Lesson Video: Lewis Acids and Bases Chemistry

In this video, we will learn about Lewis acids and bases and their properties, and identify them in chemical reactions.


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.

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