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September 17

light slowing down in a material

After doing quite abit of undergrad work on light,(reflection, refraction, the various velocities etc) no one ever explained why light slows down in a material. I mean at a fundamental level what are the photons and (electrons i assume?) getting up to that means the light velocity changes and why therefore material X has a different refractive index to material Y. Thanks --81.147.107.44 (talk) 12:14, 17 September 2010 (UTC)[reply]

Does this paragraph from Wikipedia's article on refraction help?

At the microscale, an electromagnetic wave's phase speed is slowed in a material because the electric field creates a disturbance in the charges of each atom (primarily the electrons) proportional to the permittivity of the medium. The charges will, in general, oscillate slightly out of phase with respect to the driving electric field. The charges thus radiate their own electromagnetic wave that is at the same frequency but with a phase delay. The macroscopic sum of all such contributions in the material is a wave with the same frequency but shorter wavelength than the original, leading to a slowing of the wave's phase speed. Most of the radiation from oscillating material charges will modify the incoming wave, changing its velocity. However, some net energy will be radiated in other directions Refractive_index#Interaction_of_light_and_the_medium Also remember, as you move towards the blue end of the spec the quanta value gets higher. --Aspro (talk) 13:08, 17 September 2010 (UTC)[reply]

Thanks, i see my problem i only looked at the refraction page not the refractive_index page--81.147.107.44 (talk) 13:18, 17 September 2010 (UTC)[reply]

Date of a photo

If I have a photo taken in a park where you can see landmarks and thus determine the direction the camera was facing and the sky looks clear and sunny could I determine the date by measuring the position of the sun? TheFutureAwaits (talk) 12:42, 17 September 2010 (UTC)[reply]

Yes. As always, Wikipedia has an article about it and everything else. It is called Forensic astronomy.--Aspro (talk) 12:57, 17 September 2010 (UTC)[reply]

Yes, the shadows would be the best guide. If it is a familiar location, you could just visit regularly until you see matching shadows, otherwise there is a fair amount of measurement and calculation involved. I'm not convinced that an exact date can be calculated reliably (and, of course, you need other clues to identify the season), but you might be interested in Masquerade_(book)#Solution. Dbfirs 13:12, 17 September 2010 (UTC)[reply]

Try checking the metadata, that usually has the date in it Quadrupedaldiprotodont (talk) 15:25, 17 September 2010 (UTC)[reply]

If you could tell the direction of a linear shadow (like the shadow of a post or pole), you would be able to tell the time of day fairly accurately as it would act as the gnomon of a sundial. The shadow would move from west to east and be pointing north at 12 noon in your local standard time (not daylight-saving time). In theory, you might be able to tell the month by the length of the shadow, but it would be difficult as I should think you would have to do a calibration as Dbfirs suggests by going to the park to check it over a long time frame. In a park, it would be much easier to look at what the flowers and trees were doing - and probably just as accurate. Only my opinion though; I'm not much of an expert. Alansplodge (talk) 23:06, 17 September 2010 (UTC)[reply]

The answer to the original question is no. By examining shadows you can determine the position of the sun and that will give you some information about the date, but it's not enough by itself to determine the date. Today is September 19, or just before the equinox; the sun will follow almost exactly the same path through the sky at the same length of time after the next equinox, i.e. sometime around March 24. And again on approximately these dates next year, and the following year, and so on, give or take a day due to the length of the year not being an integer.

If you could measure the sun position with extreme accuracy you might be able to do better by being able to say "This position is between the sun's paths for September 18 and 19 or 2010, so it must be in March 2010 or in a different year"; but most photos won't contain sufficient detail for that sort of measurement. --Anonymous, 04:06 UTC, September 18 or 19 (okay, 19), 2010.

FRACTIONAL DISTILLATION.

