Sciencemadness Discussion Board

Everyday Chemistry (provisional)

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CharlieA - 21-2-2019 at 18:19

Quote: Originally posted by arkoma  
Squat to do with chemistry, but I be cruising already on my metal/carbon fiber leg. w00t


Isn't science wonderful (better living through chemistry).

I'm glad to hear that you are doing well. In my limited experience (16-yr cancer survivor), attitude is everything, and you seem to have the right one! :) I became much more relaxed when I finally realized that there was no point in obsessing over things that I had no control over.
Regard, Charlie

mayko - 22-2-2019 at 21:06

dehydrated some oxalic acid under toluene. It would have been a lot easier to have a Dean-Stark but otherwise it went well

mayko - 9-4-2019 at 18:47

Got some orange peels steam-distilling, no broken glassware yet!

Herr Haber - 7-8-2019 at 07:41

I routinely check on the chemicals I have in store.
Some of them are in cold storage for a variety of reason.

What a surprise to be greeted by a strong ammonia smell when I looked at the box containing solid bases (lots of NaOH, KOH even more NaHCO3 etc.).
Culprit was 1 kg of food grade Ammonium bicarbonate. Who would have guessed the decomposition temperature was so low !

Σldritch - 18-8-2019 at 09:21

Finally completed Potassium Azide synth. What a pain that was. Making the Hydrazine Sulfate was the easy part surprisingly.

I started by making fresh isopropyl nitrite with some old homemade nitrite of dubious quality and of course it failed. Made a new good looking triple size batch which i accidentally fused into the bottom of the beaker. I lazily decided i would not need to dissolve it again because it is so soluble but it ended up getting coated with Potassium Chloride and it failed again. New batch again, takes a day to boil down all the while my hotplate keeps failing triggering the fuse of the apartment repeatedly. Because i poked a hole in my wide steel container i have to use a taller one than ideal for the batch making it extremely difficult to get the water out. I had to use up all the fuel of my two burners. After two days work i get my isopropyl nitrite.

Azide synth goes really well until the ***************** hotplate fails again. Luckily it worked out anyway though my product is an embarassingly yellow moist powder. In the end i think i used too much nitrite. At least now i can go to uni with my mind at peace having got my fill of chemistry for some time.

Ubya - 18-8-2019 at 13:31

Quote: Originally posted by Σldritch  
Finally completed Potassium Azide synth. What a pain that was. Making the Hydrazine Sulfate was the easy part surprisingly.

I started by making fresh isopropyl nitrite with some old homemade nitrite of dubious quality and of course it failed. Made a new good looking triple size batch which i accidentally fused into the bottom of the beaker. I lazily decided i would not need to dissolve it again because it is so soluble but it ended up getting coated with Potassium Chloride and it failed again. New batch again, takes a day to boil down all the while my hotplate keeps failing triggering the fuse of the apartment repeatedly. Because i poked a hole in my wide steel container i have to use a taller one than ideal for the batch making it extremely difficult to get the water out. I had to use up all the fuel of my two burners. After two days work i get my isopropyl nitrite.

Azide synth goes really well until the ***************** hotplate fails again. Luckily it worked out anyway though my product is an embarassingly yellow moist powder. In the end i think i used too much nitrite. At least now i can go to uni with my mind at peace having got my fill of chemistry for some time.

tine to get a better hotplate :')

mayko - 28-12-2019 at 20:04

Currently dissolving some copper wire in dilute perchloric acid + peroxide :cool:

Abromination - 29-12-2019 at 01:17

Am experimenting with boric acid catalysts in Fischer esterifications, I have seen too many people mention how it may work and haven’t seen anyone try it.

John paul III - 29-12-2019 at 16:19

Made some calcium peroxide from calcium acetate and 3% H2O2 to see if it worked. Will upscale with 60% H2O2 I left in my rental apartment and see what this thing does mixed with sugar

Fery - 30-12-2019 at 09:16

Quote: Originally posted by mayko  
Currently dissolving some copper wire in dilute perchloric acid + peroxide :cool:

mayko you certainly had / have / will have interesting ideas which I'm missing due to my limits, but why didn't you react CuO + HClO4?
Did you want to prepare Cu(ClO4)2 ASAP and avoid waste of time to prepare/buy CuO?
Was it an attempt to prepare very pure product (metallic Cu of high purity is more available than similar grade purity CuO)?
Or did you want to dissolve Cu in anything else than already well known, boring, stinking ways HNO3, hot conc. H2SO4? HCl + H2O2 would stink too so that's why you used HClO4? Also Na2S2O8 is used for dissolving Cu layer of PCB for electronics and another etching reagent for Cu PCB is FeCl3.

j_sum1 - 1-1-2020 at 17:34

I did an ethanol extraction of some camphor laurel leaves. This was more an excuse to use the soxhlet than anything else. But given that my street is lined with these trees, camphor extraction has been on the list for a while. I doubt my efficacy. I seem to have a lot of chlorophyll even though the leaves are dry and brittle. I will leave my bag of leaves a month or two more and try again. Probably a different method too.

Today I got my 8y old son to help me distill and recover the ethanol. It was a chance to use my Graham condenser: the cool appearance surpasses its functionality in most situations but it performed beautifully in this instance. Distilling ethanol is so satisfying: it just pours out of the condenser. My son enjoyed it too. Good times.

Fery - 2-1-2020 at 00:02

j_sum1 - steam distillation?
as you have plenty of material for free, maybe you can build some medium scale apparatus from an old pressure cooker

arkoma - 2-1-2020 at 00:56

Quote: Originally posted by j_sum1  
It was a chance to use my Graham condenser: the cool appearance surpasses its functionality in most situations but it performed beautifully in this instance. Distilling ethanol is so satisfying: it just pours out of the condenser.



