Video Transcript
In this video, we will learn about
Lewis acids and bases and their properties and identify them in chemical
reactions.
Chemists like chemicals, but it’s
very hard to learn about every chemical individually. There are just so many. Instead, chemists look for
patterns. We put chemicals in groups, and we
compare chemicals with one another. Two of these groups are acids and
bases. The simplest description of an acid
could be a substance that tastes sour, while the simplest description of a base
could be a substance that reacts with an acid. Over time, descriptions of acids
and bases have changed. Each description is named after a
different scientist or group of scientists.
In 1887, Svante Arrhenius described
acids and bases on the basis of whether they produce hydrogen ions or hydroxide ions
when added to water. When we add Arrhenius acids to
water, they disassociate forming hydrogen ions, while Arrhenius bases when added to
water disassociate to produce hydroxide ions and cations. Hydrogen ions and hydroxide ions
react to make water. This accounts for the fact that
generally Arrhenius bases will react with Arrhenius acids in a consistent way.
In 1923, Danish chemist Johannes
Nicolaus Brønsted and English chemist Thomas Martin Lowry extended the description
of bases, including substances that do not yield hydroxide ions in solution. A Brønsted–Lowry base is any
substance that can accept a proton. For instance, ammonia, NH3, can
react with hydrogen ions, H+, in water forming the ammonium ion. Generally speaking, Arrhenius acids
are also Brønsted–Lowry acids. Since they have a hydrogen ion,
they can donate. However, it’s important to realize
that Arrhenius acids are those chemicals that can donate hydrogen ions, specifically
to water. There are some substances that will
act as acids in some circumstances and as bases in others; these are known as
amphoteric substances. We’d have to make a much more
complicated diagram to account for all the variations. So, for now, I’ll leave this area
blank.
Instead, we’re going to focus on
the innovations introduced by Gilbert N. Lewis also in 1923, where he
expanded the reach of acid. A Lewis acid is any substance that
can accept a lone pair of electrons forming a bond, while a Lewis base is any
substance that has a lone pair of electrons that can be donated to form a bond. In general, Brønsted–Lowry bases
are always Lewis bases and vice versa. But not all Lewis acids are
Brønsted–Lowry acids, while all Brønsted–Lowry acids are Lewis acids. This is all getting rather
complicated, so let’s look at some examples.
Let’s start with hydrogen chloride,
HCL. HCL is an Arrhenius acid. It has hydrogen ions that it can
donate to water, but it’s not an Arrhenius base, since it doesn’t have any hydroxide
ions. All Arrhenius acids are also
Brønsted–Lowry acids, since they donate protons. But HCL is not a Brønsted–Lowry
base because it does not accept more protons. And HCL is also a Lewis acid, since
all Brønsted–Lowry acids are Lewis acids. We can see HCL can be attacked by
the lone pair of ammonia, a Lewis base. But except under extreme
circumstances, HCL will not behave as a Lewis base.
Now, let’s take a quick look at
BH3, borane. BH3 doesn’t readily donate its
protons, so we wouldn’t consider it an Arrhenius acid or a Brønsted–Lowry acid. But for reasons we’ll look at soon,
BH3 will accept a lone pair from a Lewis base. And BH3 doesn’t fit the
requirements for Arrhenius, Brønsted–Lowry, or Lewis bases. We know it can’t be a Lewis base
because it doesn’t have any lone pairs at all; again, we’ll come to that soon.
What about NaOH? It’s clearly an Arrhenius base, and
we know it’s a Brønsted–Lowry base because the hydroxide ion readily accepts
protons. And it’s also a Lewis base because
it accepts protons by donating a lone pair forming a bond. But it’s not acidic under any
category.
And next, we have ammonia, NH3. Under normal circumstances, ammonia
will not donate its protons. So, we don’t consider it Arrhenius
acidic or Brønsted–Lowry acidic. It’s also not Lewis acidic. We don’t see circumstances where
ammonia accepts a lone pair. Ammonia also doesn’t fit the
definition of an Arrhenius base. It doesn’t have hydroxide ions to
donate. However, ammonia does react with
water producing hydroxide ions in solution, and so acts like an Arrhenius base. However, ammonia is definitely a
Brønsted–Lowry base since it accepts protons. And since it does this using the
lone pair to form the bond, it’s also a Lewis base.
The advantage of Lewis’s
description was this. Lewis acids and Lewis bases react
in predictable ways. Since the set of Lewis acids and
bases is bigger than the others, Lewis’s description allows us to easily compare a
greater number of chemical reactions. There’s nothing wrong with a
Brønsted–Lowry description; it’s just more specific than Lewis’s. If we see an area of one molecule
that’s electron deficient and another area that’s electron rich with a reactive lone
pair, we can guess how they might interact, meaning that we can predict where bonds
are likely to form. In practice, we need to know a
little bit more than this, but it’s a great starting point.
When a Lewis acid reacts with a
Lewis base, they form what’s called a Lewis acid–base adduct. Here’s an example of a Lewis base,
ammonia. Ammonia has a reactive lone
pair. And here’s a Lewis acid, BH3. At the center of the BH3 molecule
is a boron atom that can accept electrons. So, we predict that the nitrogen
will form a bond with boron. And the molecule we get is called
the Lewis acid–base adduct. But that leaves one problem. How do we know where lone pairs are
and are not?
One of the neat things about Lewis
acid and base theory is the strong relationship it has with Lewis structures,
otherwise known as electron dot diagrams. Lewis structures allow us to
highlight lone pairs as well as electron-deficient areas, for instance, when we
don’t have a full octet. Let’s have a look at the previous
example, the reaction between ammonia and borane.
