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Chapter 14: Q4E (page 826)

Question: Show by suitable net ionic equations that each of the following species can act as a Bronsted-Lowry acid: (a) \(HN{O_3}\) (b) \(PH_4^ + \) (c) \({H_2}S\) (d) \(C{H_3}C{H_2}COOH\) (e) \({H_2}PO_4^ - \) (f) \(H{S^ - }\)

Short Answer

Expert verified

All the given six species act as a Bronsted-lowry acid as all these species donates a proton (\({H^ + }\)) to another molecule. The required equations are as follows.

Step by step solution

01

Define the concept of Bronsted-Lowry acid

The concept states that any compound that can transfer a proton to any other compound is an acid. In other words, the proton donor in a chemical reaction is a Bronsted-Lowry acid.

02

Suitable net ionic equations which shows that given species act as Bronsted-Lowry acid

So, it is observed that each of the species donates a proton (\({H^ + }\)) to another molecule. So. all the above species are Bronsted-Lowry acid.

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Most popular questions from this chapter

Why is the hydronium ion concentration in a solution that is \(0.10M\)in \(HCl\)and \(0.10M\) in \(HCOOH\)determined by the concentration of\(HCl\)?

We can ignore the contribution of water to the concentration of\(O{H^ - }\)in a solution of the following bases:\(0.0784 M {C_6}{H_5}N{H_2}\), a weak base\(0.11M{\left( {C{H_3}} \right)_3}\;N\), a weak base but not the contribution of water to the concentration of\({H_3}{O^ + }?\)

Explain why the neutralization reaction of a weak acid and a strong base gives a weakly basic solution

Are the concentrations of hydronium ion and hydroxide ion in a solution of an acid or a base in water directly proportional or inversely proportional? Explain your answer.

Explain why the ionization constant, \({K_a}\), for\({H_2}S{O_4}\) is larger than the ionization constant for \({H_2}S{O_3}\).
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