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Chapter 4: Problem 158

Can the following decomposition reaction be characterized as an acid-base reaction? Explain. \(\mathrm{NH}_{4} \mathrm{Cl}(s) \longrightarrow \mathrm{NH}_{3}(g)+\mathrm{HCl}(g)\)

Short Answer

Expert verified
Yes, the decomposition reaction of ammonium chloride can be characterized as an acid-base reaction. In this reaction, there is a transfer of a proton from the ammonium ion to the chloride ion, a key feature of acid-base reactions.

Step by step solution

01

Understanding the Reaction

First, look at the chemical reaction provided: \(NH_4Cl(s) \rightarrow NH_3(g) + HCl(g)\). This is a decomposition reaction where solid ammonium chloride decomposes into ammonia gas and hydrogen chloride gas.
02

Identify the Possibility of Proton Transfer

In an acid-base reaction, there is a transfer of a proton (H+) from one substance to another. To know if this reaction is an acid-base reaction, identify if there is a proton transfer in the decomposition reaction of ammonium chloride. In the given reaction, one of the hydrogen atoms from the ammonium ion (NH4+) is separated and joins with a chloride ion (Cl-) to form HCl. This illustrates a transfer of a proton (H+).
03

Conclusion

Because there is a proton transfer from ammonium to chloride, this decomposition reaction can indeed be characterized as an acid-base reaction. In this case, the ammonium ion (NH4+) acts as the acid (proton donor), and the chloride ion (Cl-) acts as the base (proton acceptor).

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Most popular questions from this chapter

Chemical tests of four metals \(\mathrm{A}, \mathrm{B}, \mathrm{C},\) and \(\mathrm{D}\) show the following results. (a) Only \(\mathrm{B}\) and \(\mathrm{C}\) react with \(0.5 \mathrm{M} \mathrm{HCl}\) to give \(\mathrm{H}_{2}\) gas. (b) When \(\mathrm{B}\) is added to a solution containing the ions of the other metals, metallic \(\mathrm{A}, \mathrm{C},\) and \(\mathrm{D}\) are formed. (c) A reacts with \(6 M \mathrm{HNO}_{3}\) but \(\mathrm{D}\) does not. Arrange the metals in the increasing order as reducing agents. Suggest four metals that fit these descriptions.

How would you prepare \(60.0 \mathrm{~mL}\) of \(0.200 \mathrm{M} \mathrm{HNO}_{3}\) from a stock solution of \(4.00 \mathrm{M} \mathrm{HNO}_{3}\) ?

Would the volume of a \(0.10 M \mathrm{NaOH}\) solution needed to titrate \(25.0 \mathrm{~mL}\) of a \(0.10 \mathrm{M} \mathrm{HNO}_{2}\) (a weak acid) solution be different from that needed to titrate \(25.0 \mathrm{~mL}\) of a \(0.10 \mathrm{M} \mathrm{HCl}\) (a strong acid) solution?

The following "cycle of copper" experiment is performed in some general chemistry laboratories. The series of reactions starts with copper and ends with metallic copper. The steps are as follows: (1) A piece of copper wire of known mass is allowed to react with concentrated nitric acid [the products are copper(II) nitrate, nitrogen dioxide, and water]. (2) The copper(II) nitrate is treated with a sodium hydroxide solution to form copper(II) hydroxide precipitate. (3) On heating, copper(II) hydroxide decomposes to yield copper(II) oxide. (4) The copper(II) oxide is reacted with concentrated sulfuric acid to yield copper(II) sulfate. (5) Copper(II) sulfate is treated with an excess of zinc metal to form metallic copper. (6) The remaining zinc metal is removed by treatment with hydrochloric acid, and metallic copper is filtered, dried, and weighed. (a) Write a balanced equation for each step and classify the reactions. (b) Assuming that a student started with \(65.6 \mathrm{~g}\) of copper, calculate the theoretical yield at each step. (c) Considering the nature of the steps, comment on why it is possible to recover most of the copper used at the start.

A 5.012 -g sample of an iron chloride hydrate was dried in an oven. The mass of the anhydrous compound was \(3.195 \mathrm{~g}\). The compound was then dissolved in water and reacted with an excess of \(\mathrm{AgNO}_{3} .\) The AgCl precipitate formed weighed 7.225 g. What is the formula of the original compound?
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