Sciencemadness Discussion Board

why gold is yellow?

emrebjk - 8-5-2006 at 01:02

why gold is yellow.Do you know why?

solo - 8-5-2006 at 05:08

A fast google search indicates that gold a mineral , and a metal has a gold color because that's how it occurs in nature...............as to why? maybe because that's the color of the mineral the nature gave gold once it cools and flows to make the veins mostly found in quartz and as erosion occurs finds itself in rivers , settled in the bottom because of its density...........solo

kazaa81 - 8-5-2006 at 05:12

I think that the answer to this question is like many others:

Gold reflects/absorb some light frequencies that gave it the gold color...it's like benzene, who has no color, but its halo-compounds (chloro-,bromo-benzene) are colored, this shows that maybe benzene has is own color, just that cannot be seen.

12AX7 - 8-5-2006 at 11:01

Something about "plasma frequency" mumbled on Jim Calvert's site.

http://www.du.edu/~jcalvert/phys/copper.htm
Quote:
Copper is often described as a "red" metal, though its actual color is an orange-red of lower intensity, not a bright signal red. It is not a spectral color by any means, but a particular impure one requiring its own name, such as "copper-red." The red color is produced by the density of electrons being insufficient to cause a high plasma frequency, so the shorter wavelengths are not reflected as efficiently as the longer, redder ones. The red color is unique to copper and its alloys.


Huh, that doesn't really make sense.. copper has a mess of electrons, doesn't it?

Stuff like this turns up on Google:
http://scienceworld.wolfram.com/physics/PlasmaFrequency.html

Maybe has to do with the relatively low number of electrons given up per metallic ion?

Tim

BromicAcid - 8-5-2006 at 14:29

Gold, like copper has 1 electron in its s orbital and a full d orbital. The difference in energy between the s and d orbitals in this case is sufficently low so that electrons from the d orbital can be excited (via visible light) and jump into the s orbital. It is this absorption of light to jump the electron that gives the metals their colors.

And from here the absorption is determined by the energy required to jump that electron and is calculated as:

(lambda) = hc/E

[Edited on 5/8/2006 by BromicAcid]

12AX7 - 8-5-2006 at 14:54

No, I don't see that working. If the orbitals are "very close", it would be in the IR band or less (fractional eV energy). Going to a different shell altogether might make the 1.5-3eV visible absorbtion, but that doesn't happen the same in metals as bare atoms. If it were spectral, there should be orange and green and blue metals too, but there aren't.

Tim

guy - 8-5-2006 at 14:54

Gold has a large nucleus and as electrons travel around a large nucleus, it reaches speed near the speed of light. The reletavistic effects makes the electron act heavier and thus it moves closer to the nucleus. This changes lowers the energy barrier between d and s orbitals so it is easier for the electrons to move around.

12AX7 - 8-5-2006 at 15:03

guy, classical mechanics don't apply to atoms. If the electrons were orbiting, they would be accelerated, and an accelerated electron releases energy (synchrotron radiation, more or less). This works for very large n, but closer to the ground state, Schrodinger takes over.

BTW, gold's nuclear size has no bearing whatsoever on chemistry (that is, the electronic behavior). Electron cloud radius is around three orders of magnitude bigger than the nucleus. The particulars of charge have much more bearing: compare tungsten (4+ or 6+) at 60pm to carbide's 260 pm, four times larger despite being carbon being 15 times lighter.

Tim

guy - 8-5-2006 at 15:11

Then maybe I read this site wrong? http://www.madsci.org/posts/archives/may97/862179191.Ch.r.ht...

I know that electrons don't actually orbit around the nucleus, becuase they would have to constantly be accelerating. So I'm confused.

[Edited on 5/8/2006 by guy]

BromicAcid - 8-5-2006 at 15:11

12AX7, your inability to accept my ranting as fact has forced me to look it up:

From Descriptive Inorganic Chemistry 3rd Ed. Geoff Rayner-Canham and Tina Overton; 2003
Quote:
Copper and gold are two common yellow metals, although a thin coating of copper (I) oxide, Cu<sub>2</sub>O, often makes copper look reddish. The color of copper is caused by the filled d-band in the metal being only about 220 kJ<sub>*</sub>mol<sup>-1</sup> lower in energy then the s-p band. As a result, electrons can be excited to the higher band by photons of the corresponding energy range - the blue and green regions of the spectrum. Hence, copper reflects yellow and red. The band separation in silver is greater, and the absorption is in the ultraviolet part of the spectrum. Relativistic effects lower the s-p band energy in the case of gold, again bringing the absorption into the blue part of the visible range, resulting in the characteristic yellow color.

