Lewis Acids and Bases - Online Tutor, Practice Problems & Exam Prep

The concept of Lewis acids and bases, introduced by Gilbert N. Lewis, diverges from the traditional definitions by focusing on electron pairs rather than hydrogen ions. A Lewis acid is characterized as an electron pair acceptor, which can be represented by positively charged ions like H⁺ or metal ions, and elements with less than eight valence electrons, such as those in groups 2A, 3A, and some transition metals. Examples include magnesium chloride and aluminum bromide. Conversely, a Lewis base is an electron pair donor, often indicated by the presence of a lone pair on a central element or a negative charge, as seen in hydroxide or cyanide ions. The product of a Lewis acid-base reaction is an adduct, where the Lewis base donates an electron pair to the Lewis acid, forming a new bond. This framework is broader than the Bronsted-Lowry and Arrhenius definitions, encompassing all their acids and bases while also including additional compounds that do not release or accept hydrogen ions. Lewis's model thus significantly expands the classification of substances as acids or bases.

Characteristics of Lewis Acids and Bases

Now we're going to say that the third acid base definition was introduced in the 1920s by Gilbert and Lewis, an American chemist. Now, this one is a little bit different from what we've accustomed to seeing in terms of Arrhenius and Bronsted-Lowry. So let's pay close attention when it comes to a Lewis acid. A Lewis acid by definition is an electron pair acceptor. We're no longer talking about releasing H ion or accept or donating H ions. That's not even part of the description of an acid. Now here we're talking about an electron pair acceptor.

Now what are the characteristics of a Lewis acid? Well, if you're going to accept electrons, electrons are negatively charged. Well, you could represent an H⁺ ion or a positively charged metal. Great examples. We have of course H⁺ ion and then all of these metals that are positively charged. Boron is an exception born a metalloid, so it has characteristics of both metals and nonmetals. So let's just say that it has a +3 charge. Now what else could fit under this idea of an electron pair acceptor? Well, we could say if you have a compound or molecular structure, if the central element has less than 8 valence electrons, it's not fulfilling its octet rule. So that opens it up to the possibility of accepting an electron pair.

Now what exactly what central elements could exactly fit under this description? Well, we're going to say elements that are found in groups 2A, 3A, and some transition metals. So for example, we have magnesium chloride. Because magnesium is a group 2A, it only has two valence electrons it connect to the chlorines. Yes, the chlorines are fulfilling their octet rule. But the magnesium itself came in with its two valence electrons. It's sharing two more, so it only has four total valence electrons around it. It's not fulfilling the octet role of having eight, so there is the possibility of it being able to accept an electron pair to get closer to that magic number of eight valence electrons.

The same thing can happen for a group 3A element. So here we have aluminum bromide, same thing. If you were to draw it out, it would have a total of 6 valence electrons around it, three that it came in with, and then three more that it's picking up by sharing with the bromines that are surrounding it. So again, it's not fulfilling the octet rule. Opens up the possibility of accepting an electron pair and then transition metals like nickel or zinc, cadmium. They because of their D orbitals, they can expand themselves. That opens up to the possibilities of having even more electron pairs.

So again, Lewis acid is an electron pair receptor. So if the acid is an acceptor, what is the Lewis base? Well, a Lewis base would simply be an electron pair donor. All right, so what can fit under this idea of an electron impaired donor? What we're going to say the presence of a lone pair on your central element within a structure. That would be a key giveaway that you have a Lewis base. Here we have this structure here, and we can say that the oxygen has lone pairs on it which it could share with another element, thereby donating the electron pair and serving as a Lewis base.

Same thing with the nitrogen here, that's part of ammonia. It has a lone pair that it could technically share. And in fact, this is what happens when ammonia accepts an H⁺, it's using its lone pairs to connect with that H and that's how we get NH₄⁺. But what else could serve as a Lewis base? Well, this is even easier, the presence of a negative charge. If you have a negative charge, you have an abundance of electrons. Remember, it's best if you're neutral. If I have too many electrons, I'm negatively charged. So I'm going to seek out a way of sharing some of those electrons with something that's electron deficient, something that needs a lone pair.

So all of these we have hydroxide here, we have the azide ion, cyanide ion and then we have HS^- ion. They all can basically share a lone pair with some other element. So again, Lewis acid is an electron pair acceptor. A Lewis base is an electron pair donor.

