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Energy-efficiency aside, what are the chemical constraints on CO₂ capture and methanation?
Synthesising $\ce {CH4}$ from air and water (in a non-biological process) has been proposed as one form of energy storage. What are the chemical constraints at play here? That is to say, what sort of catalysts might be used, and what are the performance / lifetime issues with them?
More specifically:
$\ce {CO2}$ capture and recycling (CCR) has been suggested as one form of round-trip electricity storage, with a round trip being:

electricity
$\to$ something that can hang around for hours to months
$\to$ electricity some time later

The proposal goes something like this: capture the $\ce {CO2}$ produced from the combustion of $\ce {CH4}$ in air in a gas-turbine electricity plant, and store it. Then, when there's surplus electricity, use it to synthesise $\ce {CH4}$ from the $\ce {CO2}$ plus water (I'm guessing electrolysis of water and a Fischer-Trop / Sabatier type reaction). And next time there's a deficit of electricity, burn the $\ce {CH4}$ to generate electricity, and capture the $\ce {CO2}$, and go around again. So, for the sake of this question, I'm specifically interested in the non-biological synthesis of methane.
And let's put aside for one moment the issue of energy efficiency, as that's effectively an economics question. (The question being "is the round-trip efficiency less than the ratio of off-peak to peak electricity prices?") Let's just agree that round-trip energy efficiency is one of the factors that influences the economics, and leave it at that for now.
So the question is: what sort of catalysts might be used, and what are the performance / lifetime issues with them?

I added some tags I feel are appropriate. I use fixation rather than energy-storage because I feel it better encapsulates both the sequestration and energy storage aspects of the question.

I removed the markup from the title, as I think we're moving away from doing that due to how it shows up on searches, please see this Meta post for more details.

While it is interesting to pursue better chemical ways to achieve this, we could just grow plants or algae which are already pretty good at the photochemical production of useful products from carbon dioxide.

I'm only directly attacking your direct question, "what sort of catalysts might be used, and what are the performance / lifetime issues with them?" because the rest is ill-defined. I'll mention some electrochemical conversion aspects. I hope to illustrate why your question is problematic.
There are several different products one could imagine reducing carbon dioxide to:

methane
methanol
formaldehyde
carbon monoxide
formic acid
ethylene

Each will require different materials and conditions, but they all require energy for this to occur. Not to mention, these are all difficult reactions. Thermodynamically speaking carbon dioxide is at the bottom of this energy well and to get out of it requires spending energy.
One approach involves directly electrolyzing on a metal. Cu in aqueous solutions can give methane as a major product, but it also evolves hydrogen and many of the others I listed above. There are a couple of problems with this but the major problem is that it often requires large overpotentials at rates of interest for industrial purposes and this means more cost in terms of electrical input. Another major problem surrounding nearly all of these studies is the selectivity of a catalyst: in water at pH 7 evolving dihydrogen is thermodynamically more favorable than reducing carbon dioxide to everything except methane and methanol. Unfortunately these two are the most complex, as far as chemical transformations are concerned, to make. So the other products are much more likely to be produced along the way. This is why people have focused on making simpler things such as carbon monoxide or formic acid.
Where does this energy come from? Right now coal. So we would end up creating more carbon dioxide than we reduced to other species. It reminds me of how diesel-electric transmission works in freight trains: a dirty engine turns an electric generator that generates electricity that is used on motors instead of the motor directly driving the wheels; I think we can call this diesel-electric transmission an indirect drive whereas the latter is a direct drive. It turns out that making the former is far easier in engineering rather than the latter. The same is true here. It's easier to generate electricity somewhere else then reduce CO2 rather than have energy directly fed to reduce carbon dioxide. You can get efficiencies better by going the latter, but engineering/technical complexities increase immensely.
As a means to address this electrical burden via the direct-drive analogy, people have looked at these artificial photosynthesis approaches, i.e. using light to  drive these reactions or at least help ease the energy burden. This requires photocatalysts and often with these come semiconductors. Some semiconductors are naturally good at reducing carbon dioxide, such as CdTe, but this better performance for generating methanol is in turn related to an actual corrosion process of the semiconductor. Semiconductors are very complicated and expensive to manufacture well. They tend to corrode easily when they work well with sunlight. They require catalysts and this introduces huge compatibility issues and manufacturing complexities. We'll just leave the indirect approach at the state  the field is, complicated and incomplete.
As for the direct approach:
it's complicated by the fact that carbon dioxide is fundamentally limited in solubility at room temperature. Hence why other people prefer non-water systems. These unfortunately add other complexities...
Actually while typing this I came to the conclusion that I could write a much more succinct response to you rather than continuing to babble on: The best candidate will involve a transition metal, there is no doubt about that. It may end up being a molecular species, but there are so many possible structures that it is hard to say if we won't find a simpler heterogenous catalyst that is more capable first. By Murphy's law, it will be the most expensive material imaginable. But right now, considering everything, you can always look at all extensive work that has been done on copper, but realize it isn't very efficient or selective.

Thanks for this Chris: if you could write a little about why a transition metal would be the best candidate, I'd appreciate it.