Twospoons - 8-5-2006 at 19:54

Quote:
Originally posted by 12AX7
If it were spectral, there should be orange and green and blue metals too, but there aren't.

Tim


No, as any photon with energy exceeding the energy gap can cause the electron jump. To get a blue metal would require absorption of red and green wavelengths, but not blue - clearly this cannot happen, as "blue" photons are more energetic than red and green and would be absorbed too.

So for gold (=yellow) only blue is absorbed, for copper both blue and green get absorbed.

12AX7 - 8-5-2006 at 20:04

(Argh, why do people have to use kJ/mol, it's so utterly useless for atomic stuff!! :mad: )

Come to think of it, I never was told what the energy difference is between a full d band and the filled s shell (which makes Cu, Cr, Gd, etc. have that unusual arrangement). If that change does happen to correspond to visible light, then that could do it.

It doesn't explain why copper, gold and other elements with similar arrangement don't have a spectral color though (come to think of it, does someone have a (metallic form) absorbance spectrum handy?).

Tim

turd - 8-5-2006 at 22:21

Quote:
No, as any photon with energy exceeding the energy gap can cause the electron jump. To get a blue metal would require absorption of red and green wavelengths, but not blue - clearly this cannot happen, as "blue" photons are more energetic than red and green and would be absorbed too.

Actually gold clusters can be blue.
BromicAcid and guy gave the correct answer: if it wasn't for relativistic effects, gold would be silver. Meh.

12AX7 - 8-5-2006 at 22:27

Yeah, but that's either a colloidial (red or blue) or quantum effect. I seem to recall they are engineering semiconductor particles for fluorescence in specific waveforms or something to that effect.

Tim

woelen - 8-5-2006 at 22:47

Quote:
Originally posted by BromicAcid
12AX7, your inability to accept my ranting as fact has forced me to look it up:

From Descriptive Inorganic Chemistry 3rd Ed. Geoff Rayner-Canham and Tina Overton; 2003
Quote:
Copper and gold are two common yellow metals, although a thin coating of copper (I) oxide, Cu<sub>2</sub>O, often makes copper look reddish. The color of copper is caused by the filled d-band in the metal being only about 220 kJ<sub>*</sub>mol<sup>-1</sup> lower in energy then the s-p band. As a result, electrons can be excited to the higher band by photons of the corresponding energy range - the blue and green regions of the spectrum. Hence, copper reflects yellow and red. The band separation in silver is greater, and the absorption is in the ultraviolet part of the spectrum. Relativistic effects lower the s-p band energy in the case of gold, again bringing the absorption into the blue part of the visible range, resulting in the characteristic yellow color.

That one I can hardly believe. I surely can believe all things about the electrons, the s-p band and so on, but that copper is yellow seems very strange to me.

I have put copper in oxidizing complexing media quite often, I have done lots of experiment with the metal, but I never observed a yellow color.

If copper really were yellow, then either in a strongly reducing liquid, when some pure copper is put in such a liquid, it should appear yellow (any Cu2O converted to copper), or in a mildly oxidizing acidic environment, it should look the same (any Cu2O converted to soluble copper (II) species, leaving unoxidized clean copper behind). In both conditions, however, the copper simply remains reddish and shiny. Of course, I may be wrong, but then I would really like to see a mechanism for how this supposed yellow color of copper can be shown.

DrP - 9-5-2006 at 00:41

"but that copper is yellow seems very strange to me."

It doesn't say copper is yellow, it says gold is yellow and copper reflects yellow AND red - giving a copper colour. :)

woelen - 9-5-2006 at 01:12

Quote:
Copper and gold are two common yellow metals, although a thin coating of copper (I) oxide, Cu2O, often makes copper look reddish.

