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 000 (Section 8.5 ) to rationalize why one of these two atoms is a much better proton acceptor than the other.

Short answer

The conjugate acid of hydroxylamine (NH₂OH) is NH₃OH⁺. When acting as a base, the oxygen atom in hydroxylamine accepts a proton due to its greater electronegativity, which leads to better stability of the conjugate acid.

Step-by-step solution

Understanding conjugate acids and bases

A conjugate acid is formed when a base accepts a proton (H+ ion), while a conjugate base is formed when an acid donates a proton. In this question, we need to find the conjugate acid of hydroxylamine, an amphoteric compound, meaning it can act as either an acid or base.

Identify the conjugate acid of hydroxylamine

When hydroxylamine (NH₂OH) acts as a base, it can accept a proton (H⁺) to form its conjugate acid. The resulting conjugate acid would be NH₃OH⁺.

Identify the atom in hydroxylamine that accepts the proton

When hydroxylamine acts as a base, the oxygen atom will accept the proton since it has lone pairs of electrons to donate, creating NH₃OH⁺.

Analyze the Lewis structures and formal charges for proton acceptor ability

Hydroxylamine contains two atoms with nonbonding electron pairs: nitrogen and oxygen. We'll examine their Lewis structures and formal charges to determine which atom is a better proton acceptor. (a) The Lewis structure of NH₂OH: N: valence electrons = 5, non-bonding electrons = 2, bonding electrons = 6 Formal charge = (5 - (2 + 6/2)) = 0 O: valence electrons = 6, non-bonding electrons = 4, bonding electrons = 4 Formal charge = (6 - (4 + 4/2)) = 0 (b) The Lewis structure of the conjugate acid NH₃OH⁺: N: valence electrons = 5, non-bonding electrons = 0, bonding electrons = 8 Formal charge = (5 - (0 + 8/2)) = 0 O: valence electrons = 6, non-bonding electrons = 4, bonding electrons = 5 Formal charge = (6 - (4 + 5/2)) = 0 Comparing the formal charges for the nitrogen and oxygen atoms, we see that the nitrogen atom remains neutral before and after accepting the proton. The oxygen atom also remains neutral as it has a formal charge of 0 in both cases. However, an oxygen atom is more electronegative than a nitrogen atom, making it a more favorable proton acceptor as it can facilitate greater stability due to its strong electronegativity. In conclusion, when hydroxylamine acts as a base: a) The conjugate acid is NH₃OH⁺. b) The oxygen atom accepts a proton. c) The oxygen atom is a better proton acceptor because its electronegativity stabilizes the conjugate acid, leading to its higher likelihood of accepting a proton over the nitrogen atom.

Key concepts

Hydroxylamine

Hydroxylamine is a chemical compound with the formula ${\text{NH}}_2\text{OH}$. It is known as an amphoteric compound, meaning it can behave both as an acid and a base.
This versatility arises because hydroxylamine has atoms with lone pairs of electrons, enabling it to either donate electrons in a basic reaction or accept protons in an acidic reaction. In the context of acid-base reactions, hydroxylamine typically acts as a base.
When it accepts a proton (${\text{H}}^{+}$), it forms the conjugate acid ${\text{NH}}_3{\text{OH}}^{+}$.

• Hydroxylamine acts as a base when it donates its electron pair to attach a proton.
• This characteristic allows it to play a role in various chemical reactions, particularly in studies of catalysis and synthesis.

Understanding the behavior of hydroxylamine in forming its conjugate acid helps elucidate its reactivity in different reactions.

Lewis Structures

Lewis structures are diagrams used to represent the molecular structure of a compound, showing how the atoms are bonded together and the placement of nonbonded electrons.

For hydroxylamine, ${\text{NH}}_2\text{OH}$, we draw Lewis structures to determine the positions of atoms and electron pairs.

• Nitrogen (N) has five valence electrons, forming three covalent bonds: two with hydrogen atoms and one with an oxygen atom.
• Oxygen (O) has six valence electrons and forms a covalent bond with nitrogen while possessing two lone pairs of electrons.

This electron arrangement is vital for predicting chemical reactions. Lewis structures help us visualize molecules, understanding which atoms might participate in a reaction based on their electron availability. Utilizing these structures makes it easier to predict how hydroxylamine will interact when it encounters other compounds.

Formal Charge

The concept of formal charge is essential in evaluating the most stable Lewis structures and understanding molecule interactions.
A formal charge indicates the electron distribution within an atom when compared to its ground state, providing insights into which atom might attract or donate electrons.

For hydroxylamine, the formal charge is calculated for each atom using the formula:$$\text{Formal Charge}=\text{Valence Electrons}-(\text{Nonbonding Electrons}+\frac{1}{2}\text{Bonding Electrons})$$

• Nitrogen (N) in ${\text{NH}}_2\text{OH}$ has a formal charge of 0, as it retains its original five electrons in terms of distribution.
• Oxygen (O) also has a formal charge of 0 in both free and bound states, signifying electron neutrality when bonded.

Despite having identical formal charges, the electronegativity difference leads to unequal ease of electron sharing, with oxygen's higher electronegativity tipping the balance towards oxygen being a favorable proton acceptor.

Electron Pairs

In chemistry, electron pairs play a crucial role in determining molecular geometry and reactivity.
Electron pairs can be shared between atoms to form bonds or remain unshared if they belong entirely to one atom, termed nonbonding or lone pairs.

In hydroxylamine ${\text{NH}}_2\text{OH}$:

• Nitrogen has one nonbonding pair of electrons after forming bonds with oxygen and two hydrogen atoms.
• Oxygen holds two lone pairs, indicating spaces where it can interact through potential bond formation or chemical reactions.

These nonbonding electron pairs are critical as they determine the basicity of the molecule. In proton acceptance dynamics in hydroxylamine, the lone pairs on the oxygen atom are primarily responsible for its role as a base because they can easily form bonds with donating protons due to oxygen's electronegativity. Understanding electron pairs enables predictions about which parts of a molecule will most likely participate in chemical bonding.