May 28, 2022, 09:55:05 PM
Forum Rules: Read This Before Posting


Topic: Chemical Conservation Lab Design Problem  (Read 12593 times)

0 Members and 1 Guest are viewing this topic.

Offline Douglas

  • Regular Member
  • ***
  • Posts: 19
  • Mole Snacks: +0/-0
Chemical Conservation Lab Design Problem
« on: November 07, 2006, 10:48:14 PM »
Hello all -

I am trying to design a lab that will show my students some concepts of chemical conservation, percent yield, and the differences between metals and metal ions. To do so, I have them react copper with an oxidizing agent to yield copper ion. Then I react this copper ion solution with zinc metal in order to return copper ion to copper metal. In the final reaction, I react any remaining zinc metal with 6M HCl. I am having two problems with the lab.

1) Oxidizing Agent. Typically, this reaction is done with silver nitrate in order to produce silver metal and copper (II) nitrate. It reacts quickly enough and completely. However, the cost for silver nitrate (about $250/100g, and I need ~75g/lab) has lead me to look for another agent. I tried 6M Nitric acid, but the reaction has many problems, not the least of which is the NO gas that is released and converted to NO2 (it is also slow and acid tends to be left over). So my question is what soluble oxidizing agent would work quickly, completely, and do so relatively inexpensively and safely?

2) Assuming I make nice copper (II) ion in water - evidenced by a deep blue color - I add 20 mesh zinc metal (probably from the 80's, so there is likely a coating of zinc oxide). Very exothermically, copper metal appears (much of it stuck on the excess zinc metal) and the solution goes clear as zinc ion in water is clear. I filter the copper out and immediately wash it down with 6M HCl to react the remaining zinc. To speed drying, I rinse the product with acetone and dry it in a 60?C oven. The product, however, is much more red-brown than copper metal would be, suggesting incomplete reduction to copper (I) oxide. Why is this happening?

Thanks so much for your thoughts and ideas!
Douglas Weittenhiller

Offline Borek

  • Mr. pH
  • Administrator
  • Deity Member
  • *
  • Posts: 27090
  • Mole Snacks: +1759/-405
  • Gender: Male
  • I am known to be occasionally wrong.
    • Chembuddy
Re: Chemical Conservation Lab Design Problem
« Reply #1 on: November 08, 2006, 04:09:49 AM »
Check FeCl3 as oxidizing agent.
ChemBuddy chemical calculators - stoichiometry, pH, concentration, buffer preparation, titrations.info, pH-meter.info

Offline Douglas

  • Regular Member
  • ***
  • Posts: 19
  • Mole Snacks: +0/-0
Re: Chemical Conservation Lab Design Problem
« Reply #2 on: November 08, 2006, 12:45:22 PM »
Thought about FeCl3 and will try that (electron potential of 0.771 as compared to silver's 0.799 and nitric acid's 0.96). Also thought about Hg(NO3)2 as the electron potential is 0.891, but I don't like using mercury in the lab. Unsure of the proper safe handling and disposal.

Offline Borek

  • Mr. pH
  • Administrator
  • Deity Member
  • *
  • Posts: 27090
  • Mole Snacks: +1759/-405
  • Gender: Male
  • I am known to be occasionally wrong.
    • Chembuddy
Re: Chemical Conservation Lab Design Problem
« Reply #3 on: November 08, 2006, 01:38:34 PM »
Not sure about correct FeCl3 handling but it is relatively safe reagent and it is (or at least was several years ago) routinely used for copper etching from printed circuits so the technology is well known and shoul be easy to locate on the web.
ChemBuddy chemical calculators - stoichiometry, pH, concentration, buffer preparation, titrations.info, pH-meter.info

Offline Dan

  • Retired Staff
  • Sr. Member
  • *
  • Posts: 4716
  • Mole Snacks: +469/-72
  • Gender: Male
  • Organic Chemist
    • My research
Re: Chemical Conservation Lab Design Problem
« Reply #4 on: November 08, 2006, 01:51:25 PM »
Yeah, FeCl3 is a cheap, safe option and should work, but might be a little slow.

Hydrogen peroxide in acid should work too, and it's not that expensive or pants-wettingly dangerous.

Definately avoid mercury!
My research: Google Scholar and Researchgate

Offline Douglas

  • Regular Member
  • ***
  • Posts: 19
  • Mole Snacks: +0/-0
Re: Chemical Conservation Lab Design Problem
« Reply #5 on: November 08, 2006, 02:10:03 PM »
Yeah, FeCl3 is a cheap, safe option and should work, but might be a little slow.

Hydrogen peroxide in acid should work too, and it's not that expensive or pants-wettingly dangerous.

Definately avoid mercury!

What do you suggested for acidified hydrogen peroxide? Say for every 10mL 30% H2O2, add 2-3 drops of conc. sulfuric acid when dissolving ~1-2g of copper (ignoring stoich)?

Thanks

Offline Dan

  • Retired Staff
  • Sr. Member
  • *
  • Posts: 4716
  • Mole Snacks: +469/-72
  • Gender: Male
  • Organic Chemist
    • My research
Re: Chemical Conservation Lab Design Problem
« Reply #6 on: November 08, 2006, 02:48:55 PM »
Do a calculation using stoichiometry. I think you will need more acid.
My research: Google Scholar and Researchgate

Offline Douglas

  • Regular Member
  • ***
  • Posts: 19
  • Mole Snacks: +0/-0
Re: Chemical Conservation Lab Design Problem
« Reply #7 on: November 08, 2006, 03:09:28 PM »
Ok, I'll try both and let you know the results. Any ideas on issue #2?

