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Thread: Oxidization of urea

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    Oxidization of urea

    Split off of Is shocking necessary.......... JasonLion

    Urea is the primary bather waste. And, although urea can form chlorinated derivatives such as monochlorourea, dichlorourea, trichlorourea and tetrachlorourea, I don't think that these derivatives contribute significantly to the combined chlorine level. The oxidation process for urea is much slower than that for ammonia and creates dichloramine and trichloramine as intermediates in the process. Therefore, urea can create persistent levels of combined chlorine as it breaks down. Urea may take a day or two to fully break down.

    The ammonia based chloramines absorb UV light and will be about 2/3 oxidized by the UV alone even if there is no free chlorine.

    3NH2Cl + 3(uv) --> N2 + NH4+ + 2H+ + 3Cl-

    monochloramine + uv --> Nitrogen + ammonium ion + hydrogen ions + chloride ions.

    Indoor pools with heavy bather loads and persistent chloramines should consider installing a proper ozone system or UV.

    The oxidation of ammonia by chlorine produces about 90 % nitrogen and about 10 % nitrate. If the only product were nitrogen, then the molar ratio would be 1.5 chlorine to ammonia. The formation of nitrate increases the needed ratio to about 1.75. Since chlorine is about 5.07 times the mass per mole of nitrogen, the needed ratio of ppm chlorine to ppm ammonia nitrogen is about 8.9.

    Since combined chlorine is reported in units of chlorine, the total chlorine needs to be 1.75 times the combined chlorine, assuming that all of the ammonia is combined. As long as the total chlorine is at least 1.75 times the combined chlorine, then all of the chloramines can be broken down.

    For example, if the CC were 2.0, then the TC would need to be at least 3.5 and the FC would need to be 3.5 - 2.0 = 1.5 ppm. Using more chlorine will make the process go faster. Using the minimum will oxidize the ammonia in about 1 hour. Using a higher chlorine level can oxidize the ammonia in about 15 minutes.

    To calculate the best level of chlorine to add, you should determine the maximum level you want to swim in (say 20 % of the cyanuric acid) and add 1.75 times the CC level, and then increase the TC to that number.

    For example:

    FC = 1.0
    CC = 2.0
    TC = 3.0
    CYA = 60

    .20 x 60 + 2 x 1.75 = 15.5
    15.5 (target TC) - 3.0 (existing TC)= 12.5 ppm of chlorine to add.

    High levels of combined chlorine (from ammonia), especially after heavy bather loads, indicate that there are probably also high levels of urea (which won't show up in the CC test). As noted above, urea takes time to break down and produces ammonia chloramines in the process. This can make CC levels seem to take extended amounts of time to eliminate.

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    Re: Is shocking necessary.........

    This latter discussion could be moved to The Deep End (post above and this one).

    The chlorinated urea, such as monochlorourea, most certainly get reported as Combined Chlorine (CC) and in fact are the most dominant CC reported in moderate-to-high bather load pools. See this paper including the supplemental information including Figure S2(a) where it is clear that chlorine combining with urea registers as CC in the DPD test (but does not show up as monochloramine in MIMS tests, which makes sense). This paper is a useful follow-on with more info on various disinfection by-products formed from chlorine oxidation of ammonia and organics. This paper describes how the first chlorination step for urea is slow as it comes from aqueous molecular chlorine and not from hypochlorous acid.

    Regarding the total cumulative amount of chlorine needed to oxidize ammonia, you need to account for the fact that one chlorine is already attached to ammonia in monochloramine which registers as CC so the amount needed to complete oxidation is not 1.75x but rather 0.75x. Of course, it's a moot point anyway since any amount of FC will make the reaction proceed and one can always add more to pick up where one left off. A higher active chlorine (proportional to FC/CYA ratio) just makes the reactions go faster, but nothing gets stuck. Urea requires more chlorine for its oxidation, but it's still not more than around 3x for a CC reading representing monochlorourea (the 3x gets one over the hump in the worst case; actual usage is between 2x and 3x depending on amount of nitrate produced, as you indicated).
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    Re: Is shocking necessary.........

