Sciencemadness Discussion Board - Pretty Pictures (1)

Quote: Originally posted by Bot0nist

Exactly!

I tried, My camera doesn't have a very good macro mode, and I am having trouble getting it to focus on the tower-like structure. I dumped it out to get a more contrasting backdrop. I will try to get a better camera this summer.

Vlassis - 12-3-2012 at 10:18

Just some caffeine extracted from coffee, having crystallized from an vaporised solution of Dichloromethane

MrTechGuy1995 - 13-3-2012 at 08:02

Quote: Originally posted by Vlassis

Just some caffeine extracted from coffee, having crystallized from an vaporised solution of Dichloromethane

How many grams did you obtain from this?

I'm doing a similar thing, and curious to see how I fair.

thesourcerer - 18-3-2012 at 15:08

Having eventually got my head round dowloading the photos, I thought this should be in the pretty picture section!
Ampoules of crystalline Na & K - for sale (they are over 50 years old)! See under shiney K & Na thread.

Dennis SK - 23-3-2012 at 07:01

High purity rubidium in ampoule.
Melted with a heatblower as close to Rb melting point as possible. This slow melting gives the surface this very nice "wave" look.
This was the first test result

And here a few of the ampoule. (definately one of my better samples in my element collection)

[Edited on 23-3-2012 by Dennis SK]

Dennis SK - 23-3-2012 at 13:32

Some of you might already have seen some of my cesium pictures. But here's one anyway.

Cesium 99,99%+ in ampoule.
Grown with temp. control. Took alot of tries to get this result.
I had absolutely no idea that cesium could make such crystals. So you could imagine the look on my face when I first succesfully made these

Adas - 23-3-2012 at 14:23

Dennis, those pics are beautiful! I wish I had some ampoules like this, too xD

offtopic: May I ask where r u from?

Dennis SK - 23-3-2012 at 14:29

Hello Adas,

Sure, I'm from Denmark

Dennis SK - 24-3-2012 at 04:41

Rubidium ampoule. Now with Cesium ampoules also.

Sorry for the picture spamming

blogfast25 - 24-3-2012 at 06:38

Wonderful.

barley81 - 24-3-2012 at 07:59

All of that cesium reminds me of Greek mythology. Ichor, the blood of the gods, was said to be golden in color and very poisonous to humans. Cesium metal exactly fits that description! And the crystals of Cs are just great!

Endimion17 - 24-3-2012 at 08:14

But caesium isn't especially poisonous. It's pyrophoric.

Dennis SK, great samples. There's a ton of money inside those ampoules...

barley81 - 24-3-2012 at 08:29

Well, if you eat it, it will probably do a lot of harm and corrode your mucosa very badly (not to mention the explosion). I meant cesium the metal, not the compounds.

Sedit - 24-3-2012 at 08:40

Quote: Originally posted by Bot0nist

Take another lenses from an old telescope or something similar and place it in front of your camera lenses, generally this allows you to make your own macro and has worked very well in the past for me to take pictures of small crystals.

Dennis SK - 24-3-2012 at 08:58

Quote: Originally posted by barley81

And the crystals of Cs are just great!

They really are!
I'll try to grow the same type of crystals with rubidium.
But its alot harder to control crystal growth in rubidium, but I'll show you the results here if succesful.

Rubidium has the "funny/interesting" fact that its slightly radioactive. The Rb in the ampoule above gives about 150k Bq, equal to about 4,2 americium fire detectors (calculated)

mr.crow - 24-3-2012 at 13:37

HOLY JESUS!

That must have been expensive. So beautiful

Dennis SK - 25-3-2012 at 12:49

Dennis SK, great samples. There's a ton of money inside those ampoules...

They wasn't cheap thats for sure.
But alkali metals have always been some of my favorite elements. So naturally I looked for the best samples I could possibly find.
It took 5 years and so many inspections of Rb/Cs ampoules.
Trust me, depending on source its relatively easy to find clean samples of Rb/Cs in ampoules below 10 gram. But to find large clean, particle free samples above 10g of especially Cs... Good luck!

For info the purity are the following:
Rb, Certified 99,95% minimum. But possibly even more pure.
Cs, Certified 99,99% minumim.

Glad you guys like the pictures. Once I get a better camera I will take high-end pictures.

