Sciencemadness Discussion Board - Is concentrated HCl yellow?

Is concentrated HCl yellow?

LargeV - 29-11-2015 at 13:06

I left out 10 ml of 10% HCl out in a beaker for a week and when I went to look at it, it was an extremely viscous yellow gook. Is this concentrated HCl? Thanks

aga - 29-11-2015 at 13:07

Short answer : No.

HCl_(aq) is clear, and moves/looks like water

LargeV - 29-11-2015 at 13:11

Thanks aga! Forgot to mention that it does turn CuSO4 green and eats through aluminium.

gdflp - 29-11-2015 at 13:14

I'm guessing this was acid from a hardware store. Most of the hydrochloric acid evaporated, so that's likely a small amount of ~20%HCl, in which the iron impurities have been concentrated, causing it to take on a yellow color.

LargeV - 29-11-2015 at 13:17

It was! Thank you gdflp.

aga - 29-11-2015 at 13:31

Muriatic acid is the common (old) name for Hydrochloric acid.

Where i live they sell 20w% HCl in the supermarkets, and it has no Iron that i can detect using potassium hexaferrocyanate.

The agricultural Nitric & Phosphoric acids are highly contaminated with Fe.

Edit:

You can reduce the w% by half, yet produce 100% pure HCl solution.

https://www.youtube.com/watch?v=jv1Ms6Subg4

In the end, there's no way to make the impure HCl stuff purer and stronger at the same time - for 37w% (the max) you need to pump pure, dry HCl gas into water (or some lower w% HCl solution).

There's a way to do that in Prepublications, len1 as i recall ...

[Edited on 29-11-2015 by aga]

LargeV - 29-11-2015 at 13:32

Isn't mixing ferrocyanide with acids a nono? I read somewhere that it generates HCN.

gdflp - 29-11-2015 at 13:38

Quote: Originally posted by aga

potassium hexaferrocyanate

I believe that you are referring to hexacyanoferrate, as hexaferrocyanate doesn't exist.

@LargeV

With 20% HCl, the danger of HCN is minimal. However, you can dissolve some sodium carbonate in water and add a small amount of sodium/potassium hexacyanoferrate. Adding suspect acid to this solution will still produce the characteristic color in the presence of iron, but has no danger of producing volatile cyanide.

LargeV - 29-11-2015 at 13:45

Thank you! I have a dropper of the hexacyanoferrate just for these metal identifying reasons

aga - 29-11-2015 at 13:52

Quote: Originally posted by gdflp

Quote: Originally posted by aga

potassium hexaferrocyanate

I believe that you are referring to hexacyanoferrate, as hexaferrocyanate doesn't exist.

Ah yes. The fact that it doesn't exist means that i got it backwards, again.

(@LargeV : listen to what gdflp says, always)

byko3y - 29-11-2015 at 19:34

I have a distilled HCl solution, and I can say it has a very slight yet noticable yellow-green color, despite the fact it has no iron contaminations.
UPD: lol, woelen actually made a picture of his HCl acid for wikipedia https://en.wikipedia.org/wiki/File:Hydrochloric_acid_30_perc... Mine one looks the same.

[Edited on 30-11-2015 by byko3y]

feacetech - 30-11-2015 at 18:51

I was thought it was Cl/Hypo contamination (I based this on absolutely nothing apart from colour) or some form of corrosion inhibitor although inhibited HCl is often brown (im not confused with sulphuric I use to use inhibited HCl for descaling industrial equipment)

byko3y - 30-11-2015 at 23:04

It's not chlorine, because it has nowhere to come from: there was no free chlorine as well as no oxidizers in the source and in the receiving flask.
Hypochlorite can't exist and this low ph.

xfusion44 - 1-12-2015 at 00:21

Quote: Originally posted by LargeV

Thanks aga! Forgot to mention that it does turn CuSO4 green and eats through aluminium.

"It turns CuSO4 green"

Doesn't that make sulfuric acid? I mean, when you mix HCl and CuSO4?

aga - 1-12-2015 at 04:54

Quote: Originally posted by xfusion44

Doesn't that make sulfuric acid? I mean, when you mix HCl and CuSO4?

I guess it does.

No idea how you'd separate the CuCl₂ from the solution though.

More effective is to electrolyse the CuSO₄ solution to 'plate' out the copper.

You can tell when it's completed by the lack of colour.

[Edited on 1-12-2015 by aga]

woelen - 1-12-2015 at 05:09

Pure HCl is totally colorless, but to my experience, even lab grade HCl always has a very faint green color. This color is not observable in a small test tube, but if you have a liter or so of the liquid in a clear colorless bottle, then it can be observed.

A similar thing is true of HBr. This liquid always appears somewhat yellow/brown, even high grade quality HBr.

