text:
How to plate zinc evenly on a threaded bar
I need to set a number of threaded bars into cement, to hold a new gatepost. When I cut the bar into shorter pieces, the metal near the cut became red hot, presumably hot enough to evaporate some or all of the zinc with which the entire rod was plated.
Now I want to plate zinc back onto the ends, and have had mixed success so far. I dissolved some zinc oxide with hydrochloric acid and hooked it up to a little PV panel I have, and after a while it seems there is a light-coloured deposit on the cathode. The pencil-lead (6B) anode is no longer as smooth as it first was, but it doesn't seem to be losing (much) of its bulk. I assume it to be zinc, because... well what else can it be? It plated thick enough that I had to (ab)use a nut to re-cut the thread, leaving a light-coloured shiny surface that gives me more confidence that it is, in fact, zinc.
Problem: the edges of the threads seem to be taking more zinc than the valleys. In some cases the thread's valley doesn't show any evidence of plating. The fractal-like deposit on the edges is quite pretty actually, pity I can't take a (good) closeup photo.
Is there anything I can do (perhaps using some other common(ish) kitchen chemicals?) that will result in a more uniform layer of zinc? Is it worth trying lower / higher current or pulsed current? Am I overthinking things - is it even necessary for the whole stud to have a layer of zinc?

Cool experiment! Not an answer to your question, but what you're experiencing in regards to differential rates of deposition on the edges versus the valleys is likely to be a type of diffusion limited aggregation. This is just a wild guess, but rapid agitation of your zinc solution (say, by a stirrer bar or sonication) might help prevent depletion of the solution in the thread valleys.

Yes, I discovered that stirring the rod now and again would disturb the bubbles (hydrogen?) that get trapped in the threads - which couldn't have been helping. That seemed to help, as the zinc then seemed to deposit over the whole surface.

I think the easiest way is to sputter coat the bars, if you can access one. One that can rotate your bars such that the rotation is vertical to the direction of sputtering would be the best. If not, switch sides manually.
Depends on the composition of your bars, you might want to coat something else before coating zinc. A wild guess: nickel.  I suggest so because I don't recall zinc being sticky. So the coating might come off easily.
And I'd try sputter slow but for a long time. I think low sputtering rate tends to produce more densely packed and uniform coating. I am not sure. But dense and uniform coating provides good foundation of the zinc coat.
For electric coating, I think low voltage and uniform coating is still nice. I don't think you want formation of gas at the bar electrode, as that disturbs the structure of the coating. Not sure what's influencing the nucleation sites on the bars. I would guess that the current is not uniform on the bars, say there more on the hill tops. Still consider coating something else first. I would try whatever is available. I would try to get the solutions as pure as possible though. Organic compounds can act as surfactants which tend to aggregate the particles.
Whichever your method is, search for some journal articles characterizing similar processes. That might help.

I see this is an old question ; however, in the real world steel rebar is protected with polymer coating or stainless steel rebar is used. Generally protection is unnecessary except for things like seacoast bridges.

Threaded bar is not rebar. This question is about threaded bar, and about zinc plating, not about rebar and epoxy / polymer coatings. Also not about concrete structures where rebar is deeply embedded and not directly exposed to the elements.

Oops ! Threaded rod and studs are routinely hot dipped and electrogalvanized ,  a thread die may be run over them as necessary.  But not in sizes under 1/2 in. 12mm , they are more commonly polymer coated.

Lookup "throwing power" on the net to possibly get help plating into recessed areas.


text:
Differences between KOH and phosphorus pentoxide as drying agents?
The most common drying agent I encountered when using a vacuum dessicator was phosphorus pentoxide in the form of Sicapent. But I've recently encountered people using KOH as a cheaper alternative (in a practical course for students, where such costs are amplified due to the high number of students).
What are the differences between KOH and phosphorus pentoxide as a general purpose drying agents to remove water? Is there a difference in effectiveness? Are there classes of substances where you can't use one of those drying agents?

short answer: KOH will cheaply dry alcohols, KBr pellets,  and most organic liquids.  Phosphorous pentoxide—a bit more expensive—will dry sulfuric acid into sulfur trioxide and anything else it contacts.

I found this page, which recreates information from Purification of Laboratory Chemicals, by Perrin and Armarego. It provides a qualitative ranking of drying agents, listing phosphorous pentoxide higher than KOH. The two drying agents work by different means:
KOH
Anhydrous KOH is hygroscopic. It absorbs water from the atmosphere. It much better at doing so than more common hygroscopic drying agents like calcium chloride or calcium sulfate (drierite). A pellet of KOH left out in a beaker or dish open to the air will have absorbed so much water vapor within an hour that it will now be a solution of KOH (that will continue to be hygroscopic, at least until the solution is 50% water). Calcium chloride and calcium sulfate only continue to absorb water until they reach their maximum hydration.
P4O10
Phosphorous pentoxide (really P4O10) is an anhydride of phosphoric acid. Phosphorous pentoxide reacts with water to form phosphoric acid:
$$\ce{P4O10 + 6H2O -> 4H3PO4}$$
Phosphoric acid is also hygroscopic and absorbs water to form a solution.
Since P4O10 is water reactive it is more difficult and dangerous to handle, especially for large groups of students. I might not use P4O10 if the substance being dried is very acid sensitive. Eventually, the vacuum will start volatilizing the acid. I might not use KOH for compounds that are incredibly base sensitive, although it is unlikely that KOH will sublime. In an undergraduate lab with multiple students, KOH should be fine if you want to save money and hassle.

This answer is exactly right and deserves to be marked as correct.

Usage of drying agent is ruled by four factors: cost, effectiveness, reactions it can provoke in substance being dried, level of discomfort in working process.
$\ce{P2O5}$ is much more effective, $\ce{KOH}$ is cheaper and easier to handle. $\ce{P2O5}$ is also acidic and $\ce{KOH}$ is basic. There are substances that die in acidic environment (acetales) and those that die in basic environment (halogen-anhydrides).

$\ce{P2O5}$ is an acidic oxide while $\ce{KOH}$ is a base. So if you want to dry substances which are acidic you must use $\ce{P2O5}$, while if you want to dry a substance that is basic in nature, you have to use $\ce{KOH}$.

source: Stack Exchange

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authors: نورجان البياتي

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Why does Chloride attack CrO?
In certain applications some sorts of stainless steel (V2A, 1.4301) corrode when exposed to $\small{\ce{Cl^-}}$ in water, or at least it is strongly advised to use other steels.
The conditions that were explained to me as being especially bad were: redox potential >0, acidic environment.
Under these conditions, why do I need $\small{\ce{Cl^-}}$ to corrode steel? I would think that any electrolyte would have a bad effect.
Edit to add 26.7:
Most stainless steel contains chrome, a layer of $\ce{CrO}$ is what's protecting the steel. Most articles mention that $\ce{Cl^-}$ damages this layer. More resistant steels use $\ce{Ni}$ (and are more expensive because of that). So, to narrow down the question, how does $\ce{Cl^-}$ attack the passive layer?

This is an area of corrosion science. Chloride ions do not aid general corrosion  of stainless but rather the special processes called crevice corrosion and pitting corrosion. General corrosion is of large areas. Crevice corrosion is like　a hole in the metal and local corrosion effects occur. Pitting corrosion is corrosive fluid forms its own hole and similar local corrosion effects occur. The local corrosion effects for chloride crevice and pitting corrosion corrosion are complex and dependant on the fluid but a major one is the increase in acidity from  ferric ions (which normal complex with water to be acid) as the chloride ions displace a water molecule and the new complex is more acidic. This acidity damages the complex chromium iron oxide layers. Another effect is that the chloride penetrates the oxide layer and opens it up to allow corrosive species mainly oxygen to penetrate to base metal.
Nickel gives some benefit to corrosion protection of  stainless steels by substituting for iron in the iron-chromium-oxide but that's not main reason its included but rather allows the stainless steel to be austenitic- basically more ductile and easier to weld.
Molybdenum is main element that is added to improve pitting/corrosion resistance. 304 is the most common basic number stainless steel 16-18% Chromium and 8-10% Nickel  (they add letters to give more information) . 316 is second as it contain  16-18% Chromium, 8-10% Nickel and 2-3% Molybdenum- greatly reduces pitting corrosion- conditions need to be more severe to have pitting and crevice corrosion.

Perhaps chromium can form a chloride-containing complex e.g. $ \left[ \ce{Cr \cdot 2 Cl \cdot 4 H_2O}\right] $. (I know copper forms an analogous complex, but only at high concentrations of $ \ce{Cl^-} $.) The possibility of forming this complex would make chromium oxide more soluble in water in the presence of chloride than in its absence.

You need 50 rep to add comments. To me, this seems OK as an answer since it attempts to answer the question. Guess-answers can always be improved with some backing, but they need not  be converted to comments. :)

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Samsung's S-Plasma Ion - scientific basis
Has anyone come across a scientific study of the S-Plasma Ions that Samsung claims will purify air from microbes, allergens and free radicals? I would prefer publicly available research findings.
The following is a quote from http://www.samsungspi.com:

The Samsung SPi generates Hydrogen atom ($\ce{H+}$) and Oxygen ion ($\ce{O_2-}$) which reacts together to form Hydroperoxyl radicals which the positively charged biological contaminants then reacts with three hydrogen elements of its protein structure before forming water vapor. Also, the Hydrogen atom ($\ce{H}$) will combine with $\ce{OH}$-radical to form a water vapor and neutralize harmful $\ce{OH}$-radical.

It also matches the description found on the product itself. See http://winniekepala.com/2011/03/ionizer-review-samsung-virus-doctor/ .
Here's a quote from the video on the same site:

$\ce{H+}$ is born with the help of that plasma in this plus electrode. $\ce{H+}$ is hydrogen with a positive electric charge. Through the air, the $\ce{H+}$ moves to the minus electrode which generates electrons. In here, electron meets $\ce{H+}$ and active hydrogen like me are born.... The electron that is generated in minus electrode in the SPi ion device where [I] am born is naturally combined with oxygen. That is oxygen ion ($\ce{O2-}$). (Active hydrogen and oxygen ion together transform into $\ce{(H-O-O)-}$). $\ce{HOO}$($\ce{HOO-}$) ion envelops the positively charged airborne bio-contaminants and react with three hydrogen constituted of protein components, forming water vapor. This modifies the surface protein or breaks the cell membrane to remove bio-contaminants. $\ce{HOO}$($\ce{HOO-}$) ion reacts only with single cells such as virus or germ but does not react with the human body... $\ce{OH}$ radical is an active oxygen species and one of the highly reactive and strong oxidants. It can be harmful to human's body if exposed for a long time. (Active hydrogen combines with OH radical to form water vapor).