What are the exact weight and volume percentages of products (Petrol, Diesel, Naphtha, Kerosene , Natural Gas, Fuel Oil) obtained by Fractional distillation of crude oil???(considering normal US standards) —Preceding unsigned comment added by 180.149.48.66 (talk) 13:02, 17 September 2010 (UTC)[reply]

That will depend on the crude and whether its light heavy, tar-sands etc. --Aspro (talk) 13:11, 17 September 2010 (UTC)[reply]

Consider light crude oil. thanks —Preceding unsigned comment added by 180.149.48.66 (talk) 13:30, 17 September 2010 (UTC)[reply]

An exact answer is not possible. These fractions vary with each and every oil well, and also, the fractions change throughout the production life of a single well. Googlemeister (talk) 13:36, 17 September 2010 (UTC)[reply]

Representative answers, however, can be provided. See these Dept of Energy sites for refined proportions and monthly US refinery yield. Note that the above are still correct -- we probably can't provide exact values, but the monthly yields will let you calculate the average values. — Lomn 13:51, 17 September 2010 (UTC)[reply]

I'd be a bit cautious in interpreting those numbers. The final refinery outputs will almost certainly include products of cracking to increase the yields of lighter fractions like gasoline. In other words, the refinery numbers won't match what you would get from straight fractional distillation of the raw crude, because they involve further processing. TenOfAllTrades(talk) 14:07, 17 September 2010 (UTC)[reply]

I'm not even sure that two different refineries would give you exactly the same proportions from the same crude oil. Fractional distillation of a complex mixture like crude oil is extremely sensitive to process parameters and plant design. Even in the U.S., I doubt that every refinery is optimized for maximum gasoline production (even if most are). Physchim62 (talk) 14:18, 17 September 2010 (UTC)[reply]

aids

I remember watching a tv show a long time ago about some guy who got aids and took a load of pills and magically the aids was cured and he was a media mystery. Who was that guy? —Preceding unsigned comment added by 8947tn904 (talk • contribs) 14:52, 17 September 2010 (UTC)[reply]

No, AIDS cannot be cured in this manner. However, there are medications right now that will suppress the progress of the disease to the point where many people can lead symptom-free lives for a very long time. AIDS#Treatment and blue links you can follow from there go over how the disease is treated. --Jayron32 15:00, 17 September 2010 (UTC)[reply]

That's nice but not what I asked. I saw a tv show about a guy who took a load of pills and his aids was gone. I want to know the name of the guy, or the name of the tv show. It was some sort of documentary. I don't care if it's scientifically impossible or not, I just want to know the name of the guy so I can look it up 8947tn904 (talk) 15:04, 17 September 2010 (UTC)[reply]

Try posting on the entertainment desk, they are better at that. Ariel. (talk) 16:54, 17 September 2010 (UTC)[reply]

Agreed. It would also be helpful to specify whatever additional details you can remember like a time frame, what sort of show, the language of the show, if you can remember what country the show was from or at least in what country you saw it etc etc Nil Einne (talk) 19:49, 17 September 2010 (UTC)[reply]

I think I remember there being a very low, but not nonexistent rate of spontaneous remission of the body from the AIDS in a few individuals (and in addition a very small part of a population may be immune). It could be possible that the subject took the drugs then spontaneously healed. ~A H 1^(TCU) 21:28, 17 September 2010 (UTC)[reply]

There is also a small but not nonexistent chance of a misdiagnosis.Sjö (talk) 11:44, 21 September 2010 (UTC)[reply]

You might be looking for the patient mentioned in the last paragraph of AIDS#Experimental and proposed treatments. He didn't take pills though, but received bone marrow from someone with a CCR5 mutation, which has a similar effect as Sickle Cell Anemia has with Malaria. -- 78.43.71.155 (talk) 21:04, 18 September 2010 (UTC)[reply]

wireless transmission of electricity

--Angadi Siddhartha (talk) 15:04, 17 September 2010 (UTC)[reply]
Is it possible to transmit electricity without using wires ?

Yes, see Wireless energy transfer 1230049-0012394-C (talk) 15:07, 17 September 2010 (UTC)[reply]

(ec) See wireless energy transfer and take note that making your question big big bold isn't going to get it answered better or faster. So, I removed that formatting. -- kainaw ™ 15:08, 17 September 2010 (UTC)[reply]

See Crystal radio for a rather old application of wireless energy transfer. These radios drew their power directly from the radio wave. You can even build one of these yourself. :-)

Fun fact: In the early days of radio, people living near broadcasting towers would put a lightbulb between two pieces of wire tuned to the tower's wavelength, drawing power from the radio signal to light up the bulb. That was soon outlawed - but technically, it would still work today. -- 78.43.71.155 (talk) 20:53, 18 September 2010 (UTC)[reply]

