Lewis Acids and Bases: Videos & Practice Problems

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.

Lewis Acids and Bases

The Lewis acid-base theory, introduced by Gilbert N. Lewis in the 1920s, expands the understanding of acids and bases beyond the traditional definitions provided by Arrhenius and Brønsted-Lowry. In this framework, a Lewis acid is defined as an electron pair acceptor, which shifts the focus from hydrogen ions (H⁺) to the ability to accept electrons.

Characteristics of Lewis acids include the presence of positively charged ions, such as H⁺ or various metal cations. Notably, elements that do not fulfill the octet rule—those with fewer than eight valence electrons—are also considered Lewis acids. For instance, magnesium in magnesium chloride (MgCl₂) has only four valence electrons when bonded, making it capable of accepting an electron pair. Similarly, aluminum in aluminum bromide (AlBr₃) has six valence electrons, allowing it to act as a Lewis acid. Transition metals, such as nickel and zinc, can also function as Lewis acids due to their ability to expand their d orbitals, accommodating additional electron pairs.

Conversely, a Lewis base is defined as an electron pair donor. The presence of a lone pair on the central atom of a molecule is a key indicator of a Lewis base. For example, oxygen in water (H₂O) and nitrogen in ammonia (NH₃) both possess lone pairs that can be donated to electron-deficient species. When ammonia donates a lone pair to an H⁺ ion, it forms the ammonium ion (NH₄⁺).

Additionally, negatively charged species, such as hydroxide (OH^-), azide (N₃^-), cyanide (CN^-), and hydrosulfide (HS^-), are also Lewis bases due to their abundance of electrons. These species seek to share their excess electrons with electron-deficient atoms or molecules.

In summary, the Lewis theory provides a broader perspective on acid-base interactions, emphasizing the roles of electron pairs in chemical bonding. A Lewis acid accepts electron pairs, while a Lewis base donates them, facilitating various chemical reactions and interactions.

The Adduct

The concept of an adduct arises from the interaction between a Lewis acid and a Lewis base, resulting in a new compound formed by the addition of two reactants. In this case, acetone, represented in blue, acts as the Lewis base, while boron trichloride serves as the Lewis acid. A Lewis acid is defined as an electron pair acceptor, which can be identified by elements that have fewer than eight electrons in their valence shell. Boron, located in group 3A of the periodic table, has three valence electrons and, upon bonding with three chlorine atoms, ends up with only six valence electrons. This deficiency allows boron to accept an electron pair, fulfilling the octet rule.

In this reaction, acetone donates a lone pair of electrons to boron, forming a bond. Unlike the Bronsted-Lowry definition of acids and bases, where an acid donates a proton (H⁺), the Lewis definition emphasizes the donation of electron pairs. This means that when acetone shares its lone pair with boron, it is not merely giving away electrons; it is participating in the formation of the bond, as the electrons are integral to its structure.

It is essential to note that while all Bronsted-Lowry acids and bases can be classified as Lewis acids and bases, the reverse is not true. The Lewis definition is broader, encompassing a wider range of compounds that may not fit the stricter criteria of Bronsted-Lowry or Arrhenius definitions. For instance, boron trichloride qualifies as a Lewis acid due to its electron-accepting ability, but it does not meet the criteria of a Bronsted-Lowry acid since it does not donate an H⁺ ion. Thus, the Lewis framework provides a more inclusive understanding of acid-base chemistry, with Bronsted-Lowry being more specific, and Arrhenius being the most restrictive in its classifications.

Lewis Acid-Base Identification Example

In the study of acids and bases, understanding the distinctions between Lewis and Brønsted-Lowry definitions is crucial. A Lewis acid is defined as a substance that can accept an electron pair, while a Lewis base donates an electron pair. Conversely, a Brønsted-Lowry acid donates a proton (H⁺), and a Brønsted-Lowry base accepts a proton.

For the cobalt(II) ion (Co²⁺), its positive charge indicates that it can accept an electron pair, categorizing it as a Lewis acid. However, it does not have an H⁺ group, so it does not qualify as a Brønsted-Lowry acid.

Nitric acid (HNO₃) contains an H⁺ ion, allowing it to act as a Brønsted-Lowry acid. However, it lacks the characteristics of a Lewis acid since it does not have a central atom with fewer than eight electrons or a positive charge that would facilitate electron pair acceptance.

Considering the compound with the structure CH₃-C(=O)-NH₂, the presence of lone pairs on the nitrogen and oxygen atoms enables it to donate electron pairs, classifying it as a Lewis base. Additionally, it can accept H⁺ ions, making it a Brønsted-Lowry base as well. Thus, this compound can be categorized as both a Lewis base and a Brønsted-Lowry base.

Lastly, the chlorate ion (ClO₃^-) has a negative charge, indicating an excess of electrons available for donation, which qualifies it as a Lewis base. It can also accept a proton to form HClO₃, thus acting as a Brønsted-Lowry base. Therefore, the chlorate ion can be classified as both types of bases.

In summary, cobalt(II) ion is a Lewis acid, nitric acid is a Brønsted-Lowry acid, the compound CH₃-C(=O)-NH₂ is both a Lewis and Brønsted-Lowry base, and the chlorate ion is also both a Lewis and Brønsted-Lowry 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

Answer: (D)

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.