Acid-Base Definitions

Let's talk about the acid-base
definitions for Bronsted-Lowry

and, also, Lewis.

And we'll start
with Bronsted-Lowry.

So, a Bronsted-Lowry
Acid is a proton donor,

and a Bronsted-Lowry Base
is a proton acceptor.

So let's, really
quickly, review what

this definition means by proton.

So if I look at this
diagram, right here,

I'm going to draw the hydrogen
atom, or the most common

isotope.

So hydrogen has one proton in
the nucleus and one electron,

somewhere around our nucleus.

So a negative charge, like that.

And so, we would say
this is hydrogen.

All right?

And then we put, it's one
valence electron, right there,

to represent the hydrogen atom,
or the most common isotope.

If we were to, somehow,
take away this electron,

we would only be left
with the proton here.

We'd only be left with
the proton in the nucleus.

And so, when we're
talking about a proton,

we're talking about the
nucleus of a hydrogen

atom, which is equal to H plus.

So, no longer are we
talking about the electron.

So let's see how this applies
to an acid-base reaction.

And so we start over
here with water.

And then we have HCl
over here on the right.

Now, in this bond, between the
H the Cl, one of those electrons

came from the hydrogen and one
of them came from the chlorine.

So let me just go ahead
and draw those in.

So the one from the hydrogen,
I'm going to put in blue here.

And that's this electron from
hydrogen, right here in blue.

And then for chlorine, I'm going
to make that electron green.

So right in here, like that.

And so for this
acid-base reaction,

a lone pair of
electrons in the oxygen

is going to take this proton.

So just the nucleus
of the hydrogen

atom leaving the
electron in blue behind.

And that electron
in blue stays behind

and ends up on the chlorine.

So let's go ahead and draw
what we would form from that.

We would have oxygen here.

The oxygen had two
bonds to hydrogen.

And the oxygen just picked
up another bond to hydrogen.

And so, let me go ahead
and mark those electrons.

So these electrons
in here, in magenta,

formed a new bond
with that proton.

So that's this bond right here.

And then we had some
electrons on oxygen.

Let me go ahead and
make those in red.

So these electrons in red on
the oxygen didn't do anything.

So they're still there.

So they're right here.

And that's going
to give that oxygen

a plus one, a formal charge.

And so this is the
hydronium ion, H3O plus.

Our other product,
we would also make--

we would have our
chlorine, which

had three lone pairs of
electrons around it already.

And then it picked up
both of those electrons.

Let me go ahead and mark them.

The one in green that
it had originally

brought to the dot structure.

And also, the one
in blue, the one

it took from hydrogen like that.

So chlorine now has
a negative charge.

So it's really the
chloride anion.

So this would be Cl
minus, like that.

So let's identify our
Bronsted-Lowry Acid

and our Bronsted-Lowry
Base for this reaction.

So let's go back over here
and see what happened.

So the H20, the water,
acted as a proton acceptor.

It accepted a proton from HCl.

So water would be our
Bronsted-Lowry Base.

And HCl donated a
proton to water.

So HCl would therefore be
our Bronsted-Lowry Acid.

So let's go ahead and identify
conjugate acid-base pairs here.

So if HCl is our
Bronsted-Lowry Acid,

I could think about its
conjugate base over here

would be the chloride anions.

So this would be the
conjugate base over here.

So H2O was our
Bronsted-Lowry Base,

and then over here, we can
find its conjugate acid, that's

H3O plus.

So this would be the
conjugate acid, over here.

So when you're looking for
conjugate acid-base pairs,

you're looking for
one proton difference.

So H2O and H3O plus are a
conjugate acid-base pair.

And HCl and Cl minus are a
conjugate acid-base pair.

And if we look at what we
have in the right here,

we are now saying H3O plus is
an acid, and Cl minus is a base.

And so, one thing
you'd think about

is H3O plus donating
a proton to Cl minus.

And so, we'll draw a little,
tiny arrow going back

to the left.

Because the equilibrium for this
reaction lies far to the right.

So we're going to get a lot more
of your products on the right

here.

But just thinking about
these definitions,

right, H3O plus would
be donating a proton,

and Cl minus would be
accepting a proton.

The chloride anion would
be accepting a proton.

But again, we know
HCl is a strong acid,

so we know the equilibrium
lies far to the right.

So that's the idea
about Bronsted-Lowry.

Let's look at
another definition,

which is actually a
little bit more broad.

So this is a Lewis
Acid and Lewis Base.

So a Lewis Acid is an
electron pair acceptor.

And so, an easy way to remember
this is, acid acceptor.

And a Lewis Base is an
electron pair donor.

And so, one way to remember
that this Lewis Base is

an electron pair donor is
to, if you think about this b

being lowercase.

And then just flipping it
around, you would get a d here.

So you get d.

So a base is a donor.

So let's look at
this reaction here.

And we have this cyclic
ether, over here on the left.

And then we have borine
over here on the right.

Now, notice there's no octet of
electrons around boron, right?

Boron is only surrounded
by six electrons here.

And that makes it very reactive.

Boron is SP2
hybridized, which means

it has an empty p orbital.

And so, let me go ahead and
represent the empty p orbital

like this.

It's able to accept
a pair of electrons.

And the ether over here is going
to donate a pair of electrons.

And so, let's go ahead
and show what happens.

The oxygen here is going to
donate a pair of electrons

into the empty orbital.

And there's going
to be a bond that

forms between the
oxygen and the boron.

So the ether over here is
donating a pair of electrons.

So that must be our Lewis Base.

And borine, over here, is
accepting a pair of electrons.

So that's our Lewis Acid.

Let's go ahead and
draw the product

for our Lewis acid-base
reaction here.

So we have our oxygen is
now bonded to the boron.

The boron is still bonded
to three hydrogens,

so we draw those
in there like that.

And let's follow
some of our electrons

here before we finish
drawing everything in.

So these electrons in
magenta formed this bond

between the oxygen
and the boron.

And then we also had some
other electrons on that oxygen.

Let me identify those.

So these electrons right here
in red are still on that oxygen.

So they are right
here on that oxygen.

That oxygen therefore, has
a plus one, a formal charge.

So plus one formal
charge on oxygen.

And boron gets a negative one
formal charge now like that.

And so, that's one Lewis
acid-base reaction here.

Now the Lewis
acid-base definition

is, once again, more
inclusive than Bronsted-Lowry.

If we actually go up here
to the previous reaction,

we can actually classify
these using the definition

for Lewis Acid and Lewis Base.

So let's look again at
what's happening here.

So water is donating
a pair of electrons.

Well, according to Lewis
Base, electron pair donor.

So we could say that water, we
could say this is a Lewis Base.

And HCl is accepting
a pair of electrons.

So electron pair
acceptor is Lewis Acid.

So we could call
this a Lewis Acid.

So notice, it doesn't matter
what definition you use.

If you use Bronsted-Lowry,
this is your acid.

If you use Lewis,
this is your acid.

Or if you use,
over here for base,

this is your base,
according to Bronsted-Lowry.

This is also a base
according to Lewis.

And Lewis Acid and Base also
have particular importance

in organic chemistry because
you can talk about the term

Lewis Acid as being
synonymous with electrophiles.

So you could say this
is an electrophile.

And then, you could say a Lewis
Base is an electron pair donor.

That's a nucleophile.

And nucleophile,
electrophile are

extremely important
concepts to understand

when you're talking
about organic chemistry.