Sciencemadness Discussion Board

A few of questions about HCl(aq.)

HydroCarbon - 27-7-2008 at 20:34

I just started doing chemistry as a hobby and have been doing experiments with HCl from the hardware store. While doing these I have come up with a few questions. I would appreciate it greatly if someone had some answers.

1) The bottle of muratic acid I have is very impure, yellow in color. Is it possible to purify it through a set-up like this:

http://img300.imageshack.us/img300/7780/tutlesgw7.jpg
(sorry for shitty drawing)

Would the HCl attack the rubber tubing? Could the rubber stoppers impart impurties? Will the NaOH neutralize any excess gas well?

2) What is the best kind of container to store the HCl in? As of now all I have is erlenmeyer flasks with rubber stoppers, glass beverage bottles, and plastic bottles. What is a good general concentration to standardize the solution at for both ease of storage and practical use?

3) And finally when I worked with the 10M hardware solution I have I noticed a fair amount of HCl coming out of the solution. Is it possible that this gas leaving the solution weakens the concentration and can therefore mess up predicted titration results? i.e. causing less NaOH to be needed than predicted.

[Edited on 28-7-2008 by HydroCarbon]

ShadowWarrior4444 - 27-7-2008 at 21:21

The distillation apparatus you posted will likely work, unless the impurities are volatile, such as Cl2.

HCl gas is quite corrosive to many things and should it escape it will decrease the concentration. It is usually best to store conc. HCl in a container with a threaded Teflon cap. Yellow (gas line) Teflon tape is very effective for sealing threads against HCl leak. Most laboratory grade plastics should be suitable for storing it, (LDPE, HDPE, PP, etc) but you should look at the chemical resistance ratings for them online.

The most important aspect is that you seal the cap from gas leak.

chloric1 - 28-7-2008 at 02:17

hydrocarbon-don't really know where you are located but this might be some good advice. Since you are just starting out, don't ever buy the same brand of HC twice. When you run out, try a different brand of muriatic acid. Most of the muriatic acid I buy is suffiently white for most applications.

Some of your chemistry may help to remove impurites. An example is if you use pottery grade lithium carbonate to make lithium chloride, add a few grams extra carbonate after all acid is neutralized, it will fall to the bottom of the container and leave a crystal clear lithium chloride solution. This is only an example. Study the products of your reactions for solubility trens etc. to see how you can benefit.

kclo4 - 28-7-2008 at 02:44

Quote:
Originally posted by HydroCarbon
I just started doing chemistry as a hobby and have been doing experiments with HCl from the hardware store. While doing these I have come up with a few questions. I would appreciate it greatly if someone had some answers.

1) The bottle of muratic acid I have is very impure, yellow in color. Is it possible to purify it through a set-up like this:

http://img300.imageshack.us/img300/7780/tutlesgw7.jpg
(sorry for shitty drawing)

Glad your interested in such an exciting hobby! Hope to see you read and post more on the board!

It looks like it will work, but just be careful not to allow the HCl gas to suck back the NaOH, if it does, and the concentrated of the HCl is high enough, it could be disastrous. Adding a strong base to a strong acid can produce a lot of heat. I would probably use a dilute solution of NaOH, since hopefully you won't be losing to much HCl anyways.
But for what you are doing with it, is it really necessary to purify it?
Perhaps the impurity could be reacted to turn it back into HCl, pr precipitate, but this would require you knowing what the impurity was, and that would be difficult to figure out.
Quote:

Would the HCl attack the rubber tubing? Could the rubber stoppers impart impurties? Will the NaOH neutralize any excess gas well?

I highly doubt would react with the rubber tubbing, but in order to say we would have to know what type of tubing it was. Could always give a test run or try out a small piece of the tubing.
It might add a few impurities, and IF you really need your HCl that pure, you could perhaps try washing it with HCl or water before hand.
Quote:

2) What is the best kind of container to store the HCl in? As of now all I have is Erlenmeyer flasks with rubber stoppers, glass beverage bottles, and plastic bottles. What is a good general concentration to standardize the solution at for both ease of storage and practical use?

I personally don't know what it would react with, the HCl i got seems to have came in a slightly thicker then normal milk jug, and it is also yellow. I thought the coloring was a dye.
Quote:

3) And finally when I worked with the 10M hardware solution I have I noticed a fair amount of HCl coming out of the solution. Is it possible that this gas leaving the solution weakens the concentration and can therefore mess up predicted titration results? i.e. causing less NaOH to be needed than predicted.


Yes it could, but it probably would take quite some time, so assuming you keep the lid sealed, it should be basically the same. Gases have a much greater volume then when they are in a solution, think about the Ammonia fountain demonstration.

:D

Klute - 28-7-2008 at 05:54

The yellow impurity is most likely Fe3+ contamination, but of very slight contamination. In a number a cases, this isn't problematic at all. If you try out coordiantion chemistry, it will have to be removed though....

The rubber tubing will not last long, and "dry" up to a darker, harder, cracked tubing. One way iof retarding this to to thoroughly wash the tubings with tap water when done, and leave to dry suspended. But they will not last for ever!

As KClO4 mentionned, suck back of the NaOh solution is most probable, leave the surface of the solution just at the end of the tubing, so that if any suck back occurs, the level of water will get lower than the tube. I think a 5-10% NaOH solution should be suiteable.

Good luck!

not_important - 28-7-2008 at 06:03

Quote:
Originally posted by HydroCarbon

3) And finally when I worked with the 10M hardware solution I have I noticed a fair amount of HCl coming out of the solution. Is it possible that this gas leaving the solution weakens the concentration and can therefore mess up predicted titration results? i.e. causing less NaOH to be needed than predicted.


A concentration above the constant boiling acid, ~6N 20.2%, will readily lose HCl. I'm not sure which way you are trying to go in your titrations, but alkali hydroxides absorb CO2 readily and the concentration of technical grade hydrochloric acid is rather approximate. One way to get known strength HCL aq is to distill it in a proper glass distillation apparatus, collecting a portion that boils at a fixed temperature while noting the atmospheric pressure, then using tables such as that attached below, calculate the actual concentration of HCl in the distillate.

Traces of iron are generally the cause of the yellow tinge in technical grade hydrochloric acid.

constant_boiling_HCl_ref.png - 177kB

Klute - 28-7-2008 at 06:21

and if you want to titrate the obtained solution, dilute them to 0.1-0.5 N, with conc. solution you get very approximate results because of what Not_Important mentionned (losses by gas, absorption of CO2, a single drop out of reach will cause a high experimental error), etc

not_important - 28-7-2008 at 07:07

Ah, very good point, Klute.

Investing in a few pieces of true volumetric glassware would be wise if you are looking to do titrations.

HydroCarbon - 31-7-2008 at 17:58

Thanks very much for the responses everyone.