[Edited on 25-3-2012 by Dennis SK]

Mailinmypocket - 25-3-2012 at 13:26

Tried a bit of experimenting with Cobalt complexes, I attempted to make trans-dichlorobis(ethylenediamine) cobalt(III)chloride.

The procedure I used instructed(as the before last step) to wash with 6M HCl and vacuum filter, repeating if required until the brown or blue crystals turn green. Mine came out of solution as dark red/orange crystals that did not change color at all despite 4x washes with 6M HCl

Any ideas what happened? Regardless... Pretty pictures

[Edit: Initially in the process, after adding the ethylenediamine to the CoCl2 solution, the instructions are to stir vigourously while heating in order to oxidize the Co(II) to Co(III). I wanted to see what would happen if I stirred slowly with a magnetic stirrer and bubble pure oxygen into the solution at the same time over the course of 40 minutes. I am thinking that maybe in doing this the complex was oxidized further?

Also, some tris(ethylenediamine)nickel (II) chloride made this afternoon

[Edited on 25-3-2012 by Mailinmypocket]

nora_summers - 25-3-2012 at 18:32

oh that purple is such a pretty color!

ItalianChemist - 29-3-2012 at 05:51

Quote: Originally posted by Mailinmypocket

trans-dichlorobis(ethylenediamine) cobalt(III)chloride.

I've synthesized the same complex, but i've obtained a green solid

Endimion17 - 29-3-2012 at 05:58

Quote: Originally posted by Dennis SK

Your camera is perfectly fine. You just need to take an artistic approach. Some black cloth for the background and a diffused light source, together with photography rules should do the trick.

Dennis SK - 29-3-2012 at 12:37

Your camera is perfectly fine. You just need to take an artistic approach. Some black cloth for the background and a diffused light source, together with photography rules should do the trick.

That "Camera" is actually a mobile phone camera

I didnt think much of artistics when I took the pictures. I just melted the Rb/Cs, thought it looked really nice and tried to catch it.

I have a friend who work as photographer. He has a "photo box" or a "light diffuser box" I think its called. He will help me take very good pictures. Those pictures will also be used instead of some of my current wikipedia pictures.

Really looking forward to show even better pictures of completely oxide free Rb/Cs. Of course, one can only see the true beauty "In real life"

dann2 - 29-3-2012 at 15:39

There are some mighty pretty pictures here.
Two Sodium Perchlorate shards.
They are the same shard. One was photographed with a polarizing filter on camera.

What colour is Cobalt Perchlorate anyone know?

[Edited on 29-3-2012 by dann2]

Attachment: shards in cell.zip (337kB)
This file has been downloaded 434 times

barley81 - 29-3-2012 at 16:05

Most cobalt (II) salts are red-purplish. Cobalt (II) perchlorate is likely to be similar in color to the nitrate or sulfate.

Such a large shard of a single salt! Isn't it really fragile? How did you make a crystal like that?

dann2 - 30-3-2012 at 12:57

A saturated solution (3.8 liters) of Sodium Perchlorate was cooled from 50°C to around 10°C and lots of shards formed (see zip attached to last post).
The shards are perhaps two mm thick and not that delicate. If you were to let them fall on a hard surface they would break.
There are more shards in the thread "More on PbO2 Anodes" stuck to an LD Anode.

I have some Barium Perchlorate and Cobalt Sulphate. I might try making some Cobalt Perchlorate by adding the the two together. Ba Sulphate will ppt as it is insoluble leaving me with a solution of Co Perchlorate.
Dann2

ItalianChemist - 1-4-2012 at 02:18

Chemiluminescence of Ru(bipy)3(BF4)2

Satan - 1-4-2012 at 11:18

Quote: Originally posted by ItalianChemist

Chemiluminescence of Ru(bipy)3(BF4)2

Did you make a OLED with this complex?

White Yeti - 1-4-2012 at 12:24

@ItalianChemist, that picture is impressive! How did you take it? I avoid taking pictures of things in the dark because they always come out awful. That picture might even be good enough for a screen saver considering the size and quality.

Eddygp - 1-4-2012 at 13:04

Just have been looking at those GIANT and INCREDIBLE samples of Cs and Rb and... I took a look at my homemade sodium silicate 'crystals', and left the camera.