When you first look at the periodic table you can immediately notice the alkali metals on the left and the halogens to the right. Noble gases are normally nonreactive and are being ignored here. Halides as a group tend to be very good at accepting electrons and alkali metals tend to be very good at donating electrons. Reactions such as carbon dioxide reduction require, for a single site, many steps that can be donating or accepting. Transition metals, as a group, tend to have the ability to be very adjustable in their ability to donate or accept e-s and as such are highly useful in catalysis.

Your question was about constraints, but see mostly the ways this could happen:
Methanogenic archae can metabolize Hydrogen and Carbondioxide to Methane, the input of electricity would be needed for Hydrolysis.
The $H_2$ path of methanogenesis is highly relevant in agricultural biogas plants, so you can probably get the right sort of archae from there or from grass eating animals. For anaerobic digestion in Wastewater treatment plant, the path via acetic acid seems to be far more prelevant, so that's maybe a worse place to look for your methane building workforce. One text on the subject: http://www.zdnet.com/blog/green/h2o-co2-ch4-thanks-to-archaeans/3534
The biological way of storing electricity as Methane is investigated by these guys, I think: http://www.bioferm-energy.com/bf/de_de/company.html
However, I found zero info about how exactly the nearest commercial application ( http://www.solar-fuel.net/ have a pilot plant, the technology doens't handle load changes very well - pity, since that's the entire point) exactly works. I have to assume they use some Sabatier process.

Thanks mart: the information about archaea is interesting as a side-issue, but I'm focussing here on non-biological synthesis. I've updated the question to emphasise that.

Yes. It seems to me, that the producers (at least the ones I found, all in germany) are very tight-lipped about the specifiy process and catalysts they use.

Thanks for the answer, @mart.  For future readers of the question, it would be helpful for you to adjust it to @EnergyNumbers current changes, though.

One of the earliest reports on the photochemical reduction of carbon dioxide was published about 30 years ago by Jean-Marie Lehn and Raymond Ziessel:
Photochemical generation of carbon monoxide and hydrogen by reduction of carbon dioxide and water under visible light irradiation
Proc. Natl. Acad. Sci., 1982, 79, 701-704
They used Ru(bipy)2Cl2 (the ruthenium complex absorbs the visible light) and CoCl2 as the catalytic system.


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How does Polytetrafluoroethylene (Teflon) differ from Polyvinylidene fluoride (PVDF) as a protein binding membrane material?
PTFE and PVDF (durapore) are both used in protein binding filter membranes (Millipore specifically). Chemically speaking the two polymers differ quite significantly due to the additional fluorides and molecular weight. Their protein-binding properties also differ. I'm curious about how the chemical differences affect their properties as a membrane material.
From what I understand, the durability between the two is quite different. The PTFE membranes are used for the filtration of organic solvents where as the PVDF membranes are generally for low protein binding. Furthermore the PTFE membranes typically require a polyethylene support.
I don't believe that the degree of polymerization or the architecture; if we were dealing with a copolymer I expect differently. Based on my molecular intuition, I would imagine that this is completely based on the chemistry of repeat unit.
My question concerns a few things. What makes the PTFE a better solvent resistant polymer? The other is how does the reduced number of fluroides result in reduced protein binding?

Teflon/PTFE

PVDF

Alright. From the standpoint of polymer chemistry, your question is vague. We could compare the physical properties of the bulk polymers: glass transition temperature, melting point, shear/young's modulus, crystallinity, etc. Or we could (more to the point) compare two different membranes, that have been processed differently and made from different polymers with known physical properties. How big are the holes in your membrane? Is the PVDF very regioregular? I just want you to narrow your question.

I won't be able to answer this question, but I agree that it's not yet specific enough. When you say reliability, what properties are you concerned about? And simply identifying teflon and PVDF isn't quie enough; what is the degree of polymerization of the two samples? Polydispersity index? What is the context in which you are using the membrane?

Is protein binding really that complicated of a concept? For high-throughput screens, common assays would be to use these type of filterplates for protein binding. However, the hydrophobicities of two different yet similar molecules are drastically different. The hydrophobicity of polymers are independent of the PDI.

bobthejoe, I see that you've been trying very hard to make this answerable, and the two remaining comments are giving you some guidelines as to specifics that should be in the question.  I'm not sure how to help the two parties meet here, though.  As I had mentioned in chat to @CHM, I think that perhaps guiding your definitions for 'wetting' and 'durability' might help the polymer camp better meet your needs in terms of physical properties.  I understand your frustration, but keep at it a bit and I think you will get the answer you are looking for.

Essentially you are asking why about the different surface properties of PTFE vs. PVDF.  Generally PTFE is more 'non stick' than PVDF, which is more non-stick than polyethene (PE).  PTFE is much more chemically resistant to corrosion than PVDF or PE.  You are correct to look at the chemical structure.  The more Fluorine atoms on the backbone, the more proteins will not be able to chemically latch onto the surface.  Measuring things like contact angle will confirm this.  I don't know how much the pore structure of the membranes affect things, but from a chemical standpoint, there are clear differences.