Yup. Only use i've ever had for my Graham. LOL

Abromination - 2-1-2020 at 13:57

Well for all of those who have wondered, boric acid does not catalyze fischer esterification alone. Ill add a dehydrating agent to see if that can drive it away from hydrolysis.

Tsjerk - 2-1-2020 at 14:45

I have been looking around on the webs today a bit after finding some publication describing use strong acid ion exchange resin as a catalyst for esterfications...

After being done you only have to filter and the catalyst is gone and reusable.

Appearantly any old "Brita" kitchen drinking water purifier can be opened to get the resin which can be regenerated by soaking in acid.

DraconicAcid - 2-1-2020 at 14:46

Really? I thought the Brita filters only had activated charcoal.....

Tsjerk - 2-1-2020 at 16:30

At least not the Dutch ones.. I haven't taken one apart yet but the brand Brita cartridges should have a resin compartment next to the charcoal compartment.

Edit: as long as it states to de-harden water it should have the sulfonated styrene resin.

[Edited on 3-1-2020 by Tsjerk]

Abromination - 2-1-2020 at 18:34

Quote: Originally posted by Tsjerk  
At least not the Dutch ones.. I haven't taken one apart yet but the brand Brita cartridges should have a resin compartment next to the charcoal compartment.

Edit: as long as it states to de-harden water it should have the sulfonated styrene resin.

[Edited on 3-1-2020 by Tsjerk]

That sounds awesome, I have one I can use!
EDIT:
Heres the box, it contains both activated carbon and ion exchange resin, which I assume is a sulfunated styrene.
In a bit I will take it apart and attempt an esterification with it. I have never tried using a sulfunated resin but I have heard they are excellent catalysts.

[Edited on 1-3-20 by Abromination]

Abromination - 2-1-2020 at 18:49

Sorry for the double post, wouldnt let me include pic

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arkoma - 19-1-2020 at 12:58

Isopropyl nitrite. Per Rhodium.

IMG_20200119_132211.jpg - 927kB

*edit* photos of yield


IMG_20200119_185613.jpg - 880kB IMG_20200119_185755.jpg - 865kB

[Edited on 1-20-2020 by arkoma]

DraconicAcid - 6-2-2020 at 12:34

We happened to get some dry ice in with a sample the other day, so I did a Grignard reaction with benzyl chloride to get phenylacetic acid. I messed it up, but still managed to get a small yield of white, clean-looking product. I've washed my hands a dozen times, but I can still smell it.

mayko - 11-2-2020 at 21:29

I ran out of concentrated ammonia solution a while back and I've had some crappy cold-pack ammonium nitrate languishing on the shelf, so I'm liberating the gas with lye and redissolving. The sodium nitrate will go back on the shelf until I run out of nitric acid. I was surprised at just how much NH4NO3 it took (~360 g) to make so little ammonium hydroxide (~250 mL)!

CharlieA - 12-2-2020 at 18:00

Do you know the purities of your starting materials? My calculations (always iffy), show that 4 mols of ammonium nitrate should yield 4 mols of ammonium hydroxide. 4 mols/250 mL = 16 M which seems to be high. Did you titrate the ammonium hydroxide solution to determine its molarity?

mayko - 12-2-2020 at 18:42

I based my calcs off this passage from Wikipedia:

Quote:

At 15.6 °C (60.1 °F), the density of a saturated solution is 0.88 g/ml and contains 35.6% ammonia by mass, 308 grams of ammonia per litre of solution, and has a molarity of approximately 18 mol/L. At higher temperatures, the molarity of the saturated solution decreases and the density increases.


So it looks like your math is as good as the starting numbers were! I haven't titrated, but most of the expected mass increase is there, and I haven't had the huge cloud of excess ammonia gas I'd expect if there was too little water. Apparently ammonia is just ludicrously soluble! :o



Attachment: ammonia_concentration.xlsx (5kB)
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CharlieA - 13-2-2020 at 18:18

quote: So it looks like your math is as good as the starting numbers were!

What do you know! Apparently it is true that a blind hog finds an acorn every now and then! Obviously I should have checked the solubility of ammonia in water. I suppose an accurate pH measurement of the solution would allow a calculation of the concentration also. I am very gratified to see someone report a result that they have checked against theory. This doesn't seem to be the case as often as it should be, and I confess that I am guilty of this too often.


Abromination - 15-2-2020 at 10:38

My experiment using boric acid and calcium chloride to catalyze a Fischer esterification is going quite well. The reflux time has taken quite a while but the smell of ethyl salicylate is strong now and a spot test leaves an oily, wintergreen residue. Boric acid alone was unsuccessful, but with the dry calcium chloride things are looking good.

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Tsjerk - 15-2-2020 at 11:49

Nice synth! I just saw the monohydrate of CaCl2 only loses water at 260 degrees and the dihydrate at 175... That is interesting... I guess that could be used for other esterfications.

Abromination - 15-2-2020 at 16:28

The idea was to remove water from the reaction. Reports of boric acid catalysed esterification have been floating around for a while, and after limited success with that alone I figured the drying agent would help. Seems to work!

j_sum1 - 15-2-2020 at 18:43

Today I prepared a buffer solution to pH 7.4 on a 4400 litre scale.



The pool is full of fresh clean water and looks great.

AvBaeyer - 15-2-2020 at 19:41

Abromination:

Your esterification experiment is interesting but I believe demands a bit more rigor before you can conclude that the CaCl2/B(OH)3 combination is an esterification catalyst. The long reaction time that your post implies raises the possibility that you are simply reaching an equilibrium between salicyclic acid and its ester. Salicylic acid is a strong enough acid to catalyze its own partial esterification (pKa = 2.98) given enough time and high enough temperature. I suggest that this may be what you are seeing that is giving you your product. Also keep in mind that CaCl2 is not a good water scavenger in alcohol solvents as the alcohol can compete with water for coordination sites on calcium, especially when a large excess of alcohol is present relative to the water produced. Following the scientific method, I suggest the experiments outlined below:

1. React salicylic acid and ethanol without any catalyst under the exact same conditions you are using in the experiment you described (molar ratios, time, temperature). Compare ester production versus your current experiment as best you can.

