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I tried Fe/ Acetic acid and Fe/ Ammonium chloride/EtOH and in both the cases I ended up with 0.6 % of the Des halo compound. My specification is less than 0.1 %. Can anybody suggest the best method for this problem.
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I have replied this question many times in different platforms. But now I have at least three different videos on it. How to protect halogen like bromo and Iodo or chloro while I am doing nitro to amine. Here I am putting a link of youtube. You can look for Nitro reduction with Zn, Nitro reduction with Iron-Beckhamp reduction and Finally you can see the video on Tertiary carbon nitro reduction to amine while sensitive functional groups are present.
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Hi everyone!
Unfortunately, we're missing our sodium thioglycolate stock and i urgently need to make some fresh medium for my Desulfovibrio vulgaris. If I understand correct, sodium thioglycolate is used as a reducing agent and therefore may be replaced by another (similar) reducing agent.
Any thoguhts on this?
By the way, we have already left out the iron sulfate from the original medium because we wanted to avoid FeS2 precipitation.
Maybe you could suggest some reading sources that would help me solve this?
Thanks a lot!
artur
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L-Cysteine-Sulfide Reducing Agent:
Composition per 20.0mL:
L-Cysteine·HCl·H2O ..................................................................... 0.3g
Na2S·9H2O.................................................................................... 0.3g
Preparation of L-Cysteine-Sulfide Reducing Agent: Add Lcysteine
·HCl·H2O to 10.0mL of distilled/deionized water. Mix thoroughly.
In a separate tube, add Na2S·9H2O to 10.0mL of distilled/deionized
water. Mix thoroughly. Gas both solutions with 100% N2 and
cap tubes. Autoclave both solutions for 15 min at 15 psi pressure–
121°C using fast exhaust. Cool to 50°C. Aseptically combine the two
solutions under 100% N2.
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I am working on Vanadium redox flow batteries (VRFB). In other redox flow batteries there is a problem of cross contamination of redox species across membrane.
I am using nafion 112 in VRFB and using V(IV) as catholyte and V(III) as anolyte both in 3M H2SO4 as supporting electrolyte.
So my question was that although in both compartment, vanadium ions are present across membrane and also they are in different oxidation states then why do cross contamination is less as compared to other redox flow batteries ?
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In vanadium redox flow battery only one element acts as active material for both positive and negative side. So, due to crossover if catholyte moves to anolyte and/or vice-versa then it doesn't contaminate the system. In stead of vanadium if other element would be used, then it could contaminate the electrolyte. For example, contamination might be some unwanted or irreversible side reaction. Due to crossover of active material, one will loose capacity which can be retrieved further. That's why VRFB system is called cross-contamination free.
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Hello colleagues,
We are trying to induce a liner consequences reaction of nitrification follows de nitrification processes. The main challenge is the redox potential. we need to reduce it from 250 mv by the nitrification chamber into -50 mv by the anoxic denitrification chamber. We are using a potassium carbonate to reduce the redox potential. Any new ideas how to reduce the redox potential?
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Thank you Dr. Murphy!
I just sent you an e-mail for a discussion.
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Dear Researchers,
We have found a hydrothermal association of biotite-sulfide (chalcopyrite-pyrrhotite-Co-rich sulfides)-uraninite-graphite-allanite with minor tourmaline and xenotime. We are trying to understand the redox conditions (semi quantitative or quantitative measurements) for this association. We will appreciate and acknowledge your insights and suggestions on this. Following are the relevant details.
Ti-in-biotite thermometry suggest that the crystallization temperature of this assemblage is 370-470 oC (average = 435 oC). Textural evidence, biotite Fe3+ contents, and uraninite geochemistry (e.g., high Th) support for reduced conditions and we think that graphite precipitated from reduced fluids most probably by the following mechanism: CH4 = C + H2O i.e., cooling of a CH4-bearing reduced fluid. However, we are not completely ignoring the mixing of a CH4 and a CO2-rich fluids as a graphite precipitating mechanism. We are wondering if this association formed below the Fayalite-magnetite-quartz buffer. How can we quantify the redox conditions? What are requirements for such oxygen fugacity quantification?
Thank you in advance,
Sincerely,
Abu Saeed Baidya
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Dear Mr. Baidya,
I have been studying hydrothermal uranium mineral assemblages located in fault- and shearzones at different temperatures and metamorphic conditions and associated with a wide variety of sulfides and arsenides from molybdenite to arsenopyrite and pyrite different in composition with regard to its trace element content.
First and foremost it is important to determine precisely the C-bearing compound which could be in this case kata-impsonite (metamorphic bitumen), semi-graphite and to a lesser extent graphite by crystallographic definition. You must use spectrometric methods in combination with coal petrography to determine the true nature of the C-bearing matter which runs the gamut from epi-impsonite and meta-anthracite to “full-blown” graphite.
For your information I attach Table 1 from
DILL, H.G., KUS J., GOLDMANN S., SUÁREZ-RUIZ I., NEUMANN T., and KAUFHOLD S. (2019) The physical-chemical regime of a sulfide-bearing semi-graphite mineral assemblage in metabasic rocks (SE Germany) – A multidisciplinary study of the missing link between impsonite and graphite. International Journal of Coal Geology 214:
available on the RG server
H.G.D.
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How to prepare Acetylferrocenium salt from Acetylferrocene. A detailed procedure, please.
I have tried with AgBF4 but not successful.
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Dear Narottam, thank you for ypur response and explanation. To my knowledge the ferrocenium salts are not really sensitive to dry oxygen, because they are prepared by oxidation reactions and already contain trivalent iron. The main problem with these compounds is their sensitivity towards moisture. The observation that you first obtained a greenish-blue solid shows that the reaction worked, but then the product decomposed due to the presence of moisture (hydrolysis).So what you should do is make sure that both the diethyl ether and the dichloromethane used in the reaction are rigorously dried prior to use. Diethyl ether should be dried over sodium / benzophenone and dichloromethane over calcium hydride. In any case it's worth trying the reaction again with carefully dried solvents under inert atmosphere.
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I have (at least) one organic compound in my aqueous samples which has a marked peak at ~0.1 ppm (see attached spectrum). I need to find out what it is. It is not TMS, since I use another internal standard at ~7.5 ppm. It's not a contamination from silicon grease (as some suggested in my previous question) since extensive negative controls performed on my experimental setup show absence of this peak. Other blanks show no other reagents I use have this peak.
CH4 is one of the potential products of my CO2-reduction (with H2) experiments. I did a spike test with this sample, and when I dissolve commercial methane into it the peak at 0.08 ppm does indeed become larger; supporting the idea that it's methane. I did GC-FID to see if I saw methane (see attached chromatogram). We don't have a suitable column for GC-MS unfortunately. The FID results show that whatever the peak at 1.35 min is, it isn't methane which has a peak at 1.52 min. The peak at 1.35 is absent in room air blanks, and I presume it's the same organic I see in the 1H-NMR results.
Another possible product of my experiments are (Ni or Fe) bound methyl groups: e.g. Fe-CH3, Fe-C3H9, etc. Which chemical shift should I expect from methyl protons attached to a metal atom? I suspect it would be similar to TMS, since that's exactly what TMS is, right? See the attached example of a Pt organometal chemical shift showing at 0.6 ppm.