Offline Dan

  • Retired Staff
  • Sr. Member
  • *
  • Posts: 4716
  • Mole Snacks: +469/-72
  • Gender: Male
  • Organic Chemist
    • My research
Re: Chemical Conservation Lab Design Problem
« Reply #8 on: November 08, 2006, 03:26:34 PM »
Hmmm. Not sure. This may be a problem with the old zinc. Try using new zinc powder, or even just sanding off the zinc oxide layer.

Does the colour appear unusual before you heat it in the oven?
« Last Edit: November 08, 2006, 06:18:27 PM by Dan »
My research: Google Scholar and Researchgate

Offline hmx9123

  • Retired Staff
  • Full Member
  • *
  • Posts: 897
  • Mole Snacks: +59/-18
Re: Chemical Conservation Lab Design Problem
« Reply #9 on: November 08, 2006, 06:32:11 PM »
I would avoid using hydrogen peroxide and sulfuric acid together; it leads to peroxydisulfuric acid, which is pretty bad if you have any organics around.  The pyro world likes to add hydrogen peroxide to hydrochloric acid.  I haven't tried it myself, but it is supposed to allow the HCl to dissolve the copper.  That would be one way of getting the copper in solution.

As for the zinc, you should use as small a particle size as possible for the experiment, otherwise the surface will be passivated with the copper and the zinc inside won't be able to react.

Copper can oxidize when heated in an aerobic atmosphere; it depends on heat, etc., but perhaps the fact that you're using non-degassed solvents and heating while it's wet, plus the thinness of the layer make a difference.  You also may just be looking at a reddish hue to the copper since it hasn't been polished.

Offline Douglas

  • Regular Member
  • ***
  • Posts: 19
  • Mole Snacks: +0/-0
Re: Chemical Conservation Lab Design Problem
« Reply #10 on: November 08, 2006, 07:26:02 PM »
I would avoid using hydrogen peroxide and sulfuric acid together; it leads to peroxydisulfuric acid, which is pretty bad if you have any organics around.  The pyro world likes to add hydrogen peroxide to hydrochloric acid.  I haven't tried it myself, but it is supposed to allow the HCl to dissolve the copper.  That would be one way of getting the copper in solution.

As for the zinc, you should use as small a particle size as possible for the experiment, otherwise the surface will be passivated with the copper and the zinc inside won't be able to react.

Copper can oxidize when heated in an aerobic atmosphere; it depends on heat, etc., but perhaps the fact that you're using non-degassed solvents and heating while it's wet, plus the thinness of the layer make a difference.  You also may just be looking at a reddish hue to the copper since it hasn't been polished.

The hydrogen peroxide and sulfuric acid produced poor result, with a green-oxided copper in solution and a black coating on the copper metal, presumably copper (II) oxide. I'll try the hydrochloric acid, but I still think the hydrogen peroxide is too powerful of an oxidizer (electron potential 1.777).

I really like the idea of ferric chloride just for the simplicity of an single-exchange reaction. Since I know how nitrates in concentration behave with copper, I will also try ferric nitrate.

I do have the passivation problem you describe, so I will look into newer/smaller zinc metal. To control the oxidation, I will try the reaction on ice with slow additions of zinc metal. Whether it's copper or copper (I) oxide, I'll characterize it with ammonia (should go clear if copper (I) oxide and purple if copper metal) and heating (if copper (I) oxide, it should turn the black color of copper (II) oxide.)

Thanks for the ideas! I'll let you know the results.

Offline Borek

  • Mr. pH
  • Administrator
  • Deity Member
  • *
  • Posts: 27090
  • Mole Snacks: +1759/-405
  • Gender: Male
  • I am known to be occasionally wrong.
    • Chembuddy
Re: Chemical Conservation Lab Design Problem
« Reply #11 on: November 08, 2006, 07:40:04 PM »
Since I know how nitrates in concentration behave with copper, I will also try ferric nitrate.

Standard procedure as I know it calls for ferric chloride, supposedly for a purpose, although I don't know details. Note that copper oxidation is done in low pH in the hydrochloric acid, and thus in the presence of chloride ions that are complexing agent.
ChemBuddy chemical calculators - stoichiometry, pH, concentration, buffer preparation, titrations.info, pH-meter.info

Offline Douglas

  • Regular Member
  • ***
  • Posts: 19
  • Mole Snacks: +0/-0
Re: Chemical Conservation Lab Design Problem
« Reply #12 on: November 08, 2006, 11:16:14 PM »
Upon attempting hydrogen peroxide and drops of conc. HCl, the reaction did not produce the copper (II) oxide found with drops of conc. sulfuric acid. The reaction bubbled like crazy upon the addition of the acid, and quickly the solution turned green ([CuCl4]2-, perhaps). The copper turned brilliant red and then black-red, suggesting some form of oxidation occurring. The reaction also took an long time to complete (>1h), and thus I am remaining hopeful for ferric chloride when it arrives next week.

Thanks again for all the ideas. I also purchase zinc powder instead of 20 mesh, and will try the reaction of copper ion with zinc over ice.

Douglas Weittenhiller

Offline hmx9123

  • Retired Staff
  • Full Member
  • *
  • Posts: 897
  • Mole Snacks: +59/-18
Re: Chemical Conservation Lab Design Problem
« Reply #13 on: November 10, 2006, 04:47:22 AM »
Glad to hear things are working OK.  Sounds like you probably have copper as a layer on the outside.  As for the zinc, 20 mesh is HUGE (at least it seems that way to me, being a pyro, where -325 mesh is 'fine').  Zinc powder should suit your purposes much better.  Good luck with your experimenting.

Sponsored Links