DrP, if I read this, then I can only conclude that this tells me that copper is yellow. A compound, which reflects both yellow light and red light I would not call "yellow". It probably has some orange/red appearance and the color at which it appears to me is the color I would name it.

DrP - 9-5-2006 at 01:44

Hmm, yea, I see your point - I wouldn't call copper yellow either (because it's not).

I tend to agree with the whole absorbing and re-emiting of the e's to give the colour though. Perhaps it's the description of the copper colour that is a bit skewed - yellow/red. As you said, I too would have called it orangey red or just copper.

turd - 9-5-2006 at 09:22

Quote:
Originally posted by 12AX7
Yeah, but that's either a colloidial (red or blue) or quantum effect.

It's an electronic effect. The same effect that gives bulk gold its characteristic color. It has nothing to do with light scattering (which is a quantum effect too).

vulture - 9-5-2006 at 12:32

Quote:

BTW, gold's nuclear size has no bearing whatsoever on chemistry (that is, the electronic behavior).


Yes it does. The core electrons of heavy nuclei are subject to relativistic effects because of their high speeds.

Anyway, this gold effect is observed with conducting polymers too, but only when they're partially oxidized. I vaguely remember something about surface plasmons (NOT PLASMA).

unionised - 9-5-2006 at 12:36

"Gold reflects/absorb some light frequencies that gave it the gold color...it's like benzene, who has no color, but its halo-compounds (chloro-,bromo-benzene) are colored, this shows that maybe benzene has is own color, just that cannot be seen. "
This would be a remarkable insight if clorobenzene or bromobenzene were coloured. Unfortunately, they aren't.

12AX7 - 9-5-2006 at 12:44

Quote:
Originally posted by vulture
Quote:

BTW, gold's nuclear size has no bearing whatsoever on chemistry (that is, the electronic behavior).


Yes it does. The core electrons of heavy nuclei are subject to relativistic effects because of their high speeds.


Nuclear charge yes (those inner electrons are bound to the tune of 80keV+ according to (roughly) 13.6eV * Z^2/n^2), but not size.

Tim

guy - 9-5-2006 at 13:39

Quote:

Nuclear charge yes (those inner electrons are bound to the tune of 80keV+ according to (roughly) 13.6eV * Z^2/n^2), but not size.


Nuclear charge affects all atomic radii, but in atoms with large a large nucleus, it causes the electrons to behave differently because ofreletavistic effects, and that pulls the s orbital even closer.

12AX7 - 9-5-2006 at 14:40

Hence "roughly". As I recall, the 13.6eV * Z^2/n^2 isn't derived with relativity in mind (although I seem to remember hearing Scrhodinger's equation is inherently relativistic? I might be confusing something there).

Tim

BromicAcid - 9-5-2006 at 15:09

Speaking of colloidal gold. We did a q-dots lab in my pChem class some time back. We made black, blue, and red gold solutions. This was due to the size of the gold clusters being comparable to the wavelength of the photons or something of that nature. It was very interesting to say the least. I thought that some metals were blue or at least bluish. I know I remember a few described in this way such a cobalt and such, but I have to wonder if this is the description of the pure metal or just the usual cast that is a result of its creation.

turd - 9-5-2006 at 16:14

Quote:
We did a q-dots lab in my pChem class some time back. We made black, blue, and red gold solutions. This was due to the size of the gold clusters being comparable to the wavelength of the photons or something of that nature.

These clusters are typically in the range of 1-100nm, far below the wavelength of visible light.

12AX7 - 9-5-2006 at 16:18

AFAIK, metallic color in terms of steel gray to bluish zinc (it really does look kinda bluish, come to remember it) is usually the surface oxide at work.

Hmm, now I need to know zinc's absorbtion spectrum...

Tim

unionised - 10-5-2006 at 10:40

There's a gold indium alloy that's really blue, not an oxide film.

pyrochem - 9-7-2006 at 09:46

Earlier in this thread, it was mentioned that gold and copper absorb the wavelengths they do not reflect. It seems like some of the light is transmitted. Hold a bright light behind a piece of gold leaf and the transmitted light will look blue/green. I think the same effect occurs with thin films of copper. Is the light really transmitted or is this the same effect as with the quantum dots?

