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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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