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Chapter 3: Problem 39

Write an equation that describes the equilibrium that exists when nitric acid dissolves in water. Identify each of the four species in solution as either Bronsted acids or Bronsted bases. Does the equilibrium favor the products or the reactants?

Short Answer

Expert verified
The equilibrium favors the products, with HNO鈧� as a Br酶nsted acid and NO鈧冣伝 as a base.

Step by step solution

01

Write the Dissolution Reaction

When nitric acid (HNO鈧�) dissolves in water, it donates a proton (H鈦�) to water, producing hydronium ions (H鈧僌鈦�) and nitrate ions (NO鈧冣伝). The equilibrium equation for this reaction is: \[\text{HNO鈧儅 + \text{H鈧侽} \rightleftharpoons \text{H鈧僌}^+ + \text{NO鈧儅^-\] This reaction shows nitric acid dissociating into hydronium ions and nitrate ions.
02

Identify the Species as Br酶nsted Acids or Bases

In the equation, HNO鈧� acts as a Br酶nsted acid because it donates a proton (H鈦�) to water. Water ( H鈧侽) accepts the proton to become H鈧僌鈦� and acts as a Br酶nsted base. H鈧僌鈦�, the hydronium ion, is the conjugate acid that forms, while NO鈧冣伝 is the conjugate base that results from the deprotonation of HNO鈧�.
03

Determine the Equilibrium Direction

Nitric acid is a strong acid, meaning it ionizes completely in water. This implies that the equilibrium lies far to the right, favoring the formation of products (H鈧僌鈦� and NO鈧冣伝) over the reactants (HNO鈧� and H鈧侽). Therefore, the equilibrium significantly favors the product side.

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Key Concepts

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

Br酶nsted Acid
When discussing the concept of a Br酶nsted acid, we are focusing on a substance that donates a proton, which is essentially a hydrogen ion (H鈦�), in a chemical reaction.
In the dissolution of nitric acid (HNO鈧�) in water, HNO鈧� acts as the Br酶nsted acid, since it donates a proton to water.
  • The formula for this reaction is: \[\text{HNO鈧儅 + \text{H鈧侽} \rightleftharpoons \text{H鈧僌}^+ + \text{NO鈧儅^-\]
  • After donating a proton, HNO鈧� becomes NO鈧冣伝, the nitrate ion.
  • This transformation demonstrates the principal behavior of a Br酶nsted acid鈥攑roton donation.
Understanding the role of HNO鈧� as a Br酶nsted acid is crucial, as it sets up the basic framework for evaluating acid-base equilibriums.If you imagine acids as proton donors, it becomes easier to identify other acids by observing which substances donate hydrogen ions in reactions.
Br酶nsted Base
While the Br酶nsted acid in a reaction is the donor of H鈦� , a Br酶nsted base, in contrast, accepts a hydrogen ion. In the example of nitric acid dissolving in water, water ( H鈧侽 ) functions as the Br酶nsted base.
This is because it accepts the hydrogen ion from HNO鈧� and forms H鈧僌鈦� , the hydronium ion.
  • During this process, H鈧侽 acts as a proton acceptor.
  • This proton acceptance transforms water into H鈧僌鈦� , linking the conversion to its role as a base.
  • The ability of a substance to accept a proton defines it as a Br酶nsted base regardless of its original state.
Recognizing H鈧侽 as a Br酶nsted base in this scenario not only helps classify the substances involved, but also clarifies the direction of proton movement in acid-base reactions. Having a clear understanding of Br酶nsted bases is especially important for solving equilibrium problems where both acids and bases must be identified to ascertain how reactions proceed.
Equilibrium Direction
In chemical reactions, equilibrium direction refers to the tendency of a reaction to favor either the reactants or products once the reaction has reached a stable state. This concept is vital in understanding how strong acids like nitric acid behave in solution.
Since HNO鈧� is a strong acid, it dissociates almost completely in water.
This has a significant impact on the equilibrium of the reaction.
  • In this scenario: \[\text{HNO鈧儅 + \text{H鈧侽} \rightleftharpoons \text{H鈧僌}^+ + \text{NO鈧儅^-\]
  • The equilibrium heavily favors the right side, meaning a majority of HNO鈧� converts to H鈧僌鈦� and NO鈧冣伝.
  • This is indicative of the strong acid鈥檚 propensity to release its protons almost completely.
Understanding the equilibrium direction helps predict the extent of acid ionization in solution and influences how we measure the strength of acids and equally, the weakness of associated bases in any given chemical environment.

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Most popular questions from this chapter

Aqueous solutions of ammonium sulfide, \(\left(\mathrm{NH}_{4}\right)_{2} \mathrm{S}\) and \(\mathrm{Hg}\left(\mathrm{NO}_{3}\right)_{2}\) react to produce \(\mathrm{HgS}\) and \(\mathrm{NH}_{4} \mathrm{NO}_{3}\) (a) Write the overall, balanced equation for the reaction. Indicate the state \((s,\) aq ) for each compound. (b) Name each compound. (c) What type of reaction is this?

Predict the products of each precipitation reaction. Balance the equation, and then write the net ionic equation. (a) \(\mathrm{Pb}\left(\mathrm{NO}_{3}\right)_{2}(\mathrm{aq})+\mathrm{KBr}(\mathrm{aq}) \rightarrow\) (b) \(\mathrm{Ca}\left(\mathrm{NO}_{3}\right)_{2}(\mathrm{aq})+\mathrm{KF}(\mathrm{aq}) \rightarrow\) (c) \(\mathrm{Ca}\left(\mathrm{NO}_{3}\right)_{2}(\mathrm{aq})+\mathrm{Na}_{2} \mathrm{C}_{2} \mathrm{O}_{4}(\mathrm{aq}) \rightarrow\)

Most naturally occurring acids are weak acids. Lactic acid is one example. $$\mathrm{CH}_{3} \mathrm{CH}(\mathrm{OH}) \mathrm{CO}_{2} \mathrm{H}(\mathrm{s})+\mathrm{H}_{2} \mathrm{O}(\ell) \rightleftharpoons \mathrm{H}_{3} \mathrm{O}^{+}(\mathrm{aq})+\mathrm{CH}_{3} \mathrm{CH}(\mathrm{OH}) \mathrm{CO}_{2}^{-}(\mathrm{aq})$$ If you place some lactic acid in water, it will ionize to a small extent, and an equilibrium will be established. Suggest some experiments to prove that this is a weak acid and that the establishment of equilibrium is a reversible process.

Write an overall, balanced equation for the reaction of \(\mathrm{Na}_{2} \mathrm{SO}_{3}\) with \(\mathrm{CH}_{3} \mathrm{CO}_{2} \mathrm{H},\) and name the reactants and products.

Gas evolution was observed when a solution of \(\mathrm{Na}_{2} \mathrm{S}\) was treated with acid. The gas was bubbled into a solution containing \(\mathrm{Pb}\left(\mathrm{NO}_{3}\right)_{2},\) and a black precipitate formed. Write net ionic equations for the two reactions.
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