Ag + H2O2 Reactions
Mike Monett
I'm trying to identify the residue after evaporating a solution of
silver ions in water.
I use H2O2 to dissolve pieces of 12 ga silver wire to produce silver
ions in solution. Ordinary 3% drugstore H2O2 gives what appears to be
60 ppm Ag(+) according to the following equation:
2Ag + 2H2O2 --> 2Ag(+) + O2(g) + 2H2O
To test for silver ions, adding pickling salt gives a very strong
white dispersion of silver chloride:
Ag(+) + NaCl --> AgCl(ppt) + Na(+)
and adding a small amount of ammonia causes the dispersion to
disappear:
AgCl + 2NH3 --> Ag(NH3)2(+) + Cl(-)
So it's clear the solution contains silver ions.
Now, if I generate another solution and let it evaporate, a small
residue remains in the bottom of the glass. It is difficult to tell
the color of the residue, but it might be gray.
Adding H2O2 to the glass produces bubbles at a slow rate. You have to
look under a microscope to see them, and it takes a while before the
bubbles stop.
Adding salt to the resulting solution produces a strong dispersion,
and adding ammonia clears it again. So the evaporated residue appears
to be silver, or a silver compound.
Now if it's silver, where did it get the electrons needed during
evaporation to become elemental silver again? I believe the only
source of electrons is the water, but it's not clear what reaction
could free an electron from H2O.
Another possible explanation is the evaporated residue is a silver
oxide, AgO. This is gray and decomposes to silver and oxygen at 100C:
http://tinyurl.com/bkm2p (webelements: AgO)
The reaction with H2O2 might be
AgO + H2O2 --> Ag(+) + O2 + 2H2O
but I'd expect it to go much faster.
It doesn't appear to be Ag2O, since this oxide is black or dark brown:
http://tinyurl.com/7g7sm (webelements: Ag2O)
but it's difficult to tell the color even under a microscope.
Any ideas on how to identify this residue?
Thanks,
Mike Monett
Wilco Oelen
Mike Monett wrote:
> Hi,
>
> I'm trying to identify the residue after evaporating a solution of
> silver ions in water.
>
> I use H2O2 to dissolve pieces of 12 ga silver wire to produce silver
> ions in solution. Ordinary 3% drugstore H2O2 gives what appears to be
> 60 ppm Ag(+) according to the following equation:
Why 60 ppm?
>
> 2Ag + 2H2O2 --> 2Ag(+) + O2(g) + 2H2O
No, this cannot be a correct equation. Charge is not balanced!
>
> To test for silver ions, adding pickling salt gives a very strong
> white dispersion of silver chloride:
>
> Ag(+) + NaCl --> AgCl(ppt) + Na(+)
>
> and adding a small amount of ammonia causes the dispersion to
> disappear:
>
> AgCl + 2NH3 --> Ag(NH3)2(+) + Cl(-)
>
> So it's clear the solution contains silver ions.
OK, assuming that you have silver ions in the liquid (you convincingly
have shown that with NaCl and NH3) and the solution of H2O2 is fairly
pure, then the only thing I can imagine is the following:
2Ag + H2O2 --> 2Ag(+) + 2OH(-)
The bubbles of O2 you see are due to a side reaction. H2O2 is
decomposed easily by catalytic processes. Probably Ag(+) or OH(-)
somewhat enhances the rate of decomposition of H2O2.
The compound AgOH is somewhat soluble. It does not exist in the solid
state (not as a dry compound at least), but in solution in pure water
it can exist at a concentration of approximately 13 ppm.
Frequently, commercial H2O2 does contain a small amount of an acidic
stabilizer (e.g. H3PO4). With this, the solubility of Ag(+) may even
increase somewhat and the above reaction will be as follows:
2Ag + H2O2 + 2H(+) --> 2Ag(+) + 2H2O
The H(+) is from the stabilizer of the H2O2.
>
> Now, if I generate another solution and let it evaporate, a small
> residue remains in the bottom of the glass. It is difficult to tell
> the color of the residue, but it might be gray.
>
> Adding H2O2 to the glass produces bubbles at a slow rate. You have to
> look under a microscope to see them, and it takes a while before the
> bubbles stop.
Again, probably catalytic decomposition of H2O2. No real reaction
between the residue and the H2O2. The residue just slowly dissolves.
>
> Adding salt to the resulting solution produces a strong dispersion,
> and adding ammonia clears it again. So the evaporated residue appears
> to be silver, or a silver compound.
OK, I can agree with you if you say it is a silver compound.
>
> Now if it's silver, where did it get the electrons needed during
> evaporation to become elemental silver again? I believe the only
> source of electrons is the water, but it's not clear what reaction
> could free an electron from H2O.
I don't think that the gray stuff you got is elemental silver. My guess
is that it is a mix of Ag2O and the silver salt of the stabilizer in
the H2O2, e.g. AgH2PO4/Ag2HPO4. There may be traces of elemental silver
in the mix, but only traces, because many silver compounds are somewhat
light sensitive.
If you really would get elemental silver, then you would get a deep
black residue. Finely divided silver is black. This is the basis of
black and white photography. The photo-images consist of finely divided
metallic silver, dispersed in a gelatin layer.