Sounds bogus. How did they "tune two pieces of wire to the tower's wavelength?" How long were those pieces of wire? Edison (talk) 02:13, 19 September 2010 (UTC)[reply]

A dipole antenna is a common "homemade" antenna. It is very short compared to the wavelength and has very good gain. I used dipole antennas in an experiment a few years ago to see how much power I could get from cell towers using a series of tiny (very tiny) dipole antennas. It was enough to measure, but not enough to power anything - perhaps enough to trickle charge a battery. -- kainaw ™ 03:41, 21 September 2010 (UTC)[reply]

Sulfur oxidation state discrepancy

The oxidation state of sulfur in sulfur dichloride is presumably +2 . The oxidation state of sulfur in sodium sulfate is presumably +6. Why is sulfur dichloride a so much stronger oxidizing agent? It can release chlorine gas upon standing, while sulfates can hardly do anything. Thanks. --Chemicalinterest (talk) 15:10, 17 September 2010 (UTC)[reply]

Is SCl2 necessarily a stronger oxidising agent? I think it's kinetically more unstable. Sulfates need to be protonated to be oxidised (just like nitrates) -- and of course you need a very low pH for that to happen (like, you need fuming sulfuric acid). John Riemann Soong (talk) 15:48, 17 September 2010 (UTC)[reply]

I wouldn't really say that SCl₂ is a strong oxidizing agent: I think Chem.int. is just focussing on one single reaction rather than the wider chemistry. Take the industrial manufacture of thionyl chloride, SOCl₂, for example

SO₃ + SCl₂ → SOCl₂ + SO₂

That's a classic conproportionation reaction, where S(VI) + S(II) = 2S(IV). SCl₂ is the reducing agent for sulfur trioxide (which is not a very strong oxidizing agent). Now if you look at the decomposition of SCl₂

2SCl₂ ⇌ S₂Cl₂ + Cl₂

that's an endothermic reaction, but not by much. The sulfur–sulfur bond that you form is about the same strength as the sulfur–chlorine bond you break, while the Cl–Cl bond is slightly weaker. But chlorine is a gas, which tends to push things towards the right (especially if you raise the temperature). Physchim62 (talk) 23:21, 17 September 2010 (UTC)[reply]

Are you sure its comproportionation? Isn't it symproportionation? Create a redirect if the term is really used. --Chemicalinterest (talk) 00:49, 18 September 2010 (UTC)[reply]

IUPAC says "comproportionation" [1]. In Latin, the prefix con- is the opposite of dis-, while syn- is the opposite of anti- ("n" assimilates to "m" before "p", as I forgot above). Physchim62 (talk) 01:43, 18 September 2010 (UTC)[reply]

To give some perspective, +40 kJ/mol (the heat of decomposition of SCl2) doesn't even exceed 2 OH-H hydrogen bonds). And each water molecule AFAIK forms on average 2 and half in the liquid phase. John Riemann Soong (talk) 00:56, 18 September 2010 (UTC)[reply]

Indeed. The Gibbs free energy change is only +18 kJ/mol (in the direction of SCl₂ decomposition). To put it another way, the mean bond energy of the S–Cl bonds is about 270 kJ/mol each, which is respectable for a covalent single bond but nothing amazing: carbon–chlorine bonds, for example are about 50 kJ/mol stronger. Physchim62 (talk) 01:43, 18 September 2010 (UTC)[reply]

why does blood smell like a spoon?

I understand blood is full of iron, but I didn't realise that the iron was prevalent in that high a concentration. For that matter, a drop of blood has more spoon-smell than a spoon itself. Wth? I'm pretty sure blood is not made of solid metal. John Riemann Soong (talk) 15:47, 17 September 2010 (UTC)[reply]

Metalic iron has a very low vapor pressure. Why would it smell at all? In order to smell a substance, it needs to be a gas. The ability to become a gas has nothing to do with the relative concentrations of substances. Furthermore, the smell of blood is not necessarily the smell of iron. --Jayron32 15:56, 17 September 2010 (UTC)[reply]

See this. You do not actually smell metal. What you smell is the body's reaction to metal. So, your nose is reacting to both blood and a spoon in a similar manner. -- kainaw ™ 16:04, 17 September 2010 (UTC)[reply]

