Differences Between Iron and Copper

[onBalance]

Various metals in pool water should not be lumped together in how they react in pool water conditions. The following is my limited understanding regarding metals in pool water, and would welcome comments with different information.

I agree that copper is less soluble at a higher pH than at a lower pH, and therefore, will precipitate out some copper as the pH rises and solubility lowers. However, I do not believe there is a major solubility difference between a pH of 7.4 and 7.8 in terms of how much dissolved copper will precipitate out. My concern is that at a lower pH, it is more likely that copper will be dissolved from a heat exchanger and place copper into solution, thereby increasing it.

I do not believe that adding chlorine causes copper to precipitate out (as long as the pH remains the same). In fact, I believe that high chlorine residuals actually makes copper slightly more soluble.

If the copper content is well below its' solubility, raising the pH may not cause any precipitation of copper.

I suggest that "Iron" behaves much differently as compared to copper. Any dose of chlorine will begin to cause dissolved iron in water to become oxidized and precipitate out. In fact, iron will also become oxidized and precipitate out due to oxygen in the water, whether or not chlorine is added, and regardless of the pH (within reason). A high or low pH (in normal pool pH water conditions) does not play a significant role in how it reacts to chlorine and oxygen in water.

I believe that manganese and cobalt behave similar to iron, and silver and zinc behave similar to copper.

Generally, there are usually pros and cons to everything. Sometimes, one just has to decide what their priorities are. For me, I weighed the above and have decided that I prefer a slightly higher pH than what the pool industry's ideal of 7.4 to 7.6 suggests. Even though copper is less soluble at a higher pH, it is therefore less likely to be dissolved from copper heat exchangers.

As far as adding sequestrants, if there are no metals in the pool water or fill water, I see them mostly as not needed. If there is iron or copper in the fill water or pool water, then adding them can beneficial, but will need periodic additions because they may break down in chlorinated water and stop working. For start-ups especially, adding a sequestrant makes sense. Much of the iron in well or tap water can be oxidized by chlorine and get filtered out before it stains pool surfaces.

 

[chem geek]

A pH of 7.8 has about 2.5 times higher hydroxyl ion concentration than a pH of 7.4 so if you are on the edge of precipitating copper at 7.4, you could precipitate it at 7.8. It all depends on the concentration of copper in the water. Because the saturation point for copper is dependent on the square of the hydroxyl ion concentration, the amount of copper needed to precipitate is over 6 times lower at pH 7.8 than at pH 7.4. For iron, it is over 15 times lower.

Ferrous hydroxide is soluble at several hundred ppm at usual pool pH, but ferric hydroxide is very insoluble even at lower pH. Oxidizers convert ferrous ion to ferric ion -- chlorine does this readily, but as you point out oxygen does this as well though more slowly. So it's with copper where one is much closer to the edge with copper and is how at the somewhat lower pH one is able to use copper algaecide without it precipitating.

So I would say it isn't true that pH plays an important role in copper precipitation. It does, at least when the copper is in the ballpark of saturation. For iron, it's more about having chlorine in the water. It takes fairly low amounts of iron to get iron to precipitate. Whether it stains or not is more complicated. So you are largely correct that the pool pH level isn't as critical with iron as it is with copper. Of course, if one is at the saturation point with iron, then raising the pH can result in more precipitation, but if that iron is then removed, then the pH changes shouldn't make much difference.

 

[onBalance]

Thanks for the additional information Chem geek.

In regards to staining from iron and copper, I have not been able to determine for certainty whether low or high pH helps to prevent staining or not. I tend to believe that high pH helps to prevent staining as compared to low pH's. But since iron and copper react and behave differently in chlorinated water, perhaps they also react differently to pH as far as staining pool plaster?

It is my understanding that copper can or will combine (react) with carbonate. On that basis, I am thinking that dissolved copper (in low pH) may react with a plaster surface and attaching to the plaster (calcium carbonate) surface and turning a turquoise color. Perhaps if a high pH causes copper to precipitate out and not attach to the plaster, then that copper precipitate may just get filtered out.

And perhaps a similar reaction scenario for iron?

