Bleach vs Acid

Randy Homes-Farly said:
I just don't see any reason to think that any of these are appreciably more effective than bleach at removing organics.

But why? If you believe that then you must have a reason for believing it.

Randy Holmes-Farly said:
I also don't think that lack of effectiveness is any concern. I've never heard of anyone who could not remove all the crud from rock with bleach.

I find it hard to believe that I'm the first one. I was peeling gunk off of my rocks even after a bleach and acid bath. There are a number of photos by people in this thread who have performed the same process that show noticeable discoloration and gunk on the rock even after this process.

The acid actually seems to do a better job than the bleach. For example if I take a HA infested rock and bleach it the algae just turns white or brown. When I put it back in the tank a bacterial slime coats it and begins breaking it down. Whereas the acid dissolves it completely.

I tried the NaOH solution yesterday. 1.25% concentration. It actually seemed to be less effective somehow than either the acid or the bleach. Going to try the permanganate later this week or next.
 
But why? If you believe that then you must have a reason for believing it.

Must I? :lol:

We can start with the fact the electrochemical potential for HClO is similar to MnO4- (about 1.6-1.7 V). obviously, pH and concentrations impact the actual potential, but bleach can be rather strong.

We can add to that the reason that people use bleach rather than permanganate to whiten clothes (that is, oxidize organics): it works well.

If you want to try permanganate, and are not worried about excessive manganese on the rock when you are done, then by all means try and and let us know how it works. It might turn your rock brown with MnO2.
 
Randy Holmes-Farley said:

Well it would be kind of strange for a man of science to offer a confident conclusion based solely on a gut feeling. Evidence and an understanding of the underlying principles are what guide scientific reason. If I wanted gut feelings I could ask anyone. But with your background you can offer accurate conclusions based on real evidence and unique insight on what's happening since you have the knowledge needed. That's why I came to you.

Seriously though I'm very interested in gaining a detailed understanding of the chemical side of what's going on here. If I needed anything else I could easily just consult google. But for this topic I was unable to find the information I needed, forcing me to post here. If you want to just shoot me some links to relevant reading material that would also work. I was surprised that I couldn't really find anything useful on google about decomposing organic matter chemically.

Randy Holmes-Farley said:
We can add to that the reason that people use bleach rather than permanganate to whiten clothes (that is, oxidize organics): it works well.

More often industry standards are chemicals that are "good enough" and extremely cheap. Rather than the best option. KMnO4 whitens clothes even more than clorox does but it's several times the cost. I would imagine that's the main reason that people don't use it for laundry. But yes it does work well for that purpose. But how does it fair at actually decomposing solid organic matter? Not just whitening it.

Randy Holmes-Farley said:
It might turn your rock brown with MnO2.

How come?

Edit: Oh and another question. I notice some black spots on my rock after bleaching. Any idea what these might be? I think it's a mineral or salt of some kind. I'm wondering if there is a way to remove it, which of course would first require identifying it. I can also get oxalic acid and phosphoric acid for cheap. Oxalic might be a decent choice due to its lower acidity and reduction potential. Which brings me to another question. Why are oxidizing agents used for bleaching rather than reducing agents?
 
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If I wanted gut feelings I could ask anyone.

I suspect my gut feelings on the properties of chemicals are better than average.

It might turn your rock brown with MnO2.


How come?


Because it is dark brown? I'm not saying it will, just that it might. Fresh CaCO3 surfaces bind many things and MnO2 is not all that soluble.. :)

Why are oxidizing agents used for bleaching rather than reducing agents?

