An Introduction to pH
Acids and bases occupy a very important place in the study of chemistry. Industrial processes
that take place in aqueous solution and all life processes, which also take place in aqueous
solution, depend on maintaining close control of the concentrations of acids and bases in
solution. Acids are substances that produce H+ ions in aqueous solution and bases are substances
that produce OH- ions in aqueous solution. The following relationship exists between the
concentrations of H+ and OH- in all aqueous solutions (the brackets represent concentrations in
molarity units):
[H+][OH-] = 1.0 x 10-14
(Eqn. 1)
This relationship shows that as the concentration of H+ increases in a solution, the concentration
of OH- must decrease, since the product of the two concentrations must be 1.0 x 10-14. Neutral
solutions are aqueous solutions that are neither acidic nor basic, meaning that the concentrations
of H+ and OH- must be equal to one another. According to Eqn. 1 above, this
means that the concentrations of H+ and OH- are both equal to 1.0 x 10-7 M in a
neutral solution.
In 1909, the Danish scientist Søren Sørenson proposed a method for describing
the concentration of H+ in aqueous solutions that he called pH. Sørenson
defined pH as:
pH = -log[H+]
(Eqn. 2)
The pH scale allows us to express the often very small concentration of H+ in solution as a
simple number without having to resort to the use of scientific notation. We said previously that
the concentration of H+ in a neutral solution is 1.0 x 10-7 M. From Eqn. 2, we can see that a
neutral solution must have a pH value of 7.00.
pH = -log(1.0 x 10-7) = 7.00
(neutral solution)
Solutions in which the concentration of H+ is greater than 1.0 x 10-7 M (acidic solutions) will
have pH values less than 7.00, while solutions with H+ concentrations less than 1.0 x 10-7 M
(basic solutions) will have pH values greater than 7.00.
pH
< 7.00
7.00
> 7.00
[H+] (M)
> 1.0 x 10-7
= 1.0 x 10-7
< 1.0 x10-7
[OH-] (M)
< 1.0 x10-7
= 1.0 x 10-7
> 1.0 x 10-7
Description
acidic solution
neutral solution
basic solution
The solutions we commonly work with in General Chemistry have pH values that range between
1.0 and 14.0, although it is worth noting that pH values can be larger than 14.0 or smaller than
1.0. The following table shows the approximate pH values for some common aqueous solutions
arranged from the most acidic to the most basic.
Solution
Stomach acid
Lemon juice
Cola drinks, Vinegar
Black coffee
Blood, Milk, Tears
Milk of magnesia
Household ammonia, bleach
Approximate
pH
1.0
2.0
3.0
5.0
7.0
10.0
12.0
[H+] (M)
[OH-] (M)
1 x 10-1
1 x 10-2
1 x 10-3
1 x 10-5
1 x 10-7
1 x 10-10
1 x 10-12
1 x 10-13
1 x 10-12
1 x 10-11
1 x 10-9
1 x 10-7
1 x 10-4
1 x 10-2
It is important to understand that pH is a measure of how acidic or basic a solution is, not a
measure of the strength of the acid or base used to prepare the solution. For example, you could
prepare a solution with a pH value of 3.0 by using a small amount of a strong acid or a larger
amount of a weak acid. The pH value of a solution simply tells you the concentration of H+
present in the solution.
pH values are most commonly measured using an instrument called a pH meter. A pH meter
consists of an electrode capable of measuring very small voltages, an electronic circuit that
converts the voltage measurements to pH values, and a display device. The first pH meter was
invented by Arnold Beckman in 1934.
Beckman’s Original 1934 pH Meter
Modern pH Meter
Since strong acids and bases dissociate completely in water, we can calculate the pH of solutions
containing these substances. For example, for hydrochloric acid, we can write
HCl (g) + H2O (l) → Cl- (aq) + H3O+ (aq),
or, more simply:
HCl (aq) → Cl- (aq) + H+ (aq).
Thus, a 0.055 M solution of hydrochloric acid contains 0.055 M H+ and 0.055 M Cl- ions. The
pH of the solution is then:
pH = -log [H+] = -log[HCl] = -log(0.055) = 1.26
Similarly, the pH of a solution of a strong acid provides the concentration of the substance.
Suppose, for example, that the pH of a solution of perchloric acid is found to be 2.76.
pH = -log[H+]
10-pH = [H+]
[H+] = 10-2.76 = 0.00174 M = 1.74 mM
Perchloric acid, like all monoprotic (acids containing a single proton) strong acids, contributes
one mole of protons per mole of acid:
HClO4 (aq) → ClO4- (aq) + H+ (aq).
Therefore,
[HClO4] = [H+] = 1.74 mM.