Sciencemadness Discussion Board - The Short Questions Thread (4)

Question: Slightly confused about carbon zinc battery.

I have disassembled a few lantern batteries -- for electrodes and other useful things.
First step in purifying the MnO2 paste is to mix with water to form a slurry and then filter. The intention is to dissolve the ZnCl2 or NH4Cl that acts as an electrolyte.
The other day I ran a current through it and discovered to my surprise that a nice grey metal plated out on the cathode. I thought I had a solution of ZnCl2 but of course Zn cannot be electrolytically reduced in aqueous solution. So, what is it that plated out? Do some batteries contain tin compounds or bismuth or something else strange? What is it that I have got?

Bump – cos it disappeared off the bottom of the page.

Molecular Manipulations - 23-2-2015 at 16:07

Umm who told you zinc cannot be plated in aqueous solution? They lied. Have you not heard of galvanization?
Just because the activity series says zinc displaces hydrogen doesn't mean it can't be electroplated, because it can and is, by the ton.

j_sum1 - 23-2-2015 at 16:33

Bloggy mentioned it a while ago. I may have my wired crossed. I'll come back to this one.

xfusion44 - 23-2-2015 at 18:28

Hi! I would need some help...

Would it be possible to make aluminium citrate by reacting aluminium, hydrogen peroxide and citric acid? I think that first, there should be reaction between Al and H2O2, to make Al2O3, which would (or wouldn't?) then react with citric acid, to make C6H5AlO7. I don't need aluminium citrate for any particular purpose, I would make it just for my compound collection... And if anyone could give me some data on C6H5AlO7, that would be really helpful (looks like wikipedia doesn't yet have the page about aluminium citrate). Thanks in advance!

sparkgap - 23-2-2015 at 18:38

Quote: Originally posted by Bot0nist

I was discussing some homemade stain removers on another site and this came up as a stain remover. Is it just the fact that cream of tartar is a common OTC acid salt, and the drop in pH is what matters, or if there is some other effect that can't be achieved with a similar salt, or the lowering of pH with an acid solution? In regards to the reported stain removing effects, of course.
Thanks in advance.

Tartrate is a pretty good chelating agent, so you'd want to use cream of tartar on rust stains and the like. The peroxide will of course bleach anything the cream of tartar can't deal with.

sparky (~_~)

Bot0nist - 23-2-2015 at 19:57

Thank you! I was too focused on how it may react with affect the peroxide.

j_sum1 - 24-2-2015 at 01:31

Bloggy mentioned it a while ago. I may have my wired crossed. I'll come back to this one.

Let's just put this one down to a brain-fart.

I ran a current through the solution a few weeks ago and found what I thought was zinc.
Maybe a week ago I read something here, I think from Blogfast that made me question it. Or maybe I misremembered something. Anyway, I cannot relocate the comment so never mind.
A couple of days ago I was looking up my redox potentials and was comparing Zn with the wrong line and convinced myself that it wasn't possible.

Classic brain fart material. Nothing to see here. Carry on.

xfusion44 - 25-2-2015 at 20:13

Quote: Originally posted by xfusion44

Hi! I would need some help...

Would it be possible to make aluminium citrate by reacting aluminium, hydrogen peroxide and citric acid? I think that first, there should be reaction between Al and H2O2, to make Al2O3, which would (or wouldn't?) then react with citric acid, to make C6H5AlO7. I don't need aluminium citrate for any particular purpose, I would make it just for my compound collection... And if anyone could give me some data on C6H5AlO7, that would be really helpful (looks like wikipedia doesn't yet have the page about aluminium citrate). Thanks in advance!

I tried that and it didn't work, so thanks anyway

Jylliana - 26-2-2015 at 04:35

What type of glass are (large) petridishes usually made of? Soda glass or borosilicate?
I was rather surprised that it shattered so easily earlier today from a relatively weak heat source(very small Bunsen flame, no larger than a lit candle). I couldn't find a Pyrex or some other brand 'stamp' anywhere...

It didn't just crack and break btw... it shattered 'explosively'. Glass everywhere. Luckily I was at the other side of the lab doing other stuff... O_O

Zombie - 26-2-2015 at 14:33

Flint glass is most common in the China made Petri's.

I broke quite a few cheap dishes I bought on Ebay in the pressure cooker. I've since been buying only Pyrex / Corning.

Depending on what you are doing I have found the plastic Gosslin dishes from corning do take the heat of a 15psi pressure cooker.

Bottom line is if it is not a marked name brand, it is most likely Flint Glass.

Came back to add on this answer...

"Product Description
The American Educational Flint Glass Culture Petri Dish is a 98mm outside diameter (OD), 90mm inside diameter (ID), 18mm high petri dish bottom for culturing cells and other scientific applications, and is clear for viewing contents. Flint glass may contain an occasional small bubble or surface blemish and is for unheated use. Petri dishes are commonly used in laboratories and other scientific applications."

http://www.amazon.com/American-Educational-Culture-Height-Bu...

Just one example. :cool:

[Edited on 2-27-2015 by Zombie]

xfusion44 - 26-2-2015 at 20:18

Hi, I saw an experiment on youtube when H2O2 reacts with KMnO4... I decided to try it myself too and it's pretty vigorous. But both of the reactants are oxidizers, so even oxidizers can be oxidized if there is another, stronger oxidizer, like KMnO4?

[Edited on 27-2-2015 by xfusion44]

j_sum1 - 26-2-2015 at 21:06

We have had this discussion recently. http://www.sciencemadness.org/talk/viewthread.php?tid=61637
As it turns out H2O2 is a stronger oxidiser than KMnO4 so you have that bit the wrong way around.

To add to the confusion, H2O2 is a strange beast. There are situations where it acts as a reducer. Its behaviour is pH dependent. There are four separate equilibria at play here.

