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Backyard Chemistry

December 30, 2009
1:30
It's not the end of backyard chemistry
» ‎ Backyard Chemistry

Well, if you have visited my site at www.backyardchem.com recently, you may have noticed that I am far from my heydays of amateur experimentation which occurred in 2007.
This is largely because as a college student, I have not had time to perform experiments. More importantly, I am away from my home chemistry lab except during vacations. However, I still think about how to make chemicals from household materials quite often (I still read ingredients on household products compulsively!), so I plan to continue adding to this website. I have decided to include this new ideas and tidbits page in order to give myself more flexible space to add new content. My chemistry classes at school and my years of research in a professional chemistry laboratory have vastly improved the breadth and depth of my understanding of chemistry. As a result, I am now able to tackle the problems facing amateur chemists from a new angle. I plan to share some of my new ideas here on this page without necessarily doing much more experimentation. I will include both amateur chemistry project ideas and other interesting chemistry experiments and knowledge.
August 30, 2009
0:51
How to Survive Organic Chemistry Lab
» ‎ Backyard Chemistry

As a chemistry major, I have had experience both taking and teaching undergraduate organic chemistry lab courses. I have noticed what it takes for a student to be successful in what can be a very stressful lab environment so I thought I would share my knowledge with you.

How to Survive Organic Chemistry Lab
Backyard Chemistry

As a chemistry major, I have had experience both taking and teaching undergraduate organic chemistry lab courses. I have noticed what it takes for a student to be successful in what can be a very stressful lab environment so I thought I would share my knowledge with you.

First let's start out with some common anxieties that students have to deal with in chemistry lab:

1) Finishing the experiment on time. Depending on the day, this can be a difficult task. I will give you many tips and tricks as to how work swiftly but carefully.

2) Getting a good yield. This is a major concern of students who are overly anal. There are some common practices that can increase your yield by a couple of percentage points, but most of the yield is determined by the nature of the reaction. More on this later.

3) Spilling or breaking equipment. I'll teach you only to be super careful when you need to be super careful. This will reduce much of your stress in the chemistry lab and save you time as well.

4) The reaction not working. Some students consistently think that their experiment is going to be the only one that is not going to work. The first thing you should know is that all the labs are designed and optimized to work as well as they possibly can. But nevertheless, I can give you some advice that will help you decrease your chances of complete failure.

5) Writing a good lab report. There are some basic things in an organic chemistry lab report that most students always seem to forget.

Ok so now I'll try my best to address each of these concerns specifically.

1) Finishing the experiment on time

a. Read and digest the contents of the pre-lab. Read the pre-lab the night before you do the experiment and organize in your head what different steps you will be doing. For your very first lab, you will most likely have to think about the steps in greater detail, but once you gain familiarity in various techniques, you will be able to think about your upcoming experiment in simpler terms. It is very important that when you read the experiment you think about what chemicals and what glassware and other supplies you will need. This will save you tons of time as you will not have to be constantly referring to the lab instruction manual during the lab period. During the class, you should mostly only be going back to the instruction manual to look at specific numbers that are given (mass of reactants, volume of solvents, temperature of reaction, etc.).

b. Minimize the number of trips you make. In organic chemistry lab, much of what separates the fast students from the slow students is the rate at which the experiment is set-up. In many experiments such as distillation, setting up the glassware and all of the proper equipment can take an hour or longer. There then might be a period of downtime where all of the students have to wait for a reaction to occur (let's say another hour). So it is super important to get the reaction started as soon as possible. Therefore, setting up your glassware in a timely fashion is key. In order to do this, when you go to get your supplies, carry more than one thing at once! This might seem like a silly piece of advice, but you would be surprised to know how many students I see carrying back a single Erlynmeyer flask to their hood and nothing else! When you race to the glassware storage areas to pick up your flask, think about everything else that you might need to go with it. Graduated cylinder to measure out what is going into the flask? If yes, well then you probably need a pipette and a pipette bulb to dispense liquid into the graduated cylinder. Buchner funnel to filter into the flask? If yes, well then you need filter paper to go on top of your funnel.

c. Make use of downtime. Often times, there is a lot of time when you will be waiting for a reaction to occur. These can be 10 minutes blocks of time or even several hours. This is a good time to catch up and write down things in your lab notebook. Writing everything in your notebook as you go can slow you down substantially. Instead look for blocks of time that you can use to write a lot of information down at once. If you still have extra time after that, don't just lollygag around! Plan out what you will need for the next steps. You can use this time to grab the necessary glassware and even to measure out the appropriate amount of reagents and solvents that you will need next.

d. Estimate when you can. This is one thing that I make sure to emphasize to my students. Think about if you really need an exact amount of a certain chemical even though the instructions may say so. A prime example deals with the use of drying agents. Let's say the instruction manual tells you to dry the organic layer with 2.50g anhydrous sodium sulfate. Now, the slow student would spend 5 minutes at the scale weighing out exactly 2.50g of sodium sulfate or as close as he or she could manage. The medium-paced student would realize that the amount of drying agent you use does not have to exact and would go and weigh out anything from 2-3g. The fast student, however, would know that it is entirely unnecessary to weigh out the amount of sodium sulfate and that once the drying agent stops clumping, enough has been added.