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Dear Eloi, no guarantee, but I assume that the chemical shifts will not change to a great extent. However, I have no idea if the NMR spectra of the complexes can be measured in water / D2O mixtures. Surprisingly there are rare publications describing water-stable and (partially) water soluble nickel(II)-sigma-methyl complexes. For example, please have a look at the paper entitled "Nickel(II)-Methyl Complexes with Water-Soluble Ligands L [(salicylaldiminato-K2N,O)NiMe(L)] and Their Catalytic Properties in Disperse Aqueous Systems", S. Mecking et al., Organometallics 2007, 26, 1311-1316. In this paper the authors report the synthesis of Ni-CH3 complexes stabilized by bulky, water-soluble ligands. The CH3 resonances in the 1H NMR spectra were found in the range of –1.00 to –1.50 ppm. The complexes are reported to be stable towards water. Three of the complexes were either insoluble or slightly soluble in water. One of the compounds was even found to be water-soluble. The 1H NMR spectra were measured in C6D6, CD3OD, and acetone-d6.
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hello Scientific community,
How can we distinguish oxydant from reducer gases ? what's the case for Ammonium NH3 ?
Best regards,
A. CHETOUI
----------------------------------------------------------------------
State engineer in semiconductors 
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Usually metallic fluorosilicates are reacted with Ca/Na/K etc metal or alloys to obtain fluoride salts and silicon. But These metals are not very easily produced other than electrolysis. Is it possible to reduce H2SiF6 chemically or with aid of electrolysis rather than their metallic salt? Can cheap reductants like Hydrogen/Carbon/Hydrocarbon do the job? (I do not know redox potential of hexafluorosilicate ion).
By the way, the reactant is from fertilizer industry, and Si is to be used in electronics.
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You can decompose H2SiF6 by heating into HF and SiF4 you can separate them by fractional distillation.
Then you can extract silicon by reducing SiF4 at high temperature into suitable reactors as:
SiF4 +2H2 = Si + 4 HF
Best wishes
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It is hypothesized that the nature/energies/electron distribution of the frontier orbitals of a molecule changes under external field condition. Under applied bias, the molecule can be oxidised/reduced and this changes its electronic distribution and eventually the molecules ability to conduct current. Can we model such an hypothesis using DFT in Turbomole? Some information in this regard is very welcome.
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Almost all quantum chemistry codes support applying external electric field in calculations. My recent work about ultrastrong regulation effect of electric field on various properties of cyclo[18]carbon is a typical example: https://doi.org/10.1002/cphc.202000903.
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I have several question regarding the MOR; hope anyone help me better understand.
1- in CV using (KOH+MEOH) what is the indication of absence of the redox peaks? shouldn't they exist as a sign that the electrode is being regenerated and able to carry the MOR?
2-why some papers don't operate in potential windows to reach the Methanol Oxidation peak and they use the onset potential as an indication to MOR? isn't important to know where the peak is?
3- what is the indication of the area under curve in Cyclic voltammetry (large vs small) for both (KOH alone) and (KOH+Methanol)
any resources are very appreciated.
Thank you.
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Generaly, the absence of peak means that methanol is not electrochemically active under the conditions of your experiment such as potentials window, scanning speed, pH, temperature...
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I am working with a material which is sensible to water and oxygen, I would like to oxidize this powder either chemical or electromagnetically. Does anyone knows a method to electrochemically oxidize the material in its powder form?
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Dear Yanlong Zhou please also see this potentially useful article entitled
"Thermal preparation and characterization of nanodispersed copper-containing powders produced by non-equilibrium electrochemical oxidation of metals"
Unfortunately I don't have access to the full text but you can easily request it directly from the authors.
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Bechamp process involves iron and hydrochloric acid for reducing nitro containing aromatic rings to aniline. In Bechamp process reduction, aromatic nitro compound is adsorbed on Fe0 surface. Here, it reacts with hydrogen ion H+ produced from water and acid reaction, which results into production of amine counterpart of aromatic nitro compound. Does reaction in between involves any free radical mechanism to proceed?
Any literature and suggestions pertaining to same will be duly acknowledged.
Thank you
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The Bechamp reduction is used to reduce aromatic nitro compounds to primary amines by Fe and Conc.HCl.It is an ionic mechanism.First step is protonation of oxygen followed by step electron donation by Fe (zero oxidation state) to Fe2+ . In the First step R-NHOH is formed where as in the second step in acidic medium , it gives amine salt. The neutralization of the salt by alkali gives free amine.
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The ORR polarization plot in the paper (the one with many colors) is very sharp but that of my catalyst is really smooth. I guessed that it is because all the activity sites in my catalyst have high reaction potential (bad thing for a catalyst). Is it right?
And how to improve my catalyst?
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Hello Mr. Huang, In addition to the above suggestions like 1) the film on active geometrical area of RDE is not uniform 2) Roughness is very high, you can think in slightly different line too. Whenever, the mixed and the diffusion controlled regions are not well separated this is another indication of non-uniform distribution of electroactive centers in electrocatalyst it self. In case of commercial Pt-C sometimes the Pt nano particles (d ~ 5nm) may have nonuniform distribution across the carbon support (d~30 nm).
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The energy density of most redox flow batteries is too low for transport solutions but there is a Swiss company: NanoFlowcell that claims their QUANTiNO vehicle accomplished the impossible.
“The low-voltage system in the QUANTiNO works with the nanoFlowcell® drive to form an exceptionally efficient symbiosis that promises a significantly greater potential range compared with the HV systems commonly used in electric vehicles. Furthermore, it eliminates the danger of electric shocks that can be caused by a high-voltage system in the event of an accident or when undertaking repair work to HV components.”
Only data provided by the company are the test run results. Calculations based on this data gives following results:
Energy demand: 13 kWh per 100 km
Energy density: 600 Wh per liter or kilogram
Average test run speed = around 23 kph
A full tank range at this average speed: 1,400 kilometers.
Nowadays, conventional flow batteries based on vanadium sulfate theoretically achieve energy densities of up to 19 to 38 Wh/L with respect to pure electrolyte solutions. Bromine-based systems would achieve 70 Wh/kg.
It sounds like redox flow cells are not an attractive option, as not one of these cells reaches the levels of lithium-ion ones. Maybe I’m missing something? Did anyone had contact with this compay and can can confirm their results?
Maybe somene who works with redox batteries could share his opinion?
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I wish Tesla adopts nanoflowcell technology .
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A Co(II) high spin complex is coordinated with two ligand radicals, which are antiferromagetically coupled with metal d-orbitals resulting one unpaired d-electron on Co(II) center. Now as reported in literature, Isat/(Isat+Imain) ratio for 2p3/2 peak in XPS spectrum is near 0.4 in high spin Co(II) complexes and 0.2 for low spin Co(II). But my complex having high spin Co(II) gives the value of ratio 0.26. Is this due to coupling on two d-electrons with ligand radicals, as now there is only one unpaired d-electron in Co(II) center like low spin Co(II) ?
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Thanks Dr. Ben F. Spencer . It will help me a lot. One thing to add, this is a Co(II) high spin complex. So, I think It may be better not to conclude anything from this ratio. Some of the other high spin complexes, we get the value about 0.33 which are also quite lower.
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For introductory chemistry textbooks, they occur simultaneously. But unless there is a direct transfer between charge species, (say between sodium atom and chlorine gas) especially in solution, the charge transfer between chemical species can be mediated by ions of solution. In this case, does oxidation occur first or reduction first? on which factors this sequence vary? This is important in the case of solution actively participating into solution and recycling itself.
I am well aware that, absolutely deterministic time would be not found in this case (but rather a probability) , since such chemical interactions are bounded by rules of quantum mechanics.