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Flammability (NFPA) - how is it defined?
I was reading up on a wonderful little chemical compound known as chlorine trifluoride ($\ce{ClF3}$). For a primer, check out Dr. Derek Lowe's blog post here: Sand Won't Save You This Time. The title of the post is quite telling; this compound is so reactive with other compounds - any other compound, really - that it will set sand on fire, making this old standby for extinguishing lab fires useless and putting the chemical in very rare company indeed. Obviously, water is a huge no-no for chlorine trifluoride; it will exothermically react with the water to produce clouds of steam containing two nasty acids (hydrochloric and hydrofluoric) and a host of chlorate and fluoride compounds. Somewhere in the middle of all the heat, fluorine and organic-ish compounds, you could probably even get some fluorine oxides, which would just LOVE to rapidly reduce their oxidation state (an event normally accompanied by shrapnel).
I was intrigued; I went and looked up the MSDS for this chemical to see what the instructions were for dealing with industrial spills of this stuff (apparently, there's a decent market for it in the semiconductor industry, and probably also for plastics to produce highly fluorinated polymers like Teflon). "A good pair of running shoes" can't be the only answer to a $\ce{ClF3}$ fire. (well, actually, it pretty much is; get everyone in a one mile radius and two miles downwind the hell away from the area)
Now, here's the question. While this stuff has some serious "synergy" when it comes to being an oxidizing agent (it's a better oxidizer than pure oxygen, loads better than pure chlorine, better even than pure fluorine gas), the NFPA flammability rating of the chemical is 0. This, despite reports from those who originally studied it that the chemical is hypergolic with every known fuel substance, and most things you wouldn't consider fuel (like asbestos, sand, concrete, brick, earth, test engineers). The resulting reaction certainly looks like fire; intense heat, bright light, sparks, smoke, the works. So, if chlorine trifluoride doesn't "burn" by the NFPA's definition, what is the NFPA's definition of "burning"?

You have found one of biggest flaws in the  NFPA "Fire Diamond". The NFPA rates the flammability of the compound as a fuel. ClF3 is not a fuel. It is an oxidant. ClF3 is playing the role that oxygen normally does. It causes fuels (reductants) to burn by oxidizing them rapidly and exothermically. Oxidants are incapable of creating fire in absence of fuel regardless of temperature. Fires started by ClF3 continue to burn after being smothered because oxygen isn't necessary to continue the oxidation process. And since ClF3 is denser than air, it gets trapped under the sand. Many other powerful oxidants, like fluorine and perchloric acid, have low flammability ratings. Oxygen also has a flammability rating of zero. They do not burn; they cause other compounds to burn.
The scale for the flammability rating has to do with flash point, which is the temperature at which a fuel forms an ignitable mixture with air. Since ClF3 doesn't burn (isn't a fuel), and doesn't react with air (N2, O2, CO2, Ar, etc, which are all inert towards oxidation), it does not have a measurable flash point.
The danger posed by ClF3 is summed in the yellow (reactivity or stability) diamond and the white (physical and chemical hazards) diamond. Your msds rates ClF3 as a 3 in reactivity/stability, which is appropriate. Wikipedia's NPFA rating article, even mentions chlorine trifluoride:

3 -  Capable of detonation or explosive decomposition but requires a strong initiating source, must be heated under confinement before initiation, reacts explosively with water, or will detonate if severely shocked (e.g. ammonium nitrate, chlorine trifluoride).

The white triangle will have appropriate hazard information. For ClF3, it has the symbols OXY, which means it is a strong oxidizer, and W, which means it reacts with water.
I have not found the specific reaction that ClF3 undergoes with sand (mostly SiO2). I expect that it involves ${\ce{SiO2 -> SiF}}_6^{2-}$, except that this half reaction is not an oxidation.

You may want to mention that it's basically the same as oxygen qualitatively. Oxygen also has a fire rating 0, but it supports combustion (while not being combustible itself)

""The scale for the flammability rating has to do with flash point, which is the temperature at which a fuel forms an ignitable mixture with air."" Aha, but what about solids like wood or coal? What is their flash point? There is a definition in Wikipedia!

There's a lot of intermediates that could happen here. Garden-variety sand actually contains a lot of hydrated silica, providing water to produce HF. That would in turn react with the silicon dioxide (hydrated or otherwise) to produce a host of silicon fluorides, from silicon tetrafluoride to hexafluorosilicate anion, which would have plenty of hydrogen around to form hexafluorosilicic acid and probably some trace alkali metals to form salts. If the sand was treated as a flame retardant with MgOH or AlOH there's plenty of that to go around too.

Then, there's all the oxygen that's displaced when the fluorine barges in to the dioxide, which is a secondary threat only because there's fluorine around; it can reform water (always a good exothermic spectacle) or find something else more fuel-like in the area to oxidize.

And don't forget the chlorine; again, reactions with chlorine and any nearby alkali metals are pretty much guaranteed to be spectacular, and with the heat of the other reactions around, and the liberated oxygen, I wouldn't put it past this stuff to form some nice explosive perchlorate byproducts (they wouldn't last long, but...).

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Test presence of oxygen > 1%
I want to qualitatively test for the presence of oxygen within argon gas. I would like the test to be positive if approx $1\%$ or more is oxygen, but definitely negative if less than $0.1\%$. Anywhere in between is fine—A quantitative test is a bonus. I was thinking of heating iron filings in the gas to deduce it. I'm not sure what apparatus and setup  I would need since heating gas expands, I would like to use ultra low tech equipment. I'm thinking of using water or cooking oil displacement to hold the extra gas, but I don't know if argon would dissolve, or how much oxygen would dissolve. Preferably a nonflammable, certainly nontoxic liquids (or tests).