You could do a test. Crunch a tablet of vitamin C and add to water.
Shake for a while and let the white stuff settle. Decant some of the
clear solution (take care not to get the white stuff with it) in a
beaker and add some NaOH (or Na2CO3) and dissolve all. Add this
solution with NaOH/vitamin C to the gray residue. If this is a silver
compound it will turn black at once. The black stuff is elemental
silver.
>
> Another possible explanation is the evaporated residue is a silver
> oxide, AgO. This is gray and decomposes to silver and oxygen at 100C:
AgO is non-existent. There is a compound Ag2O2, which can best be
described as Ag(+)Ag(3+)O2. This compound can only be formed under very
strong oxidizing conditions. When persulfate, S2O8(2-), is added to a
solution containing Ag(+), then a brown solution is obtained (which
contains Ag(3+), either as free ion, or comhined with Ag(+), I do not
know that). On addition of some alkali, a black precipitate of Ag2O2 is
obtained. Indeed, this precipitate is quite unstable and gives off
oxygen slowly.
>
> http://tinyurl.com/bkm2p (webelements: AgO)
>
> The reaction with H2O2 might be
>
> AgO + H2O2 --> Ag(+) + O2 + 2H2O
Again, not possible, no charge balance. Ag2O2 probably will oxidize
H2O2 to water and oxygen:
Ag2O2 + H2O2 --> Ag2O + H2O + O2 or
Ag2O2 + H2O2 --> 2AgOH(dissolved) + O2
>
> but I'd expect it to go much faster.
Not, if the slow bubbling is just due to slight catalytic enhancement
of the decomposition of H2O2: 2H2O2 --> 2H2O + O2.
>
> It doesn't appear to be Ag2O, since this oxide is black or dark
brown:
No, the grey stuff will not be pure Ag2O, that indeed is almost black
(I have some from a photography supplier and it is a very dark brown
powder).
>
> http://tinyurl.com/7g7sm (webelements: Ag2O)
>
> but it's difficult to tell the color even under a microscope.
>
> Any ideas on how to identify this residue?
See above.
>
> Thanks,
>
> Mike Monett
Hope this helps somewhat. Success,
Wilco
Mike Monett
> Mike Monett wrote:
>> Hi,
>> I'm trying to identify the residue after evaporating a solution
>> of silver ions in water.
>> I use H2O2 to dissolve pieces of 12 ga silver wire to produce
>> silver ions in solution. Ordinary 3% drugstore H2O2 gives what
>> appears to be 60 ppm Ag(+) according to the following equation:
> Why 60 ppm?
Using a Hanna 98308 PWT (Pure Water Tester), the original
conductivity of the H2O2 measured 85.5uS, probably due to the
stabilizers.
After the bubbles stopped, the conductivity was 141.5uS, which is an
increase of 56uS. (The 98308 PWT is spec'd to 99uS, but it
overranges quite well.)
From previous work on the electrolysis of silver, a solution of
Ag(+) and OH(-) ions has a conductivity very close to 1uS per ppm.
In this case, the OH(-) ions are missing, so the PWT reads low. The
actual silver concentration is higher than 56 ppm, but I don't know
how much.
>> 2Ag + 2H2O2 --> 2Ag(+) + O2(g) + 2H2O
> No, this cannot be a correct equation. Charge is not balanced!
Yes. I know. But as you will see, it is the only one that makes
sense:)
>> To test for silver ions, adding pickling salt gives a very strong
>> white dispersion of silver chloride:
>> Ag(+) + NaCl --> AgCl(ppt) + Na(+)
>> and adding a small amount of ammonia causes the dispersion to
>> disappear:
>> AgCl + 2NH3 --> Ag(NH3)2(+) + Cl(-)
>> So it's clear the solution contains silver ions.
> OK, assuming that you have silver ions in the liquid (you
> convincingly have shown that with NaCl and NH3) and the solution
> of H2O2 is fairly pure, then the only thing I can imagine is the
> following:
> 2Ag + H2O2 --> 2Ag(+) + 2OH(-)
> The bubbles of O2 you see are due to a side reaction. H2O2 is
> decomposed easily by catalytic processes. Probably Ag(+) or OH(-)
> somewhat enhances the rate of decomposition of H2O2.
Well, I am very confident there is no OH(-) in the solution.
Otherwise, the Ag(+) would combine to form AgOH (silver hydroxide)
when it dries. I see this quite often in my work.
AgOH is dark brown/black and reacts very vigorously with H2O2. The
residue I get in the above is sort of gray, and it reacts very
slowly with H2O2.
> The compound AgOH is somewhat soluble. It does not exist in the
> solid state (not as a dry compound at least), but in solution in
> pure water it can exist at a concentration of approximately 13
> ppm.
I know the references state AgOH is soluble to between 13.3 ppm and
17 ppm. However, I have done considerable work on this, and it turns
out AgOH is insoluble in pure water. It is very stable in a wet or
dry state and decomposes to silver at 100C. I prove this in
http://escribe.com/health/thesilverlist/m78851.html
and
http://escribe.com/health/thesilverlist/m79117.html
So the solubility listed in the references is wrong.