ItalianChemist - 2-4-2012 at 07:03

@ItalianChemist, that picture is impressive! How did you take it? I avoid taking pictures of things in the dark because they always come out awful. That picture might even be good enough for a screen saver considering the size and quality.

Thank you!
I used a normal camera with macro function!
@Satan I haven't done OLED yet!

[Edited on 2-4-2012 by ItalianChemist]

ItalianChemist - 3-4-2012 at 12:21

Other pictures

1. Europium nitrate hydrate under UV-A light!

2.Ru(bipy)3(BF4)2 under UV-A light!

3.Melted iodine

kristofvagyok - 6-4-2012 at 14:07

Hi everyone,

This is my first post on the board so let's start it: I work in an organic chemistry lab somewhere in Europe and I am also a freelance photographer so I got some pictures...
They are not special, no Europium, Ruthenium or any special. just simple organic and inorganic chemistry, but I hope You like them:

More pictures here: http://labphoto.tumblr.com/

-Kristof

White Yeti - 6-4-2012 at 14:28

Hey Kristof.

Nice pictures! They're quite stunning, what kind of camera did you use? Mind if I ask what these pictures are of?

Welcome to the forum!

kristofvagyok - 6-4-2012 at 14:41

what kind of camera did you use? Mind if I ask what these pictures are of?

I use Nikon D5000 with wide variety of lenses.

The pictures are from my current research, what is about nitrogen containing heterocyclic ring formations... And some else thing, what I do in the lab (last picture is inorganic chemistry).

That cute red thing in the breaker (what is an awesome dye and impossible to remove from your skin

) was a nitro-phenylpyvuric acid ester synthesis from nitrotoluene, dialkly oxalate and some sodium alkoxyde. When this material contacts any base it turns blood red, otherwise it is yellow.

The separating funnel contained some ether and some phenylpyvuric acid in the water what didn't reacted in an other kind of reaction.

And the last picture is molten silver under some borax with a methane/oxygen torch.

On the blog there is a lot other picture (updating daily, or as often as I could

)

[Edited on 6-4-2012 by kristofvagyok]

Adas - 7-4-2012 at 00:14

That methane/oxygen torch is just amazing! The temperature must be extremely high!

Bot0nist - 7-4-2012 at 06:53

Very nice kristofvagyok! Welcome to SciMad. Those pictures are beautiful. Glad your here.

Mailinmypocket - 9-4-2012 at 17:01

Finally did the Belousov-Zhabotinsky reaction in a petri dish

too many time lapse pictures to attach, but amazing to watch happen... hypnotizing! It repeated itself, sort of like a beating heart with the same recurring pattern. If i swirled the dish new patterns appeared and started the same sequence. Eventually it slowed down and became clear ~15 mins

The ferroin solution used was prepared by the reaction of 1,10- phenathroline hydrate (0.88g) and ferrous sulfate (0.35g) in 50 ml distilled H2O.

*Edit* The pictures were attached in reverse order, it goes red to blue.

See section 5.2.3 for procedure here

[Edited on 10-4-2012 by Mailinmypocket]

White Yeti - 9-4-2012 at 17:28

Those are quite nice! For everyone's viewing pleasure, I'll attach a video of the process (not mine).
<iframe sandbox width="560" height="315" src="http://www.youtube.com/embed/3JAqrRnKFHo" frameborder="0" allowfullscreen></iframe>

Mailinmypocket - 9-4-2012 at 18:04

Thanks! That was my first try, Ive seen that video many times too while trying to figure out the ideal mix. Maybe more concentrated ferroin and/or a catalyst... I didn't add any Ce or Mn, time to tweak the formula! I am going to try with a methyl orange formulation... Hopefully it is as good if not better

[Edited on 10-4-2012 by Mailinmypocket]

inspector071 - 9-4-2012 at 22:11

I've been synthesizing some energetics recently, and in particular, working with recrystallization for stability. It's given me an opportunity to drag out the microscope and check out some crystals.

Here are some crystals of picric acid recrystallized from boiling water at 100x magnification.

These are quite flat crystals of ETN, recrystallized from methanol, also at 100x magnification. This is what the batch of ETN looks like.