2. Repeat (1) using boric acid as a catalyst alone. It appears that you have tried this but you need to do it as in (1) above.

3. Repeat (1) using CaCl2 alone.

I hazard a guess that experiments 1, 2,and 3 will give results the same as your current observation. These kinds of things make chemistry interesting and can lead to new discoveries.

AvB

Tsjerk - 15-2-2020 at 22:22

Does anyone have data about the temperature where CaCl2 starts to lose methanol / ethanol as adduct? This is interesting because if that temperature is below the boiling point of the given alcohol, it should work as a catalyst with a minimal amount of acid.

Abromination - 16-2-2020 at 00:19

Quote: Originally posted by AvBaeyer  
Abromination:

Your esterification experiment is interesting but I believe demands a bit more rigor before you can conclude that the CaCl2/B(OH)3 combination is an esterification catalyst. The long reaction time that your post implies raises the possibility that you are simply reaching an equilibrium between salicyclic acid and its ester. Salicylic acid is a strong enough acid to catalyze its own partial esterification (pKa = 2.98) given enough time and high enough temperature. I suggest that this may be what you are seeing that is giving you your product. Also keep in mind that CaCl2 is not a good water scavenger in alcohol solvents as the alcohol can compete with water for coordination sites on calcium, especially when a large excess of alcohol is present relative to the water produced. Following the scientific method, I suggest the experiments outlined below:

1. React salicylic acid and ethanol without any catalyst under the exact same conditions you are using in the experiment you described (molar ratios, time, temperature). Compare ester production versus your current experiment as best you can.

2. Repeat (1) using boric acid as a catalyst alone. It appears that you have tried this but you need to do it as in (1) above.

3. Repeat (1) using CaCl2 alone.

I hazard a guess that experiments 1, 2,and 3 will give results the same as your current observation. These kinds of things make chemistry interesting and can lead to new discoveries.

AvB

Just now saw your post, so sorry for posting twice.
Boric acid was tried before under the same conditions. This resulted in the smell but negative results with a spot test.

My next step was already to try calcium chloride alone as well as just the acid and alcohol for a control. I am aware of the slight reaction between the two.

The point of the experiment is to try different reagents as a catalyst, so failure does not upset me. When each method has been attempted, I will repeat each of them again in the same conditions to stay true to the scientific method and isolate as many variables as possible. For right now its more of a matter of feasibility and curiosity.

SWIM - 16-2-2020 at 20:30

Aren't you worried about the boric acid forming esters with the alcohol?

This would both tie up your acid and add water to the reaction mixture.

You might deal with the water generated by using metaboric acid or boron trioxide mixed in with your boric acid ( edit:Or the CaCl you did use), but it might make the ethyl borate formation worse.

That's the thing: You want to soak up the water to drive the reaction but boric acid + alcohol makes water, and drying up that water also drives that equilibrium to the right as well as the equilibrium of the reaction you want.

If I'm totally wrong about boric acid forming esters in such conditions just ignore this.

[Edited on 17-2-2020 by SWIM]

Abromination - 16-2-2020 at 20:59

Quote: Originally posted by SWIM  
Aren't you worried about the boric acid forming esters with the alcohol?

This would both tie up your acid and add water to the reaction mixture.

You might deal with the water generated by using metaboric acid or boron trioxide mixed in with your boric acid ( edit:Or the CaCl you did use), but it might make the ethyl borate formation worse.

That's the thing: You want to soak up the water to drive the reaction but boric acid + alcohol makes water, and drying up that water also drives that equilibrium to the right as well as the equilibrium of the reaction you want.

If I'm totally wrong about boric acid forming esters in such conditions just ignore this.

[Edited on 17-2-2020 by SWIM]

I’m sure its possible in some quantity, but I beleive such a reaction requires a decent amount of a very strong dehydrating agent such as sulfuric acid or phosphorus pentoxide to occur. It is a factor that could also lower yield. The idea with the calcium chloride was to remove water from the reaction without taking place in the reaction itself to drive it right, but I am questioning its effectiveness in an alcoholic environment.

As stated before, this is not to find a better way to do an esterification, it is merely a project for my own curiosity and for fun. Its not particularly costly or time consuming, either.

mayko - 25-2-2020 at 18:45

As I was wrapping up the ammonia project, I ran across something in Merck that I hadn't heard before... apparently concentrated ammonia dissolves copper metal? I scrubbed a bit of copper wire until it was shiny and put it in a corked test tube with a couple mL of concentrated ammonia. Sure enough, over a couple of days it turned the dark blue of an ammonia-copper complex!

woelen - 25-2-2020 at 23:34

Concentrated ammonia only dissolves copper if also oxygen is present. The oxygen oxidizes the copper and the ammonia makes it possible for the copper to dissolve as the [Cu(NH3)4](2+) complex. Your solution turns strongly basic, the oxide ends up as hydroxide ion.

This effect actually is quite common. If a suitable coordinating agent is present in solution, then many metals can be oxidized much more easily than without the coordinating agent. Best example is the solubility of gold in a solution of KCN through which air is passed. A very nice example in the home lab is the dissolving of copper in dilute acid (e.g. 10% HCl) to which some thiourea is added. The copper then dissolves, with formation of hydrogen! Without the thiourea, copper does not dissolve in non-oxidizing dilute acid.

Bedlasky - 26-2-2020 at 03:50

That's really interesting! I have some silver cleaner containing thiourea, so I definitely try it.

fusso - 26-2-2020 at 05:14

So that means thiourea reverse the redox potentials for the rxn?

woelen - 26-2-2020 at 05:31

The thiourea does not reverse any potentials, it shifts redox potentials. In general, nearly all coordinating agents shift redox potentials, but some compounds have a strong effect, while others have less effect.