    Quote Originally Posted by chem geek
    The chlorinated urea, such as monochlorourea, most certainly get reported as Combined Chlorine (CC) and in fact are the most dominant CC reported in moderate-to-high bather load pools.
    Quote Originally Posted by John A. Wojtowicz
    Surprisingly, urea does not itself form combined chlorine and also does not appear to affect disinfection. However, urea has to be destroyed by oxidation because it is a nutrient for bacteria and algae and are a potential source of ammonia chloramines. Oxidation of urea by free chlorine is a slow process that gives rise to transient ammonia chloramines (e.g., di– and trichloramine).

    Although urea can form chlorinated compounds in aqueous solution, they apparently readily hydrolyze since they do not contribute significantly to combined chlorine (Palin 1959). In addition, since they do not affect disinfection (Fitzgerald and DerVartanian 1967), their concentration must be quite low.
    http://jspsi.poolhelp.com/ARTICLES/JSPS ... p30-40.pdf
    Quote Originally Posted by chem geek
    Regarding the total cumulative amount of chlorine needed to oxidize ammonia, you need to account for the fact that one chlorine is already attached to ammonia in monochloramine which registers as CC so the amount needed to complete oxidation is not 1.75x but rather 0.75x.
    I noted that the total chlorine needed to be 1.75 times the CC. The TC does account for the 1 chlorine already in the CC.
    Quote Originally Posted by chem geek
    Of course, it's a moot point anyway since any amount of FC will make the reaction proceed and one can always add more to pick up where one left off.
    However, it won't make it proceed to completion. My example gives the amount needed for the reaction to proceed to completion and still have the desired amount of FC upon completion.

    Regarding the papers that you cite, I can only get the abstracts for the papers unless I buy the full document, which I'd rather not have to do unless I had to.

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    Re: Oxidization of urea

    As far as I can tell, neither of you disagreed with the other. You are just describing the same situation from different points of view. Monochlorourea reads as CC, but the monochlorourea level is commonly close enough to zero that it's effect on the CC level is not important in normal situations.

    The problem with the "Raise TC to desired level plus 1.75*CC", or alternatively "Raise FC to desired level plus 0.75 times CC", rule, is that it assumes that "all of the ammonia is combined". That is almost never true in practice if you are counting all of the ammonia and not just the ammonia currently free to bind to chlorine. There is nearly always some ammonia bound up in urea, or other complex compounds, which won't show up until later. In practice, it is also possible to have excess ammonia, without enough chlorine available to form CC, and thus not showing on the test results, though that is less common.

    That is why chem geeks comment about "one can always add more" is so important. Any attempt to say "this is the CC level now, so this is what I need to add, and then I am done" is doomed if real people have been in the water in last few days. In a real pool, there will nearly always be more CC showing up in the relatively near future, due to the break down of complex chemicals that are not normally measured, so any attempt to hit a target FC level 12 hours latter will fail. Thus you will end up adding more chlorine to continue trying to hit/maintain your target FC level.

    I find it crucial to view things with the understanding that in real pools CC is both being continuously created and continuously destroyed. The question becomes what does the FC level need to be in order to get the rate of CC destruction to be higher than the rate of CC creation, which is the essential first step in maintaining minimal ongoing CC levels.

    In high bather load pools one key constraint is the highest FC level you are willing/allowed to have while people are in the pool. That often ends up forcing you to have the FC level too low during the day with rising CC levels while people are in the pool, putting you into a race to get things back into balance overnight before people want to swim tomorrow.

    In practice that kind of situation often forces you to say things like, this is the highest I can raise FC tonight and still have FC be below X tomorrow when the pool opens. That kind of thinking results in bringing up ideas like FC should be target FC + 0.75*CC, but it is important to understand that even through the formula is "true" it doesn't completely apply to real situations and you can often actually raise the FC level higher than that and still be down to your FC target tomorrow because other things will be using up chlorine overnight.
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    Re: Is shocking necessary.........

    Quote Originally Posted by John A. Wojtowicz
    Surprisingly, urea does not itself form combined chlorine and also does not appear to affect disinfection.
    :
    Although urea can form chlorinated compounds in aqueous solution, they apparently readily hydrolyze since they do not contribute significantly to combined chlorine (Palin 1959).
    John is wrong and so is his interpretation of Palin. I think John was fooled by the fact that the formation of monochlorourea is so slow so mixing chlorine and urea together won't show any CC for a while (hours) unless the urea concentrations are quite high. However, eventually it will show up and that's what Blatchley's experiments showed. It will build up until the rate of creation equals the rate of destruction. The net result is that the concentration of urea builds up in pools until a steady-state is reached. As Jason points out, in most residential pools that steady-state does not result in any significantly measurable CC because the bather load is so low and the urea is oxidized fast enough in outdoor pools (possibly due to UV in sunlight through indirect mechanisms such as hydroxyl radical formation from chlorine breakdown -- a guess on my part).

