pH - The Acid Test
Untitled
Within a cell, dynamic balance (homeostasis) is maintained when essential chemicals
exist in sufficient concentrations that vital cellular functions and structures are preserved.
Small changes in chemical concentrations can cause major shifts in cellular function and
structure. The concentration of hydrogen ions within the cytoplasm is a particular case
in point.
The concentration of hydrogen ions in solution is expressed as the pH. Maintaining
pH within narrow tolerance limits is vital for the functioning of the cells, tissues, organs,
and systems. Within the aqueous environment of the cell, molecules may separate into charged
fragments or particles called ions. This separation is termed dissociation or ionization.
One of the fragments often produced from such dissociation is a hydrogen ion (H+).
The pH of a solution is a measure of the concentration of hydrogen ions in that solution.
This is often written as [H+] where [ ] means "concentration of." The
pH is often described as a measure of the alkalinity or acidity of a solution. While this
usage is very common it is not technically correct.
"The tendency to link pH to acidity is understandable, but not justified since
acidity and pH are defined very differently. The pH is equal to the negative logarithm of
the molar hydrogen ion concentration: pH = -log[H+]. Acidity is properly defined
as the concentration of acid determined by titration with a strong base to a specified end
point (often the phenolphthalein end point). Acidity refers to the concentration of titratable
acid, but not to the H+ concentration produced by that acid."
Hydrogen ionizes when a hydrogen atom that is covalently bonded to the oxygen of one
water molecule leaves its electron behind. Then, as a hydrogen ion (H+), it bonds
with a different water molecule (H+ + H2O ® H3O+). This reaction produces two ions, a hydroxide ion (OH-) and a hydronium ion (H3O+). The reaction may be expressed as: 2H2O _ H3O+ + OH-.
We usually express the ionization of water in simpler form as H2O _ H+ + OH-.
In pure water, or in any water-based solution, a small but constant number of water molecules
become ionized. As would be expected, the number of H+ ions exactly equals the
number of OH- ions when pure water ionizes. The molar concentration or moles per
liter (M) of hydrogen ions equals 1 x 10-7 in pure water. The hydroxide ion molar
concentration is also 1 x 10-7. The product of the molar concentrations of these
two ions in pure water is 1 x 10-14. These numbers can also be expressed as base 10
logarithms. Thus the log of 1 x 10-7 is -7. pH is defined, as noted earlier, as
the negative logarithm (-log10) of the molar hydrogen ion concentration in a
solution. By using the negative logarithm, the numbers, and thus the scale on which pH is
expressed, is positive. Thus, it can be expressed as -log(1 x 10-7) = +7.
The "p" in pH stands for the "negative logarithm of." Since the product
of [H+] and [OH-] is always constant (1 x 10-14) the pOH
also equals 7 in pure water where pOH = -log [OH-]. Consider what will happen if
a substance added to a volume of water ionizes and releases more H+ into the
solution. The relative H+ concentration increases and the relative OH-
concentration decreases. The total molar concentration of H+ and OH-
remains constant at 1 x 10-14.
We can use our knowledge of pH and pOH to come up with the pH scale. Since
pH + pOH = 14 then it follows that pH = 14 - pOH and pOH = 14 - pH. For pure water the
pH = 7 and the pOH = 7. Thus a solution of pH = 7 is neutral and at the midpoint of the
scale. Solutions with pH values lower than 7 are said to be acidic. In an acidic
solution, the number of H+ ions exceeds the number of OH- ions.
Solutions with a pH above 7 are basic or alkaline: the number of OH-
ions exceeds the number of H+ ions. In other words, the more acidic a solution,
the higher the number of H+ ions and the lower the number of OH- ions.
The more basic a solution, the lower the number of H+ ions and the higher the
number of OH- ions. Remember that the pH scale is logarithmic, not arithmetic.
If two solutions differ by 1 pH unit, then one solution has 10 times the hydrogen ion
concentration of the other. While this may be logically correct it is not often correct
under natural conditions. The "one-for-ten" rule is true only for strong monoprotic
acids such as HCl which ionize com-pletely. It does not hold for weak acids such as acetic
acid (CH3COOH) or carbonic acid (H2CO3) which do not ionize
fully. In addition it should be noted that almost all pH calculations are based on the standard
temperature reference of 25° Celsius (298 Kelvin) and water
at lower or higher temperatures will dissociate less or more than at the standard reference
point.
In common usage, an acid is a substance that causes an increase in the number
of H+ ions and a decrease in the number of OH- ions in solution. This
increase is most often the result of an ionization that produces H+. Some common
acids and their ioni-zation products are:
hydrochloric acid HCl ® H+ + Cl- strong acid 100% ionized
acetic acid CH3COOH _ H+ + CH3COO- weak acid
carbonic acid H2CO3 _ H+ + HCO3- weak acid
phosphoric acid H3PO4 ® H+ + H2PO4- ® H+ + HPO4-2
The more completely the acid ionizes, the more H+ is released, and the
stronger the acid.
