that’s our extra hydrogen floating around here. Alright.
So those are our three definitions of acids and bases. Um,
the Arrhenius definition, an acid is something that increases the concentration of H plus
in water; a base increases the hydroxide concentration. Um, works really well in water, but pretty much just in water.
Um, the Arrhenius definition we have, um, our acids that donate protons, our bases that accept protons and for
those we have our conjugate acid-base pairs that we have to worry about, and our amphiprotic species; their most generic definition is the
Lewis acid and base definition that looks at the electrons. Um, acids accept lone pair electrons, bases donate lone pair electrons.
Alright? And so those are our definitions of acids and bases.