What is the idea of this analysis? Do You want to monitor a argon stream? Do You want to check the argon in a cylinder You purchased?

Yes I would like to test argon in a cylinder I purchased (sorry for delay responding)

This should do it
http://www.sciencebuddies.org/science-fair-projects/project_ideas/Weather_p004.shtml


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67 failures at extracting acetaminophen from Excedrin(R)
I have adapted the lab experiment described in this J. Chem. Ed. article for the organic chemistry lab course I teach. In the past year, the first year of implementation, I have had 67 failures (100% failure rate) on the extraction part of this experiment.
The experiment involves the isolation/separation of the active components in Excedrin(R): acetylsalicylic acid, acetaminophen, and caffeine.

The experiment is in two parts. Part 1 involves analysis by TLC and separation by column chromatography. This part works beautifully. Part 2 involves separation by successive extractions with increasingly strong bases. Acetylsalicylic acid was extracted with 10% aqueous Na2CO3. Then, acetaminophen was extracted with 3 M NaOH. Finally, caffeine was recovered by evaporation of the remaining organic layer.
In Part 2 We recovered lots of acetylsalicylic acid and caffeine, but no-one recovered any acetaminophen. Below is a list of things I tried in order to improve the results. None of them worked. I know from TLC analysis of the initial organic solution that acetaminophen was present. What happened to it?
Things I have tried:

Changing the concentration of both bases (up and down).
Using a weaker base for the first extraction (like NaHCO3 or Na3PO4).
Increasing the HCl concentration used to work up the second extraction.
No acetaminophen was extracted with the weak base, nor was any left
behind after the second extraction.

Where has my acetaminophen gone?

How much of the total quantity of each did you recover? Was it a 90% yield for ASA and Caffeine or something like 40%?

Since these were undergraduates doing extractions for the first time, the recoveries of ASA ranged from 50-90%. Recovery of caffeine was lower, like 20-50%, but that matches the results from the paper I linked. Acetaminophen was universally 0% recovery.

I've provided a partial answer below describing a ridiculously easy procedure. It can even be done at home (it's a classic drug addict's/poor man's extraction of opiates from OTC drugs)!

they note that you should use KOH and not NaOH.

@Chris They do note that. I will be trying that variation this coming semester and if it works, I will provide that as an answer to my own question.

100% failure rate suggestive for 0% of substance to be extracted. Could it be the manufacturer lie us and does not include that component? You should investigate it, who know what you will discover :)

@Ilan - note that we do TLC analysis before the extraction. The drug contained acetaminophen at that time.

Did you notice any change of colour (e.g. to brown) and/or the formation of gunk upon adding the 3M NaOH?
I'd assume one would have to boil it, but is the following scenario conceivable under your conditions:

alkaline hydrolysis of the acetaminophen (paracetamol) to p-aminophenol
aerial oxidation to the corresponding quinone imine
partial polymerisation of the latter

You have confirmed that acetaminophen was initially there. A a change in the constitution of Excedrine since the publication of the article can thus be excluded. Isn't a reaction (hydrolysis or whatever) under your extraction not the most likely explanation for the vanished acetaminophen?

This is the closest to what I figure is happening. When we switched to KOH, the results improved, but not nearly to what they should be. The potassium salt of acetaminophen is apparently more soluble in water, and thus the extraction is more efficient. The sodium salt comes out of solution as a precipitate before redissolving (presumably due to hydrolysis).

Here's what I suggest:

Mash up the pills in a mortar
Dissolve the powder in cold distilled water, use a beaker
Shake well
Put the beaker in a freezer and wait until right before the water freezes to remove it
Carefully decant the solution
Check what's left both in the water and the powder residue

I don't know exactly how much water is optimal for the separation, that you will have to figure out.
The powder should contain acetaminophen.

This is a neat technique. However, the only component in Excedrin that has appreciable water solubility is caffeine. Acetaminophen is less soluble, and ASA is basically insoluble in cold water.

Also, I was more hoping for a way that I could fix or salvage the original experiment.

@BenNorris yes, ASA and acetaminophen will both stay out of solution. How do you measure the presence of each component? FTIR?

TLC against standards. FTIR with 16 students competing for the same instrument in one hour is a little tricky.

@BenNorris I reckon you're stuck with a KBr disc spectrometer ;)

Actually, we have a pretty nice Nicolet Avatar with an ATR module. The limit isn't the speed of the technique, but the speed of an undergraduate student taking his or her second IR spectrum ever.

You could fix/salvage the experiment by first extracting the acetaminophen with DCM, and then perform your base extraction.. although it deviates a bit.
Here's an example lab: http://www.odinity.com/separation-of-the-components-of-an-analgesic-tablet/


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How does temperature affect the concentration of flavonoids in orange juice?
I am in high school, and I would like to do an experiment on this for a project.
I've done quite a bit of research into the different methods I could use (see 2.4.2 here) but I still cannot find enough information to formulate a hypothesis. It would be useful iff you could point me in the right direction. So, how does heat affect how much flavonoids are present in commercially available orange juice? (without pulp)

That's an nteresting idea, but you still have to narrow your research. Do you want to study the influence of weather patterns on the production of flavonoids by a particular variety of oranges? Or how the extraction temperature influences flavonoid concentration in the juice? Or how finished product storage influences degradation of flavonoids?