This is a crystal of RDX, created via the direct nitration of hexamine from pure nitric acid, and recrystallized from acetone, yet again at 100x magnification. The crystals are a bit bigger than table sugar.

Endimion17 - 10-4-2012 at 03:29

kristofvagyok, those photos are amazing. It really is great what proper lightning, a sense of perspective from a photographer's eye and macro focusing can do with something that seems so usual.

I was browsing through your tumblr page and noticed you described bromine as having a really bad smell. Why is that?

DJF90 - 10-4-2012 at 04:29

Because it does. Bromine comes from the greek Bromos, which means stench. At least if I remember correctly.

White Yeti - 10-4-2012 at 06:43

Isn't bromine a halogen? A volatile one too... Not only do halogens have a smell, but they also destroy your mucous membranes. That's why bromine is usually chilled before use and handled with utmost precaution.

mr.crow - 10-4-2012 at 07:00

Hi everyone,

This is my first post on the board so let's start it: I work in an organic chemistry lab somewhere in Europe and I am also a freelance photographer so I got some pictures...
They are not special, no Europium, Ruthenium or any special. just simple organic and inorganic chemistry, but I hope You like them:

More pictures here: http://labphoto.tumblr.com/

-Kristof

Incredible!!

1 L of bromine, that thing is going to weigh 3 kg! I have always wanted to hold that much Br to feel the weight. Its also going to destroy the bottle cap in no time.

Whats behind the bromine? also looks dangerous

kristofvagyok - 10-4-2012 at 07:50

Quote: Originally posted by mr.crow

The main problem with the bromine is to pour it out from the bottle and weight it, because of the fumes what comes out from it, smells really bad and causes not so good things. The positive thing is the fume hood helps a lot(:

This amount of bromine was used to make dimethoxy alkyls from alkenes by one of my friends.
The scheme was like this: RHC=CHR + Br2 = RHBrC-CBrHR
RHBrC-CBrHR + 2R'OH = RH(OR')C-C(OR')HR + 2HBr

Easy reaction with high yields even in this amounts.

And the cap of the bromine is coated with teflon, so it wont be damaged, I hope(:

[Edited on 10-4-2012 by kristofvagyok]

Endimion17 - 10-4-2012 at 10:02

Quote: Originally posted by DJF90

Because it does. Bromine comes from the greek Bromos, which means stench. At least if I remember correctly.

So, the argument is essentially: "It smells bad because it's named after the Greek word for stench."
That's a fallacy.

Bromine does not smell disgusting at all. Irritating, yes. Disgusting and smelly, no. It resemblies other halogens, nitrogen dioxide and ozone, all being strong oxidizers. Of course, with increasing concentration in unit volume of air, it gets more and more irritating, until it becomes agonizing. But it doesn't have a stench.
Smelly things include impure carbon disulphide, hydrogen sulphide, scatol, etc. Bromine isn't smelly and disgusting.

I've never read any remotely solid argument on why did bromine get that incredibly stupid name. Either something's wrong with my olfactory nerve and/or brain, or the one from the guy proposing the name.

inspector071 - 10-4-2012 at 10:07

So it has a choking odor like chlorine, but not a disgusting stench from something like putrescine?

kristofvagyok - 10-4-2012 at 10:18

Quote: Originally posted by inspector071

So it has a choking odor like chlorine, but not a disgusting stench from something like putrescine?

It is a bit choking, but not that bad as sulfur-dioxide, nitrosyl-chloride or simply chlorine is much worse.

The odor of bromine is really characteristic, similar a bit to chlorine, but not that "harsh", and completely different from any amine (like pustrescine) but I think the odor of the chemicals is somewhere in an other topic, this is about the pictures(:

[Edited on 10-4-2012 by kristofvagyok]

inspector071 - 10-4-2012 at 10:25

Quite so, and the pictures you have posted are fantastic! I checked out the ones on your tumblr too. Really makes me want to invest in a nice DSLR to take pretty pictures with. For now, my iPhone + microscope will have to suffice.

White Yeti - 10-4-2012 at 11:42

I've never read any remotely solid argument on why did bromine get that incredibly stupid name. Either something's wrong with my olfactory nerve and/or brain, or the one from the guy proposing the name.

Smell can be subjective to a certain extent (although I think everyone will agree that H2S and mercaptans smell abominable). Maybe the scientist who first isolated it had an acute sense of smell...