 

[chem geek]

Well for copper at least most copper staining occurs from use of copper sulfate algaecide products where the copper concentration is either too high or the pH gets high. So it looks like the copper will stain the plaster from high pH in that case. Both carbonate ion and hydroxyl ion are in higher concentration at higher pH. At a TA of 80 ppm and a pH of 7.5, copper carbonate is over 370 times more likely to precipitate or stain compared to copper hydroxide. So you are correct that copper staining is more likely to be copper carbonate than copper hydroxide, though it theoretically would require around 3 ppm copper to precipitate. At a TA of 120 ppm and a pH of 8.0, however, it could start to precipitate/stain at 0.65 ppm. In practice, the pH and carbonate concentrations are higher close to plaster surfaces since circulation is not perfect there.

It's with iron (in the ferric form that occurs when chlorine is present) that it precipitates so readily (i.e. it's much less soluble) that it may be more likely to get caught in a filter than to stain a plaster surface. Obviously one can have rust show up with exposed rebar, but for iron in the bulk pool water it's not so clear that it would tend to stain first. Nevertheless, we have seen high iron content fill water from wells that resulted in stains and we've had people doing ascorbic acid treatments where if they raise the pH too quickly and didn't add enough metal sequestrant, then they can get re-staining.

 

[onBalance]

Again, thanks for the good information.
In regards to copper staining, I am wondering that since copper sulfate is soluble and generally a blue color, and copper carbonate is not very soluble and is blue-green (turquoise) color, perhaps the copper from soluble copper sulfate dissociates and then the copper ions actually combines with (or reacts) with the calcium carbonate plaster surface and become copper carbonate as part of the plaster?

On the other hand, if there is a precipitation of copper carbonate occurring in the water (not on the surface), then I can see that as not staining the plaster surface, and the copper carbonate gets filtered out. Am I wrong in considering these two possible scenarios?

In regards to the iron, I agree and have experienced times when the iron is precipitated by chlorine addition (and water turns yellow-green) and then stains the plaster, and sometimes I have seen where precipitated iron (iron colored water) simply gets filtered out without staining. Haven't completely determined why the difference.

 

[onBalance]

I should add that the reason I believe dissolved copper is reacting with the plaster (to form copper carbonate) is because it requires a lot of sanding or acid washing to remove copper stains. Generally, it seems that a layer of the plaster surface has to be removed in order to get all of the copper stain removed, especially on older copper stained pools.

Whereas, iron seems to be a thin deposit and is always easily removed by simple and quick sanding or a light acid or ascorbic treatment. It seems that ferrous iron is being precipitated into ferric iron and sticking to the plaster surface, whereas I suspect copper is often reacting chemically with a calcium carbonate (plaster) surface.

 

[chem geek]

Copper ions have a pale blue color in water (technically, it's a copper hydrate ion -- see this link for a more detailed explanation). You don't need a molecular precipitate to have a color. So don't think that a color in the water means there is a precipitate -- it could just means there are transition metal ions in the water.

Yes you are correct that because copper carbonate is the likely source of the stain and because calcium carbonate is in the plaster surface, this can make the stain more difficult to remove since it is in a sense bound into the plaster itself where copper replaces calcium (see this paper for more technical details). And yes, if it were to form copper carbonate in the water away from the plaster, it would be more likely to get filtered out, but according to the paper, copper carbonate is not a common precipitate because copper hydroxide is more stable (not sure if this is due to ion pairs that increase copper carbonate effective solubility). This implies that with copper, you're more likely to get staining of plaster than precipitation into the filter.

With ferric iron, it will form the hydroxide (or an oxide hydrate) before it forms a carbonate so it is not bound into the plaster in the same way and is more of a surface stain.

 

[chem geek]

I just ran into this link describing how paint bonds to materials and thought it appropriate for this discussion since I believe that iron staining is a form of adhesion by adsorption via secondary or van der Waals forces. There might also be some mechanical interlocking for thick stains such as rust from rebar. I don't see it as likely for ferric iron chemically bonding to the plaster.

Copper stains, on the other hand, may be more from adhesion by chemical bonding if the copper displaces calcium and bonds to the carbonate directly. The chemical bonding makes them harder to remove.

 

[onBalance]

I agree that ferric iron isn't reacting (binding) with a plaster surface. It seems to simply "stick" to the surface as its own precipitated deposit.

