Neal
Hazard to Others
Posts: 116
Registered: 24-12-2021
Location: Chicago, IL, USA.
Member Is Offline
|
|
Colored inorganic liquids.
In the world of inorganic liquids, at room temperature, the vast majority of colored are yellow, especially pale yellow.
Other than that, then we have.
SCl2
S2Cl2
VCl4 (dark red)
CrO2Cl2
And Mn2O7
Any reasons why that is? The web is full of info on KMnO4 is purple, but I don't find any for these guys. But as a newbie question, what does crystal
field theory explain why KMnO4 is purple, and what does ligand theory explain why KMnO4 is purple, assuming they are 2 different reasons.
Would love to see a naturally blue colored inorganic liquid. And purple too if we can find something more stable than that purple acid.
|
|
|
Sulaiman
International Hazard
Posts: 3558
Registered: 8-2-2015
Location: 3rd rock from the sun
Member Is Offline
|
|
You could start here :
https://en.wikipedia.org/wiki/List_of_inorganic_pigments
CAUTION : Hobby Chemist, not Professional or even Amateur
|
|
|
Texium
|
Thread Moved
26-1-2023 at 17:22 |
DraconicAcid
International Hazard
Posts: 4278
Registered: 1-2-2013
Location: The tiniest college campus ever....
Member Is Offline
Mood: Semi-victorious.
|
|
KMnO4 is purple due to charge transfer electron transitions, in which an electron goes from an orbital which is mainly on a ligand to one mainly on
the metal. CFT and ligand theory only deals with d-d transitions. Same with the colours of chromate, dichromate, VO2(+), PbI2, etc.
Isn't ozone a blue liquid at very low temperatures?
I think N2O3 is also blue when liquid.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
|
|
|
j_sum1
Administrator
Posts: 6218
Registered: 4-10-2014
Location: Unmoved
Member Is Offline
Mood: Organised
|
|
Nice list. I did not know that existed.
However, the OP was specifically looking at liquids.
Chromyl chloride comes to mind.
As does bromine, iodine, and various interhalogen compounds.
Chemical Force just released a lovely video on vanadium oxychloride which is a lovely lemon yellow and a bit fresher and cleaner-looking than
chlorine.
I think stating that they tend to be shades of yellow is a bit of an oversimplification.
As for the theory behind the colours -- that is a bit beyond me. I mean I know it is all to do with electron transitions and interaction with
photons. But calculating the specifics is a bit of a niche area and I have not got the background for this.
|
|
|
DraconicAcid
International Hazard
Posts: 4278
Registered: 1-2-2013
Location: The tiniest college campus ever....
Member Is Offline
Mood: Semi-victorious.
|
|
Quote: Originally posted by j_sum1  |
As for the theory behind the colours -- that is a bit beyond me. I mean I know it is all to do with electron transitions and interaction with
photons. But calculating the specifics is a bit of a niche area and I have not got the background for this. |
I JUST TOLD you! It's a charge-transfer transition!! Why don't you ever listen????
Sorry- my engineering students have me irritable.....
But seriously, to get a coloured compound, you either need conjugation, partially-filled d orbitals, or a way to have an internal redox reaction.
So things like SiCl4, SnCl4, PbCl4, and TiCl4, while liquid, won't have colours.
VCl4 has one d electron, and should be coloured- this would also be true of CrCl5, VBr4, NbCl4, TiBr3, CrOCl3, etc, if such compounds existed and
would be liquids. I seem to recall that anhydrous copper(II) nitrate is a liquid at room temp, with a very weird structure.
CrO2Cl2, VOCl3, Mn2O7 have d-zero configurations, but have the metals in high oxidation states, so they can have internal redox reactions, giving you
charge-transfer transitions. I'd expect CrO2F2, VOF3, VOBr3, MnO3Cl and MnO3F to also be coloured, if such compounds exist. I know CrO2Br2 doesn't
exist- attempts to make it just give Br2.
Conjugation is mostly important for organic things, but that's why ozone is blue and oxygen isn't. Alas, it tends to detonate when liquid.
Now, none of what I've said applies to Br2, ICl, or BrCl, al of which are liquid at or near room temp, and are colourful. Which just goes to show
that I don't know everything.
