Backyard Chemistry

College has completely killed my lab time

backyardchem — Mon, 24 Sep 2007 21:02:18 -0700

Well, I'm attending a wonderful university, but no matter how remarkable the chemistry classes may be in the future, nothing will replace the lab. This quarter I am in an accelerated general chemistry class, which will most likely be extremely boring given that I know all of the material.

There is also no lab that goes with this course- I won't get a lab until the spring quarter of organic. I'm even worried about organic being boring, but at least I've never learned that formally and I only read the parts that particularly interest me so there will be some new knowledge gained.

Since I will only be able to conduct experiments during vacations and the summer, it will be difficult for me to continously update this blog and my website of my experiments at Backyard Chemistry

Nevertheless, I hope to still do research and plan some experiments while I'm here during my free time. I now have access to tens of thousands of books at the chemistry and chemical engineering library here along with scores of scientific journals that might be of interest. I look forward to exploring the resources that are offered here.

I am a bit concerned since my passion for chemistry is derived almost entirely from my home lab. I hope that while being away from it I'll still be able to keep it alive. Participating in undergraduate research early may be of help here.

Sodium permanganate?

backyardchem — Fri, 07 Sep 2007 11:03:00 -0700

Sodium permanganate! Yes, it exists, but why doesn't anybody talk of it? I've always wondered this and I finally have myself a legitimate answer. And with this new knowledge, I may just be able to synthesize in my backyard.

Generally speaking, sodium salts are cheaper than their potassium analogs, but often they are not as useful because either they 1) are hygroscopic, 2) have less favorable solubility paramaters, or 3) don't crystallize easily. In the case of sodium permanganate, all three of these cases apply.

Still, it is odd that there is so little information available about sodium permanganate. A quick google search of "potassium permanganate" gives 1,080,000 results, while "sodium permanganate" only yields 41,100 results.

Potassium permanganate is most commonly produced by fusing molten potassium hydroxide and manganese dioxide together. An optional oxidizer such as potassium nitrate or potassium chlorate may be added to speed this process up. The fusion results in the production of the green solid, potassium manganate. Potassium manganate is then further oxidized to potassium permanganate either through a disproportionation reaction with a weak acid that is resistant to oxidation (usually carbonic acid is employed here) or via electrolysis.

Not an terribly complex affair! In fact, several preparative inorganic chemistry books describe the entire process in great length.

Here's where things get interesting. I had always assumed that sodium permanganate would be produced in an analogous manner; namely by reacting manganese dioxide with molten sodium hydroxide. And in this, some literature, albeit sparse, seemingly backed me up.

"Sodium manganate (Na₂MnO₄), prepared by fusion of a mixture of natural manganese dioxide and sodium hydroxide; green crystals, soluble in cold water, decomposed by hot water."

"Sodium permanganate, NaMnO₄, is obtained in a similar way to the potassium salt, and is distinguished from it by being deliquescent, and therefore, crystallizing with difficulty."

So off I went, performing nearly a dozen of experiments that all had the same underlying principle: fusing solid sodium hydroxide with manganese dioxide. I encountered problems right away. At first, I thought it was because I wasn't reaching the required temperatures, then because of carbon impurities and later zinc impurities in my manganese dioxide extracted from old batteries, and then because I discovered there was fumaric acid in my potassium chloride salt substitute! After purifying both my manganese dioxide and potassium chloride, I was still met with failure. I once did somehow manage to obtain the characteristic purple color of the permanganate ion by adding bleach to my sludge, but as you can see from the photograph of it to your right, it is extremely dilute. I had no chance to crystallize anything out and as soon as I began to boil the solution down it completely turned into the sad manganese dioxide.

This last experiment gave me hope at least, and I thought that maybe I just needed to increase my yields by heating this sucker at still higher temperatures and for longer periods of time. Out of frustration, I took an extended hiatus from the project and during this interim, I discovered what may be the key to this mystery!

"The price of sodium permanganate is about 5 to 8 times that of KMnO₄. This is mainly due to the fact that NaMnO₄ cannot be made in the same way as KMnO₄, because the oxidation of MnO₂ in a NaOH melt does not lead to the required Na₂MnO₄ (with hexavalent Mn) but only to Na₃MnO₄ with pentavalent Mn. The latter is very unstable in dilute NaOH solution (and therefore cannot be converted electrolytically to the desired NaMnO₄). Even if electrolytic oxidation were possible, there would still be the difficult problem of isolating the extremely soluble NaMnO₄ from the alkaline mother liquor."

Aha! This excerpt from Ullmann's Encyclopedia seems to answer all of my questions! You cannot produce sodium permanganate in the same way that you can produce potassium permanganate! Chemically, I still don't understand why. If anyone has an inkling as to why molten potassium hydroxide is a more potent oxidizer than molten sodium hydroxide in this case, please leave me a comment!

