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Chapter 18: Problem 105

Analyze and Conclude Is it possible that an a Arrhenius acid is not a Br酶nsted-Lowry acid? Is it possible that an acid according to the Br酶nsted- Lowry model is not an Arrhenius acid? Is it possible that a Lewis acid could not be classified as either an Arrhenius or a Br酶nsted-Lowry acid? Explain and give examples.

Short Answer

Expert verified
In summary, it is not possible for an Arrhenius acid to not be a Br酶nsted-Lowry acid, as both involve the release of H+ ions. It is possible for a Br酶nsted-Lowry acid to not be an Arrhenius acid, as seen with gaseous ammonia (NH3). Finally, it is possible for a Lewis acid to not be classified as either an Arrhenius or Br酶nsted-Lowry acid, as demonstrated by boron trifluoride (BF3), which accepts electron pairs but does not donate protons or increase H+ concentration in water.

Step by step solution

01

Case 1: An Arrhenius acid that is not a Br酶nsted-Lowry acid

According to the Arrhenius definition, an acid is a substance that, when dissolved in water, increases the concentration of hydrogen ions (H+). The Br酶nsted-Lowry definition states that an acid is a substance that donates a proton (H+) to another substance (a base). Both definitions involve the presence of H+ ions, so it can be concluded that any Arrhenius acid must also be a Br酶nsted-Lowry acid since it involves the release of H+ ions into the solution.
02

Case 2: A Br酶nsted-Lowry acid that is not an Arrhenius acid

In this case, we have a substance that donates a proton (H+) according to the Br酶nsted-Lowry definition, but we need to determine if it's possible for the substance not to increase the concentration of hydrogen ions (H+) in water (Arrhenius definition). An example of this is gaseous ammonia (NH3). It can act as a Br酶nsted-Lowry base by accepting a proton (H+), forming the ammonium ion (NH4+). But since ammonia doesn't increase the concentration of hydrogen ions in water, it is not an Arrhenius acid. Thus, it is possible for an acid according to the Br酶nsted-Lowry model not to be an Arrhenius acid.
03

Case 3: A Lewis acid that is not an Arrhenius or Br酶nsted-Lowry acid

According to the Lewis definition, an acid is a substance that can accept a lone electron pair from another substance (a base). This definition is broader than the other two and can include substances that do not release H+ ions or donate protons. An example of a Lewis acid that is not an Arrhenius or Br酶nsted-Lowry acid is the boron trifluoride (BF3) molecule. It can accept a lone pair of electrons from a base (e.g., ammonia, NH3) due to the empty orbital in the boron atom but it doesn't donate a proton (H+) or increase the concentration of hydrogen ions (H+) in water. Thus, it is possible for a Lewis acid not to be classified as either an Arrhenius or a Br酶nsted-Lowry acid.

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Key Concepts

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

Arrhenius Acid
Arrhenius acid refers to a substance that, when dissolved in water, increases the concentration of hydrogen ions \( H^+ \). This is one of the earliest definitions of acids in chemistry and is quite specific to aqueous solutions.
For example, hydrochloric acid \( \text{HCl} \) dissociates in water to produce \( H^+ \) ions and chloride ions \( \text{Cl}^- \).
The Arrhenius model is limited to substances in water but forms the basis of basic acid-base chemistry knowledge.
  • Focuses on the increase of \( H^+ \) ions in solution.
  • Specific to aqueous solutions, not applicable to non-aqueous environments.
  • It's foundational for understanding other acid definitions.
Br酶nsted-Lowry Acid
The Br酶nsted-Lowry definition broadens the concept of acids by describing them as proton donors. This means that a Br酶nsted-Lowry acid donates a hydrogen ion \( H^+ \) to a base.

Unlike the Arrhenius model, this definition is not limited to water as the medium. For example, hydrochloric acid \( \text{HCl} \) donates a proton to ammonia \( \text{NH}_3 \), forming ammonium \( \text{NH}_4^+ \) and chloride ions.
  • Accommodates reactions outside of aqueous solutions.
  • More versatile than the Arrhenius definition.
  • Focuses on the ability of a substance to donate \( H^+ \) ions.
Lewis Acid
The Lewis acid concept further expands what can be considered an acid by describing an acid as an electron pair acceptor.
Unlike the Arrhenius or Br酶nsted-Lowry models, Lewis acids are not limited to proton transfer or hydrogen ion production.

A classic example of a Lewis acid is boron trifluoride \( \text{BF}_3 \), which doesn't have any hydrogen ions to donate but can accept an electron pair from a base like ammonia.
  • Definition based on electron pair acceptance, not proton donation.
  • Includes a wide range of chemical reactions beyond aqueous or proton-based systems.
  • Helps explain acid behavior in non-traditional or organic chemistry.
Proton Transfer
Proton transfer is a central concept in acid-base chemistry, particularly in the Br酶nsted-Lowry model.
This process includes the movement of \( H^+ \) ions from an acid to a base.

By donating its proton, the acid transforms into its conjugate base, while the base accepting the proton becomes its conjugate acid. For instance, when hydrochloric acid \( \text{HCl} \) donates a \( H^+ \) to water, it forms \( \text{HO}^- \) and \( \text{H}_3\text{O}^+ \).
  • Fundamental to the Br酶nsted-Lowry model of acids and bases.
  • Enables understanding of how acids and bases transform into their conjugate pairs.
  • Can occur in various mediums, not just water.
Electron Pair Acceptance
Electron pair acceptance is key to understanding Lewis acid behavior.
In a chemical reaction, a Lewis acid accepts an electron pair from a base, forming a coordinate covalent bond.

For example, when boron trifluoride \( \text{BF}_3 \) bonds with ammonia \( \text{NH}_3 \), it accepts the lone pair electrons on nitrogen, creating a stable complex.
  • Broadens the scope of acid-base reactions beyond proton involvement.
  • Essential in understanding reactions in organometallic and coordination chemistry.
  • Offers a detailed view of molecular interactions and bond formation.

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Most popular questions from this chapter

Identify the Lewis acids,and bases in the following reactions. a. \(\mathrm{H}^{+}+\mathrm{OH}^{-} \rightleftharpoons \mathrm{H}_{2} \mathrm{O}\) b. \(\mathrm{Cl}^{-}+\mathrm{BCl}_{3} \rightleftharpoons \mathrm{BCl}_{4}-\) c. \(\mathrm{SO}_{3}+\mathrm{H}_{2} \mathrm{O} \rightleftharpoons \mathrm{H}_{2} \mathrm{SO}_{4}\)

Calculate the molarity of a solution of hydrobromic acid \((\mathrm{HBr})\) if 30.35 \(\mathrm{mL}\) of 0.1000 \(\mathrm{M}\) NaOH is required to titrate 25.00 \(\mathrm{mL}\) of the acid to the equivalence point.

In terms of ion concentrations, distinguish between acidic, neutral, and basic solutions.

Challenge Write an equation for a base equilibrium in which the base in the forward reaction is \(\mathrm{PO}_{4}^{3-}\) and the base in the reverse reaction is \(\mathrm{OH}^{-} .\)

Write equations for the salt hydrolysis reactions occuring when the following salts dissolve in water. Classify each as acidic, basic, or neutral. a. ammonium nitrate c. rubidium acetate b. potassium sulfate d. calcium carbonate
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