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Why are peroxides explosive?
Why do peroxides tend to explode with little provocation? In general, what are typical products of their explosion?

In general, molecules possessing two highly electronegative atoms that are directly bonded (in your case, the two adjacent oxygens) are rather unruly substances; the weak -O-O- bond is easily cleaved.

Thanks! Would the bond tend to split homo- or heterolytically?

OH radicals are fairly stable, and I can't see why charges should be separated, so a homolytical cleavage at the O-O bond sounds reasonable.

I don't like to speak in generalities if I can help it.
But we can talk, for instance, about hydrogen peroxide. One of the most familiar schemes is a disproportionation where,
$$2 \ce{H2O2} \rightarrow 2\ce{H2O} + \ce{O2}$$
and as you study this particular species, you will see this reaction more and more often. Actually in solutions of hydrogen peroxide you can take a 30% concentrated solution and add some Mn and subsequently measure a significant increase in oxygen concentration over time. Although you should know the solution itself decomposes slowly in practice when exposed to light or due to trace impurities, but this pales in comparison to the case involving a catalyst. This reaction is, in part, why researchers believe a peroxide-type species is important for the Oxygen Evolving Center of Photosystem II. Another fun demonstration of this process is often where foam is generated from a large graduation cylinder, see here, upon addition of KI.
Now 30% is weak, if instead you look at 80% solutions you will notice that the decomposition will evolve steam because it is so exothermic. In fact, the first liquid-fueled rocket in 1939 used an 80% concentrated solution and decomposed it with calcium permanganate. Later engines utilized another important property of hydrogen peroxide for its decomposition until around 1980: the fact that it is a stronger oxidizer than chlorine/chloride.
Thermodynamically we say the disproportionation is favorable since it has a fairly negative (~ -100 kJ/mol) enthalpy change and a considerable change in entropy ( ~ + 70 J/(mol K)). The basis of this is that it is highly favorable to form dioxygen and this is a large driving force of this reaction in terms of enthalpy and entropy. You'll also see this danger in species that are really close to forming dinitrogen too, see here and here. It only takes a catalyst for this down-hill process to take off. Which could be quite deadly if unintentionally concentrated and then introduced to a catalyst in a closed system. Perhaps the story is similar for other peroxides.


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Why is mercury a "fancy proton"?
I remember that one of my old organic chemistry profs referred to mercury as being a "fancy proton". I think it was in the context of oxymercuration reduction.
What might have been his rationale for offering this comparison?

It is a large, soft, Lewis acid. One can use protic acid and water to promote the same reaction, but the product distribution can sometimes be different.
 In the oxymercuration reaction Hg can form a pi complex with alkenes or alkynes, and these "bridged" cations will not undergo rearrangments to the extent that those produced by protic acids might.
 Bottom line - they're both essentially performing the same function in this reaction - activating the pi bond towards attack - but the oxymercuration is milder.

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What is the correct name for this lab technique in crystallization?
Take a glass rod and rub vigorously the wall of the flask, the substance will crystallize out of the solution.
Take a fire polished stirring rod and etch (scratch) the glass of your beaker. The small pieces of glass that are etched off of the beaker serve as nuclei for crystal formation
But reaching back to my halcyon years of organic synthesis and crystallization, recalling vivid memories of scratching the sides of Erlenmeyer flasks with a glass rod to induce crystallization
It's a common lab technique. But does it have a name?

I don't remember it having a name, but the idea is that the bits of glass from the scratching will serve as "seeds" (nuclei) from which crystals can form, i.e., heterogeneous nucleation.

Could you be a bit more verbose in your question? It makes it easier to understand.. Provide as many details as you know (do a bit of research yourself to get more details, that makes it a really good question).

OK, I have extensively googled it but I can only find descriptions, not the name. I'll add more descriptions

I've always just heard this technique described. The purpose is to provide nucleation sites which then seed the formation of crystals.

@Joseph - regarding the notion that scratching abrades glass particles off the beaker which act as nucleation sites, I'm not convinced. My suggestion would be that the points of contact between the stirring rod and the beaker are nucleation sites.

@Richard: I suppose it depends on what the rod and the container are made of. Borosilicate glass has Mohs hardness ~ 6, and flint glass in stirring rods has Mohs hardness ~ 7, so it is likely that you can scratch off bits of Pyrex from a container if your rod is hard enough. (Rods might also be made from borosilicate, in which case both the rod and the container might contribute glass bits.)

@J.M. - Regarding the Moh's hardness difference - interesting. Shouldn't borosilicate glassware acquire a 'memory' of where it has been scratched (due to the ablation of material from that point) which would obviate future scratching, however?

@Richard Terret, You are absolutely right, and btw. the need of a fire polished glass rod is nonsense. A old and scratched one is sufficient. I see the nucleation from the (ultra)sound or cavitation happening at the rubbed surfaces. Think of the "starters" used in those hand warmer bags with sodium acetate. A name for the lab technik is "rub". :=)

@Richard: I've seen flasks acquiring streaks on the bottom from their continued use in recrystallizations over the years by students in lab classes. It doesn't seem like the scratches are initially apparent (unless one rubs really hard), but the effects of the scratching will show up over time.

@Georg (and Richard)- I agree. I'm not so sure that it is the "microscopic flecks of glass" so much as the sonic disruption of a supersaturated solution. I have accomplished something similar by flicking the flask hard with my fingernail. BTW, I've always called (and heard called) this technique "scratching".

@Ben right, "scratching"! My English was so-so in grammar school, and did not improve much later :=( What do You think about "itch"? :=)

I think the most specific you can get on this phenomenon is nucleation , specifically heterogenous nucleation / surface nucleation.
The rough surface formed creates scope for nucleation to occur, and  this creates the crystals.
I don't think there is a specific name for this technique as a lab tenchique .

In my experience, my senior lecture has called this method 'scratching' and said that this is a widely used term.
The term nucleation is used to aid the description of the technique 'seeding' where a couple of grains of the pure substance you are creating is added to act as a nuclei for crystallisation to take place upon.

I believe the term you are looking for is trituration.  The word itself has other meanings, but this method of inducing crystallisation is one.

No trituration exploits relative solubilities for purification. It isn't about forming the initial crystals from solution.

If you're correct, why would one be attempting to exploit these solubilities in the first place? The terminology applies to both adding solvent to dissolve impurities vs product (or vice versa ) as well as using a solvent mix or volatile solvent to dissolve and subsequently attempt crystallisation

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Does ozone (O₃) gas have a color?
Does ozone have a color or is it colorless? I have heard that ozone has a bluish color.
Could someone explain this observation scientifically?

This is pretty Googleable, and I'm closing it as such. If anyone disagrees, please let me know (with a reason)

Well, of course this is googlable, but why closing for "too localized"?

I googled this. I got two answers. I just say it in question. Bluish and colorless. Also a specific explanation not getting. Also I think a question closing because of its Googleable not a good practise and discouraging.Thanks.. :(

@Manishearth see this meta question.. http://meta.stackexchange.com/questions/8724/how-should-we-deal-with-google-questions

@georg best fit. Unfortunately, the "general reference" option is only enabled as a trial on some sites. Still, current consensus seems to be that such questions ought to be discouraged.

Hmm, that meta post is from the earlier days of SO, when rules were much more lax. Issue is, we want good, conceptual questions, not stuff that can be easily Googled. Such stuff is better asked on [chat]. in this case, it did get a conceptual answer, though the question itself still remains not that good. And don't let closes discourage you--if you can improve it, it may be reopened. If you still feel that this shouldn't have been closed, feel free to post on [meta], we can see what the community thinks. (as it is I was unsure about closing this, but another mod agreed)

You may not realize, but Google has picked up this question and answer as the featured link at the top of the page for "colour of ozone".  The advice to google it leads back here. What's more, another chemistry stack exchange answer similarly closed also is in the first page. Also a Wikipedia article on Ozone which says almost nothing on the topic.
Luckily Quora comes second. I've added an answer there for those wondering where they can find the colour of ozone https://www.quora.com/What-is-the-colour-of-ozone/answer/Robert-Walker-5. Others here might like to add answers there as well.

Ozone gas does indeed have a slight blue color. It is dark blue in a liquid form and purple-black in solid form.
The color comes from the fact that it absorbs all wavelengths of light other than those in the blue region. As the phase changes, the molecules get closer together, shifting the area of absorption, changing the color of the material.


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Which cyclohexane conformation is more stable?
I got this question wrong on a test and I don't know why:

Question: Which conformation is more stable?

Answer: the chair on the left is more stable.

I thought that the ring conformation has zero ring strain so I'm confused.

The conformation on the left is the one where all the bond angles can assume the natural tetrahedral angle, so...

Last not least the plane thing is not a conformer! This plane thing would "sit" on top of a local maximum of the conformational hypersurface, thus it does not "exist".

Yeah, the other main conformation of cyclohexane is boat, not flat. I recommend getting a molecular modelling kit and building them; That was how my orgo prof taught us, and it really lets you 'see' the bond angels and why it exists as a chair and boat.

By symmetry, the cyclohexane on the right requires each carbon atom to have a $\ce{C-C-C}$ angle of 120 degrees. From VSEPR theory, this is not optimal for tetrahedral carbon. A regular tetrahedron has angles between vertices of approximately 109.5 degrees and the equilibrium bond angles of a secondary carbon won't be too discrepant from this.

Pointing out the stable conformations of cyclohexane was a major part of Barton's Nobel Prize. His original review is here (paywalled, but you can see the first page).
The topic is discussed by Henry Rzepa on his blog starting with a little history:

Like benzene, its fully saturated version cyclohexane represents an icon of organic chemistry. By 1890, the structure of planar benzene was pretty much understood, but organic chemistry was still struggling somewhat to fully embrace three rather than two dimensions. A grand-old-man of organic chemistry at the time, Adolf von Baeyer, believed that cyclohexane too was flat, and what he said went. So when a young upstart named Hermann Sachse suggested it was not flat, and furthermore could exist in two forms, which we now call chair and boat, no-one believed him. His was a trigonometric proof, deriving from the tetrahedral angle of 109.47 at carbon, and producing what he termed strainless rings.
  henry rzepa blog

Some related calculations are available here which show the relative energy levels of the different conformers.
Some of these results are easy to visualise with simple molecular modelling kit, but you need space-filling models to really appreciate some of the subtle detail (eg why a twist-boat conformation is more stable than a boat).
It is also worth noting that, at least at room temperature, the conformers will interconvert rapidly as there is enough thermal energy to overcome the barriers between them.
And, there is even a Youtube visualisation showing some of this.