    I misunderstood your TC rule, thinking it was FC; sorry about that. I agree with Jason, that such rules aren't very useful because it's really more about continual creation and destruction and having a high enough FC to control the overall net rate or lowering CC. The main problem is with some indoor pools where the CC becomes persistent where even shock levels of FC don't seem to be able to get rid of it. Without UV from sunlight, some CC is simply too slow to oxidize and needs help which can be obtained by using non-chlorine shock (MPS), for example.

    As for paying for articles, I hate that science isn't free but have just bit the bullet and have downloaded dozens of paid-for articles costing hundreds of dollars. Such is the price for finding out the truth these days. The "Supporting Info" is free, however, so look at this PDF file where you can see Figure S2(a) I was talking about. Notice that a reasonably detectable amount of CC (which looks like dichloramine if using DPD tests that distinguish between them) shows up after an hour if one has a rather high urea level (1.8x10-5 M) and uses at least a 6:1 molar ratio of chlorine to urea. This implies around 7.7 ppm FC.
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    Re: Oxidization of urea

    Here is a reference that contains this information:

    Urea, for example, is added to pool water in urine and sweat and, according to Dr. Blatchley, is a significant source of trichloramine (NCl3), the primary cause of “chlorine” odors and irritation in pool environments. In Dr. Blatchley’s experiments, it takes 12 hours for half the urea to react with the chlorine and over 5 days for the combined chlorine to drop to less than 0.1 ppm (measured with DPD) at a chlorine dose over 22 times the nitrogen concentration.
    Here is an interesting article that contains the following:

    Combined chlorine is defined as the sum of the following compounds [1]: derivatives of ammonia in which one, two or three hydrogen atoms have been replaced by chlorine atoms (monochloramine, NH2Cl; dichloramine, NHCl2; trichloramine, NCl3); and all chlorinated derivatives of urea and organic nitrogen compounds such as creatinine and amino acids.

    The decisive mechanism for the formation of trichloramine in pool water is the step-by-step reaction of the urea introduced by pool users with free chlorine to 1,1,3,3-tetrachlorourea and finally to trichloramine, as described in the literature [15].

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    Re: Oxidization of urea

    You really need to buy the actual scientific papers because all of these summaries only give partial explanations of what is going on and don't have the full context or conditions described in detail. However, the main points are correct -- chlorine combining with urea to from chlorourea that measures as CC is a slow process (so in practice, urea builds up over time until that rate gets faster) and the oxidation of urea is a main source of nitrogen trichloride. What wasn't said in those snippets, but what comes from the detailed papers, is that dichloramine and monochloramine are also formed from the oxidation of urea so the proposed mechanism can't be the only thing that is going on. In fact, the proposal from Samples via an NOH intermediate can't be right either since it doesn't explain the dichloramine and monochloramine.

    I think a reasonable clue as to what might be really going on comes from this link describing different pathways for urea decomposition in water where the (a) mechanism involves urease (via a rather unique mechanism) while the (b) mechanism is regular hydrolysis (though with a nickel catalyst). There are several possible pathways that could be occurring starting with quadchlorourea and assuming the (b) mechanism which might be enhanced after chlorination of the amine groups. [EDIT] In the following, the "***" following an equation shows those that are needed to more closely match the Blatchley & Cheng paper. [END-EDIT] [EDIT] The rate of hydrolysis of isocyanic acid (HNCO which is a tautomer of cyanic acid HOCN) comes from this paper. [END-EDIT]

    CO(NCl2)2 ---> NCl3 + ClN=C=O (HOCl is a catalyst for this reaction, i.e. a reactant and a product)
    [EDIT] CO(NCl2)2 + H2O ---> NHCl2 + ClN=C=O + HOCl *** [END-EDIT]

    ClN=C=O + OCl- ---> Cl2NCOO- (relatively slower; unsure of actual rate)
    Cl2NCOO- + H+ --> NHCl2 + CO2 (fast)
    Cl2NCOO- + HOCl --> NCl3 + CO2 + OH- (fast)

