Is Copper the Only Metal With a Blue Color
Why are there only two colorful metals?
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I went to a science exhibition where samples of almost all metals where kept, I think I could only see Copper and Gold having different colors.
Why is it so ? why not other metals show colors ?
Answers and Replies
I think there are only two metals, Copper and Gold , that exhibits some color other than silver or gray.
Osmium is blueish and tungsten is brownish.
So in full spectrum white light they appear to be some variation on grey/silver.
A few metals absorb light more at some wavelengths than others.
Copper looks reddish because it is absorbing more of blue light wavelengths and reflecting more of red.
Oh this is nice but I did not saw anything like this.Google pictures of bismuth. If this isn't colorful, then nothing is. E.g. on http://unitednuclear.com/index.php?main_page=product_info&products_id=1169
Most metals reflect and absorb visible light at all wavelengths.
So in full spectrum white light they appear to be some variation on grey/silver.
A few metals absorb light more at some wavelengths than others.
Copper looks reddish because it is absorbing more of blue light wavelengths and reflecting more of red.
Why only these two or three (counting caesium) only, why not other 115 elements ?
More or less silverish.Osmium is blueish and tungsten is brownish.
i can't get access to the paper. I can only read abstract.Here's a paper on it:http://pubs.acs.org/doi/pdf/10.1021/ed076p200
Its somewhat technical but you should find your answer there.
why not other 115 elements
Most aren't metals. You add the others in and you have iodine (purple), sulfur (yellow), phosphorus (lots of different colors), chlorine (green) and so on.
Maybe a better question would be why are most metals white in the first place?
My guess is because they're big atoms with a lot of free electrons so there is a lot of possible energy levels to jump to and from, so lots of spectral lines?
Color is also based on our very limited view of the Em spwctrum
Wouldn't this contradict the fact, that really many of metal-oxides are clearly defined colors which have been often used as such in former times before we had petrochemistry? How can some oxygen atoms change the entire situation?My guess is because they're big atoms with a lot of free electrons so there is a lot of possible energy levels to jump to and from, so lots of spectral lines?
Colors can also come from structure. Shapes can trap certain wavelengths. Some butterflies have visibly blue wings that have no blue pigment at all: http://www.webexhibits.org/causesofcolor/15A.htmlWouldn't this contradict the fact, that really many of metal-oxides are clearly defined colors which have been often used as such in former times before we had petrochemistry? How can some oxygen atoms change the entire situation?
I know gold shines the way it does because the outer electrons approach the speed of light.
This is actually a bit oversimplified. The relativistic correction affects all of the s orbitals, not just the outermost one (which in gold is 6s). The reason for gold's color is that the correction makes the energy of the 5d-6s transition in gold significantly smaller than the energy of, for example, the 4d-5s transition in silver, so the latter corresponds to a UV photon whereas the former corresponds to a visible light photon. The relativistic correction to the inner s orbitals doesn't have a similar effect because all of the inner shells are filled, so there are no available transitions anyway.
Maybe a better question would be why are most metals white in the first place?
Because in most metals there are no available energy level transitions that correspond to visible light photons; most of them have energies that correspond to UV photons (i.e., too high energy to be visible).
My guess is because they're big atoms with a lot of free electrons so there is a lot of possible energy levels to jump to and from, so lots of spectral lines?
"Big atoms" does not mean "lots of free electrons in the atom". Almost all of the electrons are in filled shells with no energy level transitions available. In the case of transition elements like gold, there will be one, two, or in rare cases three electrons per atom that can be considered "free". (In gold it's just one.)
The vast majority of elements on the periodic table are metals.Most aren't metals. You add the others in and you have iodine (purple), sulfur (yellow), phosphorus (lots of different colors), chlorine (green) and so on.
