electrochemistry

[budullewraagh]

i decided that since my personal lab is severely underfunded i may as well use it to produce some kind of income.

 

first, i wish to electroplate silver onto various items. of course, AgNO3 and AgClO4 are two of the very few Ag salts are reasonably soluble at any reasonable temperature.

 

AgNO3 would be a beast to try to make

AgClO4 would as well, but i have an idea; tell me if you think this is possible and/or sane

 

mix clorox with ammonia, recover gas evolved. burn the gas. problem is attempting to add the Ag. if i bubble it into a solution i'll get plenty of HClO4...

 

 

beyond electroplating, there is electrolysis. hey, it's cheap, no? anybody know of the most efficient way to recover metals from electrolysis? anybody know how long it should take? should i use a car battery in a 500mL beaker of solution?

[Mr_Mediocre]

G'day,

 

Not too sure what you are after here. What source of silver do you have? What are you trying to plate?

 

Avoid perchlorates wherever possible. They ain't good. If you have a source of metallic silver then silver nitrate would be pretty easily produced by dissolving in conc HNO3 and rotovapping down. Give me a few more details about what you want to do and I'll see if I can help you out further.

 

Cheers,

 

MM

[YT2095]

for a start, I`m not altogether certain that your method of preparation for perchloric acid will work to the degree you`de require, there would be other side reaction too, making HCl and HClO3 as well as a small amount of HClO4, and free chlorine.

you`re best bet would be to distil some KClO4 in sulphuric acid, making the real deal

but as MM stated, it`s NOT a chem you want to mess with! it can and does spontaniuosly explode on contact with organic materials, and is in my opinion in the top 3 of strongest acids!

 

go the Nitric acid way, it`s just as easy and you don`t have the explosion or Toxicity risk and you`ll not make insoluble and light sensitive AgCl

[raivo]

Silvering is possible with AgNO3 solution but you need to use some silver piece as anode or your electrlyte will soon be unuseable. Speed of electrolytic reactions is directly proportional to current that passes through solution and is generally quite slow if you do not have tens or hundreds ampers of current.

 

When plating something you can not use more current than some ampers per square decimeter of the plateable thing. Otherwise metal particles will not settle properly on surface and you will get very ugly coating ( in best case).

 

One amper of current must flow through electrlyte during 26.8 hours to move one mole of substance ( in case ions are charged +1 or -1, otherwise speed is proportionally lower)

[budullewraagh]

the lot of you speak as if i were a professional chemist. alas, i am not even 16 yet and of course i have extremely limited resources for chemicals. sadly, in the states, chemicals are more controlled than in england. there are no nitrate sources available commercially, nor is H2SO4 unless i open many batteries.

 

unfortunately, i have recently found that it would be impossible for me to obtain or synthesize silver salts.

 

yt, you spoke of perchloric acid. just out of curiosity, what is the highest molar concentration one can create with it? how does it compare with HF:SbF5?

[Mr_Mediocre]

I know you said `just out of curiosity' but DO NOT TOUCH HF!!! OR SbF5!!! You will be dead in hours. High concentrations of F- are extremely lethal. It is my social duty to say so, even though you may know this already. To satisfy your curiosity, 70 % is the highest concentration that one can get perchloric acid to. 33 % (10 M) is the highest HCl can get to. HF can be prepared as a 73 % (44 M) solution. I don't know about SbF5 but you prolly don't wanna mess with it.

 

http://www.chm.bris.ac.uk/safety/hf.htm

 

Cheers,

 

MM

[budullewraagh]

I know you said `just out of curiosity' but DO NOT TOUCH HF!!! OR SbF5!!! You will be dead in hours. High concentrations of F- are extremely lethal.

please see my above post regarding my lack of resources :\

 

also, i've seen 12M HCl at room temp. that is saturated. i have a 25 year old container of "muriatic acid" that is extremely concentrated, and i understand the value in not bathing in it. everytime i open it, HCl(g) is released. i understand NOT to breathe it in, since it is similar to mustard gas. thanks for the warning, however.

 

i think you misread the HF:SbF5 thing. a 50% solution of HF:SbF5 is 10^18 x as acidic as conc H2SO4. this is because every molecule of HF yields H+ and the F- makes SbF6- since Sb+5 is so electropositive.

[Mr_Mediocre]

Are you SURE you have seen 12M HCl? Or did someone just write that on the label. I would go as far as to guarantee that you didn't see it in aqueous solution (and I can't think of any other solvent that would hold more, except POSSIBLY ionic liquid salts, but I doubt it). The ABSOLUTE most concentrated you can get an HCl aqueous solution is 33-35%, which is just under 10M. I will bet my life on it....and I am pretty fond of the 'ol `life'. You have a link to chemfinder which will prove it to you

 

I see what you mean about the HF/SbF5 thing now, my apologies. Where did you get that from out of curiosity? 10^18 more acidic than H2SO4 doesn't SOUND right but I am always willing to learn .

 

Cheers,

 

MM

[budullewraagh]

i did a bit of research:

according to my lange's handbook of chemistry (actually it's hardly a handbook being thousands of pages), at 0 degrees C, 512g HCl (aq) in 1L H2O is saturated. at 10 C, 475g. at 20 C, 442g. at 30 C, 412g. at 40 C, 385g.

molecular mass of HCl=36.453 AMU

1 mole HCl=36.453g

12*36.453=437.436g

using my graphing calculator, i figured that 437.436g would be soluble at approx 22.6 C, which is equal to 72.68 F, which is approx room temp. turns out that 12M HCl at room temp is approx saturated.