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Actually the two processes occur in the same time. It known that sodium alone will not oxidized and chlorine alone will not reduced. When both elements become near each other, the sodium atom starts to loose the election while in the same time chlorine starts to accept the election.
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CV setup is glassy carbon WE, Pt wire CE, and Ag/AgNO3 RE (no salt bridge). My solvent is DMSO with tetraethylammonium tetrafluoroborate (TEAFB) as the electrolyte. The reference electrode contains 10mM AgNO3 with 100mM TEAFB in DMSO. Solutions are made in a dry anaerobic chamber and the electrochemical cell is continuously purged with argon.
Before performing CV on the redox couple analyte, I condition the electrodes by doing CV with just electrolyte solution, which should generate a flat line (I think). Instead, I am consistently observing very small peaks near 0V.
I was able to measure a reasonable duck plot for ferrocene a week ago, but with fluorene I only see the tiny peaks around 0V, however the signal strength is about 1.5x in that solution (which has 3x the concentration of electrolyte).
How do I troubleshoot this? If it were O2 or some other contaminant, shouldn't the peaks appear somewhere else? Currently I am thinking that something is going on with the electrodes but I don't know what.
Please help, thank you
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Hi;
I totally agree with Chris Gunderson.
With my best regards
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Nafion membranes are used to separate the two electrochemical cells in the CO2 reduction experiments. Can anyone explain what is its role in CO2 reduction experiments?
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As the others said, separation of products is the main purpose, to prevent their loss due to crossover.
A critical requirements is also to support the transport of H+ from the anode to the cathode, where H+ is a reactant in the CO2 reduction reactions. This completes the electro-chemical circuit in a sustainable way. However, this is often overlooked. Many researchers use Nafion which is a cation exchange membrane. But in the pH ~7 solutions usually employed, the effective H+ concentration is ~10-7 M, whereas the electrolyte cation concentration (K+, Na+, etc) is usually far far higher (often 0.1-1 M, depending on the concentration used). Depending on the relative transference numbers, the primary species transported through the membrane is likely the alkali cation, not H+, and therefore the cell will be dialysed under operation, an unsustainable configuration.
An anion exchange membrane could be more suitable, but then organic anions produced by CO2 reduction (formate, acetate) could be lost across the membrane as well.
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I am a phD student and since I started, I was using ferrocene carboxylic acid for CV and impedance scans, but now I started to use Ferri ferrocyanide in order to try the redox couple with a new polymer. The problem that I am having at the moment is that I am not be able to get a proper CV shape (as you can see in the Fig 1) in the bare gold. I tried:
1. Different current and potential
2. Checked the holders and connections
3. Use external counter, external reference and external working electrode to check the reaction in the internal reference and counter on chip.
4. Different scan rates,
5. Remove oxygen in the solution.
6. Use different concentration of ferri ferro and PBS.
7. Clean the surface with acid
8. Square wave voltammetry.
After try all this options is happening the same all time but just in the internal working electrode (if I use external one the scans are perfect. Another people in my lab were trying ferri ferrocyanide in the past and looks like they were having the same problem as me and them decided to use another type of ferrocene. But at the moment I think that if with a commercial electrode the solution works (fig 3), it should be the same with the internal electrode on the chips that we use.
When I did Square wave voltammetry, I Could realise that there are appear different peaks when I use the same electrode in different time (fig 4). And each one of the electrodes has different peaks. But what is weird is that each time its different on the same electrode, which means that there is a type of reaction happening on the surface.
Can anyone help me to understand what is happening here? I think that in the surface of the electrode there is ‘’something’’ that can be detected using ferri ferro but not with ferrocene carboxylic acid (because the reaction is mask for the groups OH during the oxidation of carboxylic groups).
I use a concentration of 1.5 mM of ferri ferrocyanide in PBS, but as I wrote before I tried different concentrations (of ferri ferro and PBS) and nothing works with the internal electrode but it works with the external working electrode (so the solution it’s not the problem in my opinion). Also I was reading about electro oxidation of methanol or ethanol or some surface with wax layer and the behaviour is a little bit similar, but I am not an expert on this…
Thanks in advance for your time and I hope get help on this.
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Thanks a million for the answer @dan-marianjoita . I will try more things, but I will check which types of membranes have the electrodes that come from the fab.
Regards
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I am looking for a nucleobase (utmost its derivative) which shows reversible redox activity. I went through some of the literature and found out that the redox features of nucleobases are very poorly reversible, in general.
As I am novice in this line of research, any relevant reading suggestions are also much appreciated. Thanks.
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Abel J S C Vieira Thanks for the Clue.
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For the diffusion-controlled redox reactions in a battery or capacitor system, except for the ion intercalation, are there any other reactions can be account as diffusion controlled redox reactions? I very appreciate that if you can provide the related references. Thanks a lot!
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This online link can help you to understand what you are looking for...https://en.wikipedia.org/wiki/Diffusion-controlled_reaction
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The standard reduction potential of Fe3+ is 0.77V vs SHE, and the standard reduction potential of [Fe(CN)6]3- is 0.37V vs SHE.
Prussian Blue, or Iron Hexacyanoferrate, is a coordination polymer made of interconnecting Fe cations and [Fe(CN)6] anions. Cations such as K+ or Na+ may fill the interstitial sites, along with zeolitic water. Additionally, due to the strong ligand field (CN)-, the C-coordinated Fe of [Fe(CN)6] groups are in the low-spin state, while N-coordinated Fe remains in the high-spin state.
It has been proven that the redox potential of low-spin [Fe(CN)6] is higher than that of high-spin Fe in Prussian blue.
If the standard reduction potential of high-spin Fe3+ in water is higher than low-spin [Fe(CN)6]3- in water, why is it the opposite in Prussian Blue? My explanation has always been that the oxidized electron in low-spin [Fe(CN)6] is at lower energy than high-spin Fe, but that also assumes the barycenter is the same. However, this explanation doesn't work for the redox potential of aqueous ions.
Edit: For clarification, this is an example of the intercalation and reduction of both Fe. Redox potential is for aqueous Na+ battery from this source:
  1. FeIII[FeIII(CN)6] + Na+ + e- <=> NaFeIII[FeII(CN)6] 1.2V vs SCE
  2. NaFeIII[FeII(CN)6] + Na+ + e- <=> Na2FeII[FeII(CN)6] 0.2V vs SCE
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Adrian Fortuin , thank you for your reply. I am not sure if catalytic surface is the right word to describe Prussian blue in this paper, but I understand that you are saying it is different than ions in solution.
Yes, Prussian blue is mixed valence, but in its fully oxidized state (Berlin Green) there is only Fe3+, FeIII[FeIII(CN)6]. The low-spin [Fe(CN)6] is always reduced first. I have added some clarification to the discussion description.
I agree with your response about thermodynamics, the chemical potential of the system must change which changes the redox potential. I think Volodymyr Tkach is correct in saying the stable complex formed will affect the redox potential, but specifically why or how does it change. The compound may be more stable than its ions in solution, which lowers its chemical energy (higher voltage) for low-spin [Fe(CN)6], but actually high-spin Fe has a lower redox potential than when it is an ion in solution.
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I am currently trying to use the xanthine-xanthine oxidase method of the determination of the redox potential of flavoproteins as first described by Massey in Flavins and Flavoproteins 1990. I am having trouble getting the experiment to work and was wondering if anyone had any tips or recommendations for me to get this thing to function properly. I'll briefly explain how the experiment works in the next paragraph for anyone who is not familiar with it but feels like they might be able to offer some insight. For anyone who has done this experiment, you can just skip the next paragraph.
