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Chapter 15: Problem 47

Which of the following can behave both as a Bronsted-Lowry acid and as a Bronsted-Lowry base? (a) \(\mathrm{HCO}_{3}^{-}\) (b) CN (c) \(\mathrm{H}_{2} \mathrm{O}\) (d) \(\mathrm{H}_{2} \mathrm{CO}_{3}\)

Short Answer

Expert verified
(a) \\(\text{HCO}_3^-\\) and (c) \\(\text{H}_2\text{O}\\) can act as both.

Step by step solution

01

Understanding Bronsted-Lowry Theory

The Bronsted-Lowry theory defines acids as substances that can donate a proton ( ext{H}^+), while bases are substances that can accept a proton. To identify a compound that can act as both, we look for compounds that contain both a proton to donate and a site to accept a proton.
02

Analyzing \\( ext{HCO}_3^-\\)

The bicarbonate ion \( ext{HCO}_3^-\) can donate a proton to form \( ext{CO}_3^{2-}\), acting as an acid. It can also accept a proton to form \( ext{H}_2 ext{CO}_3\), acting as a base. Therefore, \( ext{HCO}_3^-\) can behave as both an acid and a base.
03

Analyzing CN

The cyanide ion (CN^-) generally acts as a base because it can accept a proton to form HCN, but it lacks a proton to donate, so it cannot act as an acid.
04

Analyzing \\( ext{H}_2 ext{O}\\)

Water \( ext{H}_2 ext{O}\) can donate a proton to become \( ext{OH}^-\), acting as an acid, and can accept a proton to form \( ext{H}_3 ext{O}^+\), acting as a base. Thus, water can act as both an acid and a base.
05

Analyzing \\( ext{H}_2 ext{CO}_3\\)

Carbonic acid \( ext{H}_2 ext{CO}_3\) predominantly acts as an acid because it can donate a proton to become \( ext{HCO}_3^-\). It easily donates but does not have the same ability to accept a proton, so it's not typically considered a base under the Bronsted-Lowry theory.
06

Conclusion

Among the given options, \( ext{HCO}_3^-\) and \( ext{H}_2 ext{O}\) can both act as Bronsted-Lowry acids and bases.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Acid and Base Behavior
The Bronsted-Lowry theory is a fundamental concept in chemistry that defines acids and bases by their ability to donate or accept protons. An acid is a substance that can donate a proton \( (\text{H}^+) \), while a base is one that can accept a proton. This theory helps us understand how different substances interact in chemical reactions, highlighting their dual nature in certain cases where a compound can act as both an acid and a base, such as water. By analyzing these interactions, we identify compounds with both acidic and basic behaviors.
Bicarbonate Ion
The bicarbonate ion \( (\mathrm{HCO}_3^-) \) is an interesting example of a compound that can act as both a Bronsted-Lowry acid and base.
  • As an acid: It donates a proton to become carbonate ion \( (\mathrm{CO}_3^{2-}) \).
  • As a base: It accepts a proton to form carbonic acid \( (\mathrm{H}_2\mathrm{CO}_3) \).
This ability makes it amphoteric, playing a vital role in buffer systems like the blood, where it helps maintain pH balance.
Cyanide Ion
The cyanide ion \( (\mathrm{CN}^-) \), in contrast, mainly acts as a base. It can accept a proton to form hydrogen cyanide \( (\mathrm{HCN}) \),which is a simple acid-base reaction. However, because it lacks a proton to donate, it cannot function as an acid in the Bronsted-Lowry sense. Its behavior is a good illustration of substances that have a predominant single-role function in acid-base chemistry.
Carbonic Acid
Carbonic acid \( (\mathrm{H}_2\mathrm{CO}_3) \) predominantly acts as a Bronsted-Lowry acid. It has the capacity to donate protons readily, forming bicarbonate ion\( (\mathrm{HCO}_3^-) \) in the process. While it acts well as an acid, it generally doesn鈥檛 accept protons, so it isn鈥檛 recognized as a base under this theory. Its primary role involves balancing blood pH by donating protons when necessary to interact with bicarbonate ions.
Water as an Acid and Base
Water \( (\mathrm{H}_2\mathrm{O}) \) is a quintessential example of a compound that can behave both as a Bronsted-Lowry acid and base.
  • As an acid: It donates a proton to become hydroxide ion\( (\mathrm{OH}^-) \).
  • As a base: It accepts a proton to become hydronium ion\( (\mathrm{H}_3\mathrm{O}^+) \).
This property of water not only supports its role in acid-base chemistry but also in sustaining the dynamic equilibrium in various biological and chemical systems.

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Most popular questions from this chapter

For each of the following reactions, identify the Bronsted-Lowry acids and bases and the conjugate acid-base pairs: (a) \(\mathrm{CN}^{-}(a q)+\mathrm{H}_{2} \mathrm{O}(l) \rightleftharpoons \mathrm{OH}^{-}(a q)+\mathrm{HCN}(a q)\) (b) \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}(a q)+\mathrm{H}_{2} \mathrm{O}(l) \rightleftharpoons \mathrm{H}_{3} \mathrm{O}^{+}(a q)+\mathrm{HPO}_{4}^{2-}(a q)\) (c) \(\mathrm{HPO}_{4}^{2-}(a q)+\mathrm{H}_{2} \mathrm{O}(l) \rightleftharpoons \mathrm{OH}(a q)+\mathrm{H}_{2} \mathrm{PO}_{4}^{-}(a q)\) (d) \(\mathrm{NH}_{4}^{+}(a q)+\mathrm{NO}_{2}^{-}(a q) \rightleftharpoons \mathrm{HNO}_{2}(a q)+\mathrm{NH}_{3}(a q)\)

Calculate the \(\mathrm{pH}\) of solutions prepared by: (a) Dissolving \(0.20 \mathrm{~g}\) of sodium oxide in water to give \(100.0 \mathrm{~mL}\) of solution (b) Dissolving \(1.26 \mathrm{~g}\) of pure nitric acid in water to give \(0.500 \mathrm{~L}\) of solution (c) Diluting \(40.0 \mathrm{~mL}\) of \(0.075 \mathrm{M} \mathrm{Ba}(\mathrm{OH})_{2}\) to a volume of \(300.0 \mathrm{~mL}\) (d) Mixing equal volumes of \(0.20 \mathrm{M} \mathrm{HCl}\) and \(0.50 \mathrm{M} \mathrm{HNO}_{3}\) (Assume that volumes are additive.)

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-

$$ \begin{aligned} &\text { Calculate the percent dissociation of } \mathrm{HNO}_{2}\left(K_{\mathrm{a}}=4.5 \times 10^{-4}\right)\\\ &\text { in: (a) } 0.010 \mathrm{MHNO}_{2} \text { (b) } 1.00 \mathrm{M} \mathrm{HNO}_{2} \end{aligned} $$

Give the formula for the conjugate acid of each of the following Bronsted- Lowry bases: (a) \(\mathrm{SO}_{3}^{2-}\) (b) \(\mathrm{H}_{2} \mathrm{O}\) (c) \(\mathrm{CH}_{3} \mathrm{NH}_{2}\) (d) \(\mathrm{OH}\) (e) \(\mathrm{HCO}_{3}^{-}\) (f) \(\mathrm{H}\)
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