(It has numerous grammatical mistakes which I just copied verbatim).
I would like to know what are the various compounds and radicals mentioned in the explanation, and how the reactions above take place. What is the plasma in the positive electrode, and what is the potential difference between the electrodes that allows the generation of the active hydrogen and oxygen ions? What are the real chemical states of the active hydrogen and the oxygen ions? How do they combine to form the $\ce{HOO-}$ ion? What is the $\ce{HOO-}$ ion and how is it stable enough to go around the room? How does the active hydrogen combine with the $\ce{OH}$ radical?

Could you elaborate the question a little bit based on what you know about these ions?

The point is that I don't know anything about these ions! And I'm not sure if I'm taking a risk now by copying this from Samsung's website. It's only for educational purposes though.

I don't think there should be any problem with you copying that, you have cited the website it came from and based on fair use, it should be fine.  I think this will likely help it tremendously to get an answer, tbh, as it's a lot more info for someone to go on.

if you can take a look at the main web site, then it seems they make $\ce{HO2}$ and not $\ce{HO2-}$. This a peroxy radical which exists at trace levels in the ambient atmosphere, usually from reactions of light with oxides of nitrogen and volatile organic compounds. Its really reactive and I don't know for sure how long it will last in the air of the room. It will however react with $\ce{OH}$ radical and make water and $\ce{O2}$. $\ce{HO2 + OH -> O2 + H2O}$. $\ce{HO2}$ will also react with $\ce{NO}$ to form $\ce{NO2}$ and then depending if there is sunlight could make ozone. One of the issues for the machine will be that $\ce{HO2}$ will need to made at some point and therefore one of the first molecules it will encounter will be itself and therefore make $\ce{H2O2}$ (hydrogen peroxide). I think getting one into a lab and testing it would be a start.

Welcome to Chemistry.SE @fumifugium. We have MathJax imbedded into our site to help you format math and chemistry. I have edited your answer to make the formulas look like formulas. Check out our FAQ to learn how to do this for yourself.

@BenNorris: Just a note, you can link to the FAQ just by using [faq] in a comment (another useful one is [ask] for [ask]) See this post for more. :)

@Manishearth - thanks for the tip. How do I get that greyed background for the links to show the markup?

Never mind. I figured it out.

Thanks for pointing out the possible inaccuracy of $\ce{(H-O-O)-}$. However, where on the main web site did you look? This link also says the same thing.


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Difference between lab-grade and food-grade purity?
What is the difference in purity betwen lab-grade and food-grade when talking about potentially consumable compounds?
As an example let's take ascorbic acid powder: if it's marked as lab-grade, does that mean it is also fit for consumption?
In particular, I'm wondering whether lab-grade supersedes (and/or improves) a food-grade labeling, or if they are two totally different criteria? Can I say lab-grade > food-grade?
If I can't say lab-grade > food-grade, then what is the actual difference between the two?

So-called food-grade stuff is meant to be consumed by humans. It's important that every single compound of a foodstuff is food-grade or at least GRAS. In the European Union the law giver has defined the state when a compound is good for consumption as "safe".
Lab-grade chemicals often contain toxic materials, so when something is called lab-grade it is not safe per definition.
For example: $\ce{NaOH}$ is widely used in preparation of foodstuff as well as in the lab. Now it's a fact that lab-grade (standard) $\ce{NaOH}$ (Sigma Aldrich) contains a certain amount of heavy metals. On the other hand, food-grade $\ce{NaOH}$ will not contain any heavy metals (usually food-grade $\ce{NaOH}$ will have a purity of more than $99\%$ while standard $\ce{NaOH}$ pellets are sold with a purity of less than $97\%$). So using lab-grade $\ce{NaOH}$ in foodstuff is a crime (at least in the EU).
So my advice is: Do not use lab-grade materials for any food-related applications.

Perfect. Nick's answer was very good, but I think you make it even clearer. The impurities are what matters here. So lab grade is not > food grade. It's a completely different classification and both classifications have their own (unrelated) goals with regards to the impurities left in the material.

Generally laboratory grade reagents have not been prepared with human consumption as a consideration, but with regard to their use as a chemical reagent.
For example, $95\%$ ethanol is mostly ethanol and $5\%$ water. This is fine for cleaning equipment, or using as a solvent for a TLC. If you want to use ethanol as a reagent, however, you want to use absolute ethanol (which is much more expensive). Absolute ethanol is NOT safe for human consumption though, as the water has been removed from the $95\%$ (which is the azeotrope, i.e. as far as you can get through distillation of water/ethanol alone) through a process which involves adding additives such as benzene. Benzene won't cause side reactions in the way that water would, and so absolute ethanol is much better for a lab reagent than $95\%$, but worse for consumption.
In the same way, any chemical which is prepared for use as a reagent will contain impurities which don't generally cause side-reactions, while for consumption you want the impurities to be non-toxic.

Most laboratory grade reagents will also say "not for human, food, or pharmaceutical use".

Background: I'm a chemist and have worked for companies in Pharma, Research Chemical Supply, and Dietary Supplements.
Lab Grade is different than Food Grade, yes. Normally Food Grade is a LOWER quality than Lab Grade.
Lab Grade normally has exhaustive testing on it to determine all impurities. The lab is using this as a reagent and they need to know all that.
Food grade means that it won't cause human disease. So really all they need to know is if there are too many heavy metals or too many bacterial contaminants. The limits are typically higher for food grade.
For instance, when I worked at a Pharmaceutical facility we worked to get a material accepted by FDA as a food grade material and not a Pharma Grade or Reagent Grade. The main reasons were insect and metal parts. The FDA regulations are much more relaxed for food grade and you can have quite a few more of each in the product for food grade than either of the other two.

Welcome to Chem.SE! Thanks for posting; please stick around! We'd be glad for you to contribute your expertise to the site.  Check out the [tour] and [help] for more information about the philosophy of the site and how it works and once you reach 20 reputation feel free to drop by the main chatroom and say hello. Enjoy!

I purchased lab-grade Potassium Bicarbonate with a grade of 99.5% and then compared it to the purity of food-grade and found food-grade is listed as 99.0% purity. Very interesting.

Food grade might usually have lower purity but the point of being food grade is which impurities. Some impurities that matter when people are going to eat the substance are very different to the ones that matter in the lab. So absolute ethanol is the lab is very pure compared to whiskey but it often contains benzene which people really should not consume.

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How can I be safe and protect myself using trimethyl tin chloride?
I will be using trimethyl tin chloride (1M in THF) in a few days for the first time. Looking at the MSDS, it looks like it is very toxic and has noxious vapors. Are there any more standard precautions for working with it other than the usual?
E.g.,

Check for proper airflow in the hood
Wear lab coat, safety glasses, and double gloves

I've never worked with this (which is why this is a comment rather than an answer), but if you're worried that your fumehood and other safety preparations aren't going to cut it, you may want to see if there's either a proper glovebox installation or a portable sealable glovebox available. They can be a bit of a pain to work in, but hey, it's better than tin poisoning.

Use a glove bag in a hood, you being double gloved. Properly dispose of all contacted materials, and both sets of gloves.  Don't get contaminated during disposal- bag the bag in the hood.  I used powdered anhydrous sodium saccharine in a synthesis.  I was very careful and it still tasted it.

In addition to the safety precautions mentioned in the question, the following can increase the safety for the experimentator:

Keep the fume hood as closed as possible with the head outside and the arms inside. Usually this keeps fumes from entering the body, since the fume hood airflow goes the other direction.
Work quickly. While this not only minimizes side reactions of the reactant with the air and other things, it also ensures less exposure time, minimizing vaporization.
As Aesin mentioned in the comment, if you have concerns about the infrastructure, you could always perform the experiment in a glovebox.

In the end, you have to feel confident that the safety measures are effective. If you don't, you will have fear of the reaction, and inevitably be less concentrated and more prone to errors. In combination with "work quickly" this may well become a disaster.
Maintain a healthy portion of respect, keep quenching agents close by and execute the reagent transfer with a steady hand.

I suggest you to read carefully this sheet nj.gov I'm quoting the main infos:

WAYS OF REDUCING EXPOSURE

Where possible, enclose operations and use local exhaust ventilation at the site of chemical release. If local exhaust ventilation or
  enclosure is not used, respirators should be worn.
Wear protective work clothing.
Wash thoroughly immediately after exposure to Trimethyltin Chloride and at the end of the workshift.
Post hazard and warning information in the work area. In addition, as part of an ongoing education and training effort, communicate all
  information on the health and safety hazards of Trimethyltin Chloride
  to potentially exposed workers.

WORKPLACE CONTROLS AND PRACTICES

Where possible, automatically transfer Trimethyltin Chloride from drums or other storage containers to process containers. Good WORK
  PRACTICES can help to reduce hazardous exposures. The following work
  practices are recommended:
Workers whose clothing has been contaminated by Trimethyltin Chloride should change into clean clothing promptly.
Do not take contaminated work clothes home. Family members could be exposed.
Contaminated work clothes should be laundered by individuals who have been informed of the hazards of exposure to Trimethyltin
  Chloride.


text:
What does the ZZE tag in an Chemical Table file refer to?
I have a number of Chemical Table files which include a "M  ZZE" tag in the properties block.
For example, in Galactobuxin, I have:
M  ZZE  1  31  23
In Rubrobrassicin there is:
M  ZZE  3  28  23  36  23  48  23
In R-hexanoyl CoA there is:
M  ZZE  1  55  23
The typical format for such tags is that it starts with an M, followed by two spaces. Then there is a three-letter code, in this case ZZE. Other codes include ISO (for isotope) or CHG (for charged atoms in the compound), then a number which indicates how many records on that line. This is followed by the actual data.
So
M  ZZE  3  28  23  36  23  48  23
would probably mean that there are three sets of ZZE data, being "28 23" "36 23" and "48 23". In every record I have the second number is 23.
The ZZE tags seem to occur only for compounds containing ribose rings, and the first number is the index of one of the bonds inside the ribose ring. Specifically, of three three bonds adjacent bonds with a stereochemistry marker, it always refers to the one in the middle:

I'm guessing the 23 is some property of the bond: so in rubrobrassicin the 28th, 36th and 48th bonds (there's one, and only one, ZZE record for each of the ribose rings in rubrobrassicin).
Does anyone know what the ZZE property in a chemical table file refers to, or what it indicates? I can't find any information on the format.

Jivlain, i'm sorry but i don't see the ZZE entries in the SDF files on the PubChem files you're linking to.

Usually, i'd use Z (= zusammen) and E (= entgegen) to denote the stereochemistry at double bonds but that doesn't play a role for the structures you mentioned. Furthermore, to my knowledge, E and Z as described above are not part of the MDL molfile specification.