What you might be observing is a catalytic effect and the situation is simple decomposition of H2O2 to H2O and O2. I don't recall if the Mn in KMnO4 catalyses this reaction but it sure wouldn't surprise me. (KI does in the classic elephant toothpase reaction.) If it is a decomposition then we are not really talking about redox at all.

xfusion44 - 27-2-2015 at 12:21

We have had this discussion recently. http://www.sciencemadness.org/talk/viewthread.php?tid=61637
As it turns out H2O2 is a stronger oxidiser than KMnO4 so you have that bit the wrong way around.

To add to the confusion, H2O2 is a strange beast. There are situations where it acts as a reducer. Its behaviour is pH dependent. There are four separate equilibria at play here.

What you might be observing is a catalytic effect and the situation is simple decomposition of H2O2 to H2O and O2. I don't recall if the Mn in KMnO4 catalyses this reaction but it sure wouldn't surprise me. (KI does in the classic elephant toothpase reaction.) If it is a decomposition then we are not really talking about redox at all.

Hmm, well that's surprising. In the video on yt it says that H2O2 is being oxidised and decomposed.

Thanks for your answer!

quantumcorespacealchemyst - 1-3-2015 at 17:55

If washing arsenic acid and/or arsenous acid crystals in diethyl ether, is there any peroxide formation or other reactions to be aware of?

Arsenic acid, H3AsO4, is an oxidizer. If the contact is brief, only to dislodge them from where they crystallized, will peroxides form?

[Edited on 2-3-2015 by quantumcorespacealchemyst]

Sulaiman - 2-3-2015 at 05:23

I have slowly collected (eBay) a Quickfit 14/23 distillation setup
hopefully for volatiles, acids, hydrocarbons etc.

How do I use the thermometer pocket (SH 4A) correctly?

I had a quick Google search but I'd like to hear any suggestions,
I'm thinking about oil for rapid temperature response
or sand (I have 400 mesh silicon dioxide) to protect the bulb from my clumsiness.

How do you guys use a thermometer pocket?

EDIT:
The thermometer goes into a pocket that looks like this
http://www.scilabware.com/Thermometers/Pocket/p-130-508/
which fits into something similar to this
http://www.scilabware.com/Flasks/Pear-shaped/Claisen-Vigreux...
except instead of a vigreux section there is a vacuum jacket section.

Initially I will use one of my mercury thermometers
later one of my stainless steel clad thermocouples.

[Edited on 2-3-2015 by Sulaiman]

Haloform products + Iodine powder

quantumcorespacealchemyst - 2-3-2015 at 20:56

What is the reaction of I2 with NaOCl haloform products? NaOCl(bleach) + CH3COCHCH3 + (NaOH in bleach)---->CHCl3 + NaOH + NaOOCH3 and then adding equimolar mols of I (as Cl) to the solution?

[Edited on 3-3-2015 by quantumcorespacealchemyst]

if 3NaOCl + CH3COCH3 + NaOH ---> 3NaOH + NaOOCCH3 + CHCl3

does the sodium acetate affect the pH too much effecting the NaOH reacting with the I2?
There is a purple solution after adding half the mols equivalent of Cl, of I (1/2I2). It was heated under boiling and cooled.

[Edited on 3-3-2015 by quantumcorespacealchemyst]

Sulphuric acid from thiosulfate?

Argentum - 3-3-2015 at 03:16

Just saw in another forum that H2SO4 was made by sodium thiosulfate and H2O2. He wrote the following equation

Na2S2O3 + H2O2 → H2SO4 + Na2SO4 + H2O

The equation can be balanced but it doesn't mean that the reaction will work that way.

I tried it myself yesterday, added one drop of H2O2 to some crystals of Na2S2O3, a lot of fizzing and a pleasant smell of H2S... I disposed the products and decided to ask before trying again...

I'm thinking that would be an easy way for producing sulphuric acid and someone would have tried it before

Is this possible? Or just another way of making H2S...

PHILOU Zrealone - 4-3-2015 at 05:29

Quote: Originally posted by quantumcorespacealchemyst

What is the reaction of I2 with NaOCl haloform products? NaOCl(bleach) + CH3COCHCH3 + (NaOH in bleach)---->CHCl3 + NaOH + NaOOCH3 and then adding equimolar mols of I (as Cl) to the solution?

[Edited on 3-3-2015 by quantumcorespacealchemyst]

if 3NaOCl + CH3COCH3 + NaOH ---> 3NaOH + NaOOCCH3 + CHCl3

does the sodium acetate affect the pH too much effecting the NaOH reacting with the I2?
There is a purple solution after adding half the mols equivalent of Cl, of I (1/2I2). It was heated under boiling and cooled.

[Edited on 3-3-2015 by quantumcorespacealchemyst]

The purple colour may come from I2 dissolved into CHCl3....
It is not very clear to me what you intend to do?
-NaOH + I2 will produce some NaOI and NaOIO2 (NaIO3) by disproportionation and depending on heat.
-I2 and NaOCl will also lead to NaOI and NaOIO2 (NaIO3) (and maybe NaOIO3(NaIO4)).
-I2 and aceton in basic media usually lead to orange precipitate of CHI3 (iodoform), it is a characteristic test for methyl-ketonic group.

Just as a remark, be carreful with what you write:
NaOCl + CH3COCHCH3 + (NaOH in bleach)---->CHCl3 + NaOH + NaOOCH3
==> CH3-CO-CH-CH3 is uncorrect ... You miss one H atom to get tetravalent carbon -CH- is trivalent... so better write CH3-CO-CH2-CH3 is Methyl-Ethyl-Keton (MEK) or (2-)butanone.
I suspect you wished to write aceton /propanone CH3-CO-CH3
==> NaOOCH3 is uncorrect... would be sodium peroxomethylate! I suspect you wished to write NaO2C-CH3 (sodium acetate)

In your reaction you may get orange CHI3 dissolved into CHCl3 or intermediary CHICl2, CHI2Cl since you may have a certain level of competition between the two halogens...