I want to test the effect of storage temperature on the degradation of flavonoids. My teacher recently bought an electronic water bath that I can use which accurately sets the temperature to +/- 0.1C. I was thinking I could heat up the orange juice for about 10-15 mins(in a test tube, ensuring the same time period for the temperatures I set) and test for the amount of flavonoids after that period of time using colorimetry.

This should be of interest: http://dx.doi.org/10.1016/j.chroma.2006.01.019

@Adi, determination of flavonoid content seems like it should be a pretty standard food chemistry technique. The method described in the paper you linked seems unusually complex. It would help to know what instruments or equipment you have access to. Based on CHM's comment, you should be able to do this more simply with access to an HPLC with a UV-Vis detector.

Trouble is I don't have access to HPLC. I'm on a budget of a high school student, so this is the best I can do. I have access to a spectrophotometer and an electronic water bath.

Orange juice is often subjected to temperature variation during processing, pasteurization, storage and refrigeration at 4$^{\circ}$C. According to some research, there appears to be a dramatic decrease in the concentration of certain flavonoids at storage temperatures especially above 20$^{\circ}$C.
A variety of flavonoids found in orange juice includes didymin, eriocitrin, hespertin, narirutin, poncirin, apigenin and neoeriocitrin.
Effective analytical methods used to identify and quantify individual flavonoids include gas chromatography mass spectroscopy (GCMS) and high-performance liquid chromatography (HPLC).
Whilst chromatography is the most effective method of identifying specific flavonoids, you may be interested to know that ultra-violet spectrophotometry has been used to measure total flavonoid content before & after treatment.


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Efficiency of various Ewald-type schemes for the summation of electrostatic interactions
I am looking into implementation of Ewald summation techniques for a number of point charges in periodic 3D space (i.e. a molecular simulation box with periodic boundary conditions). The “mainstream” Ewald-type techniques I know are:

original Ewald scheme
particle-mesh Ewald (PME)
smooth particle-mesh Ewald (SPME)

However, while I can find papers describing each of those, I have not been able to find a comprehensive article/white paper/technique note comparing their implementations and their CPU timings on a typical dense system (liquid, protein in solvent, whatever…). What resources do you know of on this topic?

Why not give it a go and just benchmark it? I know DLPOLY Classic implements Ewald and SPME. Run some of their test ensembles with both and see how it behaves and scales

I also unsuccessfully tried to look for such a paper in the past but below are some useful resources I've found. The only definitive answer I can give you is that traditional Ewald summation will be slower than a particle-mesh technique for any relevant system sizes (this is backed up in several of the references below).
Chapter 4.9 of the GROMACS user manual has some very nice discussion on their various implemenations of long-range electrostatics and appropriate references.
The LAMMPS documentation also has quite a few references and some bits of helpful discussion.
You can also try digging through the AMBER manual which uses a PME implementation but I haven't read any of this myself.
Also here you can find implementation details for the PPPM method in HOOMD (helpful if you're interested in GPGPU implementations).
I think the most relevant bit of discussion is probably this excerpt in the GROMACS manual:

The Particle-Particle Particle-Mesh methods of Hockney & Eastwood can
  also be applied in GROMACS for the treatment of long range
  electrostatic interactions [99]. Although the P3M method was the ﬁrst
  efﬁcient long-range electrostatics method for molecular simulation,
  the smooth PME (SPME) method has largely replaced P3M as the method of
  choice in atomistic simulations. One performance disadvantage of the
  original P3M method was that it required 3 3D-FFT back transforms to
  obtain the forces on the particles. But this is not required for P3M
  and the forces can be derived through analytical differentiation of
  the potential, as done in PME. The resulting method is termed P3M-AD.
  The only remaining difference between P3M-AD and PME is the
  optimization of the lattice Green inﬂuence function for error
  minimization that P3M uses. However, in 2012 it has been shown that
  the SPME inﬂuence function can be modiﬁed to obtain P3M [100]. This
  means that the advantage of error minimization in P3M-AD can be used
  at the same computational cost and with the same code as PME, just by
  adding a few lines to modify the inﬂuence function. However, at
  optimal parameter setting the effect of error minimization in P3M-AD
  is less than 10%. P3M-AD does show large accuracy gains with
  interlaced (also known as staggered) grids, but that is not supported
  in GROMACS (yet).

If you're really looking to dive in, you can look at the source code for these packages (I believe you need a license for AMBER) to see the actual implementations.


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Fundamental forces behind covalent bonding
I understand that covalent bonding is an equilibrium state between attractive and repulsive forces, but which one of fundamental forces actually causes atoms to attract each other?
Also, am I right to think that "repulsion occurs when atoms are too close together" comes from electrostatic interaction?

I understand that covalent bonding is an equilibrium state between attractive and repulsive forces, but which one of fundamental forces actually causes atoms to attract each other?