Stupid things like that happen sometimes. Example: Fahrenheit wanted to base his temperature scale on human body temperature (100F = normal human body temperature), but he probably had a cold that day because normal body temperature is 98F

To give another example, we based all our formulas for electricity on the wrong charge. Now we have the conventional/true current ambiguity.

ItalianChemist - 10-4-2012 at 13:22

I like very much smell of H2S!
Bromine is like Cl2! I like it because it isn't as stifling as Cl2 smell!

dann2 - 10-4-2012 at 15:03

...........
The main problem with the bromine is to pour it out from the bottle and weight it, because of the fumes what comes out from it, smells really bad and causes not so good things. The positive thing is the fume hood helps a lot(:.........

[Edited on 10-4-2012 by kristofvagyok]

When pouring a liquid from a bottle to another container a trick that usually works nicely it to obtain a glass rod and hold it against the opening of the bottle and use it to guide the liquid into the other container. Liquids usually 'jump' to the rod as opposed to dribbling down where you do not want them to dribble.
When pouring Bromine from one bottle to another the glass rod trick will NOT work. The Bromine refused to jump to the rod. In fact it's as if the rod repells the Bromine. A mess when you have no fume hood.

White Yeti - 10-4-2012 at 15:54

Quote: Originally posted by dann2

When pouring Bromine from one bottle to another the glass rod trick will NOT work. The Bromine refused to jump to the rod. In fact it's as if the rod repells the Bromine. A mess when you have no fume hood.

I'm sorry to be brusque, but DUH! Bromine does not adhere to glass the same way as other liquids (polar liquids notably). I haven't worked with bromine or benzene, but I'm willing to bet that benzene will not trickle down a glass rod either because of its polarity, or lack thereof.

Mailinmypocket - 10-4-2012 at 18:52

Quote: Originally posted by Mailinmypocket

Thanks! That was my first try, Ive seen that video many times too while trying to figure out the ideal mix. Maybe more concentrated ferroin and/or a catalyst... I didn't add any Ce or Mn, time to tweak the formula! I am going to try with a methyl orange formulation... Hopefully it is as good if not better

[Edited on 10-4-2012 by Mailinmypocket]

Tried the Belousov-Zhabotinsky again, this time by adding methyl orange along with ferroin to the mixture. The waves were much more pronounced and more frequent and organized. The first time the petri dish more or less went through a rapid color change, this time it was a slow creeping wave formation that repeated itself every 15 seconds or so.

This is a fun reaction to do and try modifying the concentrations on different reagent and time/monitor the effect of the changes.

This is the latest try with methyl orange:

[Edited on 11-4-2012 by Mailinmypocket]

RonPaul2012 - 11-4-2012 at 05:33

I have wanted to contribute to this forum in some way for quite some time.

Here are some pictures I took of a flaming boric acid denatured ethanol mix.

Enjoy

kristofvagyok - 11-4-2012 at 12:13

Quote: Originally posted by RonPaul2012

Here are some pictures I took of a flaming boric acid denatured ethanol mix.

Enjoy

When I made the same experiment and the ethanol was with high excess (95% EtOH and just 5% MeOH) the flame was mainly yellow and had just a slight green tint, just with pure MeOH was a really green flame produced. But these results are fantastic! Congrats!

Today I tried our with some benzene and it didn't "tricked" down the glass rod. And I won't try it our with bromine, ok?

ItalianChemist - 11-4-2012 at 12:46

Hi!
I've made this nice compound mixing FeCl2, 2,2'-Bipyridine and NaBF4
I think that this is Fe[bipy]3(BF4)2, but I'm not sure

Mailinmypocket - 11-4-2012 at 15:24

Quote: Originally posted by ItalianChemist

Hi!
I've made this nice compound mixing FeCl2, 2,2'-Bipyridine and NaBF4
I think that this is Fe[bipy]3(BF4)2, but I'm not sure

Nice job! It's so satisfying to see big crystals in the bottom of your vessel/filter paper etc! I'm not certain of what compound you made either but it looks cool

RonPaul2012 - 11-4-2012 at 15:52

Quote: Originally posted by RonPaul2012

Here are some pictures I took of a flaming boric acid denatured ethanol mix.