In the example of thiourea and copper, the redox potential of the reaction from metallic copper to copper(I) is shifted, such that it goes below the reaction from H2 to H(+) ion. Copper is somewhat noble, but not much so, and it is shifted to the range, where metals can dissolve in non-oxidizing acids with production of H2. The copper itself in this reaction then forms a thiourea-complex of copper(I).

Medyc - 28-2-2020 at 09:38

I know we've moved past the boric acid catalyzed esterification, but i do have experience with this.

DOI: 10.1021/ol036123g

I have used this method to prepare the alpha hydroxy ester of threonine. Good yields, and largely chemoselective between acids.


[Edited on 28-2-2020 by Medyc]

mayko - 9-4-2020 at 21:40

Stores are still largely empty of cleaning supplies here! I had a bucket of calcium hypochlorite granules in the shed so I set about converting it into chlorine bleach with sodium carbonate. Filtering the calcium carbonate was the hardest part - it is very fine and the concentrated bleach quickly destroys paper and cotton! :o Next time I might try to set up a large scale vacuum filtration system.

I measured concentration by density and wound up with 5.5 gallons of ~6% bleach. I think I could do it more efficiently next time. This was way more than I needed so I donated some to the local makerspace (which is staying open making face shields) and the rest to a mutual aid group for distribution :cool:

arkoma - 10-4-2020 at 12:28

Got a 1943 Mercury dime the other evening and promptly set it to soak in 68% HNO3. A 1968 Canadian nickel also.

The silver solution has a slight blue cast from the 10% Cu, and the nickel solution is a lovely green.

wg48temp9 - 11-4-2020 at 04:06

Quote: Originally posted by Medyc  
I know we've moved past the boric acid catalyzed esterification, but i do have experience with this.

DOI: 10.1021/ol036123g

I have used this method to prepare the alpha hydroxy ester of threonine. Good yields, and largely chemoselective between acids.


Interesting paper thanks. It mentions a 1971 paper claiming it was the first paper suggesting the use of boric acid for esterification but only for phenol esters. The paper is attached below

Attachment: boric-acid-lawrence1971.pdf (69kB)
This file has been downloaded 146 times

fusso - 11-4-2020 at 07:55

Quote: Originally posted by arkoma  
Got a 1943 Mercury dime the other evening and promptly set it to soak in 68% HNO3. A 1968 Canadian nickel also.
Such a waste arkoma, shouldve save a few more for bulk rxn to save reagents

arkoma - 11-4-2020 at 11:25

Quote: Originally posted by fusso  
Quote: Originally posted by arkoma  
Got a 1943 Mercury dime the other evening and promptly set it to soak in 68% HNO3. A 1968 Canadian nickel also.
Such a waste arkoma, shouldve save a few more for bulk rxn to save reagents


Nah. I've spent $100 making something I could buy for $3 Only other silver i have is a Kookaburra from Oz, and AIN'T using it.

*edit* the HNO3 was home sourced also.

*edit* Some diethyl ether



[Edited on 4-11-2020 by arkoma]

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[Edited on 4-11-2020 by arkoma]

mayko - 9-5-2020 at 11:24

I've been experimenting with using the paper straws that are popular these days as casings for DIY fireworks. I've had a few successes, but I think the best part so far has been lighting Zn/S at night and shining the shortwave UV Tdep gave me through the plume of phosphorescent smoke. It's like incense smoke in a sunbeam, except it's ghostly green and everything is dark. It doesn't show up well on my camera, unfortunately

Abromination - 11-5-2020 at 12:18

Ran some sulfur dioxide through concentrated hydrogen peroxide to restock my sulfuric acid (so inconvenient, really wish you could buy it here), was a great test for my new fumehood. I had never noticed how endothermic the reaction of HCl and sodium metabisulfite is, I had always ended up having to do it in freezing weather outside before now.

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j_sum1 - 30-5-2020 at 21:10

Distilled some otc isopropyl alcohol. Seemed to be already high purity. Head temp held rock steady and I discarded only a couple of mL of heads and tails.

IPA is not quite the ubiquitous substance here that it is in the US. Nice to find such good quality.

mayko - 1-6-2020 at 15:09

I'm trying my first medium-scale brew & still.... 3 gallons of what the yeast pack promises is 20% alcohol are in the fridge and I'm setting up to try distilling in the backyard tomorrow :)

arkoma - 1-6-2020 at 19:02

I find it quite fulfilling distilling small (gallon size batch of 55% EtOH) amounts of alcohol. And besides, I'm chipping in on the Sars-Co2 abatement. Hand sanitizer is mostly EtOH.

DraconicAcid - 28-7-2020 at 21:37

I've got my son (a 16-year-old version of me) helping me in the lab, videoing some experiments so that this semester's social-distancing rules won't kill off our college's chemistry labs. I am very proud to be able to give him the lab experiences that I wished I could have had at that age.

In addition to lots of esterification reactions, TLC, the copper carnival, and distillations, we also prepared some hydrated copper compounds (for the students to determine the degree of hydration in). Just for fun, we decided we'd make copper sulphamate (following this prep: https://prepchem.com/synthesis-copper-sulfamate/ ), since I saw a lovely sample of it on Instagram, and copper(II) sulphanilate. The latter did not work at all- after reducing in volume, it started to crystallize and solidify as it cooled, so we didn't need to triturate it with acetone. Recrystallizing it from water/methanol gave pale blue crystals which turned white when washed with methanol- clearly not very coppery. According to this: https://patents.google.com/patent/US3321273A/en it's probably ammonium bisulphate.

The copper(II) sulphanilate seems to have worked, giving a nice green precipitate from the reaction of copper(II) acetate with sulphanilic acid (eventually- some boiling down of the solution was required. I'm hoping no hydrolysis of the anion took place).