The color here comes from interference effects from the variable thickness passivating oxide layer that builds up on bismuth. Bismuth itself is dull gray.Google pictures of bismuth. If this isn't colorful, then nothing is. E.g. on http://unitednuclear.com/index.php?main_page=product_info&products_id=1169
This website:
http://www.webexhibits.org/causesofcolor/9.html
Gives a decent explanation. Basically, for gold, its filled d band has a huge density of states compared to the valence s band. This means (from Fermi's golden rule) that electrons from the d band have an appreciable likelihood of being excited above the Fermi level. The d band happens to lie at an energy below the Fermi level that corresponds to blue photons. By comparison, for silver, the analogous band is much further below the Fermi level, meaning that the analogous excitation lies in the UV range. For copper, it's the same case as with gold, the difference being that copper has a smaller lattice constant than gold or silver, while gold and silver comparable lattice constants due to relativistic orbital contraction.
Other transition metals don't exhibit visible range colors because their Fermi levels are located inside the d bands. The coinage metals ("d9 metals" in organometallic chemistry) are oddities in that their Fermi levels appear in the s bands.
So basically most metals are colorful, just not to our eyes.Because in most metals there are no available energy level transitions that correspond to visible light photons; most of them have energies that correspond to UV photons (i.e., too high energy to be visible).
I was actually saying two different things: that they're big atoms so they have lots to hit (which I now see was incorrect since the energy levels are already filled,) but the free electrons I mentioned was not those. I thought the outer shells in metals were constantly swapping electrons with their neighbors, isn't that what makes metal conductive? I assumed that they weren't even in orbits, I'm not sure why, it doesn't make any sense now that I think about it."Big atoms" does not mean "lots of free electrons in the atom". Almost all of the electrons are in filled shells with no energy level transitions available. In the case of transition elements like gold, there will be one, two, or in rare cases three electrons per atom that can be considered "free". (In gold it's just one.)
most metals are colorful, just not to our eyes.
All elements are "colorful" if you don't restrict yourself to visible light. Every substance has available energy levels somewhere that electrons can transition between.
I thought the outer shells in metals were constantly swapping electrons with their neighbors, isn't that what makes metal conductive?
It would be more accurate to say that in metals, there are some small number of electrons per atom (in gold it's one) whose wave functions can "spread out" spatially over a distance much larger than the size of an atom, if the atoms are close together (as they are in a metal in solid, or in the case of mercury liquid, state). The spread is what allows the metal to be conductive.
Correct except there are only ~17 non metals compared to 118 - ~17 metals, So metals are the majority. But this is completely irrelevant to my question.Most aren't metals. You add the others in and you have iodine (purple), sulfur (yellow), phosphorus (lots of different colors), chlorine (green) and so on.
Thanks for the article.Colors can also come from structure. Shapes can trap certain wavelengths. Some butterflies have visibly blue wings that have no blue pigment at all: http://www.webexhibits.org/causesofcolor/15A.html
This is actually a bit oversimplified. The relativistic correction affects all of the s orbitals, not just the outermost one (which in gold is 6s). The reason for gold's color is that the correction makes the energy of the 5d-6s transition in gold significantly smaller than the energy of, for example, the 4d-5s transition in silver, so the latter corresponds to a UV photon whereas the former corresponds to a visible light photon. The relativistic correction to the inner s orbitals doesn't have a similar effect because all of the inner shells are filled, so there are no available transitions anyway.Because in most metals there are no available energy level transitions that correspond to visible light photons; most of them have energies that correspond to UV photons (i.e., too high energy to be visible).
"Big atoms" does not mean "lots of free electrons in the atom". Almost all of the electrons are in filled shells with no energy level transitions available. In the case of transition elements like gold, there will be one, two, or in rare cases three electrons per atom that can be considered "free". (In gold it's just one.)
That explains gold but is this the case for Cesium and copper ?Two minor gripes:The vast majority of elements on the periodic table are metals.
The color here comes from interference effects from the variable thickness passivating oxide layer that builds up on bismuth. Bismuth itself is dull gray.
This website:
http://www.webexhibits.org/causesofcolor/9.html
Gives a decent explanation. Basically, for gold, its filled d band has a huge density of states compared to the valence s band. This means (from Fermi's golden rule) that electrons from the d band have an appreciable likelihood of being excited above the Fermi level. The d band happens to lie at an energy below the Fermi level that corresponds to blue photons. By comparison, for silver, the analogous band is much further below the Fermi level, meaning that the analogous excitation lies in the UV range. For copper, it's the same case as with gold, the difference being that copper has a smaller lattice constant than gold or silver, while gold and silver comparable lattice constants due to relativistic orbital contraction.Other transition metals don't exhibit visible range colors because their Fermi levels are located inside the d bands. The coinage metals ("d9 metals" in organometallic chemistry) are oddities in that their Fermi levels appear in the s bands.
