Question

Consider the base hydroxylamine, \(\mathrm{NH}_{2} \mathrm{OH}\). (a) What is the conjugate acid of hydroxylamine? (b) When it acts as a base, which atom in hydroxylamine accepts a proton? (c) There are two atoms in hydroxylamine that have nonbonding electron pairs that could act as proton acceptors. Use Lewis structures and formal charges ers (Section 8.5) to rationalize why one of these two atoms is a much better proton acceptor than the other.

Short answer

(a) The conjugate acid of hydroxylamine (NH₂OH) is NH₃O⁺. (b) The oxygen atom in hydroxylamine accepts a proton when it acts as a base. (c) Although both nitrogen and oxygen atoms have nonbonding electron pairs and formal charges of 0, the oxygen atom is a better proton acceptor due to its higher electronegativity.

Step-by-step solution

(a) Determine the conjugate acid of hydroxylamine

In order to determine the conjugate acid of hydroxylamine, we need to add a proton (H⁺) to the molecule. A conjugate acid is formed when a base accepts a proton. Hydroxylamine (NH₂OH) can accept a proton on its oxygen atom, leading to the formation of the conjugate acid with the chemical formula NH₃O⁺.

(b) Identify the atom in hydroxylamine that accepts a proton

When hydroxylamine acts as a base, it donates an electron pair to a proton (H⁺). In the case of hydroxylamine (NH₂OH), the oxygen atom contains a lone pair of electrons, which allows it to accept a proton, forming the conjugate acid NH₃O⁺. Therefore, the oxygen atom in hydroxylamine is the one that accepts a proton when it acts as a base.

(c) Use Lewis structures and formal charges to determine the better proton acceptor

There are two atoms in hydroxylamine that have nonbonding electron pairs - the nitrogen atom and the oxygen atom. To determine which of these two atoms is a better proton acceptor, we need to examine the Lewis structures and formal charges of both. Here are the Lewis structures for both atoms in hydroxylamine: -> For nitrogen: [N] = 1s²2s²2p³ -> For oxygen: [O] = 1s²2s²2p⁴ Formal charges are calculated using the formula: Formal charge = (Valence electrons) - (Nonbonding electrons) - 0.5*(Bonding electrons) -> For nitrogen: Formal charge = 5 - 2 - 0.5*(6) = 0 -> For oxygen: Formal charge = 6 - 4 - 0.5*(4) = 0 Both atoms have formal charges of 0, which means that neither atom is more likely to accept a proton based on formal charge alone. However, oxygen is more electronegative than nitrogen, meaning it has a greater attraction for electrons. This makes the oxygen atom more likely to accept a proton, making it a better proton acceptor than the nitrogen atom in hydroxylamine.

Key concepts

Conjugate Acid of Hydroxylamine

Understanding the concept of a conjugate acid is crucial in acid-base chemistry. In the case of hydroxylamine (\textbf{NH}\(_2\)\textbf{OH}), to find the conjugate acid, we envision the molecule accepting an extra hydrogen ion (\textbf{H}\(^+\)). This process transforms hydroxylamine into its conjugate acid, resulting in the new chemical formula \textbf{NH}\(_3\)\textbf{O}\(^+\).

When a base gains a proton, it forms its conjugate acid, which is essentially its partner in the acid-base pair. The chemical behavior of the conjugate acid of hydroxylamine is quite different from hydroxylamine itself, showcasing the versatile nature of such molecules in chemical reactions. The understanding of conjugate acids is pivotal for students delving into topics such as titration, buffer solutions, and equilibrium in aqueous solutions.

Proton Acceptor in Hydroxylamine

The role of a proton acceptor in a base is integral to the concept of acid-base reactions. Hydroxylamine (\textbf{NH}\(_2\)\textbf{OH}) acts as a base by offering an electron pair to bond with a free hydrogen ion (\textbf{H}\(^+\)). This behavior is due to the presence of nonbonding electron pairs, often referred to as lone pairs, on the molecule.

Specifically, the oxygen atom in hydroxylamine is the site of this action. Oxygen has two lone pairs of electrons and it readily uses one of these pairs to bond with the hydrogen ion, converting hydroxylamine to its conjugate acid, \textbf{NH}\(_3\)\textbf{O}\(^+\). This capacity to accept a proton is a defining characteristic of base functionality, and students are encouraged to identify the potential sites for proton acceptance when analyzing different molecules.

Lewis Structures and Formal Charges

Lewis structures are visual representations that illustrate the bonding between atoms within a molecule, as well as any nonbonding electron pairs. Formal charges, on the other hand, provide insight into the electronegativity and distribution of electrons in the molecule.

Understanding Lewis Structures

To create a Lewis structure, we arrange the atoms to show how they're bonded and then allocate dots around the atoms, showcasing the nonbonding (lone) pairs.

Determining Formal Charges

The formula for calculating formal charges is: Formal charge = (Valence electrons) - (Nonbonding electrons) - 0.5*(Bonding electrons).

In hydroxylamine (\textbf{NH}\(_2\)\textbf{OH}), both the nitrogen and oxygen atoms have a formal charge of zero, indicating a stable configuration. However, considering the greater electronegativity of oxygen, it is preferred over nitrogen for accepting a proton. This detail can impact chemical reactivity and is particularly relevant in the study of acid-base chemistry, resonance structures, and reaction mechanisms. Recognizing these properties allows students to predict molecular behavior and the outcomes of chemical reactions more accurately.