4.7: Acid-Base Equilibria
 show that you understand the ideas about acidity developed in the nineteenth and twentieth centuries
from generalisation about substances with a sour taste to the theory that acids produce an excess of
hydrogen ions in solution (Arrhenius) and then to the Brønsted–Lowry theory
 show that you understand that according to the Brønsted–Lowry theory:
i) an acid is a proton donor
ii) a base is a proton acceptor
iii) acid–base equilibria involve the transfer of protons
 show that you understand the Brønsted–Lowry theory of acid–base behaviour, and are able to identify
conjugate acid–base pairs
Historical
Jabir ibn-Hayyan(c722-c815)
The mineral acids that we know today were discovered by the Arab
alchemist. His pioneering experimental techniques and inventions of
equipment are still commonly used in laboratories. He used
crystallization and distillation to discover sulphuric, hydrochloric and
nitric acid. He also studied ethanoic acid in vinegar and tartaric acid
in wine.
Antoine Lavoisier(1743-1794) and Humphry Davy
Lavoisier, a French nobleman, laid the foundations of modern
chemical theory. His experiments led to the oxygen theory of burning.
The name oxygen means ‘acid former’ he gave the gas this name
because he thought that all acids contained oxygen. However this
was disproved by Davy(1809-1810) when he heated hydrogen
chloride to high temperatures with a range of metals and non metals.
He could find no trace of oxygen in the compound. After further
work(1816) he proposed that all acids have hydrogen in common!
Svante Arrhenius
In 1884 the young Swede in his 20’s wrote a doctoral thesis
proposing that some compounds are ionized in solution all the time.
He was bitterly disappointed when he was awarded the bottom
grade for his paper! In 1903 he was vindicated and awarded the
Nobel prize for chemistry.
He used his ionic theory to explain why all acids have similar
properties when dissolved in water. His theory could account for
what happens when an acid is neutralized by an alkali and could
explain the difference between a strong and weak acid.
In 1887 he defined an acid as a compound that could produce
hydrogen ions when dissolved in water and an alkali as a compound
that could produce hydroxide ions in water.
HA(aq) ↔ H+(aq) +A-(aq)
and an alkali as a compound that could produce hydroxide ions in
water.
B(aq) +H2O(aq)
↔
BH+(aq) + OH-(aq)
According to his theory:
Hydrochloric acid is a strong acid which is fully ionized when
dissolved in water.
HCl(aq)
H+(aq) +
Cl-(aq)
Ethanoic acid is a weak acid which is only slightly ionized
CH3COOH(aq) ↔
CH3CO2-(aq) + H+(aq)
His theory could account for the typical reactions of dilute acids in
water eg
Acids + metals
salt + hydrogen
Acids + carbonates
Acids + bases
salt + carbon dioxide + water
salt + water
When exactly equal volumes of hydrochloric acid and sodium
hydroxide of the same molar concentrations are mixed together they
form a solution that has no effect on litmus paper. We say that the
acid has neutralized the base or vice versa.
According to Arrhenius’ theory neutralization occurs because the
number of hydrogen ions (H+) is exactly equal to the number of
hydroxide ions(OH-) and the two react with each other to form
water(H2O)
H+ (aq) +
OH-(aq)
H2O(aq)
His theory is still useful but limited to aqueous solutions.
The oxonium ion (H3O+)
Once we had a greater understanding of the structure of the atom it
was inconceivable that the hydrogen ion could exist independently.
The H+ ion is really a proton with a diameter 70 000 times smaller
than a Li+ ion!
It was suggested that a H+ ion exists in association with a water
molecule as the H3O+ ion, the oxonium ion.
When talking about acids and bases we sometimes refer to
hydrogen ions as protons we should refer to them as oxonium ions
and write H3O+ instead of H+.
The importance of water in the behavior of acids was recognized from
observations of hydrogen chloride(HCl) and ethanoic acid(CH3COOH)
dissolved in organic solvents….they do not conduct electricity and do
not effect dry litmus paper….since the solution does not contain H+
ions.
The Bronsted-Lowry Definition of acids and bases
The preferred theory for discussing acid base equilibria was put
forward by the Danish chemist Johannes Bronsted and the English
chemist Thomas Lowry (1923). They described acids as proton
donors and bases as proton acceptors.
In aqueous solution ammonia (a base) and hydrogen chloride (an
acid) react to form a solution of ammonium chloride (a salt).
However the reaction between ammonia and hydrogen chloride does
not need water or any other solvent to happen. The white fumes are
tiny crystals of ammonium chloride, formed as ammonia gas and
hydrogen chloride react together.
NH3(g) + HCl(g)
NH4Cl(s)
This is clearly the same reaction as the one that occurs in water
and ought to be an acid base reaction. However the Arrhenius
definition does not allow us to do this as there is no reaction
between H3O+ and OH-.
From the Bronsted- Lowry definition the HCl can be seen as a
proton donor and the NH3 molecule as a proton acceptor.
NH3(g) + HCl(g)
NH4+Cl-(s)
Acid or Base
Classically described by their taste as sour…..acid and soapy
feel…..bases. The taste and feel test are not recommended as strong
acids and bases can cause considerable harm!
Some substances can act as acids and bases eg: water
It acts as an acid when it reacts with substances like ammonia
donating a proton during the reaction:
H2O(l) + NH3(aq)
OH-(aq) + NH4+(aq)
It behaves as base when it reacts with hydrogen chloride accepting
a proton in the reaction
H2O(l) + HCl(aq)
H3O+(aq) + Cl-(aq)
Acids as proton donors
According to Bronsted-Lowry, hydrogen chloride molecules give
hydrogen ions (protons) to water molecules when they dissolve in
water producing hydrated hydrogen ions called oxonium ions. The
water acts as a base.
Hydrogen chloride is a strong acid…..it readily gives up its protons
to water molecules and its equilibrium lies well to the right
effectively making it completely ionized in solution. Other strong
acids are sulphuric and nitric acid.
Ethanoic acid is a weak acid it is not completely ionized in water an
equilibrium is established
CH3COOH(aq) + H2O(aq) ↔
CH3COO-(aq) + H3O+(aq)
Bases as proton acceptors
According to Bronsted-Lowry, a base is a molecule or ion that can
accept a hydrogen ion from an acid. A base has a lone pair of
electrons which can form a dative covalent bond with a proton.
An ionic oxide such as calcium oxide reacts completely with water
to from calcium hydroxide . The calcium ions do not change; but
the oxide ions, which are powerful proton acceptors, all take the
protons from the water molecules. An oxide ion is a strong base as
well as hydroxide ions, ammonia, amines, carbonates and hydrogen
carbonate ions.
Conjugate acid-base pairs
From the Bronsted Lowry definition when an acid dissociates we
can consider it as a acid-base reaction which is in equilibrium
HA(aq) + H2O(aq) ↔ H3O+(aq) +A-(aq)
In the forward reaction HA acts as an acid, donating a proton to a
water molecule. Water accepts a proton and acts as a base. In the
reverse reaction, the H3O+ acts as an acid, donating a proton to Awhich in turn is acting as a base.
In acid-base reactions it is always possible
to find 2 acids:
HA and H3O+
and 2 bases:
H2O and AIn each case, the acid on one side of the equation is formed from
the base on the other side of the equation. They are called acid-base
conjugate pairs.
HA is the conjugate acid of A-, which means that A- is the conjugate
base of HA etc.
NH4+(aq) + H2O(aq) ↔
A1
B2
NH3(aq) + H30+(aq)
B1
A2
The equilibrium involves 2 conjugate acid-base pairs.
The p H scale