I discuss that in this ozone article:

Ozone and the Reef Aquarium, Part 1: Chemistry and Biochemistry
http://reefkeeping.com/issues/2006-03/rhf/index.php

from it:

Oxidation of Organics by Ozone: Decoloration
The oxidation of organics is, it turns out, the primary reason that reef aquarists use ozone because it is the organic material in seawater that causes clarity and color issues. Its impact on organic materials is also why ozonation impacts skimming. While most organic compounds that are exposed to enough ozone for a long enough period will be oxidized in some way, some are very much more sensitive than others. In fact, at the levels of ozone attained in a typical reef aquarium contact chamber (less than about 0.3 ppm ozone) or even disinfection applications where the doses are much higher, the total dissolved carbon does not appreciably change during the ozone exposure (although it may later if bacteria find the newly oxidized organics more bioavailable; see below).

In a marine mammal pool,18 for example, it was found that disinfection with 4 ppm ozone with a 30 minute contact time (a disinfection level much higher than is typically used in reef aquaria) did not reduce the pool's total organic carbon (TOC) (~13 ppm TOC), while the use of granular activated carbon (GAC) did reduce it by 37%. Interestingly, the combination of ozone and GAC was even more effective, removing 60-78% of the TOC, suggesting that the ozonation may have altered some of the molecules in a way that made them bind more strongly (or more rapidly) to GAC. An alternative explanation that cannot be ruled out involves biological transformations of the organic compounds taking place on the GAC surface as it became colonized with bacteria).

One research group19 studying the reaction between a variety of organic compounds and ozone concluded:

"…comparisons of rate constants with chemical structures of the reacting groups show that all reactions of O3 are highly selective…"

Fortunately, many of the organic compounds that are most reactive with ozone coincidently are those that aquarists want to eliminate from aquaria. As seawater ages in marine aquaria, the water often becomes yellow as a wide variety of different organic pigments build up. Because of the ozone's reaction with many natural pigments, it is often used in drinking water purification for the purpose of "decoloration;" not organic removal per se, but decoloration.20

In order to understand this effect, it is first instructive to understand which organic molecules lead to coloration, because not all of them do. In fact, most organic molecules are not colored. That is, they do not absorb visible light. Looking through bottles of purified organic compounds, the vast majority are white powders. Organisms, however, have a significant need to absorb light, for example, to photosynthesize or to see.

In order to generate molecules that absorb visible light, natural systems often turn to conjugated carbon-carbon double bonds. Figures 1 and 2, for example, show the structures of chlorophyll and b-carotene. Both of these molecules are widespread in organisms, and both contain conjugated double bonds that lead to the absorption of visible light. These figures do not show the hydrogen atoms (there are dozens of them), but all of the other atoms are shown, and there is a carbon at each intersection of two or more lines. This is how chemists often show structures, allowing the important features to stand out and not get lost in a clutter of atomic letters. What is important here is each segment with a C═C (shown in red). Without going into ridiculous chemical detail for a reef article, having a bunch of C═C bonds arranged together with a single C─C bond between them can lead to the absorption of visible light. That is why organisms have developed such chemical structures for the absorption of light despite their instability toward oxidation (see below).

It is just that instability, however, that aquarists take advantage of when employing ozone. Figure 3 shows, for example, where ozone first attacks oleic acid (a dietary fatty acid).21,22 It is attacked at its double bond, breaking it apart into smaller fragments that no longer have a C═C bond. Consequently, while a huge dose of ozone lasting a very long time will break down these bits even more, even a small dose will remove the C═C bond.

Translating that reactivity to the pigments shown in Figures 1 and 2 makes it apparent why ozone is so good at reducing seawater's coloration and increasing its clarity: it reasonably selectively targets many of the structures that nature uses to absorb light, and converts them to nonabsorbing chemical structures.

A second type of colored organic compound that accumulates in seawater (in both the ocean and aquaria) is one of the functional groups in humic and fulvic acids (the compounds often identified as the yellowing agents in aquaria).20 These "compounds" are complex mixtures of many compounds, but among them is the phenol functional group (Figure 4). Phenol can be attacked by ozone,23,26 with breakdown products shown in Figure 4. It is the Ring-OH group that is colored when in the Ring-O- ionized form, and many of these breakdown products lack such a functional group. Hence the oxidation of such phenolates in humic acids with ozone will reduce color in aquarium water.
The various chemical products described in this section are, of course, not the only reaction products of ozone, hypobromous acid and hypobromite with organic compounds. Other products include brominated organic compounds and many other chemical structures. These have not been fully elucidated, a fact which is not surprising since even in the absence of ozone, the nature of all of the organics in natural seawater or reef aquarium water remains poorly defined.
 
