Question

Using curved arrows, show how the species in part (a) can act as Lewis bases in their reactions with HCl, and show how the species in part (b) can act as Lewis acids in their reaction with \(\mathrm{OH}^{-}\). (a) \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OH}, \mathrm{HN}\left(\mathrm{CH}_{3}\right)_{2}, \mathrm{P}\left(\mathrm{CH}_{3}\right)_{3}\) (b) \(\mathrm{H}_{3} \mathrm{C}^{+}, \mathrm{B}\left(\mathrm{CH}_{3}\right)_{3}, \mathrm{MgBr}_{2}\)

Short answer

Lewis bases donate electron pairs to HCl; Lewis acids accept electron pairs from \( \mathrm{OH}^{-} \).

Step-by-step solution

Identify Lewis Bases (a)

The compounds listed in part (a) - \( \mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OH} \), \( \mathrm{HN}\left(\mathrm{CH}_{3}\right)_{2} \), and \( \mathrm{P}\left(\mathrm{CH}_{3}\right)_{3} \) - can donate a pair of electrons to form a bond with an electron-deficient atom, acting as Lewis bases. Ethanol (\( \mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OH} \)) has an oxygen with lone pairs, dimethylamine (\( \mathrm{HN}\left(\mathrm{CH}_{3}\right)_{2} \)) has a nitrogen atom with lone pairs, and trimethylphosphine (\( \mathrm{P}\left(\mathrm{CH}_{3}\right)_{3} \)) also has a phosphorus atom with lone pairs.

Draw Curved Arrows for Lewis Bases (a)

For \( \mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OH} \), draw an arrow from the lone pair on the oxygen to the hydrogen in HCl to form bonding with H. Similarly, for \( \mathrm{HN}\left(\mathrm{CH}_{3}\right)_{2} \), draw an arrow from the lone pair on nitrogen to the hydrogen of HCl. For \( \mathrm{P}\left(\mathrm{CH}_{3}\right)_{3} \), draw an arrow from the lone pair on phosphorus to the hydrogen in \( \mathrm{HCl} \).

Identify Lewis Acids (b)

The compounds in part (b) - \( \mathrm{H}_{3} \mathrm{C}^{+} \), \( \mathrm{B}\left(\mathrm{CH}_{3}\right)_{3} \), and \( \mathrm{MgBr}_{2} \) - can accept a pair of electrons, acting as Lewis acids. \( \mathrm{H}_{3} \mathrm{C}^{+} \) is electron-deficient, \( \mathrm{B}\left(\mathrm{CH}_{3}\right)_{3} \) has an incomplete octet, and \( \mathrm{MgBr}_{2} \) can accept electrons due to magnesium's positive charge.

Draw Curved Arrows for Lewis Acids (b)

For \( \mathrm{H}_{3} \mathrm{C}^{+} \), draw an arrow from the lone pair on \( \mathrm{OH}^{-} \) to the carbon atom to show electron pair donation. For \( \mathrm{B}\left(\mathrm{CH}_{3}\right)_{3} \), draw an arrow from the lone pair on \( \mathrm{OH}^{-} \) to the boron atom. For \( \mathrm{MgBr}_{2} \), draw an arrow from the lone pair on \( \mathrm{OH}^{-} \) to the magnesium atom.

Key concepts

Understanding Curved Arrows

Curved arrows in chemistry serve as a visual guide to illustrate how electron pairs move during chemical reactions. This concept is frequently used to describe the mechanism of how Lewis acids and bases interact. Whenever you see a curved arrow, it represents the path that an electron pair takes as it moves from the Lewis base to the Lewis acid.
For instance, in the case of a Lewis base, which can donate an electron pair, you'd see a curved arrow starting from the electron donor (the atom with the lone pair) moving towards the electron acceptor which is usually an atom with a high affinity for electrons. This electron acceptor is the Lewis acid.

• In Lewis base reactions: The arrow starts at the lone electron pair on the base and points to the electron-deficient center.
• In Lewis acid reactions: The arrow points towards the atom that can accept electrons.

This visual representation simplifies the understanding of bonding interactions by clearly indicating where the electrons are transferring and how new bonds are formed.

The Role of Electron Pairs

Electron pairs play a crucial role in chemistry, particularly in the formation of chemical bonds. A pair of electrons can be shared by atoms to form a covalent bond. However, when discussing Lewis acids and bases, we're usually talking about lone pairs, which are not shared and are available for bonding.
In the context of Lewis bases, these lone electron pairs are key players. For instance, in ethanol (\( \mathrm{CH}_{3}\mathrm{CH}_{2}\mathrm{OH} \)), the oxygen atom has a lone pair, which it can use to form a bond with a proton from hydrochloric acid (HCl). This is the role of the Lewis base: donating these lone pairs to form new bonds.
Conversely, Lewis acids, such as \( \mathrm{H}_{3}\mathrm{C}^{+} \), are electron pair acceptors. They are often electron-deficient species, which makes them eager to accept an electron pair from a Lewis base. This acceptance allows them to fill their valence shell, completing their electronic configuration. Thus, understanding how electron pairs move is vital for mastering the concept of bonding interactions.

Bonding Interactions in Lewis Acids and Bases

Bonding interactions between Lewis acids and bases are foundational in chemistry, explaining how certain reactions occur. When a Lewis base donates an electron pair to a Lewis acid, a coordinate covalent bond (or dative bond) is formed.
A few key points to understand about these interactions:

• Lewis Bases: These are usually molecules with lone electron pairs, capable of donating those pairs to electron-deficient atoms.
• Lewis Acids: These molecules have an incomplete electron octet or are positively charged, making them eager to accept electron pairs.

Consider the interaction of OH⁻ with \( \mathrm{H}_{3} \mathrm{C}^{+} \). The OH⁻ ion uses its lone pair to form a bond with the electron-deficient carbon in \( \mathrm{H}_{3} \mathrm{C}^{+} \), stabilizing the molecule by compensating for its positive charge.
Moreover, these interactions are depicted through curved arrows, which show the electron flow from the donor (Lewis base) to the acceptor (Lewis acid), simplifying the visualization of chemical reactions and promoting a better understanding of molecular transformations.