@KlausWarzecha: Indeed - I have been sent some SDF files which contain ZZE tags (I just linked Pubchem for convenience's sake). Your zusammen/entgegen suggestion seems quite plausible, however - I had a closer look and noticed that the ZZE tag always seems to refer to the bond in the middle one of the three bonds with a sterochemistry marker.

Could you possibly upload your SDF file somewhere (for example, filedropper.com) and share the link? Also, it may be instructive to know where/how the individual who sent you the SDF file generated them. It may be a designation specific to the program used to generate the file. As @KlausWarzecha indicated, the SDF files on pubchem don't contain the ZZE entries, so we need more information about the origin of the SDF files or the actual files themselves.

Jivlain, i had a look at the CT file specification as provided by Accelrys and i found nothing on Z or E. @JM is right, concerning the origin of the SDF data. Knowing how the files were generated might really help to solve the puzzle.

@JM I've linked some SDFs and also a picture with the relevant bonds highlighted (the ones that I've drawn a circle on).

The current version of the CTFile spec can be downloaded for free from Accelrys. Based on the files you have, you'll want to look at the "V2000" section.
Short answer: I believe [M  ZZE] is a vendor-specific property key, possibly introduced by ACD Labs, given it's the package cited in the header on line 2. The CTFile format spec for V2000 files says nothing about an "M  ZZE" property key.
The CTFile spec says that property keys that are not recognized by parsing software are to be ignored. This makes vendor-specific extensions feasible.
You seem to be on the right track in your interpretation. The Galactobuxin entry has this line in the properties block:

M  ZZE  1  31  23

Breaking this down:
{M  ZZE} is the mystery properties key.
{1} is the number of data entries on the line.
{31} is the 1-based index of either an atom or bond.
Atom 31 is one of the rather boring methoxy groups.
On the other hand, bond 31 is a real problem child:

This file is abusing the meaning of 'wedge bonds' by trying to use them to indicate chair conformation perspective, rather that what they're intended for - flat stereoprojection. My guess is that the file is trying to annotate this bond.
The only question left is - what does {23} refer to?
Bond 23 is unremarkable. Atom 23 is one of the origins of a wonky wedge bond. But it's not clear why it should be associated with bond 31.
My best guess is that the software that created this file knew it was abusing wedge bonds and wanted to include a flag on bond 31 so that software reading the file wouldn't misinterpret.
Given that you're seeing the same {23} in every file, this could be a code to specially process bond 31's stereo flag.


text:
Chemical synthesis via 3D printing?
I happened to stumble upon this Nature Chemistry article today. It claims that

[...] use a 3D printer to initiate chemical reactions by printing the reagents directly into a 3D reactionware matrix, and so put reactionware design, construction and operation under digital control.

Unfortunately, I cannot obtain the full article to read, but even the abstract does not seem plausible in my opinion. My questions are...

How can one "print reagents"? What is meant by this?
Most synthetic reactions require hours of refluxing and whatnot, not to mention isolation/purification afterwards. How can all of these techniques be performed with a printer?

Perhaps I'm not "getting it", but the abstract of the article isn't providing me with the explanations and theory I'm looking for. How can something like this be possible?

References

Symes, Mark D., Philip J. Kitson, Jun Yan, Craig J. Richmond, Geoffrey J. T. Cooper, Richard W. Bowman, Turlif Vilbrandt, and Leroy Cronin. “Integrated 3D-printed Reactionware for Chemical Synthesis and Analysis.” Nat Chem 4, no. 5 (May 2012): 349–354.

The authors use a 3D printer to generate the labware, and to place the reagents in its inlets.
So we're not printing standard ink on paper but using a the print head to exactly place very amounts of monomers at certain positions. These monomers cure (polymerize) upon heat or irradiation to a solid resin. By this, small reactors of arbitrary design (think in 3-neck reactors) with a volume of ~ 10 mL may be printed layer by layer.
Apparently, it is even possible to create reactors with windows to allow for later reaction monitoring or print and assemble electrodes for spectroelectrochemistry.

How can one "print reagents"? What is meant by this?

This is all about precise positioning of the printhead! Thus same printhead used for the ink can be used to fill the inlets (think in dropping funnels of normal labware) of the reactor with componends whose reaction you want to monitor. Instead of dropping, material transfer is achieved by applying a vacuum to the reactor through a self-healing hole.

Most synthetic reactions require hours of refluxing and whatnot [...]

Well, a lot of reactions do proceed quite well at room temperature, think in

Diels-Alder reactions of furanes and maleimides
electrochemically triggered reactions
photoinduced electron transfer reactions

just to mention a few.

[...] not to mention isolation/purification afterwards. How can all of these techniques be performed with a printer?

Oh, you got that one slightly wrong. The "printer" is just used to build the reaction sphere (the flask, if you like it) and to place the reagents in the right inlet of your vessel. The rest is "just chemistry" ;)
I hope i made myself clear enough - at least after the n-th edit ;)
I think the concept is pretty cool.

Great answer! But still curious: does this mean that the printer contains ALL reagents that it may ever need in order to synthesize a specific drug?

Thanks :) The printer thing really got you :) But as far as i understood: yes, at least for the cases reported. The printer can serve as the glass blower to build the apparatus and the experimentator to fill the inlets/reservoirs with the reagents necessary. On the other hand, i guess this isn't a must. One could probably also fill the reactor with a syringe or deliver the reagents with a piston or peristaltic pump. It is seemingly just very convenient (and clever) to use the precise positioning and digital flow control of the printer head for that task too.

I'm a mechanical engineer, not a chemical engineer nor chemist.  I don't think they used the printer to actually fill the reaction vessel.  I think they did that with a pipette.  This reactionware is somewhat larger than microfluidic devices.  They 3d print a structural vessel with polypropylene and line it with silicone rubber IIRC.  The reagents are in seperate areas of the vessel and are moved into the main reactor by various means.  A previous poster suggested the vacuum technique.  I think they also did it by rotating the vessel from one position to the next.  There was also a bit about creating catalyst surfaces by mixing a catalytic material in with the silicone and printing a wall of the vessel with it.  Very interesting stuff.

Good point! I might need to read the article again after 5 yrs ;-) Btw, welcome to Chemistry.SE! Did you consider to register (free, no obligations)? If you planned to come and stay, which would be a good decision, these unregistered cookie-based accounts start to suck pretty soon.


text:
What is the molecular structure of the coating on cast iron cookware known as seasoning?
When a piece of cast iron is repeatedly covered in fat and exposed to heat it develops a hard, black, non-stick coating referred to as "seasoning".
What exactly is this coating made of molecularly? How is it bonded to the cast iron substrate? What is the chemical process that transforms a liquid fat into this coating? And what if any characteristics of the fatty acids or triglycerides used have an effect on the final coating or how well the it bonds to the iron?

Mixture of various black oxides (approx. Fe3O4), phosphates(?), carbonised oil of any sort. Sticks like hell, and scratches repair themselves during further use. :-)

Some hints:

if you mean that "what exactly" in the original sense of the words,
then the answer is: nobody knows and will ever know.
If you are content with some more vague answer, then
look at the pyrolysis chemistry occuring from 150 °C
up with fats, proteins and starch/sugars, and all that
at once. This is a vast field of chemistry, and all
is gives to a trained chemist is a idea of the structures
of that coating.
Bonds to iron will not occur, first because on cast iron and
on steel pans there is always a layer of iron oxide,
on cast iron  you find a glaze formed by fusion of the
molds finish layer (often containing graphite and clay)

Cast iron pans that develop seasoning are not enameled. There is definitely iron bonding, because the same seasoning doesn't develop on aluminum or other cookware, and the seasoning is developed before you let a pan rust.

@yincrash Why definitely? He explains that the not even is iron in the first place in the third item. But iron oxide, and molten remains of the mold in which the pan was cast into. Similar to molten stone or so.

Mostly- all the oils coated are of DRYIG oil -meaning they are polymeric. eg- the Linseed oil etc -they polymerize on heating. Since the cast iron itself is slightly porous- these oils penetrate along with the iron or , could be reacting with the iron salts, once they break down to fatty acids, even partly.
This also explains - why this is not happening in pure steel cooking vessels.


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Can nanotechnology reduce electrode losses in electrochemical cells?
I'm just curious about this topic, as most portable, "clean" energy sources or carriers (in the sense of not being a primary energy source -fuel cell hydrogen, for example) are in essence electrochemical reactors, and most of the current work in the field is aimed at reducing irreversible losses at the electrodes, be it because of discharge overpotential, or polarization because of mass transfer along the membrane electrolyte; and of course, in making this structural elements cheaper.
Also, a lot of proposed and existing waste treatment and recycling processes aimed at recovering rare and/or active metals (such as lithium from Li-ion batteries) involve at least one electrochemical unit op.
I wonder if nanotechnological advances, such as increasing the surface area of electrode materials or increasing the effecive diffusivity of relevant species in solid porous electrolytes, can reduce this irreversibilities, perhaps making use of cheaper raw materials than noble metals (i.e. platinum). I would also appreciate any references to current work in this applied field.

if it can be reduced, then why not.

I was wondering if someone here could give me a more comprehensive answer and/or references. But what you wrote is more than I could ask for :)

Nanotechnology offers enormous opportunities for improving a huge range of electrochemical processes.
The use of platinum as a catalyst is been widespread throughout the electrochemical industry, helping reduce the activation energy of reactions such as in hydrogen and methanol fuel cells. However, platinum is a rare and expensive metallic element. Given that the reactions only take place on the surface of the catalyst, the use of nanoparticles of platinum reduces the amount of platinum needed to achieve the same surface area, thereby achieving the same performance at a fraction of the cost.
Another interesting application is in lowering the operating temperature in solid oxide fuel cells.
A good reference on Nanotechnology in Electrocatalysis for Energy, which covers background fundamentals and some applications, can be read on Google Books for free.

Speaking specifically on lithium-ion batteries, it depends. Smaller cathode particles tend to allow more lithium diffusion due to reduced diffusion path length, which means better kinetics. However, smaller particles also have larger surface energies, meaning you might have to pay a high penalty in the form of overpotential.
source article


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Identity of a green substance formed during electrolysis reaction?
I did a short experiment in which I connected a $1.5\:\mathrm{V}$ battery with 2 paper clips into a bowl of water to which I added salt. After a while, I saw a yellowish-greenish cloud inside the water around one of the paper clips, and when I took it out, the part that was in the water was more black than the rest of the paper clip. After a while, some of it turned into a greenish color.
What caused the paper clip to change its color?

More information would be helpful. Were they both steel paper clips? What is the electrolyte?

I think the paper clips are made out of steel, and the electrolyte is NaCl.