If you look at a Newman projection of "flat" cyclohexane, you'll see that along each C-C bond, the hydrogens and carbons are eclipsed. This leads to considerable torsional strain (imagine twisting a spring - it naturally wants to spring back) favoring relief of these unfavorable interactions. By contrast, in the chair form of cyclohexane, the Newman reveals that all of the hydrogens and carbons are staggered as you look along each C-C bond, which is a considerably more stable conformation.
The difference between the two forms is about 102 kJ/mol. For more references see here http://www.ch.ic.ac.uk/local/organic/tutorial/cyclohexane/index.html

Bond dihedral effects are something I overlooked and on reflection (with reference to the energies of eclipsed/gauche conformations of the central bond of n-butane) 102 kJ/mol sounds about right for 6 bonds. How much water does my answer hold, so to speak?

I don't have a precise answer for you, but the ring strain effect (120 degrees) would be minimal when compared with torsional strain effects.


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How were elements and molecular formulæ determined in the old days?
Before modern chemical techniques, say in the 19th century and earlier, how did (al)chemists determine whether a substance was an element or a compound? If it was a compound, how did they determine its molecular formula and structure?

"...how did (al)chemists determine whether a substance was an element or a compound?" - It wasn't always foolproof. Chlorine was once thought to be a compound, and lime (calcium oxide) was once thought to be an element.

This should be a community wiki, because there is no single correct answer. It won't let me do it because I'm new to the chemistry stack exchange, but anyone else is welcome to.

First thing, they couldn't tell elements and compounds apart. See Dalton's list of elements with their symbols.

Lime is right there--and it's not actually an element. Neither is Baryte(s). Soda and Potash probably refer to the oxides of Sodium and potassium, though terminology may have been different back then.

In those days, elements were things that could not be synthesized by addition reactions. If you have a reaction $\ce{A + B ->C}$, then $\ce{C}$ is a compound (can't say much about the other two). Alternatively, if the substance can be decomposed (via anaerobic heating/etc), then it's a compound.
Most probably they would stuff things into the list "innocent until proven guilty" (I don't know about this).

Molecular mass can be determined by osmotic pressure, ebullioscopy, Cryoscopy, etc. (Usually stuff delendant on colligative properties)
Determination of molecular formula is much harder.
For organic compounds, it can be done by burning in a controlled environment. The ratios of oxides produced lets us know the relative quantities of all elements but oxygen. The quantity of oxygen can be determined by finding the decrease in Oxygen content of the chamber, and subtracting it from the Oxygen content of the oxides. The molecular formula can be found if you use the emperical formula and determine molecular mass.
For inorganic compounds, they probably had various tests for ions/etc (like the silver nitrate test for halides). They also had flame tests (and the borax bead test), which could help identify certain ions. I'm not too sure here, though.

I once read through a book of chemical identification techniques from the 50s, it was quite interesting. Yes, there were a lot of precipitate tests, and a lot of branching series of experiments to nail down the exact identity.

@Canageek: some schools still use precipitation methods to test the skills of their undergraduates.

I've done a couple of those, but this was mostly spot tests you could do on less them a miligram of compound, all really easy to perform as well.

One very useful characterization tool would've been the sodium fusion test.

It is a general test for the detection of halogens, nitrogen and
  sulfur in an organic compound. These elements are covalently bonded
  to the organic compounds. In order to detect them, these have to be
  converted into their ionic forms. This is done by fusing the organic
  compound with sodium metal. The ionic compounds formed during the
  fusion are extracted in aqueous solution and can be detected by simple
  chemical tests. The extract is called sodium fusion extract or
  Lassaigne's extract.
When an organic compound is heated strongly with sodium, any halogens,
  nitrogen, and sulfur will be converted into inorganic sodium salts
  such as sodium halide (for halides), sodium cyanide (for nitrogen),
  sodium sulfide (for sulfur), and sodium thiocyanate (for sulfur and
  nitrogen).The nitrogen is confirmed with ferrous sulfate i.e. iron
  sulfate.

An example procedure for halogens: Melt a small pellet of sodium (first strip the oil off with hexanes) in a test tube using a Bunsen burner. When it gets red hot, remove from heat, add a small amount of sample (1 to 2 drops if liquid), allow it to react for a bit and then place back on flame for a few seconds. When adding the sample, I would make sure that the test tube is contained in another piece of glassware, should it break. After allowing the sample to cool, quench it with a few mL of water and then neutralize the mixture using dilute nitric acid. If there is sulfur or nitrogen present you should boil this solution for a minute. Now, add silver nitrate and you should get out a nice precipitate that varies slightly in colour depending on the halogen in question. There are milder procedures but this is the only one I've had experience with.


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Grease & Thermal resistant cardboard for the foods industry?
I am on the lookout for, preferably, a cardboard (ish), that is both heat & grease resistant, for usage in microwave popcorn. As such, the material must be appropriate for direct contact with food.
The idea is that a tray is placed on the bottom of a box, that will shield the body from any grease that will go liquid when heated (microwaved). Obviously, the most widely available material, and the cheaper the better is the quest that I have undertaken, and I hope that you folks can point me in the right direction, as I can't quite afford professional help for this project, as it stands.

Hi Alamari, welcome to the site.  Can you give us a better sense of what type of temperature this material would need to withstand?  It may help for you to reframe this in terms of actual chemical properties rather than looking for a specific product, as "shopping recommendations" are technically off-topic on SE, and I think this would fall under that designation.

I don't want to answer your question, because it seems like it might be off-topic (and too narrow) as written. I will give you some advice, instead. Parchment paper has most of the properties that you want. It is paper embedded with silicone to provide heat resistance, a non-stick surface, and grease resistance.


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A quick method for separating liquids?
I'm looking for a way to take 3 separate liquids (liquids which function as paint -- they will all be different colors), put them together into a container, and then rapidly separate them back to distinct colors.
I can handle all the electrical engineering aspects of my project, but this portion is beyond me. Does anyone have any ideas?
When I say "a quick method", I mean something that can be done easily without an excess of time or equipment. Ideally an way to make the paints intrinsically immiscible, but I need to be able to separate them entirely, so that when they are poured together into a container, I can then separate and reuse them.

Column chromatography, but you will need the appropriate glass tube, powdered silicate and an abundance of the appropriate solvent. This depends highly on what your paints actually are, but this is my "quick" response.

This question should be returned to physics, better theoretical physics, due to the questions ethereal nature.

@Chris Paints regularly contain pigments! So any kind of chromatography will fail.

@Georg, any chromatography will fail on any pigment? That's incorrect, the classical chromatography experiment is to separate leaf pigments.

Really think that the question needs expansion. Please explain what kind of experiment this is or what kind of effect you want to observe.

I wish to place three colored substances (which act as paint, regardless of whether they be liquid, gel, spray, etc) onto a surface, freely overlapping, then almost immediately remove them, leaving not a spot behind, and separate them back into separate entities. I do realize this will take some professional-grade chemical engineering, but I just wanted to verify that it was even possible.

You need to explain your "colored substance" better.  If they are not a single component, but a mixture of components that make up each paint (colored substance), then I think this is a very difficult if not unsolvable problem.  If these "colored substance" are single, individual components, then you have some hope.  Please provide more thorough detail of this problem.

The content of the substances is irrelevant for other purposes; they can consist of whatever they need to in order to have vibrant color and meet the above criteria.

A stable 3-phase liquid system can be created, as Spießbürger mentioned with three mutually immiscible liquids. The best one that I know of is the organic-aqueous-fluorous 3-phase system.
Fluorocarbons (hydrocarbons in which every hydrogen atom is replaced by a fluorine atom), are interestingly immiscible in both water and common organic solvents. This property has added a dimension to chemical separations using fluorous phase extractions.
If you have three paints, one oil-based, one water-based, and one fluorous-based, then you could separate them easily using a separatory funnel or similar device. You would have to make sure that the pigments in each paint was only significantly soluble in that solvent, which would take experimentation.

You could use three immiscible liquids with different densities. Gravity will separate them for you. Here is a typical example of one of these experiments using water, cooking oil and corn syrup:
http://sabrinaschool.wordpress.com/2008/02/08/liquid-density/
For faster separation you could build a centrifuge.

You have three miscible colored liquids you want to separate quick, cheap, and easy at will? There must be two orthogonal properties.
1) Couple one chromophore to a superparamagnetic particle, e.g., a ferrofluid.  It is pulled out in a strong and strongly divergent magnetic field'
2) Couple the second chromophore to a strong acid, e.g, as a sulfonic acid.  It stays in water, certainly to pH 7 and more basic, resistant to immiscible solvent extraction.
3) Couple the third chromophore to a strong base, like guanidinium.  It stays in water until you go strongly alkaline, and then then be extracted into an immiscible solvent.
If you don't mind spending a year or ten in the lab, case (1) plus a temperature-driven critical phase change will do a pretty, polystyrene- and poly(vinyl methylene)-linked chromophores, http://www.doitpoms.ac.uk/tlplib/solid-solutions/printall.php (bottom),

Better still, hook one chromophore to poly(N-isopropylacrylamide),  http://en.wikipedia.org/wiki/Poly(N-isopropylacrylamide).  One comes out with magnet, once comes out with temperature, the third remains.

The only thing that comes to my mind is laminar flow.
Watch this video and tell me if that's something like what you're look for.
Problem is, you're not really separating "three liquids"...

i dont think their is any easy method(in a single step) for doing this kind of separation. if even one of the three liquids are immisible(not soluble) to others then you can try simple separation  using funnel (most paints are emulsions )and also try to learn about different kinds of separation techniques. if possible find out the solubility of those paints(which are polar and nonpolar) and resubmit here.

Sorry to say no experience but I would think a polymer base cheap achrylic mixed with a 3/4  Mercury metal heated to bond, a symple oil base paint and liquified water color would do the trick in a more than precise time


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Why does oxygen not condense onto open dewers of liquid nitrogen?
I had a discussion with my supervisor today after I was a touch careless pumping down a flask with a liquid nitrogen cold trap attached. I checked, and there are reports of explosions after running air through a liquid nitrogen trap.
Given this, why does liquid oxygen not condense when you leave an open dewer of liquid nitrogen out? I mean, the air touching the surface of the nitrogen on top of the trap is the same as the air pumping through it. Is there too much nitrogen evaporating for the air to get close, or is the surface of the nitrogen too warm, or some other reasons?