The Adduct

Now the adduct represents the product of a Lewis based acid reaction. Here in this example we have two reactants combining together to give us one product. We are taking those two reactants and adding them together. That's why we say adduct, we're adding them together. The first structure in blue here, it's called acetone. It is reacting with boron trichloride and together they add to create this adduct.

Yeah, Now remember that we set a Lewis acid is an electron pair receptor. We said that this can be represented by H⁺ positive metals or where the center element has less than 8 electrons around it. If we look at boron boron's and group 3A. So it has three valence electrons and it picks up three more by forming bonds with these chlorines. But in total it only has six valence electrons around it, which opens up the possibility of accepting 1 electron pair in order to fill the octet rule.

So here we're going to say that this represents the Lewis asset and if that's Lewis acid, then acetone represents the Lewis base, it is the electron pair donor and what it's donating is this lone pair. Now this is a bit different from Bronsted-lowry acid definitions because in Bronsted-lowry acid definitions, the acid donates in H⁺. It kind of just gives away the H⁺ and stays behind. Here, if you're donating your electron pair, you're coming along for the ride because you're electrons are a part of your core as an element.

So here, this oxygen is going to share this lone pair with the boron to help make this bond. Remember, a bond itself is just two electrons and look how the whole structure came along with it. So we're adding together the Lewis acid and the Lewis base to give us our adduct. Now this is important. All Bronsted-lowry acids and bases are also Lewis acids and bases, and that's because Lewis is the broadest of the three types of acids and bases. It encompasses all types of acids and bases.

Bronsted-lowry is a little bit more narrowed, focused and then Arrhenius even more so we're going to say. However, many Lewis acids and bases are not classified as acids or bases Under Bronsted-lowry or Arrhenius models This here is a Lewis acid, but under Bronsted-lowry and Arrhenius it wouldn't be an acid because of the lack of an H⁺ ion. So as you can see, Lewis definition really expands as far out as it can in terms of collecting as many different compounds under the acid and base banners.

All right, so just keep this in mind. Lewis is the most broad, followed by Bronsted-lowry, and then Arrhenius seems to be the most narrow in its description of what classifieds an acid, What classifieds a base.

Lewis Acids and Bases Example

Here it says to identify each of the following as a Lewis acid or base. Bronsted-lowry acid or base or both. So for the first one we have cobalt 2 ion. Here it is a positively charged metal ion, so we could easily accept an electron pair. Therefore represents a Lewis acid. Here it wouldn't represent a Bronsted-lowry acid because of the lack of an H plus group. So here we're going to say that this is a Lewis acid.

For B we have nitric acid. Now nitric acid has the presence of an H plus group, so it could serve as a Bronsted-lowry acid now, but could it serve as a Lewis acid? Well, it doesn't have the presence of an H plus immediately doesn't have a positive metal. It doesn't have a central element that has less than 8 electrons around it. So we wouldn't exactly say that it it's a Lewis acid. We're just going to say here it's a Bronsted-lowry acid.

For the next one, this is. If we were to draw it out, it'd be Chapter 3, which is connected to a carbon double bonded to an oxygen and then bonded 2 and H₂. So this is our structure. We're going to say The presence of these lone pairs here could mean that they can serve as electron pair donors. So this could serve as a Lewis base. Also, because of the presence of lone pairs, they could accept H plus ions, so that's also a possibility. So it could serve as a Bronsted-lowry base. So we're going to say it could serve both types of bases. So it's both.

Then finally we have CO₃^2-. The presence of a negative charge means that it has extra electrons that it could share, so it could act as an electron pair donor and therefore be a Lewis base. At the same time it is CO₃^2-. It could accept an H plus to become HCO₃^-, which is bicarbonate, so it could also serve as a Bronsted-lowry base. So it could be both types of bases.

Now this is how we describe each of the following compounds on whether they're a Lewis acid or base or Bronsted-lowry acid or base.

Identify the Lewis acid and Lewis base in the following reaction.

CaO (s) + CO₂ (g) → CaCO₃ (g)

• A. CO₂ Lewis Base, CaO Lewis Acid
• B. CaCO₃ Lewis Base, CO₂ Lewis Acid
• C. CO₂ Lewis Base, O²⁻ Lewis Acid
• D. O²⁻ Lewis Base, CO₂ Lewis Acid

Here’s what students ask on this topic:

Which option best describes the definition of Lewis acids and bases?