    ClN=C=O + H2O ---> HClNCOOH ---> HClNCOO- + H+ (relatively slower; possibly at about 9% per minute) ***
    HClNCOO- + H+ --> NH2Cl + CO2 (fast) ***
    HClNCOO- + HOCl --> NHCl2 + CO2 + OH- (fast)

    [EDIT]
    OR alternatively, the reaction may proceed by the mechanism that occurs with urease but instead of water it is HOCl that may be the electrophile:
    Cl2NCONCl2 + HOCl ---> Cl2NC(O-)(OH)(NCl3+) ---> Cl2NCOO- + H+ + NCl3
    Cl2NCONCl2 + H2O ---> Cl2NC(O-)(OH)(NHCl2+) ---> Cl2NCOO- + H+ + NHCl2
    Cl2NCOO- + H+ ---> NHCl2 + CO2 (fast)
    Cl2NCOO- + HOCl ---> NCl3 + CO2 + OH- (fast)
    [END-EDIT]

    The above could be a plausible mechanism that allows for differential amounts of dichloramine and nitrogen trichloride which is what Blatchley's experiments were seeing. In fact, he was even seeing some monochloramine which the above would allow. I show other possibilities in my spreadsheet such as dichlorourea getting oxidized by hypochlorous acid to monochloramine and dichloramine. The mechanism proposed by Samples going to NOH, however, won't create the monochloramine and dichloramine that are seen right away so I don't think his proposal is correct.

    One could also start with a dichlorourea and have the following possible pathway:

    H2NCONCl2 ---> NHCl2 + HN=C=O ***

    HN=C=O + OCl- ---> ClHNCOO- (relatively slower; unsure of actual rate)
    ClHNCOO- + H+ --> NH2Cl + CO2 (fast)
    ClHNCOO- + HOCl --> NHCl2 + CO2 + OH- (fast)

    HN=C=O + H2O ---> H2NCOOH ---> H2NCOO- + H+ (relatively slower; roughly at about 9% per minute) ***
    H2NCOO- + H+ --> NH3 + CO2 (fast) ***
    NH3 + HOCl ---> NH2Cl + H2O (very fast) ***
    H2NCOO- + HOCl --> NH2Cl + CO2 + OH- (fast)

    [EDIT]
    OR alternatively, the reaction may proceed by the mechanism that occurs with urease but instead of water it is HOCl that may be the electrophile:
    H2NCONCl2 + HOCl ---> H2NC(O-)(OH)(NCl3+) ---> H2NCOO- + H+ + NCl3
    H2NCONCl2 + H2O ---> H2NC(O-)(OH)(NHCl2+) ---> H2NCOO- + H+ + NHCl2
    H2NCOO- + H+ ---> NH3 + CO2 (fast)
    NH3 + HOCl ---> NH2Cl + H2O (very fast)
    H2NCOO- + HOCl ---> NH2Cl + CO2 + OH- (fast)
    [END-EDIT]

    Depending on the various reaction rates for each of the above, one can end up with differing amounts of monochoramine, dichloramine and nitrogen trichloride. Note that to get nitrogen trichloride in the large quantities that Blatchley was seeing, the formation of quadchlorourea seems to be the dominant mechanism. [EDIT] OR if the reaction proceeds by the mechanism that occurs with urease but instead of water it is HOCl that is the electrophile, then nitrogen trichloride is produced from either quadchlorourea or dichlorourea, but the former produces no monochlorurea while the latter produces no dichlorourea. [END-EDIT]
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    Re: Oxidization of urea

    Quote Originally Posted by chem geek
    I think a reasonable clue as to what might be really going on comes from this link describing different pathways for urea decomposition in water where the (a) mechanism involves urease (via a rather unique mechanism) while the (b) mechanism is regular hydrolysis (though with a nickel catalyst).
    In the scientific literature, the mechanism of trichloramine formation from urea is discussed from three different directions:

    • Enzymatic degradation of urea, by the enzyme urease which is contained in various bacteria, to ammonia or ammonium, and reaction of the latter with free chlorine to trichloramine. According to Jessen and Gunkel [13], this process does not occur in chlorinated pool water;

    • Hydrolysis (cleavage by the action of water) of urea, with formation of ammonia or ammonium, and subsequent reaction with free chlorine to trichloramine. This occurs only at temperatures of more than 65C (149 F) and is not, therefore, relevant to pool water;

    • The decisive mechanism for the formation of trichloramine in pool water is the step-by-step reaction of the urea introduced by pool users with free chlorine to 1,1,3,3-tetrachlorourea and finally to trichloramine, as described in the literature [15].