There's an additional factor in a liquid nitrogen trap that makes it dangerous, and that is that you often condense organic solvents in it, which can be ignited by the liquid oxygen.

There is a worrying tendency in answers here to explain why something doesn't happen when, in fact, it does happen.

@matt_black We leave our glovebox cold trap up for hours at a time, and I've never seen any blue when pouring it out.

If you look at the vapour pressure graph for oxygen you’ll see that at its partial pressure in air, it can only liquefy below about 77 K, which by coincidence is the temperature of boiling liquid nitrogen. The air immediately above the liquid nitrogen / air interface cannot be any colder than that 77 K (slightly less, because nitrogen is only 4/5ths of the air), so the 1/5th that is oxygen cannot actually condense – it doesn’t get cold enough. Adding to that is the fact that the evaporating nitrogen further dilutes the oxygen that is present, further lowering the temperature that would be needed to condense oxygen at that pressure.

At atmospheric pressure, oxygen liquifies at 90 K, not 77 K.

@ColinMcFaul, where do you find oxygen at 1 bar other than in, say, the Apollo 1 capsule? If you think of oxygen as "water", then just above the liquid nitrogen surface, where it is just a tad under 77K, the "relative humidity" is at most (thumbsuck) 97% - still under 100%. That means it can't "rain" liquid oxygen. Since evaporating nitrogen dilutes what oxygen there is, the air might be even "drier" than that 97%. Even less chance of oxygen rain.

Why does oxygen condense in vapour traps then?

The link is broken...

The boiling point of nitrogen is 77 K, and the boiling point of oxygen is 90 K, so oxygen is liquid when liquid nitrogen is around. Wikipedia repeats the warning you mentioned:

Liquid nitrogen has a lower boiling point at −196 °C (77 K) than oxygen's −183 °C (90 K), and vessels containing liquid nitrogen can condense oxygen from air: when most of the nitrogen has evaporated from such a vessel there is a risk that liquid oxygen remaining can react violently with organic material.

So it's not because the surface of the nitrogen is too warm.
My guess is that the pressure due to the evaporating nitrogen is pushing air away from the opening. LN2 has an expansion factor of 694, which you can compute by dividing the density of the liquid (0.808 g/mL) by the density of the gas at 20 °C (1.17 g/L). Of that expansion factor, a factor of 303 K/77 K = 3.94 comes from warming the gas from 77 K to 20 °C. The other factor of 176 comes from turning the liquid into a gas. That expansion is pushing air out of the headspace of the dewar.

Even if a little liquid oxygen does condense in the dewer, the partial pressure of oxygen will still be lower than normal air unless significant quantities of oxygen condense, which seems unlikely given the expansion. There is also typically nothing combustible inside a dewer, unlike a trap which you just condensed solvent in!

The fact that liquid oxygen can exist at 77K doesn't mean that oxygen at 77K and a partial pressure of 0.21bar will condense. In fact, your liquid oxygen at 77K will (very slowly) evaporate if I'm reading that vapour pressure graph correctly.

@BerndJendrissek, my understanding of that vapor pressure graph is that it is the total pressure that is relevant, not the partial pressure of oxygen. As a comparison, water boils when its vapor pressure reaches 1 bar.

@ColinMcFaul Pure water boils when its vapor pressure reaches 1 bar.  Dilute it with, for instance, ethanol, and the boiling point is depressed.

The boiling point occurs when the vapor pressure equals the ambient total pressure. That's why the BP of water decreases with elevation- lower atmospheric pressure. The boiling temperature doesn't change from day to day depending on how humid it is.

Actually, it does.
Normally, though, you won't notice it as it will be heavily diluted by the nitrogen and won't ever reach a concentration large enough to worry about. It's different if you pull air through a trap as the oxygen content of the flowing gas will be condensed preferentially leaving a liquid that is concentrated in oxygen. This is really dangerous if there is any organic material in the trap as well which is why those of who have worked with such systems are so wary of pale blue liquids in our traps.
But, if you like dangerous experiments, you can do some that illustrate the presence of oxygen. If you cut of the top of an aluminium can and half fill it with liquid nitrogen, then wait until most has evaporated, the remaining liquid will be mostly oxygen. The safe way to show this is to use a magnet on a stick to pull some out (oxygen is fairly strongly paramagnetic). Or, much less safely, you could drop a lit match in (this is very dangerous as the can will probably burn violently). There is even a youtube demonstration of this (I think it is number 5 of 7 "super cool" demonstrations with liquid nitrogen). I think they used to have a version with burning but may have taken it off for safety reasons.

Oxygen can and does condense into liquid nitrogen dewars. Good lab practice dumps vacuum line dewars Friday afternoon, refills and restarts Monday morning. Remember – the universe hates you.
Dumping old liquid nitrogen in the parking lot can be fun. LOX plus asphalt paving (or wooden floors) is a contact explosive for days. Charcoal briquettes plus LOX was a blasting combo in coal mines. NEVER pour LOX on burning charcoal briquettes. The fireball burns so hot you can get UV burns, assuming you have any skin remaining. Steel wool plus LOX is a thermal lance, minimum 2730 °C, sometimes claimed to be 4500 °C.
Take care that anything you wear when working with cryogens can be removed very fast – especially porous insulating gloves.

Vapor pressure is property of liquid and partial pressure is property of a gas in a mixture.
The normal boiling point for the pure component in liquid phase is the temperature at which vapor pressure exerted by the liquid becomes equal to 1 atm.
In a gas mixture, as long as partial pressure of a component remains less than its liquid phase vapor pressure at the temperature of the mixture, the component will not condense. Condensation starts only from the point when $PP=VP$ and starting from the case where $PP>VP$, condensation continues only till $PP=VP$. When $PP<VP$, condensation stops.
Air at 1 bar total pressure will have 0.21 bar partial pressure of $\ce{O2}$ and say it is at -183C (the normal boiling point of $\ce{O2}$). Vapor pressure of liquid $\ce{O2}$ at this temperature is 1.01 bar, but partial pressure is 0.21 bar  which is less than 1.01 bar so $\ce{O2}$ should not condense. But if Air is at 5 bar total pressure, this will have $0.21\times 5=1.05$ bar partial pressure of $\ce{O2}$ and this air is at -183C then $\ce{O2}$ will condense till its partial pressure has gone below 1.01 bar.
Vapor pressure is independent of composition in liquid phase but dependent on temperature. As temperature increases, the vapor pressure of a liquid component increases. Vapor pressure does not depend on gas phase mixture properties.
At a given temperature, the component having the higher vapour pressure in liquid phase is the more volatile components and will have lower boiling point.

Welcome to Chem SE. I am guessing your answer is in response to one of the other answers. Typically, that activity should be in the form of a comment instead of another answer. You no doubt posted it as an answer because you don't have enough reputation to comment. It only takes 50 rep to be able to comment (on others' posts), which you can get to pretty quickly by asking, answering and editing. Each vote up is 10 rep, so you'd only need to get voted up five times. Good luck!


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What is the effect on humans of using dry ice and liquid nitrogen as food additives?
I have seen on a TV show that some  guy uses dry ice (solid $\ce{CO2}$) to make ice cream and on an another TV show, liquid nitrogen ($\ce{N2}$) is used with tequila.  Both are edible.
I know that both gases are inert, but could these gases (perhaps in larger concentrations) be poisonous or create any side-effects in our body?  If these gases don't create any side effects, what mechanism(s) does the body use to remove these gases from the body?

To clarify, the ice cream and tequila are editable. Ingesting dry ice or liquid nitrogen would be life threatening.  Recently there was a woman in the news that drank a cocktail with liquid nitrogen in it and literately damaged her esophagus and blew her stomach apart.  http://wgntv.com/2015/09/22/woman-has-stomach-removed-after-drinking-liquid-nitrogen-cocktail-uk-bar-fined/

Nitrogen has a boiling point of 77K (-196 °C), CO2 melts at 195K (-78 °C). You don't cool down the ice cream or whatever else you want to eat to those temperatures, that would be very unsafe. Both will evaporate or sublimate quickly when they are in contact with warmer material, so there won't be much left once you eat e.g. the ice cream.
Air is composed of 78% N2 and 0.04% CO2, nitrogen gas can only be dangerous when it replaces all the oxygen in the air. Carbon dioxide is toxic at around 7-10% concentration. Both conditions are not realistically achievable by the methods you describe.
The biggest danger I see would be ingesting a piece of dry ice shortly after it was put in a drink, which would be cold enough to hurt.

The take-home message is that both liquid nitrogen and dry ice have such high vapour pressures that in any sensible scenario you would not actually ingest any above and beyond what is already present in the atmosphere. Less of a food additive and more of a food process. Similarly, supercritical CO2 or dichloromethane is used to decaffeinate coffee but none is retained by the final product.

Liquid nitrogen is a dangerous substance to handle without proper knowledge. While food created using it will be perfectly safe (79% of the air you breathe is $N_2$, so no risk there), there are dangerous in the perpetration process.
Someone in Germany lost both their arms while trying to use Liquid Nitrogen to produce food at home when they poured the left-over liquid nitrogen down the sink, causing an explosion which took their arm off. (See Bluementhal Style Chef Blows Off His Hands)
$CO_2$ is produced by your body in respiration and breathed out normally. It is also present in carbonated fizzy drinks and sparking wine. The only risk for dry-ice would be if you ate a lump of it, where you might suffer from the cold if it got stuck in your throat.

The sillyness of liquid nitrogen as a coolant lies in the fact that it has a low heat of vaporisation. "Chef" use is for show, nothing else, a problem in America predominantly. With respect to that ominous explosion in Germany, do You have some reference (date, names?)

