Just as you say, chlorine dissolved in water likes to be Cl- or ClO- anions.
My suspicion after more thought is that dissolving Cl2 in water produces excess metastable ClO, which breaks down into Cl over a few hours, and that out-of-equilibrium ClO is the antimicrobial and also a reasonable chlorinating agent to be concerned about.
Cl- and ClO- are two different things, the latter is an oxidant while the former is not. It seems odd to me to bundle them together.
I don’t know what particular oxidants would cause more or worse biological effects. So, not sure whether it would matter if there’s “excess metastable” of one particular oxidant or not. But, “excess metastable ClO” seems an odd thing to expect—it sounds like you’re expecting a reaction to go past equilibrium, why?
Cl- and ClO- are two different things, the latter is an oxidant while the former is not. It seems odd to me to bundle them together.
Truly, shame on bhauth three comments above. (sarcasm) I figure we’re just mentioning them together because they’re what you get when reacting chlorine gas with water.
(Note I meant “excess metastable ClO” to be parsed as “an excess of metastable ClO,” not as if “excess metastable” was an extra fancy kind of excess—un excès métastable. Although I guess that also kind of works.)
it sounds like you’re expecting a reaction to go past equilibrium, why?
I do not. So maybe I should explain the three steps I’m thinking of:
State 1 is water and chlorine gas. This is out of equilibrium—the chlorine will react with the water.
State 2 is water, H3O+, and a somewhat-even mixture of Cl- and ClO- ions (and trace Cl2). This is also out of equilibrium, the ClO- will react with the H3O+. But it’s less out of equilibrium than state 1.
State 3 is water, H3O+, Cl-, and trace amounts of ClO- (and even tracer Cl2). This is the putative equilibrium.
I expect treated water to go 1->2->3. Not 1->3->2->3 or any such thing. If it just goes 1->3, never 2 at all, then I’m wrong. Also, it if goes 1->2->3 but 1->2 is slow and 2->3 is fast, then I’m wrong.
When dissolved in water, chlorine converts to an equilibrium mixture of chlorine, hypochlorous acid (HOCl), and hydrochloric acid (HCl):
Cl2 + H2O ⇌ HOCl + HCl
In acidic solution, the major species are Cl2 and HOCl, whereas in alkaline solution, effectively only ClO− (hypochlorite ion) is present. Very small concentrations of ClO2−, ClO3−, ClO4− are also found.[18]
So the putative equilibrium is the above (and also including some H3O+ and ClO− and Cl− from dissociation of the stuff on the right) and not this:
State 3 is water, H3O+, Cl-, and trace amounts of ClO- (and even tracer Cl2). This is the putative equilibrium.
Note that the total quantity of Cl2 and HOCl (or in more basic solution ClO− ) is conserved in the above reaction. You do not get to a point where both are trace, if that’s not what you started with. In natural stomach acid, you would presumably not start with either of them, but in chlorinated tap water you do.
Regarding the specific intermediate steps in a reaction and how fast they are, perhaps you could post a specific equation for what you think will happen.
FWIW (though I really have no expertise on this) intuitively it wouldn’t seem surprising to me if the reaction actually used OH- like this: Cl2 + OH- ⇌ HOCl + Cl- , whereas it’s harder for me to visualize how it would work with OCl- as an intermediate (like, why is your intermediate reaction stripping both H’s off the O, while your final reaction puts them back on?).
Note that the total quantity of Cl2 and HOCl (or in more basic solution ClO− ) is conserved in the above reaction
Ah, nice, that reaction is actually an explanation of why you might get roughly equal amounts at first.
So the putative equilibrium is the above (and also including some H3O+ and ClO− and Cl− from dissociation of the stuff on the right) and not this:
So can you tell me what the equilibrium is at pH 7? That wikipedia quote solely mentions environments of unspecified levels of acidity / alkali.
It might be important to note that I’m “cheating,” because I know as an empirical fact that ClO- is unstable in water—dilute bleach eventually turns into salt water, while salt water does not turn into dilute bleach.