The experiment works by using the oxidation of xanthine to uric acid by xanthine oxidase under anaerobic conditions. Methyl vionlogen is added to act as the electron mediator and will be reduced by xanthine oxidase when the xanthine is oxidized (when oxygen is not present). You add in your flavoprotein along with a reference redox dye with a known redox potential. The now reduced methyl viologen will then be oxidized and reduce both the flavin and the reference dye. The absorbance wavelengths are recorded for 1-2 hours as this reaction takes place, and through some math magic (not really important here since I haven't made it to this step) you can determine the redox potential of the flavin (since the redox potential of the dye is known). To achieve anaerobic conditions, many papers use argon sparging to remove oxygen from the sample in the cuvette and some papers include the use of glucose and glucose oxidase in the mix to consume any left over oxygen.
I have tried many variations of this experiment with a range of concentrations for each component, but here is what one of the typical experiments looks like: 100 mM HEPES, pH 7.5, 200 mM KCl, 400 µM xanthine, 20-30 µM methyl viologen, 20 µM flavin, 20 µM reference dye, 450 µg/mL Xanthine oxidase (Calzyme - 1.3 U/mg), and then I am also adding glucose at 5-10 mM and glucose oxidase at 100 µg/mL to achieve anaerobic conditions. Everything except for the enzymes are added to the cuvette. We do not have the money to spend on a fancy sealed cuvette, but I found a bunch of these UV-Vis disposable cuvettes (BRAND) that have a circular opening that a size 11 septum fits perfectly into. The septum is put on and the sample is degassed with argon for 30-60 minutes. Enzyme is then added to the sample using a gastight Hamilton syringe. The cuvette is gently mixed and the septum is parafilmed and placed in the UV-Vis for measurements.
The problem I am running into is that I am getting no reduction of the flavin or reference dye. Since I am just trying to get this to work, I have just been using free FAD and either safranin or phenosafranin as the reference dye. I have found that the degassing technique and the cuvette/septum does work for removal of oxygen. I did an experiment with just xanthine and xanthine oxidase and I monitored the formation of uric acid at 290 nm (xanthine is at 271 nm). Normal conditions (non-degassed) showed enzyme activity. I then did the same with a degassed sample and saw huge (probably 90%) reduction in enzyme activity. I then tried with the addition of glucose and glucose oxidase to remove any remaining oxygen and saw 0% enzyme activity. We just bought new methyl viologen and that still is not working. I have been using fresh xanthine and xanthine oxidase each day that I do tests, but am having no luck. Any help would be greatly appreciated as I have run this experiment over 30 times with no results and I have run out of ideas as to what could possibly be going wrong.
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Just wanted to let anyone who was wondering know that I finally got this reaction to work. Long story short, I ran an activity assay on the xanthine oxidase that we had and it was only 0.5% of the reported activity. Immediately bought new enzyme, from Sigma this time, and the reaction worked perfectly. So, to anyone who is having trouble with this assay, check your enzyme activity, and just save yourself some time and sanity and order it from Sigma.
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Dear Researchers,
Based on my current knowledge, I am capable of calculating peak current values when it is a well-shaped CV curve like a duck. (I am working with EC-Lab software- Biologic)
Unfortunately, when it comes to such graphs that does not include any possible part on them for me to be able to locate any regression line, it gets really hard to make any assumption, especially for the oxidation peak as you could see.
As far as I am familiar with the concept, there is no information about how to detect these values on such shaped graphs with no possibility allowing you to make a straight line through the oxidation peak to calculate anodic peak current value.
Could someone please explain based on their experience/knowledge?
I am even not sure whether or not I would be able to find out the values from these graphs.
I would greatly appreciate any help!
Best,
Ali
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Hi Ali,
How to choose the values for the scan rate depends on the interval. It has to be representative in an interval of values. 3, 5, 8, 12, 15 are not common scan rate values. If you have to choose one, usually is 100 mV/s or 50 mV/s. Then possible values are 10, 25, 50, 75, 100, 150, 200 or 10, 20, 40, 60, 80, 100, 120...
Regarding the measurement, if you go one by one you will see a change in the slope, you can try to draw the tangent by this point and measure the peak current. Even if the peak is "sluggish" you have to measure. Do always the same,
Best,
Teresa
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I would like a simple method to measure in situ soil reduction potential. Does anyone know of a durable sensor/probe that can be inserted into the soil profile for a reading of ORP? I'm reluctant to make a slurry or paste and measure with an Eh/ORP sensor as this will introduce atmospheric oxygen. 
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I have a YSI (ORP 15A) which costs quite a bit more than the ExStik. Dr Rezai, do you know the type of Electrode the ExStik uses ?
I have to ask YSI about the type of electrode used in the ORP 15A. Thanks
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For example, Fe3+ in octahedral sites of ZnFe2O4 normal spinel have unpaired electrons aligned anti-parallel to each other by the superexchange force. In contrast, Fe3+ in octahedral sites of NiFe2O4 inverse spinel are parallel to each other, and anti-parallel to Fe3+ in tetrahedral sites.
1. Does the energy level of unpaired electrons differ between magnetically coupled and uncoupled states? Assume the same ion and same coordination.
Redox potential is related to the energy level of the electron being taken or given to the metal ion orbital.
2. Could we expect the redox potential of metal ions to differ due to a difference in magnetic coupling? I understand that redox potential can also be affected by other changes to local environment such as changing lattice parameters, induction effect, etc, so I am just asking if magnetic coupling could contribute to different redox potential, not solely responsible for it.
I appreciate your responses and any related literature recommended.
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In most ferrites we can predict that they have normal or reverse spinel structure but the alignment of the dipoles in these oxides is an intrinsic property and can not pprdicted.
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How can i calculate quantity of the reducing agent(Copper powder) in to silver nitrate to complete the recovery of silver ,and what parameter should i look while reaction that will indicate the completion of process of redox or cementation of silver
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If you are using copper to recover Ag, how are you going to separate the excess amount of Cu from silver?
You better add zinc instead of copper, that you can easily remove the unreacted zinc by adding HCl.
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I'm working on a five organic dyads, consisting of an electron donor and electron accepter unit. I want to calculate the charge seperation and charge recombination states energies of these compounds. I got the values of three compounds with Weller's equation using redox experiments. However due to poor solubility of two of the dyads, I failed to get information form cyclovoltametry experiments. I tried to get these values by normalized absorbance and emission of CT band. But its not accurate. Is their any other way to find these energies with maximum accuracy.
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T. H. Al-Noor, Thank you very much for your literature sharing. Actually I have already measured two compounds with the method described in this literature. However the compounds with poor solubility are the reason I'm looking for another method to find these values.
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Could anyone please tell me the names of some common irreversible redox systems/molecules?
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Richard A Durst here on RG seems to be the one to suggest
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Hi everyone,
I computed the HOMO-LUMO of a compound to be -5.79 and -3.64 eV, using DFT.
But I've seen works where these values are reported as mV vs the normal hydrogen electrode or ferrocene/ferricinium.
How do I convert/reference my values to this?
Thanks in advance
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Have you took a sight at the Farrell's work?. He works on electro-oxidation and made some DFT simulations and related them to electrode potentials.
Understanding anodic wear at boron doped diamond film electrodes
Chaplin, B. P., Hubler, D. K. & Farrell, J. Feb 1 2013 In : Electrochimica Acta. 89, p. 122-131 10 p.