On the ammonia side, we have three
hydrogen atoms, contributing one valence electron each, and a nitrogen atom, which
contributes five. That’s eight electrons in
total. That’s enough for a single bond
between the nitrogen and each hydrogen and two left over for the lone pair of
nitrogen completing its octet. Now, what about BH3? Each hydrogen contributes one
valence electron and a boron atom contributes three. That’s six electrons in total. Six is just enough to form single
bonds between boron and each of the hydrogen atoms. The boron atom could be more stable
with one extra bond, but it doesn’t have the electrons to do it. Instead, we need both electrons to
be provided by a Lewis base. The nitrogen donates both electrons
in its lone pair into the nitrogen–boron bond.
However, not all lone pairs are
reactive. The reasons for this are a bit too
complex for this video, so here’s a list. These are the main ones. We tend to see more reactive lone
pairs on more electronegative elements. And the electron-deficient species
that will act as Lewis acids include boron and aluminum compounds and various
transition metal ions. You don’t need to remember all
these examples. You’ll just need to remember how to
draw a Lewis structure, find a lone pair, and find an incomplete octet. Now, let’s have a look at some
examples in reactions.
BF3 can react with F− ions, forming
BF4−. If we draw the Lewis structures, we
can easily see where the lone pair that forms the bond comes from. The fluoride ion with its reactive
lone pair forms a bond with the electron-deficient boron. And here’s an example where we have
a transition metal ion, silver plus, reacting with a Lewis base, ammonia. With positive metal ions, we tend
to leave out any valence electrons, particularly with transition metals where it
becomes much more complicated. The lone pairs on the nitrogens of
the ammonia molecules form bonds with the silver. This particular chemical is active
in the Tollens reagent test for aldehydes, where we produce a silver mirror. And Lewis acids and bases are also
very important in organic synthesis.
Now the last area we’re going to
look at are Lewis amphoteric substances, substances that are both Lewis acids under
some circumstances and Lewis bases under others. When hydrogen chloride dissolves in
water, a water molecule reacts with the hydrogen ion in HCL forming the hydronium
ion and the chloride ion. We often treat the hydronium ion,
H3O+, like H+. But actually, whenever we see H+,
we’re really dealing with a Lewis acid–base adduct. One of the lone pairs on the oxygen
in water donates to form a bond with H+. In this example, water is acting as
a Lewis base, donating a lone pair.
But let’s have a look at this
example, the formation of ammonium hydroxide when we add ammonia to water. In this case, the lone pair that’s
going to steal an H+ ion is on the nitrogen. What we form is the ammonium ion
NH4+ and the hydroxide ion OH−. In this example, it’s water that’s
behaving as the Lewis acid, giving up an H+ ion to the ammonia. So, water is amphoteric. It can be a Lewis acid or, in other
circumstances, a Lewis base. Now, let’s take a look at a
question.
Which of the following statements
best defines a Lewis acid? (A) A substance that can donate a
pair of electrons, (B) a substance that can accept a pair of electrons, (C) a
substance that can donate an H+ ion, (D) a substance that can accept an H+ ion, or
(E) a substance that produces OH− ions.
The first thing that’s important to
identify is that we’re not looking for a simply correct answer. We’re looking for the best
statement out of the five. Lewis described acidity and
basicity in terms of the acceptance or donation of lone pairs of electrons. He describes substances like
ammonia with reactive lone pairs as bases and electron-deficient substances, like
BH3, as acids. When the two react, a bond is
formed between the area with the lone pair and the electron-deficient area.
The first statement suggests that a
Lewis acid is a substance that can donate a pair of electrons. This corresponds to the description
of a Lewis base, not a Lewis acid. The second statement suggests that
a Lewis acid accepts a pair of electrons. This fits nicely with what we’ve
already seen. So, let’s hold on to this answer
and check the other three.
Our third option is that a Lewis
acid is a substance that can donate a hydrogen ion. These would be substances like
hydrochloric acid and nitric acid. Here’s the reaction of a common
base sodium hydroxide with hydrochloric acid, forming sodium chloride and water. This reaction does conform to the
description of Lewis acids and bases because a lone pair from the hydroxide ion is
attacking the hydrogen ion.
Statement (C) is not our correct
answer because even though a substance that can donate a hydrogen ion is a Lewis
acid, not all Lewis acids donate hydrogen ions. So, this statement is not the best
description. Instead, this type of acid is
commonly referred to as a Brønsted–Lowry acid. The next statement changes the game
by talking about accepting hydrogen ions. This is a good definition of a
Brønsted–Lowry base and not a Lewis acid.
Finally, the last statement
suggests that a Lewis acid is a substance that produces OH− ions, hydroxide
ions. A good example of this is the
addition of solid sodium hydroxide to water, producing hydroxide ions in
solution. Sodium hydroxide is an Arrhenius
base. It dissociates in water to produce
hydroxide ions. The general term that covers
statement (E) is alkali because there are some substances that are not Arrhenius
bases that still react with water to produce hydroxide ions. Either way, this is definitely not
an acid and definitely not a Lewis acid, meaning of the five statements we’ve been
given, the one that best defines a Lewis acid is a substance that can accept a pair
of electrons.
Let’s have a look at the key
points. Arrhenius, Brønsted and Lowry, and
Lewis described acids and bases slightly differently. Lewis’s description covers the
biggest range of substances. A Lewis acid is simply any
substance that can accept a lone pair of electrons, while a Lewis base is the
reverse, a substance that can donate a pair of electrons. Lewis acids accept lone pairs,
while Lewis bases donate lone pairs. You can remember this by taking a b
and flipping it and making it look like a d.
Lewis acids and bases react to form
Lewis acid–base adducts, for short Lewis adducts. And finally, we can use Lewis
structures to work out what’s a Lewis base and what’s a Lewis acid by highlighting
lone pairs or areas which are electron deficient.