Enjoy

Today I tried our with some benzene and it didn't "tricked" down the glass rod. And I won't try it our with bromine, ok?

Thanks , I'm glad you liked it

ItalianChemist - 12-4-2012 at 02:43

Quote: Originally posted by Mailinmypocket

Nice job! It's so satisfying to see big crystals in the bottom of your vessel/filter paper etc! I'm not certain of what compound you made either but it looks cool

Thank you!
They aren't very big, in the photo they are enlarged!

Endimion17 - 12-4-2012 at 03:15

This is not something I've exactly made, but it's a chemical compound nonetheless, and it's pretty.

link to all sizes

It's vanadinite, Pb5(VO4

3Cl.
Although it does look somewhat edited, it really does look like this in strong sunlight. The back of the crystals have an opaque deposit, so the light is reflected a lot more.

Squall181 - 12-4-2012 at 07:45

Thanks Endimion that's a beautiful picture and worthy of being my desktop background

dann2 - 13-4-2012 at 04:32

Quote: Originally posted by White Yet

The 'rod trick' can be very useful when transferring liquids.
Benzene runs down a rod OK.
Perhaps the very dense Br has very high surface tension, like Hg. It's as if the glass rod repells it.

plastics - 13-4-2012 at 12:36

A solution of perylene in acetone fluorescing under UV light on a watch glass, and some acting as a fluorophore in a bottle demonstrating chemisluminescence with TCPO.

Credit to garage chemist for pointing me to a synthesis of perylene starting from 2-naphthol on versuchschemie.de

The chemiluminescence was so intense it could be seen in daylight

garage chemist - 13-4-2012 at 12:43

Hey plastics, didn't you want to describe how you made the 2,4,6-trichlorophenol? Doesn't have to be a detailed procedure, just tell me how you did it, and how you determined that the product was the desired substance and whether it was pure.

Dennis SK - 20-4-2012 at 04:43

High purity 99,999%+ Niobium metal crystals

nyll - 20-4-2012 at 09:18

This is not something I've exactly made, but it's a chemical compound nonetheless, and it's pretty.

link to all sizes

It's vanadinite, Pb5(VO4

3Cl.
Although it does look somewhat edited, it really does look like this in strong sunlight. The back of the crystals have an opaque deposit, so the light is reflected a lot more.

Looks photoshoped, but great still.

dann2 - 20-4-2012 at 12:48

@Dessik SK

It the Nb you own, did you grow them yourself?

Dennis SK - 20-4-2012 at 13:50

Quote: Originally posted by dann2

@Dessik SK

It the Nb you own, did you grow them yourself?

Hello.

Yes It's from my collection.

No, It was made in Russia many years ago. I dont think this type of Niobium is made anymore.

barley81 - 25-4-2012 at 17:31

Iodine ingot made by melting iodine under concentrated sulfuric acid. Credit goes to woelen for suggesting this.

Suspended iodine.

[Edited on 26-4-2012 by barley81]

hkparker - 25-4-2012 at 17:50

Quote: Originally posted by barley81

Iodine ingot made by melting iodine under concentrated sulfuric acid. Credit goes to woelen for suggesting this.

Oh cool, I'll have to try this.

Here's some Tris(acetylacetonato)iron(III) I made in chem lab:

Not a great photo but its pretty cool in person.
<a target="tab" href="http://sciencemadness.org/scipics/IMG_20120425_183748.jpg"><img height="237" width="372" src="http://sciencemadness.org/scipics/IMG_20120425_183748.jpg"></a>

[Edited on 26-4-2012 by hkparker]

barley81 - 25-4-2012 at 18:18

What a beautiful compound! I wonder where you can get acetylacetone, and if other transition metals will make such lovely crystalline complexes.

hkparker - 25-4-2012 at 18:37

In lab other groups did cobalt, copper, and chromium. They all looked pretty cool, but iron was probably the prettiest due to its reflective properties. Copper was a pale blue but I remember it being a dull powder. Chromium was green but very dark, cobalt was very dark.

I haven't looked for acetylacetone but I'm sure its around.

Teen Chemist - 28-4-2012 at 08:40

Some crude triethyl borate a relative of trimethyl borate the camera doesnt do the green flame justice.