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Which acetonitrile is the best to use for CV? In the one I'm using from SPS I can see redox waves from impurities in the blank.
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Oxygen is a big problem in acetonitrile. The solubility is 8.1 mM/ atm, around eight times higher than in water. Worse, the one electron reduction to superoxide dominates (and this is chemically reversible, so you get both forward and backward waves), and even worse, traces of water will protonate the superoxide leading to a range of redox active pH-dependent products. Outgassing under vacuum (with the MeCN over molecular sieve) is a good start, then bubbling with high purity argon (best) or oxygen-free nitrogen (which needs scrubbing, since it's never really oxygen free). The pipework needs to be top-notch, pipe runs as short as possible, and you need to pre-moisten the gas before sparging by blowing it through a Dreschel bottle of backgroud electrolyte, or your cell electrolyte will evaporate and also cool down (further worsening the O2 solubility issue). all of this means working in a fume cupboard, or you'll fill the lab with lachrymatory MeCN vapour and your co-workers will be weeping more than usual. Your electrolyte salts also need to be dried before use. The quats in particular pick up a lot of water even under careful storage.
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For example, in a 3-electrode test of a non-aqueous Na-ion battery. The working electrode is being reduced or oxidized during Na-ion intercalation/deintercalation, therefore oxidation or reduction must be taking place at the platinum counter electrode in order to maintain charge balance in the electrolyte.
If my electrolyte is 1M NaClO4 in propylene carbonate, what reaction is happening at the counter electrode?
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Have you got some electrochemical results (e.g. CV and GCD) of this three electrode system? During the test, have you observed some gas generation on the counter electrode? In my opinion, the electrolyte ions would accumulate on the surface of the counter electrode first to balance the charge. Then there might be decomposition of the solvent (PC).
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Hello all,
I have been trying to modify a plain printed carbon electrode with cobalt(II) phthalocyanine and I have found my CV does not look like any of the CVs in literature. I have attached a copy of my CV to demonstrate.
Protocol:
Suspended 10mM CoPC in 1:1 H20:MeOH, probe sonicate for 10 minutes. Solvent cast 5uL onto the WE and leave to dry at room temp.Wash with MeOH and run in pH7.2 PBS (0.1M).
Most literature shows redox peaks for CoPC at ~-0.3 - -0.1V, however my electrode using a Pt count, AG/AgCl ref shows 1 giant reduction peak at -0.22V but no oxidation peak occurs.
Any suggestions?
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I agree with the previous answers and suggest you to add cobalt (III) oxyhydroxide to phthalocyanine. It would enhance the catalytic efficiency of the catalyst.
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NAD is involved in most cellular redox reactions as an oxidant (NAD+) or reductant (NADH) and it changes between both states multiple times in every fraction of second in every cell. But it is still not clear to me the following:
  1. Why is it called dinucleotide when I see only one nucleotide as its component (adenine)? or is it nicotinamide a nucleotide too?
  2. Why it is not so used something like Nicotinamide Guanine Dinucleotide (NGD) instead? or with Cytosine or Thymine? what makes Adenine appropriate for this function?
NAD recently recalled my attention since I read a very interesting article in which the authors managed to decrease age-associated diseases in mice by administration of NAD intermediates:
Mills, Kathryn F., et al. "Long-term administration of nicotinamide mononucleotide mitigates age-associated physiological decline in mice." Cell metabolism 24.6 (2016): 795-806.
Thanks for your time.
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Hi,
Living systems is a result of thousands of consecutive biological reactions, catalyzed by specific enzymes, in orchasetrated manners. NAD/NADH & NADP/NADPH are very important biological cofactors, participating mainly in numerous cellular oxidation/reduction reactions, as an electron acceptor(to be reduced) or donner(to be oxidized). There is another biological cofactor known as FAD(Flavine Adenine Dinucleotide) also participates in biological reactions. And the particular enzyme involved in each reaction, has strong choice (specificity) for its cofactor of either NAD or NADP or FAD, to catalyze the reaction optimally. This specific choice for cofactors of the catalyzing enzymes, makes other analogues unsuitable for biological reaction. The specificity depends on many factors as three dimensional conformity of the subustrates, enzymes, cocfactors and redox(reduction/oxidation) potential of the reaction & reactants etc.
Although there may be some artificial or synthetic analogues that may somehow support a biological reaction in lab conditions, the reation rate may not be optimum as in biological system. So, if cellular reactions do not run optimally, there must be some traffic congestion, thus making the whole system go-slow or idle (ageing?). All These make NAD/NADP so important. Good luck for you.
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In a photo-catalyzed system based on Ru(bpy)3 complex, is it possible to prevent electron-transfer back-reactions and returning of the generated radical to the ground state due to presence of a radical scavenger like DTT?
UPDATE:
These info are also necessary for the question. The intended system is used for generation of radicals from  and crosslinking of a protein (Tyrosine). In this way, formation of [Ru(bpy)3]3+ from the excited [Ru(bpy)3]2+* is necessary. But as I mentioned in the original question, is it possible to drive the system down the oxidative quenching pathway of [Ru(bpy)3]2+ by an oxidative quencher like ammonium/sodium/potassium persulfate, and at the same time, have a strong reducing agent like DTT present in the system? The presence of DTT is due to the need for activation of a Michael addition reaction after photo-activation of [Ru(bpy)3]2+.
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Yes, you have problems with terminology. I don't understand well your question. The chemistry of Ru2/Ru3 is the electron transfer ET) processes. The concept "Nucleophilic" is  not used in this area. A donor and acceptor of electrons are the right terms. The reaction can't be "stabilized." The back reaction can be replaced by another faster reaction using a sacrificial electron donor or acceptor. Ru2* can be reduced by an electron donor to form Ru1 and "oxidized" electron donor, which can recombine back.  Commonly, you can direct the reaction in the desired way. 
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Hello all,
I found a research paper detailing the oxidation potentials of some 50 quinones obtained via CV, and I'd like to replicate this technique using benzyl alcohol and similar compounds (ex. 4-Br-BzOH, 4-methyl-BzOH, etc.).  I am doing research with an organic Te dye capable of oxidizing small molecules (like thiols). I have determined the oxidation potential of this dye and wish to compare it to the respective oxidation potentials of these alcoholic substrates so that I can predict which experiments will yield a successfully oxidized product (aldehyde).
I want to find the oxidation potentials of these substrates using CV preferably, but after looking through literature I fear I may have to resort to chemical oxidation and measure a delta-G value somehow (and convert to V). I have found that benzyl alcohol oxidizes at >2.00 V (this is the only lit value I could find), but since my Pt working electrode itself becomes oxidized around 1.8V, I don't think it is possible to find the oxidation potential this way. Oxidation would occur not electrochemically but via chemical oxidation by PtO2 or Pt(OH)2 or some other species.
How can I determine the oxidation potential of aromatic alcohols?
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Oxidation of benzylic (and) other) alcohols commonly proceeds via C-H bond cleavage. Therefore, the C-H bond dissociation energy is more important than the reduction potential, e.g. read 
Mayer et al, Chem. Rev. 2010, 110, 6961
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 The cyclic voltammetry study of the electrochemical oxidation of a particular analyte showed only a single oxidation peaks in the forward direction. No cathodic peaks were observed in the reverse direction. I believe that the absence of any reduction peaks on the reverse scan confirms that the charge transfer is electrochemically irreversible at the electrode surface. Is it possible to study the effect of pH vs Epa using Nernstian expression for such systems?