[Edited on 4-3-2015 by PHILOU Zrealone]

Molecular Manipulations - 4-3-2015 at 06:59

Just saw in another forum that H2SO4 was made by sodium thiosulfate and H2O2. He wrote the following equation
Na2S2O3 + H2O2 → H2SO4 + Na2SO4 + H2O

It's always best to give a reference to what forum or paper you're citing.
I think you mean this one. I don't see why you didn't bother balancing the equation or even copy/pasting it from your source, where it was already balanced:
Na₂S₂O₃ + 4 H₂O₂ (aq) → Na₂SO₄ (aq) + H₂SO₄ (aq) + 3 H₂O (l).

I tried it myself yesterday, added one drop of H2O2 to some crystals of Na2S2O3, a lot of fizzing and a pleasant smell of H2S... I disposed the products and decided to ask before trying again...

Well asking isn't going to change the nature of this reaction, but disposing of that mess was a good idea. I'm guessing that the equation above predominates over the formation of hydrogen sulfide. I'm nearly certain that the formation of hydrogen sulfide was caused by an excess of sodium thiosulfate. In other words, don't add hydrogen peroxide to sodium thiosulfate, add a solution of the latter to hydrogen peroxide, and add slightly less than the stoichiometric quantity. Hydrogen sulfide is a reducing agent and would not likely form with enough of a strong oxidizer present.

I'm thinking that would be an easy way for producing sulphuric acid and someone would have tried it before

I'm thinking of all the methods of making the easily available sulfuric acid, this is one of the lousiest, but each to his own. It probably will work, but definitely do this outside or in a fumehood if at all.

[Edited on 4-3-2015 by Molecular Manipulations]

j_sum1 - 5-3-2015 at 04:01

I know this question has come up before, but I don't recall seeing an actual answer.
This product, is it actually pure molybdenum? If not, then what else is in it?
And what do mobile phone repairers use it for anyway? It is all over eBay and other places. I would like to know some details before I get some.

Zombie - 5-3-2015 at 06:33

It's used as a knife. Like pulling out a car windshield.

https://www.youtube.com/watch?v=qowrEiDsdTo

They use Molybdenum because of its high tensile strength.

No idea on the purity tho. I guess you would have to find the manufacturer, and verify that yourself.

I wouldn't trust asking a seller. unless the label or packaging clearly states this.

Sulaiman - 5-3-2015 at 19:59

Since I am buying some sodium metal for the 'transparent sodium' experiment
I would like to know of other uses for sodium metal
I may make some sodium methoxide to try small scale transesterification

other than these two are there any other interesting uses for sodium metal?

Zombie - 5-3-2015 at 20:12

https://www.youtube.com/watch?v=RAFcZo8dTcU

Sulaiman - 5-3-2015 at 21:02

as a bearded white muslim convert that lived in a muslim country for 20 years
I'm already paranoid enough,

I'll give that one a miss for now, thanks.

Zombie - 5-3-2015 at 21:47

You poor man! Where's all the fun in being normal?

quantumcorespacealchemyst - 5-3-2015 at 21:55

Quote: Originally posted by The Volatile Chemist

Quote: Originally posted by quantumcorespacealchemyst

I get it all the time on things in my lab when evaporating HCl solutions. I think it's condensed HCl, or a salt of Cl^-, that is, if you have ammonia vapors nearby or something.

isn't that funny though, I thought I resupplied the NH3 after I got some HCL (and it was a while since having NH3 in the lab). Possibility is urine, I collect in sealed bottles. The smell of the white salts doesn't smell like anything I know, having made urea and Ammonium chloride (presumably) and smelling the piney scent. I don't know.

[Edited on 6-3-2015 by quantumcorespacealchemyst]

quantumcorespacealchemyst - 5-3-2015 at 22:10

PHILOU Zrealone thanks, you are right about the chemical formulas, i got it wrong, meaning CH3COCH3 and NaOOCCH3.

[Edited on 6-3-2015 by quantumcorespacealchemyst]

j_sum1 - 5-3-2015 at 22:36

Quote: Originally posted by Zombie

It's used as a knife. Like pulling out a car windshield.

https://www.youtube.com/watch?v=qowrEiDsdTo

They use Molybdenum because of its high tensile strength.

No idea on the purity tho. I guess you would have to find the manufacturer, and verify that yourself.

I wouldn't trust asking a seller. unless the label or packaging clearly states this.

"The easy way??? That video looked awkward to say the least.

If Mo is udes for tensile strength then it is likely that it contains a reasonably high proportion. It will probably suit my purposes.
I don't have a fine wire in my element collection so it will add something there. And It will be more than adequate for my current project. O am making gas discharge tubes of various elements. I want something with a high MP and that will resist erosion by sparking. I decided I wanted something a bit thinner and more economical than the tungsten wire I have experimented with. An Mo phone slicer will do just fine.

peroxide formation?

quantumcorespacealchemyst - 9-3-2015 at 21:58

Does washing/dislodging crystals of Arsenic acids, with Diethyl ether, carry risks of peroxide formation?

j_sum1 - 9-3-2015 at 22:54

Quote: Originally posted by quantumcorespacealchemyst

Does washing/dislodging crystals of Arsenic acids, with Diethyl ether, carry risks of peroxide formation?

You have asked that question before.
You might well be one of very few people on this board who are prepared to play with arsenic.

j_sum1 - 11-3-2015 at 03:35

Just how nasty is CS2?

I wanted to demonstrate to my students the dissolving of sulfur. It was in the context of a series of experiments on sulfur and sulfur compounds. The aim was to build up an understanding of the relationship between intermolecular forces and physical properties. Solubility in various solvents of course being a physical property.

