The pH of water rarely indicates how much
acid or alkali is needed to change the pH.
For example:
It is not uncommon to have 2 different samples of water of ‘equal’ pH
where one requires 4 times more pH adjustment than the other!
This phenomenon is due to the concentrations
of ‘bicarbonate’ and ‘carbon dioxide’ present in the water. It is
particularly pronounced with bore waters.
Bicarbonate
Bicarbonate (HCO3-) is
alkaline and therefore elevates pH. Its concentration is normally
expressed as “alkalinity”. It is one of the main factors causing pH to
rise in nutrient solutions and also confuses growers in their attempt to
estimate how much pH Down will be required to lower pH.
Unlike ‘hydroxide’ (i.e. common ingredient
for pH Up), bicarbonate is only weakly alkaline and therefore unable to
elevate pH above ~8.3, regardless of its concentration. As a consequence
of this, unlike hydroxide, bicarbonate has a strong pH buffering capacity
which means it resists pH change when acid (pH Down) is added. For
example, a weak solution of hydroxide can have a pH of 14 whereas a
bicarbonate solution 10 times more concentrated has a pH lower than 8.3!
Now, the interesting fact is, to lower the pH down to 4.5 the bicarbonate
solution requires ~10 times more acid than the hydroxide solution - even
though its initial pH was so much lower.
Hence the presence of bicarbonate is
deceiving because unlike hydroxide it is not detectable from pH readings
and is only noticeable once you attempt to lower the pH.
Carbon dioxide
Ever wondered why pH fluctuates (i.e.
typically upwards) after it is lowered? This behaviour is actually a
consequence of adjusting the pH. Lowering pH via adding acid, removes
bicarbonate and produces carbon dioxide. The presence of this free (i.e.
uncombined) carbon dioxide (CO2) tends to lower the pH because it reacts
(only weakly) with water to form carbonic acid. However, CO2
concentrations above about 0.5 mg/L in water are unstable when such waters
are exposed to the atmosphere (at sea level pressures). Under that
condition CO2 in excess of 0.5 mg/L will slowly escape from the water into
the atmosphere. Consequently this loss of acidity causes a corresponding
rise in pH.
This subsequent rise in pH is particularly
noticeable with ground waters (i.e. bore water) which typically have CO2
contents around 50 - 200 mg/L (due to biological activity within the
aquifer). When these waters are pumped to the surface, the pH rises with
time because the excess (acidic) CO2 gradually escapes (Fig 1.16). The pH
will then rise to a stable value solely dependent on the water's
bicarbonate content.
Example: A bore water with 100 mg/L
bicarbonate and 100 mg/L of free CO2 will have an initial pH of 6.3. Its
pH will gradually rise to 8.2 after it has been exposed to the atmosphere
for sufficient time to allow the CO2 content to drop to around 0.5 mg/L.
The same phenomenon (although to a much
lesser extent due to lower CO2 contents) can occur with scheme (tap)
water. Thus the conclusion – because the pH of waters is only stable after
aeration, it is only the "after aeration" pH value that has any
interpretative significance. To determine that value, aerate the water by
tumbling a sample of it from one container to another, 30-40 times prior
to measuring its pH.
Conclusion: Interpret pH values
with caution because a water with a lower pH than another may produce the
higher pH after both are aerated!!