Mono's answer and the comments are mostly on the money. If you want to try this again, you need both a different electrolyte and electrodes. Simplest would probably be magnesium sulfate (Epson salts) and carbon rods, respectively. Make certain to wrap the connection of the carbon rods and the wires to your battery with waterproof tape to avoid reactions with the wire (i.e. the comments about your paper clips). Good luck--I love this experiment!

I once did this experiment, and I too got the same result. My bet would be that it's due to the $\ce{Fe^{3+}}$ ions that were removed out of the paperclip.

Without further details of your equipment, this might be a possible answer. I assume the electrolyte you added was table salt (NaCl).
What you saw as a yellowish cloud was gaseous chlorine which discharged on the clip you connected to the positive electrode (cathode) of your battery. The reason why you did see this, instead of seeing just colorless bubbles of oxygen, is because of something called "overpotential". It's a joint property of the species in your cell and the material and surface termination of the electrode; in your case, the clip (made of wire steel) has a lower overpotential for the discharge of chloride anions (from the salt you added) than oxygen.
Elemental chlorine is a very reactive oxidant, and you had a thin chlorine film around your clip electrode. The reaction between chlorine and iron in aqueous solution most probably yielded iron (II) chloride tetrahydrate which is the reason you saw the greenish colour of the previously submerged fraction of the clip.

How can the chlorine be in gaseous state if it's in the water?

Ferrous chloride solutions are yellowish; I thought you also had seen a "cloud" of chlorine bubbles discharging from the electrode into the air above.

http://www.youtube.com/watch?v=iWpfHkWr5DY

By "cloud" I meant a cloud inside the water, a yellow liquid but only around the paper clip.

I used a low voltage battery so I barely saw any bubbles, but I could see the yellow liquid clearly.

Then it's most probably the dissolution of corroded iron from the clip into ferrous chloride.
This is the main reason why carbon steel electrodes aren't used in industrial electrolysis of aqueous solutions.

Elemental chlorine is also soluble in water at a level of about 3 g/L, which could also account for the yellow color in solution.

Why are you so sure it is chlorine, not oxigen? For low NaCl concentration I would first elementar oxygen to oxidise the anode, resulting in some form of iron hydroxide.


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Copper Chloride: neutralizing and disposal
I plan on etching some PCBs with hydrochloric acid and hydrogen peroxide, which will therefore produce some kind of copper chloride (green color) which is highly toxic if released into the environnement (it may be useful as an herbicide, but I don't have any garden right now).
While cupric acid (obtained from hydrochloric acid, hydrogen peroxide and some PCBs etching) may be bubbled with air/oxygen to re-activate it to etch again other PCBs I'm not sure this may what I want and I seek suggestions in case I need to dispose of this solution.
Basically I heard (though none were actually on chemistry sites, except the 1st one):

Acid/Base neutralization: pour some baking soda to take care of the chloride
Pour some alluminium balls to get back some of the copper (didn't really understand how this works)
Copper electrolysis: basically get solid copper at the cathode, creating gaz at the anode and leaving (?) some residues that can be neutralized with baking soda (?)

Any insights on this? I think once the copper and chloride are taken care of, it shouldn't pose a huge problem to pour it down the drain or leave it outside to evaporate.
Specifically:

Which metal should I choose for the anode / cathode ? In the first source it speaks about reactivity (and since copper is less reactive than hydrogen, copper is produced at the cathode). No idea about the metal I should use though.
Would adding alluminium balls into the solution help make it more "eco-friendly" ?
Any other suggestion is welcome

Sources:

Electrolysis
How to etch PCBS - see comments
Recovery of copper using electrolysis

Regarding #2: Adding $\ce{Al}$ or $\ce{Mg}$ metal to $\ce{CuCl_2}$ solution reduces the $\ce{Cu_2^+}$ to $\ce{Cu}$, and you then have $\ce{Cu}$ metal (filter out and recover) and $\ce{AlCl_3}$ or $\ce{MgCl_2}$, respectively, in solution. Both solutions are safer in small quantities to flush down the drain.
Regarding #1: $\ce{2NaHCO3 + CuCl2 -> CuCO3 + 2NaCl + CO2 + H2O}$. $\ce{CuCO3}$ is insoluble, but should still be disposed of at an appropriate facility.
In both cases, you get an insoluble copper compound (or metal) rather than a soluble one.

Can I filter CuCO3 too or is it too small to filter?

Sure, you can filter the CuCO3.

Thank you. Would I need some special kind of filter or something like a coffee filter is enough for this job (Cu / CuCO3)?

I'd try a coffee filter, should work.

Reducing copper by aluminium or magnesium is nonsense. Thes metals would react much too vigoously! That process is done with eiger Zinc or, cheaper with the aid of iron (scrap iron). The handling of used pcb liquids is discussed since years ad nauseam in electronics literature!

@Georg: Well, one of my lab textbooks has as a very standard experiment of the oxidation of Cu to various states, and then back to Cu. The last step uses Mg and does not mention great danger in doing this.  As for all reactions, caution is in order.  Your alternatives are certainly correct in terms of reactivity (and are a good answer for OP's #3), but Al and Mg are not wrong.

Depending on where you are, there may be places where you can hand over hazardous waste for proper treatment. E.g. in Germany the communal waste collection centres take properly labelled hazardous waste. For private stuff this is paid for by the general waste fees, so no extra fees are charged. If you do this on a commercial basis (= large scale), you'll have to pay for it, of course (but in this case you shouldn't have to ask how to deal with this in a forum...).
In any case, I'd recommend that you let (most of) the water in your solution evaporate so you get either a concentrated solution or the salt (chloride or carbonate) in order to reduce volume and mass of the hazardous waste.

Commercial users will sell that liquors to someone reprocessing the copper. Or after cementing the copper will sell the copper-sludge to an affinery.

@Georg: you'll need to be quite above the baking-soda-for-neutralization scale in order to reach Cu amounts where you can ask that someone who comes and gets the stuff (under German/EU environmental regulations) should pay you... From the OP I'm not even sure he reaches the g Cu scale.

The best way to dispose of this is to use ascorbic acid (orange juice even?) to reduce the copper, causing it to precipitate as a metal. To deal with the corrosive iron chloride that is left you should add baking soda or calcium carbonate (limestone) to neutralize the liquid until the addition of either produces no bubbles.

I made blue water via electrolysis of water (with salt, using carbon rods wrapped in copper wire as electrodes), which I guess is copper chloride. I thought this was just going to be a fun little experiment; now I have to stay up late googling how to dispose of this. Anyway, I followed your instructions--squeezed a lime into the blue water and added baking soda until no more bubbles. Can I pour it down the drain now? It's changed to a light green color.


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Radiation time vs. half life period
Is radiation time the same thing as half life period?
Also, if I'm talking about excited species in a gas, how would the gas interaction affect the radiation time, are there some formulas/approaches?

Can you please clarify what you mean by 'radiation time'?

Radiation time might have to do something with Mean Half time of the radioactive decay?

In the context of photochemistry, radiation time or half life period are not used, as far as I know.
The only relevant technical term i'm aware of -in that context- is half-life (time), which describes the time needed for the concentration of an excited species to decrease to 50% of the initial value, cf. S. E. Braslavsky, Glossary of terms used in photochemistry,
Pure Appl. Chem., 2007, 79, 293-465
[DOI].
For longer-lived excited triplet states in the gas phase, radiationless deactivation upon collision with a ground-state molecules will play a role.
Thus, the lifetime might be correlated with the mean free path of the molecules.

You missed the lifetime of a radioactive species, which works on the natural logarithm (e) instead of base 2. It is normally represented by tau.

@Canageek You're absolutely right! I only thought in electronically excited species. I guess i'm still a bit biased ;)

I worked in a radiochem lab last summer, so I had to work with both systems when looking things up.

In the context of radioactivity (which might be what you meant by “radiation”), the half-life of a radionuclide is a well-defined term and commonly used. “Radiation time”, on the other hand, is not used to describe radionuclides.


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How can I compute the temperature change from the de-excitation of excited particle species?
The motivation for the question in the title is that I have some formula given,
$$\Delta T=\frac{3410\ \gamma_{O_2}}{2.5+\gamma_{O_2}},$$
where $\gamma_{O_2}$ is the molar fraction of oxygen in a binary micture and the ratio of oxygen and rate gas heat capacities are supposed to be 7.5, if that helps. And there is another one with a factor $\frac{1-\eta}{\eta}$ multiplied, where $\eta$ is a singlet delta oxigen efficiency.
Where does this formula come from?

Could you provide the reference for that formula given. Is it from a journal article or a textbook? If you could provide a little more context and some evidence of what you have been able to come up with on your own, it would help folks generate a better answer.

@BenNorris: equation (9) and the one before that.


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Copper substituting silver in Tollen's reagent
I am an artist attempting to use Tollen's Reagent in a number of my projects.
Is there any reason I could not substitute Copper Nitrate for the Silver Nitrate in the reaction? Would the copper adhere to the glass?

Manishearth points to the right direction, the crucial parameters to look at are the standard electrode potentials. One should however take into account that the relevant species that get reduced under the condition of the Tollens' reagent are not the aquo- but the amine complexes, namely $\ce{[Ag(NH3)2]+}$. and $\ce{[Cu(NH3)4]^{2+}}$, respectively.
Unfortunately i don't have the data at hand.
Notably, a copper-based reagent comparable to the Tollens' reagent exists:
Fehling's solution. However, in this case, Cu(II) is reduced to Cu(I) only but not to copper metal.
In the end, you just might give it a try and post the results :)

(Note that if you're going to give it a try, copper sulphate is usually easier to get hold of and shouldn't make a difference to the reaction as far as I can tell.)

I got 'copper mirror' with Fehling's solution once, but adhesion was not best. I didn't care, thought, so this might be my fault. Anyway, Fehling's solution is much, much easier to handle then Tollen's reagent.

For the particular reaction where you get a silver mirror with an aldehyde, I doubt it'll work for copper, due to the reduction potentials. I'll take acetaldehyde $\ce{<->}$ acetate as an example here, but I believe that the reduction potentials of most aldehydes will be similar. The nitrate is a spectator ion here, so we need not consider its presence.
$$E^o_{\ce{Ag^+,Ag}}=+0.80$$
$$E^o_{\ce{Cu^{2+},Cu}}=+0.34$$
$$E^o_{\ce{CH3COO^{-},CH3CHO}}=+0.580$$
Our reaction involves the metal ($\ce{Ag}$ or $\ce{Cu}$ being reduced, so it must have a higher reduction potential than the acetate/acetaldehyde pair. Of the two metals, only silver has a higher SRP, thus the reaction involving silver nitrate is spontaneous--whereas the copper nitrate reaction is  infeasible.
Of course, you could find some substrate other than an aldehyde which reacts with copper nitrate to achieve the effect you want--though I'll have to look up some stuff to figure this out. I think the Tollen's reaction is a single step everyday redox reaction, which means that we just need to look for a stronger reducing agent which doesn't harm the glass.
Actually, it looks like the standard process for plating nonconductors with metals is to give them a conducting layer by dipping them in something, and then using electroplating. Not sure about this, though.