While completing my risk assessment documents I found that CS2 was a prohibited substance -- that is according to the software being used (which admittedly belongs to a different jurisdiction but we use it anyway.) Restating this, it is not allowed for state schools to use the stuff at all. I am at a private school so I have a loophole.
Risks obviously are toxicity -- ingested, fumes and skin. Also volatility and explosive combustion.

I decided to proceed with the dissolving as a demonstration on a micro scale. About 0.5mL of CS2 dissolving some sulfur in a test tube. Full protective kit and completed in a fume cupboard. Interesting contrast between the carbon disulfide and xylene and some good learning happened.

But it leaves open for me the question of how dangerous this stuff is. Any thoughts and comments?

j_sum1 - 11-3-2015 at 04:02

Quote: Originally posted by Sulaiman

it appears that CS2 is a risk in industrial environments but low risk in small ammounts.
http://www.epa.gov/airtoxics/hlthef/carbondi.html

unless you dissolve yellow phosphorous in it then pour the liquid over a small pile of sodium chlorate and wait .....

I only did this experiment once, a teaspoon full cracked the concrete slab that it was on!

Ok. On reading that, the RfC is 0.7 mg/m^3. That is, regular exposure to anything greater than that is estimated to have chronic effects.
The smell threshold is only 0.05 mg/m^3. This suggests that as far as inhalation goes, if you can't smell it then you are below the level where damage is going to be caused. I am happy with that. I like it where the smell threshold is lower than the toxicity threshold. Little chance for nasty surprises.

The procedure I used meant that no smell was detected and there was no chance of it contacting skin. I conclude that our risk assessment process is a little over-cautious in this particular case.

Bert - 11-3-2015 at 04:18

I wanted to demonstrate to my students the dissolving of sulfur. It was in the context of a series of experiments on sulfur and sulfur compounds. The aim was to build up an understanding of the relationship between intermolecular forces and physical properties. Solubility in various solvents of course being a physical property.

Perhaps substitute hot toluene? You can recrystalize without needing to evaporate the solvent then-

http://www.sciencemadness.org/talk/viewthread.php?tid=3902

j_sum1 - 11-3-2015 at 04:45

Thanks for that Bert.
Unfortunately no toluene available and only a few mL of xylene. (I have ordered a bottle of each because I think a lab without these is deficient. Even if it is just a school lab and does very little organic and (at the moment) few separations involving solvents.)

I found that xylene dissolved a little sulfur but that amount increased dramatically with a little heating -- just the same as GarageChemist reported for the toluene. It would be good to do a comparison of the two solvents (although on reading down the thread you linked I can see that it has been done at least in part.) CS2 however was a much better solvent. A lot of sulfur dissolved.

As an aside, I love the idea of growing some larg(ish) sulfur crystals. They would be great in the element collection.

I love the way that around here you ask one question and you open a box full of interesting possibilities. Thanks Bert.

Bert - 11-3-2015 at 05:07

I like the look of the large Sulfur crystals from toluene too-

In USA, I can just go to the paint store and buy quarts of xylene or toluene- No questions asked. They can be re-distilled if purity is an issue-

Australia controls these?

The idea of adding Phosphorus dissolved in a flammable and fast evaporating solvent to potassium chlorate had not occurred to me... Probably because I spend my time with chlorate mixtures trying to keep them from going off prematurely, rather than encouraging it!

Was it spontaneously explosive on drying, or did it require heating/shock/friction to explode?

Phenolphthalein indicator preparation

Sulaiman - 11-3-2015 at 05:17

I have some Phenolphthalein powder from an old chemistry set.
Normally it is dissolved in alcohol for use as an indicator, but at the moment I have
methylated/de-natured alcohol with a purple dye in it, or pure acetone.
Is there any problem using acetone?
or should I distil the methylated spirit to get colourless alcohol?

The tube that it is in reads : diluted with salt, 1:14
I don't know why that is done but I assume that very little sodium chloride will dissolve in pure acetone,
and some may dissolve in the alcohol due to the water content?

EDIT: just answered my own question,
added 3.5g of the powder to 50 mL acetone, mixed thoroughly, decanted the liquid off of the undissolved salt, tested ..OK.
...............................................................................................................

Bert, it explodes after the CS2 has evaporated, leaving the phosphorus and oxidiser in intimate contact,
it took about 10 to 20 minutes to spontaneously detonate (decades ago, in my youth, memory fading)

[Edited on 11-3-2015 by Sulaiman]

j_sum1 - 11-3-2015 at 05:18

Quote: Originally posted by Bert

I like the look of the large Sulfur crystals from toluene too-

In USA, I can just go to the paint store and buy quarts of xylene or toluene- No questions asked. They can be re-distilled if purity is an issue-

Australia controls these?

I was referring to the stocks in my school lab.
Xylene is available at hardware stores although it is not especially cheap. Not sure of the purity. Toluene I haven't seen but that doesn't mean that it is unavailable. There are a huge number of paint thinner products that aren't fully labelled. I have no idea what is in them. (Ditto for oils and paraffins. They all say "petroleum distillate", which is not that useful.)

I have a litre of xylene at home and so that might do just fine for a sulfur crystallisation. I might work out a way of lowering the temperature slowly first.

DrMario - 14-3-2015 at 15:14

What is the best glue to attach nylon to steel (or other metals)?

Sulaiman - 14-3-2015 at 15:34

Due to the significantly different coefficients of thermal expansion for steel and nylon
and the stretchiness of nylon I think it will be difficult,
physical methods such as screws, bolts, clamps or straps would be the most reliable,
'stretchy' adhesives such as silicone or rubber may work for light loading.

I'm not an expert, just based on experience.

Zombie - 14-3-2015 at 16:53

Quote: Originally posted by DrMario

What is the best glue to attach nylon to steel (or other metals)?

Depending on the application... There are many adhesives that will work.

The key is to rough up the mating surfaces with a grinding wheel or 60 - 80 gritt sand paper. The rougher the better.

The best overall adhesive I can think of for MOST applications would be 3M's 5200 Marine adhesive.