A substance need not give up hydrogen ions itself to cause an increase in the
[H+] in a solution. An important molecule in biological systems is CO2
(carbon dioxide) which combines with water to form carbonic acid.
CO2 + H2O « H2CO3 « H+ + HCO3-
This reaction has been considered the basis for assuming a pH of normal rain at 5.6.
However, even trace amounts of natural gases such as ammonia (NH3) can sharply
alter rain pH. It takes an air concentration of only 0.0039 parts per million of NH3
to raise the pH of pure water by 1.4 units, whereas the same size decrease in pH requires
an air-CO2 level of 320 ppm. Thus NH3 is roughly 80,000 times as
effective as CO2 in altering the pH, and an extremely small amount of ammonia
in normal rain will produce a pH significantly higher than 5.6.
The reaction of SO2 (sulfur dioxide) with atmospheric water is, in part,
responsible for acid rain.
SO2 + H2O « H2SO3
(sulfurous acid) « H+ + HSO3-
If the water droplets contain oxidants such as oxygen, ozone, or hydrogen peroxide and
trace amounts of certain metals which act as catalysts, the SO2 can rapidly oxidize
to SO3. The sulfur tri-oxide hydrolyzes to form sulfuric acid (H2SO4),
a strong acid.
A base is a substance that causes a decrease in the concen-tration of
[H+] in solution and an increase in [OH-]. In many cases this is
achieved by the ionization of the molecule to produce OH- (hydroxyl ion) which
not only adds to the OH- in solution but also removes H+ from solution
by combining with it to form water, thus raising the pH.
KOH + HCl « K+ + OH- + H+ + Cl- « KCl + H2O
Thus, neutralization of a strong acid by a strong base produces a salt
(an ionic compound composed of a negative ion from an acid and a positive ion from a base)
and water. Some common bases that ionize to produce OH- are:
sodium hydroxide NaOH ® Na+ + OH-
magnesium hydroxide Mg(OH)2 ® Mg+2 + 2 OH-
potassium hydroxide KOH ® K+ + OH-
These are all strong bases because they ionize completely in solution. Ammonia (NH3)
dissolved in water, is also basic; it can remove H+ from solution: (NH3 + H2O ® NH4+ + OH-) Here water acts as an acid and NH3 acts as a base. Water is amphoteric (amphipro-
tic) in that it can function as an acid (in the presence of a strong base) and as a base
(in the presence of a strong acid).
The bicarbonate ion (HCO3-) is also basic: it, too, can accept
H+:
H+ + HCO3- « H2CO3
This is an important reaction in maintaining proper blood pH.
pH EXAMPLE
0 -|- 1 N HCl
|
|
1 -|- Gastric fluids
|
|
2 -|-
|- Lemon Juice, Vinegar
|- Acid Mine Drainage
3 -|- Peaches, Rhubarb, Pineapple,
|- Ginger Ale, Coke
|
4 -|- Acid Soil, Tomatoes
|- Tomato Juice
|- Black Coffee
5 -|- Cheese, Cabbage, Beans
|
|
6 -|- Peas, Corn, Salmon, Shrimp,
| Milk Begins to Taste Sour
|- Milk
7 -|- Pure Water
|- Blood
|- Bile
8 -|-
|- Sea Water
|
9 -|- Very Alkaline Natural Soils
|- Borax
|
10 -|- Alkaline Lakes, Soap Solutions
|- Milk of Magnesia
|
11 -|- Household Ammonia Concentrate
|
|
12 -|-
|- Lime, (Saturated Solution)
|
13 -|- Oven Cleaner
|
|
14 -|- Lye, 1 N NaOH
BUFFERS
Physiological processes require that pH remain relatively constant. The pH of blood
in our bodies, particularly in the arterial system, is usually maintained between 7.3 and 7.5.
However, blood returning to the heart contains CO2 picked up from the tissues
(recall that CO2 combines with water to form carbonic acid), and our diets as well
as the normal metabolic reactions in cells may contribute an excess of hydrogen ions. The pH
must be kept constant by several buffer systems.
A buffer is defined as a solution that resists change in pH when small amounts
of acid or base are added. Bicarbonate, phosphate, and protein buffer systems maintain our
blood pH. We will use the phosphate buffer system as an illustration.
A buffer is made by mixing a weak acid with its salt in order to have in solution
something that can act as an acid (give up hydrogen ions) and something that can act
as a base (accept hydrogen ions). In a potassium phosphate buffer, the weak acid
H2PO4-, is supplied as KH2PO4
(monobasic potassium phosphate) and its salt as K2HPO4 (dibasic potassium phosphate).
At equilibrium, these substances are ionized to some degree:
KH2PO4 _ K+ + H+ + HPO4-2
K2HPO4 _ 2 K+ + HPO4-2
If hydrogen ions are added to the solution, they can be picked up by
HPO4-2, which acts as a base:
H+ + HPO4-2 « H2PO4-
If hydroxyl (OH-) ions are added to the solution, they can be picked up by
H+:
OH- + H+ « H2O
In summary:
OH-
®
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/ \
H2PO4- HPO4-2 + H2O
\ /
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H+
Email: emmeluth@gateway.net
