Lewis Acids and Bases: Videos & Practice Problems

Lewis acids are defined as electron pair acceptors, including positively charged ions like hydronium and metals, as well as elements with fewer than eight electrons, such as beryllium and aluminum. Conversely, Lewis bases are electron pair donors, typically containing lone pairs or negative charges, like hydroxide and cyanide ions. When a Lewis acid reacts with a Lewis base, they form an adduct, where charge distribution changes, highlighting the importance of electron sharing in acid-base reactions.

Lewis Acids and Bases

Our final understanding of what constitutes an acid, what constitutes a base was covered by Lewis. Now according to the Gilbert Lewis definition, a Lewis acid is an electron pair acceptor. Now we're no longer talking about donating H⁺ or accepting H⁺, but in fact we're talking about how the use of electrons in the form of lone pairs can be used to abstract or remove acidic hydrogens from one compound to the next, or how those lone pairs can be used to attach themselves to another compound. Now here, what fits into the idea of a Lewis acid? Well, the first category for Lewis acids are positively charged hydrogen or metals.

We can have our hydronium or a hydrogen ion. This would constitute a Lewis acid. It's positive, so it'll be willing to accept negative electrons from a lone pair. We could also have a positive metal. These would also fall under the idea of being Lewis acids.

They're all positively charged so they would freely accept negative electrons from lone pairs. Next, if your central element within a given compound has less than 8 electrons. In this case, they're not following the octet rule, so they have room accept a lone pair from a Lewis base. Here, elements with less than 8 electrons are typically from groups 2 or 3. Elements from group 2 only have 2 valence electrons, so think of beryllium.

And let's say it was connected to 2 chlorines. Chlorine is in group 7a, so it has 7 valence electrons. Remember, when you make a covalent bond, you're going to be sharing electrons. Now although beryllium is a metal, because of its position on the periodic table, it does share some characteristics that are common with other nonmetals. So we'd say here that the beryllium chlorine bond kind of mimics a covalent bond.

Here, beryllium came with 2 lone electrons, and it picks up 2 more by sharing them with chlorine. Right now, it only has 4 electrons around it, so it has room to accept another lone pair to get closer to the octet rule of being close to having 8 electrons. Remember, elements do this in order to get electron configuration as close to a noble gas as possible. That would be an example of a group 2 element doing this. In group 3, we could have aluminum.

Aluminum is in the center. It's in group 3a, so it has 3 valence electrons. Let's say it connects to hydrogens. Hydrogen is in group 1a, so it only has one valence electron. Aluminum came with 3 valence electrons but it picks up 3 more by sharing them with hydrogen, so it has 6 around it.

It can still accept a lone pair to get closer to the octet rule of 8 electrons. These are typical examples of Lewis acids. Now another sub-category which we usually don't talk about because this isn't really our organic chemistry course, but it does fit the profile. So we can say here that sometimes if you have a central element with pi bonds. A great common example, you could have CO₂.

Here, that carbon is making pi bonds which means that it could accept a lone pair from an outside compound. By accepting that lone pair, we'd have the movement of one of these pi bonds to the oxygen. So this is typical of nonmetal oxides when I'm talking about a central element with pi bonds. So nonmetal oxides just means a nonmetal with oxygen.

CO₂ is a good example, SO₂, SO₃, these types of nonmetals, with oxygens involved. NO₃^- is that's a harder case, but that could also fit this characteristic. So these are the 3 general groupings for Lewis acids. Now that you've looked at Lewis acids, click on to the next video to look at Lewis bases.

Lewis Acids and Bases

So if a Lewis acid is an electron pair acceptor, then a Lewis base must be an electron pair donor. Now to be able to donate those lone pairs, you must have extra electrons. So we're going to say what fits into this definition of a Lewis base. Well, compounds with lone pairs on the central element. Here we have examples.

We have ammonia where nitrogen has a lone pair, water, methanol, the alcohol here, and then we have this ether, dimethyl ether. So, all of them have in common a central element with lone pairs which can freely donate them to a Lewis acid. Now what else would constitute a Lewis base? Well, even easier, compounds with a negative charge: CN^- cyanide ion, hydroxide ion, methoxide ion, and your azide ion.

Here, this thing has lone pairs which you don't see, which can donate. So these basically reference basic layouts for Lewis bases. Now here, if we're thinking of a Lewis base and a Lewis acid reacting with one another. So let's say we had for our Lewis acid, we had aluminum with fluorines. So remember, aluminum is in group 3a, so it only has 3 valence electrons.

It picks it gets to share 3 more from the fluorine, so it has 6. So it hasn't reached its octet number of 8 yet. Here, I'm gonna react it with let's say we react it with hydroxide ion. It's negative, so it's a Lewis base for sure. It has an abundance of lone pairs on the oxygen.

So what happens here is this oxygen can donate a lone pair over to the aluminum. Now with the Bronsted-Lowry definition, we donate H⁺. A compound can donate an H⁺ and be left behind as a conjugate base. But when we're sharing our lone pairs with another compound, we come along with them. These two would combine together.

We'd have the aluminum still with its 3 fluorines attached. And then the oxygen uses its lone pair to make a new bridge, a new bond with the aluminum in the center. It still has 2 of the lone pairs it's not using plus the hydrogen. Now oxygen was negative, but now it's sharing its lone pair with the aluminum. So now it's neutral.

Aluminum, on the other hand, was neutral, but now it has extra electrons it's gained from sharing the new bond with oxygen. So now it's negative. So it becomes important to remember charge distribution. What happens when I share electrons with a new compound? I become more positive.