[Edited on 27-1-2023 by DraconicAcid]
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
|
|
|
Bedlasky
International Hazard
Posts: 1219
Registered: 15-4-2019
Location: Period 5, group 6
Member Is Offline
Mood: Volatile
|
|
| Quote: Originally posted by DraconicAcid | | Quote: Originally posted by j_sum1 |
As for the theory behind the colours -- that is a bit beyond me. I mean I know it is all to do with electron transitions and interaction with
photons. But calculating the specifics is a bit of a niche area and I have not got the background for this. |
I JUST TOLD you! It's a charge-transfer transition!! Why don't you ever listen????
|
But he told that he doesn't know how to CALCULATE what colour some particular compound would have. From j_sum post is pretty clear that he knows about
charge-transfer transition.
NbCl4 and TiBr3 exist, they are very dark violet solids.
They exist, Chemical force made recently video about MnO3F. Woelen also did experiment with MnO3F, but green vapor isn't clearly visible in his
experiment (probably due to low concentration).
https://www.youtube.com/watch?v=a_hiwU67HpM&t=241s
https://woelen.homescience.net/science/chem/exps/KMnO4+NaF+H...
MnO3Cl also exist but it is highly unstable at room temperature.
https://woelen.homescience.net/science/chem/exps/HSO3Cl/inde...
CrO2F2 is in the pure state red coloured solid
https://en.wikipedia.org/wiki/Chromyl_fluoride
https://woelen.homescience.net/science/chem/exps/volatile_ch...
VOF3 is orange solid:
https://en.wikipedia.org/wiki/Vanadium(V)_oxytrifluoride
VOBr3 is dark red liquid which can be made from VOCl3 and HBr or by bromination of V2O3-carbon mixture:
https://sci-hub.ru/https://pubs.rsc.org/en/content/articlela...
https://sci-hub.ru/https://pubs.rsc.org/en/content/articlela...
CrO2Br2 does exist, it's just unstable at room temperature. Strangely enough it doesn't resemble CrO2Cl2 and CrO2F2 in colour, CrO2Br2 is purple.
https://chemistry.stackexchange.com/questions/95286/why-do-c...
|
|
|
j_sum1
Administrator
Posts: 6218
Registered: 4-10-2014
Location: Unmoved
Member Is Offline
Mood: Organised
|
|
| Quote: Originally posted by Bedlasky | | Quote: Originally posted by DraconicAcid | | Quote: Originally posted by j_sum1 |
As for the theory behind the colours -- that is a bit beyond me. I mean I know it is all to do with electron transitions and interaction with
photons. But calculating the specifics is a bit of a niche area and I have not got the background for this. |
I JUST TOLD you! It's a charge-transfer transition!! Why don't you ever listen????
|
But he told that he doesn't know how to CALCULATE what colour some particular compound would have. From j_sum post is pretty clear that he knows about
charge-transfer transition.
|
I think the snark was tongue in cheek. It is necessary to vent occasionally when your students are frustrating.
No. I can't calculate. And to be accurate, I only know the mechanism of colour observance in broad terms: enough to guess when a compound may be
coloured, but not enough to guess what the coliur may be.
I teach my students that transition metal compounds are likely to be coloured (except group 3 and 12) and that alternating double bonds give rise to
colour. This is sufficient for 15 year olds and is actually beyond the curriculum in every high school I have encountered.
The late (and accidental) discovery of YInMn blue suggests that deep analysis and calculation is done retroactively rather than predictively. My
impression is yhat the specifics are incredibly complex. But it is an area of interest since vibrant colour is so much a part of what promotes
curiosity in chemistry.
I love conversations like this since I akways learn something. I had never considered internal redox as a source of colour, but it makes sense. Time
to do some thinking about N2O3.
|
|
|
chornedsnorkack
National Hazard
Posts: 521
Registered: 16-2-2012
Member Is Offline
Mood: No Mood
|
|
Simple halogens have a nice progression of spectrum:
F2 condenses at -188 C, and is very pale yellow in thick layers. The maximum is actually around 280 nm
Cl2 condenses at -35, and is stronger yellow
Br2 condenses at +59, freezes at -7 and is brown
I2 contenses at +184, freezes at +114 and is violet
The spectral transitions match. How do the spectra compare?