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Can I use one alkaline battery to recharge another?
Unfortunately, I don't have batteries to waste to test this myself, but I am curious...
Let's say you have two identical batteries. AA Duracell.
Now let's say you run one of the batteries through an LED until the LED starts emitting at half the intensity it was emitting at when you first started (effectively draining the battery to half voltage)
Now, lets' say you touched the positive ends of both batteries for a reasonable amount of time, and then did the same for the negative ends.
Given that the used battery would have a much lower voltage, raw electrical current would travel from the unused battery into the used battery.
So this brings me to my question -- Would the batteries now effectively have the same amount of energy in them? I am not just talking voltage, but chemically too, or would the similarity in voltage on both batteries immediately disappear if either of them were used (because while, yes, either end of the battery does contain equal electrical charge to its counterpart, either end of each battery does NOT contain equal chemical potential)
To simplify this, would touching two batteries as described above effectively drain the unused one, and charge the used one?

One thing is clear: you can't charge batteries that aren't rechargeable. And since you can hardly "waste" a rechargeable battery, you can safely test this yourself as long as you have any rechargeable batteries available.

@GeorgesOatesLarsen: Actually, it was an answer, but I converted it to a comment since it didn't really answer the question from a chem point of view (plain yes/no answers aren't really allowed without some scientific explanation)

First, some batteries are rechargeable, but some are not. However, even if we talk about rechargeable one, then
1) some degradation during recharge usually occurs
2) deep equality will require very long time to achieve

In fact, accumulators oscillate when charging each other. The result is always two discharged batteries.


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How did Michael Faraday melt Sodium Chloride?
Michael Faraday melted sodium chloride in a silver spoon and ran an electric current through the molten salt, which released chlorine as a gas and left pure sodium behind.
Sodium chloride melts at $800\ \pu{^\circ C}$. How did Michael Faraday melt the salt?
*I know he invented the bunsen burner, but did he use that?

If you're incredulous that such temperatures were achievable those days, you may want to check out the average flame temperature of a candle-- 800 degrees isn't a big deal. Or are you asking for the exact setup he used in his electrolysis apparatus?

I am asking for the exact setup. I found an open flame to be insufficient.

I came across an explanation of why NOT to use this method: http://yarchive.net/explosives/sodium_make.html

The answer is given in Faraday’s own work, Experimental Researches in Electricity, Volume 1, freely available thanks to The Project Gutenberg:

399. Other substances, which could not be melted on glass, were fused by the lamp and blowpipe on platina connected with one pole of the battery, and then a wire, connected with the other, dipped into them. In this way chloride of sodium, sulphate of soda, protoxide of lead, mixed carbonates of potash and soda, &c. &c., exhibited exactly the same phenomena as those already described: whilst liquid, they conducted and were decomposed; whilst solid, though very hot, they insulated the battery current even when four troughs were used.

The text above shows that the lamp in question is the spirit lamp, and the whole aparatus is described here:

The crux of it is that an open flame might not have enough heat, so a conventional flame is simply contained into a more restricted space.

license: Creative Commons - Attribution Share-Alike - https://creativecommons.org/licenses/by-sa/4.0/


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How would I neutralize a PEP (Phosphoenolpyruvic) acid?
PEP (Phosphoenolpyruvic acid) is acidic but I need to make it neutral. I need a solution (pH 6-8) of PEP in a final 1 mL volume with a concentration of 100 mM.
How would I make this?

For our work, we also need to make 100 mM PEP and we neutralize it using KOH. Traditionally, using Na+ based salts isn't ideal for biological assays so K+ salts are better for our applications.
Alternatively, if you're adding the PEP to a different solution, you can buffer the reaction solution. While the Phosphate does have a good amount of buffering capacity, effective concentrations of the molecule will be low so you ultimately won't be adding much.


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Why add water first then acid?
From school, I remember a very important rule: first you need to pour the water and then the acid (when you need to mix them) not vice-versa. This is because otherwise the aсid becomes very hot and splashing may happen.
So, why does it get hotter when water is poured into it? What reaction takes place?

Henry's resting long and placid He added water to the acid He didn't do what he oughta and add the acid to the water - not specific to the question, but relevant to those needing to remember what order to add things!

German mnemonic "Erst das Wasser, dann die Säure, sonst geschieht das Ungeheure" ("First water, then acid, otherwise a monster will appear") indeed exists for a good reason.

This is mostly the case for sulfuric acid. Commercially available sulfuric acid is dense (~1.8 g/ml) and when water is added, it may not mix. In this case a layer of hot weak acid solution is formed, which boils and sprays around. When acid is poured into water, it flows down the flask and mixes much better, so no boiling occurs.
The reason this occurs is due to the large amount of energy released in the hydration reaction of sulfuric acid ions. Do not believe that heat comes from dissociation, as the dissociation of acids, bases, and salts always consumes energy. The energy is released from subsequent hydration, and the release may be high, especially if $\ce{H+}$ or $\ce{OH-}$ ions are hydrated.

The big difference is the boiling point of the minor phase.

This happens with strong acids and bases. Strong means the acid almost complete dissociates into ions; e.g. $\ce{HCl}$ to $\ce{H+}$ and $\ce{Cl-}$.  When the dissociation occurs, it releases a great deal of heat (i.e. exothermic) because of the subsequent hydration of $\ce{H+}$ to $\ce{H3O+}$, as the other comments have indicated. The reason you add acid to water is that if you add water to acid, the first drop of water reacts immediately, and the heat might be high enough to boil the water instantly, which could spray acid out of the container.

The same argument could be made about the first drop of acid.

@DmitryGrigoryev while I'm no expert chemist, I think that the first drop of acid in to water dilutes quickly, such that there is no longer a strong acid present, and the rate of reaction is slowed. The other way around, it takes many drops of water added to a strong acid to dilute it enough to slow the reaction, so all of that initial water added to the acid reacts strongly, producing a lot of heat quickly. I think the same amount of total heating occurs either way, just the speed is different. I think the first 2 paragraphs of Matt's answer says something similar, but with more science.

@DmitryGrigoryev Also, in general, the heat of reaction depends on the actual molar ratio. There is no a priori reason to expect a symmetry. I find this a satisfactory answer, tough it should be combined with those by permeakra and by Matt.

This is more a problem with thermal mass. Sulfuric acid releases a lot of thermal energy when water is introduced to typical stock sulfuric acid which is about 90% acid by weight. When dilution is performed, a small mass of the acid is combined with a much larger mass of water.
If the acid mass is in the receiving flask, there is initially a small mass of acid and water as the dilution is started. The energy released will heat the low mass mixture, with four joules of energy heating up one gram of the mixture by one degree Celsius. If there isn't much mass, then the temperature change is dramatic and can easily cause the solution to boil, throwing droplets of concentrated sulfuric acid in the air.
Sulfuric acid is extremely dangerous. It is intensely hygroscopic and will readily dehydrate carbohydrates and proteins by removing oxygens and hydrogens to make water molecules. This process will destroy the carbohydrates and proteins, including the materials that make up your skin and muscle. Concentrated sulfuric acid has disfigured many people who have been unlucky enough to come in contact with it. Even when sulfuric acid is diluted the proper way, adding the acid to a volume of water can still result in an extremely hot mixture that might even boil.
Good practice involves using a receiving flask that can contain a splashing acid water mixture, such as an Erlenmeyer flask, and even chilling the mixture during the dilution. A splash shield can also be placed between the worker and the dilution mixture and it could be done in a fume hood. Heavy neoprene lab gloves, as well as eye protection,  are recommended. Students in introductory chemistry classes rarely perform this dilution procedure, because of the high likelihood of accidents.

What makes it actually bad is imo the high viscosity of conc. H2SO4. Water around a drop of H2SO4 immediately has an onset of turbulent convection, distributing the heat. A drop on water in H2SO4 just sits there for half a second until it boils in itself.

A large amount of heat is released when strong acids like sulfuric acid and HCl are mixed with water. First, adding more acid releases more heat, as it is exothermic.
If you add water to acid it forms an extremely concentrated solution of acid initially. So much heat is released that the solution may boil very violently, splashing concentrated acid out of the container and all this because the reaction is exothermic.
On the other hand, adding acid to water, the solution that forms is very dilute and the small amount of heat released is not enough to vaporize it.   So it is always safer to add acid to water, and not water to acid.


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Adding acid to hydrates
What would happen if I added $\ce{Na2CO3 \cdot n(H2O)}$ to a solution of $\ce{HCl}$? Could I consider the reaction the same as just $\ce{Na2CO3 + HCl -> NaCl + CO2 + H2O}$?

There is no difference in properties of salt hydrates and non-hydrates after they are dissolved in water, as they are hydrated to same state. So, yes, you can.
However, there is difference in reaction of salt hydrates and non-hydrates outside water solutions.


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What is the chemical structure of the cocoate ion vs. the olivate ion?
I am wondering how the cocoate ion (in natural soaps, in the form of sodium or potassium cocoate and from the saponification of coconut oil) differs structurally from the olivate ion (from the saponification of olive oil)?
My family is attempting to implement the methods in solveeczema.org and believe that our children have detergent-reactive eczema and our baby is most sensitive.  In moving away from detergents and to natural soaps, we have started using soaps that result from the saponification of vegetable oils - mostly coconut oil, palm kernel oil, and olive oil.  We discovered that our other child (a toddler) seems to react to these soaps in a similar way that he would react to a cleaner containing much stronger/harsher surfactants (such as sodium lauryl sulfate).
Our toddler seems to get bad eczema flares using soaps that contain cocoate, but he has no reaction to soaps that contain 100% olivate.  I have no idea why this would be, since I would assume that both cocoate and olivate have the R-COO- structure.  Is cocoate much more hydrophilic than olivate?  Is the functional group different?  Is the cocoate ion a longer molecule that looks much more like a harsh detergent than a salt resulting from saponification of olive oil or tallow?
Please note that this question is related to two questions already asked:
1)  Is soap the natural form of Sodium Lauryl Sulfate?
2)  https://parenting.stackexchange.com/questions/5047/how-to-test-whether-my-toddler-is-allergic-to-detergent/5049#comment7644_5049

Note 1. Both coconut oil and olive oil are natural oils. The soap, produced from them, is mixture of several compounds.
The difference of olive and coconut oil may be easily found in wikipedia: http://en.wikipedia.org/wiki/Coconut_oil#Composition_and_comparison
In short, coconut oil contains mostly saturated acids esters, while olive contains mostly unsaturated ones. The unsaturated acid esters are responsible for olive oil ability to solidify on air. But here is
Note 2: In my knowledge, allergy is usually caused by carbohydrates or proteins. Fatty acids and their ions have too few binding sites to be recognized by immune system as something distinct from our own. The real source of reaction may be

Minor components of initial oil. There are quite many of them and removing them completely is not easy.
Other components of soap, such as perfume or pigment.
Soap solution pH, that may differs from soap kind to kind, but this is unlikely.