It appears that dissolved copper ions in the water chemically binds to the calcium carbonate plaster surface. In fact, I have observed that older copper stains often show a deterioration of the plaster surface. The copper color goes very deep and the plaster is somewhat soft wherever the turquoise color is. That suggests a break-down of the harder plaster surface caused by the copper/carbonate chemical binding reaction. I have not seen Iron stains do the same to a plaster surface.

 

[onBalance]

A couple interesting things about iron staining. My experience has been that precipitated iron generally "sticks" to plastic or PVC fittings first before it does to a plaster surface. Much of the iron residue can be rub off of plastic, but usually needs some diluted acid to remove all of it. I concluded years ago that iron, having a positive charge (cations), is attracted to plastic due to its negative charge.

Another thing I remember going to back to the 1970's when the Arneson Pool Sweep (which floated on top of the water) was the major pool cleaner, is that the Pool Sweep unit would travel around and rub along a plaster pool wall (one foot below the water line), and we would sometimes see a light yellow stain about two inches wide around the entire pool. Apparently the plastic "rim" of the Pool Sweep would transfer its own negative charge (electrons) to the plaster surface, which then attracted precipitated iron.

 

[Patrick_B]

I think that makes two people on this board who remember the "Pool Sweep" I swam in a pool during the early/mid 80's and the big thing was avoiding that dang thing when using the diving board. the other was using the tails to squirt everyone in and outside the pool.

Interesting theory about the transfer. I don't think I would have thought of that.

 

[onBalance]

Yea, I also had a lot of fun squirting people with those water whips (or tails) many years ago. Of course, my kids often returned fire and got the best of me.

It seems to me that the concept of higher pH's causes metal staining may not always be true. Especially with iron, since chlorine and oxygen will oxidize all dissolved iron in the pool water at any pH fairly quick. Therefore, raising the pH should not be an issue with iron.

With copper, it may be that dissolved copper (perhaps copper sulfate) at normal pool water pH's may be attaching and reacting with the plaster surface (primarily calcium carbonate) and turning the plaster surface blue/green. Therefore, keeping the pH lower may not prevent that type of copper staining.

 

[chem geek]

It all depends on concentration. At a given copper ion concentration and pool pH, even a more alkaline (higher pH) plaster surface isn't going to stain with copper unless the copper ion concentration is high enough. If one raises the pool pH, then the likelihood of such staining increases dramatically. So regardless of the fact one may not be at copper carbonate saturation in the bulk pool water, it would be wrong to assume that the likelihood of copper staining is not related to pH. Higher pH in the bulk pool water will result in even higher pH at the pool plaster surface since the pH there is related to the pH in the bulk pool water via diffusion.

 

[onBalance]

Chem geek, perhaps I am wrong, but it seems are you talking more in terms of solubility and precipitation of copper carbonate from the water? I understand that copper (and iron) become less soluble as the pH rises. And I understand that the higher the concentration, the more likely precipitation occurs due to the lack of solubility.

But what about reactions with calcium carbonate in the plaster as opposed to in the water? (There is virtually no carbonate existing in water with a pH of 7.5).

If dissolved copper (any low concentration) in pool water has an attraction or affinity for carbonate, it is not possible that the dissolved copper can simply attach or react with the carbonate and become copper carbonate and become past of the plaster, as opposed to a deposit on top of the plaster surface as a precipitate?

 

[chem geek]

If the water is not saturated with copper carbonate, then any copper displacing a calcium in the calcium carbonate crystal structure will be temporary. The reason is that at the surface of the solid crystal there is dissolving and reforming constantly going on but with the water not saturated in copper carbonate then statistically over time there will be less and less copper in the crystal. Just as calcium carbonate scale can dissolve when the water is not saturated with calcium carbonate, the same is true for copper carbonate stains but even more true with what you describe which is a surface dislodging of calcium with copper. That copper will likely get dislodged again by a calcium. The definition of saturation is when there is an equilibrium between the solid dissolving vs. reforming.

 

[onBalance]

I mentioned above that older copper stains are deep into the plaster matrix, not easily removed with acid washing or sanding (as opposed to iron deposit stains which remains at surface). To me, that suggests that the copper is transferring (or traveling) deeper and deeper into the plaster surface over time. I know that bicarbonate and carbon dioxide reacts with calcium hydroxide in the plaster surface and forms calcium carbonate (known as carbonation), and does not easily re-dissolve or reverse itself. But I know that formation is different from copper carbonate forming.