How about BrCl, ICl and IBr?
|
|
|
DraconicAcid
International Hazard
Posts: 4278
Registered: 1-2-2013
Location: The tiniest college campus ever....
Member Is Offline
Mood: Semi-victorious.
|
|
I looked up ICl, and it looks just like bromine.
(And thanks, j_sum1, for recognizing my brand of humour.)
I did make N2O3 once, crudely, by condensing the fumes from copper-in-nitric acid onto a piece of dry ice. A glorious dark blue. I need to learn to
make ampules....
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
|
|
|
Neal
Hazard to Others
Posts: 116
Registered: 24-12-2021
Location: Chicago, IL, USA.
Member Is Offline
|
|
So are there any substances that are different colors between different states of matter? And let's make at least 1 (freezing point and boiling point)
be near room temperature.
We already know there is a phenomenon called thermochromism where the temperature can change the color, within the same state.
|
|
|
Neal
Hazard to Others
Posts: 116
Registered: 24-12-2021
Location: Chicago, IL, USA.
Member Is Offline
|
|
Hm what is the equivalent for for f-f- transitions? As well as d-f and f-d transitions.
|
|
|
Texium
Administrator
Posts: 4508
Registered: 11-1-2014
Location: Salt Lake City
Member Is Offline
Mood: PhD candidate!
|
|
That’s called
“stupidly complicated stuff that nobody really comprehends besides the people who dedicate their careers to studying it.” If this stuff really
interests you, I’d recommend getting yourself an inorganic chemistry textbook and reading up on CFT and LFT.
|
|
|
Neal
Hazard to Others
Posts: 116
Registered: 24-12-2021
Location: Chicago, IL, USA.
Member Is Offline
|
|
Here's some answers I got elsewhere.
Oh god you’re asking a cursed question (and my PhD specializes in the lanthanides). There are definitely f-related transitions that can happen, but
they aren’t as influenced by geometry as they are for d orbitals since the bonding is essentially ionic in nature by the time you invoke f orbitals
in a meaningful context
-
Energy is very small for lanthanides so you don’t worry too much about ligand field effects. The radial distribution of the 4f orbitals is so small
that it is treated as “core-like” - the interesting photophysics for lanthanides come from electron-electron repulsion and spin orbit coupling,
ligand field effects only weakly modify that picture unlike d-orbitals. Ligand field can be stronger for actinides because 5f has larger radial
distribution comparing to 4f elements - but biggest effects are still the e-e repulsion and spin orbit coupling.
You can get some more info with some googling As a starting point see;
http://www.chem.helsinki.fi/~sundholm/winterschool/lecture_n...
-
(Reply to someone else)
Ligand field is established for f orbitals also, it’s just a much weaker perturbation so not given too much consideration.
-
The other answers are pretty good. I'll add as a PhD who works extensively with the actinides...
If you look at radial extension of orbitals and compare 4f and 5f orbitals, you'll see that 4f orbitals are all closer to the nucleus than the xenon
orbitals before it, while the 5f overlap and extend beyond the radon core.
Any contribution from d-orbitals can greatly increase the transition intensity. f-f transitions are forbidden (as are d-d transitions without ligand
field effects), but f-d transitions are not Laporte forbidden. Meanwhile, since ligand effects are minimal for pure 4f systems, there is very very
little relaxation for these elements.
The early actinides (Pa-Pu) have significant d-f degeneracy, and thus experience much more intense coloration than congener lanthanides. Moreover, as
5f electrons are not fully shielded by the radon core of electrons (due to relativistic effects), they experience much greater ligand field effects
than 4f elements, though still modest by comparison to many transition metals.
In short, ligand/crystal field effects still influence f-electrons, but to a lesser extent than d-electrons.
|
|
|
DraconicAcid
International Hazard
Posts: 4278
Registered: 1-2-2013
Location: The tiniest college campus ever....
Member Is Offline
Mood: Semi-victorious.
|
|
I happened to have some dry ice left over from work, so I made some N2O3 to show my students.
[Edited on 26-3-2023 by DraconicAcid]
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
