Question

Boric acid \(\left(\mathrm{H}_{3} \mathrm{BO}_{3}\right)\) is a weak monoprotic acid that yields \(\mathrm{H}_{3} \mathrm{O}^{+}\) ions in water. \(\mathrm{H}_{3} \mathrm{BO}_{3}\) might behave either as a Bronsted-Lowry acid or as a Lewis acid, though it is, in fact, a Lewis acid. (a) Write a balanced equation for the reaction with water in which \(\mathrm{H}_{3} \mathrm{BO}_{3}\) behaves as a Bronsted-Lowry acid. (b) Write a balanced equation for the reaction with water in which \(\mathrm{H}_{3} \mathrm{BO}_{3}\) behaves as a Lewis acid. Hint: One of the reac-

Short answer

(a) \(\mathrm{H}_{3} \mathrm{BO}_{3} + \mathrm{H}_{2} \mathrm{O} \rightarrow \mathrm{H}_{2} \mathrm{BO}_{3}^{-} + \mathrm{H}_{3} \mathrm{O}^{+}\); (b) \(\mathrm{H}_{3} \mathrm{BO}_{3} + \mathrm{2}\mathrm{H}_{2} \mathrm{O} \rightarrow \mathrm{B(OH)_{4}^{-}} + \mathrm{H}_{3} \mathrm{O}^{+}\).

Step-by-step solution

Identify The Bronsted-Lowry Acid Reaction

A Bronsted-Lowry acid donates a proton (H⁺) to a base. In this reaction, boric acid donates an H⁺ to water. The equation for this reaction is:\[\mathrm{H}_{3} \mathrm{BO}_{3} + \mathrm{H}_{2} \mathrm{O} \rightarrow \mathrm{H}_{2} \mathrm{BO}_{3}^{-} + \mathrm{H}_{3} \mathrm{O}^{+}.\]

Identify The Lewis Acid Reaction

A Lewis acid accepts an electron pair. In the case of boric acid, it can accept a hydroxide ion (OH⁻) from water, forming the tetrahydroxyborate ion (\[\mathrm{B(OH)_{4}^{-}}\]). The equation for this reaction is:\[\mathrm{H}_{3} \mathrm{BO}_{3} + \mathrm{2}\mathrm{H}_{2} \mathrm{O} \rightarrow \mathrm{B(OH)_{4}^{-}} + \mathrm{H}_{3} \mathrm{O}^{+}.\]

Key concepts

Bronsted-Lowry acid

In the world of chemistry, a Bronsted-Lowry acid is anything that donates a proton, or a hydrogen ion (\( H^+ \)), to another substance. This definition emphasizes the acid's role in a chemical reaction and focuses on the transfer of protons. This type of acid-base reaction takes place between an acid (proton donor) and a base (proton acceptor). For example, when boric acid (\( \mathrm{H}_{3} \mathrm{BO}_{3} \)) reacts with water, it donates a proton to form a hydronium ion (\( \mathrm{H}_{3} \mathrm{O}^{+} \)). In this case:

• Boric acid acts as the Bronsted-Lowry acid by donating a proton.
• Water acts as a base by accepting the proton, forming hydronium ion.

Understanding this concept provides insight into how certain substances can behave as acids, specifically in aqueous solutions.

Lewis acid

A Lewis acid, in contrast to the Bronsted-Lowry acid, is defined based on its ability to accept an electron pair rather than donate a proton. This concept focuses on the electron interactions rather than proton transfers. In basic terms, when a molecule like boric acid accepts an electron pair from another molecule, it acts as a Lewis acid.
For instance, when boric acid (\( \mathrm{H}_{3} \mathrm{BO}_{3} \)) reacts with water, it may accept a hydroxide ion (\( \mathrm{OH}^- \)), resulting in the formation of \( \mathrm{B(OH)_{4}^{-}} \). In this reaction:

• Boric acid accepts an electron pair from the hydroxide ion, showcasing Lewis acid behavior.
• This expands the notion of acidity beyond just proton donation, reflecting a broader understanding of chemical reactivity.

Recognizing a Lewis acid thus requires a focus on the acceptance of electron pairs in a chemical reaction.

boric acid reactions

Boric acid is known as a weak acid, specifically a weak monoprotic acid, which means it can release only one proton per molecule. Its interesting trait is that it can behave in different ways, showing the dual nature of acid-base chemistry.
In its reaction with water:

• As a Bronsted-Lowry acid, it donates a proton to water, leading to the formation of a hydronium ion (\( \mathrm{H}_{3} \mathrm{O}^{+} \)).
• As a Lewis acid, it accepts a hydroxide ion (\( \mathrm{OH}^- \)) to form a tetrahydroxyborate ion (\( \mathrm{B(OH)_{4}^{-}} \)).

These reactions demonstrate the versatile chemistry of boric acid, and help us appreciate the complex and diverse nature of chemical reactions involving acid and bases.

balanced chemical equations

Writing balanced chemical equations is an essential aspect of understanding chemical reactions. A balanced equation ensures that the number of atoms for each element is equal on both the reactant and the product sides.
For boric acid:

• As a Bronsted-Lowry acid with water, the balanced equation is \( \mathrm{H}_{3} \mathrm{BO}_{3} + \mathrm{H}_{2} \mathrm{O} \rightarrow \mathrm{H}_{2} \mathrm{BO}_{3}^{-} + \mathrm{H}_{3} \mathrm{O}^{+} \), maintaining balance by equalizing the number of hydrogen and oxygen atoms.
• As a Lewis acid with water, the balanced equation is \( \mathrm{H}_{3} \mathrm{BO}_{3} + 2 \mathrm{H}_{2} \mathrm{O} \rightarrow \mathrm{B(OH)_{4}^{-}} + \mathrm{H}_{3} \mathrm{O}^{+} \). This represents the proper balance of all atoms involved.

Balanced equations are crucial as they provide accurate representations of chemical processes, allowing chemists to predict the outcomes of reactions and the quantities of substances consumed or produced.