> Frequently, commercial H2O2 does contain a small amount of an
> acidic stabilizer (e.g. H3PO4). With this, the solubility of Ag(+)
> may even increase somewhat and the above reaction will be as
> follows:
> 2Ag + H2O2 + 2H(+) --> 2Ag(+) + 2H2O
> The H(+) is from the stabilizer of the H2O2.
Interesting. What are the complete equations, and how is hydrogen
released from the compound?
I understand different vendors may use different stabilizers. How do
we know which stabilizer is used in the Walmart H2O2 that I buy?
>> Now, if I generate another solution and let it evaporate, a small
>> residue remains in the bottom of the glass. It is difficult to
>> tell the color of the residue, but it might be gray.
>> Adding H2O2 to the glass produces bubbles at a slow rate. You
>> have to look under a microscope to see them, and it takes a while
>> before the bubbles stop.
> Again, probably catalytic decomposition of H2O2. No real reaction
> between the residue and the H2O2. The residue just slowly
> dissolves.
>> Adding salt to the resulting solution produces a strong
>> dispersion, and adding ammonia clears it again. So the evaporated
>> residue appears to be silver, or a silver compound.
> OK, I can agree with you if you say it is a silver compound.
>> Now if it's silver, where did it get the electrons needed during
>> evaporation to become elemental silver again? I believe the only
>> source of electrons is the water, but it's not clear what
>> reaction could free an electron from H2O.
> I don't think that the gray stuff you got is elemental silver. My
> guess is that it is a mix of Ag2O
Ag2O is black:
http://tinyurl.com/7g7sm (webelements: Ag2O)
> and the silver salt of the stabilizer in the H2O2, e.g.
> AgH2PO4/Ag2HPO4.
That would make a lot of sense. What would happen when you add H2O2?
> There may be traces of elemental silver in the mix, but only
> traces, because many silver compounds are somewhat light
> sensitive.
> If you really would get elemental silver, then you would get a
> deep black residue. Finely divided silver is black. This is the
> basis of black and white photography. The photo-images consist of
> finely divided metallic silver, dispersed in a gelatin layer.
I realize photography relies on silver becoming black. However, when
you heat AgOH, it decomposes to finely divided pure silver, and it
is gray in color.
> You could do a test. Crunch a tablet of vitamin C and add to
> water. Shake for a while and let the white stuff settle. Decant
> some of the clear solution (take care not to get the white stuff
> with it) in a beaker and add some NaOH (or Na2CO3) and dissolve
> all. Add this solution with NaOH/vitamin C to the gray residue. If
> this is a silver compound it will turn black at once. The black
> stuff is elemental silver.
Excellent - can you list the equations so I don't make a mistake in
understanding the process?
>> Another possible explanation is the evaporated residue is a
>> silver oxide, AgO. This is gray and decomposes to silver and
>> oxygen at 100C:
> AgO is non-existent. There is a compound Ag2O2, which can best be
> described as Ag(+)Ag(3+)O2. This compound can only be formed under
> very strong oxidizing conditions. When persulfate, S2O8(2-), is
> added to a solution containing Ag(+), then a brown solution is
> obtained (which contains Ag(3+), either as free ion, or comhined
> with Ag(+), I do not know that). On addition of some alkali, a
> black precipitate of Ag2O2 is obtained. Indeed, this precipitate
> is quite unstable and gives off oxygen slowly.
Very interesting. But webelements lists AgO:
>> http://tinyurl.com/bkm2p (webelements: AgO)
How could they do that if it doesn't exist?
>> The reaction with H2O2 might be
>> AgO + H2O2 --> Ag(+) + O2 + 2H2O
> Again, not possible, no charge balance. Ag2O2 probably will
> oxidize H2O2 to water and oxygen:
> Ag2O2 + H2O2 --> Ag2O + H2O + O2 or
> Ag2O2 + H2O2 --> 2AgOH(dissolved) + O2
AgOH is insoluble.
Where is the Ag(+)? The solution gives an AgCl dispersion when salt
is added, so we need silver ions.
>> but I'd expect it to go much faster.
> Not, if the slow bubbling is just due to slight catalytic
> enhancement of the decomposition of H2O2: 2H2O2 --> 2H2O + O2.
Again, adding H2O2 to the gray residue makes silver ions.
>> It doesn't appear to be Ag2O, since this oxide is black or dark
>> brown:
> No, the grey stuff will not be pure Ag2O, that indeed is almost
> black (I have some from a photography supplier and it is a very
> dark brown powder).
>> http://tinyurl.com/7g7sm (webelements: Ag2O)
>> but it's difficult to tell the color even under a microscope.
>> Any ideas on how to identify this residue?
> See above.
>> Thanks,
>> Mike Monett
> Hope this helps somewhat. Success,
> Wilco
Excellent, Wilco. Thanks for your help.
Best Wishes
Mike Monett
Wilco Oelen
Mike Monett wrote:
<snipped>
> Using a Hanna 98308 PWT (Pure Water Tester), the
original
> conductivity of the H2O2 measured 85.5uS, probably due to
the
> stabilizers.
>
> After the bubbles stopped, the conductivity was 141.5uS, which is
an
> increase of 56uS. (The 98308 PWT is spec'd to 99uS, but
it
> overranges quite well.)