It's the one "glue" I despise in the marine trades. You can not cut / pry, this stuff off of anything.

Argentum - 19-3-2015 at 15:22

Something I've seen but can't find anywhere

When mixing H2O2 with pool water treatement pills it gave a red chemiluminiscence, am I right?

If I am, what is the reagent in the pills? (I think trichloroisocianuric acid, but can't remember)

How to choose a stir bar?

Sulaiman - 20-3-2015 at 11:53

I am going to make ionic silver solution for medicinal use
and I want to stir the solution whilst the silver wires are electrolytically ionised
my questions are a bit vague so any explanations of general principles appreciated

1) which is best, a small stir bar rotating quickly or a large bar rotating slower?

2) which is the best shape?

2.1) ovoid, e.g http://www.ebay.co.uk/itm/2pcs-Model-A-Teflon-PTFE-Magnetic-...

2.2) rod, e.g. http://www.ebay.co.uk/itm/2pcs-Model-C-PTFE-Teflon-Magnetic-...

2.3) rod with ring, e.g http://www.ebay.co.uk/itm/2pcs-Model-B-PTFE-Teflon-Magnetic-...

I can make the motor virtually any speed/power
Later I will also mix various aqueous fluids for other general purpose chemistry, e.g. titration, nitration ......

pneumatician - 22-3-2015 at 08:56

antibacterial properties of spices
hi, checking the old post I see a link to a page with the antibacterial properties of spices. mozilla crash and bye bye. I found 1 but exist another with the page background of color red or orange. I put a lot of search terms and see all messages with spices, oils... and does not appear.
someone remenber the link?
TIA

[Edited on 22-3-2015 by pneumatician]

j_sum1 - 24-3-2015 at 15:58

Tin(II)chloride solubility.
References I have looked at indicate that SnCL2, both the anhydrous and dihydrate are very soluble in water. However, I have had a dog of a time getting them to dissolve. Wikipedia reports 83.9g/100mL @100°C. I have been unable to get a tenth of that. I have been working at RT however.
I have also tried reacting tin granules with HCl which has proved an extremely slow process but after a couple of days appears to be working.

I know that SnCl2 solutions are used in a lot of different applications. So what is the procedure? And why the discrepancy that I have noted between my results and the literature?

Volanschemia - 24-3-2015 at 21:44

Quote: Originally posted by Sulaiman

I am going to make ionic silver solution for medicinal use
and I want to stir the solution whilst the silver wires are electrolytically ionised
my questions are a bit vague so any explanations of general principles appreciated

1) which is best, a small stir bar rotating quickly or a large bar rotating slower?

2) which is the best shape?

2.1) ovoid, e.g http://www.ebay.co.uk/itm/2pcs-Model-A-Teflon-PTFE-Magnetic-...

2.2) rod, e.g. http://www.ebay.co.uk/itm/2pcs-Model-C-PTFE-Teflon-Magnetic-...

2.3) rod with ring, e.g http://www.ebay.co.uk/itm/2pcs-Model-B-PTFE-Teflon-Magnetic-...

I can make the motor virtually any speed/power
Later I will also mix various aqueous fluids for other general purpose chemistry, e.g. titration, nitration ......

I have always used the rod with ring type stirbar, and I think they give you the best stir power so to speak. As for small/fast or large/slow, I would choose large/slow because it gives you a better idea of how fast the liquid is rotating. Just my opinions.

Sulaiman - 25-3-2015 at 01:58

The Australian Scientist, Thanks for your info. I'll go for the rod with ring types.

i_sum1, when I reacted tin metal (scrapings from a solid ingot) with 36% HCl, I had to warm it to get a rapid reaction,
at r.t. it was very slow, as you found.
I think that the solubility of SnCl2 is very temperature dependant (a quick google gave no solubility curves)
but from the Wikipedia article on SnCl2;

"Tin(II) chloride can dissolve in less than its own mass of water without apparent decomposition, but as the solution is diluted hydrolysis occurs to form an insoluble basic salt:

SnCl2 (aq) + H2O (l) is in equilibrium with Sn(OH)Cl (s) + HCl (aq)
Therefore if clear solutions of tin(II) chloride are to be used, it must be dissolved in hydrochloric acid (typically of the same or greater molarity as the stannous chloride) to maintain the equilibrium towards the left-hand side (using Le Chatelier's principle). Solutions of SnCl2 are also unstable towards oxidation by the air:"

so it looks like an HCl acidic solution is required?

For my tin crystal growing https://www.youtube.com/watch?v=G1sq4hnrBgM I used excess HCl.
The tin crystals easily and quickly react with the solution they grew from,

Molecular Manipulations - 25-3-2015 at 09:55

j_sum1, Tin(II) chloride can dissolve in less than its own mass of water without apparent decomposition, but as the solution is diluted hydrolysis occurs to form an insoluble basic salt:

SnCl₂ (aq) + H₂O (l) <--> Sn(OH)Cl (s) + HCl (aq)
If you add a little hydrochloride acid it should shift the equilibrium to form less chlorohydroxide.

The Volatile Chemist - 26-3-2015 at 13:30

Something I've seen but can't find anywhere

When mixing H2O2 with pool water treatement pills it gave a red chemiluminiscence, am I right?

If I am, what is the reagent in the pills? (I think trichloroisocianuric acid, but can't remember)

Is it possibly Pyrogallic acid?

I'm sure it isn't. That's the only source of red chemoluminescence I know of that uses H₂O₂. I'd be interested in knowing what does if/when you find out.

Volanschemia - 26-3-2015 at 14:54

The pool treatment tablets are either Calcium Hypochlorite or Trichloroisocyanuric Acid, and I think most are the latter.

I seem to recall reading somewhere about this. I think the reaction of the acid and the Hydrogen Peroxide evolves Singlet Oxygen, which is the electronically excited state of O₂. When it comes down off it's high, so to speak, it emits red light.