Lewis acids and bases are a concept introduced by Gilbert N. Lewis. A Lewis acid is a species that can accept a pair of electrons, whereas a Lewis base is a species that can donate a pair of electrons. This definition is broader than the traditional Arrhenius and Bronsted-Lowry definitions, as it doesn't require the acid to donate a proton (H⁺). Instead, it focuses on the ability to accept or donate a pair of electrons during a chemical reaction.

For example, in the reaction between ammonia (NH₃) and boron trifluoride (BF₃), ammonia acts as a Lewis base by donating its lone pair of electrons to boron trifluoride, which acts as a Lewis acid by accepting the electron pair. This results in the formation of a coordinate covalent bond, where both electrons in the bond come from the same atom, the Lewis base.

How can one identify Lewis acids and bases?

To identify Lewis acids and bases, you need to look at the way they interact with electron pairs. A Lewis acid is a species that can accept an electron pair, while a Lewis base is a species that can donate an electron pair.

Here's a straightforward way to recognize them:

• Lewis Acid: It often has a positive charge or is a neutral molecule with a central atom that has an incomplete octet, meaning it doesn't have a full set of electrons in its outer shell. This makes it eager to accept an electron pair to achieve stability. For example, ${\mathrm{BF}}_3$ is a Lewis acid because the boron atom only has six valence electrons and can accept a pair to complete its octet.
• Lewis Base: A Lewis base typically has a negative charge or contains at least one lone pair of electrons that it can donate. It's usually a molecule or ion with a central atom that has a full valent shell that can share electrons. For instance, ${\mathrm{NH}}_3$ (ammonia) is a Lewis base because the nitrogen atom has a lone pair of electrons that can be donated to an acid.

Remember, the interaction between a Lewis acid and a Lewis base results in the formation of a coordinate covalent bond where the electron pair is donated from the base to the acid.

How can Lewis acids be identified?

Lewis acids can be identified by their ability to accept a pair of electrons. In chemical terms, a Lewis acid is any substance, such as a cation or an electron-deficient molecule, that can accept an electron pair from a Lewis base, which is an electron pair donor. To identify a Lewis acid, look for the following characteristics:

1. Electron Deficiency: Atoms with an incomplete octet of electrons, such as boron in BF₃ or aluminum in AlCl₃, are typical Lewis acids because they can accept electron pairs to complete their valence shell.
2. Positive Charge: Cations, especially those with a high charge and small size like Fe³⁺ or H⁺, are often Lewis acids because they have a strong tendency to attract electrons to balance their positive charge.
3. Electrophilic Centers: Molecules with polarized bonds where the electron density is low, typically due to the presence of electronegative atoms, can act as Lewis acids. For example, the carbon in carbon dioxide (CO₂) has a partial positive charge and can accept an electron pair.
4. Vacant Orbitals: Atoms or molecules with vacant orbitals can accept electrons into these empty spaces. Transition metal ions often serve as Lewis acids because they have empty d-orbitals that can accommodate electron pairs.

By looking for these characteristics, you can identify Lewis acids.

Which is a characteristic of a Lewis base?

A Lewis base is a chemical species that donates an electron pair to form a covalent bond. This characteristic is central to the Lewis theory of acid-base reactions, where a Lewis base pairs with a Lewis acid, which is an electron pair acceptor. The ability to donate an electron pair can be attributed to the presence of lone pairs of electrons on the Lewis base. These lone pairs are often found on atoms like nitrogen, oxygen, or sulfur, which have available electrons that are not engaged in bonding. The Lewis base can share these electrons with a Lewis acid, thereby forming a new bond, which is called a coordinate covalent bond or dative bond. This electron pair donation capability is what defines a substance as a Lewis base in chemical reactions.

Which of the following is a Lewis base?

A Lewis base is a compound or ion that can donate a pair of electrons to form a covalent bond. In the context of Lewis acid-base theory, the base is the electron pair donor. To determine if a substance is a Lewis base, you should look for species with lone pairs of electrons that are not involved in bonding and can be donated to another atom, typically a Lewis acid, which is an electron pair acceptor.

For example, ammonia (NH₃) is a common Lewis base because it has a lone pair of electrons on the nitrogen atom that can be donated to form a bond with a Lewis acid. Other molecules like water (H₂O), hydroxide ion (OH^-), and chloride ion (Cl^-) can also act as Lewis bases due to the presence of lone pairs of electrons that can be shared. Organic molecules with lone pairs on nitrogen, oxygen, or sulfur atoms, such as amines or ethers, can also be Lewis bases.