I can imagine that the highly concentrated pure liquids contain trace amounts of free halogen or trace amounts of hypohalous acids, just enough to give a visible color. It might be that this is due to the presence of dissolved oxygen. This is speculation from my side, however, I never read about this, nor did I investigate.

Detonationology - 1-12-2015 at 08:37

Quote: Originally posted by woelen

Lab grade HCl always has a very faint green color.

I found a picture of fairly concentrated HCl to support this.

Amos - 1-12-2015 at 10:26

Quote: Originally posted by xfusion44

Quote: Originally posted by LargeV

Thanks aga! Forgot to mention that it does turn CuSO4 green and eats through aluminium.

"It turns CuSO4 green"

Doesn't that make sulfuric acid? I mean, when you mix HCl and CuSO4?

No, it doesn't "make" sulfuric acid. You just get a mixed solution of the ions of both component compounds. The green color arises from the formation of [CuCl₄]^2-, also called tetrachlorocuprate. This is just sort of the default state of any reasonably concentrated solution containing both Cu²⁺ and an excess of Cl^- ions.

AJKOER - 3-12-2015 at 12:13

Quote: Originally posted by LargeV

I left out 10 ml of 10% HCl out in a beaker for a week and when I went to look at it, it was an extremely viscous yellow gook. Is this concentrated HCl? Thanks

So what could occur on leaving an impure HCl (I will assume transition metal contaminated) in an open vessel in the presence of oxygen and possibly sunlight for a week?

Well, let's start with some possible Fenton based reactions creating the hydroxyl radicals, .OH and the superoxide anion, .O2- . As a reference, see, for example, "Generation of Hydroxyl Radicals from Dissolved Transition Metals in Surrogate Lung Fluid Solutions" by Edgar Vidrio, et al at http://www.ncbi.nlm.nih.gov/pmc/articles/PMC2626252/ . Cited reactions :

Cu(l)/Fe(II) + O2(aq) --) Cu(ll)/Fe(III) + .O2-

As an alternate reference for the above reaction (which I have personally performed on Cuprous citrate using an air pump from an old fish tank), see for example, https://books.google.com/books?id=WjReuSXxl4YC&pg=PA17&a...

The reaction chain continues as:

Cu(l)/Fe(II) + .O2- +2 H+ --) Cu(ll)/Fe(III) + HOOH

Cu(l)/Fe(II) + HOOH --) Cu(ll)/Fe(III) + .OH + OH-

Net of the last three reactions:

3 Cu(l)/Fe(II) + O2(aq) +2 H+ --) 3 Cu(ll)/Fe(III) + .OH + OH-

And, in the presence of sunlight (or a reductant like Citric or Ascorbic acid), a cyclic reaction could ensue in the case of sunlight:

Cu(ll)/Fe(lll) (aq) + hv → Cu(l)/Fe(ll) (aq) + HO• + H+

Furthermore, in the presence of the hydroxide radical and a halide ion (X-), some additional possible reactions include:

.OH + X− → HOX.− (an unstable intermediate)

HOX.− → X. + OH-

X. + X.→ X2

X2 + H2O = HX + HOX

Reference, see for example: https://www.google.com/url?sa=t&source=web&rct=j&...

In the case of X being Cl, any formed HOCl could also be consumed in yet another, Fenton-type reaction, which is seen in biological systems. As a reference, see "Fenton chemistry in biology and medicine*" by Josef Prousek available at https://www.google.com/url?sa=t&source=web&rct=j&... . To quote reaction 15 on page 2330:

"For Fe(II) and Cu(I), this situation can be generally depicted as follows [20,39],

Fe2+/Cu+ + HOX → Fe3+/Cu2+ + .OH + X- (15)

where X = Cl, ONO, and SCN. "

The net process so far thus indicates a possible consumption of water, the formation of the hydroxyl anion, the hydroxyl radical and the chlorine radical and chlorine water (with the associated formation of chlorides/hydroxyl chlorides), all of which would be consistent with the observed results of leaving transition metal contaminated HCl in an open vessel in the presence of oxygen and possibly sunlight.

[Edit] So, the answer to your question is that it may contain some HCl only if the all those formed OH- anions left some, otherwise:

OH- + HCl = H2O + Cl-

which could then react with the chlorine radical (created above per the action of the hydroxyl radical) as follows, forming the so called pseudo-halide radical anion (see, for example, discussion, equations and Table I on page 392 at
https://books.google.com/books?id=cKjhBwAAQBAJ&pg=PA392&... ):

Cl- + Cl. = .Cl2-

which may be present in solution with chloride salts and such. I would not be surprised if the resulting mix was reactive (at least as long as the .OH generation was kept active), but not largely due to the concentration of the HCl.

[Edited on 3-12-2015 by AJKOER]