I thinks it is unlikely for copper as it is a lot smaller that both cesium and gold.
I mentioned the copper case briefly in my post, but I'll explain a little more: The placement of the energy bands in the coinage metals (Cu, Ag, Au) is affected by many things, including lattice constant (what you refer to as the "size" of the atom). If copper, silver, and gold were all the same size (let's say they're all the size of gold), then gold would absorb in the blue, silver in the UV, and copper further into the UV. As it turns out, gold and silver are the same size (because of relativistic contraction of the gold orbitals). So gold absorbs in the blue and silver absorbs in the UV. However, copper is much smaller. This means that the filled d energy band in copper gets pushed up toward the Fermi level. As it turns out, it just happens to get pushed up enough that the predominant excitations now lie in the blue/green range, so copper appears orange/red.That explains gold but is this the case for Cesium and copper ?
Cesium is another oddity. All of the alkali metals have half-filled s bands as their valence bands. Below these s bands are filled p and d bands with a high density of states. For lithium through rubidium, the energy to excite electrons from these bands is in the UV, but for cesium, it's just in the visible, giving cesium the gold color you see.
That's a whole different ball of wax: the band structure for elemental metals is much easier to intuitively predict than the band structure of alloys (which may or may not even be crystalline).Not to forget the coloured alloys.
So metals are the majority
There are maybe 80 or 85 metals. But ~20 of them are colored bright red because they are so radioactive that they would glow red-hot if you got enough of them together to see. So out of ~60 metals where we could tell the color, a handful are not silver or grey. Is that a huge number?
Yes that is the question, Why only a handful show different colors ?There are maybe 80 or 85 metals. But ~20 of them are colored bright red because they are so radioactive that they would glow red-hot if you got enough of them together to see. So out of ~60 metals where we could tell the color, a handful are not silver or grey. Is that a huge number?
Thanks for the answer.I mentioned the copper case briefly in my post, but I'll explain a little more: The placement of the energy bands in the coinage metals (Cu, Ag, Au) is affected by many things, including lattice constant (what you refer to as the "size" of the atom). If copper, silver, and gold were all the same size (let's say they're all the size of gold), then gold would absorb in the blue, silver in the UV, and copper further into the UV. As it turns out, gold and silver are the same size (because of relativistic contraction of the gold orbitals). So gold absorbs in the blue and silver absorbs in the UV. However, copper is much smaller. This means that the filled d energy band in copper gets pushed up toward the Fermi level. As it turns out, it just happens to get pushed up enough that the predominant excitations now lie in the blue/green range, so copper appears orange/red.Cesium is another oddity. All of the alkali metals have half-filled s bands as their valence bands. Below these s bands are filled p and d bands with a high density of states. For lithium through rubidium, the energy to excite electrons from these bands is in the UV, but for cesium, it's just in the visible, giving cesium the gold color you see.
That's a whole different ball of wax: the band structure for elemental metals is much easier to intuitively predict than the band structure of alloys (which may or may not even be crystalline).
One more question,
Can we predict color of a undiscovered metal with QM model of a atom ?
In principle, yes. For example, people do quite a bit of theoretical work on the chemical and physical properties of ultraheavy elements. They don't have a choice: those elements aren't stable enough in most cases to get any experimental physical or chemical data.Can we predict color of a undiscovered metal with QM model of a atom ?
Thank you very much.In principle, yes. For example, people do quite a bit of theoretical work on the chemical and physical properties of ultraheavy elements. They don't have a choice: those elements aren't stable enough in most cases to get any experimental physical or chemical data.
This only shows that crystal structure is not very important for the colour of a metal. I.e. the colour of cesium does not change on melting. Also the colour of brass changes but little with composition, but crystal structure does.That's a whole different ball of wax: the band structure for elemental metals is much easier to intuitively predict than the band structure of alloys (which may or may not even be crystalline).
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