HeYBrO - 28-7-2020 at 21:48

Quote: Originally posted by j_sum1  
Distilled some otc isopropyl alcohol. Seemed to be already high purity. Head temp held rock steady and I discarded only a couple of mL of heads and tails.

IPA is not quite the ubiquitous substance here that it is in the US. Nice to find such good quality.


Nice, did you dry with NaOH?

You can get relatively large amounts from eBay and if you want smaller quantities (ca. 250 ml) you can buy 99% isopropanol from jay car (at least in vic...). I also noticed it at pharmacies as a 60% solution but it was stupidly expensive.

j_sum1 - 29-7-2020 at 02:33

Quote: Originally posted by HeYBrO  

Nice, did you dry with NaOH?

I could say that I didn't feel the need to. But the truth is that it did not occur to me at the time. But I am not doing anything water sensitive with it anyway.

My main reason for getting it was because I needed a secondary alcohol for teaching. I had a really nice lab session with a handful of kids just over a month ago. Alcohol oxidation with acidified dichromate was one of the things covered.

j_sum1 - 14-8-2020 at 06:39

Low on ammonia solution. So I spent an evening making a batch. Finished up ith 222mL at 22%.
Nitric acid tomorrow.

Lion850 - 15-8-2020 at 02:50

J_sum1 I am not quite ready to make nitric acid. I’m sticking to simple reactions within my limits :) Fortunately 70% nitric acid and ammonia solution are both available locally.

I’ve returned to cyanurates while waiting for more glycine; nickel carbonate and cyanuric acid have been stirring in the shed since Friday.

arkoma - 31-8-2020 at 22:37

do believe my CH3NH3 has turned out. No MP test, but slightly "fishy" and VERY hygroscopic. Also give a blue reaction to Simon's reagent.

DraconicAcid - 10-9-2020 at 21:05

So I was trying to get some nice crystals of copper(II) lactate, but it seems that it forms a mixture of different hydrates and/or crystallizes with excess lactic acid. So I said to heck with this- I'll precipitate it out as the carbonate. Added potassium carbonate, and got a very dark blue solution, matching the description of an anionic lactate complex that I recently read about. Bother.

Bedlasky - 11-9-2020 at 12:05

Just now I trying to make potassium iodate by permanganate method according to recipe from my high school. Right now reaction is done and I filter manganese dioxide from solution. And it is pain in the ass. I filter it through coffee filter and solution goes through incredibly slowly. Filtrate is still brown, but most of MnO2 remains in filter. After filtration I'll do another filtration through finer filter which retain fine particles of MnO2.

MidLifeChemist - 11-9-2020 at 17:08

Today was Cobalt day for me and the family.

I added NaOH to CoSO4, and got a nice blue precipitate.

In a separate container, I added CoSO4 to excess NaOH although the precipitate was blue for a moment, it quickly turned pink.

Gravity filtered both of them with coffee filters (blue solid filtered quickly, the pink solid filtered extremely slowly) and now letting them both dry. Will heat them both to convert to Cobalt Oxide, then I'l weigh and then add to my compound collection. I'm curious if either the pink or blue form "converts" to the other during these processes.

I also added CoSO4 to excess NH4OH, and the solution turned brown. Adding some NH4OH to CoSO4 gave a blue precipitate, like adding NaOH did.

Bedlasky - 11-9-2020 at 20:53

Co(OH)2 will react with atmospheric oxygen to produce brown CoO(OH).

Something about blue and pink form is on wiki:

https://en.wikipedia.org/wiki/Cobalt(II)_hydroxide

Brown soluble complex with ammonia is [Co(NH3)6]2+ and it's slowly oxidized by oxygen in to [Co(NH3)6]3+.

arkoma - 5-11-2020 at 23:18

Figured i'd see what the heck my semi-crappy vacuum gauge was doing for real, so I set up for simple vac disty with h2o.

results:
(temps are B.P,'s)("inch" figures what my gauge read)

66C showed 25"=286mmHG
73C showed 21"=354.2mmHG
88C showed 15"=535.4mmHG

Ain't been real scientific lately, so its something......

Cou - 5-11-2020 at 23:56

I just evaporated hexane and boiled off any remaining 1-pentanol from about 30 mL of crude pentyl nonanoate. Next I need to wash with bicarb, water, and NaCl, then purify by vacuum distillation.

DraconicAcid - 6-11-2020 at 00:02

My son and I are doing titration curves for the students to study next semester. Got a surprisingly smooth titration curve from vanillin after adding a bit of ethanol.

Also enjoying the smoothest, fruitiest crabapple wine I've ever made.

Sulaiman - 6-11-2020 at 05:22

Distilled some sulphuric acid, now I have
C40ml Az. (cloudy - remainder from the pot, not distilled)
25ml azeotropic, 350 ml c98% and 350ml c28% clean distilled sulphuric acid.

I was surprised at the difference in the dehydrating power of azeotropic sulphuric acid and the 93% to 96% drain unblocker that I'm used to,
and when hot, the dehydrating power of az. sulphuric acid is awesome - and terrifying.

Distilled using Chinese glassware and a cheap portable butane gas stove.
During tidy up I noticed a crack in the side joint of my distillation head, no other problems.

hasani10 - 6-11-2020 at 12:45

finally made some lead carbonate from the acetate i had from ages ago. Back home from uni so got access to my home lab, whapped out the silly looking janky vacuum pump i made to filter it.

Also ordered a bunch of stuff for future projects! Soon i'll be set for sodium benzoate -> benzene -> nitrobenzene -> aniline -> dimethylaniline -> crystal violet.


beta4 - 6-11-2020 at 13:04

Quote: Originally posted by hasani10  

Also ordered a bunch of stuff for future projects! Soon i'll be set for sodium benzoate -> benzene -> nitrobenzene -> aniline -> dimethylaniline -> crystal violet.