>
> From previous work on the electrolysis of silver, a solution
of
> Ag(+) and OH(-) ions has a conductivity very close to 1uS per ppm.
>
> In this case, the OH(-) ions are missing, so the PWT reads low.
The
> actual silver concentration is higher than 56 ppm, but I don't
know
> how much.
Probably the anion of the stabilizer. So, I do not think there is much
difference in conductivity (of course the other anion may have other
mobility, resulting in other conductivity, but you cannot state that
there is no counterion).
>
> >> 2Ag + 2H2O2 --> 2Ag(+) + O2(g) + 2H2O
>
> > No, this cannot be a correct equation. Charge is not balanced!
>
> Yes. I know. But as you will see, it is the only one that
makes
> sense:)
No, this cannot make sense, whatever way you look at it. There MUST be
somewhere, where the electrons are going.
<snipped for brevity>
> > The bubbles of O2 you see are due to a side reaction. H2O2
is
> > decomposed easily by catalytic processes. Probably Ag(+) or
OH(-)
> > somewhat enhances the rate of decomposition of H2O2.
>
> Well, I am very confident there is no OH(-) in the
solution.
> Otherwise, the Ag(+) would combine to form AgOH (silver
hydroxide)
> when it dries. I see this quite often in my work.
>
> AgOH is dark brown/black and reacts very vigorously with H2O2.
The
> residue I get in the above is sort of gray, and it reacts
very
> slowly with H2O2.
OK, if these are your observations, then I think you are right. So
let's assume that there is no OH(-). This can only be the case if there
is an appreciable amount of stabilizer in the H2O2. Then, the gray
residue you get is the silver salt of the stabilizer (possible silver
(di)hydrogenphosophate).
>
> > The compound AgOH is somewhat soluble. It does not exist in
the
> > solid state (not as a dry compound at least), but in solution
in
> > pure water it can exist at a concentration of approximately
13
> > ppm.
>
> I know the references state AgOH is soluble to between 13.3 ppm
and
> 17 ppm. However, I have done considerable work on this, and it
turns
> out AgOH is insoluble in pure water. It is very stable in a wet
or
> dry state and decomposes to silver at 100C. I prove this in
>
> http://escribe.com/health/thesilverlist/m78851.html
>
> and
>
> http://escribe.com/health/thesilverlist/m79117.html
>
> So the solubility listed in the references is wrong.
>
Are you sure? I'm not yet fully convinced. In the experiments, the ions
simply may go too fast from anode to cathode, so that no such buildup
of concentration occurs. Why would all references, books etc. be wrong
with such familiar compounds like Ag(+)/OH(-)?
> > Frequently, commercial H2O2 does contain a small amount of
an
> > acidic stabilizer (e.g. H3PO4). With this, the solubility of
Ag(+)
> > may even increase somewhat and the above reaction will be
as
> > follows:
>
> > 2Ag + H2O2 + 2H(+) --> 2Ag(+) + 2H2O
>
> > The H(+) is from the stabilizer of the H2O2.
>
> Interesting. What are the complete equations, and how is
hydrogen
> released from the compound?
Hydrogen peroxide is more stable in somewhat acidic media than in
alkaline media. This is why frequently some acid is added to hydrogen
peroxide. Most common is the addition of a small amount of phosphoric
acid: H3PO4.
H3PO4 is a moderately strong acid and the first proton H+ is split off
farily easily:
H3PO4 <--> H(+) + H2PO4(-) The equalibrium is fairly much to the right
at low concentration, as is the case in stabilized H2O2.
If the amount of stabilizer is appreciable (more than a few tens of
ppm), then indeed you will not get much (if any at all) AgOH, but then
you will get AgH2PO4 (assuming that the stabilizer is H3PO4).
>
> I understand different vendors may use different stabilizers. How
do
> we know which stabilizer is used in the Walmart H2O2 that I buy?
This will be very hard, if you do not have reagents for analysing this
H2O2. A very common stabilizer is phosphoric acid, combined with a
small amount of sodium pyrophosphate. The phosphoric acid makes the
liquid a little bit acidic and the pyrophosphate is used as a
sequestering agent, which forms complexes with many metal ions, which
might increase the decomposition rate by means of catalytic processes.
If you add H2O2, then I do not think that much will happen. The stuff
may slowly dissolve and enhance the decomposition rate of H2O2 at
little bit.
>
> > There may be traces of elemental silver in the mix, but
only
> > traces, because many silver compounds are somewhat
light
> > sensitive.
>
> > If you really would get elemental silver, then you would get
a
> > deep black residue. Finely divided silver is black. This is
the
> > basis of black and white photography. The photo-images consist
of
> > finely divided metallic silver, dispersed in a gelatin layer.
>
> I realize photography relies on silver becoming black. However,
when
> you heat AgOH, it decomposes to finely divided pure silver, and
it
> is gray in color.
Are you sure? How pure was the AgOH you heated. Are you sure that no
other silver salts were in it. If it is contaminated with other silver
salts, then heating will result in something gray. Metallic silver,
mixed with the other silver salt, which probably will be white. So, the
mix will be gray.
>
> > You could do a test. Crunch a tablet of vitamin C and add
to
> > water. Shake for a while and let the white stuff settle.