To add to the answer from Nick about the dangers and from Georg about the use of liquid nitrogen for flair, a woman recently had her stomach removed to save her life after she consumed a liquid nitrogen alcoholic beverage.  You can read about it at on the Yahoo! News web site.
One must be careful with these things and even dry ice can be dangerous enough to cause burns.


text:
Deliquescence/efflorescence: How can I convert between salt concentration and RH?
I am looking at two phase diagrams, one has axes of $T$ and $\%\ce{MgClO4}$ in a $\ce{H2O-MgClO4}$ mix, and the other has axes of $T$ and relative humidity. The authors of the second phase diagram say "The stability diagram is modiﬁed from [the earlier phase diagram] by converting salt concentration to relative humidity using the Pitzer model and parameters." How does changing the salt concentration affect the relative humidity, and how can I reproduce this calculation?
(I am a geologist with minimal chemistry background).

Top figure from Chevrier et al., GRL, 2009; bottom figure from Gough et al., EPSL, 2011.

You can't. The two phase diagrams are for two different parts of the system. There isn't enough data to go between the two phase diagrams. You have to use the two diagrams together to estimate the missing data (Relative humidity as a function of temperature over wt% magnesium percolate).
The top figure (Chevrier et al., GRL, 2009) is for a salt in solution. Think of having a sealed bottle with air over solution. The phase diagram shows  the relative humidity of the gas as a function of temperature. So at about 240 degrees the equilibrium system has about 45 wt% magnesium percolate in solution.
The solid circles in the lower figure (Gough et al., EPSL, 2011.) show relative humidity as a function of temperature. So at 240 degrees the relative humidity is about 50% for the the equilibrium system.
So at a temperature of equilibrium 240 degrees a solution of magnesium percolate will have about wt 45% of salt and the relative humidity of the air over the solution will be about 50%. You have to combine data from the two phase diagrams to figure out the whole system.
The open circles in the bottom bottom figure (Gough et al., EPSL, 2011.) shows a solid salt in various hydration states (5 vs 6 waters). This sort of detector would be like a crystal oscillator with a coating. As the hydration in the coating changes, then mass of the oscillator changes which changes its fundamental frequency. The chemical reaction (between hydrates with 5 and 6 molecules of water) is not temperature dependent hence the near vertical line.


text:
Remove formalin from food using chemical found in kitchen
In some parts of the world food is stored with formalin so it looks fresh forever! This is astonishing, but true (See refs here). As formalin is very bad for health and damages kidneys and liver very quickly, I am looking for a way to remove it.
When I buy fish, I know it has formalin in it. Every fruit I bought has formalin. I remember I bought 2 kilgrams of apples once. And it lasted for 1 month, looking fresh whole time. Then I threw it away.
My question is: how can you remove formalin from food? (Especially fish and fruit). It's best if we can use a chemical found in the kitchen. If not, I'd think one that is found at a chemical store.

This is astonishing but true. Do you have a reference for that information?

Firstly, this formalin business sounds a lot like one of those hoax chain emails (though I can't find anything to support that). Secondly, are you sure that it is formalin? There are hundreds of other preservatives which can keep stuff fresh for a long time if added in sufficient quantity. Thirdly, I think that any chemical method of removing formalin will involve adding more harmful chemicals to the object, so that won't work either. :\

@Manishearth - Consult this WP page and references therein. Outrageous, but then people put melamine in milk to fool protein tests. :(

@richard Oh, there is a lot of this where I live (fortunately, it's avoidable). Thanks for the WP ref--note that this doesn't ensure that the observed preservative is indeed formalin. But that's irrelevant to the question

@shiplu Remember, Googling stuff may not give reputable sources. I've seen hoaxes that get backed up by a quick Google Search. In this case, though, the refs on the WO article are good enough. Note that it's not just Bangladesh--Indonesia, Vietnam, and probably some other countries as well.

There is no known method of removing formalin from fish and fish products. A chemist may be to tell if it can by repeated washing of fish or after cooking for certain time. My suggestions are to buy fresh local fish, preferably live one. Develop awareness among the value chain not to sell formalin added fishes. Reduce import of fish from the countries who add formalin. All consignments should be tested compulsorily and certified at the entry point. The fish traders found with formalin added fish should be punished with 5 years imprisonment. [contd]

DFO, UFO and other relevant officers should be made responsible for their area of jurisdiction for providing safe food to the people. They should be brought under accountability. Care should be taken that small fishes like mola, kaski etc are not adultered with formalin.

Sorry to tell you, but real, healthy apples can easily have a month shelf life.

While the idea of filtering formaldehyde from food might sound naive, the motivation is understandable and the problem apparently real.
I've found at least one publication on the topic:
T Yeasmin, *M S Reza, M N A Khan, F H Shikha and M Kamal
PRESENT STATUS OF MARKETING OF FORMALIN TREATED FISHES IN
DOMESTIC MARKETS AT MYMENSINGH DISTRICT IN BANGLADESH
Int. J. BioRes. 2010, 1, 21-24
From the abstract:

[...]Study showed that formalin was not detected in any
  fish produced locally, but was detected in the imported ones of catla and rohu ranging from 0.5% to 1% which was
  sold in different markets of Mymensingh with comparatively lower price than those produced locally. The shelf life of
  the locally produced fish was much longer than those of imported fish which may be related to the loss of shelf life
  during its transport from importing countries. [...]

Concerning the initial question:
I'm afraid but i don't see any chance to remove formaldehyde from food without turning it to another toxic cocktail.
But since formaldehyde was mostly used to increase the shelf time of imported fish, buying from markets predominantly selling local products might be a way--if available and affordable.

""ranging from 0.5% to 1% which was sold"" did You ever handle/smell formaldehyde? .5 % would render a fish unsellable due to the pungent smell.

I do prefer the smell of geranial too ;) Seriously, whenever possible, i rather used paraformaldehyde in synthesis, e.g. for m-meconine on the route to fluorescent phthalimides. Even then, the synthesis isn't particular nice and thus best left to undergraduates :D

Sorry, but formaline is often used to kill parasites in shellfishes and other farmed animals, so the whole conclusion "But since formaldehyde was mostly used to increase the shelf time of imported fish, buying from markets predominantly selling local products might be a way--if available and affordable." is pretty questionable.


text:
What is the difference between temperature and thermodynamic temperature?
My course book says that kelvin is the unit of thermodynamic temperature, then what is the difference between temperature and thermodynamic temperature.

Temperature is more of a relative notion--heat flows from higher temperature to lower temperature, but that's about it. The zero in a temperature scale is meaningless; it is an arbitrary convention. Usually, we measure temperature in Celsius, Fahrenheit, Rankine, etc (note that Kelvin measures temperature as well--but it also measures thermodynamic temperature)
Thermodynamic temperature is such that zero temperature is absolute zero (unlike Celsius, where -273.15°C is absolute zero). It is related to the energy (specifically the kinetic energy) of the molecules of a substance-- for example, in a gas, kinetic energy per molecule is $\frac12k_BT$ (rotational energy, etc can be obtained if you know the type of molecule--but temperature is principally a measure of kinetic energy). It also has strong ties with entropy.
We generally don't use "temperature" in thermodynamics, thermodynamic temperature is more convenient. Temperature is only useful in equations involving temperature difference--whereas equations involving the absolute value of temperature (Arrhenius equation in kinetics, entropy, Maxwell distribution equation, etc) demand thermodynamic temperature.

Whenever you use a temperature in any calculation rather than a temperature difference, you have to use Kelvin! (Or another thermodynamic temperature scale.) If you are working out chemical reaction rates, for example (Which aren't normally thought of as thermodynamics but kinetics.)

Do these equations involve general gas equation? @ManishEarth

There is no difference between temperature and thermodynamic temperature. Wikipedia has a nice writeup. The book likely uses the adjective in order to differentiate from ill-defined conventional connotations. Temperature is an ensemble property, e.g., an ideal gas has an average energy of $\frac{1}{2}k_BT$. Outside of equilibrium, temperature is ill-defined.

https://en.wikipedia.org/wiki/Thermodynamic_temperature is the absolute temperature
https://en.wikipedia.org/wiki/Temperature is the ensemble propery
Heat radiation is based on fourth power the absolute Temperature (e.g. Kelvin), here you cant take Celcius or Farenheit (without convertion)
Also the thermal energy is approximate proportional to the absolute Temperature, not the relative Temperture. So there is a difference.

"Thermodynamic temperature" is a misnomer.  It arises from the misinterpretation of the double adjective used in English.  For example, if we say "Jane is a Chinese History Professor," does this mean that "Jane is a Professor of History who happens to be Chinese" or that "Jane is  Professor of Chinese History"?  For the case of temperature, does "thermodynamic temperature scale" mean "thermodynamic scale of temperature" or "scale of thermodynamic temperature"?  It appears that when the kelvin was introduced into the SI, the former was intended but the physical quantity was (falsely) interpreted as the latter.


text:
What is the exact nomenclature of ionization?
I'm a little confused as to what reactions are considered ionizations. Is there any requirement for the educts, for example?
Some references suggest that it's only considered an ionization if you produce additional charged atoms/molecules from the educts (and some processes are dubbed detachment, which intuitively look like ionizations to me). Then "double ionization" wouldn't be possible, strictly speaking. To what extend have electrons to be produced. At the moment I think that it's legal to call most processes that look like ionization ionization.
There is a remark in the wikipedia article regarding the distiction from dissociation, but I don't really understand it.

The reference for chemical terminology is the IUPAC Gold Book, or IUPAC Compendium of Chemical Terminology. The entry for ionization is as follows:

The generation of one or more ions. It may occur, e.g. by loss of an electron from a neutral molecular entity, by the unimolecular heterolysis of such an entity into two or more ions, or by a heterolytic substitution reaction involving neutral molecules […] The loss of an electron from a singly, doubly, etc. charged cation is called second, third, etc. ionization. This terminology is used especially in mass spectroscopy.

It includes examples that you might find instructive. It confirms your intuition that the term has a broad understanding, covering a wide variety of reactions.


text:
Extracting elemental ratios from X-ray photoelectron spectroscopy (XPS)
Is it possible to extract information on elemental ratios from X-ray photoelectron spectroscopy (XPS)?  For example, I am reading a paper and in it they say that the carbon to oxygen (C:O) ratios of the four samples are 2.8, 2.9, 5.1, and 23.3.  However, the authors do not provide the actual XPS spectra, not even in a Supporting Information-type addendum.
I am a computational chemist and have no experience with XPS.  I do know that elements, and even particular hybridizations, give characteristic peaks at characteristic binding energies (in units of eV, for example).  This, I know, enables qualitative analysis of compounds.  But is there a way to extract quantitative information from XPS?  Suppose I want to determine the C:O ratio in a compound containing only C, O, and H.  Do I take the ratio of the C and O peak heights, or do I take the ratio of the area under the peaks corresponding to C and O?
Do you know of any good review articles that would help me to understand the basics of XPS?