Hi and thanks for your response.  I don't think my toddler is allergic to the soap with cocoate but rather believe his skin is damaged by the cocoate ions in a similar way as SLS and harsher detergents relating to increasing membrane permeability (see http://www.solveeczema.org/thesolution.html#membrane_permeability).  Many other children whose parents have implemented the methods have have had their eczema cured by using soaps (not detergents) even when these soaps contain cocoate but my son seem to have an extra (rare) sensitivity.

The soaps we are using containing cocoate had no other perfumes/pigments that he could be reacting to (one he reacted to contained only saponified coconut oil, palm oil, and olive oil!).  The pH thing is a possibility since some soaps are more drying than others and I wonder if more alkaline soaps cause irritation to the extent that it causes eczema on my toddler (because it seems to cause excessive drying to the point of discomfort, but intact skin) on me.

This is why I wondered whether what the cocoate ion looks like compared with the olivate ion, and if the cocoate ion is better at producing similar results that affect membrane permeability on an eczematic person (similar to other coconut-derived surfactants like SLS) than olivate?

Soap are kind of detergent. Technically. They are little less dangerous for skin, then LSL, but not completely safe.

LSL and soap are so known 'anionic surfactants'. They are less costly then any alternative. They are, however, most aggressive. There are cationic surfactants and non-ionic surfactants as well. Traditional soaps also gives slightly basic solutions, that is not best for skin. This is the reason one should look skeptically on natural soaps, despite their, ergh, 'naturality'.

there is class of non-ionic surfactants/detergents, that used in cosmetics and sometimes washing products (liquid soaps, shampoo etc). They are much less irritating, but have more potential as allergen. You can try cationic surfactants as well. Problem: even if they model naturally occurring compounds, I doubt any of them is completely green, so I do not think you'll find one in green cosmetic markets.

@Julie Leung - Soaps made by saponification of natural oils can have residual sodium hydroxide (which is necessary for the hydrolysis) in them, especially those labeled "lye soap". Sodium hydroxide is caustic, and it could be in low enough concentrations that it only is affecting the more sensitive skin of the baby.


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What are the alternatives to the Periodic Table of the Elements?
What are the different diagrams/tables used to organize the elements other than the Periodic Table? What are the advantages/disadvantages of each?

There is an interesting article with lots of table variants: http://www.chemistry-blog.com/2009/04/26/alternative-periodic-tables/

Hmmm, I've never seen anything but the PT used neither in Highschool nor University, up to my degree, in any text book or class I took. I can't really see any advatages any alternative representations have over the classic PT, or we would see them used in mainstream chemistry.

Check the Wikipedia page on this: https://en.wikipedia.org/wiki/Alternative_periodic_tables

One periodic table that apparently physicists find useful is ADOMAH periodic table, for its usefulness is finding electron configuration of a particular element. More information at this guide. A disadvantage to this table, along with any new table configuration you have not experienced, will be the lack of knowledge in the trends on your new table. There are many spiral looking tables, but the trends such as the change in atomic radii, or electronegative characteristics, might not be as obvious, and you might have to find a guide or relearn old tricks to methods you would learn in your first year of chemistry. At least the ADOMAH Table is built in such a way that these trends should still be recognizable, although clearly not the same as the normal table you find in any high-school science classroom.
Of particular interest to me are dynamic periodic tables. These are tables that are basically computer applications which will change according to what you are looking for. One outstanding example is the widely popular ptable.com which will maintain all the common trends you probably know already, and dynamically show you various states of the elements when changing the temperature. I learned from this table that rhenium has the highest known boiling temperature, while I thought it was tungsten (although tungsten has the highest melting temperature of the transition metals). And better yet, carbon has the highest melting temperature of all elements, very cool table! The disadvantage is obviously that you cannot carry it around in your notebook without some flavor of a mobile computer and/or internet connection.
I am sure there are many other great tables still to be found, but I hope you have learned something new from my post.

I wrote this long answer to a question which was deemed "duplicate" but really asked a more specific question, why there are the gaps in the periodic table etc.

Organisation in Periodic table

The elements are organized in groups (the columns) because the elements in each group have very similar chemical properties.
For example, the noble gases are inert: today, we know that it is because all the electron shells up to a rather big gap are occupied, so there is no free electron that can create bonds with other atoms; and it is not easy for the atom of the noble gas to adopt an additional electron, either.
On the opposite side, the group starting from hydrogen and lithium contains highly reactive elements because they have exactly one electron on top of the full shells above.
If one considers elements with electrons in the shell $1s$ only, there are only two elements, hydrogen and helium (the maximally reactive; or the maximally inert), and nothing in between. There is nothing in between because the shell $1s$ only contains two "slots" for electrons, so the number of electrons may only be $0,1,2$ (zero is not an atom).
However, as we go down the table, towards elements with a higher number of electrons (and protons in the nucleus), there are more possibilities. For example, in the second row, we have 8 elements. Why? It's because the shells $2s$ and $2p$ contain two and six states, respectively – eight in total.
So if we want to have at least one electron in the $n=2$ shell, there are eight options, from $1$ to $8$ electrons. Again, the element with $1$ electron in the $n=2$ shell is analogous to the hydrogen, and the element with the maximum number $8$ electrons, i.e. neon, is analogous to helium because the shells are fully occupied.
But unless the first group which only consisted of hydrogen and helium, there are 6 new elements in between, beryllium, boron, carbon, nitrogen, and fluorine. Their occupation is something in between the minimal and maximal. Those have no counterparts among the simplest elements that only work with the $1s$, $n=1$ shell because there are simply not six extra integers in between one (hydrogen) and two (helium). That's why the gap appears.
When we work with the more complicated shells, we always reveal some new "intermediate" elements that had no counterparts in the simpler periods. That's why one needs to create increasingly wide "gaps" in the upper periods – because the lower periods have some additional "intermediate" elements because there is a greater number of ways to fill more complicated, larger shells.

This doesn't really answer the OP's question about whether there  are alternatives to the standard periodic table. It just explains how the periodic table is laid out.

This is an answer to another question that was labeled a duplicate and closed. It's not my mistake if it was labeled a duplicate even though it is not.

"This is an answer to another question" - Are you serious?! Getting closed means a question shouldn't get answered. Now attempting to answer what community didn't want to do so is wrong.

No, getting closed as a duplicate means that the questions are merged as duplicates.

@LubošMotl Mods can merge dupes only if they both contain good answers. That's rare because we don't want it to happen. Merging doesn't happen to dupes. If it did, some questions in bigger SEs would look like gibberish. After some time, that question is going to get closed.


text:
What is the bandgap value of SiO2?
Using Diffuse Reflectance Spectroscopy, I have obtained the band gap for $\small\ce{SiO2}$ (purchased from a chemical retailer) in the range of 3.25 to 4.75 eV.
I would like to know if the values are in the right ballpark and, if not, would like to review the DRS procedure that I carried out.  I would also like to know if the band gap determined through DRS would have such a wide range.

there are several polymorphic modifications of SiO2, that have different gaps. Most stable alfa-quartz has ~9 ev band gap according to quickly found article http://dx.doi.org/10.1016/j.bbr.2011.03.031 . Which one you mean?

@permeakra The DOI links to a document in "Behavioural Brain Research" ;)

@KlausWarzecha try this http://dx.doi.org/10.1109/16.870550 . Sorry

@permeakra No need to be sorry. It was just a funny coincidence.

For thermally grown amorphous $\ce{SiO2}$ films, a bandgap of 9.3 eV was measured by Weinberg.
Measurements on $\alpha$-quartz by Calabrese give $E_g(\Gamma)$ = 6.3 eV, compared to 9.2 eV calculated by Chelikowsky and Schlüter.
Calculations by Ramos on $\beta$-cristobalite with the space groups $Fd3m$ (face-centred cubic, fcc), $I\bar{4}2c$ (body-centred tetragonal, bt), and $P2_13$ (simple cubic, sc) gave direct bandgaps of $E_g(\Gamma)$ = 5.48, 5.79 and 5.73 eV, respectively.


text:
Why do lipids in a cell hydrolyze into glycerol and fatty acids?
Most Lipids in a cell have ester bonds (made up of glycerol and fatty acids).

I have learned in my chemistry class that we can hydrolyze an ester by simply adding water to it.
$$\ce{CH3COOCH2CH3 + H2O <=>[H^+ (aq)] CH3COOH + EtOH}$$
Then why do lipids which  have an ester bond don't hydrolyze inside a cell which is 70 percent  water? While searching throught web I have found this link which explains of different ways to hydrolyze an ester. In that it explains about how to hydrolyze esters using dilute acids like HCl. In our stomach HCl is present so logically HCl should destroy the ester in stomach lining (which are made up of cells) instantaneonusly, but this is not happening. Why?
$%edit$

Organisms control rates of chemical reactions in two ways. First, by controlling the conditions, and second, by controlling the location of reagents.
Conditions
Hydrolysis of esters is very slow at normal physiological conditions: 37 oC and pH 7.4. A mixture of water and simple esters like ethyl acetate remains stable for years. At higher or lower pH, the reaction becomes faster. As you mention, the HCl in the stomach would definitely increase the rate of hydrolysis. The reaction is also accelerated by enzymes, and organisms use enzymes for all intentional ester hydrolysis.
Location
For a reaction to occur, the water molecule and the acid catalyst needs to get to the ester. In cell membranes, that ester functional group is not on the surface of the membrance; the ionic phosphate part is. The ester group is just below the surface, in the hydrophobic interior of the membrane. Neither water, nor acid can penetrate into this hydrophobic region.
In the stomach, there is an extra layer of protection: mucus, which prevents the stomach acid from reaching the cells.