But how does it fair at actually decomposing solid organic matter? Not just whitening it.

It won't break every bond, but it breaks many of given enough time and concentration. I've stripped coral skeletons totally using concentrated bleach, even when the whole LPS coral tissue was still there at the beginning.
 
That still doesn't really answer any of my questions. After reading that I still don't know why oxidation is preferred over reduction. Or why the breakdown occurs.

What about using solid forms of chlorine like calcium hypochlorite (used for pool chlorination) as a more economical form? A $3.50 1 pound bag of Ca(ClO)2 could make 4 gallons of 5% concentration solution. And each molecule has two hypochlorite ions instead of one. And at a lower PH (which reduces stability but increases oxidation potential).

What about using a chlorite (ClO2), chlorate (ClO3), or perchlorate (ClO4) instead? These are much stronger oxidizers. Potassium chlorate is quite cheap and it along with sodium chlorate are widely used in weed killers.

Any ideas about the black mineral on my rock?
 
From everything I have read, and my own personal experience, I would bleach it first, then acid it if you want that really pearly white look.
 
That still doesn't really answer any of my questions. After reading that I still don't know why oxidation is preferred over reduction. Or why the breakdown occurs.

What about using solid forms of chlorine like calcium hypochlorite (used for pool chlorination) as a more economical form? A $3.50 1 pound bag of Ca(ClO)2 could make 4 gallons of 5% concentration solution. And each molecule has two hypochlorite ions instead of one. And at a lower PH (which reduces stability but increases oxidation potential).

What about using a chlorite (ClO2), chlorate (ClO3), or perchlorate (ClO4) instead? These are much stronger oxidizers. Potassium chlorate is quite cheap and it along with sodium chlorate are widely used in weed killers.

Any ideas about the black mineral on my rock?

I don't understand what you mean by reduction. What would you use and what would the resulting product be if the goal were to reduce a protein? Reducing most organic molecules does not break them down in size and makes them less water soluble. oxidizing breaks them down in size and more water soluble. That's why we'd elect oxidizing.

Calcium hypochlorite would probably work OK, although you may not be able to make it as concentrated as bleach because you may begin to reach the solubility limit, if only for calcium hydroxide precipitating. It's oxidizing capability basically the same as bleach (sodium hypochlorite) if equally concentrated, but it is not much cheaper nor is it a stronger oxidizer. Your math is flawed. Having two hypochlorite ions per calcium does not increase the potency of a 5% solution, since it is weight based already. A one pound bag of pure calcium hypochlorite would make about 2.1 gallons of bleach equivalent, but most one pound containers of that for pool treatment are not 100%, but are diluted solids. The pH will only be lower if the concentration is lower, or calcium hydroxide precipitates (which means it is above 12).

If you want to investigate with all those other oxidizers, go for it. I see no reason to think are are better than ordinary, inexpensive, and readily available bleach, but they might be.

I would guess the black stuff may be iron oxide, but I'm not at all sure.
 
Ok so I've ran through all the steps with the pukani rock....it didn't come out pearly white but is a lot nicer then before I didn't wanna do a second batch because it ate quite a bit of the rock as is....now what have you guys been doing after these steps
Right now I have the rock sitting in a batch of fresh r/o water. Been about 6 hours and no smell... Phosphates are at 0
Any advice
Great thread by the way
 
The last step was bleach before the RO water? No bleach smell at all?

The choice is to let it dry and for bleach residue to disappear that way, to use a dechlorinator, or rinse it well. You do not want residual bleach in the tank. :)

Are you starting a new tank or adding to an existing one? The concern is somewhat greater in the second case (IMO), since the first will need to cycle anyway.
 
