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Carbon Dioxide and Water
Name: Kelsey
Status: student
Grade: 9-12
Location: N/A
Country: N/A
Date: 6/4/2005
Question:
I am trying to use tapwater in one of my biology
experiments. I have noticed that whenever i use tapwater the pH
inexplicably increases. I have done some tests to confirm that it is in
fact the tapwater. The pH is originaly 7.2, and when i leave a beaker of
tapwater to sit for a 24hr period of time i check the pH and it has risen
to 8.2. That is a large increase considering I left the water alone and
didn't do anything to it. Would you possibly have an answer to why this
happens? Is it a common occurence? My hypothesis is that the chlorine
present in the water in the form of Hypo and hydrochloric acid evaporates
as it sits out in the open. Removing those 2 acids would explain the
increase of pH. I would appreciate any reply to this question.
Replies:
Kelsey-
The pH of pure water in the middle range (5-9) is an extremely sensitive
function of
any small impurities. It only takes 0.1ppm of acid to lower pH 7.0->6.0,
or 0.1pm of a weak base to raise it to 8.2 as you saw.
So your theory about losing HCl and/or HClO isn't bad.
But it could be anything else, too.
I would have expected your pH 7 to go to pH 5 due to absorption of
atmospheric CO2, making H+ + HCO3-.
Apparently there is a stronger base either in the tap water or in the air
where you left your beaker out.
Ammonia or similar amine vapors may be present in your kitchen.
Another possible source: oxidation reactions often use up acidity, as in:
H2O2 + 2 H+ + 2 e- --> 2 H2O ,
or ClO- + 2 H+ + 2 e- --> Cl- + H2O .
At pH 7 or so, HCl is highly dissociated into H+ and Cl-, which are firmly
trapped in the water,
so I'd guess it's not HCL that is leaving. HClO is a weaker acid, so it
might have more ability to leave.
Your water company tries to deliberately eliminate it's chlorine before
sending the water down the pipes to you,
so I doubt there was 1.0ppm HClO before any evaporation. But O.1ppm may be
plausible.
If 0.1ppm chlorine did remain,
and it found more dissolved dirt to oxidize while you watched,
a pH shift from 7 to 8 would be likely.
This would happen even under an air-tight lid.
Maybe this is why chemists use distilled or de-ionized water...
Jim Swenson
The "answer" to what is going on is trickier that might appear at first
glance. Tap water contains a certain amount of dissolved CO2. The amount
depends upon the "history" of the water from your tap. The amount of
dissolved CO2 is governed by Henry's Law, which states that P(CO2) = Kh *
C(CO2) where P(CO2) is the partial pressure of CO2 in the ambient air, Kh
is Henry's Law constant, and C(CO2) is the concentration of dissolved CO2
in the water. Now, the solubility of gases decreases with increasing
temperature. Stated another way the Henry's Law constant decreases with
increasing temperature. In addition, CO2 reacts with water: (CO2)aq + H2O
== H2CO3 == H(+1) + HCO3(-1) and the equilibrium constant for this
reaction is also a sensitive function of temperature. The net effect is
that there is less CO2 dissolved in water as the temperature increases.
I'm assuming that the temperature of the tap water is colder than ambient
conditions that equilibrate when you let the water stand for a while
(hence the pH increases due to the smaller concentration of CO2 present).
The multiple equilibria make a detailed quantitative calculation more
involved however. But the trend would be for the pH to increase, however.
One way to test this hypothesis would be to bring the tap water to a boil
for a minute or so, then cover the beaker and let it cool to ambient
temperature. Let another beaker just warm up from tap water temperature to
room temperature in the usual way. Heating the water should boil out any
dissolved CO2, and keeping the beaker covered would prevent it from
absorbing any CO2 from the atmosphere while equilibrating to ambient
temperature. The pH of the heated water beaker should start off higher
than the tap water sample, and may even remain higher than the tap water
sample. You could also check the result by breathing through both samples
with a soda straw, which will supply CO2 from your exhaled breath.
Whatever the results, congratulations on making the observation and
seeking an "explanation". I can't guarantee that the one I propose is
correct, but it works in the right direction. You can see from the attached
reference site that quantitatively verifying the result can get pretty
involved. See:
http://www.thuisexperimenteren.nl/science/carbonaatkinetiek/Carbondioxide%20in%20water%20equilibrium.doc
Vince Calder
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Update: June 2012
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