Stomach wall are made of mucous membrane, that secrets mucus like crazy. Mucous is made of glycoproteins -- sugar+protein complex, that makes diffusion slower, and protein structure make it acid hydrolysis slow. The mucous membrane cells maintains normal physiological pH, and while they are isolated from stomach acid by mucus they secret, they are fine, as hydrolysis is very slow without catalyst.
here is wikipedia article on the question http://en.wikipedia.org/wiki/Gastric_mucosal_barrier
The barrier may be broken by some chemical agents or if mucus secreting cells are suppressed by illness. In latter case stomach damages itself routinely and peptic ulcer decease occurs.


text:
Hierarchy of electronic wavefunctions
The previous question contained too much unnecessary information and was edited.
I am wondering about the "hierarchy" of wavefunctions. If one can combine atomic orbitals (AO) into molecular orbitals (MO) through the LCAO method, could one combine MOs to form a supra wavefunction, or an LCMO?
I've done some research and found a paper on an approach to finding the MO of molecules by effectively taking a LCMO of smaller molecules or constituents. While interesting, this is not what I'm referring to as a "supra" wavefunction, if there's such a thing.

I do not know enough about the particulars of the field to give an answer, but any set of orbitals that can be constructed as linear combinations of molecular orbitals should be valid. In principle, there are more than one way to construct LCAOs. For example, there is a way that resembles the localized valence bond orbitals. In theory LCMOs are still a different set of LCAOs.

I've been reading up on this, and will take a shot at the answer. It is going to be an interesting answer, and probably not what you expect. There has been an interesting conversation going on in the Journal of Chemical Education in response to this letter advocating the retirement of the localized hybrid orbitals because they "aren't real" and cannot be used to predict "all properties" of molecules. Rebuttals included comments that delocalized molecular orbitals also fail to predict "all properties". However, the final word in the discussion came from this letter, which reminds us what orbitals are.
Orbitals are not real. The electrons that populate orbitals are real. Orbitals are mathematical constructs created to describe the curious wave-like probabilistic quantum mechanical properties of electrons. When we talk about atomic orbitals, we generally mean the 1-electron hydrogenic atomic orbitals (HAOs), which are mathematically represented as a wavefunction ($\Psi$) and graphically represented by the probability density function ($\Psi ^ * \Psi$) at either 90% or 95% level. Despite valiant and massive efforts, wavefunctions for multi-electron atoms cannot be solved exactly. They are approximated as a parameterized HAOs. When we talk about the atomic orbitals of carbon, we usually mean the HAOs that closely approximate what we think the AOs of carbon should be.
Molecules and their orbitals are even more troublesome. The LCAO approach creates a bunch of nice 1-electron wavefunctions out of the set of HAOs. These methods give rise to both the canonical delocalized molecules orbitals (CMOs), and the localized valence bond orbitals (VBOs). As the authors of that letter point out:

The term canonical implies that CMOs are not the only possible MOs, and there are many orbitals that generate an identical “chemical reality”.

The approximate polyelectron wavefunction of a molecule (at least in HF theory) is the Slater determinant of a series of 1-electron CMOs $\psi _i$:
$$  \psi _1 \overline{\psi} _1 \psi _2 \overline{\psi} _2 \psi _3 \overline{\psi} _3 \psi _4 \overline{\psi} _4 ... $$
An interesting property of Slater determinants is that unitary operations (like rotations) on the set of CMOs return a different set of MOs, which produce the same expectation values for the polyelectron wave function as original set:
$$ \psi _1 \overline{\psi} _1 \psi _2 \overline{\psi} _2 \psi _3 \overline{\psi} _3 \psi _4 \overline{\psi} _4 ...~~ \underrightarrow{unitary~~transformation} ~~\psi _a \overline{\psi} _a \psi _b \overline{\psi} _b \psi _c \overline{\psi} _c \psi _d \overline{\psi} _d ...$$
Mathematically, it seems to me, that a linear combination of MOs that are themselves linear combinations of AOs would still be a linear combination of AOs, just a different one than you started with. If that is true, then all you have is a different set of MOs than you started with, and according to HF theory, that should give you the same expectation values.
Thus a set of "LCMOs" are no more or less valid than the original set of LCAOs. They likely do not produce any additional insights into the structure of the molecule. However, the approach might be computationally more efficient when trying to determine the MOs of very large molecules and polymers.

I just have to say.. 'yeowch!' regarding the rebuttal to Grushow. They laid it on pretty thick. Thanks for the links.

I wish Richard Bader (of QTAIM fame) were still alive.  He would have said: "See, I told you Schroedinger himself warned about the perils of directly interpretting wave functions." Indeed, one can probably find one (of the infinite number) of unitary transformations of wave functions that whose qualitative MO picture supports virtually any chemical hypothesis.

@BenNorris I agree with everything you have written. Just to make a few comments: MO and VB are equivalent once they are taken to their Configuration Interaction limits: MO-CI = VB-CI. Hiberty, et al. essentially invoke a "delocalized MO" picture to explain why localized orbitals aren't "wrong" -- and in doing so provide a rebuttal against the very purpose for doing localization in the first place.  It simply doesn't matter when one must sum/product over all Molecular Orbitals (VB configurations, etc.).  P.S. Thanks for writing the complex conjugate of psi.

complex conjugate of psi:  reminder that psi may be complex, and so what's the purpose of directly analyzing a complex entitity?  It just isn't real. (Like the electron density.)

As a supplement to Ben Norris's answer, I thought I'd add the following:
When you use a basis to construct an orbital, your original basis functions can already be represented as linear combinations of your whole set:
$\chi(x,y,z,s)=\chi(\mathbf{x})= c_1\varphi_1 + c_2\varphi_2+c_3\varphi_3\dots \\
\varphi_1 = 1\varphi_1 + 0\varphi_2 + 0 \varphi_3\dots$
(where $\chi$ is an MO and $\varphi$ is a basis function)
So, altering the coefficients or linearly combining MOs doesn't increase the complexity of your result:
$\chi_a + \chi_b = (c_{a,1} + c_{b,1})\varphi_1 + (c_{a,2}+c_{b,2})\varphi_2 + (c_{a,3}+c_{b,3})\varphi_3\dots$
When you construct a wave function from basis functions, the usual method is to use a Slater determinant - the determinant of a matrix as follows:
$\Phi(\mathbf{x}_1,\mathbf{x}_2,\dots\mathbf{x}_N) = \frac{1}{\sqrt{N!}}\left|
   \begin{matrix} \chi_1(\mathbf{x}_1) & \chi_2(\mathbf{x}_1) & \cdots & \chi_N(\mathbf{x}_1) \\
                      \chi_1(\mathbf{x}_2) & \chi_2(\mathbf{x}_2) & \cdots & \chi_N(\mathbf{x}_2) \\
                      \vdots & \vdots & \ddots & \vdots \\
                      \chi_1(\mathbf{x}_N) & \chi_2(\mathbf{x}_N) & \cdots & \chi_N(\mathbf{x}_N)
   \end{matrix} \right|$
(from the Wikipedia article)
This gives you an expression involving products of your MOs rather than just linear sums, increasing the complexity. This is kind of awkward looking for the n-function case, so instead, for a 3 MO case:
$\Phi(\mathbf{x}_1,\mathbf{x}_2,\mathbf{x}_3) = \frac{1}{\sqrt{3!}} ( -\chi_3(\mathbf{x}_1)\chi_{2}(\mathbf{x}_2)\chi_1(\mathbf{x}_3) + \chi_2(\mathbf{x}_1)\chi_{3}(\mathbf{x}_2)\chi_3(\mathbf{x}_1) + \chi_3(\mathbf{x}_1)\chi_{1}(\mathbf{x}_2)\chi_2(\mathbf{x}_3) + \chi_1(\mathbf{x}_1)\chi_{1}(\mathbf{x}_2)\chi_3(\mathbf{x}_3) + \chi_2(\mathbf{x}_1)\chi_{1}(\mathbf{x}_2)\chi_3(\mathbf{x}_3) ) $
Then, just as you can use AOs to build MOs, you can either use that alone as your wave function, or you can use linear combinations of them to obtain a multiconfigurational wave function - the MCSCF methods mentioned in jjj's answer:
$\Psi = c_1\Phi_1 + c_2\Phi_2 + c_3\Phi_3 \dots$
I'm not sure if this is the sort of thing you were looking for.

I am writing this as a comment, since all the replies are fine, but I cannot write comments. I hope though that the answer is not fully redundant.
Quantum Mechanics is linear. In the space of a particle you can write the wave function as a linear combination of whatever functions. So you can write MOs (molecular orbitals) as LC (linear combination) of AOs (atomic orbitals) or AOs as LC of MOs. The distinction between these set of basis is physical, not mathematical.
One should realise that the electronic wave function of a molecule is not an orbital: it "dwells" in a 3N space (not including spin) where N is the number of electrons, and it must by antisymmetrized. Again, since Quantum Mechanics is linear, the LC of wave functions of 3N electrons is an electronic wave function of 3N electrons. You can always approximate one in terms of others that are simpler and indeed this is the basis of all methods beyond Hartree-Fock.
Therefore, the LC of MOs is another MO of a single electron in the molecule, not the wave function of a "molecule" or "super molecule". The LC process does not move you away from the initial space where the function is defined.

(atomic orbitals -> molecular orbitals) -> ("Slater determinats" -> MCSCF wave functions) -> mixed state N-electron density matrices -> grand canonical ensemble (Fock space) density matrices
the first two are 1-electron objects, the next two are N-electron objects representing pure states at different level of approximations


text:
Reaction of soap with water
I've seen it many times... Why does soap become foam-like when it reacts with water (moisture) or when we touch it with wet hands.. What compound actually causes this physical change? And why does this physical change occur?

So you're thinking of solid 'bar' soap?

yeah,..! Why are there any other different species in your mind.>?

Sorry, what I meant is that I guess you're in particular thinking about the physical change of the properties of solid soap when you get water on it.

Yeah man.. I mentioned the word 'physical change' in the question

The simplest kinds of soap are just a fatty acid chain with something more ionic on the end - something like the ubiquitous sodium laureth sulphate:

When it's in bar form, it behaves a bit like a solid block of plastic, because chemically speaking, those long chains of carbon are quite a lot like a common plastic such as polythene. These chains attract each other and repel and exclude water, keeping the block together. The sulphate group on the end is hydrophilic, though - it's chemically stable surrounded by water. This means that when you add water, it infiltrates and teases apart the structure of strands, getting to the ionic ends on each chain and forming a "solvation shell" of waters around each one. There's no real chemical reaction here - as you say, it's a physical process, which is why rubbing the soap helps - you're applying shear to the structure and helping to separate the strands.
Once they're separated from the bulk, the soap forms tiny bubbles called "micelles", which have the chains pointed inwards towards each other and the ionic sulphate groups meeting the water at the outside.