What I did was
1)a wash in tap and a scrub down
2) bleach wash 1-5 ratio
3)fresh water with a Chlorine remover(prime)
4) acid wash 1-10 ratio then added some baking soda
5)final scrub and was with hose
6) been keeping in ro water for the last 24h
Now I was gonna let it dry out and start building g my new structure
This is a new tank build and I will be adding dr toms to speed up my cycle I've got live stock in a friends that and my tangs are not behaving and wanna get my live stock back
 
I suspect the rock is fine as is. I'd get to work on assembling the system. If there's any bleach odor, there's some bleach on the rock, but the acid will have destroyed the bleach, so I don't know how that could happen.
 
There's no smell but there was a few pieces if shelf i didn't acid because the pieces were awesome looking and didn't wanna change them but the there is no smell actually smells fresh and clean if there is such a smell
 
I'm back.

BRS had a free shipping sale recently and I used the opportunity to get some more pukani rock. I used this opportunity to test some of the chemicals that I was not able to try last time. I ended up trying oxalic acid (H2C2O4), sodium percarbonate (2Na2CO3*3H2O2), and sodium metabisulfite (Na2S2O5). Let's go over each of these real quick and why I picked them.

Oxalic acid is extremely cheap ($2.60 per pound on ebay). Even cheaper per pound than vinegar. A gallon of vinegar costs roughly $3 and contains about 0.4 pounds of acetic acid. That means that oxalic acid is nearly 3x as cheap by weight. It is a much stronger acid than acetic acid but much weaker than hydrochloric acid. Which could be a good thing if you don't want to dissolve too much rock but don't want to wait days for vinegar to work. It is also a strong reducing agent and chelating agent. Which means it should be effective at removing metal and other similar anions if I am not mistaken. Randy mentioned that the stains I was referring to earlier might be iron stains. I am beginning to suspect this more and more myself. Oxalic acids main use is removing iron stains from minerals and metals and as a wood bleach. If these stains are indeed iron then it sounds like this would be the perfect chemical to use to get rid of them. It breaks down into acid H(+1) and oxalate C2O4(-2). Which when reacted with CaCO3 will yield CO2 and CaC2O4 (calcium oxalate). Both harmless.

Sodium percarbonate is an oxygen bleach. It's the active ingredient in oxiclean. You can think of it like solid hydrogen peroxide. It breaks down into sodium carbonate and hydrogen peroxide when dissolved. Basically it provides the same effect as hydrogen peroxide but is much cheaper and has a much better shelf life. It acts as a strong oxidizer like bleach but without the chlorine. This means it produces no harmful byproducts when it breaks down, just water, oxygen, and sodium carbonate (soda ash). Like bleach it is very alkaline (10.5 for a 1% solution if wikipedia is right) and a strong oxidizer. It will oxidize certain things better than hypochlorite bleach and other things worse. Overall hypochlorite bleach is still considered stronger and is more cost effective. It breaks down much faster than hypochlorite bleach does. Full neutralization of a full strength solution occurs within 6 hours usually while hypochlorite bleach can take several days. Oxygen bleach is "color safe" with laundry and actually brightens some pigments (I have no idea why). It is also ridiculously cheap (around $2 a pound).