It's a pretty interesting reaction which I will have to try out one day.

blargish - 26-3-2015 at 16:04

Almost certainly TCCA or NaDCCA. The red glow is given off by singlet oxygen formed from the decomposition of a peroxyhypochlorous acid species produced in the reaction.

If the effect was quite strong and not much foam was produced, it is most likely NaDCCA. TCCA produces the same effect, but with more foaming (from my experience) and a weaker glow.

Edit: I forgot that calcium hypochlorite produces this effect too. However, I think the effect is short-lived as compared to that from the other two chemicals.

[Edited on 27-3-2015 by blargish]

The Volatile Chemist - 27-3-2015 at 08:23

I'm pretty sure W. Oelen has some stuff on his site about it, maybe in the physics section.

badboy39560 - 5-4-2015 at 18:02

Please explain how to convert a sulfate into a hydrochloride?

smaerd - 6-4-2015 at 03:25

@Badboy39560 - It might help to be more specific. Different oxidations states of metals can make this more or less a different operation.

The Volatile Chemist - 10-4-2015 at 16:00

For a general answer, add a hypochlorite salt to a solution of sulfate, but there are so many exceptions, that that generalization isn't useful. What's the salt?

DraconicAcid - 10-4-2015 at 16:16

Quote: Originally posted by The Volatile Chemist

For a general answer, add a hypochlorite salt to a solution of sulfate, but there are so many exceptions, that that generalization isn't useful. What's the salt?

He said "hydrochloride", not "hypochlorite". Methinks it is the hydrochloride of some alkaloid he's trying to make.

Deprotonate it with some base, extract the neutral alkaloid into some organic solvent, then add hydrochloric acid.

The Volatile Chemist - 10-4-2015 at 17:26

No, no, see, you're suppose to give them two 'possible routes, one being the legitimate route, and one obviously not (i.e. heat with dilute bromine-water)

Texium - 26-4-2015 at 19:54

Is camphor readily soluble in xylenes? I searched, but wasn't able to find anything on the subject. Due to its high solubility in most common organic solvents, it seems like it would be, I'd just like to know for certain. Thanks.

Metacelsus - 27-4-2015 at 05:10

According to an old 1905 book, <i>Transactions of the Wisconsin Academy of Sciences, Arts, and Letters, Volume 15, Part 1</i> (from Google Books), "camphor is very soluble in toluene," so I assume it would be in xylenes.

from 36-37 question, you are correct

quantumcorespacealchemyst - 30-4-2015 at 15:39

Quote: Originally posted by The Volatile Chemist

Quote: Originally posted by quantumcorespacealchemyst

I get it all the time on things in my lab when evaporating HCl solutions. I think it's condensed HCl, or a salt of Cl^-, that is, if you have ammonia vapors nearby or something.

yes, you are correct.
i could not understand how it happened, not thinking about the blue solution of ammonia and copper salts i had been saving sealed in a yogurt container with saran wrap under the lid . i just checked it and most of the water is evaporated.

thanks

pneumatician - 1-5-2015 at 07:23

what salt from human urine is insoluble in water and pure etanol?

gdflp - 7-5-2015 at 09:29

what salt from human urine is insoluble in water and pure etanol?

None normally, hence the reason why urine, which is water based, is transparent.

ATX power supply question

Metacelsus - 7-5-2015 at 10:57

Can I use the +12V and +3.3V outputs on an ATX power supply to make an 8.7V output? Current would flow out of the 12V output and into the 3.3V output. (I have a Peltier cooler that I want to run that only accepts 8.3 to 8.8 volts, and draws 4 amps.) I don't want to ruin my ATX power supply by experimentation.

Edit: More research says that my power supply should be able to handle it.

Edit 2: Testing shows that the power supply just shuts off when I try to do it (probably some protection mechanism).

Edit 3: Even more research says that it is only possible if the 3.3V output is driving something that uses more than 4 amps. (The 3.3V output cannot sink current, only source it.) However, running the Peltier cooler off 5 volts works well enough for my purposes, I've found out.

[Edited on 7-5-2015 by Cheddite Cheese]
[Edited on 7-5-2015 by Cheddite Cheese]

[Edited on 8-5-2015 by Cheddite Cheese]

pneumatician - 9-5-2015 at 17:04

Quote: Originally posted by gdflp

what salt from human urine is insoluble in water and pure etanol?

None normally, hence the reason why urine, which is water based, is transparent.

so now I have one and one problem!

pneumatician - 9-5-2015 at 17:08

Quote: Originally posted by Cheddite Cheese

Can I use the +12V and +3.3V outputs on an ATX power supply to make an 8.7V output?

the fast, best and cheap is to buy a voltage reductor, this come in a integrated circuit like a transistor with 3 pins. perhaps not of 8,7v but 9v... but in reality electronics have a wide range of tolerance!!! :cool:

Ramium - 11-5-2015 at 22:34

question

is there any feasible way I can convert dichloroisocyanuric acid to sodium dichloroisocyanurate?

thanks

j_sum1 - 11-5-2015 at 23:13

Before I had any idea what I was doing, I managed to make a very small amount of potassium dichlorocyanurate from TCCA and KOH. I got some of the characteristic purple copper complex I was aiming for but extremely small amounts. It turns out this family of chemicals is a bit weird and there are multiple side reactions and alternate reactions. The dichloro milecule was merely an intermediate on the way to something else.
So, I guess it would be possible but might be tricky. woelen explained it rather well when I asked a similar question. He did some extensive work on these chemicals about 6-8 years ago. Good info is on the board but you will have to search. You might start by searching on me (oct last year) and see where that takes you.

