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14.1 Brønsted-lowry acids and bases  (Page 3/14)

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It is important to realize that the autoionization equilibrium for water is established in all aqueous solutions. Adding an acid or base to water will not change the position of the equilibrium. [link] demonstrates the quantitative aspects of this relation between hydronium and hydroxide ion concentrations.

The inverse proportionality of [h 3 O + ] and [oh − ]

A solution of carbon dioxide in water has a hydronium ion concentration of 2.0 × 10 −6 M . What is the concentration of hydroxide ion at 25 °C?

Solution

We know the value of the ion-product constant for water at 25 °C:

2 H 2 O ( l ) ⇌ H 3 O + ( a q ) + OH − ( a q ) K w = [ H 3 O + ] [ OH − ] = 1.0 × 10 −14

Thus, we can calculate the missing equilibrium concentration.

Rearrangement of the K w expression yields that [OH − ] is directly proportional to the inverse of [H 3 O + ]:

[ OH − ] = K w [ H 3 O + ] = 1.0 × 10 −14 2.0 × 10 −6 = 5.0 × 10 −9

The hydroxide ion concentration in water is reduced to 5.0 × 10 −9 M as the hydrogen ion concentration increases to 2.0 × 10 −6 M . This is expected from Le Châtelier’s principle; the autoionization reaction shifts to the left to reduce the stress of the increased hydronium ion concentration and the [OH − ] is reduced relative to that in pure water.

A check of these concentrations confirms that our arithmetic is correct:

K w = [ H 3 O + ] [ OH − ] = ( 2.0 × 10 −6 ) ( 5.0 × 10 −9 ) = 1.0 × 10 −14

Check your learning

What is the hydronium ion concentration in an aqueous solution with a hydroxide ion concentration of 0.001 M at 25 °C?

Answer:

[H 3 O + ] = 1 × 10 −11 M

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Amphiprotic species

Like water, many molecules and ions may either gain or lose a proton under the appropriate conditions. Such species are said to be amphiprotic    . Another term used to describe such species is amphoteric    , which is a more general term for a species that may act either as an acid or a base by any definition (not just the Brønsted-Lowry one). Consider for example the bicarbonate ion, which may either donate or accept a proton as shown here:

HCO 3 − ( a q ) + H 2 O( l ) ⇌ CO 3 2− ( a q ) + H 3 O + ( a q ) HCO 3 − ( a q ) + H 2 O( l ) ⇌ H 2 CO 3 ( a q ) + OH − ( a q )

Representing the acid-base behavior of an amphoteric substance

Write separate equations representing the reaction of HSO 3 −

(a) as an acid with OH −

(b) as a base with HI

Solution

(a) HSO 3 − ( a q ) + OH − ( a q ) ⇌ SO 3 2− ( a q ) + H 2 O ( l )

(b) HSO 3 − ( a q ) + HI ( a q ) ⇌ H 2 SO 3 ( a q ) + I − ( a q )

Check your learning

Write separate equations representing the reaction of H 2 PO 4 −

(a) as a base with HBr

(b) as an acid with OH −

Answer:

(a) H 2 PO 4 − ( a q ) + HBr ( a q ) ⇌ H 3 PO 4 ( a q ) + Br − ( a q ) ; (b) H 2 PO 4 − ( a q ) + OH − ( a q ) ⇌ HPO 4 2− ( a q ) + H 2 O ( l )

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Key concepts and summary

A compound that can donate a proton (a hydrogen ion) to another compound is called a Brønsted-Lowry acid. The compound that accepts the proton is called a Brønsted-Lowry base. The species remaining after a Brønsted-Lowry acid has lost a proton is the conjugate base of the acid. The species formed when a Brønsted-Lowry base gains a proton is the conjugate acid of the base. Thus, an acid-base reaction occurs when a proton is transferred from an acid to a base, with formation of the conjugate base of the reactant acid and formation of the conjugate acid of the reactant base. Amphiprotic species can act as both proton donors and proton acceptors. Water is the most important amphiprotic species. It can form both the hydronium ion, H 3 O + , and the hydroxide ion, OH − when it undergoes autoionization:

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Source:  OpenStax, Chemistry. OpenStax CNX. May 20, 2015 Download for free at http://legacy.cnx.org/content/col11760/1.9
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