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Hi Birhanu,
The Nernst equation is only valid in equilibrium. If the reaction is slow, then accounting for the kinetics may be important (through the Butler-Volmer equation for example).
In addition, adsorption effects may contribute to the absence of the peak in the reverse direction. We observed similar results in our experiments and successfully captured the physics with a model taking into account diffusion, kinetics, adsorption effects and morphology changes on the electrode surface.
Here is the link to our paper, I hope it helps:
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I am doing CV for ferrocene and the deuterated analogue of ferrocene. What  might be expected from the CV with regards to shifts in the oxidation and reduction potentials?
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In the case of ferrocene, electron transfer is very fast thus on CV you could only see thermodynamic effect. Moreover even if you could see the kinetics, it is a simple outersphere electron transfer without any bond breaking or formation.
From a thermodynamical point of view, I don't see why an isotopic effect would change the thermodynamics of the system.
Thus I would expect no change between both molecules, if you have one that could be very interesting.
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Since i'm not familiar with cyclic voltammetry i'm wondering how to do an experimental set up to analyse laccases for the dependence on distinct concentrations of ethanol ...
How to arrange the cell (electrodes) and which solvents and electrolytes are appropriate ?
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Dear researcher
i am working on electrochemical detection of DNA and proteins after adsorption of Fc-DNA probe or Fc-aptamers on reduced graphene oxide modified electrodes. But, the sqaure wave voltammetry signal is not stable with time. always there is 5 to 10% decrease in the swv signal which cannot be used for further measurement. does anyone saw this issue before and how can we figure out it. many publications have been published in this field bit none talk about the stability of the adsorbed DNA containing redox tag.
regards
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Dear,
Although my main experience is with gold-thiol chemistry, the problem seems very similar. After modification of the electrode surface there will always be loss due to non specific adsorption. This may be due to weak interactions or crowding of the electrode surface. Depending on your modification method, I would consider a loss of 5-10% reasonable. Normally I used to include "desorption" step in the modification method. This is essentially an additional incubation which mimics the conditions of the test (but without the analyte).  Normally losses should become smaller after this. To distinguish between stability of Fc or the desorption of DNA probe, you could  use a unmodified DNA probe  and use  for example Co(phen) to characterize the electrode. If the loss is similar under these conditions then the problem originates in the probe adsorption not the ferrocene labeling or Fc stability. Hope this helps.
Kind regards,                          Mike
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I was trying to use an alternative way to measure the I-V of RGO on copper immersed in NaCl. I got three peaks. (-50 mV, 175mV, and 275 mV) roughly. What do they account for ?
are they because of the reaction with the functional groups attached to RGO ?
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Thanks Uddin for the link.
Zhu thank you for your answer, I am looking forward to knowing how far you have been with that.
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Please send equation to convert.
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I would friendly advise you to learn what is the reduction potential, oxidation potential, and what are the different Ag/Ag+ reference electrodes, and then google. I sincerely encourage you to learn how to learn and be minimally professional in scientific forum. I would friendly recommend to remove your question or at least edit it. My comments are not intended to insult you.  
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We are interested in a brain mimetic solution or material to record voltammograms using microelectrodes.
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Hi Alex
We would routinely use brain tissue homogenate, however, if you don't have access to this then I would recommend looking at individual constituents of brain tissue such as proteins and lipids. These poisons effect electrode characteristics. We use Bovine Serum Albumin (BSA) and Phosphatidylethanolamine (PEA) to mimic the effect of proteins and lipids respectively. TritonX is a surfactant that we would use occasionally too. 
However, if you are just wanting to investigate the voltammogram in solution that mimics the ionic composition of the brain then I would suggest artificial cerebrospinal fluid (aCSF). There are a number of recipes for this that you can avail of.
hope this helps.
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I have seemingly random samples in which the tin surface appears to have an oxidation layer after being allowed to sit with a hydrogel on it. These samples also show a voltage when connected to the multimeter that is much higher than expected. I have done some experiments in which I introduced different amounts of oxygen into the gel solution prior to UV curing. They appear to show poorer cross-linking at higher O2 content, as expected, and a higher voltage read out. I hypothesize that a poorly formed gel may result in high voltages (on the order of 100 mV). The samples that perform poorly seem to 'ooze' over time. Removal of the gel from the tin yields a whitish haze that cannot be wiped off (for the bad samples) and the shine expected for elemental tin on samples with good voltage readings.
The hydrogel primarily consists of polyacrylate/polyacrylamide, glycerol, water, polyvinylpyrrolidone, KCl, and a few other species.
Can anyone offer insights into what may be happening at the interface between the hydrogel and the tin surface? Obviously, there appears to be a redox reaction between tin and something else, but I don't have much experience with hydrogels.
Best,
Jack
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 Titus, Salts should enhance this, but I'm left with the question why do some samples show oxidation while others don't? The ones that do, appear to be leaching shorter chain polyacrylate. I imagined that a localized high pH generated at the surface is causing enhanced oxidation. My question is more along the lines of if anyone has experience with observing a potential one the order of 120 mv and if this was concluded to be due to a base mediated Sn O2 reaction or something else. 
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My professor and I were using sulfuric acid and perchloric acid with 0.2M concentration to study the quartz crystal by EQCM technique, but we could not see any changes in mass when we increased or decreased the concentration. The mass changes data did not show any relations to the concentration changes. We were thinking about the other acids such as perbromic acid or periodic acid.
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My professor and I follow this experiment design with carbon electrode and quartz crystal gold working electrode in sulfuric soltuion. The nitrogen gas was purging inside the solution before the electrochemical process. Our data showed that there was 16ng difference on the quartz crystal, but the studies could not go further when we change the concentration of sulfuric soltuion below or above 0.2M. The mass change was still around 16ng, and we assumed from the reference that the oxygen layer was adsorped on the quartz crystal surface.
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While using RRDE in oxygen reduction reaction. We have to calculate %age of H2O2, What will be N value?
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Voltammety of H2O2 on Pt-electrode allows to determine H2O2 concentration in mineral asid's solutions with help of calibration of the limiting diffusion current of H2O2 oxidation (at potential appr. 0,9-1,2 v (RHE). 
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In my graph only oxidation peak exists and reduction peak was not observed. How can I calculate E1/2 from that graph?
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Thank you @François L L Muller.
Your explainations have given me some new ideas.
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In cellular respiration, O2 is used as a final electron receptor an turns into 2H2O. To do that, it should be turned into O+O (Not sure). So what is the enzyme or the biophysical process that do this? All steps of cellular respiration are known, but I was told that the answer of this question is unknown, but I am not sure about that.
Tip (I think it may help): NADPH oxidase, an enzyme that can be found in the plasma membrane as well as in the membranes of phagosomes used by neutrophil white blood cells to engulf microorganisms. It generates superoxide by transferring electrons from NADPH inside the cell across the membrane and coupling these to molecular oxygen to produce superoxide anion, a reactive free-radical.
Thanks a lot in advance.
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Look at the following review for cytochrome c oxidase, the site of O2 reduction.
Reaction Mechanism of Cytochrome c Oxidase
Shinya Yoshikawa and Atsuhiro Shimada
Chem. Rev., 2015, 115 (4), pp 1936–1989 
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Ellman's (DTNB) had been reduced to TNB- on the surface of thiol functionalised graphene, and resulted in two adjacent peaks at around -0.5 to -0.8 V. I was expecting only one peak, and cant think of a reason for the second one. 
Any help would be greatly appreciated! 
