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Chemical Analyses: Water: pH
pH is the most frequently used water chemistry test. Practically
every phase of water chemistry, including, acid-base neutralization,
metal ion solubility, wastewater treatment and corrosivity is pH
dependent. pH is a measure of the balance between the concentration
of hydrogen (H+) ions and hydroxyl (OH-) ions in solution. In acidic
conditions, there are more H+ ions than OH- ions (pH < 7). In alkaline
(basic) conditions, there are more OH- ions than H+ ions (pH > 7).
When conditions are neutral, the concentration of H+ and OH- ions
are equal (pH = 7). pH is defined as -log[H+], where [H+] is defined
as the H+ concentration. pH values are on a scale from 0 to 14.
This scale is based on the product of the H+ and the OH- concentrations
which is equal to 10-14 M2, the equilibrium constant, K:
[H+] [OH-] = 10-14 M2
M represents molarity which is molar concentration (moles of solute
/ liters of solution). Where does that number 10-14 M2 come from
anyway? It is based on the theory of equilibrium and the energy
required for reactions to take place. Equilibrium is based on chemical
reactions (such as the dissociation of H2O):
H2O <–-> H+ + OH-
which are reversible and proceed until the forward and backward
reactions occur at the same rate. The reaction is then said to be
at chemical equilibrium, and no further change in the concentration
of products and reactants occurs. At equilibrium, the chemical reaction
requires no energy to be maintained. Movement away from equilibrium
is nonspontaneous; which means that an outside energy source is
required. The actual numbers that are derived as equilibrium constants
(i.e. 10-14 M2 for the dissociation reaction of H2O) originate from
the concept of the energy required for a reaction to proceed depending
on how far the reaction is away from equilibrium (Campbell,
1996).
For a more in depth understanding of the relationship between equilibrium,
energy and the equations that the equilibrium constant, K, is based
on, see Stumm
and Morgan (1996).
In a neutral solution, [H+] = 10-7 M and [OH-] = 10-7 M, and the
product is 10-14 M2 (10-7 M * 10-7 M). If enough acid is added to
a solution to increase [H+] to 10-5 M, then [OH-] will decline by
an equivalent amount to 10-9 M (10-5 M * 10-9 M = 10-14 M2). Whenever
we know the concentration of either H+ or OH- in a solution, we
can deduce the concentration of the other ion (Campbell,
1996).
Because the H+ and OH- concentrations of solutions can vary by
a factor of 100 trillion or more, scientists developed a way to
express this variation more conveniently, by use of the pH scale,
which ranges from 0 to 14. By taking the negative log of the [H+]:
pH = -log [H+],
pH will always be between 0 and 14, which gives an easier scale
for understanding the balance between H+ and OH- ions in solution.
It is important to remember that each pH unit represents a tenfold
difference in H+ and OH- concentrations. It is this mathematical
feature that makes the pH scale so compact. A solution of pH 3 is
not twice as acidic as a solution of pH 6, but a thousand times
more acidic. When the pH of a solution changes slightly, the actual
concentrations of H+ and OH- in solution change substantially (Campbell,
1996).
pH measurement is conducted using a standard hydrogen electrode
and a reference electrode in order to determine the activity (concentration)
of the hydrogen ions by potentiometric measurement. Modern electrodes
combine the reference and standard electrodes into one unit. The
electrode is calibrated over the measuring range using certified
standard buffers standardized at specific temperature prior to measuring
the sample. The temperature of the sample should be the same as
the buffers; otherwise the temperature of the sample should be recorded.
Many pH meters have a temperature dial to adjust the potentiometric
slope to compensate for temperature (MEND,
2001). For specifics on the pH potentiometric method, see section
4500-H+ (Standard
Methods, 1998). In addition, there are numerous color indicators
for pH, which are useful when accuracy requirements of 0.5 pH units
are acceptable (Kirk-Othmer,
1995). The most common is litmus paper which turns red in acidic
solutions and blue in alkali solutions. Color indicators are a rapid
trouble free method to obtain an approximate pH value.
Chemical
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