Thanks an interesting read. I've noticed before my hands smell like 'metal' when I've been holding keys or similar things but it's not something I've ever thought to look in to Nil Einne (talk) 19:35, 17 September 2010 (UTC)[reply]

Thanks so much. I was really puzzled at why blood would smell so strongly of metal when I was sure metal had no substantial vapor pressure. John Riemann Soong (talk) 19:45, 17 September 2010 (UTC)[reply]

It tastes like spoons too. I used to get nosebleeds in my sleep, and I have vivid memories of waking up with the taste of spoons in my mouth, and thinking "Oh no, another bloody pillow". DuncanHill (talk) 16:19, 17 September 2010 (UTC)[reply]

Are you sure it's a spoon? I just had some of my own and it tastes like more like a spork to me. DRosenbach ^{(Talk | Contribs)} 16:38, 17 September 2010 (UTC)[reply]

The hemoglobin in your blood contains iron oxide compounds, much the same substance that makes up rust. Metal spoons also contain iron and are in contact with air. ~A H 1^(TCU) 21:23, 17 September 2010 (UTC)[reply]

You are smelling Oct-1-en-3-one. Graeme Bartlett (talk) 11:36, 18 September 2010 (UTC)[reply]

...which is completely devoid of iron altogether! SpinningSpark 11:57, 18 September 2010 (UTC)[reply]

liquid becoming solid at higher temperature?

There was something on the radio this morning about a substance that might be used for medical treatments that is a liquid when chilled but a solid at body temperature. Is this actually a reversable phase change or is it more like a thermoset glue? (If it's the former, is there a name for this property?) RJFJR (talk) 18:40, 17 September 2010 (UTC)[reply]

This news item sounds like what you are remembering. DMacks (talk) 19:16, 17 September 2010 (UTC)[reply]

That's it: "a injectable formulation derived from natural sugar chains that exists as a liquid when cooled and becomes a solid at body temperature." RJFJR (talk) 19:19, 17 September 2010 (UTC)[reply]

Thermoreversibilty I think could be the word. Poloxamer gels exhibit this property. They are nowhere as strong as the gels we are used to though. Perhaps that why they want to develop them further. They would have the property of keeping damaged cells from leaking out their contents out all over the place but are not up to the job just yet. --Aspro (talk) 19:46, 17 September 2010 (UTC)[reply]

The article Thermosetting polymer deals with the process in general terms; this sounds like the use of bioorganic molecules like sugars and polysacharrides to develop materials with similar properties. --Jayron32 19:47, 17 September 2010 (UTC)[reply]

Here we go: Temperature Induced Gelation. It is probably the next sec of 5.2.2. Cellulose derivatives which is applicable because cellulose is only sugar. Here's a patent as well.[2]--Aspro (talk) 20:11, 17 September 2010 (UTC)[reply]

Oh and I forgot to say: I found the patent because as always Wikipedia has an article about it (Poloxamer) and it was under references. Thermoreversible also appease to be the commonly accepted term for this property.Here is another example of its use. [3] Does this fully answer your question?--Aspro (talk) 21:22, 17 September 2010 (UTC)[reply]

hydrogen peroxide, oxidants and Lewis acids

I see we don't have any articles on Lewis acid catalysis, especially for organic redox reactions. I would eventually like to start an article on one (with appropriate research). I'm especially interested in some of these reactions as proof of concept things...when I search the literature I get really complex asymmetric catalysts -- while that's really useful I'd like to start from the basics. If I get papers from the 1870s -- even better.

H2O2 + Fe(III) + THF. Do I get an alpha-substitution? Do I get an acetal (assuming no cleavage), lactone or an ether peroxide? Or do I just get a catalytic decomposition of the hydrogen peroxide?

Cu(II) chloride + H2O2 + THF. If I get substitution, is it chlorination, peroxidation or hydroxylation? Or does the hydrogen peroxide reduce the Cu(II) to Cu(I), with the evolution of O2? Or just disproportionation of the hydrogen peroxide?