Interesting pathway, I considered trying it, but after looking into it in detail, I gave up due to:
- toxicity of nitrobenzene
- difficulty in finding information on a good and readily available methylating agent for aniline

How do you plan to make the dimethylaniline?

Heptylene - 14-11-2020 at 07:18

I distilled some fuming nitric acid this afternoon! I decided to go with a relatively small scale: 2 moles in a 1 liter flask to manage the foaming. This worked okay, but I found out that my apparatus leaks near the receiver and it turns out kek clips do not tolerate fuming nitric acid, even cold. They're made of polyacetal and I strongly suspect they hydrolyse to formaldehyde or some other nasty product.

I think I'll get a new apparatus made of a single piece of glass instead of head+condenser+receiver, it's not the first time it has become leaky at an inconvenient time. This should make distilling nitric acid much less of hassle if I don't have clouds NO2 and acid coming out of the joints.

In the end the product was obtained in 83% yield with a density of 1.518(5), which corresponds to about 95-97% concentration or 23 M.

arkoma - 14-11-2020 at 09:37

soxhlet extracted a 3oz can of Dollar General ground black pepper with acetone (IPA is HARD to find ATM). Obviously didn't use enough KOH to get the oils out, because i've got some nice big globs of orange-ish crap in my (poor) workup, along with some lovely yellow piperine. Might have been slightly inebriated. Haven't thrown out any of the extraction solvent yet though.

Have a beaker with some interesting stuff growing in it too--some ostensibly Benzyl alcohol I dropped a bit of K Dichromate in thats been sitting on a shelf. Smells as good as crazy it looks.

Will post pics.

Heptylene - 14-11-2020 at 09:59

Oh Arkoma I feel for you, pepper extract is a bitch to work up. I've done it 3 or 4 times and everytime there was bunch of resinous crap in the product. I can't figure out how to avoid this completely. Maybe using dry solvents and dry pepper, using more KOH, or something else...

DraconicAcid - 14-11-2020 at 10:24

Agreed. I've had my students run it through a column and it still turns out resiny.

arkoma - 15-11-2020 at 09:26

I have a small amount of piperine i extracted 5(?) years ago that came out so lovely. Before mechanical breakage in the ampoule they are in (from handling) had nice yellow needles.



piperine.jpg - 677kB

This is everyday chemistry too.

Quote:
I've got my son (a 16-year-old version of me) helping me in the lab, videoing some experiments so that this semester's social-distancing rules won't kill off our college's chemistry labs. I am very proud to be able to give him the lab experiences that I wished I could have had at that age.


First off, KUDOS to all teachers here. So important. I REMEMBER the names of my sixth, seventh, ninth, and eleventh grade science teachers NAMES. All the others? Not so much.

Had a moment yesterday.

This 14 year old kid that's been working at our shop (I don't do anything except take up space there these days) was with me at my house for a while yesterday, and my lab is right on front street so to speak. He was looking at it, and I started talking about it, as we are won't to do (we all like to share our favorite "drugs" <<metaphor). I started losing him quick with the O Chem shit. Had copper sulfate solution and a repurposed cell charger at hand. Said kid, you got a quarter on you? he said why. i said, give me a quarter. He had a dime, and as he watched copper deposit on it before his eyes and wanted to know the WTF? of it, i knew i had hit a home run. Also (mua-ha-ha) had tin chloride solution. Dendritic tin is even MORE amazing visually.

Yeah, that kid is hooked I hope.

Everyday chemistry at it's best!

teachers that touched my life:
Miss Astor
Mr. Gadd
Miss McEvoy
Mr. Gerofsky

Thank You all.


[Edited on 11-15-2020 by arkoma]

[Edited on 11-15-2020 by arkoma]

beta4 - 15-11-2020 at 10:15

Quote: Originally posted by Heptylene  
I think I'll get a new apparatus made of a single piece of glass instead of head+condenser+receiver, it's not the first time it has become leaky at an inconvenient time.


There's a simpler solution: buy stainless steel keck clips. They can be found easily on ebay,
Honestly I don't know why people bother with the plastic ones. They break after a few uses and you have to keep replacing them, other than being prone to breaking/melting in the middle of a distillation when distilling something aggressive/with a high boiling point.

[Edited on 15-11-2020 by beta4]

woelen - 15-11-2020 at 11:00

Did some crystallization of a mixed ethylenediamine/ammine complex of nickel(II). Something like Ni(en)2(NH3)2(ClO4)2. Beautiful dark blue crystals. Of course I made a picture of the crystals.

In the past I have done experiments with Ni-(en) complexis, and the not fully coordinated Ni(en)2(H2O)2 complex is dark blue. I did not manage to crystallize that. I only could get the purple Ni(en)3(ClO4)2 complex. But now I think I have managed to make the ammonia analogue Ni(en)2(NH3)2(ClO4)2 from a dark blue solution of nickel perchlorate and ethylene diamine to which I added some ammonia.

DraconicAcid - 15-11-2020 at 11:16

Quote: Originally posted by woelen  
Did some crystallization of a mixed ethylenediamine/ammine complex of nickel(II). Something like Ni(en)2(NH3)2(ClO4)2. Beautiful dark blue crystals. Of course I made a picture of the crystals.

In the past I have done experiments with Ni-(en) complexis, and the not fully coordinated Ni(en)2(H2O)2 complex is dark blue. I did not manage to crystallize that. I only could get the purple Ni(en)3(ClO4)2 complex. But now I think I have managed to make the ammonia analogue Ni(en)2(NH3)2(ClO4)2 from a dark blue solution of nickel perchlorate and ethylene diamine to which I added some ammonia.


Nice. I've tried making Ni(en)2 complexes with anthranilate and oxalate, but never had much luck in getting anything but the tris(en) complexes.

Sulaiman - 15-11-2020 at 11:32

Quote: Originally posted by beta4  
There's a simpler solution: buy stainless steel keck clips.