""and even particular hybridizations, give characteristic peaks at characteristic binding energies (in units of eV, for example)"" You mix up photoelectron spectroscopy and X-Ray-PES

@Georg Thanks. What is the difference between photoelectron spectroscopy and X-ray-PES?

The latter uses Röntgen-rays, the other vacuum-UV to set free those electrons. That means that only the highest level electrons are expelled (carrying information on chemical bonds) or electrons "deep" in the atoms are expelled by Röntgen-rays, giving information on elemetary composition.

@Georg XPS can also give chemical information (through less sensitive than UPS).

I found some information in Skoog, Principles of Instrumental Analysis, 6th ed.  On page 597:

Quantitative Applications.  Once, XPS was not considered to be a very useful quantitative technique. However, there has been increasing
  use of XPS for determining the chemical composition of the surface
  region of solids.  If the solid is homogenous to a depth of several
  electron mean free paths, we can express the number of photoelectrons
  detected each second $I$ as $$I = n \phi \sigma \epsilon \eta A T \ell$$ where $n$ is the number density of atoms ($\text{atoms cm}^{-3}$) of the sample, $\phi$ is the flux of the incident X-ray
  beam ($\text{photons cm}^{-2} \text{ s}^{-1}$), $\sigma$ is the
  photoelectric cross section for the transition
  ($\text{cm}^{2}/\text{atom}$), $\epsilon$ is the angular efficiency
  factor for the instrument, $\eta$ is the efficiency of producing photo
  electrons ($\text{photoelectrons/photon}$), $A$ is the area of the
  sample from which photoelectrons are detected ($\text{cm}^{2}$), $T$
  is the efficiency of detection of the photoelectrons, and $\ell$ is
  the mean free path of the photoelectrons in the sample ($\text{cm}$).
For a given transition, the last six terms are constant, and we can
  write the atomic sensitivity factor $S$ as $$S = \sigma \epsilon \eta A T \ell$$  For a given spectrometer, a set of relative values of $S$
  can be developed for the elements of interest.  Note that the $I/S$ is
  directly porportional to the concentration $n$ on the surface.  The
  quantity $I$ is usually taken as the peak area, although peak heights
  are also used.  Often, for quantitative work, internal standards are
  used.  Relative precisions of about 5% are typical.  For the analysis
  of solids and liquids, it is necessary to assume that the surface
  composition of the sample is the same as its bulk composition.  For
  many applications this assumption can lead to significant errors.
  Detection of an element by XPS requires that it be present at a level
  of at least 0.1%.  Quantitative analysis can usually be performed if
  5% of the element is present.


text:
Cyanogenic glycosides and HCN
I'm researching cyanide for a biology project and noted that one of the largest sources of carbohydrates in the tropics is the cassava root -- which I previously researched as a food that contains a cyanogenic glycoside that is metabolized into hydrogen cyanide.
Most of the sources I reviewed claimed that, while boiling the cassava, the cyanogenic glycosides were broken down into components such hydrogen cyanide, which made the plant safer to consume. I have a limited understanding of chemistry and can't resolve why this is considered a safe way to cook the plant.
My research indicates that inhalation of hydrogen cyanide is significantly more toxic than the oral route, so why would this method be considered beneficial? Not sure why this is considered safer, unless the information I've found is not correct. Any help in putting this together would be appreciated. Thanks!

OK, I was too lazy to look up LD50 and LC50 of $\ce{HCN}$ for you ;)
But if you have a chance, have a look at

G. Padmaja & Keith H. Steinkraus,
Cyanide detoxification in cassava for food and feed uses,
Critical Reviews in Food Science and Nutrition, 1995, 35(4), 299-339 [DOI]

Seemingly, the processing of cassava (manioc) involves several steps, such as drying in the sun, grating, watering and cooking, each of which either facilitates the enzymatic or thermal cleavage of the linamarin (i.e. the toxic compound) to glucose or supports the release of $\ce{HCN}$.
In total, the processing minimizes the risk of exposition through inhalation. None of the steps is perfomed in confined space, watering helps to wash out the $\ce{HCN}$ which in water dissociates a bit ($\ce{HCN}$ is a weak acid).
Thus, digestion of the untreated roots is likely to be fatal, while the slow processing is apparently not.

Ok, I digged a bit and maybe you have you have access to this one:

Bala Nambisan
Strategies for elimination of cyanogens from cassava for reducing toxicity and improving food safety Food and Chemical Toxicology, 2011, 49, 690–693 [DOI]

Thanks for the great answer. I hadn't seen any reference to the drying and grating process which would most likely release the HCN. I don't have access to this article and the ones I've read referenced the boiling process. It does make me wonder whether the boiling process would be enough to facilitate the thermal cleavage of HCN, since the enzymatic compounds may not be available to the linamarin without the grating process. Anything from the article give insight on that question? Thanks again for your answer.

My pleasure :) I'm afraid i don't have full access to the article either.

But everybody has access to Wikipedia

Thanks Georg. If you can find any relevant information on Wikipedia that would be helpful, please feel free to include it here. I've already read the pages on hydrogen cyanide, linamarin, and cassava, but if there's anything I'm missing let me know.

If You read the forum rules, You will find that asking questiones easily anwered in wiki as well as links/quotes from Wiki are not welcome. You read the article on Cassava in English? I doubt that!

And we can use google too @Greg !!
"The cyanide ion (CN-) is the toxic moiety in hydrogen cyanide. The toxicity of simple cyanide salts, such as potassium and sodium cyanide is, therefore, similar to that of hydrogen cyanide." and from link "Toxicities: Cyanide is the least toxic of the "lethal" chemical agents. The LCt50s of AC and CK by inhalation have been estimated to be 2500-5000 mg"min/m3 for AC and about 11,000 mg"min/m3 for CK. LD50s for hydrogen cyanide have been estimated to be 1.1 mg/kg for intravenous administration and 100 mg/kg after skin exposure. The oral LD50s for sodium and potassium cyanide are about 100 and 200 mg/kg respectively." AC - the term refers to the anion CN-, or to its acidic form, hydrocyanic acid (HCN).
Have a look at link absorption, metabolism and toxicity of cyanides

Thanks, I appreciate your response. Like I mentioned in my question, I don't have a background in chemistry so I'm not familiar with some of the terminology, such as the LCt50 ratings. The toxicity thresholds mentioned in the second link suggest 150 mg/m3 as a lethal exposure via inhalation to HCN gas. I haven't found data on the oral LD50 of linamarin in humans -- but in rats it is around 500 mg/kg, so I think we can assume it is substantially higher in humans. I'll look into estimating the concentration of the vapors in a typical kitchen and see if I can find the answer there.

This "answer" is off topic, nobody spoke of toxicity of various cyanic chemicals. Who is that "Greg"? Me?

The answer is as actual as could be - If u read all the resources provided in links u would realize that, but who cares, the better way is to send one answer down. is that correct Georg ?


text:
How to find the band gap for a compound?
I have processed a si/polymer hybrid using a silane coupling agent which I need to determine the energy gap / band gap.  Would you please let me know the procedure to go about determining the band gap?

The energy gap being asked about the band gap between the valence band and the conductance band of the material. It is roughly analogous to the HOMO-LUMO energy gap of small molecules. I am also slightly confused about the fabrication process, but the methods of determining these energy gaps are generally independent of the material and how it was made.

People might be more familiar with the term "band gap" than they are with the term "energy gap," so I suggested an edit to that effect.

"Band" is appropriate in case of bands, i e. for metals and semiconductors or insulators in general crystalline solids. user 423 has done some si/polymer thing, nobody knows what that is, not even whether the answer below is wisedom or speculation. I had asked here for more information from user423, but somebody erased my comment. Manisherth?

Georg, polymers are usually considered to have "bands" also. In a conducting polymer, by the time you get to a degree of polymerization of 10 or more, the pi orbitals become so close in energy so as to begin to behave like bands in semiconductors. I agree that user433 could be clearer about what the material in question is, but that doesn't change the methodology used to determine a band (or HOMO-LUMO) gap, which is what the question asks.

There is no information on the type of polymer and what that "si" means. Maybe some silicon polymer? or some polymer bound to a silicon wafer? Answerinf such vague questions reveals "eggboundedness".

There are two ways in which the band gap of a material: using optical spectroscopy and using electrochemistry. The way you go about analyzing the data to determine the band gap is independent of material, but the process of collecting the data is very specific to your material, and you should consult papers where others have made similar measurements on similar systems.

Optical spectroscopy.

You need to collect an absorbance spectrum of your material. Given what I can gather from your fabrication description, you will probably need to collect said spectrum in the solid state. The exact means of doing that vary somewhat by spectrometer model. Find someone who knows your spectrometer very well and ask them how to do it (or read the manual).
The absorbance process measures the energy difference between the ground state and the excited state as electromagnetic radiation is absorbed. Since there are a number of vibrational states associated with both the ground state and the excited state, the absorbance spectrum appears as a blob instead of a nice line. The "true" band gap, or $E_{00}$, the energy difference between the lowest vibrational states of the ground and excited states, is difficult to measure. However, the most commonly accepted approximation is the onset of absorption from the low energy side. The onset value is the intersection of the extrapolations of the linear parts of the spectrum. See the image below.