White Yeti - 28-4-2012 at 09:47

Since everyone seems to like green, I'll make a humble contribution along these lines. The emission spectrum of copper:

Copper acetate inserted into the flame of my home made isopropyl burner. The picture's a little fuzzy because of the low light conditions.

Dennis SK - 30-4-2012 at 06:11

Nice flametests!

I have a picture of a cesium flametest I made a few years ago. Will find it and post it soon (its on another computer)

Cesium can with some effort produce a beautiful and quite unique light-blue or "Skyblue" flame

Endimion17 - 30-4-2012 at 07:17

Guys, it seems to me almost none of you really put any effort in making a nice photograph. Those are just snapshots.
I understand, those flames/chemicals are very pretty, but the photos aren't. I don't see their beauty captured in an image. The thread is about pretty pictures, and the theme are experiments, compounds, etc.

Please don't say your camera is bad. That's not an excuse. People do amazing things with shitty cameras all the time, and I'm not talking about the Instagram crap.
Come on, put some effort in it.

kristofvagyok - 30-4-2012 at 08:44

Come on, put some effort in it.

You meant something like this?

Endimion17 - 30-4-2012 at 08:54

Exactly.

My complaint wasn't pointed to your work, which is really great, as I've mentioned earlier. Some photos you've taken could be used for making fancy chemistry textbooks.

White Yeti - 30-4-2012 at 12:35

Granted, I'm not a talented photographer, and I'm not going to put a blame on my camera. However, there are some phenomena that are just difficult to capture. I challenge you to take a picture of the flame test of an element. You have to remember that a flame test is a fleeting moment in time and that the flame you choose to use is in constant motion. Although you can choose to use a higher shutter speed and get something like what kristof got, you can also choose a regular shutter speed and risk getting a fuzzy picture (which is what I got). The problem with a short exposure is that you upset the balance between the brightness of whatever phenomenon you choose to take a picture of, and its background.

I'll try again and see if I can take a better picture of a flame test of copper.

[edit]

I'm not sure if it's better or worse, I'll leave it up to you to decide. I tried to use macro, but it doesn't work with a moving target.

[Edited on 4-30-2012 by White Yeti]

Endimion17 - 30-4-2012 at 14:08

I'm well aware of the technical difficulties, however, that wasn't my point. What I meant was to work on the details. For example, why are we seeing parts of your ceramic wall tiles? Get it now?

Photography is an art, too. If you want to focus the viewer onto a chemical reaction, don't distract him with things like tiles and dirt spots (that's what photo editing software is for).

As for the capturing of the dynamics of the flame, copper gives a very luminous blue-green flame. I'm sure it could be caught with a digital camera in all its glory if you use 1/800 s and ISO-200 or 400.
Yes, digital macro shooting can be a pain in the ass. That's why there's something called focus locking. You put your finger where you expect the flame to be, press the first step of the button, and then while holding it, start the flame and then press the whole way down. Timer can help.
Dimming lights helps a lot. And if you want some equipment to be visible, you can always take two photos without moving the camera, and then merging them in a software. One lit normally, one with the flame.

Here's an example of a flame less luminous than copper flame.

It was probably 1 ml of acetone sprinkled towards the camera in dark conditions. I think the shutter speed was 1/800 s.
It's one of many similar photos. Here's butane flame deflagration, edited for lulz.

Use your imagination.

cyanureeves - 1-5-2012 at 04:09

damn! dude that is like time-life quality, heck even an old donkey will come out as a pretty picture.nice pic and good thing it's about chemistry not an old mule.

[Edited on 1-5-2012 by cyanureeves]

[Edited on 1-5-2012 by cyanureeves]

Arthur Dent - 1-5-2012 at 04:18

Our friend Woelen has a nice little primer on his website entitled "Guidelines for successful chemistry-related photography"... very interesting and simple techniques for photography using proper lighting of the subject and the background.

Robert

BaCl2

sargent1015 - 1-5-2012 at 19:10

Fresh barium chloride crystals! Just made them last night and they are slowly crystallizing out

Copper II acetate crystals with what appears to be copper carbonate residue on them. Anyone have a solution to get rid of that? More acetic acid? The crystals are gorgeous, minus that small flaw

btw, sorry about the poor picture quality, trying to get within the post-able limits is tougher than I expected!