The role of Pauli Exclusion in bonding
It is an unfortunate accident of history that because chemistry has a very convenient and predictive set of approximations for understanding bonding, some of the details of why those bonds exist can become a bit hard to discern. It's not that they aren't there -- they most emphatically are! -- but you often have to dig a bit deeper to find them. They are found in physics, in particular in the concept of Pauli exclusion.
Chemistry as avoiding black holes
Let's take your attraction question first. What causes that? Well, in one sense that question is easy: it's electrostatic attraction, the interplay of pulls between positively charged nuclei and negatively charged electrons.
But even in saying that, something is wrong. Here's the question that points that out: If nothing else was involved except electrostatic attraction, what would be the most stable configuration of two or more atoms with a mix of positive and negative charges?
The answer to that is a bit surprising. If the charges are balanced, the only stable, non-decaying answer for conventional (classical) particles is always the same: "a very, very small black hole." Of course you could modify that a bit by assuming that the strong force is for some reason stable, in which case the answer becomes "a bigger atomic nucleus," one with no electrons around it.
Or maybe atoms as Get Fuzzy?
At this point some of you reading this should be thinking loudly "Now wait a minute! Electrons don't behave like point particles in atoms, because quantum uncertainty makes them 'fuzz out' as they get close to the nucleus." And that is exactly correct -- I'm fond of quoting that point myself in other contexts!
However, the issue here is a bit different, since even "fuzzed out" electrons provide a poor barrier for keeping other electrons away by electrostatic repulsion alone, precisely because their charge is so diffuse. The case of electrons that lack Pauli exclusion is nicely captured by Richard Feynman in his Lectures on Physics, in Volume III, Chapter 4, page 4-13, Figure 4-11 at the top of the page. The outcome Feynman describes is pretty boring, since atoms would remain simple, smoothly spherical, and about the same size as more and more protons and electrons get added in.
While Feynman does not get into how such atoms would interact, there's a problem there too. Because the electron charges would be so diffuse in comparison to the nuclei, the atoms would pose no real barrier to each other until the nuclei themselves begin to repel each other. The result would be a very dense material that would have more in common with neutronium than with conventional matter.
For now I'll just forge ahead with a more classical description, and capture the idea of the electron cloud simply by asserting that each electron is selfish and likes to capture as much "address space" (see below) as possible.
Charge-only is boring!
So, while you can finagle with funny configurations of charges that might prevent the inevitable for a while by pitting positive against positive and negative against negative, positively charged nuclei and negatively charged electrons with nothing much else in play will always wind up in the same bad spot: either as very puny black holes, or as tiny boring atoms that lack anything resembling chemistry.
A universe full of nothing but various sizes of black holes or simple homogenous neutronium is not very interesting!
Preventing the collapse
So, to understand atomic electrostatic attraction properly, you must start with the inverse issue: What in the world is keeping these things from simply collapsing down to zero size -- that is, where is the repulsion coming from?
And that is your next question:

Also, am I right to think that "repulsion occurs when atoms are too close together" comes from electrostatic interaction?