Once you have the wavelength (x) at that point, you can convert it into energy using the Planck relation:
$$E(\text{J})=\dfrac{hc}{\lambda(\text{m})}$$
If you want your energy in electronvolts, you can convert directly:
$$E(\text{eV})=\dfrac{1239.84187 \text{ eV nm}}{\lambda(\text{nm})}$$

Electrochemistry, specifically voltammetry

Voltammetry gives you the ability to determine the peak oxidation potential, $V_{\text{ox}}$, and the peak reduction potential, $V_{\text{red}}$, of your material. Both cyclic voltammetry and differential pulse voltammetry can give you these data, and both should work on solids, though you may want to find someone who has experience in such things.
The assumption that we operate under when using electrochemistry to determine HOMO-LUMO or bad gap, is that when we oxidize a molecule or material, we take the first electron out of the HOMO (or valence band). When we reduce the molecule or material, we put an electron into the LUMO (or conductance band). Thus, $V_{\text{ox}}$ tells us something about the HOMO (valence band) and $V_{\text{red}}$ tells us something about the LUMO (conductance band). The energy gap (in electronvolts) is simply the difference between the two potentials (in volts) multiplied the charge on the electron $e=-1 \text{ eV/V}$.
$$E(\text{eV})=e(V_{\text{red}}-V_{\text{ox}})$$


text:
Algebra of conserved quantities for chemical reactions
So there is the standard formation enthalpy $\Delta H^0_f$ and there is also formation entropy.
Are there more (independent) quantities? What about the heat capacity? Is there a general theoretical structure where one can show one got all the conserved quantities associated with a chemical reaction? It seems to me that one could this on a quite rigid structure and invesigate the properties of chemical processes from that point of view.
And to what extend are the procedures like using Hess's law just rules of thumb or what are some insights about he deeper structure where they come from.

Oh yes of course! Ever read about Lavoisier?

No really algebra, but electron spin is conserved, this is why singlet oxygen is very reactive.

What I think you are asking about is the number of thermodynamic parameters which characterize a chemical reaction. Those parameters are used to make quantitative predictions of the final state of the reactive system (at least in the context I'm familiar with, i.e. chemical engineering); for example, the amount of heat exchanged with the reactor's environment, the composition of the mixture at equilibrium, it's material phase, overall density, etc.
In classical thermodynamics, the way to address this problem is through the potentials defined on the system of interest, and studying equilibrium as the extremum of one these. The chosen potential depends on the particular configuration of the system under study; and in all cases their relation with equilibrium is a consequence of the Second Law. For example, if the system is restricted to evolve at constant volume and entropy, the minimized potential is the internal energy; if it evolves at constant energy and volume (an isolated simple compressible system) it maximizes entropy; etc.
From a practical point of view of natural systems and industrial processes -which evolve at nearly constant temperature and pressure fixed by the environment- the involved minimized potential is the Gibbs free energy. That the minimization of $G$ yields an equilibrium state for a closed system at constant pressure and temperature is straightforward (it can be found in any physical chemistry textbook). Moreover, when we say here that the system at it's minum $G$ is at equilibrium, we are saying it's at overall equilibrium: it's simultaneously at mechanical equilibrium (because of null pressure gradients), thermal equilibrium (because of null temperature gradients) and in chemical equilibrium (because of null chemical potential gradients); so there's no "force" (in a generalized sense) to drive any further process.
So, going back to your question, which are the independent variables of a reactive system as such? Well, suppose that in your system you have $P$ chemical species which can be both reactives or products of $Q$ reactions taking place in a monophase. For every species, the balance on the number of moles of each one yields:
$n_{i}=n_{0i}+\sum_{j=1}^Q{\nu_{i}^j\cdot \xi_{j}}$
Where $\xi_{j}$ is the extent of the j-th reaction.
Then, we can define the chemical potential as the partial molar Gibbs free energy per species:
$\mu_{i}=\displaystyle\frac{\partial G}{\partial n_{i}}$
And also, given that $G$ is an extensive property (or better phrased, a homogeneous function of first degree in mass) by Euler's Theorem it follows that:
$G=\sum_{i=1}^P\mu_{i}n_{i}=\sum_{i=1}^P\mu_{i}\cdot (n_{0i}+\sum_{j=1}^Q{\nu_{i}^j\cdot \xi_{j}})$
Every chemical potential can be expressed as the sum of it's specific Gibbs free energy of it's pure form, at the temperature and pressure of the mixture, plus a term which ponderates the influence of concentration and chemical affinity of each particular species towards the rest of the present compounds which constitute the mixture (for example, the solvent matrix, if one species is so concentrated that it can be considered so). The latter term contains an adimensional parameter known as the activity of the species:
$\mu_{i}=\mu_{0i}(P,T)+RTln(a_{i})$
The zero index in the first term just means that it's referred to a pure system. This specific free energies can be related to the formation enthalpies and entropies which are tabulated through:
$\mu_{0i}=h_{0i}(P,T)-T\cdot s_{0i}(P,T)\\\mu_{0i}=\left\{{h_{fi}^0(P_{0},T_{0})+\displaystyle\int_{(P_{0},T_{0})}^{(P,T)}dh_{0i}}\right\}-T\cdot \left\{{s_{fi}^0(P_{0},T_{0})+\displaystyle\int_{(P_{0},T_{0})}^{(P,T)}ds_{0i}}\right\}$
Where $h_{f}^0(P_{0},T_{0})$ and $s_{f}^0(P_{0},T_{0})$ are the standard formation enthalpy and entropy, respectively. The integrals evaluate the variation in enthalpy and entropy of the pure compounds due to a difference in temperature and pressure in reference to standard values (25 ºC and 1 bar). The evaluation of the integrals in terms of measurable properties must be done using Maxwell relations and definitions of potentials:
$dh_{0i}=\displaystyle\frac{\partial h_{0i}}{\partial {T}}dT+\displaystyle\frac{\partial h_{0i}}{\partial {P}}dP=C_{p}dT+\frac{dP}{\rho}\\ds_{0i}=\displaystyle\frac{\partial s_{0i}}{\partial {T}}dT+\displaystyle\frac{\partial s_{0i}}{\partial {P}}dP=Cp\frac{dT}{T}-\frac{\alpha}{\rho}dP$
What about the activities? These are non-linear functions of the composition, temperature and pressure of the mixture, and usually that non-linearity is expressed through an activity coefficient (it may be the case that further coefficients are introduced theoretically, but all they do is to "partition" the contributions to non-ideality into a product of these, and so here I talk about a unique coefficient without loss of generality). So:
$a_{i}=x_i\cdot\gamma_{i}(P,T,x_1,x_2,...,x_P)$
And:
$x_i=\displaystyle\frac{n_i}{\sum_{k=1}^P n_k}=\displaystyle\frac{n_{0i}+\sum_{j=1}^Q{\nu_{i}^j\cdot \xi_{j}}}{\sum_{k=1}^P\left\{{n_{0k}+\sum_{j=1}^Q{\nu_{k}^j\cdot \xi_{j}}}\right\} }$
Finally, the (enormously complicated) expression for G turns out to be dependent of the variables:
$G=G(P,T,h_{fi}^0,s_{fi}^0,Cp_{i},\rho_{i},\alpha_{i},n_{0i},\xi_j) \ \ \ \ \ \ i: 1 < i < P \ \ \textrm{and}\ \ \ j: 1 < j < Q$
Those are, a priori, the "independent quantities" involved in the overall reaction's equilibrium. Once you've fixed the temperature, pressure, the original amount of substances and get the intensive properties stated for all the chemical species present in the system along it's evolution, G becomes only a function of the extents of the reactions:
$G=G(\xi_j) \ \ \ \ \ \  j: 1 < j < Q$
And the equilibrium is reached when $G$ is at minimum; that's a problem solvable through optimization. The values of the extent of every reaction at equilibrium is enough to calculate every other useful parameter of the system: it's composition, the heat involved, the change in overall volume, etc.
Summary: the conserved quantities in a chemical reaction are physical constraints (conservation of total mass, balance of moles according to stecheometry, conservation of energy exchanged with the environment, strict increase in entropy, etc) but that's not enough information to uniquely evaluate the equilibrium parameters of a general reaction (or reactions). Most of the expressions I left in an undefined closed form in this answer (such as the activity coefficients, all of them functions of compositions) must be experimentally correlated, or modelled with some theoretical formalisms, generally highly complex. The standard thermodynamic potentials might be tabulated though, as it's the case with the intensive properties of pure substances (specific heats, thermal expansion coefficients, etc).


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Electrical conductivity of graphite
On this Wikipedia page, the electrical conductivity of various materials are given in the third column ($\sigma \text{ (S/m) at 20}^\circ \text{C}$).  I am interested in the entry for Carbon (graphite):

$2$ to $3 \times 10^5 \text{ S/m} \perp \text{basal plane}$
$3.3 \times 10^2 \text{ S/m} \parallel \text{basal plane}$

Why is conductivity lower parallel to the plane than in the direction perpendicular to the plane?
I am surprised by the relative magnitudes of these $\sigma$ values because I thought that one of the amazing things about graphene -- an atomic layer of graphite -- was that its conductivity in the plane (i.e., parallel to the basal plane) is very high.  I would expect a similar trend parallel to the basal plane in graphite.
Do you have any suggestions for reading that I can do to understand this?

Good catch - you are absolutely correct. This is a mistake in the Wikipedia page, as a quick check with the reference listed there confirms. The values are the wrong way around.
Changing it now...
Edited to add: this of course raises the question of why it should be this way around. A somewhat hand-wavy answer is that, just like in benzene, the $p_z$ orbitals overlap within each graphene sheet so that the electron contributed by each carbon is delocalised in this plane (the other three are bound up in the three $\sigma$ bonds). So in a sense this is conjugation, but over a whole 2D plane rather than around a ring or along a single chain of atoms. On the other hand, the graphene sheets are rather further apart from one another than the covalently bonded atoms within a sheet (as you might expect since the sheets are only weakly bound through van der Waals interactions), and so the overlap is greatly reduced and conduction is much more difficult.
A more complete answer would rely on a study of graphite's electronic band structure, which is an active research area and beyond the scope of this response!

Just as a follow-up, why is the electrical conductivity higher in the parallel direction?

You assumed that already in Your question, why?


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What type(s) of gloves are effective against DCM and acetone?
I know that standard latex and nitrile gloves don't stop common organic solvents such as DCM and acetone. However, I was reading a Reddit thread and there were comments that if you are coming into contact with solvents you are just wearing the wrong type of gloves. In a non-manufacturer specific fashion, what types of gloves will stop DCM and acetone? What will it slow down long enough for me to have a chance to finish what I'm doing and change my gloves in a safe manner? I am looking for a guide on what gloves to use. I just had DCM on my mind since nitrile doesn't slow it down, and it burns. And acetone, since it is hard to stop.

There is a very common (and wrong) thought in this question: All or nothing, black and white etc. No glove made from any rubber will "stop" a solvent. Solvents (the smaller the molecule the faster) will diffuse through all rubbers, the diffusion rates may differ a bit, but not much. The problem to use some rubber or a different type is resistance of the rubber against swelling! Depending on polarity of rubber/solvent and degree of crosslinking, the rubbber will swell a little bit, or often so excessivelly, that it will tear from a touch.