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the buffer was purged with Ar
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Can anyone tell what is the oxidation potential of arsenic (III)? I need a different approach to get the value...not the Eh-pH diagram.
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Thanks for your suggestion Dr Geletti.
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Antioxidative property of extract from any biological material will be elicited through either single electron transfer or hydrogen atom transfer from the reducing agent to oxidizing agent. What is the chemistry behind the antioxidative property of Maillard reaction products?
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Thank you all for sharing your helping hand. Dr.Abdallah, thank you for the offer of your informative dissertation paper.
Kindly,  
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In a capacitor or battery system, there is reversible redox reaction at the electrode/electrolyte interface during cycling, the redox species come from electrolyte and the electrode does not get involved into the redox reactions. So, during cycling, the electrode itself just provides the active sites for the occurring of redox reactions which the electrolyte species are involved. Does anyone familiar with this case, I very appreciate that if you can provide me the related references. Thanks so much.
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Actually, the electrode surface gets involved in the redox reaction. for the redox reaction to occur in electrochemical energy storage, the electrode should either positively or negatively polarised. That's why transition metal complexes (oxide, sulphide etc) or conducting polymer or hetero-atom doped carbon can take part in redox reaction, which is called pseudocapacitance. 
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Compared with the acidic electrolyte where H+ is converted to H, it is likely that the conversion of H2O into H is more sluggish, so the activity of HER in acid is usually higher than that in alkaline. But how about the comparison within basic electrolyte,i.e., increasing the concentration of OH- in electrolyte usually leads to a higher HER activity? Why? Is there any qualitative explanation besides that PH can influence the Hydrogen binding energy? 
Thx
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Water reduction is a half reaction of water splitting to H2 and O2. The reaction thermodynamics of water reduction is described by a simple formula:
E = 0.059*pH (V). 
The reaction rate is not controlled exclusively by thermodynamics but also by activation barrier (kinetics) The activation barrier can be lowed by the catalyst. The catalyst  activity depends on pH. 
In addition, water oxidation might be a rate limiting step for the overall water splitting. Water oxidation is more favorable under basic conditions. 
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An organic moeity reduced at Ep = -0.497V, Eh = -0.405V with Ip = 5.187uA in CV. I need Hpw for this voltammogram. How to find this? 
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see enclosed file
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if anybody have reference papers, please suggest me 
I would like to know my Molecules are transferring whether one electron or more and their reversibility property.
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It might help. 
Hyperelectronic Metal−Carborane Analogues of Cymantrene (MnCp(CO)3) Anions: Electronic and Structural Noninnocence of the Tricarbadecaboranyl Ligand
Article in Organometallics 26(18) · August 2007
DOI: 10.1021/om700496v
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Thank you so much for your help
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Thank You very much Mr. Saunak Das
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Hallo:
I want to see the exchange current in charge transfer reactions at TiO2 films -0.6M NaCl solution interface.I want to know about the redox potential of my solution as its very important to find relation.
How can I find the redox potential.
Thanks
Best Regards
Qaiser Ali Khan
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Dear Professor Cornel radu:
Thanks once again, I have sent you detailed answer as message.
Thanks
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Hi,
I have a nickel oxyhydroxide and it is found to oxidize ethanol when treated electrochemically. But I am unaware about the mechanism and intermediates produced during the process. I want to know what changes are taking place in my catalyst. Several literature reported the formation of Intermediate 1 and Intermediate 2. But I have no clue about the intermediates. Please help.
Thanks,
Saurav
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The details of how ethanol is oxidized by Ni(III)O(OH) is unknown. There are several Ni(III) oxides and hydroxides. Their stability depends on pH and temperature. There are at least two different forms of black Ni(O)(OH). A mixed Ni(II)Ni(III) hydroxide of stoichiometry Ni3O2(OH)4 forms on aging of NiO(OH). It is assumed that the so-called nickel-peroxide (nobody knows what that is) works through a free radical mechanism. See Chem. Rev., 1975, 75 (4), pp 491–519 and references cited therein.
By the way your title "...metal oxyhydroxide (MOOH)..." gives the impression that you believe that other metals beside nickel  behave the same. This is certainly not so.
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It is said that 0.5M H2SO4 acts as a stop solution to stop the reaction and turns the green color to yellow.I would like to know how it halts the reaction and how TMB and HRP play their role in the reaction in a bioassay.
Thank you,
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H2SO4 makes the solution acidic and thus inactivates horseradish peroxidase, stopping the enzymatic reaction. It also stabilizes the colorful product of the reaction.
TMB is the hydrogen donor for the reduction of H2O2 to water, in the reaction, TMB is converted to its blue counterpart (see link).
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I tried to use H2O2 to make the orp go up in my solution. Unfortunately the orp went down. After that i tried to put hydrogen peroxyde in water to see what will happen and the orp went down again.
Is not peroxyde an oxidizing agent?
Thank you
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It's dependable to method conditions such as, PH , temp. , etc., so the hydrogen peroxide was reported to redox reaction that mean the reaction may be down or upgrading...
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The concentration of equine cytochrome c is calculated using the extinction coefficient at a given wavelength.
Reduced cytochrome c at 550 nm using e29.5 x 103 xM-1 cm-1
Reduced minus oxidised using e21.1 x 103 xM-1 cm-1
Oxidised cyto c at 550 nm using e8.4 x 103 xM-1 cm-1
Complex 4 activity was analysed using cytochrome c (reduced using ascorbate and purified via Sephadex column) in phosphate buffer (10 mm pH 7.4) and mitochondria at 50ug/ml. The spectrum of reduced cytochrome c was read from 400-600 nm before the addition of mitochondria (to confirm it was reduced) and repeated following the assay to confirm that oxidation had taken place. Absorbance at 550 nm was usually ~ 0.7
30ul of Reduced cytochrome c had an absorbance of 0.702 at 550 nm = 0.0332M (before adding mitochondria) and an absorbance of 0.135 at 550 nm =0.0063M (following the assay). By calculating the slope of the reaction and the extinction of reduced minus oxidised cytochrome c I can estimate how much oxidation of reduced cytochrome c occurred.
My question is this: Why would I need to calculate the isosbestic point of cytochrome c? My reduced cytochrome c was made in batches and stored at -80 for ~2 months at a time. I reduced the same cytochrome c for all assays and calculated from the absorbance at 550 nm what volume I required to get 100um in a 600ul cuvette (30ul was the usual vol. used).
My understanding of the Beer-Lambert law where A=e l c says that I know the concentration of both oxidised and reduced cytochrome c. However I am advised that this is not total cytochrome c.
Are my missing some unknown extinction coefficient for total (as opposed to oxidised and reduced)?
I have attempted to reduce cytochrome c by starting with oxidised and adding ascorbate at different concentrations however my isosbestic points are not in agreement with the literature possible because my reduced cytochrome c now contains ascorbate, and is thus not a suitable method for determining some unknown in cytochrome c. It is not comparable to the cytochrome c I used for assays as this was purified via Sephadex column.
Despite a trawl through the journals beginning from 1959 I am none the wiser.
Any advice/suggestions /references/job offers are gratefully received.
