Questions related to Redox Chemistry
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.
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!
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 ?
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?
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,
Abu Saeed Baidya
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.
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.
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.
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.
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?
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.
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?
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?
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) ?
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.
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
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?
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.
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.
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!
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:
- FeIII[FeIII(CN)6] + Na+ + e- <=> NaFeIII[FeII(CN)6] 1.2V vs SCE
- NaFeIII[FeII(CN)6] + Na+ + e- <=> Na2FeII[FeII(CN)6] 0.2V vs SCE
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.
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!
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.
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.
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
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.
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
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?
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.
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.
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:
- Why is it called dinucleotide when I see only one nucleotide as its component (adenine)? or is it nicotinamide a nucleotide too?
- 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.
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?
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+.
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?
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?
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?
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 ?
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.
We are interested in a brain mimetic solution or material to record voltammograms using microelectrodes.
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.
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.
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.
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!
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?
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.
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?
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.
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.
Qaiser Ali Khan
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.
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.
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?
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.
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
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.
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 .
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.
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
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
anyone could you explain like order of redox potential of the material present in the compound/ composites.