Sodium metabisulfite is a powerful reducing agent. It is moderately acidic (PH 4.3 at 1% concentration). In water it breaks down into sodiom bisulfite (NaHSO3) which when dissolving metals/minerals breaks down into sodium sulfite (Na2SO3), another metal sulfide (in this case CaSO3), and acid (H+). Sulfites are unstable and will oxidize into sulfates (SO3->SO4) in the presence of oxygen. Bleaches are divided into three main categories, chlorine based bleaches, oxygen based bleaches, and reducing bleaches (which are always based on sulfur compounds). The first two categories can be grouped into one category called "oxidizing bleaches". Textiles and other sensitive products that might be harmed by oxidizing bleaches are sometimes bleached with reducing bleaches instead. I've seen several websites state that oxidizing bleaches are better for removing "natural stains" while reducing bleaches are better for removing "manmade stains". Although I'm not sure what leads them to that conclusion chemically speaking. It is easy to confirm however that each is better/worse at tackling different stains. These reducing compounds are also commonly used to remove chlorine and oxygen from water. All of the dechlorinating products that you've ever used were likely using one of these sulfur based reducing agents. I was curious to see what if anything a reducing bleach would do to the rock. My choices of chemicals were sodium sulfite (Na2SO3), sodium metabisulfite (Na2S2O5), and sodium thiosulfate (Na2S2O3). Sodium dithionite (Na2S2O4) is a very popular and much more potent reducing agent than what's listed above but it was far too expensive for me to use. I chose sodium metabisulfite because it was the strongest reducing agent of the three and they all cost about the same (about $2 a pound). In acidic solutions all three of these compounds produce sulfur dioxide, an extremely powerful reducing agent and very toxic gas. You can think of SO2 as the opposite of Cl2 in terms of reactions. Sodium thiosulfate is the safest of the three. Sodium thiosulfate is neutral and sodium sulfite is alkaline (ph 9 at full saturation).

So what were the results? The sodium percarbonate did an amazing job. To my surprise it far outperformed the bleach. I had two rocks, a small 2 pound and a larger 7 pound, and I treated one with each. I used 1/4 dilution of 5% sodium hypochlorite bleach, brand new. So 1.25% concentration. Just like last time it caused a lot of the surface dirt and organic crust to peel off and some of it to dissolve. The water was definitely dirty and had a strong greenish brown tint. I let it soak for 24 hours in warm circulating water. A second treatment did nothing, the water simply remained green. The rock definitely looked whiter but nothing amazing. The white parts of the rock have a slightly yellow tint to them and the reddish brown parts (which I am currently assuming are iron stains) turn black/dark green. Since iron chloride is black/dark green this further backs up the theory that these stains are iron.

Now onto the percarbonate solution. First off this salt is pretty difficult to dissolve. You need to keep stirring it for quite some time and hot water is strongly recommended. It's a bit hard to tell when it has finished dissolving because the small bubbles that begin forming immediately make the solution cloudy until the reaction is finished. Once I put the rock in an eruption of bubbles began on the rocks surface. Like someone had stuck a strong air pump with a very fine airstone in the bucket. The solution quickly turned a thick dark brown. The surface began to foam just like it does in the hydrochloric acid bath. At exactly the 6 hour mark the reaction finished. I washed and I washed and I washed and the water kept coming back brown. Eventually I just stuck it in a bucket in the yard and left the hose on for a few hours until the rock was finally fully clean. The results were amazing. The brightest white I have ever seen on dry rock before. The hypochlorite bleached rock looks dirty by comparison. The iron stains turned a bright reddish brown and the corraline pigments came back! Patches of smooth bright pink and purple appeared right where the corraline patches were, as if the corraline was still alive. I can't explain how it happened but it looks really cool. I ran the chlorine bleached rock through a percarbonate treatment and the results were the same. Indicating that the percarbonate was removing a lot of organic material that the hypochlorite was not. I ran both rocks through a second treatment, this time at full strength (the first treatment was done at half strength). The water came back brown and foamy although not as bad as the first time. Washed them and then did it a third time. On the third treatment the solution came back slightly brown for the big rock and white for the small rock. Indicating that the oxygen bubbles being generated by the percarbonate deep in the rock build up in pressure and blast their way through to the surface causing some fine CaCO3 dust to come out of the rock and cloud the solution a bit even if their was no organic material to dissolve. So it seems that 2 treatments is recommended at the minimum. I believe the motion of the oxygen bubbles through the rocks pore network is part of the reason why it is so effective. It circulates fresh solution deep within the rock and blasts little bits of organic material off the surface area off the pores. It also seems to dissolve a lot of organic material that the bleach has no effect on deep in the rock and opens the pores up a bit more. I was never happy with the results I was getting with regular bleach since it didn't do too much to any dry organic material, just whitened them slightly without dissolving them. But I am very happy with this stuff. And it's much safer to use since it decomposes into soda ash, water, and oxygen.