Ramium - 12-5-2015 at 00:05

thanks. it sounds tricky so i'll probably just buy the sodium dichloroisocyanurate

Brain&Force - 16-5-2015 at 22:55

europium(III) 6,6,7,7,8,8,8-heptafluoro-2,2-methyl-3,5-octanedionate‏

Is this the correct systematic name for EuFOD?

smaerd - 17-5-2015 at 04:01

Brain&Force - I don't think so. Ligands have special naming conventions and that doesn't appear to follow them (to me). I never cared much for hugely esoteric molecules and the names of them though. Anyways, if you wanna get inundated, http://en.wikipedia.org/wiki/IUPAC_nomenclature_of_inorganic...

learningChem - 19-5-2015 at 11:16

I reacted 2g of sulphur and 1g of aluminium - thermite reaction - and got 1.9g of slag/product. I'm wondering what happened to the 1.1g that's missing. Maybe part of the sulphur is vaporizing instead of reacting with the aluminium?

diggafromdover - 19-5-2015 at 11:58

Some is vaporizing. Some may be combining with oxygen. Suggestion: Figure out the stoichiometry of the reaction. Likely aluminum is the limiting factor. Find out how many moles you started with, and how many moles of Al2S3 you would expect at 100% yield.

I would guess that any unreacted sulfur did volatilize and the slag is alumina.

learningChem - 19-5-2015 at 12:33

Thanks Digga!

I was assuming 2Al + 3S → Al2S3
roughly 2x27 + 3x32 ~ 54/96

so I used 1/3 Al 2/3 S - that's where the 1g/2g mix came from. Now I was thinking that if part of the S vaporizes, maybe I should use an excess of it?

blogfast25 - 19-5-2015 at 12:46

Quote: Originally posted by learningChem

Now I was thinking that if part of the S vaporizes, maybe I should use an excess of it?

That doesn't happen, in my experience. Stoichiometric ratio is good.

learningChem - 19-5-2015 at 16:52

Thanks Blogfast! Maybe I should add that I'm not using a container proper. I put the mix in a small piece of folded paper. That's probably not the best way to stop oxygen from interfering, I'm guesing.

blogfast25 - 19-5-2015 at 17:06

Quote: Originally posted by learningChem

Thanks Blogfast! Maybe I should add that I'm not using a container proper. I put the mix in a small piece of folded paper. That's probably not the best way to stop oxygen from interfering, I'm guesing.

Oxygen won't interfere.

Why not use an old coffee cup or egg cup or similar? The ceramic will not melt (but it will crack). Embed it in some dry sand. That's what I used to do. Works well.

[Edited on 20-5-2015 by blogfast25]

Brain&Force - 21-5-2015 at 22:56

Anyone have any tips for filing small pieces of hard distilled metal?

Also, will the use of a steel file impart any significant amount of steel contamination to the powder?

xfusion44 - 24-5-2015 at 10:42

Quote: Originally posted by Brain&Force

Anyone have any tips for filing small pieces of hard distilled metal?

Also, will the use of a steel file impart any significant amount of steel contamination to the powder?

File is made of very hard steel, so I wouldn't say that there will be much, if any contamination.

Texium - 25-5-2015 at 18:45

Would putting damp manganese sulfate in a desiccator with calcium chloride leave me with the anhydrous salt or the monohydrate?

j_sum1 - 25-5-2015 at 21:12

Question —ous and —ic
This feels like a real newbie one, but here goes.
Is there any rhyme or reason to the use of the suffixes ic and ous for ions and acids?

Consider
Ferric / ferrous: 3+ and 2+
Cupric / cuprous: 2+ and 1+
Stannic / stannous: 4+ and 2+
Plumbic / plumbous: 4+ and 2+
Mercuric / mercurous: 2+ and 1+
Hydrochloric acid / hypochlorous acid: 1- and 1+ oxidation states for Cl
Sulfuric acid / sulfurous acid: +6 and +4 oxidation states for S
Phosphoric acid / phosphorous acid: +5 and +3 oxidation stated for P

The only commonality I see is that the —ic is a more oxidised state. But then hydrochloric (and the other halides) buck the trend. And in the case of simple cations, how does the nomenclature stack up when there is more than two predominant oxidation states (as is true of many transition metals)?

While we are at it, is there a rule for —ate and —ite on anions: nitrate, nitrite etc?
And for bicarbonate, bisulfate, bisulfite, bitartrate etc... my simplistic reasoning merely modifies the anion by throwing an extra H+ at it. Is this the full story however?

Deathunter88 - 25-5-2015 at 21:21

Question —ous and —ic
This feels like a real newbie one, but here goes.
Is there any rhyme or reason to the use of the suffixes ic and ous for ions and acids?

Consider
Ferric / ferrous: 3+ and 2+
Cupric / cuprous: 2+ and 1+
Stannic / stannous: 4+ and 2+
Plumbic / plumbous: 4+ and 2+
Mercuric / mercurous: 2+ and 1+
Hydrochloric acid / hypochlorous acid: 1- and 1+ oxidation states for Cl
Sulfuric acid / sulfurous acid: +6 and +4 oxidation states for S
Phosphoric acid / phosphorous acid: +5 and +3 oxidation stated for P

The only commonality I see is that the —ic is a more oxidised state. But then hydrochloric (and the other halides) buck the trend. And in the case of simple cations, how does the nomenclature stack up when there is more than two predominant oxidation states (as is true of many transition metals)?

While we are at it, is there a rule for —ate and —ite on anions: nitrate, nitrite etc?
And for bicarbonate, bisulfate, bisulfite, bitartrate etc... my simplistic reasoning merely modifies the anion by throwing an extra H+ at it. Is this the full story however?

Here is a crash course video on nomenclature, might give some insights to your question: https://www.youtube.com/watch?v=U7wavimfNFE

byko3y - 25-5-2015 at 21:31

Perchloric/ Chloric / Chlorous / Hypochlorous acid: +7, +5, +3, +1 for Cl.
Nitric / Nitrous acid: +5 and +3 for N.
per-ate or di-ate / -ate / -ite / hypo-ite / -ide
per-ic or di-ic / -ic / -ous / hypo-ous
permanganate, manganate.
persulfate, sulfate, dithionate, sulfite, dithionite, silfide.

gdflp - 25-5-2015 at 21:42

Essentially, yes, ions and acids ending in -ic contain a central atom which is at a higher oxidation state than those containing the -ous suffix.

