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.