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emrebjk
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why gold is yellow?
why gold is yellow.Do you know why?
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solo
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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
It's better to die on your feet, than live on your knees....Emiliano Zapata.
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kazaa81
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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.
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12AX7
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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
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BromicAcid
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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]
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12AX7
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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
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guy
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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.
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12AX7
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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
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guy
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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]
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BromicAcid
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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. |
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Twospoons
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| 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.
Helicopter: "helico" -> spiral, "pter" -> with wings
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12AX7
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(Argh, why do people have to use kJ/mol, it's so utterly useless for atomic stuff!!  )
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
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turd
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| 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.
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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.
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12AX7
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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
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woelen
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| 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. |
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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.
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DrP
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"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. 
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woelen
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| Quote: | | Copper and gold are two common yellow metals, although a thin coating of copper (I) oxide, Cu2O, often makes copper look reddish.
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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.
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DrP
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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.
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turd
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| 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).
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vulture
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| 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).
One shouldn't accept or resort to the mutilation of science to appease the mentally impaired.
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unionised
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"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.
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12AX7
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| 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
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guy
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| 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.
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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.
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12AX7
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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
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BromicAcid
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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.
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