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- Jun 22, 2009
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I was reading Sandy King's article "Vandyke Brown: A Method for Making Permanent Prints in Gold Metal" about Vandyke Brown prints on unblinkingeye.com : https://unblinkingeye.com/Articles/GTV/gtv.html ...
As usual, great article, but something caught my eye...
Sandy King writes the following :
I of course totally agree with this, the chemical reactions involved being the following (correct me if I'm wrong) :
(in French, ferriques means ferric and ferreux means ferrous...)
Then Sandy King writes the following :
This is what caught my eye... : I don't think this is possible for ferrous iron to oxydize the silver metal, but once again correct me if I'm wrong...
A redox reaction between Fe²⁺ and metallic silver Ag cannot occur spontaneously under standard conditions. If we examine the standard reduction potentials of the species involved:
- Fe²⁺ + 2e⁻ → Fe, with E° = -0.44 V
- Ag⁺ + e⁻ → Ag, with E° = +0.80 V
The higher the reduction potential, the more likely the species is to gain electrons (i.e., be reduced), right ? Here, silver ions Ag⁺ are much more easily reduced to Ag than Fe²⁺ is to Fe.
For Fe²⁺ to oxidize metallic silver Ag, silver would have to be oxidized to Ag⁺ and Fe²⁺ would have to be reduced to Fe. However, metallic silver is very stable, and because of the high reduction potential of Ag⁺, it does not easily oxidize back to Ag⁺.
Thus, if I understand correctly, there is no driving force for Fe²⁺ to be reduced to Fe while oxidizing Ag to Ag⁺ and I come to the conclusion that no spontaneous redox reaction can occur between Fe²⁺ and metallic silver Ag. The potential difference between the species indicates that Fe²⁺ cannot oxidize silver metal.
On the contrary, a redox reaction can occur between (ferric) Fe³⁺ and metallic silver (Ag). If we look at the standard reduction potentials of the species involved:
- Fe³⁺ + e⁻ → Fe²⁺, with E° = +0.77 V
- Ag⁺ + e⁻ → Ag, with E° = +0.80 V
In this case, the reduction potential of Ag⁺ is slightly higher than that of Fe³⁺, meaning Ag metal can be oxidized to Ag⁺, and Fe³⁺ can be reduced to Fe²⁺.
The reaction would look like this:
Fe³⁺ + Ag → Fe²⁺ + Ag⁺
Since the reduction of Fe³⁺ to Fe²⁺ and the oxidation of Ag to Ag⁺ are close in potential, this reaction can happen spontaneously.
So, is Sandy King's article wrong, with confusion between ferrous and ferric ions ?
I'd be tempted to answer yes, since one can read in Mike Ware's article about "The Argyrotype Process" (another silver-iron process) : https://www.mikeware.co.uk/mikeware/Argyrotype_Process.html :
Your thougths ?
As usual, great article, but something caught my eye...
Sandy King writes the following :
I of course totally agree with this, the chemical reactions involved being the following (correct me if I'm wrong) :
(in French, ferriques means ferric and ferreux means ferrous...)
Then Sandy King writes the following :
This is what caught my eye... : I don't think this is possible for ferrous iron to oxydize the silver metal, but once again correct me if I'm wrong...
A redox reaction between Fe²⁺ and metallic silver Ag cannot occur spontaneously under standard conditions. If we examine the standard reduction potentials of the species involved:
- Fe²⁺ + 2e⁻ → Fe, with E° = -0.44 V
- Ag⁺ + e⁻ → Ag, with E° = +0.80 V
The higher the reduction potential, the more likely the species is to gain electrons (i.e., be reduced), right ? Here, silver ions Ag⁺ are much more easily reduced to Ag than Fe²⁺ is to Fe.
For Fe²⁺ to oxidize metallic silver Ag, silver would have to be oxidized to Ag⁺ and Fe²⁺ would have to be reduced to Fe. However, metallic silver is very stable, and because of the high reduction potential of Ag⁺, it does not easily oxidize back to Ag⁺.
Thus, if I understand correctly, there is no driving force for Fe²⁺ to be reduced to Fe while oxidizing Ag to Ag⁺ and I come to the conclusion that no spontaneous redox reaction can occur between Fe²⁺ and metallic silver Ag. The potential difference between the species indicates that Fe²⁺ cannot oxidize silver metal.
On the contrary, a redox reaction can occur between (ferric) Fe³⁺ and metallic silver (Ag). If we look at the standard reduction potentials of the species involved:
- Fe³⁺ + e⁻ → Fe²⁺, with E° = +0.77 V
- Ag⁺ + e⁻ → Ag, with E° = +0.80 V
In this case, the reduction potential of Ag⁺ is slightly higher than that of Fe³⁺, meaning Ag metal can be oxidized to Ag⁺, and Fe³⁺ can be reduced to Fe²⁺.
The reaction would look like this:
Fe³⁺ + Ag → Fe²⁺ + Ag⁺
Since the reduction of Fe³⁺ to Fe²⁺ and the oxidation of Ag to Ag⁺ are close in potential, this reaction can happen spontaneously.
So, is Sandy King's article wrong, with confusion between ferrous and ferric ions ?
I'd be tempted to answer yes, since one can read in Mike Ware's article about "The Argyrotype Process" (another silver-iron process) : https://www.mikeware.co.uk/mikeware/Argyrotype_Process.html :
Your thougths ?