    [15] Robson, H.L.: Chloramines. In: Encyclopedia of Chemical Technology, Kirk, R.; Othmer, D.F. ed., 2nd ed., Vol. 4, 908-928, John Wiley & Sons, New York, 1993.

    Reference
    This reference that you cited indicates that the monochloramine and dichloramine are the result of the hydrolysis of trichloramine.
    N-Chlorourea then appears to undergo further chlorine substitution; the fully N-chlorinated urea molecule is hypothesized to undergo hydrolysis and additional chlorination to yield NCl3 as an intermediate. NCl3 is hydrolyzed to yield NH2Cl and NHCl2, with subsequent decay to stable end products, including N2 and NO3.

    http://pubs.acs.org/doi/abs/10.1021/es102423u
    The oxidation of ammonia nitrogen by chlorine to gaseous nitrogen at the breakpoint would theoretically require 1.5 mol of chlorine (Cl2) per mole of nitrogen oxidized according to Equation 14.16.

    The observed stoichiometric molar ratio between chlorine added and ammonia nitrogen consumed at breakpoint is typically about 2:1, suggesting that more oxidized nitrogen compounds are produced at breakpoint rather than N2 gas. Experimental evidence (Saunier and Selleck, 1979) indicates that the principal additional oxidized product may be nitrate formed via Equation 14.17:

    NH4+ + 4HOCl = NO3- + 4 Cl- + 6 H+ + H2O (14.17)

    Depending upon the relative amount of nitrate formed in comparison to nitrogen at breakpoint, between 1.5 and 4.0 mol of available chlorine may be required, which is consistent with the available data.

    Below the breakpoint, inorganic chloramines decompose by direct reactions with several compounds. For example, monochloramine may react with bromide ions to form monobromamine (Trofe, 1980).

    If trichloramine is formed, as would be the case for applied chlorine doses in excess of that required for breakpoint, it may decompose either directly to form nitrogen gas and hypochlorous acid or by reaction with ammonia to form monochloramine and dichloramine (Saguinsin and Morris, 1975).

    http://203.158.253.140/media/e-Book/Eng ... 593_14.pdf

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    Re: Oxidization of urea

    Quote Originally Posted by JamesW
    This reference that you cited indicates that the monochloramine and dichloramine are the result of the hydrolysis of trichloramine.
    N-Chlorourea then appears to undergo further chlorine substitution; the fully N-chlorinated urea molecule is hypothesized to undergo hydrolysis and additional chlorination to yield NCl3 as an intermediate. NCl3 is hydrolyzed to yield NH2Cl and NHCl2, with subsequent decay to stable end products, including N2 and NO3.

    http://pubs.acs.org/doi/abs/10.1021/es102423u
    This is the problem with only looking at the abstracts. The detailed paper itself doesn't talk at all about hydrolysis of NCl3 and the mechanisms proposed by Samples referenced in the paper were speculative and since shown to not be accurate in other papers. The Samples reference in the Blatchley paper has NCl3 go to nitrate. The best models for the inorganic chloramines are Jafvert & Valentine (1992), Vikesland, Ozekin & Valentine (2000) and Kumar, Shinness & Margerum (1987) where the latter does have a SLOW hydrolysis of nitrogen trichloride to produce dichloramine. I have a spreadsheet with these inorganic chloramine models plus some early stabs at a urea model but the latter is very incomplete. Again, the Blatchley data shows monochloramine and dichloramine forming almost immediately along with nitrogen trichloride (in Figure 3), much faster than could occur from hydrolysis of nitrogen trichloride which is roughly a creation rate of 3% per hour (initial creation rate of dichloramine in Figure 3 is over 30% of the nitrogen trichloride amount per hour).

    Again, you really need to read the detailed papers to see that some of the summary info is either misleading or incorrect. Papers tend to focus their accuracy on the main parts of their experiment, but when they wander off into tying things to early models, they sometimes don't reference newer papers or attempt to speculate on models that would be more consistent with their data. Blatchley does briefly reference Jafvert & Valentine right after listing the very old Samples (1959) proposal when Blatchley writes "Note that the unstable intermediate NOH has also appeared in other mechanisms that address the behavior of reduced-N and free chlorine (15)." where (15) is a reference to Jafvert & Valentine (1992).