I have corroded stainless steel Keck clips with the nitric acid vapours invisibly escaping at joints.
I don't like using conc. H2SO4 to seal joints and grease is attacked by the hot nitric acid
My best option so far is ptfe tape instead of grease.
(and use ss Keck clips)

artemov - 15-11-2020 at 19:39

Quote: Originally posted by Sulaiman  

I don't like using conc. H2SO4 to seal joints


Why do you not like this? I use 60% H2SO4 when distilling HCl and low boiling alcohols. I hope I'm not doing things wrong ...

Sulaiman - 15-11-2020 at 22:36

nothing wrong, its just me, so that I do not forget an get it on me ... trivial really.

Bezaleel - 16-11-2020 at 14:15

Made some Ni(en)3I2. Dissolved this in dilute KI to recrystallise.

20201114_162514_adj_30pct_small.jpg - 78kB

mayko - 16-11-2020 at 21:40

I've been trying to extract ibuprofen from OTC pills. should be a simple base/acid extraction, but the filler made the first extraction completely insoluble and when I tried to recrystallize the acid precipitate from IPA, it has instead formed an oily bilayer. What a mess.

MidLifeChemist - 16-11-2020 at 21:59

Quote: Originally posted by Bezaleel  
Made some Ni(en)3I2. Dissolved this in dilute KI to recrystallise.



Did you do that to reduce the solubility via the common ion effect?

Bezaleel - 17-11-2020 at 06:15

Quote: Originally posted by MidLifeChemist  
Quote: Originally posted by Bezaleel  
Made some Ni(en)3I2. Dissolved this in dilute KI to recrystallise.



Did you do that to reduce the solubility via the common ion effect?
No, it's just a final purification step, to get rid of assumed co-crystallised chloride.

MidLifeChemist - 17-11-2020 at 08:07

Got it - ok I'm assuming the idea is that the iodide ions from the KI will replace the chloride ions. So why do you need to do this second step - to get the Ni(en)3I2 the first time, did you simply boil away or evaporate the solution, is that why you need to recrystallize? And how do you know this is the iodide and not the chloride? I couldn't find any references to Tris(ethylenediamine) nickel (II) iodide or to its solubility. It looks like an interesting compound that I may want to try to make one day. Thanks in advance for the info!

Quote: Originally posted by Bezaleel  
Quote: Originally posted by MidLifeChemist  
Quote: Originally posted by Bezaleel  
Made some Ni(en)3I2. Dissolved this in dilute KI to recrystallise.



Did you do that to reduce the solubility via the common ion effect?
No, it's just a final purification step, to get rid of assumed co-crystallised chloride.

Bezaleel - 17-11-2020 at 10:03

Posted my answer here.
Quote: Originally posted by MidLifeChemist  
Got it - ok I'm assuming the idea is that the iodide ions from the KI will replace the chloride ions. So why do you need to do this second step - to get the Ni(en)3I2 the first time, did you simply boil away or evaporate the solution, is that why you need to recrystallize? And how do you know this is the iodide and not the chloride? I couldn't find any references to Tris(ethylenediamine) nickel (II) iodide or to its solubility. It looks like an interesting compound that I may want to try to make one day. Thanks in advance for the info!
Quote: Originally posted by Bezaleel  
Quote: Originally posted by MidLifeChemist  
Quote: Originally posted by Bezaleel  
Made some Ni(en)3I2. Dissolved this in dilute KI to recrystallise.

Did you do that to reduce the solubility via the common ion effect?
No, it's just a final purification step, to get rid of assumed co-crystallised chloride.

woelen - 18-11-2020 at 02:16

Quote: Originally posted by woelen  
Did some crystallization of a mixed ethylenediamine/ammine complex of nickel(II). Something like Ni(en)2(NH3)2(ClO4)2. Beautiful dark blue crystals. Of course I made a picture of the crystals.

In the past I have done experiments with Ni-(en) complexis, and the not fully coordinated Ni(en)2(H2O)2 complex is dark blue. I did not manage to crystallize that. I only could get the purple Ni(en)3(ClO4)2 complex. But now I think I have managed to make the ammonia analogue Ni(en)2(NH3)2(ClO4)2 from a dark blue solution of nickel perchlorate and ethylene diamine to which I added some ammonia.

:( My nice crystals of this complex have withered in the last few days. I had put them in a dry place to have perfectly dry crystals, but I think that they lost part of their ammonia. The crystals now look ugly with a pale blue/green powder, sticking to the crystals, and some crystals are halfway gone, simply disintegrated into powder. A solution of this material in water is turbid. Most likely, besides simple loss of ammonia, also some basic carbonate salt is formed, contaminated with ammonium perchlorate, due to absorption of CO2 from air.

So, the ammine-salts cannot be dried succesfully in air. Maybe next weekend I try again, but then with rinsing first with acetone, followed with a rinse of diethyl ether and then quickly transferring the crystals to a glass vial to keep them around without decomposition. Fortunately I still have the picture of the crystals. I'll make a short write-up with these pictures and share this.

[Edited on 18-11-20 by woelen]

Bedlasky - 18-11-2020 at 05:32

Yes, these ammine complexes really easily lose ammonia. If you make tetraamminecopper(II) sulfate, you must store him in ammonia atmosphere. I still have some of it, I made it few years ago. I put it in the glass jar, added few drops of concentrated ammonia and closed it. It's still perfectly dark blue.

DraconicAcid - 18-11-2020 at 08:59

Nickel ammonia complexes are particularly prone to losing ammonia.

woelen - 18-11-2020 at 11:06

Here follow pictures of the complex. One picture was made just after carefully pressing the crystals between filter paper and tissue to get most of the adhering absorbed in the paper. The other picture was made two days later. After that period the crystals were perfectly dry, but they do not look attractive anymore.



