Harry_-

Sorry I don't have lots of time to write a reply at the moment I'm working on a lab. However if you can find the redox potentials for the cell you could use: ΔGo= nFE n is the number of moles used in the balanced equation F is faraday's constant of 96500 Cmol-1 E is your cell potential under standard conditions (reduction - oxidation potentials) If you need help ask or there are lots of good online videos on that equation. Regards, Harry

Thanks so much. I'm preparing a lab experiment and can't really reference that, are there any papers that I could use? Also I notice that the yield for that experiment is 72% whereas without ammonium nitrate and with a lot more cupric acetate monohydrate the yields on the other paper are 90%. Why is this?

They say nitric acid can be used but nothing about ammonium nitrate

What about the ammonium Nitrate. I can't find any papers on it being used in this reaction

Does anyone know how this reaction proceeds with the following reactants? Glacial acetic acid (i presume solvent) Cupric acetate monohydrate Ammonium Nitrate

Hello, I'm researching into the 2019 Nobel Prize for chemistry, and therefore lithium ion batteries. Wittingham's battery was a lithium anode, a titanium disulfide cathode and a lithium perchlorate in dioxolane electrolyte. I have always thought that batteries rely on a displacement reaction split into two halves. However in animations that I have seen show the electrons at the cathode being picked up by the very lithium ions that released them and that have just moved from the anode to the cathode. This doesn't make sense to me as if the lithium is picking up the electrons it just gave away what was the cause for this reaction to take place? I would have expected the electrons to be picked up by the electrolyte. What is happening here? Any help would be greatly appreciated, Harry

Yes, I did take these into account, I can't remember what the exact values were for each compound but I used the value on the bottle. However if the bottles were contaminated/ not pure (very possible as I was using school chemicals which students in practicals can just scoop out what they need directly from the bottle using a spatula) wouldn't this cause my entire gradient to be off as in dilution I wouldn't be making the concentrations of the solutions I thought I was making. I also was only using a 2 d.p scale (only one I have) so I cannot expect each solution to be the exact same concentration.

This is also weird to me as if the electrometer was in parallel, I would expect a slightly lower value for standard emf (due to the lost volts not completely being cancelled out by the resistor) however if the electrometer was in series I would have expected a very precise value of E0. What if the ZnSO4 or my CuSO4 wasn't perfectly pure therefore each of my subsequent solutions were slightly off so even my standard emf (if [reducing agent] was larger than [oxidising agent]) would be lower than expected. As I added water to my solutions to make them more dilute this could have changed the shape of my graph. I'm also having a hard time believing I would put a voltmeter in series 😂

Unfortunately I do not have the equipment anymore. My circuit drawing is setup for measuring current and not voltage, It has been a while since I did this experiment so I'm not sure if I put the electrometer in series with the resistors or if I set it up the correct way. Its embarrassing I know. However, even if my electrometer was in series with the resistor (100M Ohm) the voltmeter's impedance is 5x1013. Compared with 100x106 of my resistor in series, shouldn't this mean that my results should still be accurate to the degree that I have.

My concentrations that I wanted to use started at molarity 2 and were tested against the other electrolyte of molarity 1.5 and 1. I had an issue with creating 2 molarity solutions and the solutions were saturated, so I doubled the volume to make molarity 1 solutions as a start. I could be possible that I made up the solutions wrongly as I was working very fast

Standard Electrode potentials: Zinc= -0.76 copper= 0.34 (both of these should account for the reaction using/ taking 2 moles of electrons) 0.34--0.74=1.1v It is weird that I am pretty much exactly a factor of 2 off, which is why it makes me believe that iit is a calculation error rather than a practical one. For example, if I was only using 1 mole of electrons in the balanced ionic equation my value for F would be perfect. Here is the electrometer I used's manual: http://physics-astronomy-manuals.wwu.edu/Keithley 614 Electrometer Manual.pdf Here is a diagram of the apparatus setup, the electrometer being in series shouldn't be a problem should it? I assumed it had internal parallel circuits to record voltage as it also records current. Finally, my x-axis was ln([Oxidising Agent] / [Reducing Agent])

I'm not sure what you mean by 'zero point' please could you expand. I used my electrometer to measure the voltage, it was zero-checked before the experiment and then connected to my circuit. What do you believe to be off in the recorded voltages? From my calculations the standard emf should be 1.1v, the value I got experimentally was 1.06 (mainly due to insufficient resistance in the circuit to negate internal lost volts).

Thanks

@studiot

Hi, I know it has been a while, so much has been going on. I have just been writing up my results from this experiment and have come across some trouble with my calculated value for Faraday's constant. It is almost perfecty half what it should be, one of two problems could have occured: 1) In my stupidity I created the wrong concentrations of solution (not very probable as all sorts of anomalous points would be created when I then made my more diluted/concentrated solutions) 2) Something is wrong in my calculations Could any off you check these calculations to see if there are any problems. I used a Zinc anode and Copper cathode so my value for number of electrons transferred in the balanced ionic equation should be 2. If you require any more info, just say :). Hope you are all well. Harry