What happens when I gain new electrons from another compound? I become more negative. Here we have our acid which accepts lone pairs. Our base which donates lone pairs; we added them together. That's why we call a Lewis acid-based product typically an adduct because we added them together.

These are the fundamental principles behind the Lewis acid-base model. Now that we've seen this, look to see if you can answer the question given below. Attempt it on your own, but if you're stuck, don't worry. Just come back and see how we approach that same question.

Lewis Acids and Bases

So this question isn't as simple as it looks. But remember, I've given us hints and I've talked about it up above to help guide us to the correct answer. So we have to identify who's the Lewis acid and who's the Lewis base. In this reaction, we have sulfur dioxide reacting with water to produce sulfurous acid. Well, sulfur dioxide fits one of the definitions we talked about above.

It is a nonmetal oxide. It's a nonmetal and sulfur connected to oxygen. And we said that those typically represent Lewis acids where the central element is double bonded. This is what sulfur dioxide would look like. It'd have more of a bent structure but here, let's not worry about that.

So here, it's a nonmetal oxide so it's the Lewis acid, therefore the electron pair acceptor. Oxygen and water, the lone pairs on the central element oxygen, can be donated. So this is acting as our Lewis base. Together, they combine to give us our adduct, sulfurous acid. Later on, when we explore more and more about acids and the different types beyond Bronsted, Lowry, and Lewis, and Arrhenius, those are just classifications.

Later on, we'll learn about the 2 major types of acids. And we'll learn that one of these types, coined oxyacids, are typically created by combining a nonmetal oxide like sulfur dioxide with water. An oxyacid itself is just a compound possessing H or more than 1 H connected to a polyatomic ion with oxygens, 1 or more oxygens. Because what is this really when you break it down into its components?

It's really just H, 2 of them, connected to sulfide ion. So we have our H ions connected to our sulfide polyatomic ion, which contains oxygens. So guys, just realize that when it comes to classifications of acids and bases, there are 3 major types. We have Arrhenius, we have Bronsted-Lowry, and we have Lewis.

And as we go by each one, the definition becomes more and more broad in our explanation and understanding of acids and bases and how they react and relate to one another. Continue onward as we explore more and more about calculations dealing with acids and bases.

Here’s what students ask on this topic:

What is a Lewis acid and how does it differ from a Bronsted-Lowry acid?

A Lewis acid is defined as an electron pair acceptor, while a Bronsted-Lowry acid is defined as a proton (H⁺) donor. The key difference lies in their definitions: Lewis acids focus on the acceptance of electron pairs, whereas Bronsted-Lowry acids focus on the donation of protons. For example, AlCl₃ is a Lewis acid because it can accept an electron pair, but it is not a Bronsted-Lowry acid because it does not donate a proton. Conversely, HCl is a Bronsted-Lowry acid because it donates a proton, but it is not a Lewis acid because it does not accept an electron pair.

What are some common examples of Lewis bases?

Common examples of Lewis bases include compounds with lone pairs of electrons or negative charges. Some typical examples are:

• Ammonia (NH₃): Nitrogen has a lone pair of electrons.
• Water (H₂O): Oxygen has two lone pairs of electrons.
• Hydroxide ion (OH⁻): Oxygen has three lone pairs of electrons and a negative charge.
• Cyanide ion (CN⁻): Carbon has a lone pair of electrons and a negative charge.
• Methanol (CH₃OH): Oxygen has two lone pairs of electrons.

These compounds can donate their lone pairs to Lewis acids, forming adducts.

How do Lewis acids and bases form adducts?

Lewis acids and bases form adducts through the donation and acceptance of electron pairs. A Lewis base donates a lone pair of electrons to a Lewis acid, which accepts the electron pair. For example, when aluminum chloride (AlCl₃) reacts with ammonia (NH₃), the nitrogen in ammonia donates a lone pair to the aluminum in AlCl₃, forming an adduct:

AlCl₃ + NH₃ → AlCl₃NH₃

In this reaction, AlCl₃ acts as the Lewis acid, and NH₃ acts as the Lewis base. The resulting adduct is a complex where the electron pair is shared between the two species.

What are the characteristics of a Lewis acid?

Lewis acids are characterized by their ability to accept electron pairs. Common characteristics include:

• Positively charged ions, such as H⁺ and metal cations (e.g., Fe³⁺).
• Elements with fewer than eight electrons in their valence shell, such as beryllium (Be) and aluminum (Al).
• Compounds with central atoms that have vacant orbitals, such as BF₃ and AlCl₃.
• Nonmetal oxides with central atoms that can accept electron pairs, such as CO₂ and SO₂.

These characteristics enable Lewis acids to accept electron pairs from Lewis bases, forming adducts.

Can you explain the concept of charge distribution in Lewis acid-base reactions?

In Lewis acid-base reactions, charge distribution changes as electron pairs are shared between the acid and base. When a Lewis base donates an electron pair to a Lewis acid, the base becomes more positive, and the acid becomes more negative. For example, when hydroxide ion (OH⁻) reacts with aluminum chloride (AlCl₃), the oxygen in OH⁻ donates a lone pair to the aluminum in AlCl₃:

AlCl₃ + OH⁻ → AlCl₃OH

In this reaction, OH⁻ (the Lewis base) becomes neutral, and AlCl₃ (the Lewis acid) becomes negatively charged. This change in charge distribution is crucial for understanding the formation of adducts in Lewis acid-base chemistry.