What if I use ice bath conditions?

p.s. can I prepare a solution of CuCl2 by pouring 30% H2O2 on copper and then reacting it with HCl? I actually have access to lab facilities where I can do this stuff (I mean I handle gold(III) chloride and perchloric acid regularly for biochem lab protocols...so), and a fume hood. John Riemann Soong (talk) 20:07, 17 September 2010 (UTC)[reply]

Hydrogen peroxide isn't actually all that great to work with as a reagent. IIRC, most of the time, if you require a peroxide, you use stuff like Organic peroxide and peroxyacids. Peroxides mainly work as radical initiators for things like radical halogenation and other reactions, because of the homolytic cleavage of the O-O single bond. Their use in oxidation reactions occurs by similar mechanisms. --Jayron32 20:23, 17 September 2010 (UTC)[reply]

Hmm, I seemed to think that by two-electron transfer, peroxide worked via the same mechanism as bleach -- a sort of SN2 attack on an electrophilic oxygen that would be activated by acid. I mean, they say "the O-O bond is weak" but the O-O bond is still quite a stable covalent bond. That is, OH* formation occurred but not rapidly -- just enough to jumpstart chain reactions. John Riemann Soong (talk) 23:02, 17 September 2010 (UTC)[reply]

The O-O bond is quite weak, but is stabilized in organic (alkylic) peroxides by the presence of electron donating groups, which alkyl groups tend to be. Hydrogen is a terrible electron donator, so in hydrogen peroxide, the O-O single bond is as weak as it is going to get; which is why hydrogen peroxide degrades fairly rapidly at room temperature or in strong light. --Jayron32 03:06, 18 September 2010 (UTC)[reply]

To answer the original question, using H₂O₂ with Fe(III) or Cu(II) will give you a form of Fenton's reagent. The original Fenton's reagent was with iron, but copper works as well: you simply need an element which can undergo one-electron transfer. Fenton's reagent is very effective at oxidation by oxygen transfer, but it's rather unspecific at where the oxygens go. Unsurprising really, given than the active species in Fenton oxidations is the hydroxyl radical! Physchim62 (talk) 08:45, 18 September 2010 (UTC)[reply]

Yes, you can make CuCl2 by reacting H2O2 and HCl with copper. I have 3% H2O2 and it works, albeit slowly. H2O2+HCl is a safe oxidizing acid. --Chemicalinterest (talk) 12:03, 18 September 2010 (UTC)[reply]

That's not how you'd usually make it though (assuming you have to make it instead of buying it). A standard prep would be to dissolve the copper in nitric acid, then precipitate the copper(II) as the hydroxide or the carbonate: wash the precipitate, then dissolve in the minimum of dilute hydrochloric acid to get your copper(II) chloride. Physchim62 (talk) 12:25, 18 September 2010 (UTC)[reply]

My lab doesn't have nitric acid. I don't have nitric acid. We do have perchloric acid though. Plus I'd rather spill 30% H2O2 than nitric acid eeek. Btw, Fenton's reagent appears to potentially be an oscillatory reaction. That's just fascinating....plus it's used in organic synthesis so it must have some selectivity if you have the right conditions? Esp. for alpha-oxidation? John Riemann Soong (talk) 17:54, 18 September 2010 (UTC)[reply]

Kitchen chemistry experiment

Hi. I poured sour lemonade into a bowl, then added some baking soda. At this point, mild fizzing occurred. I then proceeded to add leftover milk and mustard to the mixture, and slow mixing resulted. Afterwards, I added vinegar to the mix, which reacted with the baking soda. Next, I added more milk, and put in some ketchup. The next step was a surprise, as it produced an unexpected reaction. When I added some table salt to this mixture, it suddenly started fizzing, producing a white floating layer with a thick layer of small bubbles. Next, I put in some oil, which quickly expanded from small drops, and continues expanding. I put the bowl into the microwave for ten seconds. After that, I tried adding more salt to the mix, and an interesting thing happened. The salt fizzed as before, but within the oil regions, some salt crystals floated on top and did not dissolve, but a denser mass sank through the oil and the liquid underneath, and then rose up to form denser bubbles (denser because these were underneath the oil bubble layer). Within the oil areas, a line of bubbles from the salt began forming at the edges, but they were different in texture (if you could call it that) than the ones outside the oil. After a few more minutes, the oil bubbles proceded to expand, and the floating salt was replaced with a few bubbles, and convection started occurring when a layer of foam started in the middle of an oil layer and floated to opposite directions in a line, causing the oil layer/bubble to expand further. By the time I dumped out the contents of the experiment, the oil layer covered in patches most of the mixture's surface, with circular regions within and outside the oil layers that appeared white. Most of these regions had a foam surface on the outside, and in the central part there was an oily mass of the same bubbles that surfaced after the salt. At the end, when I poured everything out, a glob of white foamy substance remained stuck to the bottom and the bowl was left with an unappealing smell. Please help explain the following:

What gases were produced by the reactions?
What salts, other than the ones already present, were precipitated or dissolved into the mixture by the reactions?
About how much water not already present was produced by the neutralization reactions?
What factors caused the oil to expand?
Did the microwaving have any chemical effects on the mixture?
Why did the salt produce more fizzing than the reaction of baking soda with either the lemonade or the vinegar?
Were any toxic or harmful compounds likely produced?
What caused the bottom-sinking (and thus dense) glob?
Were any plastic/polymer-like substances produced?
What compounds were responsible for the smell?

If it is possible to answer the above questions with the information provided, please do so. This is not homework, nor am I requesting medical advice! Thanks. ~A H 1^(TCU) 21:07, 17 September 2010 (UTC)[reply]

I think that's probably the most poorly designed experiment I have ever heard of, although I reviewed a paper once that was almost as bad. Not quite, though: they only had six manipulations all happening together, as opposed to the ten in your experiment. Looie496 (talk) 21:44, 17 September 2010 (UTC)[reply]

At each step, you need to test for possibilities by having diagnostic things to eliminate or confirm suspicions. That's how science is done. John Riemann Soong (talk) 23:00, 17 September 2010 (UTC)[reply]

I'll try to answer some of the questions.

Carbon dioxide. I don't know what else would be produced.
Many things contains acids, such as vinegar in ketchup, so sodium acetate should be a major salt. It is made by reaction of baking soda and vinegar. If citric acid was in anything, you may get sodium citrate.
Water is produced by many acid-base reactions like you did, so quite a lot of water should be added.
Surfectants? I don't know much in that area of chemistry so please help me out.
Probably not, other than speeding it up by heating it.
Please repeat your experiment with the salt and try some other acids. There might have been some discrepancy. Assuming you mean sodium chloride, it is neutral and does not participate significantly in an acid-base reaction.
Hmm, probably not. Most chemical compounds that are in foods are either not toxic or in too small amounts to be toxic.
Check out casein. It is a polymer that can be made from milk. You may have polymerized the milk. You also might have curdled the milk with the acids. This applies to the two questions.
Vinegar has a strong smell. Several other strong-smelling acids may have been released as well as a mist of solution by the fizzing.

Hope this helps! --Chemicalinterest (talk) 23:22, 17 September 2010 (UTC)[reply]

While salt does not participate in traditional acid-base experiments, they (as you discussed so vigourously above) participate in redox reactions and solubility reactions. Salt denatures certain proteins. Salt especially likes to denature casein already under stress by acid. Vinegar should have been added in a side reaction as a diagnostic -- milk is already acidic. Thus you would be able to confirm whether it was the milk's acidity that contributed anything, or something else in the milk. Such is the protocol to deal with any complex mixture, as foods tend to be.

When you cook Western vegetables or chicken in water, a bubbly layer of "scum" often floats to the top -- these are organic complexes (of protein and other biochemicals) that have found new ways to interact and are no longer soluble in water, but do not have the density to precipitate. Ignoring boiling point changes, adding osmotic agents such as salt often produces more scum (so the decision to add salt early in cooking can be beneficial or harmful, depending on your purposes). John Riemann Soong (talk) 00:49, 18 September 2010 (UTC)[reply]

When you acidified sodium bicarbonate, you saturated the water with carbon dioxide, i.e. carbonated it. Adding salt to any carbonated beverage makes it bubble... though I'm not sure I've tracked down the definitive explanation for it yet. Wnt (talk) 04:56, 19 September 2010 (UTC)[reply]

Why not nucleation? I'd expect other factors, like specific ion effects, to be negligible in comparison. 109.155.33.219 (talk) 21:29, 19 September 2010 (UTC)[reply]

Mmm yes, nucleation. Diet coke and mentos anyone? John Riemann Soong (talk) 19:21, 20 September 2010 (UTC)[reply]