So basically, during all of my experiments, I was essentially performing the following reaction:

4MnO₂ + 12NaOH + O₂ --> 4Na₃MnO₄ + 6H₂O

And then when I extracted the mass with water, I got:

2Na₃MnO₄ + 2H₂O --> Na₂MnO₄ + MnO₂ + 4NaOH

On one lucky occasion this occurred:

3Na₂MnO₄ + 2H₂O --> 2NaMnO₄ + MnO₂ + 4NaOH

These equations seem to accurately describe what I witnessed in my experiments. It is probably true that in all cases, my yields were extremely low.

At this point, there was still one part of the puzzle that was missing. I had discovered these passages describing the industrial production of sodium permanganate:

"Sodium manganate, Na2MnO4, is formed when a mixture of equal parts of manganese dioxide and soda-saltpetre is heated for sixteen hours; the mass is then lixiviate with a small quantity of water and the solution cooled down, when the salt separates out in small crystals isomorphous with Glauber's salt, and having the composition Na2MnO4-10H2O. These dissolve in water with partial decomposition, yielding a green solution."

"For disinfecting purposes it is not necessary to employ the pure, well-crystallized salt [potassium permanangate], which is used in the laboratory, but a commercial article consisting of a mixture, more or less pure, of manganate and permanganate of sodium is used. The substance is obtained by mixing the caustic soda obtained from 1,500 kilos of soda-ash with 350 kilos of finely-divided manganese dioxide in a flat vessel, and heating this mixture for forty-eight hours to dull redness. The product is then lixiviated with water, and the solution either boiled to the requisite degree of strength or evaporated to dryness. If the manganate is to be completely converted into permanganate it is neutralized with sulfuric acid, the solution concentrated until Glauber's salt separates out, and these crystals are then removed and the liquid further evaporated."

What? These sources point to the direct production of sodium permanganate! I thought Ullmann's Encylopedia said that was impossible!

Well, it is impossible with sodium hydroxide, but not with good old soda, sodium carbonate at elevated temperatures (probably somewhere around 600C) and with prolonged heating. The set of reactions is probably:

4MnO₂ --> 2Mn₂O₃ + O₂

2Mn₂O₃ + 4Na₂CO₃ + 3O₂ --> 4Na₂MnO₄ + 4CO₂

This proposal has instilled hope in me once again for the backyard production of sodium permanganate (and from there potassium permanganate)! Once I finish building my furnace capable of withstanding such high temperatures, I am giving this one a try!

More Fun With Oxalic Acid: Synthesis of Sodium trioxalatoferrate(III) monohydrate

backyardchem — Sun, 02 Sep 2007 00:38:53 -0700

As always, you can view this experiment and many others at http://www.backyardchem.com

I decided to indulge in some coordination chemistry today! The oxalate ion is a strong ligand and bidentate to boot so it makes for some nice colorful salts when combined with various transition metals. I was reading about oxalic acid in a 19th century encyclopedia and came across a large section which was erroneously labeled as oxalate "double salts" (as opposed to "complex"). I spotted that oxalate ions coordinate with ferric ions to form emerald green salts. This sounded pretty cool and after some more research I had myself an experiment!

the trioxalatoferrate(III) ion

Before I outline exactly what I did, I should note that in leisure I did not choose the most straightforward synthesis route. I did not particularly care about the end yield and so this allowed me to witness more chemical reactions.

Here is the specific path I took to get to Na₃[Fe(C₂O₄)₃]-H₂O, a chemical with an unsurprisingly large number of synonyms which include sodium triethanedioatoferrate(III) monohydrate, sodium (tris)ethanedioatoferrate(III) monohydrate, sodium trioxalatoferrate(III) monohydrate, sodium (tris)oxalatoferrate(III) monohydrate, sodium ferric ethanedioate monohydrate, sodium iron(III) ethanedioate monohydrate, sodium iron(III) oxalate monohydrate, and sodium ferric oxalate monohydrate.

CaCl₂ + FeSO₄ --> CaSO₄ + FeCl₂

2NaHSO₄ + NaOCl + NaCl --> 2Na₂SO₄ + H₂O + Cl₂

2FeCl₂ + Cl₂ --> 2FeCl₃

FeCl₃ + 3NaHCO₃ --> 3NaCl + Fe(OH)₃ + 3CO₂

COOHCOOH + NaOH --> COOHCOONa + H₂O

Fe(OH)₃ + 3COOHCOONa --> Na₃[Fe(COOCOO)₃] + 3H₂O

There is little information available about the preparation of sodium trioxalatoferrate(III). Preparations of the analogous potassium salt, however, abound in literature. Most sources I read instruct to mix solutions of potassium bioxalate with ferric chloride. One preparation I read refluxed a mixture of potassium oxalate, barium oxalate, and ferric sulfate, taking advantage of the extremely low solubility of barium sulfate. Most probably, the much more common calcium oxalate cannot be substituted for the barium salt in this process because calcium sulfate is comparatively much more soluble.