    The inorganic chloramine latest models show nitrogen trichloride reacting with monochloramine and dichloramine to form nitrogen gas, but not reacting with ammonia to form monochloramine or dichloramine. Look at the Jafvert & Valentine (1992) paper as his experiments were quite extensive. It's not perfect, but probably the best model out there right now if one adds the Kumar, Shinness & Margerum rate for nitrogen trichloride hydrolysis which was too slow to be of importance in the Jafvert & Valentine experiments.
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    Re: Oxidization of urea

    Since some trichloramine off gasses, taking with it 3 chlorines, which is 1 to 1.5 more than would be used to oxidize nitrogen (assuming it takes 1.5 to 2 chlorine to oxidize 1 nitrogen), how much do you think the total chlorine demand is increased by this?

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    Re: Oxidization of urea

    Not very much because the trichloramine concentrations are normally very small since anything above 20 ppb in water is considered to be irritating and though it's very volatile, it still takes times to outgas. Only extremely high urea and chlorine levels as were done in some of Blatchley's experiments produce enough nitrogen trichloride to affect mass balance by any degree. Table 1 in the "Reaction Mechanism for Chlorination of Urea" paper at a pH of 7.5 after 24 hours with 100 µM N urea concentration (5x10-5 M urea) and chlorine at 1.5x10-4 (10.6 ppm FC), 70.1 µM remained as urea, 0.266 µM was still as NH2Cl (monochloramine), 0.964 µM was still as NHCl2 (dichloramine), 1.85 µM was as NCl3 (nitrogen trichloride), 5.01 µM was as NO3[- (nitrate) and the rest, 21.8 µM is presumed to be nitrogen gas.

    At lower active chlorine levels, the nitrogen trichloride will be proportionately less while monochloramine and dichloramine will be higher (though they are intermediates). Also, typical pool urea concentrations are in the range of 0.2 to 6 µM.
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    Re: Oxidization of urea

    Since UV is so effective at eliminating chloramines, I wonder if anyone has ever considered putting UV lights over the pool to radiate the entire pool overnight?

    Safety would be a major issue, so there would need to be some way to make sure that no one was in the pool area during the time that the lights were on.

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    Re: Oxidization of urea

    Installing UV lights that were strong enough to be effective through several feet of water would be extremely expensive.
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    Re: Oxidization of urea

    A related topic though would be to have a UV in-line system with a spectrum more like the UV portion of sunlight. Though it would use up some chlorine, that's not a bad thing in that it produces hydroxyl radicals that might accelerate oxidation of urea. Of course, a more direct approach at such oxidation is to use boron-doped electrodes to do electrolysis that produces a higher quantity of hydroxyl radicals, but that's expensive (it's in Oxineo from Adamant).
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    Re: Oxidization of urea

    Quote Originally Posted by chem geek
    Of course, a more direct approach at such oxidation is to use boron-doped electrodes to do electrolysis that produces a higher quantity of hydroxyl radicals, but that's expensive (it's in Oxineo from Adamant).
    Disinfectant Production
    Depending on the composition of the water medium, several disinfecting agents are produced at diamond electrodes. The presence of chlorides, sulfates and carbonates respectively induces a very efficient generation of free chlorine, peroxydisulfate and percarbonates. But peroxides, ozone and hydroxyl radical are also directly produced from water. Most often, with diamond electrodes, a mixture of all these oxidants is responsible for the water disinfection. http://www.adamantec.com/water_applicat ... ection.php
    Their explanation is not exactly clear about which oxidants are the primary disinfectant species.

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    Re: Oxidization of urea

    Hydroxyl radicals electrochemically generated in situ on a boron-doped diamond electrode.
    Electrogeneration of Hydroxyl Radicals on Boron-Doped Diamond Electrodes.
    Essential explanation of the strong mineralization performance of boron-doped diamond electrodes..
    Boron doped diamond electrode for the wastewater treatment.
    Electrogeneration of Hydroxyl Radicals on Boron-Doped Diamond Electrodes.

    etc...

    Just do a Google search on "boron-doped diamond electrodes hydroxyl radical". Though it is true that depending on the salt level some chlorine will be generated, the whole idea of boron-doped diamond is to favor production of hydroxyl radicals.
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