No; that is simply wrong. In the absence of "something else," the charges will wiggle about and radiate until any temporary barrier posed by identical charges simply becomes irrelevant... meaning that once again you will wind up with those puny black holes.
What keeps atoms, bonds, and molecules stable is always something else entirely, a "force" that is not traditionally thought of as being a force at all, even though it is unbelievably powerful and can prevent even two nearby opposite electrical charges from merging. The electrostatic force is enormously powerful at the tiny separation distances within atoms, so anything that can stop charged particles from merging is impressive!
The "repulsive force that is not a force" is the Pauli exclusion I mentioned earlier. A simple way to think of Pauli exclusion is that identical material particles (electrons, protons, and neutrons in particular) all insist on having completely unique "addresses" to tell them apart from other particles of the same type. For an electron this address includes: where the electron is located in space, how fast and in what direction it is moving (momentum), and one last item called spin, which can only have on of two values that are usually called "up" or "down."
You can force such material particles (called fermions) into nearby addresses, but with the exception of that up-down spin part of the address, doing so always increases the energy of at least one of the electrons. That required increase in energy is a nutshell is why material objects push back when you try to squeeze them. Squeezing them requires minutely reducing the available space of many of the electrons in the object, and those electrons respond by capturing the energy of the squeeze and using it to push right back at you.
Now, take that thought and bring it back to the question about where repulsion comes from when to atoms bond at a certain distance, but no closer. They are the same mechanism!
That is, two atoms can "touch" (move so close, but no closer) only because they both have a lot of electrons that require separate space, velocity, and spin addresses. Push them together and they start hissing like cats from two households who have suddenly been forced to share the same house. (If you own multiple cats, you'll know exactly what I mean by that.)
So, what happens is that the overall set of plus-and-minus forces of the two atoms is trying really hard to crush all of the charges down into a single very tiny black hole -- not into some stable state! It is only the hissing and spitting of the overcrowded and very unhappy electrons that keeps this event from happening.
Orbitals as juggling acts
But just how does that work?
It's sort of a juggling act, frankly. Electrons are allowed to "sort of" occupy many different spots, speeds, and spins (mnemonic $s^3$, and no, that is not standard, I'm just using it for convenience in this answer only) at the same time, due to quantum uncertainty. However, it's not necessary to get into that here beyond recognizing that every electron tries to occupy as much of its local $s^3$ address space as possible.
Juggling between spots and speeds requires energy. So, since only so much energy is available, this is the part of the juggling act that gives atoms size and shapes. When all the jockeying around wraps up, the lowest energy situations keep the electrons stationed in various ways around the nucleus, not quite touching each other. We call those special solutions to the crowding problem orbitals, and they are very convenient for understanding and estimating how atoms and molecules will combine.
Orbitals as specialized solutions
However, it's still a good idea to keep in mind that orbitals are not exactly fundamental concepts, but rather outcomes of the much deeper interplay of Pauli exclusion with the unique masses, charges,  and configurations of nuclei and electrons. So, if you toss in some weird electron-like particle such as a muon or positron, standard orbital models have to be modified significantly and applied only with great care. Standard orbitals can also get pretty weird just from having unusual geometries of fully conventional atomic nuclei, with the unusual dual hydrogen bonding found in boron hydrides such as diborane probably being the best example. Such bonding is odd if viewed in terms of conventional hydrogen bonds, but less so if viewed simply as the best possible "electron juggle" for these compact cases.
"Jake! The bond!"
Now on to the part that I find delightful, something that underlies the whole concept of chemical bonding.
Recall that it takes energy to squeeze electrons together in terms of the main two parts of their "addresses," the spots (locations) and speeds (momenta)? I also mentioned that spin is different in this way: the only energy cost for adding two electrons with different spin addresses is that of conventional electrostatic repulsion. That is, there is no "forcing them closer" Pauli exclusion cost like you get for locations and velocities.
Now you might think "but electrostatic repulsion is huge!", and you would be exactly correct. However, compared to the Pauli exclusion "non-force force" cost, the energy cost of this electrostatic repulsion is actually quite small -- so small that it can usually be ignored for small atoms. So when I say that Pauli exclusion is powerful, I mean it, since it even makes the enormous repulsion of two electrons stuck inside the same tiny sector of a single atom look so insignificant that you can usually ignore its impact!
But that's secondary, because the real point is this: When two atoms approach each other closely, the electrons start fighting fierce energy-escalation battles that keep both atoms from collapsing all the way down into a black hole. But there is one exception to that energetic infighting: spin! For spin and spin alone, it become possible to get significantly closer to that final point-like collapse that all the charges want to do.
Spin thus becomes a major "hole" -- the only such major hole -- in the ferocious armor of repulsion produced by Pauli exclusion. If you interpret atomic repulsion due to Pauli exclusion as the norm, then spin-pairing two electrons becomes another example of a "force that is not a force," or a pseudo force. In this case, however, the result is a net attraction. That is, spin-pairing allows two atoms (or an atom and an electron) to approach each other more closely that Pauli exclusion would otherwise permit. The result is a significant release of electrostatic attraction energy. That release of energy in turn creates a stable bond, since it cannot be broken unless that same energy is returned.
Sharing (and stealing) is cheaper
So, if two atoms (e.g. two hydrogen atoms) each have an outer orbital that contains only one electron, those two electrons can sort of look each other over and say, "you know, if you spin downwards and I spin upwards, we could both share this space for almost no energy cost at all!" And so they do, with a net release of energy, producing a covalent bond if the resulting spin-pair cancels out positive nuclear charges equally on both atoms.
However, in some cases the "attractive force" of spin-pairing is so overwhelming greater for one of the two atoms that it can pretty much fully overcome (!) the powerful electrostatic attraction of the other atom for its own electron. When that happens, the electron is simply ripped away from the other atom. We call that an ionic bond, and we act as it if it's no big deal. But it is truly an amazing thing, one that is possible only because the pseudo force of spin-pairing.
Bottom line: Pseudo forces are important!
My apologies for having given such a long answer, but you happened to ask a question that cannot be answered correctly without adding in some version of Pauli "repulsion" and spin-pair "attraction." For that matter, the size of an atom, the shape of its orbitals, and its ability to form bonds similarly all depend on pseudo forces.

I would have expected mention of Slater's treatment using the quantum virial theorem, and the important result that it is electrostatics that keeps atoms together.  Also, your answer seems to ignore two important observations 1) nuclear fusion is possible and 2) the Pauli principle leads to the vanishing of the wave function when electrons of the same quantum number share the same region of space, again cf. Slater determinants for antisymmetrized wave functions. (I do not deny the importance of the Pauli principle in specific solutions to $H \Psi=E \Psi$, but that is not the issue here.)

I don't like this explanation. To start off with the Pauli Exclusion principle puts the cart before the horse. The Pauli Exclusion principle is in the middle of quantum mechanics. You have to explain the standard model of particle physics to explain why the Pauli Exclusion principle keeps electrons and protons apart. To me it seems that the OP doesn't understand why Coulomb's law isn't a sufficient description of atoms. To make a poor analogy the OP's question is like asking how do you install a kitchen sink before the house is built. So (1) atom (2) ionic bonding (3) covalent bonding

To continue, the OP's situation reminds me of a line from the movie "A few Good Men" - You can't handle the truth.

I always thought I have a firm grasp of quantum theory, but this left me dazed and confused.

Martin, please be specific: What do you find baffling? If Pauli exclusion is an issue, I can assure you that it is simply not possible to understand chemical bonding in a deep, non-cartoonish way without also acquiring a clear understanding of when and how electrons resist being in the same state -- that is, of Pauli exclusion. I take great pains not to give the formulaic versions, because while vital for calculating specific results, formulas can also lend an illusion of understanding without providing the full conceptual framework for how to apply them. So I try to focus on concepts.