If you're curious about solvent penetration through your gloves, I suggest pouring some solvent in to a glove, closing it at the top, and then seeing how long it takes to smell the solvent through the glove. Just be sure to waft carefully - some of the solvents will penetrate faster than you think!

When you try to find a suitable glove for acetone, don't ever let nitrile gloves come across your mind. Go and check for the permeation rate and breakthrough time in order to find the most suitable one, even neoprene gloves were not strong enough to prevent the breakthrough of nitrile.

somebody did run a test on youtube: https://www.youtube.com/watch?v=TlJj0JLmLBI

You can use tips from the site link - the table at the bottom or just google which glove can suit your needs.
As you can read from the link:

Nitrile gloves are Low cost, excellent physical properties, dexterity, but they are Poor vs. benzene, methylene chloride, trichloroethylene, many ketones. They are recommended for Oils, greases, aliphatic chemicals, xylene, perchloroethylene, trichloroethane; fair vs. toluene,
According to Agronne for dealing with acetone you should use Natural Latex/Rubber or Butyl gloves.
For DCM there are no good gloves - you could try using neoprenes

I just did a bit of an experiment. My lab uses Kimberly Clark nitrile gloves, and we have a huge jug of acetone. First, I put a glove on my hand, and poured acetone all over it. My hand felt cold, but not nearly as cold as it feels when I spill acetone on my bare hand. I allowed the acetone to evaporate from the glove (approximately 30 s), and removed the glove. There was absolutely no acetone on my hand or the inside of the glove. I then took a fresh glove, and poured enough acetone into it to fill the fingers with acetone. I held the opening shut for about 5 minutes, and poured out the acetone. During that time, the glove slowly felt thinner and easier to stretch. At the end, it was so thin that it tore with almost no effort. So for the particular pair of nitrile gloves and acetone, it looks like somewhere between 1 and 5 minutes' exposure will render the glove useless.


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What chemicals are present in burnt toast?
I'm planning to do an experiment where I extract burnt toast and test how that affects the growth of mung beans. I know burnt food contains PAHs, but what are some other chemicals present and how would they affect plant growth?

First of all, one should have look a look on the ingredients of your toast. It mainly consist of starch, proteins (2% maybe), fat (a few %), water and some inorganic compounds such as NaCl.
Well Salt normally won´t interact with anything below 1400°C so no need to look at this.
Water just evaporates, i think we must not discuss that here.
Fat: that's more interesting because fat is known to degenerate into different compounds when heated.
But the definitely most interesting reactions will be present when one heats carbon hydrates together with protein: so called Maillard Reaction(s) those kind of reactions are still issues of ongoing research although this type of reactions have been discovered long ago. As a good starter on the various compound which can occur during Maillard Reaction you might read Mottram or Monnier. Especially acrylamide could be of interest when looking a toxicity. I don't have a certain paper in mind but i know that there has been research about toxicity of acrylamide in fish and other animal. So it's likely that there was some research about plants.
When you look at starch and protein separately you will encounter heat degeneration as well as caramelisation.

I'll have to check with our resident expert, but I don't think that the MathJax rendering engine supports the small caps for some reason, so I just edited them out.

ok i just thought i would take a while. Maybe small caps would be a good idea. i think i will be quite common here (in chemistry) to cite someone.

I tried it using \small with caps, which does render, but it looks muddled, unfortunately.  I think MathJax only renders a subset of the full LaTeX.  I will poke around a bit about it, but since it is a 3rd party product, I don't think we have as much control over it.  For now, you can edit the names however you would like, but I just didn't want the \textsc in front of the name that wasn't being rendered.


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What salts are/are not birefringent?
Is there a quick way to determine if a salt (or ionic compound in general) crystal should be birefringent from its formula without having to know/look up the crystal structure?  When I'm looking for protein crystals, if I know the condition has a lot of salt in it, but it shouldn't be birefringent (e.g. NaCl), it's a hint that birefringent crystals I do see might be protein.
Is the only non-birefringent crystal type (e.g. not glasses) cubic?  Does it matter if it's primitive, face-centered, or body-centered?
Back to looking for a generalization, are most (all?) salts with polyatomic ions birefringent?    All +1/-1 (NaCl, KBr, etc...) aren't?

Why didn't You read the wikipedia article on birefringerence? http://en.wikipedia.org/wiki/Birefringence

@Georg Is there something in the article that you'd like to point out in a more constructive fashion without being accusatory?

@Georg I have absolutely no problem with you mentioning Wikipedia.  It's when you say why didn't you read. Instead of addressing a specific issue with the OP, you are accusing them of ignorance without any basis.  We have asked you repeatedly to treat other users with respect, so I again echo that sentiment.

Yes, the only crystal system that is guaranteed not to be birefringent is cubic, because all of the other crystal systems have at least one axis distinct from the others. What matters is the symmetry not the centring, so this is equally true for face-centred or body-centred as for primitive crystals.
Sadly, however, there is no way to determine the crystal structure, or even the crystal system, simply from consideration of the formula - indeed, crystal structure prediction and prediction of polymorphism are active research fields. Counterexamples to your first generalisation include ammonium chloride, bromide, and iodide; to your second include ammonium fluoride and silver iodide.

I probably should have said alkali halides instead of "+1/-1".  Do you have a good resource that lists all the systems for the common and somewhat common compounds?  WP is somewhat spotty.

With alkali halides you are probably on safer ground! For really common compounds try a book like Greenwood and Earnshaw Chemistry of the Elements. For less common crystal structures the standard resources are the Cambridge Crystallographic Database (mostly organic) and Inorganic Crystal Structure Database, both of which are subscription-based services, but the Crystallography Open Database (http://www.crystallography.net/) and American Mineralogist database (http://rruff.geo.arizona.edu/AMS/amcsd.php) are both free.


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Why does an ideal capacitor give rise to a rectangular cyclic voltammogram (CV)?
This question is sort of a sequel to my previous question about cyclic voltammetry (CV).  One of the responses made reference to the fact that an ideal capacitor gives rise to a rectangular cyclic voltammogram.  Can you please help me understand why this is the case?  In other words, why does an ideal capacitor reach a constant current $I$ as soon as a voltage $V$ is applied?
I indeed see nearly ideal CVs in many literature articles (CVs which are rather rectangular with rounded corners).  In other figures, though, I see relative deviation from "rectangles with rounded corners," in that I see abrupt peaks, spikes, or valleys.
For example, below I have plotted two figures from Khomenko, Electrochimica Acta 2005, 50, 2499-2506.  Just very roughly and "hand wavy," what might be the qualitative reason for the "rectangle with rounded corners" behavior of Figure 8 (left) and the "abrupt peaks" behavior of Figure 4 (right)?  Could it be that the sample in Figure 8 (left) is relatively unreactive toward apllied potential, whereas the sample in Figure 4 (right) undergoes redox (Faradaic) reactions -- indicating the presence of so-called pseudocapacitance -- when an external potential is applied?

Please know that I am not looking for an answer specific to the article to which I linked.  I am only asking this question in the context of basic, qualitative aspects of cyclic voltammetry.

The kind of analysis truly required is outside the scope of chemistry and would be best handled in the careful hands of an electrical engineer. I'll give a brief attempt that exploits an easy-to-use circuit simulator.
We have to first think of an equivalent circuit that we can use as a model to determine its behavior. I suggest the following:

where my variable is R1's value for its resistance, which I shall set to 0, 10 and 100 ohm. If we were to watch this happen in time then we would see the following:

Quickly converting these to Current vs Voltage and running two other simulations at different resistances and we get:

These results are due to setting up differential equations and solving them appropriately.
You can play with the circuit I made for yourself here.

This is veering into physics, but the reason for your initial observation - that an ideal capacitor has a square cyclic voltammogram, as shown in Chris's diagram - is that the charge on an ideal capacitor is proportional to the voltage across it, with the constant of proportionality being the capacitance: $q = CV$. Differentiating with respect to time, we see that the current across the capacitor is $I = \frac{\mathrm{d}q}{\mathrm{d}t} = C\frac{\mathrm{d}V}{\mathrm{d}t}$. Now when you put a sawtooth wave across the capacitor, $\frac{\mathrm{d}V}{\mathrm{d}t}$ flips back and forth between a positive and a negative value, hence so does $I$, giving the characteristic square shape.
Charge flowing for some other reason - such as a redox reaction - will indeed give the peaks shown in the second diagram.


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Lewis dot structures and exceeding the octet
I'm confused about the octet rule.  As far as I understand, you need to complete the valence shell of the elements by sharing electrons.
Now, carbon has 6 electrons which means it has 4 valence electrons, but because it has 6 electrons in total it exceeds the first shell with 4 electrons. The second shell has 4 orbitals with 2 electrons each, so to complete the second shell I need 8 valence electrons, but I had read somewhere that the octet rule holds true until down to the 4th period. How can that be when say, magnesium, has 12 total electrons and thus it exceeds the the second shell with 4 electrons but the third shell needs 18 valence electrons to be complete, not just 8?
Secondly, I read that if the central atom has more then 8 electrons, then that's ok too? What happens if I have a choice of exceeding the 8 electrons by forming another bond with an adjacent atom or sticking those electrons on the opposite side of the adjacent atom?
Is one strictly wrong/preferable?

I don't really understand your question, but you may want to see http://chemistry.stackexchange.com/questions/444/can-an-atom-have-more-than-8-valence-electrons-if-not-why-is-8-the-limit first.