Decant
> > some of the clear solution (take care not to get the white
stuff
> > with it) in a beaker and add some NaOH (or Na2CO3) and
dissolve
> > all. Add this solution with NaOH/vitamin C to the gray residue.
If
> > this is a silver compound it will turn black at once. The
black
> > stuff is elemental silver.
>
> Excellent - can you list the equations so I don't make a mistake
in
> understanding the process?
Vitamin C in alkaline environments is oxidized very easily. Ag(+) and
almost any silver salt is sufficiently strong oxidizer.
In alkaline environments, vitamin C forms so called ascorbates:
C6H8O6 + OH(-) --> C6H7O6(-) + H2O
C6H7O6(-) is a strong reductor. It hence can be regarded as electron
donor and this results in formation of silver:
Ag(+) + e --> Ag(black)
>
> >> Another possible explanation is the evaporated residue is
a
> >> silver oxide, AgO. This is gray and decomposes to silver
and
> >> oxygen at 100C:
>
> > AgO is non-existent. There is a compound Ag2O2, which can best
be
> > described as Ag(+)Ag(3+)O2. This compound can only be formed
under
> > very strong oxidizing conditions. When persulfate, S2O8(2-),
is
> > added to a solution containing Ag(+), then a brown solution
is
> > obtained (which contains Ag(3+), either as free ion, or
comhined
> > with Ag(+), I do not know that). On addition of some alkali,
a
> > black precipitate of Ag2O2 is obtained. Indeed, this
precipitate
> > is quite unstable and gives off oxygen slowly.
>
> Very interesting. But webelements lists AgO:
>
> >> http://tinyurl.com/bkm2p (webelements: AgO)
>
> How could they do that if it doesn't exist?
This formula gives just the ratio of silver and oxygen in the compound.
So, from a stoiciometric point of view, the compound can be called AgO,
but when the compound is analysed in more detail, then there is strong
evidence for presence of Ag(+) and Ag(3+) in the compound. The formula
AgO suggests that this contains Ag(2+), so that is why more accurate
descriptions of this compound preferrably use Ag2O2 as formula.
>
> >> The reaction with H2O2 might be
>
> >> AgO + H2O2 --> Ag(+) + O2 + 2H2O
>
> > Again, not possible, no charge balance. Ag2O2 probably
will
> > oxidize H2O2 to water and oxygen:
>
> > Ag2O2 + H2O2 --> Ag2O + H2O + O2 or
> > Ag2O2 + H2O2 --> 2AgOH(dissolved) + O2
>
> AgOH is insoluble.
Well, a small amount of AgOH will go into solution as Ag(+) and OH(-),
the rest remains as hydrous solid in the liquid. When hydroxides and
oxides are prepared from aqueous solutions, then the difference between
them frequently is not clear at all.
An hydroxide can also be regarded as a hydrated oxide. In many
situations, even more complex structure are formed with OH groups
bridging between metal centers. So frequently, the formulas MOH and M2O
(in case of a univalent metal) are just oversimplified approximations
of reality.
>
> Where is the Ag(+)? The solution gives an AgCl dispersion when
salt
> is added, so we need silver ions.
>
> >> but I'd expect it to go much faster.
>
> > Not, if the slow bubbling is just due to slight
catalytic
> > enhancement of the decomposition of H2O2: 2H2O2 --> 2H2O + O2.
>
> Again, adding H2O2 to the gray residue makes silver ions.
Just another suggestion. Create some of the gray residue and add some
white vinegar (without any H2O2). The vinegar added, must be absolutely
free of salt, otherwise this test does not work. If you obtain a
liquid, which gives a white precipitate with salt, then you know for
sure that the gray residue is not silver, but a silver salt.
The vinegar contains CH3COOH, which splits into CH3COO(-) and H(+) to
some extent. This H(+) helps dissolving the residue.
<snipped for brevity>
Let me know what happens,
Wilco
Mike Monett
> Mike Monett wrote:
[...]
>>> Frequently, commercial H2O2 does contain a small amount of an
>>> acidic stabilizer (e.g. H3PO4). With this, the solubility of
>>> Ag(+) may even increase somewhat and the above reaction will be
>>> as follows:
>>> 2Ag + H2O2 + 2H(+) --> 2Ag(+) + 2H2O
>>> The H(+) is from the stabilizer of the H2O2.
>> Interesting. What are the complete equations, and how is hydrogen
>> released from the compound?
> Hydrogen peroxide is more stable in somewhat acidic media than in
> alkaline media. This is why frequently some acid is added to
> hydrogen peroxide. Most common is the addition of a small amount
> of phosphoric acid: H3PO4.
> H3PO4 is a moderately strong acid and the first proton H+ is split
> off farily easily:
> H3PO4 <--> H(+) + H2PO4(-) The equalibrium is fairly much to the
> right at low concentration, as is the case in stabilized H2O2.
> If the amount of stabilizer is appreciable (more than a few tens
> of ppm), then indeed you will not get much (if any at all) AgOH,
> but then you will get AgH2PO4 (assuming that the stabilizer is
> H3PO4).
[...]
