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Chapter 16: Problem 77

What are the Lewis definitions of an acid and a base? In what way are they more general than the Bronsted definitions?

Short Answer

Expert verified
Lewis theory defines an acid as an electron pair acceptor and a base as an electron pair donor. This is more general compared to the Bronsted-Lowry definition which only includes proton transfer, as Lewis's definition incorporates reactions outside the scope of proton exchange.

Step by step solution

01

Definition of an acid and a base according to Lewis

According to the Lewis theory, an acid is any substance that can accept an electron pair. A base, in contrast, is defined as any substance that can donate an electron pair. As such, a Lewis base donates an electron pair to form a bond, while a Lewis acid accepts an electron pair from the base, resulting in the formation of a covalent bond.
02

Definition of an acid and a base according to Bronsted

The Bronsted-Lowry theory defines an acid as a substance that can donate a proton (H+) and a base as a substance that can accept a proton. In essence, acid-base reactions, according to the Bronsted-Lowry definition, are essentially proton transfer reactions.
03

Comparison of both definitions

While the Bronsted-Lowry definition is restricted to substances that can donate or accept protons, the Lewis theory is more general in scope. The Lewis theory can encompass even those reactions where no protons are involved, such as the reaction between BF3 (boron trifluoride) and NH3 (ammonia) to form an adduct, BF3NH3, where BF3 is the Lewis acid (electron pair acceptor) and NH3 is the Lewis base (electron pair donor). This shows that the Lewis theory extends beyond just proton transfer reactions, making it more universally applicable and therefore broader in scope.

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Most popular questions from this chapter

Explain what is meant by the strength of an acid.

Which of the following is a stronger base: \(\mathrm{NH}_{3}\) or \(\mathrm{PH}_{3}\) ? (Hint: The \(\mathrm{N}-\mathrm{H}\) bond is stronger than the \(\mathrm{P}-\mathrm{H}\) bond. \()\)

Classify each of these species as a weak or strong acid: (a) \(\mathrm{HNO}_{3}\) (b) HF, (c) \(\mathrm{H}_{2} \mathrm{SO}_{4}\) (d) \(\mathrm{HSO}_{4}^{-}\) (e) \(\mathrm{H}_{2} \mathrm{CO}_{3},\) (f) \(\mathrm{HCO}_{3}^{-},(\mathrm{g}) \mathrm{HCl}\) (h) HCN, (i) \(\mathrm{HNO}_{2}\)

Classify these following oxides as acidic, basic, amphoteric, or neutral: (a) \(\mathrm{CO}_{2},\) (b) \(\mathrm{K}_{2} \mathrm{O},\) (c) \(\mathrm{CaO}\) (d) \(\mathrm{N}_{2} \mathrm{O}_{5},(\mathrm{e}) \mathrm{CO},(\mathrm{f}) \mathrm{NO}\) \((\mathrm{g}) \mathrm{SnO}_{2},(\mathrm{~h}) \mathrm{SO}_{3},(\mathrm{i}) \mathrm{Al}_{2} \mathrm{O}_{3}\) (j) \(\mathrm{BaO}\).

A \(0.400 M\) formic acid (HCOOH) solution freezes at \(-0.758^{\circ} \mathrm{C} .\) Calculate the \(K_{\mathrm{a}}\) of the acid at that temperature. (Hint: Assume that molarity is equal to molality. Carry your calculations to three significant figures and round off to two for \(K_{\mathrm{a}}\) .)
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