subsecret
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Reduction of Copper (II) Compounds
I thought that this topic would be pertinent to my current work with copper compounds, specifically copper (II) compounds.
I understand that it is important not to dispose of copper compounds down the drain, as they are very toxic to aquatic life. Doing this could
potentially lead to a disaster caused by killing the bacteria in one's septic tank.
Now for my question:
Will more reactive metals (aluminum or magnesium for instance) create a faster reaction with a copper (II) solution than less reactive ones?
I've got a beaker of various copper (II) waste that's reacting with some iron in the form of steel wool. For safety, it's been placed inside a larger
beaker in case of a foam-over, though I'm not sure if this reaction would be fast enough.
Additionally, are iron (II) compounds safe to dispose of down the drain? Should they be converted to less harmful salts?
Any help is appreciated. Thank you.
Fear is what you get when caution wasn't enough.
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bfesser
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<a href="https://www.sciencemadness.org/whisper/viewthread.php?tid=25000#copper"><strong>Topical Compendium</strong>
#copper</a>
Specifically, check out the topics under "organic redox" and "toxicity." Personally, I prefer to reduce copper(II) waste with <a
href="https://en.wikipedia.org/wiki/Ascorbic_acid" target="_blank">L-ascorbic acid</a> <img src="../scipics/_wiki.png" /> to
precipitate Cu<sup>0</sup> for re-use.
[Edited on 3.1.14 by bfesser]
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Metacelsus
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Copper (II) sulfate is used to remove roots that are clogging drains, so it can't be that bad.
http://www.roebic.com/pdf/msds/K-77-MSDS.pdf
However, it does say not to use it if you have a septic tank, and that it is toxic to aquatic life.
I suggest precipitating it with sodium carbonate/bicarbonate instead of reducing it. That way, you can recover the copper (II) easily. However, I
found that this method does not work well if the copper is complexed.
If you want to reduce it, I'd go with scrap aluminum. Because it is oxidized to Al (III) and has a relatively low atomic mass, it's more effective
gram for gram than iron.
Iron (II) should pose less of a problem than copper (II), but it (at least the sulfate) is still "moderately toxic to aquatic invertebrates and
slightly toxic to fish".
http://www.epa.gov/oppsrrd1/REDs/factsheets/4058fact.pdf
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bfesser
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Quote: Originally posted by Cheddite Cheese  | Copper (II) sulfate is used to remove roots that are clogging drains, so it can't be that bad.
. . .
However, it does say not to use it if you have a septic tank, and that it is toxic to aquatic life. |
This is
one of my pet peeves; downplaying the detrimental impact of excess Cu<sup>2+</sup> on aquatic systems—it can be downright
devastating. Copper(II) sulfate is still sold to remove roots from drains, because old plumbers would throw a fit if it were to suddenly disappear
from shelves. Just something has been done a certain way for a long time does not make it right. Beside the environmental impact, there's no reason
not to recover <em>all</em> copper waste, as it's valuable, and recovery is facile.
[Edited on 9.8.13 by bfesser]
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cyanureeves
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i thought i saw copper sulphate on a bottle of algae control for koi ponds although it was in tiny amounts.
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subsecret
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Thanks, bfesser, for that very helpful page. Where do you get your ascorbic acid? Additionally, I certainly agree with your comment
on the extreme toxicity of copper to aquatic life. My house has a septic tank, so I have to be especially careful of how I manage waste.
Cheddite, I assume that you're recommending that I precipitate my copper (II) with NaHCO3 to obtain copper (II) carbonate. I've tried this,
yet the leftover solution retains a blue color, even when NaHCO3 is in excess. It does precipitate, but it's hard to separate from the
water, even when a "professional grade" filter paper. Precipitation via aluminium may be the best option.
Thank you for the quick responses.
Fear is what you get when caution wasn't enough.
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bfesser
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<strong>cyanureeves</strong>, <a href="http://en.wikipedia.org/wiki/Copper(II)_sulfate"
target="_blank">CuSO<sub>4</sub>·5H<sub>2</sub>O</a> <img src="../scipics/_wiki.png" /> is often
sold (reasonably pure) for that use. As far as I've read, it kills many microbes and bacteria, though—not just the algae. It's also
harmful or toxic to higher forms of aquatic life.
<strong>Awesomeness</strong>, I can't tell if you're being sarcastic, so I'll just give myself the benefit of the doubt and assume that
you're not. I'm glad you found that page to be helpful. I have a huge bottle of analytical reagent (AR) grade L-ascorbic acid
(C<sub>6</sub>H<sub>8</sub>O<sub>6</sub> from
Mallinkrodt, but I work in microscale (often sub-mg quantities). I'm sure you can find plenty of over the counter (OTC) sources for it, and I know
others have posted about this—sorry, I don't recall where.
I recommend Na<sub>2</sub>CO<sub>3</sub> over NaHCO<sub>3</sub>—less mess (<strong>Eq.
1</strong> & <strong>Eq. 2</strong>. I've never seen either
precipitate copper quantitatively (neither with ascorbate reduction). Aluminium is relatively cheap and could work (as would Zn and plenty of other
metals; see <strong>Eq. 3</strong>, but the Cu<sup>0</sup>
would require further purification and processing to get back a Cu<sup>2+</sup> salt.
<strong>Eq. 1. CuSO<sub>4</sub>(aq) + 2 NaHCO<sub>3</sub>(aq)
→ [Cu(HCO<sub>3</sub><sub>2</sub>] +
Na<sub>2</sub>SO<sub>4</sub>(aq) → CuCO<sub>3</sub>(s) +
Na<sub>2</sub>SO<sub>4</sub>(aq) + CO<sub>2</sub>(g) + H<sub>2</sub>O</strong>
<strong>Eq. 2. CuSO<sub>4</sub>(aq) +
Na<sub>2</sub>CO<sub>3</sub> → CuCO<sub>3</sub>(s) +
Na<sub>2</sub>SO<sub>4</sub>(aq)</strong>
<strong>Eq. 3. Zn<sup>0</sup> + Cu<sup>2+</sup>(aq)
→ Zn<sup>2+</sup>(aq) + Cu<sup>0</sup></strong>
Ca(OH)<sub>2</sub> might work well for this purpose (<strong>Eq. 4</strong>, as Ca<sup>2+</sup> is essentially harmless to the environment—limestone is a calcite
(CaCO<sub>3</sub> rock which dissolves in slightly acidic ground water,
forming caves [sorry, going off topic into my geology interest]. You could even boil off the water to recover CaSO<sub>4</sub>, if you're
so inclined. NaOH could work, but it's important to note that copper becomes soluble again under certain conditions (high pH; <a
href="http://webpages.charter.net/dawill/tmoranwms/Chem_Cu.html#Cuprate">reference</a> <img src="../scipics/_ext.png" />.
<strong>Eq. 4. CuSO<sub>4</sub>(aq) + Ca(OH)<sub>2</sub>(aq)
→ Cu(OH)<sub>2</sub>(s) + CaSO<sub>4</sub>(aq)</strong>
Of course, the aqueous chemistry of copper is not as simple as I've indicated above, but there's no need to go into that level of detail
here—abstraction is nice.
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subsecret
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Thank you. And you've assumed correctly - I was not being sarcastic.
Fear is what you get when caution wasn't enough.
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