Ni_en2_nh3_2.jpg - 1.3MB



Ni_en2_nh3_2_dried.jpg - 1.6MB

These crystals have a size of appr. 0.5 mm. They are stable, but when heated in a flame, they deflagrate. Still, this experiment was quite interesting. It is possible to isolate a dark blue Ni(en)2-complexes. One more thing to try is to isolate a skyblue/cyan Ni(en) complex, just one (en) ligand attached. The color of this is very different from the bright green of aqueous nickel(II), but its lightness is comparable.



[Edited on 18-11-20 by woelen]

Bedlasky - 21-11-2020 at 09:46

Finally I am able to reduce tungstate in to W(V), but also in to W(IV) and W(III), which I thought was not possible in aqueous solution.

If you reduce tungstate by strong reducing agents in acidic solution (I tried zinc powder, various metal ions in low oxidation state, dithionate and ascorbic acid) you always end with insoluble tungsten blue.

But in the mixture of oxalic acid/hydrochloric acid you can reduce tungsten even in to very dark blue (maybe even darker than molybdenum blue) mixture of W(V) complexes.

If you dissolve tungstate in hot conc. hydrochloric acid and add some aluminium foil, firstly is tungstate reduce in to greenish blue W(V) (reduction must take place in hot conc. HCl, in cold it doesn't work, reaction is very vigorous and produce lot of HCl gas). If you add more aluminium foil, it is reduce in to green W(IV) and than in to reddish brown (and also very unstable) W(III). These conclusions aren't final, I try to find more sources about W(IV) and W(III), but according to my observations and books and articles that I read this is probably right conclusion, but still it need more research, because there are few differences between my observations and what books say.

All reduced tungsten species can be easily oxidized even by mild oxidizing agents like Cu2+ in to tungstate. Cu2+ is reduced in to Cu+ (when I added solutions upon oxidation in to KOH solutions, Cu2O precipitated out).

[Edited on 21-11-2020 by Bedlasky]

DraconicAcid - 22-11-2020 at 23:02

Nice!

My son and I were filming the various vanadium reductions for the online version of the standard vanadium lab, but the reduction with aluminum wire isn't reliable (and mucked up, giving us grey instead of violet, and the titration said it was V(III) instead of V(II)). We'll have to do that part over again- i hate the idea of using Hg amalgam, but I might have to.

Bedlasky - 23-11-2020 at 06:34

For making V(II) is sufficient zinc powder, you don't need AlHg amalgam. I did it thousand times by this way. Just mix vanadate, HCl or H2SO4, zinc powder and mix it well. I did it just cca two weeks ago when I needed some reducing metal ions for molybdate and tungstate reduction.

njl - 23-11-2020 at 07:20

Has anybody heard of urea being converted to ammonium formate? I'm trying to find a renewable source.

DraconicAcid - 29-11-2020 at 11:31

I mentioned in another thread that I was trying to expand the number of possible metals in my college's qualitative analysis lab. I was amazed to find that copper hydroxide did indeed redissolve in excess sodium hydroxide- I've used hydroxide to precipitate copper dozens of times, and have never seen that happen before.

njl - 10-12-2020 at 09:45

Can anyone shed some light on what "FG" means in the linked paper? Also, no developments on my previous reply so any ideas would be appreciated.

https://sci-hub.st/https://doi.org/10.1021/ol8004326

B(a)P - 10-12-2020 at 11:06

Quote: Originally posted by njl  
Can anyone shed some light on what "FG" means in the linked paper? Also, no developments on my previous reply so any ideas would be appreciated.

https://sci-hub.st/https://doi.org/10.1021/ol8004326


Functional group?

Texium - 10-12-2020 at 11:54

I tried searching the paper for FG, F.G., F G etc and nothing came up... can you quote the sentence in question?

arkoma - 10-12-2020 at 17:10

Had to have been drunk. Thought I posted here about the propyl and formate esters I've been fooling with last few daze.

*edit* n-propyl alcohol and salicylic acid cooking as I type

*edit* propyl formate ain't real nose friendly

[Edited on 12-11-2020 by arkoma]

[Edited on 12-11-2020 by arkoma]

AvBaeyer - 10-12-2020 at 17:41

njl,

FG does indeed mean "functional group." The abstract graphic is not all the clear so I can see your confusion.

AvB

Purification of 925 sterling silver

B(a)P - 10-12-2020 at 17:59

18.58 g of 925 sterling silver (SS) was added to a solution of 40 mL HNO3 70%, 20 mL H2O2 and 60 mL dH2O.
The solution was left for 48 hours, at which point undissolved SS was still present and so mild heat was applied until all silver was in solution.
A slight excess (11 g) of NaCl dissolved in 50 mL of dH2O was added to the silver nitrate solution and a thick white precipitate of AgCl immediately formed.
The AgCl was then vacuum filtered in a Buchner funnel and rinsed with dH2O until the AgCl was perfectly white.
The AgCl was then added to 100 mL of dH2O and 10 g of NaOH was added slowly while stirring with a glass rod. After 5 minutes of stirring a few white flecks still persisted so another 5 g of NaOH was added under stirring until a consistent dark precipitate of Ag2O was observed in the reaction vessel.
The precipitate was rinsed once with dH2O then a further 400 mL of dH2O was added prior to adding 150 g of dextrose slowly and in portions while stirring.
No reaction was observed so the solution was heated to boiling while stirring.
After 5 minutes the precipitate started to become lighter in colour and after a further 5 minutes all traces of dark Ag2O precipitate had been converted to silver powder.
The silver powder was rinsed 10 times with dH2O until neutral with litmus.
The powder was then dried and weighed.
Total mass of silver powder was 17.82 g which corresponds to a yield of 101%. Presumably the SS was a little more than 92.5% silver.

arkoma - 11-12-2020 at 13:17

propyl salicylate smells pretty good

Bedlasky - 11-12-2020 at 14:21

Arkoma: How propyl salicylate smells like? I plan to make propyl benzoate or salicylate, I am not sure which one I'll pick.
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