Jdurg - 9-7-2006 at 17:46

Cesium metal is also a yellow color just like gold, and osmium metal has a blue hue to it. Pure bismuth and pure tantalum metal have a pinkish color to them as well. Still, Copper, Gold and Cesium are the only three metals I know of that have a solid color to it and not a hue.

neutrino - 9-7-2006 at 19:08

Is osmium really blue or is this just refraction from its oxide layer? A quick google search revealed no answers.

Nerro - 10-7-2006 at 00:53

According to my professors the high mass of the nuclei near and past gold causes relativistic contraction of the s and p orbitals and relativistic expansion of the d and f orbitals. This accounts for the chemically inert nature of Au, Pt, Hg and other elements, incidentally it also accounts for the fact that Hg is a liquid at RT. It would seem to me that it also explains why lanthanides and actinided quite readily go to very high oxidation states like +6 and +8.

The colour of gold, caesium and some other elements is said to be caused by the decreased energy requirement of the jump from 5d to 6s which is caused by this relativistic contraction.

Why the s and p orbitals contract as the mass of the electrons increases I don't know. The same goes for the relativistc expansions. Just like I'm not sure why the speed of electrons should vary according to their distance from the core.

-edit- I did some reading and apparently the electrons move in circles within their "probability clouds", as their masses increase because their speeds do (I'm still fairly clueless why the speeds would vary...) the radii of these circles decrease and are thus brought closer towards the core. It's the center of the electrons orbit that is used to calculate the properties of the particular electrons and so if the radius of the orbit decreases so does the observed distance to the nucleus. The weird shapes of the probability clouds for d and f electrons might account for their expanding rather than contracting.

This whole story makes a lot of sense to me except for one thing, why does the speed of an electron increase as it goes farther away from the nucleus?



[Edited on Mon/Jul/2006 by Nerro]

turd - 11-7-2006 at 10:35

Quote:
It would seem to me that it also explains why lanthanides and actinided quite readily go to very high oxidation states like +6 and +8.

No, this is mostly non-relativistic effect: higher energy orbitals are more easily oxidised/polarised. You see the same effect for all groups: Xe-compounds, but no He-compounds. I easily oxidised, F not so. Etc, etc...
Quote:
Why the s and p orbitals contract as the mass of the electrons increases I don't know.

Gravitation.
Quote:
The same goes for the relativistc expansions.

If the inner orbitals are contracted, the nucleus is shielded, therefore the electric field is weakened.
Quote:
Just like I'm not sure why the speed of electrons should vary according to their distance from the core.

Electric and gravitational field diminish with distance.
Quote:
The weird shapes of the probability clouds for d and f electrons might account for their expanding rather than contracting.

Kind of; it's not so much the shape as the radial probability distribution.

Nerro - 11-7-2006 at 11:21

Quote:

Electric and gravitational field diminish with distance.
Why does that make the electrons travel faster when they go farther from the core? Also, I thought gravitational pull was insignificant at the atomic level. Electrostatic force is supposed to be something like 10^27 times stronger...

Jdurg - 11-7-2006 at 18:22

Quote:
Originally posted by neutrino
Is osmium really blue or is this just refraction from its oxide layer? A quick google search revealed no answers.

Osmium only forms the simple tetroxide (OsO4) under typical conditions, and this is a VERY volatile and very nasty smelling liquid. As a result, it does not bond at all to the metal itself. Osmium metal itself is a nice bluish color and is really neat to see.

praseodym - 11-7-2006 at 20:38

something like this?


woelen - 12-7-2006 at 00:05

Just a few remarks about the past few posts in this thread.

Lanthanides do not form high-oxidation state compounds. The highest oxidation state, which can be achieved is +4 for cerium metal (and from experience I know that this is DIFFICULT to obtain from the +3 ion). The lanthanides' main feature is that they are so uniformly restricted to their +3 oxidation state. Some exist in +2 (but with great difficulty) and cerium as +4, and that's all.

Osmium does not only form OsO4. It also forms a lower, very hard, inert and high melting point oxide, OsO2. But I agree with Jdurg, that osmium has a blue hue and I also think it is due to the color of the metal itself, and not due to oxide layers.