Note that the total quantity of Cl2 and HOCl (or in more basic solution ClO− ) is conserved in the above reaction.
Sure, but it’s not the only possible reaction. Consider 2HOCl <-> 2HCl + O2.
So can you tell me what the equilibrium is at pH 7?
I unfortunately don’t know the Cl2 equilibrium at neutral pH, (I tried to calculate an overall equilibrium constant at fixed neutral pH from the equilibrium constants for acid and base context on wiki, got inconsistent results from them, and since I actually don’t know what I’m doing but am just applying half-remembered stuff from the one chemistry course I ever took at university plus looking things up on the fly, don’t understand why). But if you just want to know how much is OCl− vs HOCl, here’s a link with a graph on page 2 (basically should be the same as what you’d calculate using the acid dissociation constant for hypochlorous acid).
Anyway:
It might be important to note that I’m “cheating,” because I know as an empirical fact that ClO- is unstable in water—dilute bleach eventually turns into salt water, while salt water does not turn into dilute bleach.
So yes, you could have oxygen bubbling out too. I guess in this case we wouldn’t be as concerned, as presumably the sort of oxidation reaction done by oxygen itself is expected in an oxygen rich environment such as we live in and wouldn’t cause additional harm.
So, I guess that’s what you meant by metastable ClO−. But it sounds like this reaction is slow if not at high temperature? Also, I wouldn’t expect it so much to happen directly with HOCl, rather than ClO− , because O is the middle atom in HOCl so it seems to me it would be less likely to get pulled out in one step than if it’s one of the edge atoms.
At neutral pH HClO + ClO- can obviously react to form oxygen via peroxyhypochlorous acid. That’s not what happens to bleach over time because it’s sodium hypochlorite. It either loses Cl2 to air or ClO- disproportionates to chlorate and Cl-.
I’m trying to push frontiers. This stuff is on wikipedia.
I hadn’t heard of peroxyhypochlorous acid before, but looking it up (HOOCl) I can imagine it forming by the O’s of ClO- and HOCl meeting and kicking out one of the Cl’s as Cl-. That being said, given that Cl with O’s bonded tends to be more of a thing than oxygen-oxygen bonds (and Cl would be the more positive side (?) of the Cl/O bond having more protons and thus more likely to bond with the O than another O?), wouldn’t chlorous acid (HOClO) be more likely to be produced by those things reacting (with either of the Cl’s bonding with the O of the other and kicking out the other Cl as Cl-)? Which would then presumably lead to further oxychlorine stuff rather than pure oxygen?
Just as you say, chlorine dissolved in water likes to be Cl- or ClO- anions.
My suspicion after more thought is that dissolving Cl2 in water produces excess metastable ClO, which breaks down into Cl over a few hours, and that out-of-equilibrium ClO is the antimicrobial and also a reasonable chlorinating agent to be concerned about.
Cl- and ClO- are two different things, the latter is an oxidant while the former is not. It seems odd to me to bundle them together.
I don’t know what particular oxidants would cause more or worse biological effects. So, not sure whether it would matter if there’s “excess metastable” of one particular oxidant or not. But, “excess metastable ClO” seems an odd thing to expect—it sounds like you’re expecting a reaction to go past equilibrium, why?
Truly, shame on bhauth three comments above. (sarcasm) I figure we’re just mentioning them together because they’re what you get when reacting chlorine gas with water.
(Note I meant “excess metastable ClO” to be parsed as “an excess of metastable ClO,” not as if “excess metastable” was an extra fancy kind of excess—un excès métastable. Although I guess that also kind of works.)
I do not. So maybe I should explain the three steps I’m thinking of:
State 1 is water and chlorine gas. This is out of equilibrium—the chlorine will react with the water.
State 2 is water, H3O+, and a somewhat-even mixture of Cl- and ClO- ions (and trace Cl2). This is also out of equilibrium, the ClO- will react with the H3O+. But it’s less out of equilibrium than state 1.