>>> AgO is non-existent. There is a compound Ag2O2, which can best
>>> be described as Ag(+)Ag(3+)O2. This compound can only be formed
>>> under very strong oxidizing conditions. When persulfate,
>>> S2O8(2-), is added to a solution containing Ag(+), then a brown
>>> solution is obtained (which contains Ag(3+), either as free ion,
>>> or comhined with Ag(+), I do not know that). On addition of some
>>> alkali, a black precipitate of Ag2O2 is obtained. Indeed, this
>>> precipitate is quite unstable and gives off oxygen slowly.
>> Very interesting. But webelements lists AgO:
>>>> http://tinyurl.com/bkm2p (webelements: AgO)
>> How could they do that if it doesn't exist?
> This formula gives just the ratio of silver and oxygen in the
> compound. So, from a stoiciometric point of view, the compound can
> be called AgO, but when the compound is analysed in more detail,
> then there is strong evidence for presence of Ag(+) and Ag(3+) in
> the compound.
Interesting. Can you give some links to show how this analysis is
done?
> The formula AgO suggests that this contains Ag(2+), so that is why
> more accurate descriptions of this compound preferrably use Ag2O2
> as formula.
Why couldn't the Ag be doubly ionized? I thought we were supposed to
use the simplest common formula... Also, AgO is much more popular.
Google shows only 103 hits for Ag2O2:)
[...]
> Just another suggestion. Create some of the gray residue and add
> some white vinegar (without any H2O2). The vinegar added, must be
> absolutely free of salt, otherwise this test does not work.
How do I get salt-free vinegar?
> If you obtain a liquid, which gives a white precipitate with salt,
Do you mean to let it evaporate again?
> then you know for sure that the gray residue is not silver, but a
> silver salt. The vinegar contains CH3COOH, which splits into
> CH3COO(-) and H(+) to some extent. This H(+) helps dissolving the
> residue.
> <snipped for brevity>
> Let me know what happens,
> Wilco
Thanks, I will try your experiment when I find out how to get
salt-free vinegar.
In the meantime, I did another experiment. I added 30mL H2O2 to a
glass and added 4 short pieces of 12 ga silver wire. The bubbles
started immediately.
The I added ordinary 5.25% household bleach (sodium hypochlorite:
NaOCl). This caused a great deal of fizzing as it converted the H2O2
to oxygen and water:
H2O2 + NaOCl --> H2O + NaCl + O2
I kept adding bleach and finally had to toss about half the solution
to make room for more bleach. The silver wire bubbled a great deal
while I was adding the bleach.
Eventually, the bleach killed the H2O2 and the fizzing stopped.
But the bubbles kept coming from the silver wire!
So it appears you were right. The H2O2 is a sideshow - it is the
stabilizers that are converting the silver to ions.
Best Wishes,
Mike Monett
Wilco Oelen
Mike Monett wrote:
> > This formula gives just the ratio of silver and oxygen in
the
> > compound. So, from a stoiciometric point of view, the compound
can
> > be called AgO, but when the compound is analysed in more
detail,
> > then there is strong evidence for presence of Ag(+) and Ag(3+)
in
> > the compound.
>
> Interesting. Can you give some links to show how this analysis
is
> done?
from the book "chemistry of the elements" by Greenwood and Earnshaw,
page 1181. The diamagnetism of the oxide "AgO", and diffraction studies
showing that there are two different kinds of silver ions in this
oxide, with oxide ions grouped around them in different ways, give
evidence for Ag(I)Ag(III)O2.
>
> > The formula AgO suggests that this contains Ag(2+), so that is
why
> > more accurate descriptions of this compound preferrably use
Ag2O2
> > as formula.
>
> Why couldn't the Ag be doubly ionized?
Well it could, but in this case it isn't...
> I thought we were supposed to
> use the simplest common formula...
Who told you to use the simplest common formula? Why do you write H2O2
and not HO? I think, because you know that the smallest entity, being
hydrogen peroxide consists of *two* hydrogens and *two* oxygens. The
compound OH is a completely different compound (it is an extremely
reactive radical, called hydroxyl, which only can exist transiently as
reaction intermediate or at extremely low pressures).
A similar reasoning holds for Ag2O2. The simplest "entity" which
completely describes this compound consists of one Ag(+), one Ag(3+),
and two O(2-) ions in a single crystal lattice. This simplest
description has two Ag-ions and two O-ions. This is somewhat comparable
with H2O2, but now the compound is not molecular but ionic.
I *can* go with you some way, that Ag2O2 is not a very good formula.
Unfortunately there is not a nice notation for what this compound
really is. What about AgAgO2? The two Ag's are not equivalent. A
similar situation exists for the thiosulfate ion, S2O3(2-). The two
S-atoms are not equivalent in this ion, but in simple formulas there
simply is no better notation than S2O3(2-).
> ... Also, AgO is much more popular.
> Google shows only 103 hits for Ag2O2:)
More complex explanations (which in this case really are a better
approximation of reality) are not always the most popular
explanations...
>
> [...]
>
> > Just another suggestion. Create some of the gray residue and
add
> > some white vinegar (without any H2O2). The vinegar added, must
be
> > absolutely free of salt, otherwise this test does not work.
>
> How do I get salt-free vinegar?
