Looks like it’s time to teach some chlorine chemistry. I was really hoping it wouldn’t come to this because the vB coding is gonna be a real pain. But first I’ll present the data I dug up about the long lasting chlorine residual.
It turns out we ran measurements on five sets of water: capped in the refrigerator, open in the fridge, capped at room temp. in the light, open at room temp. in the light, and capped in a dark cupboard.
All five sets began with a Free Chlorine Residual of 0.72 mg/l (mg/l is the same as parts per million, ppm). Measurements were taken every day for the first two weeks, and every few days after that.
The jar that was left open at room temperature and in the light lost its chlorine the quickest. Between day 6 and 7 half (0.36ppm) of the residual remained, and it was undetected at day 14.
The next to go was the jar that was capped and kept at room temp. in the light. It was half depleted at day 7 and was finally undetected at day 26.
Then came the jar that was open and refrigerated. It was half depleted around day eleven and undetected at day 20.
Next was the jar that was capped and kept in the dark at room temp. It was half depleted at day nine, but made it all the way out to day 42 before finally being indetectable. Even at day 34 the residual was still above 0.1ppm.
And lastly, the water that was capped and kept in the dark refrigerator still had a 0.45ppm residual at 48 days. At that point it was pretty clear the chlorine was going to stick around for a good while yet.
Conclusions: having a lid prevents the residual from leaving the water, and light (specifically UV light) speeds up its degradation.
Now it’s time for school. Rather than address other posts point by point, it’ll go more smoothly if I just tell it like it is. I’ll still try and make it short.
Take it as fact that chlorine gas dissolves in water very easily. Chlorine gas exists as a dimer, i.e. two atoms of Cl connected. Cl[sub]2(g)[/sub] (g for gaseous). In water, it becomes Cl[sub]2(a)[/sub] (a for aqueous) but not for long. Water is plenty polar enough to split it into spearate ions, and chlorine splits unevenly into Cl[sup]+[/sup] and Cl[sup]-[/sup].
Cl[sup]-[/sup] is now an insignificant ion that doesn’t amount to much and is balanced out by other positive ions in solution, but his brother with a full positive charge is pretty reactive; so he snatches a nearby H[sub]2[/sub]O, boots off an H[sup]+[/sup] and hooks up with the remaining OH[sup]-[/sup] and becomes HOCl, which is known as hypochlorous acid. (Note here that H[sup]+[/sup] is acid and HO[sup]-[/sup] is a base. The reaction created one of each, but tied up the base…so our solution just got more acidic. pH is important a little later.)
Hypochlorous acid is the main player in the disinfection game. He is very good at disabling bacteria. But HOCl can only exist as such in a certain pH range, spanning from pH of less than one up to pH 7.5. If the solution gets less acidic and the pH approaches 7.5, HOCl will begin to shed his H[sup]+[/sup] and exist as OCl[sup]-[/sup], known as Hypochlorite ion. (Incidentially, this anion is the negative half of Sodium Hypochlorite and Calcium Hypochlorite (HTH), which are other disinfectants.) The trouble is, hypochlorite ion is 80 to 100 times less effective of a disinfectant. So what we learn here is the importance of maintaining control of pH. Otherwise we’d have to use 100x as much chlorine gas to get the same level of disinfection.
Big deal, huh? Well, yeah. It would be all too easy to maintain a nice low pH in the water treatment process to insure all the Cl[sub]2[/sub] gas we added gave us the maximum bang for the buck in the form of HOCl. In addition to the Chlorine + water reaction producing H+, other chemicals used in conventional water treatment create more H+ ions driving the pH down even further. The problem is twofold: customers don’t like water that burns their throat; and even if they did, the pipes that carry the water still hate it. Acidic water is very corrosive to pipes, and in just a few short years, the whole distribution infrastructure would be a leaky mess. So treatment plants try to make the water non corrosive (and even slightly scaly, to build a protective coating inside the pipes) by keeping the pH a little above 7.0. So a pH range of 7.0-7.5 doesn’t leave a whole lot of room for errors in the treatment process, especially since the quality of the source water can change quite rapidly.
So much for being brief. I’m like that when no one is in the room to stop me. Or flee.
[sub](Just wondering if anyone else found this ironic.)[/sub]