For the first part of your question, you are correct that Mg has 12 electrons, which are arranged in shells in the pattern 2 - 8 - 2. But magnesium is a typical metal in that it tends to lose electrons to form a cation that then forms solids by ionic bonding. The octet rule predicts that the resulting cation will be stablest if its outermost shell contains eight electrons. And indeed the magnesium ion is $\mathrm{Mg}^{2+}$, two electrons less than the atom, with an electron configuration of 2 - 8.
What would happen if you had asked about a non-metal atom in the third period, say Cl? This has an electron configuration of 2 - 8 - 7, and tends to form a single covalent bond (e.g., to form $\mathrm{Cl}_2$ or $\mathrm{HCl}$), giving a configuration of 2 - 8 - 8 rather than 2 - 8 - 18. The reason for this is that the first eight valence electrons occupy the $s$ (2 electrons) and $p$ (6 electrons) subshells, while any further electrons in the third shell must go into the $d$ subshell. These orbitals are both higher in energy than those from the first two subshells (which is why, as you go through the periodic table, the $4s$ shell is filled before the $3d$) and extend less far from the nucleus, so that they are much less likely to overlap other atoms' orbitals to form a covalent bond. So it is often energetically preferable not to put electrons into this subshell if possible, which is why the octet rule can still apply to elements in this period.
Of course, it is possible to put electrons into this subshell, which is why it is possible for the central atom to have more than an octet in compounds like $\mathrm{SF_6}$. This leads into the second part of your question, which, as Manishearth suggested, looks like a duplicate of Can an atom have more than 8 valence electrons? If not, why is 8 the limit?.


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Why is a squared cosine window function commonly used for 2D spectra?
In nearly all default parameter sets for 2D NMR experiments on Bruker spectrometers, the window function is set to a pure squared cosine function for both dimensions (QSINE with SSB = 2). In contrast to that, 1D spectra are often processed with the efp command, which uses an exponential window function.
Why is the squared cosine window function the default? Which advantages and disadvantages does it have compared to other common window functions?

please let me know if the post did not fully answer your question. Otherwise, please mark the question as answered.

No problem, I hadn't heard anything and wasn't sure if the post had left something out or not.

The difference in window function is a result of the parameter being optimized for the respective experiment.  Exponential multiplication of the Free-Induction Decay (FID) weights the beginning of the FID where the amplitude is at it's greatest.  Weighting the beginning of the FID results increases the signal-to-noise ratio, since the noise should have the least impact with the strong signal.
For 2D NMR, we are usually more interested in resolution to aid in differentiating the various peaks and cross-peaks.  We can always take more scans to try and increase the signal strength, so resolution enhancement becomes the key objective.  A pure sine function offers increased resolution at the loss of signal enhancement.  The shifted sine function (or pure cosine function with SSB=2), provides a good mixture of the two objectives.  The enhancement doesn't weight the early parts of the FID as heavily as exponential weighting, but more than a sine function.
The use of SSB=2 for a cosine function is used most often in 2D experiments where we need a usually more significant amount of signal enhancement, such as TOCSY, NOESY etc.  sine enhancement is used most often for COSY, to reduce filtering out peaks from overlap.
--UPDATE--
To answer some further questions posted in the comments.  Using a gaussian weighting for 2D-NMR data is not as preferable as cosine or sine weighting.  The gaussian weighting has a much more narrow distribution, and rapidly approaches zero.  This can prematurely truncate the FID (by essentially weighting the later portion as zero), which introduces artificats into the resulting spectrum known as "truncation wiggles".
Squaring the cosine (which remember is just the sine function shifted) is done both to slow the increase and decrease of the weighting function over the course of the FID.  At the beginning of the FID, this allows for a greater amount of sensitivity to the early parts, while sacrificing a small amount of resolution.  Additionally, the slow decay also reduces the truncation artifacts discussed in the first question of the update.

Why not use a Gaussian function, since they are often used in 1D spectra? Why is the cosine function squared?

@Canageek I've updated the answer to address your questions.

I have a question for you: http://chemistry.stackexchange.com/questions/842/why-is-the-guassian-window-function-used-to-enhance-resolution-in-1d-nmr-spectra


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Why is the Gaussian Window Function used to enhance resolution in 1D NMR spectra?
In a recent question it was explained that Gaussian window functions are not used in 2D NMR spectra as they can introduce artifacts into the spectrum baseline by truncating the FID. Why then are they used with 1D spectra instead of sin or cosine functions?

Under the conditions you measure a plain 1D proton spectrum, you usually never have any roblems with truncation of the FID. You have a relatively long relaxation delay, you don't gain anything by making your acquisition time much shorter than it can be. So a default setup might measure 64k or even 128k points, which is far, far larger than the values you use in 2D NMR.
In two-dimensional spectra there are several aspects that severely limit the number of points you measure:

Measuring more points in the indirect dimension takes time, if you want twice the points you have to measure twice as long. For that reason you only measure as much points in the indirect dimension as you have to, no more.
Processing can become a bit slow and take quite a bit of space when you increase the number of points in the direct dimension in a 2D spectrum. It's not a big problem with current hardware, but with a 3D (or higher) spectrum this gets much, much worse.
If you need to decouple during acquisition time, the acquisition time is severely limited by the probe hardware and heating effects of the decoupling.

So a typical 1D spectrum will have a lot more points measured, beyond the point where the FID has decayed below the noise. In such a case you won't get any truncation artifacts, as the FID was never truncated to begin with. You only need to worry about truncation artifacts if you actually truncate the FID.
In the following image you can see

lila: QSINE (SSB 4) 10k points in FID
green: GM (LB -1, GB 0.5) 10k points in FID
red:  QSINE (SSB 4) 2k points in FID
blue: GM (LB -1, GB 0.5) 2k points in FID

You can clearly see the truncation artifacts in the gaussian-processed spectra with 2k points, but already at 10k points they are not visible anymore. The quadratic cosine of the truncated FID doesn't look all that nice, but it suppresses the truncation artifacts much better than the gaussian window function.
The SSB 4 I used for this example is also a far more aggressive setting than I'd usually use. I chose it because it gives a relatively similar resolution as the gaussian I compared it to. Usually I'd use SSB 2 for most 2D spectra, which emphasizes signal/noise over resolution.

Now I want to add a question asking what the heck the QSINE window function does. I'm quite annoyed that none of the spectroscopy textbooks I've read go into this, when it is one of the easiest parameters to fiddle with.

@Canageek The Bruker Processing manual contains the exact formula of that one

Yeah, but I'm not good enough to tell what it does or why from the math.


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What role does H2O2 have in copper (II) acetate formation?
Mixing solid copper, 5% vinegar, and hydrogen peroxide ($\ce{H2O2}$) causes copper acetate to form. The process will occur very slowly without hydrogen peroxide. Adding $\ce{H2O2}$ speeds the formation of copper acetate.
How does Hydrogen Peroxide promote the formation of Copper (II) Acetate?

The redox potentials $E$ for $\mathrm{pH} = 0$ show that $\ce{H+}$ cannot oxidize $\ce{Cu}$ to $\ce{Cu^2+}$:
$$\begin{alignat}{2}
\ce{Cu^2+ + 2e- \;&<=> Cu}\quad &&E^\circ = +0.340\ \mathrm{V}\\
\ce{2H+ + 2e- \;&<=> H2}\quad &&E^\circ = +0.000\ \mathrm{V}
\end{alignat}$$
Thus, non-oxidizing acids such as acetic acid cannot directly oxidize copper.
However, $\ce{Cu}$ can be oxidized by $\ce{O2}$:
$$\ce{O2 + 4H+ + 4e- <=> 2H2O}\quad E^\circ = +1.229\ \mathrm{V}$$
Therefore, copper is slowly oxidized in acetic acid in contact with air.
The oxidation can be increased by addition of oxidizing agents such as hydrogen peroxide:
$$\ce{H2O2 + 2H+ + 2e- <=> 2H2O}\quad E^\circ = +1.763\ \mathrm{V}$$

Presumably hydrogen peroxide causes copper metal to donate 2 electrons to form $\ce{Cu(II)}$ -- which is soluble in water -- and we'd say the hydrogen peroxide is acting as an oxidizing agent.

authors: غسان باسل مظهر


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Diamond filtering - is it real or bogus?
I always hear that some big corporations use "diamond dust" to filter water seven times to make "diamond clear" vodka. I tried googling something about using diamonds for filtering, but I could not find anything except corporate sites that tell us only the best things about "diamond dust filtration technology".
Is this just a bogus claim? Because I can not think about diamonds as a viable filtering material.

It appears that there is some validity to this technology.  Not surprisingly, the published journal literature available is basically non-existent, except for a handful of patent applications describing something along those lines.
The title of one journal publication "Preparation of a filtering material using nanomaterials for removal of a heavy metals from wastewater" uses a diamond nanopowder to aid in filtering out heavy metals from waste water.  The citation is "Antonenko, L. P., Energotekhnologii i Resursosberezhenie, 6, 2011, 49-54"
Additionally, there is a patent appplication from France (where at least one article I read cited as the technology originating from), title "Apparatus for activation and mineralization of water for domestic use".  In this, they use a variety of powdered gemstones (diamond, ruby, sapphire) as well as precious metals (gold, silver) to filter the water.
Certainly, nano-particle scale materials are effective filtration agents, as most contaminants will not pass through such a small sized filtration device.  The use of some gemstones (certainly diamond) are beneficial for their relative chemical inertness, and certainly silver is well-known for it's antimicrobial activity, and synthesis of gold-nanoparticles or gold-silver coatings are also actively under research

so for now i can claim that in laymans terms it's somewhere between bogus and real thing. Thanks for the great answer!

@Gabriel I would say that the "terminology" they use is maybe trumped up or exaggerated, but I would certainly say the actual technology seems valid.
That is assuming we are accurately paring their phrase "diamond dust" with the appropriate technology we've located above.3

Altered properties (band gap, Stokes shift) of cristalline nanoparticles as compared to bulk have been demonstated for various semiconductors, namely CdS, CdSe, TiO2, etc. They result from a significantly different surface-to-volume ratio and the fact that atoms at the crystal surfaces simply lack the proper number of neighbours.
Where the crystal ends, the is just no other neighbour. Since there's less than normal neighbours, centres at the surface are more reactive.
As far as the properties of the diamond dust are concerned, i do however remain skeptic and to me, the whole thing mostly sounds like marketing lingo.
The term diamond dust suggests something valuable: if the manufacturer uses something as precious as diamonds for filtering, the product must be valuable too.
But why would one actually want to use something with a rather inert diamond lattice of sp3-carbon atoms instead of microporous activated carbon with well-known adsorption properties?

Yes, diamond filters are used. Usually in the filtration of ethanol and water (vodka). All drinkable ethanol must be filtered through specialized carbon filters. So in the thinking that diamonds are formed from compressed carbon, therefore a diamond filter is the extreme sports of filtration. Whether it's better? I don't know, but it is a great marketing tool.

Could you elaborate? There's nothing here that isn't in the other answers, unfortunately.