Try photography suppliers. They sell acetic acid stop bath. Sometimes
as 28% acetic acid, sometimes as glacial acetic acid. Household
cleaning vinegar (not to be confused with vinegar for consumption) may
also be a good source. It contains (at least here in the Netherlands)
approximately 10% acetic acid. It can be purchased at supermarkets.
I partially agree with you. The H2O2 is just a sideshow in dissolving
the grey residue you have after evaporating your liquid. The
stabilizers dissolve the residue.
It does however take a crucial role in your initial dissolving of the
silver wire. Here it works *together* with the stabilizers. The H2O2
does the oxidation from Ag to Ag(+) and the (acidic) stabilizers help
the Ag(+) dissolve instead of forming a precipitate of AgOH or Ag2O.
>
> Best Wishes,
>
> Mike Monett
Mike Monett
> Mike Monett wrote:
[...]
>> I thought we were supposed to use the simplest common formula...
> Who told you to use the simplest common formula? Why do you write
> H2O2 and not HO? I think, because you know that the smallest
> entity, being hydrogen peroxide consists of *two* hydrogens and
> *two* oxygens.
Very good point. No, H2O2 is already strange enough - we don't need
to change the name:)
I found your comments in the Copper acetate thread very relevant and
help illuminate this topic a great deal:
--------------------------------------------------------------------
"I've discovered that chemistry is a science of approximations.
The formulas and equations frequently are not such precise
descriptions of what happens in reality. They help us understand
the phenomena to a certain extent, but sometimes (frequently?) one
has to refine the coarse approximations."
"Notorious examples of where the high school simple chemistry does
not give sufficient explanation is in transition metal chemistry
(e.g. copper, iron, titanium) and in the formation of
precipitates."
"When you see something described like Fe(3+) + 3OH(-) -->
Fe(OH)3, then in reality a tremendously complex process is
executed, which can only be described by more advanced equations.
There simply is not such a compound like Fe(OH)3. E.g.
Fe-(u-OH)-Fe bridges can be formed or even much more complex
structures, with a lot of Fe, OH and H2O units."
"The only thing, known for sure, is that in the limit for very
large numbers of Fe(3+) units, OH(-) units and H2O, the ratio of
Fe(3+) and OH(-) must be 1:3 in order to keep charge balanced."
--------------------------------------------------------------------
Your last point makes good sense, and it already answered a question
I was going to ask about the ratios.
> The compound OH is a completely different compound (it is an
> extremely reactive radical, called hydroxyl, which only can exist
> transiently as reaction intermediate or at extremely low
> pressures).
The hydroxyl ions also form during silver electrolysis and are quite
stable. They can persist for years in solution with the silver ions,
but they cause a gradual drop in conductance as the Ag(+) ions
slowly combine with the OH(-) to form AgOH.
I made a simple demo using red cabbage juice that allows you to
visualize how the Ag(+) and OH(-) ions diffuse in the solution:
http://escribe.com/health/thesilverlist/m61491.html
The hydroxyl ions eventally turned the entire solution dark purple,
so you couldn't see anything. But another demo allowed you to trace
the path of the Ag(+) ions as they crossed the distance between the
electrodes:
http://escribe.com/health/thesilverlist/m61527.html
> A similar reasoning holds for Ag2O2. The simplest "entity" which
> completely describes this compound consists of one Ag(+), one
> Ag(3+), and two O(2-) ions in a single crystal lattice. This
> simplest description has two Ag-ions and two O-ions. This is
> somewhat comparable with H2O2, but now the compound is not
> molecular but ionic.
> I *can* go with you some way, that Ag2O2 is not a very good
> formula. Unfortunately there is not a nice notation for what this
> compound really is. What about AgAgO2? The two Ag's are not
> equivalent. A similar situation exists for the thiosulfate ion,
> S2O3(2-). The two S-atoms are not equivalent in this ion, but in
> simple formulas there simply is no better notation than S2O3(2-).
Thanks, Wilco. This is very good information.
[...]
>>> Just another suggestion. Create some of the gray residue and add
>>> some white vinegar (without any H2O2). The vinegar added, must
>>> be absolutely free of salt, otherwise this test does not work.
>> How do I get salt-free vinegar?
> Try photography suppliers. They sell acetic acid stop bath.
> Sometimes as 28% acetic acid, sometimes as glacial acetic acid.
> Household cleaning vinegar (not to be confused with vinegar for
> consumption) may also be a good source. It contains (at least here
> in the Netherlands) approximately 10% acetic acid. It can be
> purchased at supermarkets.
I will try to find some and see what happens.
[...]
>> In the meantime, I did another experiment. I added 30mL H2O2 to a
>> glass and added 4 short pieces of 12 ga silver wire. The bubbles
>> started immediately.
>> The I added ordinary 5.25% household bleach (sodium hypochlorite:
>> NaOCl). This caused a great deal of fizzing as it converted the
>> H2O2 to oxygen and water:
>> H2O2 + NaOCl - > H2O + NaCl + O2
>> I kept adding bleach and finally had to toss about half the
>> solution to make room for more bleach. The silver wire bubbled a
>> great deal while I was adding the bleach.
>> Eventually, the bleach killed the H2O2 and the fizzing stopped.
>> But the bubbles kept coming from the silver wire!