State 3 is water, H3O+, Cl-, and trace amounts of ClO- (and even tracer Cl2). This is the putative equilibrium.
I expect treated water to go 1->2->3. Not 1->3->2->3 or any such thing. If it just goes 1->3, never 2 at all, then I’m wrong. Also, it if goes 1->2->3 but 1->2 is slow and 2->3 is fast, then I’m wrong.
According to wikipedia:
So the putative equilibrium is the above (and also including some H3O+ and ClO− and Cl− from dissociation of the stuff on the right) and not this:
Note that the total quantity of Cl2 and HOCl (or in more basic solution ClO− ) is conserved in the above reaction. You do not get to a point where both are trace, if that’s not what you started with. In natural stomach acid, you would presumably not start with either of them, but in chlorinated tap water you do.
Regarding the specific intermediate steps in a reaction and how fast they are, perhaps you could post a specific equation for what you think will happen.
FWIW (though I really have no expertise on this) intuitively it wouldn’t seem surprising to me if the reaction actually used OH- like this: Cl2 + OH- ⇌ HOCl + Cl- , whereas it’s harder for me to visualize how it would work with OCl- as an intermediate (like, why is your intermediate reaction stripping both H’s off the O, while your final reaction puts them back on?).
Ah, nice, that reaction is actually an explanation of why you might get roughly equal amounts at first.
So can you tell me what the equilibrium is at pH 7? That wikipedia quote solely mentions environments of unspecified levels of acidity / alkali.
It might be important to note that I’m “cheating,” because I know as an empirical fact that ClO- is unstable in water—dilute bleach eventually turns into salt water, while salt water does not turn into dilute bleach.
Sure, but it’s not the only possible reaction. Consider 2HOCl <-> 2HCl + O2.
I unfortunately don’t know the Cl2 equilibrium at neutral pH, (I tried to calculate an overall equilibrium constant at fixed neutral pH from the equilibrium constants for acid and base context on wiki, got inconsistent results from them, and since I actually don’t know what I’m doing but am just applying half-remembered stuff from the one chemistry course I ever took at university plus looking things up on the fly, don’t understand why). But if you just want to know how much is OCl− vs HOCl, here’s a link with a graph on page 2 (basically should be the same as what you’d calculate using the acid dissociation constant for hypochlorous acid).
Anyway:
It makes sense that that could happen:
So yes, you could have oxygen bubbling out too. I guess in this case we wouldn’t be as concerned, as presumably the sort of oxidation reaction done by oxygen itself is expected in an oxygen rich environment such as we live in and wouldn’t cause additional harm.
So, I guess that’s what you meant by metastable ClO−. But it sounds like this reaction is slow if not at high temperature? Also, I wouldn’t expect it so much to happen directly with HOCl, rather than ClO− , because O is the middle atom in HOCl so it seems to me it would be less likely to get pulled out in one step than if it’s one of the edge atoms.
At neutral pH HClO + ClO- can obviously react to form oxygen via peroxyhypochlorous acid. That’s not what happens to bleach over time because it’s sodium hypochlorite. It either loses Cl2 to air or ClO- disproportionates to chlorate and Cl-.
I’m trying to push frontiers. This stuff is on wikipedia.
I hadn’t heard of peroxyhypochlorous acid before, but looking it up (HOOCl) I can imagine it forming by the O’s of ClO- and HOCl meeting and kicking out one of the Cl’s as Cl-. That being said, given that Cl with O’s bonded tends to be more of a thing than oxygen-oxygen bonds (and Cl would be the more positive side (?) of the Cl/O bond having more protons and thus more likely to bond with the O than another O?), wouldn’t chlorous acid (HOClO) be more likely to be produced by those things reacting (with either of the Cl’s bonding with the O of the other and kicking out the other Cl as Cl-)? Which would then presumably lead to further oxychlorine stuff rather than pure oxygen?