Acid-Base Equilibria (1) Definitions of acids and bases

Edexcel A level Chemistry (2017)

Topic 12: Acid–-Base Equilibria:

Here are the first learning objectives:

12/1. To know that a Brønsted–Lowry acid is a proton donor and a Brønsted–Lowry base is a proton acceptor

12/2. To know that acid-base reactions involve the transfer of protons

12/3. To be able to identify Brønsted–Lowry conjugate acid-base pairs

Definitions of acid and base:

Bronsted and Lowry defined acids as proton donors.

The definition applies to all solvents, not just aqueous systems.

So         HCl(g)      +      H₂O(l)      ⟶    H₃O⁺(aq)     +      Cl^—(aq)

When hydrogen chloride gas dissolves completely in water, all its protons (H+ ions) bond to water molecules to form a hydrated proton or oxonium ion. 

Bronsted defined acids broadly so that in the above equation the HCl is an acid in the forward direction but H₃O+ is an acid in the reverse direction. 

Bronsted in other words saw acids and bases as pairs, strictly speaking conjugate pairs. 

So     HCl(g)       +       H₂O(l)        ⟶    H₃O⁺(aq)   +        Cl^—(aq)

         Acid                    base           conjugate acid   conjugate base

The chloride ion is the conjugate base of the acid HCl.

The water molecule is the base of the conjugate acid H₃O+ which introduces an immediate conceptual leap in our understanding for who thinks of water as a base!!

Then again consider this aqueous system:

NH₃(aq)    +       H₂O(aq)     ⇌      NH₄⁺(aq)   +       OH^—  (aq)

Base                   acid                    conjugate acid   conjugate base

Here ammonia accepts a proton from a water molecule so ammonia is a base and that means that water acts here as an acid!!! (whereas in the previous example it acted as a base). 

What’s more the ammonium ion is also an acid but in the reverse direction. 

And as this is an equilibrium reaction not all the ammonia protons are donated to water so ammonia is a weak base.

Within this topic, students can consider how the historical development of theories explaining acid and base behaviour show that scientific ideas change as a result of new evidence and fresh thinking.

Historical background:

Back in the day it was Lavoisier (1777) (of discovering oxygen fame)who first attempted a theory about acids.

Lavoisier—I guess as you might expect since it might be argued he was in love with the element—suggested that all acids contain oxygen. 

When he discovered oxygen he named the gas oxygen because the name means “acid producer”.

It fell to Sir Humphrey Davy (of miners safety lamp fame) to put a spanner in Lavoisier’s theory and tell us all that “oxymuriatic acid” is nothing but hydrochloric acid and contains no oxygen.

Davy (1816) instead was the first to suggest that all acids contained hydrogen not oxygen.

The first true attempt at a theory of acidic action came with Arrhenius (yes he of the rate equation fame!!).

Arrhenius (1887) defined an acid as a compound that could produce hydrogen ions in aqueous solution. 

Conversely an alkali was a compound that could produce hydroxyl ions in aqueous solution.  

This is about where you are as a student when you study acids at GCSE or High School level. 

And you may extend Arrhenius’ definition to say that strong acids, like hydrochloric acid, produce many hydrogen ions in aqueous solution—they ionise or dissociate completely.

At GCSE or in High School you might have shown how this means that a solution of hydrochloric acid exhibits a very high conductivity.  

Bronsted (1923) as we discussed above, extended the Arrhenius definition to apply to other solvents besides water such as liquid ammonia NH_3.

In liquid ammonia there are NH₄⁺  ammonium ions and NH₂^— amide ions in an equilibrium:

NH₃   +    NH₃    ⇌     NH₄+    +   NH₂^—

This self–ionisation is analogous to that of water:

H₂O    +    H₂O      ⇌     H₃O⁺     +    OH^—

And in both cases the water molecule and the ammonia molecule is amphiprotic in that it can act as both an acid and a base accepting and/or donating protons. 

Bronsted said that acid–base behaviour involves the transfer of protons but in any solvent not just water.

Finally 1938 G.N. Lewis (of dot and cross diagram fame!!!) extended the definitions of acid and base even further.

Lewis defined an acid as an acceptor of a pair of electrons and a base as a donor of a pair of electrons. 

This broadens the definitions even further.

Not only is a hydrogen ion an acid but in this reaction below BF₃ is an acid:

H₃N:    +       BF₃     ⇌   H₃N⟶BF₃

base           acid

Since BF₃ is deficient of an electron pair in its outer shell is accepts the lone pair from the nitrogen atom. 

A dative covalent bond results between the two species.

Can you pick out then the acids and their conjugate bases in these equations?

HSO₄^—     +   H₂O   ⇌  H₃O+    +    SO₄ ^2—

CH₃COOH     +    H₂O     ⇌  H₃O+      +      CH₃COO^—