I decided to stray away for both of these preparations and for the heck of it, produce sodium trioxalatoferrate(III) by reacting dissolving ferric hydroxide in a solution of sodium bioxalate, which as I far as I can tell, has no problems in theory.

Starting Reagents:

Calcium chloride- purchased at a hardware store for use in dehumidifiers

Ferrous sulfate- purchased in the gardening section of a local hardware store

Oxalic acid- purchased at a hardware store as �Wood Bleach�

Sodium bicarbonate- purchased at a grocery store as baking soda

Sodium bisulfate- purchased at a hardware store for decreasing the pH of pools

Sodium hypochlorite- purchased at a grocery store as bleach in a 6% solution

Sodium hydroxide- purchased as a drain cleaner at a hardware store

Metathesis Formation of Ferrous Chloride:

CaCl₂ + FeSO₄ --> CaSO_4(s) + FeCl₂

11.0g of calcium chloride and 20.0g of ferrous sulfate were each separately grounded up and mixed with enough water to completely dissolve them. The solutions were then mixed and a thick precipitate of calcium sulfate immediately formed. This was filtered and washed several times yielding a dilute yellow-green solution of ferrous chloride.

White precipitate of calcium sulfate

Filtered solution of yellow-green ferrous chloride

Chlorination of Ferrous Chloride to Ferric Chloride:

2NaHSO₄ + NaOCl + NaCl --> 2Na₂SO₄ + H₂O + Cl₂

2FeCl₂ + Cl₂ --> 2FeCl₃

In this procedure, chlorine gas oxidizes the ferrous ion to the ferric ion. I made chlorine case by slowly dripping 100mL of bleach from a separatory funnel into a solution containing 20g of sodium bisulfate. The chlorine generating flask was gently heated to limit chlorine�s solubility in water. It was then led into the solution of ferrous chloride and then subsequently led into a strong solution of sodium hydroxide to effectively neutralize excess chlorine gas.

The setup: bleach, sodium bisulfate, ferrous chloride, and sodium hydroxide from left to right

The result: a deep red solution of ferric chloride!

Formation of Ferric Hydroxide:

FeCl₃ + 3NaHCO₃ --> 3NaCl + Fe(OH)₃ + 3CO₂

Next I added excess sodium bicarbonate to the solution of ferric chloride. The solution turned quickly to orange and then slowly to more of a brown, and obviously lots of frothing and foaming ensued. The stoichiometry of this product is most likely not very precise and is more aptly described as the berthollide Fe₂O₃-nH₂O where n is ranges between 2 and 3. Regardless, the important thing here is that iron is the +3 oxidation state.

Filtering precipitated ferric hydroxide

Once filtered, I let the precipitate dry for a while in the sun, but I did not worry about being able to dry it enough to accurately mass it. I was content with a ballpark estimate of 5g.

Half Neutralization of Oxalic Acid:

COOHCOOH + NaOH --> COOHCOONa + H₂O

A solution of 5g sodium oxalate was slowly added to a 12.5g solution of oxalic acid to form a solution of sodium bioxalate. I am not sure what I was thinking here because I used rather valuable lye to neutralize the oxalic acid when I could have used the benevolent sodium bicarbonate. Oh well, hopefully I will eventually be manufacturing my own lye on a large scale from baking soda and slaked lime anyways.

�Ferric hydrate� and a solution of sodium bioxalate

Formation of the trioxalatoferrate(III) ion:

The hot sodium bioxalate solution was poured on the wet ferric hydroxide and was stirred, resulting in a lime green solution.

Fractional crystallization of sodium trioxalatoferrate(III) monohydrate:

The above 200mL solution was eventually boiled to about 40mL. I first cooled the solution to room temperature when it was at about 100mL and filtered out a white solid, which may have been excess oxalic acid and/or excess sodium bioxalate. At 40mL, I cooled the solution again and crystallized a light green solid of what is presumably sodium trioxalatoferrate(III) monohydrate. I did not bother to crystallize everything out and I was left with a medium-deep green solution.

Wet yield of sodium trioxalatoferrate(III)

As always, you can view this experiment and many others at http://www.backyardchem.com

Backyard Cemistry

backyardchem — Thu, 30 Aug 2007 21:43:11 -0700

Hello All,

Today I am launching my own chemistry website entitled "Backyard Chemistry".

Current amateur experiments listed are:

Enjoy!