Randy please read this paragraph I need your input here:

Next I tried the sodium metabisulfite. Since I didn't know what would happen I again tried it only on the little rock at first. At first things looked very promising. It was bubbling very slightly and the rock turned completely white. Those mineral stains that I'm assuming are iron completely went away and turned white. The corraline patches also turned white. Note that HCl, hypochlorite bleach, and percarbonate have no effect on these mineral stains. HCl dissolved some of it but also reveals new stains underneath other sections of the rock as it dissolves the rock. Then something bad happened. The rock texture started to become smooth and it started to turn yellow. Orangish/brownish crystals started to grow at the top and the bubbling speed increased. I pulled it out. I really want to know what happened here. My best guess is that the redox reaction turned the iron into a soluble salt but then once the acid had been used up the PH rose and caused the sulfur/iron to precipitate onto the rock. But this all happened within an hour. So how could the acid have been used up so fast? I made a 5% solution. If it was still acidic then why was precipitation occurring when the acid should have kept the sulfur in soluble form? I really want to know what happened here because its ability to remove the mineral stains when nothing else would shows promise. I want to try other sulfur compounds now but I'm not sure if I should aim for something more acidic or more alkaline next time. My other options are sodium sulfite, sodium thiosulfate (which currently looks like the most promising), sulfamic acid, and sodium bisulfate (unlikely to do anything more than dissolve some of the rock). If I knew what had happened here on a chemical level then I would have a better idea of what to test next. I would not recommend using this stuff even to neutralize bleach as it might stain the rock.

Now onto the acid bath. First the oxalic acid. The reaction was much slower than hydrochloric but faster than vinegar. It wasn't as fast as I had imagined it would be though. The solution remained clear/white indicating that any organic material was either not being dissolved or had already been removed by the percarbonate bath (more likely). It dissolved a tiny bit of the rock and whitened the iron stains a bit on the big rock. It didn't completely eliminate them though like the metabisulfite did. While I was happy with the results I would still prefer HCl to this. I would however dilute it much more than what people are usually using. 1/10 is way too much. It eats away way too much rock for no real benefit. 1/30 is probably the highest strength I would go. HCl is still a lot more cost effective (a 0.3% HCl solution was much stronger than my 5% ocalic acid solution) and has almost the same effect. The oxalic acid is a bit better at removing mineral stains though since it seems to have a slight preference for transition metals (iron/lead/copper, etc.) over alkaline/alkaline earth metals. The oxalic acid dissolved most of the sulfur stains very quickly. But there were a few small patches of darker colored sulfur stains that would not dissolve no matter how long I left it in there.

Next the hydrochloric bath. I made a 1/24 diluted (therefore 1.28% concentration) solution since that's all I had left in the bottle and stuck the small rock in first. To my surprise not much happened. Normally when I stick pukani rocks into an HCl bath of this strength I get the same reaction you see in this thread. Lots of bubbling and the solution quickly turns a dark brown color that's so thick that it almost looks like syrup. Big bubble foam begins to rise about a foot off the surface Until the reaction eventually begins to die down. But this time almost nothing happened. I got a small amount of bubbling and no change in solution color or foaming. At first I thought that maybe the HCl went bad even though that doesn't make sense chemically. Those small remaining sulfur stains that wouldn't come off in the oxalic acid still wouldn't come off in the HCl no matter how long I left it in. Weird. Then I put the big rock in. And this time the bubbling was A LOT more intense. Likely indicating that the pore surface area is a lot higher on the big rock. The little rock might have had its pores clogged by the metabisulfite somehow. But again no browning or foaming in the solution. Just lots of bubbling. Indicating that after the percarbonate bath there was no organic matter left for the HCl to remove. Or at least that's my interpretation. If true this means that the percarbonate does a better job of bleaching than actual bleach, a better job of organic matter dissolution than HCl, with no harmful byproducts, no harmful crossreactions, no significant dissolving of the rock., and much better coloration and porosity.