MrHomeScientist - 3-5-2012 at 05:11

Quote: Originally posted by sargent1015

Copper II acetate crystals with what appears to be copper carbonate residue on them. Anyone have a solution to get rid of that? More acetic acid? The crystals are gorgeous, minus that small flaw

If it is indeed copper carbonate, I should think a dip in some acid would be get rid of it easily. The bubbles produced should help knock the pieces off your crystals, too. Try shaking a crystal around a bit in acetic acid and see if that helps. If not, go a step higher and try hydrochloric - that should dissolve the carbonate more quickly and shouldn't (I think) mess with your acetate crystals any. Of course, they will shrink slightly as they dissolve in whatever solution you dip them in.

Crystals

sargent1015 - 3-5-2012 at 07:09

Thanks! I will try that out and post pictures of the results. If all else fails, I'll boil them down and make a new batch.

Crystal growing is exciting even if you have to destroy some to make more

Endimion17 - 3-5-2012 at 08:47

For god sake, hydrochloric acid is going to ruin them. Stronger acid pushes the weaker one's anion out.
Just sprinkle them with concentrated acetic acid. The efflorescence and turning to carbonate is normal if they're left in the open atmosphere. They should be kept in a sealed container, slightly damp from acetic acid.

sargent1015 - 3-5-2012 at 09:25

Alright, I will try that instead!

Thanks

(Pictures still to come)

bob800 - 3-5-2012 at 13:24

Quote: Originally posted by Squall181

@ Woelen Yes those where made by electrolysis of copper sulfate solution with a copper anode and copper cathode.

Those crystals where grown in the course of a week with a pretty dilute solution of copper sulfate.
A 12 volt adapter that could supply 1 amp of current was used; I did not have a multimeter on hand to take exact measurements.

The rig is pretty simple to build: a 20 oz soda bottle is taken and the top cut off, then the top is inverted back into the bottom piece of the bottle and taped with the cathode down in the bottom of the bottle. Next a filter is made using paper towel or something similar and is placed into the inverted cone. A stiff piece of wire is used at the top as a connection from which a copper anode is hung. The power is switched on and you have to be really patient, because it is a slow process, but I think if you use a more concentrated solution it should be faster.

Here is a link to where I got the idea: http://youtu.be/kq1W-QdMsWQ
Skip to the middle to see the how the cell is constructed.

I tried replicating your experiment; however, when I raised the voltage to 12 volts, the copper began coating itself in some sort of black oxide. Lowering the voltage to ~4 volts solved this problem, but the copper only formed into a sort of amorphous sponge... Is there something I'm missing? It took a day or two until any visible copper started depositing on the cathode, so I don't think high current caused the oxide.

watson.fawkes - 3-5-2012 at 15:54

Quote: Originally posted by bob800

[...] I raised the voltage to 12 volts [...] Lowering the voltage to ~4 volts [...] Is there something I'm missing?

The kind of reactions that occur depend upon the voltage. It all has to do with the potential across the polarization region near the electrodes. Reactions that are ordinarily not favored can start happening when that voltage goes up.

bob800 - 3-5-2012 at 16:18

But would it be possible that the lower voltage is resulting in the poor copper formation (i.e. forming as a sponge instead of crystals)? Or perhaps it's due to the fact that only a trickle of current is flowing through the cell... I'm using a fairly concentrated solution of CuSO4 with some added sulfuric acid for conductivity–I guess I could just keep adding more until I get a larger current. Still, I don't understand how Squall181's crystals weren't covered in oxide (he/she did use 12 volts).

m1tanker78 - 3-5-2012 at 18:17

Various pictures of Calcium Acetate:

After gently heating a solution of acetate for a few hours, the solid begins to form at the surface:

Closer look:

Calcium acetate solution allowed to slowly evaporate produces different crystals:

Tank

kristofvagyok - 4-5-2012 at 13:51

My new fav picture, soon I will make a hi-res print from this to my wall...(:

Two parallel reactions in two 4 liter flask equipped with reflux condensers, dropping funnels, placed on IKA magnetic stirrer/hotplates. Hope that they will work…

Chemistry is awesome(:

tumblr_m3gm668E701rszmkso1_500.jpg - 367kB

sargent1015 - 4-5-2012 at 13:53