I believe the question asked in one of the classes this week was that why are metals so shiny? The explanation given in class was:
Metals have high free/conduction electron density > which means they have high plasma frequency > higher than the frequency of visible light > the dielectric constant in the wave equation is negative >the solutions to the wave equation decay exponentially> no radiation can propagate> the electromagnetic wave is reflected> metals appear shiny.
In a classical picture, we imagine all the free electrons as oscillators which upon incident em fields absorb the em field and then re emit it. Why it is not lost dissipated as heat due to collision? Its because the frequency of the incident field, w>1/tau where tau is the relaxation time. i.e. the frequency of the oscillation is faster than the frequency of collision. As pointed in class, above the plasma frequency the metals become transparent such as in the UV region for alkali metals. One way of thinking about it is that the oscillating electrons cannot respond fast enough so em radiation just travels through as if nothing is there.
Then someone asked why other non-metalic objects which also reflect light are not as shiny? It probably has something to do with insulators not having as many free electrons which can respond to the field so they cannot reflect as well. Maybe someone can give a better explanation? It was also pointed out that is also the fact that reflection of a metal surface is specular due to the smoothness of the metallic surface so it looks shinier.
Ok, that probably was boring to some of you so read this http://www.coolsciencefacts.com/2006/metal.html and tell me what you think.
Subscribe to:
Post Comments (Atom)
The science fact was pretty cool: be careful not to scratch your tools too much in space.... I wonder if it would be a problem on Mars, since there is a much lower concentration of oxidising agents? I imagine that most tools are coated pretty well anyway, considering the potential for bad stuff to affect them in most (livable) environments.
ReplyDeleteOn your post;
I thought it was interesting. I recall from 3rd year fields that the sign/imaginary status of the permittivity is pretty important for this situation. Is it metals that have an imaginary one? I think this difference shows up with the 'quality' of shininess: metals are shiny even when scratched (e.g. stainless steel) and they seem to reflect everything visible pretty well (except for gold and copper, due to cool relativistic electrons! Would be some interesting alloy physics...) but a polished rock (was going to say painted wall, but that would be too strange...) looks clearly different to a piece of metal except at some angle of maximum reflection. Upon thinking further, I recall that the angle of reflection changes for metals and dielectrics based on the permittivity.... I suppose that the 'classical' view of electrons being able to oscillate is a good representation: metals have heaps of free electrons spinning around, able to absorb lots of energy and emit it (with many possible frequencies; they're rather unbound) but an organic compound/non-metal has only some or very few delocalised electrons, and these ones are more rigidly bound by the structure nearby. Thus, for (many) metals under UV, we get shiny-ness, but for organic fluorescent compounds (like green fluorescent protein, GFP) we get a pretty colour when looking under UV.
I think the result of this discussion is that more revision of 3rd year fields (or a talk with Timo/Till) is required. :)
Josh H
Okay, interesting thought.
ReplyDeleteIf a metal has oxidised, then it has lost electrons to oxygen, then is metal+oxygen.
If bring an oxidised chunk of metal near a non-oxidised chunk of metal, then that's metal+oxygen near metal.
Would the metal then start to oxidise as well?
Then now we have two chunks of metal which won't stick to any other metal?
(not sure about this reasoning)
Sidenote: Josh, http://www.instructables.com/id/Dorodango/?ALLSTEPS
forgive me for the chemistry term, but one need only consider the Stoichiometry. Simply put, in an isolated system such as this, there are a pre-defined number of electrons/Oxygen atoms. As the oxidation of metal is an energetically favorable process, one would need to input energy to reduce the metal, and free the oxygen. Of course if you could make this happen, then the Oxygen might bind to the other metal, but that is beside the point.
ReplyDeleteThat seems more convincing.
ReplyDeleteThanks Dave.
At least Stoichiometry is a cool chemistry word! I think that Dave's line of thinking is on the right track: the layer of metal oxide is only one metal-oxygen molecule thick (so if the metal was perfectly smooth, there would be a layer of oxygen one atom thick on it), so there wouldn't be any 'spare' oxygen to transfer to the other metal. If oxidising the other metal was sufficiently energetically favourable, however, oxygen would leave metal 1 and bond to metal 2—this is similar to what happens when you coat something with a sacrificial metal. If you're interested in making explosives (and who isn't? Me, obviously, Police!), or rusting stainless steel, you can actually scrape off the oxide layer, revealing the reactive material beneath. In the case of explosives, when you grind up aluminium foil and mix it with rust, you can initiate the thermite (NOT 'termite', Google thinks this is spelt wrong :) ) reaction because there is unprotected aluminium exposed. Aluminium is actually very reactive! For stainless steel, it turns out that the chromium in it rises to the top and is oxidised, so you can remove the chromium oxide top layer and get to the rustable metal beneath if you scratch hard enough.
ReplyDeleteI also like your examples of polished brown stuff, Ann. :)
ReplyDelete