While this question should get a more elaborate answer, mine will at least give you leads.
Soaps act as surfactants, that is, they are amphiphilic (both hydrophobic and hydrophilic) molecules that lower the surface tension of a liquid.
By lowering the surface tension of water, soap bubbles can form, the surface becoming more "elastic" - the molecules at the interface experience less inward/normal force than they would were surfactant not present. On the contrary, adding soap to a pond would prevent water striders from gliding on the surface of the pond.

Thanks. But I didn't ask about surface tension. Its in my 10th grade syllabus. I'm asking 'why is this physical change?' Please read the question and its not necessary for ponds & surface tension

@CrazyBuddy Then your question is ambiguous. The question you ask in the main post and the one you ask in the comment are different. Consider clearing up your question.

One thing that should be made clear is the question you want to ask, and also what you mean by "solid-like".

sorry for that.. I told solid-like 'cause i thought soap as something like colloid due to its appearance

The soap dissolves in the water. Dissolution is not a chemical change because the same chemicals are present, they just interact with each other differently. Like CHM said, soaps act as surfactants and the difference in surface tension improves the survivability of bubbles, thus creating a "foam".

There. Is. No. Chemical. Change. We have all been trying to tell you this. We've been telling you about physical changes and you've been ignoring us.

We're talking about physical changes and you keep going "no I mean the physical change man".


text:
What is burst nucleation?
Can you tell me what is a burst nucleation process, and how can it be used for synthesising uniform nanoparticles?
Can you also provide (if possible) any links to papers or materials regarding mechanism of nanoparticles growth?

This is a part answer - an annotated short list of references regarding burst nucleation and mechanisms of nanoparticle growth.
Model of nanocrystal formation in solution by burst nucleation and diffusional growth (Robb and Privman, 2007), in this article the authors

present a model with the assumptions of instantaneous thermalization below the critical nucleus size and irreversible diffusive growth above the critical size

A main feature of this article is that

for the first time formulates LaMer's explanation of burst nucleation in a manner allowing quantitative calculations.

This is followed up by one of the authors the next year with the article Diffusional Nucleation of Nanocrystals and Their Self-Assembly into Uniform Colloids (Privman, 2008), where he

review theoretical explanation of mechanisms of control of uniformity in growth of nanosize particles and colloids.

An explanatory diagram of LaMer's explanation (referred to in the 1st article) is shown below:

The image is from the article Controlled synthesis of colloidal silver nanoparticles in organic solutions: empirical rules for nucleation engineering (Sun, 2012) - unfortunately, the article is behind a paywall.
However, a sample chapter explaining the process in full entitled Growth of Nanocrystals in Solution (Viswanatha and Sarma), provides a greater detail about the processes involved with nanoparticle growth.


text:
How can carbon dioxide be converted into carbon and oxygen?
How can $\ce{CO2}$ be converted into carbon and oxygen?
$$\ce{CO2 -> C + O2}$$
Alternatively:
$$\ce{CO2 + ? -> C + O2}$$
I'm aware that plants are capable of transforming $\ce{CO2 + H2O}$ to glucose and oxygen via photosynthesis, but I'm interested in chemical or physical means rather than biological.

Plants do not exaclty do this, they do $2n CO_2 + 2n H_2O + photons \rightarrow 2(CH_2O)_n + 2n O_2$ according to http://en.wikipedia.org/wiki/Photosynthesis In quite a complicated reaction schema.

You can use scrubbing process. Using $\ce{KNO3}$

In my opinion, the catalytic, solar-driven conversion of carbon dioxide to methanol, formic acid, etc. is much more interesting and promising, but since Enrico asked for the conversion of carbon dioxide to carbon itself:
The group around Yutaka Tamaura was/is active in this field. In one of their earlier publications,[1] they heated magnetite ($\ce{Fe3O4}$) at 290 °C for 4 hours in a stream of hydrogen to yield a material which turned out to be stable at room temperature under nitrogen. This material, $\ce{Fe_{3+\delta}O4}$ $(\delta=0.127)$, i.e. the metastable cation-excess magnetite is able to incorporate oxygen in the form of $\ce{O^2-}$.
Under a $\ce{CO2}$ atmosphere, the oxygen-deficient material converted to "ordinary" $\ce{Fe3O4}$ with carbon deposited on the surface.
This remarkable reaction however is not catalytic, but a short recherche showed that the authors have published a tad more in this field. Maybe somebody else finds a a report on a catalytic conversion among their publications.

Tamaura, Y.; Tahata, M. Complete reduction of carbon dioxide to carbon using cation-excess magnetite. Nature 1990, 346 (6281), 255–256. DOI: 10.1038/346255a0.

Aren't they using something like this in the International Space Station to provide astronauts with fresh oxygen?

@EnricoPallazzo, evidently, the ISS life support systems produce oxygen primarily by electrolysis of water. It can also be chemically generated by the actions of strong oxidizers (chlorates, perchlorates, peroxides, superoxides, etc.), though whether those technologies are used in the ISS, I don't know.

@GregE. generating oxygen is a separate concern from CO2 removal (i.e. wouter dorgelo's  'providing astronauts with fresh oxygen' ... stated in the context of CO2 conversion to C and O2). Yes, you can generate O2 via electrolysis, but it is the concentration of CO2 and its build-up that creates a hazard for humans. The EPA lays out concentration levels and their effects on humans. Problems start at 1000 ppm (mental fog) and progress from there. I believe Wouter's intended question was "isn't this how the CO2 scrubbers in the ISS work?"

Electrolysis of carbonates in anhydrous environment can produce either $\ce{CO + O_2}$ or $\ce{C + O_2}$, [ref] depending on conditions. A base, remaining in electrolyzed liquid then can capture carbon dioxide from other sources and be recirculated.
Reference: L. Massot, P. Chamelot, F. Bouyer, P. Taxil; Electrodeposition of carbon films from molten alkaline fluoride media. Electrochimica Acta, 2002, 47 (12), 1949-1957. https://doi.org/10.1016/S0013-4686(02)00047-6

Since the carbon atom is linked to two oxygen atoms via double bonds, usually more amount of energy must be supplied in order to separate it. About 94 kcal of energy is required per mol of $\ce{CO2}$ (about 44 g). This energy input could come from any source, but the major source of conversion is through photosynthesis using solar energy which is very well known by the famous equation:
$\ce{6CO2 + 6H2O -> C6H12O6 + 6O2}$
But in fact, there is a machine built by Sandia researchers known as Counter-Rotating-Ring Receiver Reactor Recuperator (CR5), which uses solar power to convert carbon dioxide and water to carbon monoxide, water, hydrogen and oxygen at a temperature of about 1500 °C using a solar concentrator. Iron oxide acts as an extractor of oxygen from $\ce{CO2}$ forming $\ce{CO}$. But, the main aim of this method is to produce fuel and not carbon. But it’ll take at least 15–20 years to come into usage because, only the prototype of this machine has been invented and tested.

Use Exploit the fact that burning magnesium continues to burn in atmosphere of carbon dioxide.
$${\ce{Mg}}+\color{\red}{\ce{CO2}}\to \ce{MgO}+\color{\red}{\ce{C}} $$
Now you can electrolyse your mixture :
$$\ce{MgO}+\ce{H2O}\to \ce{Mg(OH)2} $$
$$\ce{4OH-}\to\color{\red}{\ce{ O2}}+\ce{2H2O}+\ce{4e-}$$

How can we start burning magnesium in CO2?

React carbon-dioxide with hydrogen in the Sabatier process to get methane. 400 °C, high pressure and Nickel catalyst needed. The process is slighly exothermic so it can keep going on its own:
$$\ce{CO2} + \ce{4H2} \rightarrow \ce{CH4} + 2\ce{H2O}$$
This process have been proposed to generate fuel on Mars, and used on the ISS to process exhaled carbon-dioxide.

Split the resulting water (use electrolysis or some other thermochemical cycle) take the oxygen, bring the hydrogen back to Step 1:
$$2\ce{H2O} \rightarrow 2\ce{H2} + \ce{O2}$$
Electricity for this may come off solar panels.

Methane decomposes at high temperatures. The process goes to completion around 1200°C. Collect the condensed carbon, bring the hydrogen back to step 1.
$$\ce{CH4} \rightarrow \ce{C} + 2\ce{H2}$$
This process is proposed as an emissions free alternative to produce hydrogen from natural gas. Heat may come from concentrated solar light.

$\ce{CO2}$ into $\ce{CO + O2}$ is easier but you need a catalyst and about 1500C
http://www.rsc.org/chemistryworld/news/2008/january/03010801.asp
In short:
Fe3O2 is heated to 1500C driving off oxygen. The resulting FeO is moved to CO2 chamber where it absorbs oxygen from the CO2. Result is CO and cooled Fe3O2. Cycle repeats.

Link only answers are discouraged, could you give an explanation of the method listed there (a summary is OK if it's too long)

Fe3O2 is heated to 1500C driving off oxygen.  The resulting FeO is moved to CO2 chamber where it absorbs oxygen from the CO2. Result is CO and cooled Fe3O2.  Cycle repeats.

The electronic configuration of carbon is $\ce{1s^2 2s^2 2p^2}$. So a carbon atom requires 4 electrons to complete its outermost orbital. The electronic configuration of  oxygen is $\ce{1s^2 2s^2 2p^4}$ so an oxygen atom requires 2 electrons to complete its octet configuration.
Thus two oxygen atoms and 1 carbon atom form a double bond between them so more energy is required to break this bond. In this way $\ce{CO2}$ is formed.
For converting $\ce{CO2}$ back into carbon and oxygen atoms, $\ce{CO2}$ should be heated at almost $\pu{298 K}$ so carbon becomes gaseous and oxygen becomes part of air.

Edit from another deleted answer:
If $\ce{CO2}$ can be produced by combining $\ce{C}$ and $\ce{O2}$ under the presence of air; then oppositely $\ce{CO2}$ can be separated to $\ce{C}$ and $\ce{O2}$ by heating $\ce{CO2}$ at almost that temperature which is impossible for $\ce{CO2}$ to absorb heat.

$\ce{CO2}$ probably doesn't thermally dissociate into $\ce{C + O2}$ until the temperature is at least several thousand kelvin. It takes around 1000 K to perform $\ce{CO2 -> CO + 1/2 O2}$, and removal of the last oxygen atom from $\ce{CO}$ is an extremely difficult process, as carbon monoxide contains the strongest covalent bond known with a bond energy of $1072\ kJ/mol$.

I would bet that diatomic oxygen dissociates at much lower temperatures than CO, in which case it will be impossible to use thermal methods to get C + O2.

CO2 should be heated at room temperature so carbon becomes gaseous and oxygen becomes part of air. What the hell are you talking about? That is obvious nonsense and clearly doesn't happen. Wha the hell did you mean?

298K is room temperature, in case anyone didn't spot that.