 

 

 

Where did you get that from out of curiosity? 10^18 more acidic than H2SO4 doesn't SOUND right but I am always willing to learn

i thought the same thing when i first read it in my chem textbook. it's crazy stuff; it was used to make the first "stable" carbocation.

 

check these links:

http://www.cmm.upenn.edu/research/acid.html

http://chemserv.bc.edu/Department/Faculty/hoveyda/pka.htm

http://poohbah.cem.msu.edu/Portraits/PortraitsHH_Detail.asp?HH_LName=Olah

http://reedgroup.ucr.edu/projects/proj1.htm

[apathy]

HF:SbF5 is called "magic acid" but I doubt you'd hallucinate if you took it, (well, you might before you die)

 

Conc. HCl is 12 M ( i don't know if that got answered) when you buy it in a bottle.

[apathy]

hehe! If you aren't 16 yet, then who is that in your avatar?

This whole time I thought it was you.

 

[budullewraagh]

haha! the man in my avatar is a certain music professor i know. he's known for being uber cool and all that, and for a short time there was a website http://www.whatwouldlangdo.com and his picture (my avatar) was on it. so, i now use his picture as my avatar. i do not know why.

 

by the way, can you see my sig image?

[apathy]

what's a sig image, and no, i guess i don't see it...

[budullewraagh]

an image in one's signature

[budullewraagh]

also, can anybody here tell me the best way to isolate aqueous alkali and alkaline earth metals from their salts? of course, using electrochemistry is the best way, but i was wondering the most efficient ways to do so.

[Mr_Mediocre]

Well thats a little bizarre...I went back and worked from the value of 62g/100ml that chemfinder preaches. This works out to 1.7 mol HCl in 100ml..which is 17M HCl!! I have certainly not seen or heard of this, hell, I hadn't even heard of 12 M HCl. When we buy in Conc. from Aldrich it always comes as 10 M.

[Mr_Mediocre]

Ah well learn something new everyday!

[Mr_Mediocre]

Regarding your salt separation, it would depend greatly upon WHICH and HOW MUCH of those metals you have in solution.

[budullewraagh]

i just remembered; NH3 solutions reduce alkali metals to form anions. would an NH3 solution be able to reduce Mg+2 in a solution of NH3 and MgSO4?

[Mr_Mediocre]

I think it is the other way around...group two (alkaline, not alkali) metals (not 2+ cations, but the metals) are strong reducing agents as they easily give away the two outershell s electrons. They can't do any further reducing once they are fully oxidized to 2+, which they will be if they are in solution. I'm guessing you are just working with household chemicals in your endeavours? Where are you getting you sulfate from?

[budullewraagh]

epsom salts.

 

i understand how group 2 metals oxidize, but i was wondering if NH3 solutions have the same effects on group 2 ions as they do on group 1. also, i was wondering NH3 solutions would reduce aqueous group 1 ions since they reduce neutral group 1 metals.

[Mr_Mediocre]

NH3 doesn't reduce neutral group one metals, it oxidizes them.

 

It is also not going to be able to donate an electron to reduce a group 1 or group 2 cation to its native metal state. Even if it did, as soon as the neutral metal was produced it would react with the water in your ammonia solution. You'd need to do electrolysis of the salt, and in a non-protic solvent like dry acetonitrile or DMF. If you DO go down this path, be very careful and read up very well.

 

Best of luck

[budullewraagh]

actually i messed that up; it wouldnt work in H2O but it would work in liquid NH3...perhaps in a solution of NH3 in an alcohol?

 

NH3 doesn't reduce neutral group one metals, it oxidizes them.

my source: The Elements by John Emsley, page 104. cheesy name but it says: Oxidation states:

Li^-1 (s^2) Li solutions in liquid ammonia

[Mr_Mediocre]

Does the book say how the Li- was formed initially?

 

My source: Jonathan C. Wasse, Shusaku Hayama, Sotiris Masmanidis, Sarah L. Stebbings, and Neal T. Skipper, Journal of Chemical Physics,

 

The microscopic structures of lithium–ammonia and sodium–ammonia solutions have been measured by the technique of isotopic labeling in neutron diffraction, at and above the metal–nonmetal transition that occurs in the range 2–8 mole percent metal (MPM). Substitution of *Li by 6Li has been used to obtain the lithium-centered first-order difference function at 8 MPM and 230 K. This function shows us that the lithium cations are strongly solvated by 4 ammonia molecules. Substitution of *N by 15N has then been used to probe the nitrogen-centered structure in lithium–ammonia solutions at 4, 8, and 12.5 MPM and sodium–ammonia at 12.5 MPM. These functions give us new insight into both the disruption of hydrogen bonding as alkali metal is added to ammonia, and the solvation structure of the sodium cations. The former manifests itself through a progressive loss of the hydrogen-bonded N–D peak at ~2.4 Å. The latter appears as an N–Na shoulder at ~2.5 Å, and shows us that sodium is solvated by ~5.5 ammonia molecules. In contrast to previous data for saturated (~21 MPM) metal–ammonia solutions, we do not observe intermediate-range ordering of the solvated cations at the concentrations studied here

[Mr_Mediocre]

sorry forgot rest of the reference.

 

Jonathan C. Wasse, Shusaku Hayama, Sotiris Masmanidis, Sarah L. Stebbings, and Neal T. Skipper, Journal of chemical physics, 2003, (118), 7486

 

http://content.aip.org/JCPSA6/v118/i16/7486_1.html