But I am not so sure that copper carbonate is also very easily dissolved and reversed. I also know that calcium carbonate that precipitates onto a plaster surface is not easily dissolved. It seems to requires fairly aggressive conditions to make an impact. And as we know, that calcium scaling doesn't seem to manifest itself until the CSI gets up to at least to +0.5.

I suppose if the copper concentration is slowly increasing over time, then the copper carbonate at surface may not re-dissolve. I know that there is usually an exchange involving most chemical reactions, but is it possible that there is no displacement of the calcium ion and the copper simply reacts and attaches itself to the carbonate? I also don't understand how the water can be saturated with copper carbonate when there is very little carbonate in water below a pH of 8.0.

So what about the iron issue? Even though iron is less soluble at higher pH's, since it gets oxidized easily at any pH, pool owners shouldn't worry about a higher pH causing iron staining?

 

[chem geek]

The depth of the copper penetration is irrelevant if the product of copper and carbonate concentrations in the water exceeded the solubility product. When that is the case, they can precipitate so that means they could do so by penetration of the plaster if it is porous. That doesn't prove anything about getting copper to stain when the copper concentration is lower. If you keep the copper concentration and the pH at levels so that the product of copper and carbonate concentrations are below the solubility product, then you should not get significant staining. That is how many pools using copper sulfate work -- they try and maintain the copper level to not get too high and keep the pH from getting too high as well.

"Saturated" in a chemical equilibrium sense doesn't mean both components of a pair are high in concentration. It is the product of their concentrations that determines whether there will be a net dissolving or a net growth of solid. It's a statistical thing. Also, for chemical solids with very low solubility, it doesn't take very much to begin precipitation. Copper carbonate has much lower solubility than calcium carbonate. That is why calcium carbonate doesn't scale at usual pool pH unless the calcium level gets very high towards 1000+ ppm. For copper carbonate under similar conditions, it only takes a little more than 1 ppm copper ions to stain (though the precise levels depends on pH since that determines the carbonate concentration).

You say that you don't understand how the water can be saturated with copper carbonate when there is very little carbonate below pH 8.0, but the same can be said for calcium carbonate yet you know that precipitates in spite of the low carbonate level below a pH of 8.0. With a very high CH level, one can still get scaling at a pH of 7.5. The difference between calcium and copper is that copper carbonate has a much lower solubility so it takes a lot less copper to cause copper carbonate to precipitate.

It comes down to thermondynamics which predicts what can happen (not how quickly -- that's kinetics) where chemical systems tend toward their lowest free energy state where free energy is a combination of enthalpy (heat energy) and entropy (randomness vs. ordered). Copper and carbonate bind to each other as a solid fairly strongly and therefore have lower free energy than calcium carbonate. Nevertheless, there are ions and molecules ramming into solids dislodging atoms and because the calcium concentration is so much higher than copper, it is likely for the copper to eventually get dislodged and replaced by calcium. There's also water ramming into the solid dislodging (dissolving) it. The surface is not at all static -- it is constantly being modified at a microscopic level. When the water is reasonably saturated with calcium carbonate, the dislodging of an atom often results with it getting replaced so that there is no net change, but over time the surface will get rougher even if the average thickness does not change.

 

[onBalance]

It seems we agree that copper ions can react and combined directly to the plaster itself.
I have drained hundreds of pools and acid washed them. I see where iron staining is mostly a thin deposit on top of smooth plaster. It the plaster is severely etched or porous, the iron staining does penetrate and is more difficult to remove. Copper always seems to penetrate even smooth, hard, and dense (non-porous) plaster and become part of the plaster itself. So I believe there are differences between the conditions of copper staining versus iron staining. I also mention again that iron will deposit onto plastic pool fittings and components. It is interesting that copper does not seem to do that in normal conditions.

I have drained many pools with copper stains that don't disappear upon re-filling with copper free water. I am sure it helps to keep the amount of copper below the solubility level, and yes, it is a very small amount compared to the solubility of calcium carbonate (which is about 14 ppm). Yes, chemicals and metals are bouncing and ramming around and making some changes, but copper staining seems to be rather permanent. Even lowing the pH to 1.0 doesn't always remove it easily. Iron comes off smooth and dense plaster easily. Copper doesn't.