You also can estimate when you are adding solvents. For example if a procedure calls for 50mL of ethanol to run a reaction in, it's not going to do any harm if you only have 49mL. Estimation is okay as long as you write down exactly what you did in your lab notebook.

e. Multi-task when you are comfortable. At first, you will be wanting to give 100% of your attention to everything you are doing because of a lack of experience as to what to expect. However, once you gain more knowledge, you will be able to take on more than one task at the same time. This is especially important when you are running a column. Here, it will save you tons of time if you run your TLC analysis while you are pushing solvents through your column.

f. Store frequently used solvents. If you anticipate that you'll need lots of different solvents in your lab. There's no harm in storing a good size of these solvents in your hood. Use Erlynmeyer flasks and label them quickly. If you put the solvents in a beaker, then they will evaporate more rapidly.

g. Clean up at the end, not as you go. You should inquire as to whether or not you will be docked if you are messy when working in your fume hood. If it's alright, then it is much more efficient to store all of the glassware and equipment that you are done using in the corners of your hood where they will not get in the way. Then, after you are done for the day, you can clean everything up at once. This saves a lot of time since the instinct is to clean as you go. (Cleaning, however, is something good to do when you have downtime.)

2) Getting a good yield.

Whatever you do, don't fake your yield. I've known students to say that there reaction yielded a product of 80% when the highest I, as a teaching assistant, had ever gotten was 40%! So definitely, don't artificially inflate your yield. If you think your yield is low, then explain in as detailed terms as you can why you think so. Hopefully, your lab performance grade is not based upon yield, but about the nature of your report and your explanations.

Good habits that will increase your chances of getting a better yield include making sure that your glassware is clean especially before adding the key reagents. Another important thing is to make sure that your reaction is at the requisite temperature during the entire duration of heating.

You also can make sure that you don't loose too much of your product to physical processes by washing anything your product has touched with plenty of solvent and then evaporated it down. There is no harm in doing this but make sure you don't go overboard on how much solvent you use or else you will be evaporating for hours.

3) Spilling or breaking equipment.

Unfortunately, this is a sad reality of the game. However, there are some precautions you can take. If there is anything at all that seems jammed or stuck, then you should ask your teacher what to do because chances are you probably are using the wrong piece. In addition, if you are using any type of ground-glass joints, make sure to use grease, because this will prevent the pieces from sticking together. Don't think that this step is unnecessary. Also, if you have a complicated glassware set-up going on, make sure that everything is well clamped and that you are not moving anything that is not well-supported. As a general rule, don't force things in too place. This is usually a bad sign. Finally, I should say that I break glassware the most frequently when I have a lot of glassware in my hood and it is crowded. In this situation, it is very easy to knock something over accidentally. For this reason, it is important to either keep your hood clean or put glassware that you are not using safe away in the corner of the hood.

4) The reaction not working.

Making sure that your glassware is fairly clean is a good first step to ensuring that at least you get some product from your reaction. TLC is also an extremely useful tool in determining if you have some product. So if you can perform TLC, it is generally a good idea to do this as soon as possible so that you may monitor your reaction carefully. Another good idea is to double check that you are actually adding the correct reagents. Some chemical names can be very similar to one another. Some examples are manganese sulfate vs. magnesium sulfate or ether vs. petroleum ether.

5) Writing a good lab report

Along with any other instructions, be sure to include what your yield was and why you think what it was. For this explanation, it is best to include by physical and chemical rationales. You should try to come up with any side reactions that might be occurring. In addition, you should explain in detail all of the analytical results you used to determine the purity and identity of your product (TLC, IR, NMR, melting point, physical appearance, etc). This is important evidence that you must draw on throughout your lab report.

As a final word of encouragement, I will tell you this. If you are unsure as to whether a possible explanation you are writing in your lab report is correct, include it anyways! You can always preface it with it is possible that. In chemistry, a lot of different things can happen to some extent (everything is in equilibrium!) so chances are on your side that a certain reaction or explanation you posit is valid, at least to some extent.

Good luck with it! Enjoy your chemistry lab! Chemistry is supposed to be fun!

Backyard Chemistry
September 24, 2007
23:02
College has completely killed my lab time
» ‎ Backyard Chemistry

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.
September 07, 2007
13:03
Sodium permanganate?
» ‎ Backyard Chemistry

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!
September 02, 2007
2:38
More Fun With Oxalic Acid: Synthesis of Sodium trioxalatoferrate(III) monohydrate
» ‎ Backyard Chemistry

As always, you can view this experiment and many others at [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 --> CaSO4(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 chlorines 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
August 30, 2007
23:43
Backyard Cemistry
» ‎ Backyard Chemistry
Hello All,

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

Current amateur experiments listed are:
Enjoy!