>> So it appears you were right. The H2O2 is a sideshow - it is the
>> stabilizers that are converting the silver to ions.
> I partially agree with you. The H2O2 is just a sideshow in
> dissolving the grey residue you have after evaporating your
> liquid. The stabilizers dissolve the residue.
> It does however take a crucial role in your initial dissolving of
> the silver wire. Here it works *together* with the stabilizers.
> The H2O2 does the oxidation from Ag to Ag(+) and the (acidic)
> stabilizers help the Ag(+) dissolve instead of forming a
> precipitate of AgOH or Ag2O.
Yes, I was a bit exuberant in my first comment:)
Wilco, thank you very much for your help. You have answered my
questions on how the silver ions are produced, and I am pleased to
have a much greater understanding than before.
Best Wishes,
Mike Monett
Wilco Oelen
Mike Monett wrote:
<snipped for brevity>
> Wilco Oelen wrote:
> > The compound OH is a completely different compound (it is
an
> > extremely reactive radical, called hydroxyl, which only can
exist
> > transiently as reaction intermediate or at extremely
low
> > pressures).
>
> The hydroxyl ions also form during silver electrolysis and are
quite
> stable. They can persist for years in solution with the silver
ions,
> but they cause a gradual drop in conductance as the Ag(+)
ions
> slowly combine with the OH(-) to form AgOH.
about the hydroxide ion, OH(-), I am talking about the hydroxyl
molecule, uncharged OH. Hydroxide indeed is stable. The single charge
makes a whole world of a difference!
>
> I made a simple demo using red cabbage juice that allows you
to
> visualize how the Ag(+) and OH(-) ions diffuse in the solution:
>
> http://escribe.com/health/thesilverlist/m61491.html
>
> The hydroxyl ions eventally turned the entire solution dark
purple,
> so you couldn't see anything. But another demo allowed you to
trace
> the path of the Ag(+) ions as they crossed the distance between
the
> electrodes:
>
> http://escribe.com/health/thesilverlist/m61527.html
Looks interesting. This is a really nice way to visualize such
electrolysis processes. Just a suggestion... If you have a digital
camera, include some pictures of the most special situations. You can
have nice JPEG images at acceptable detail (e.g. 640x480) in less than
40 kByte, so, even through telephone lines at 33k6 or 56k the page can
be viewed reasonably well.
Wilco
<snipped for brevity>
Mike Monett
>
> Mike Monett wrote:
> <snipped for brevity>
> > Wilco Oelen wrote:
> > > The compound OH is a completely different compound (it is
> an
> > > extremely reactive radical, called hydroxyl, which only can
> exist
> > > transiently as reaction intermediate or at extremely
> low
> > > pressures).
> > The hydroxyl ions also form during silver electrolysis and are
> quite
> > stable. They can persist for years in solution with the silver
> ions,
> > but they cause a gradual drop in conductance as the Ag(+)
> ions
> > slowly combine with the OH(-) to form AgOH.
> We are talking about completely different entities. You are talking
> about the hydroxide ion, OH(-), I am talking about the hydroxyl
> molecule, uncharged OH. Hydroxide indeed is stable. The single charge
> makes a whole world of a difference!
That's funny. Your description seemed to fit the hydroxyl ion perfectly. I
thought the missing minus sign was just a typo:)
Just a note on terminology - isn't the OH(-) ion called the hydroxyl ion?
> > I made a simple demo using red cabbage juice that allows you
> to
> > visualize how the Ag(+) and OH(-) ions diffuse in the solution:
> >
> > http://escribe.com/health/thesilverlist/m61491.html
> >
> > The hydroxyl ions eventally turned the entire solution dark
> purple,
> > so you couldn't see anything. But another demo allowed you to
> trace
> > the path of the Ag(+) ions as they crossed the distance between
> the
> > electrodes:
> >
> > http://escribe.com/health/thesilverlist/m61527.html
> Looks interesting. This is a really nice way to visualize such
> electrolysis processes. Just a suggestion... If you have a digital
> camera, include some pictures of the most special situations. You can
> have nice JPEG images at acceptable detail (e.g. 640x480) in less than
> 40 kByte, so, even through telephone lines at 33k6 or 56k the page can
> be viewed reasonably well.
>
> Wilco
Photos would indeed be nice. The problem is the Ag(+) reaction is
extremely faint and may be difficult to photograph. It took a lot of
fiddling with different sheets of paper in the background to see it
visually. But it would be worth the effort.
Mike Monett
Wilco Oelen
Mike Monett wrote:
> Just a note on terminology - isn't the OH(-) ion called the hydroxyl
ion?
"hydroxyl ion". In documents, where both the ion and the molecule are
mentioned, the ion is called "hydroxide" and the molecule "hydroxyl".
For an example see the section on hydrogen peroxide in
http://www.astaxanthin.org/oxidation.htm
So, I think that both names can be used for the ion. For the molecule
only the name hydroxyl is suitable.
Wilco
Mike Monett
Thanks for the interesting link, Wilco.
You were right about this molecule only surviving at low pressures - like in a
vacuum. It is used in astronomy since it radiates at several frequencies near
1720MHz and is part of the "water hole" along with hydrogen. And all this time, I
thought they were talking about the OH(-) ion!
Mike Monett