I won't have another batch of tests to run until I get some more rock on sale. Which likely won't be for another year.

TL : DR: I tested three new chemicals in my chemical cleaning process this time. Sodium percarbonate is amazing. You should use it instead of sodium hypochlorite (chlorine bleach). At least 2 6 hour baths are needed at full strength for it to be fully effective and it is difficult to dissolve so hot water is recommended. You might even consider skipping the acid bath if you use the percarbonate. Sodium metabisulfite was effective at removing mineral stains but had negative side effects and so isn't really an option. Oxalic acid was a bit better at removing mineral stains than hydrochloric acid but I still prefer hydrochloric due to its strength and cost effectiveness. I will test more chemicals the next time I get some new rock which likely won't be for awhile. I've taken chlorates and permangates off the list due to their cost.

Some relevant chemistry to what I think is happening with the sodium metabisulfite:

Sodium metabisulfite reacts with water to form sodium bisulfite:
Na2S2O5 + H2O = 2NaHSO3
Sodium bisulfite dissolves and reacts with calcium carbonate to form calcium sulfite, sodium sulfite, carbon dioxide, and water:
2NaHSO3 + CaCO3 = CaSO3 + Na2SO3 + CO2 + H2O
Calcium sulfite and sodium sulfite reacts with oxygen to form calcium sulfate and sodium sulfate:
CaSO3 + Na2SO3 + O2 = CaSO4 + Na2SO4
Carbon dioxide reacts with water to form carbonic acid:
CO2 + H2O = H2CO3
Carbonic acid dissolves and reacts with calcium carbonate to form calcium bicarbonate:
H2CO3 + CaCO3 <-> Ca(HCO3)2
therefore sodium metabisulfite, water, oxygen, and calcium carbonate react to form calcium sulfate, sodium sulfate and carbon dioxide:
Na2S2O5 + H2O + O2 + CaCO3 = CaSO4 + Na2SO4 + CO2

As the reaction slows down and the solution begins to lose acidity calcium sulfate and sodium sulfate will precipitate onto the calcium carbonate surface staining it yellow. The calcium bicarbonate will remain in solution unless you evaporate it, which will yield carbon dioxide and calcium carbonate. Once it reaches an equilibrium point it will start to break apart into CO2 and CaCO3 as new Ca(HCO3O)2 is formed. The excess CO3 will evaporate out of the solution (bubbling). A similar reaction also happens with sodium carbonate and sodium bicarbonate.
 
Wow. Thanks for such a detailed breakdown. I have about 250 lbs of dry live rock that has been sitting in a crate in my shed for 5 years. Most of it is discolored Just how much of the percarbonate per gallon for a 100% strength wash?
 
Thanks for the information! That's all very interesting. I'd be interested in trying the oxalic acid and the sodium percarbonate. I've never liked the idea of using hydrochloric acid. It's a bit too toxic for me to be happy handling it. The sodium percarbonate might be more lung-friendly than the bleach, which is important to me.

I'm not enough of a chemist to comment much on the reactions involved. Maybe some of the real chemists can comment.
 
I am jumping on this thread as a newbie. I have about 100 pounds of rock in bleach currently. After i finish cleaning and rinsing etc.... DO i cure it separately with a shrimp or ammonia? or do i put it my 90 gallon that currently has 2 small clowns and ten pounds of healthy live rock?
 
I probably would use a bit of shrimp or fish food to cure it separately, but ammonia works, too. You can add it to the 90g if you're sure that the bleach has removed all the organic debris. You could test that by neutralizing the bleach and letting the rock sit with a powerhead for a few days, and see whether any ammonia shows up.
 
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