Uploaded by Jei-Kean Hernandez


“Acid-Base Equilibria” is very important not only in Chemistry but also in
related sciences such as biology, medicine, agriculture and geology. Topics
such as production of chlorophyll in plants, formation of limestone in caves,
deterioration of a marble statue in polluted air, and the activity of a lifesustaining enzyme can only be discussed intelligently if we understand acidbase behaviour.
Recalling the definitions of Acids and Bases
Early & far from ideal definitions:
Acids – substances that taste sour; turn litmus red
Bases – substances that taste bitter; turn litmus blue.
Note that not all compounds exhibiting these properties are chemically similar.
Arrhenius Theory of Ionization (1880s)
Acids – substances that dissociate into hydrogen ions and anions when
dissolved in water.
Bases – substances that dissociate into hydroxide ions and cations when
dissolved in water.
Bases react with acids to form a salt and water. The water produced in an acidbase reaction was considered to be the result of hydrogen ions combining with
hydroxide ions.
Bronsted-Lowry Concept (1923; England)
Acid - any substance capable of donating a proton in a reaction.
Base – any substance capable of accepting a proton in a reaction.
The B-L concept of acid-base (A-B) behavior has proven especially convenient
when using water and water-like (protophilic) solvents.
Role of the Solvent
One of the most significant features of the B-L concept of the A-B
behavior is the incorporation of the solvent as a vital part of the A-B system.
For example, when formic acid dissolves in water, ionization occurs because
the solvent acts as a base or proton acceptor.
H2O <==>
H3O+ +
(Reaction 1)
The H3O+ is merely a solvated proton called the hydronium ion.
When ammonia is dissolved in water, the solvent act as an acid or proton
H2O <==>
NH4+ +
(Reaction 2)
Amphiprotic solvents – solvents like water that can exhibit both acidic and
basic character, depending on the solute. Low MW alcohols and acetic acid are
other common amphiprotic solvents.
Exercise. Do this in five minutes in a one-half sheet of paper, take a photo &
send in my messenger account.
1) HCOOH + C2H5OH <==>
2) NH3 + C2H5OH <==>
Conjugate acids and bases
When an acid donates a proton in a reaction, it becomes a substance that is
capable of accepting a proton to re-form the original acid. Such a substance is
A B-L base.
When a base accepts a proton in a reaction, it becomes a substance capable of
donating a proton.
Bronsted & Lowry referred to such substances as conjugate pairs. Every B-L
acid has a conjugate base and every B-L base has a conjugate acid.
H2O +
H2O <==> H3O+ +
(Reaction 3)
acid1 & base1 are conjugates; acid2 & base2 are conjugates.
All amphiprotic solvents undergo self-ionization or autoprotolysis reactions. The
extent to which the reaction occurs is represented by the expression
Keq = [(aH3O+)(aOH-)] / (aH2O) = [(aH3O+)(aOH-)] / 1
Keq = Kw = [H3O+][ OH-]
(Equation 4)
or, taking the negative logarithms of both side
pKw = pH + pOH
(Equation 5)
autoprotolysis constant for water, Kw = 1.00 x 10-14 at 24oC (about room
According to Reaction 3, the concentrations of hydronium and hydroxide ions
in pure water must be equal and Equation 4 is used to calculate hydronium
ion concentration of pure water:
Kw = [H3O+]2
[H3O+]2 = 1.00 x 10-14
[H3O+] = 1.00 x 10-7 M
pH = -log (1.00 x 10-7) = 7.000
& pOH of pure water = 7.000
According to Equation 4,
an increase in the hydronium-ion concentration resulting from the addition of
an acid to water must be accompanied by an equal decrease in the hydroxideion concentration.
Conversely, an increase in the hydroxide-ion concentration must be
accompanied by a decrease in the hydronium ion concentration.
If the concentration of either H3O+ or OH- is known, the other can be calculated
using Eq. 4.
Exercise (Acid-Base Equilibria)
1. What is the concentration of H3O+ in a 1.00 x 10-2 M NaOH solution?
2. Calculate the pOH of a 0.020 M HCl solution.