This naming convention only holds for oxoacids, as you have already noticed. The naming trend proceeds as follows, hypo---ous acids are less oxidized than -ous acids, which are less oxidized than -ic acids, which are less oxidized than a per---ic acids. I believe that the name for each ion is determined by assigning the per---ic acid to the highest oxidation state, but I am unsure. The naming convention is separate from acids which contain solely hydrogen and another element(or a pseudohalogen such as cyanide or azide) such as HCl, HBr, H₂S etc. In this case, to distinguish the acid from oxyacids, the prefix hydro- is used and all acids which have this prefix will have the suffix -ic.

In regards to ions, those with an -ic ending are simply in a higher oxidation state than those with an -ous ending. These are considered archaic names and have had the suffixes appended with no true rhyme or reason other than the common oxidation states of the element in question, these need to be memorized since there is no rule to determine them. Even though they're considered to be archaic, they are still used quite frequently in some cases(I prefer them since they're easier to type than IUPAC names such as iron(II) and iron(III)). And yes, anions with a bi- prefix simply have an additional
acidic hydrogen from a partial deprotonation of a multiprotic acid.

In regards to your -ate, -ite question, this corresponds directly to the name of the acid. A per---ic acid will be a per---ate, an -ic acid will be an -ate, an -ous acid will be an -ite, and a hypo---ous acid will be a hypo---ite. A hydro---ic acid will be an -ide.

[Edited on 5-26-2015 by gdflp]

j_sum1 - 25-5-2015 at 22:55

Thanks gdflp (and others.) That is a clear, concise answer. Nice to know I was half wsy there.
I will revisit the table I drew up a couple of days ago to help me memorise the dozen or more sulfur-contaning anions and their various names. (Metabisulfite and two persulfates to go thanks.) Armed with a somewhat systematic framework it might make more sense.

xfusion44 - 26-5-2015 at 04:05

Does potassium nitrite really explode at 537C? But it also decomposes at 440C?

PHILOU Zrealone - 26-5-2015 at 12:37

Quote: Originally posted by badboy39560

Please explain how to convert a sulfate into a hydrochloride?

Except with alchemical or nuclear transmutation...
If the sulfate is soluble simply add one equivalent of H2SO4 to get the hydrogenosulfate and then add BaCl2 solution...
BaSO4 will precipitate and leave you with the hydrochloride or chloride.

PHILOU Zrealone - 26-5-2015 at 12:41

Quote: Originally posted by zts16

Would putting damp manganese sulfate in a desiccator with calcium chloride leave me with the anhydrous salt or the monohydrate?

What Mn sulfate? II, III, IV

If it can be dehydrated to anhydrous in the open cold dry air, then it will aswel in a dessicator with CaCl2 what simply provide almost dry air in the closed container. Otherwise it will remain with cristalization water (mono or more hydrate).

PHILOU Zrealone - 26-5-2015 at 12:48

Quote: Originally posted by gdflp

what salt from human urine is insoluble in water and pure etanol?

None normally, hence the reason why urine, which is water based, is transparent.

so now I have one and one problem!

Uric acid and calcium salt.
Calcium oxalate.
Magnesium ammonium phosphate (struvite).

All 3 responsible of urinary calculi (urinary stones).

I see your problem

...and yes... very painfull...

Loptr - 26-5-2015 at 13:18

Does anyone have a reference procedure for a halide swap of an acyl chloride to form a acyl bromide?

Metacelsus - 26-5-2015 at 18:49

You mean like a Finkelstein reaction?

The reaction procedure will depend on the substrate, but it's essentially a matter of finding a solvent which dissolves a bromide but not a chloride.

Loptr - 27-5-2015 at 09:47

Quote: Originally posted by Cheddite Cheese

You mean like a Finkelstein reaction?

The reaction procedure will depend on the substrate, but it's essentially a matter of finding a solvent which dissolves a bromide but not a chloride.

That is actually the reaction I had in mind, but couldn't recall the name of it. I also didn't know if it had applicability with acyl halides, or acetyl chloride in my particular case.

Now that you have given me the name of the reaction, I will go do some reading on the mechanism.

Thank you!

j_sum1 - 31-5-2015 at 03:55

Chelated Magnesium?
There exists at my place some magnesium dietary supplements. (Why? Don't ask.) On the box it says that it is chelated. To what purpose? Mg2+ is pretty bioavailable and doesn' t have any real tricky chemistry to my knowledge. Why wrap it up in a complex?

DraconicAcid - 31-5-2015 at 08:37

It's chelated to make it sound more nutritious.

blogfast25 - 31-5-2015 at 09:00

It's chelated to make it sound more nutritious.

That won't work with chemophobes!

DraconicAcid - 31-5-2015 at 09:12

Quote: Originally posted by blogfast25

That won't work with chemophobes!

Yes, but the chemophobes aren't buying magnesium supplements, they're buying St. John's Wort and Oil of Harmony.

blogfast25 - 31-5-2015 at 09:21

Oil of Harmony.

Does it work? If so, can they put some in the water?

DraconicAcid - 31-5-2015 at 09:30

Actually, that's just an alternate name for olive oil. But the people buying it don't know that.

blogfast25 - 31-5-2015 at 09:59

Actually, that's just an alternate name for olive oil. But the people buying it don't know that.

For all their banging on about 'Big Bad Pharma', honesty is in very short supply on the 'alter-med' scene.

It might be goat's wool they're trying to pull over people's eyes but it's still wool.

The Volatile Chemist - 31-5-2015 at 14:19

Quote: Originally posted by blogfast25