Gravity does not play ANY role at the chemistry at the atomic level. Gravity only plays an important role at super-macro scale (e.g. planets, moons, solar systems), even two bodies of 100 kg, held close together have extremely low gravitational interaction. Even the slightest static charge on these bodies will result in a larger electrolstatic force.

The relativistic contraction of orbitals in the heavier elements is not due to high mass of the nucleus, but due to high charge of the nucleus. If you solve the wave equation for a two point-charge system, and you increase the charge of one of them, then you'll see that the orbitals shrink and "velocity" increases. It is this increase of "velocity" which introduces relativistic effects. For large charges of the nucleaus, one should solve the wave equations, but not in a classical sense, but in a relativistic sense. The math involved, however, is increadibly complex.

[Edited on 12-7-06 by woelen]

turd - 13-7-2006 at 22:44

Ugh. Of course gravity doesn't play a role, what was I thinking!
Quote:
The relativistic contraction of orbitals in the heavier elements is not due to high mass of the nucleus, but due to high charge of the nucleus. If you solve the wave equation for a two point-charge system, and you increase the charge of one of them, then you'll see that the orbitals shrink and "velocity" increases. It is this increase of "velocity" which introduces relativistic effects. For large charges of the nucleaus, one should solve the wave equations, but not in a classical sense, but in a relativistic sense. The math involved, however, is increadibly complex.

For qualitative statements you don't necessarily have to solve the Dirac equation and not even the Schrödinger equation. It's still the same laws of nature you find on a macroscopic scale, just applied in a "funny" way. Isn't it as simple as more mass and same force gives less speed and thus a smaller orbital?

DrP - 14-7-2006 at 00:24

Quote:

Quote: BROMIC
We did a q-dots lab in my pChem class some time back. We made black, blue, and red gold solutions. This was due to the size of the gold clusters being comparable to the wavelength of the photons or something of that nature.
Quote TURD
These clusters are typically in the range of 1-100nm, far below the wavelength of visible light.


I think that it comes into the wavelenght of the surface plasmons though. In Raman Spectroscopy, if your metal clusters are the right size (depending on the sample and the laser wavethength) you can get the surface plasmons to resonate which enhances your Raman scattering about 50 fold. Usefull for very weak/faint signals.

Nerro - 14-7-2006 at 06:03

I always thought the colour of a colloid had something to do with the way light is diffracted inequally by the particles. The perceived colour would be the product of the bending of the light.

Can someone explain to me what plasmons are?

-edit- google really is a good friend sometimes :P

I read that plasmons are the colelctive oscillations of the "electron" gas in a plasma. I can see how the valence electrons of a metal can move about freely enough to enable similar effects on its surface. Can these collective electron-cluster oscillations also absorb and emit photons?

The coolness of theoretical chemistry never really wears off, does it? :D

[Edited on Fri/Jul/2006 by Nerro]

DrP - 14-7-2006 at 07:46

Quote:
[quote: Nerro]
I read that plasmons are the colelctive oscillations of the "electron" gas in a plasma. I can see how the valence electrons of a metal can move about freely enough to enable similar effects on its surface. Can these collective electron-cluster oscillations also absorb and emit photons?


Thats right - SURFACE plasmons. I don't actually know if the photons are absorbed and re-emmited by the surface plasmons, but, it would make sense - if the Raman scattering is increased 50 fold (it is at least 50 fold - I can't remember the exact amplification factor) and the plasmons are in resonance then I could well imagine that a photon of the right wavelength could get absorbed and then thrown out again at right angles (Raman scattered light is given of at 90 deg).
Quote:

The coolness of theoretical chemistry never really wears off, does it?


Ha Ha - I might agree, but alot of people would call you a complete geek! ;) :D As for it being theoretical - a friend of mine done his Ph.D. back in the 90's on Raman spectroscopy and the effect of surface plasmon resonances on the signal. The work was being undertaken with a view to being used for drug delivery mesurements - diffusion co-efficients were mesured across membranes using Raman spec as the detection system for extreamly small concentrations of the drug. Because the drug concs were so low - they were trying to use the surface plasmon resonances from coloidal films to amplify the signal. I thought it was all pretty cool as well.


Oops! sorry - I've just realised this is all abit off topic - although as it has been said the particle size of the clusters effect colour.



[Edited on 14-7-2006 by DrP]