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Hello guys,
I wonder if there is any assay that could be used to determine the redox state of HMGB1 protein released in culture media ? your answers will be greatly appreciated
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Dear Sir. Concerning your issue about the assays are used to determine the redox state of HMGB1. During inflammation, high-mobility group box 1 in reduced all-thiol form (at-HMGB1) takes charge of chemoattractant activity, whereas only disulfide-HMGB1 (ds-HMGB1) has cytokine activity. Also as pro-angiogenic inducer, the role of HMGB1 in different redox states has never been defined in tumour angiogenesis. To verify which redox states of HMGB1 induces angiogenesis in colorectal carcinoma. To measure the expression of VEGF-A and angiogenic properties of the endothelial cells (ECs), at-HMGB1 or ds-HMGB1 was added to cell medium, further with their special inhibitors (DPH1.1 mAb and 2G7 mAb) and antibodies of corresponding receptors (RAGE Ab and TLR4 Ab). Also, a co-culture system and conditioned medium from tumour cells were applied to mimic tumour microenvironment. HMGB1 triggered VEGF-A secretion mainly through its disulfide form interacting with TLR4, while co-operation of at-HMGB1 and RAGE mediated migratory capacity of ECs. Functional inhibition of HMGB1 and its receptors abrogated HMGB1-induced angiogenic properties of ECs co-cultured with tumour cells. HMGB1 orchestrates the key events of tumour angiogenesis, migration of ECs and their induction to secrete VEGF-A, by adopting distinct redox states. high-mobility group box protein 1 (HMGB1) is a nonhistone chromatin-binding protein involved in the regulation of transcription. Extracellularly, HMGB1 acts as a danger molecule with properties of a proinflammatory cytokine. It can be actively secreted from myeloid cells or passively leak from any type of injured, necrotic cell. Increased serum levels of active HMGB1 are often found in pathogenic inflammatory conditions and correlate with worse prognoses in cancer, sepsis, and autoimmunity. By damaging cells, superoxide and peroxynitrite promote leakage of HMGB1. Recent Advances: The activity of HMGB1 strongly depends on its redox state: Inflammatory-active HMGB1 requires an intramolecular disulfide bond (Cys23 and Cys45) and a reduced Cys106. Oxidation of the latter blocks its stimulatory activity and promotes immune tolerance. Critical Issues: Reactive oxygen and nitrogen species create an oxidative environment and can be detoxified by superoxide dismutase (SOD), catalase, and peroxidases. Modifications of the oxidative environment influence HMGB1 activity. Future Directions: In this review, we hypothesize that manipulations of an oxidative environment by SOD mimics or by hydrogen sulfide are prone to decrease tissue damage. Both the concomitant decreased HMGB1 release and its redox chemical modifications ameliorate inflammation and tissue damage. I think the following below links may help you in your analysis:
Thanks
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For H2O2 electrogeneration by 2e ORR(Oxygen Reduction Reaction), will H2O2 molecules accumulate on cathode surface?
What measures can I take to avoid accumulation of H2O2 molecules, as it will cause severe H2O2 disproportion reaction.(Path 3 in the following figure)
I just want to try pulsed power supply.
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Hi, I think by using lower current densities and porous platinized electrode and some surfactant reagent in the solution for lowering the surface tension you can overcome to your problem.
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In LiCoO2 battery, why there should be energy gap between the redox potential of Cobalt ion and the Oxygen 2p energy bands?
In General why the voltage of a cell is limited by the top of the anion p-bands of the cathode? I went through some literature, they talk about polaronic state and itinerant holes (Incase of LiCoO2, oxygen will evolve if the redox potential of Co and P-band of oxygen overlap) this concept is out of my scope. Hopefully will get some understanding in these sector . Thank you for your time and energy . 
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Appreciate it Dr. Amiery !!
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We have been working with aromatic carboxylic-amidic compounds and Cu(NO3)2·xH2O salt in various solvents, e.g. water, EtOH, MeOH, DMF and mixtures of them, these in order to generate coordination polymers, and in these trials we have obtained dark residuals. Therefore, in this conditions we aware that we are obtaining redox products instead of the stable coordination polymer. Please provide further advise in order to understand chemical transformations involving amide and Cu ions.
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You're reducing the Cu, what about starting from 4-formyl benzoic acid? if you make the reaction ie 4-amino aniline, you'll have the schiff bases which can be easily reduced to the amine group, those are very stable with copper in different solutions. Be sure to protect the acid group at the beginning. 
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Hi everyone,
I'm currently rewriting an old protocol I once did. The procedure is the following:
- Dissolve decylubiquinone in acidic EtOH
- Add a tiny amount of NaBH4
- Vortex until solution becomes colorless
- Add cyclohexane, vortex
- Add a little H2O, vortex, allow for phases to separate
- Remove organic phase, add more cyclohexane, vortex
- Allow to settle, remove organic phase and pool both organic phases
- carefully dry under N2, resolve in acidic EtOH, analyze and quantify
Now the thing is we're moving to a new lab and we've got an ample supply of n-hexane. I don't want to order a bottle of cyclohexane because there's no other use for it outside of this reaction which is very rarely done.
I see no problem in substituting the cyclohexane with n-hexane here, or am I getting something wrong?
Cheers and thank you
Yaschar
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There should be no difference as their solvation properties are basically identical. Hexane has a slightly lower boiling point so drying under N2 will be faster.
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In some modified Winkler titration method, people add NaN3 in the alkaline iodide solution to eliminate nitrite interference (Alsterberg 1925; Barnett and Hurwitz 1939; Broenkow and Cline 1969). However, in the hypoxic or anoxic zones, where the dissolved oxygen (DO) is severely deficient and hydrogen sulfide (H2S) appears, H2S would interfere the Winkler titration of DO by reacting with azide (N3-) or consequent iodine (I2). The azide in alkaline iodide reagent would somehow be reduced by H2S (Pluth 2013), and the left part of H2S would reduce iodine after acidification (S2-+I2=S+2I-). Therefore, in some cases, the H2S is expressed as ‘negative oxygen’, which is the amount of oxygen equivalent to the amount of H2S produced through reduction of sulphate (Fonselius, 1969). For our Chesapeake Bay cruise in August, 2016, we measured the DO in spectrophotometric Winkler method and determined H2S following the method by Fonselius. The data shows very low DO (<20 umol O2 /L) coexisting with H2S (< 13 umol/L) in some bottom water. Thus, we are looking into the effect of H2S on Winkler DO. As mentioned above, H2S would react with azide, but I have no idea how this redox reaction happens (NaN3+H2S=?) and its extent. Is the azide enough to eliminate the H2S before acidification? The main ions inside the BOD bottles are listed below:
1.       [Mn2+]   23936 umol/L
2.       [I-]           31596 umol/L
3.       [OH-]       63191 umol/L
4.       [N3-]        1215   umol/L
5.       [H2S]        0-13   umol/L
6.       [O2]         5-350  umol/L
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Dear Su
Please follow the article: 
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anyone could you explain like order of redox potential of the material present in the compound/ composites. 
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Do Cyclic voltammtery for your compound. You will get the redox potentials
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How to determine redox conditions in mangrove sediments?
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Michael:
This link would provide you with interesting insights:
Best
Syed
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I have been trying but unable to find some redox reactions or electrolytic reactions of pentacene. Help needed please.
Any helpful link, a tip on how to find them or a little list will be appreciated. Thanks in advance.
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Dear Ahmad,
Attached is a publication which discusses the electrochemistry of pentacene among other properties. You may read pages 37-41 as part of the section  "Electrochemistry and spectroscopy".
Hoping this will be helpful,
Rafik
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Given the structure of a molecule can we theoretically calculate the redox potential of that molecule.
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This tutorial for a specific compound might help:  http://www.chem.yale.edu/~batista/classes/tutorials/redoxpotentials.pdf