There are way too many issues I would have to address and I really don't think that the comment section is the right place to do that. Let's just say I don't understand the way you phrase it, I am certain there is some truth in it, I simply cannot find it.

@Martin, that's entirely fair; I have given a very different take that was aimed conceptually at folks who were interested but not specifically trained in quantum mechanics. You likely simply have an excellent but different set of underlying concepts for describing such phenomena.

My issue with this answer is that by starting with Pauli repulsion etc. you ignore that all covalent bonding characteristics should be exhibited in the simplest bound molecule - H2+. Thus any explanation based on Pauli repulsion - necessarily involving 2e- - can  only give detail on other forms of bonding, and not the core fundamentals of covalent bonding.

Pauli exclusion doesn't apply to one electron. It only applies to two or more electrons, since it is the electrons (only) that are excluding each other. Thus asserting that the single-electron bond in the $H^{2+}$ is the fundamental building block of all covalent bonding is like asserting that cars without steering wheels are the fundamental building blocks of how highway traffic works.

Clearly Pauli repulsion plays a major part in normal covalent bonds, but I would assert that the fundamentals of covalent bonding should be exemplified in $\ce{H2+}$ - There is nothing else going in in $\ce{H2+}$ and so the attractive interaction present must be due to the concept of covalent bonding. I think saying that $\ce{H2+}$ is a special case leads to the question why is it a special case? I'm not saying I think your answer is wrong, it was certainly an interesting viewpoint, but I disagree that covalent bonding only occurs with 2 e- or more.

Sure! You are right to point out the the case of $H_2^+$ bonding is both unique and for that reason especially interesting. It is to the best of my knowledge the one and only particle configuration that depends solely on electrostatic attraction, unmodified by the complex energy and geometric constraints (just look at orbital diagrams!) added by Pauli exclusion. So, it is a very interesting and well worth pointing out, even if it does not generalize to other situations... hmm, hold on: Do species such as $Li^{2±}$ exist, say maybe in deep space?? Those too would be examples of your bond...

@Martin-マーチン That might be because most physics textbooks on quantum mechanics end with the $\ce{H2}$ molecule.

@MaxW Gravitational , strong and even weak forces are mostly irrelevant at the scale of molecules. A particle physicist should know that to describe phenomena at some length scale you can average the behavior of the theory at shorter scales to get a simplified model at longer length scales. Another thing is that Pauli Exclusion principle does not come out of quantum mechanics postulates.

@Zythos, could you clarify what you mean by that last comment about Pauli exclusion not coming from quantum mechanical postulates? Once you accept the interchangeability of particles at the quantum level, symmetric and antisymmetric interpretations of the amplitudes pop up very easily and very quickly.

@TerryBollinger Yes, of course. Take a look at http://doi.wiley.com/10.1002/qua.10276

@TerryBollinger "Only if we reduce Pauli principle to forbid all symmetries but   totally (anti)symmetric permutations is then a consequence of quantum mechanics assumptions."

@Zythos, pardon the delay, that was a truly fascinating reference! I'm intrigued by the argument, but there is no way I can assess it on the fly. So, you've added to my queue of Must Reads. Meanwhile, more perhaps for others reading this, here is the best online example I've found of the simplicity of the symmetric-antisymmetric argument that I alluded to earlier: http://psi.phys.wits.ac.za/teaching/Connell/phys284/2005/lecture-03/lecture_03/node16.html

Argh, I see the full article is behind a paywall, and there's no preprint in arXiv or anywhere else. Only 26 citations? I'm a bit surprised, but then that happens more often if there's no easy online access.  Also, frankly, this was too narrow of a venue for an article that takes aim at this fundamental of an assumption of quantum mechanics. More quantum physicists and theorists need to see it...

Thinking about this like a physicist, there are four fundamental forces: the strong nuclear force, the weak nuclear force, the electromagnetic force, and gravity. The strong nuclear force holds the protons and neutrons in the nucleus together. The weak nuclear force causes beta decay. Those two might be considered chemistry. I don't, but some people do. Gravity is much too weak to have any effect on chemistry. So that leaves the electromagnetic force to control nearly all of chemistry.
On a simple conceptual level, that's all there is. The nuclei are both positively charged, so they repel each other. The electrons are negatively charged, so they are attracted to their respective nuclei. When the electron clouds get close enough to interact with both nuclei, then they begin to pull the nuclei together.
The deeper explanation requires quantum mechanics. When the atoms are separated, you can use the Schrödinger equation with the electric potential from the nucleus. That gives you the electron orbitals for an atom all by itself. When the two atoms get close together, you use the electric potential for both nuclei in the Schrödinger equation. The solution is then the molecular orbital rather than the atomic orbitals. Because the Schrödinger equation is impossible to solve exactly for a molecule, chemists need an approximation. The usual approximation is to build the molecular orbital out of the atomic orbitals by adding and subtracting the atomic orbitals. This is where the ideas of $sp$-, $sp^2$-, and $sp^3$-hybridizaton, and $\pi$- and $\sigma$-bonding come from.
For further information, most introductory college-level general chemistry texts should discus this. As an example, I pulled most of the above explanation from Zumdahl's Chemistry. In the 5th edition, this is in chapters 8 and 9 (the current edition appears to be the 7th). This is a much more important idea in organic chemistry, so those textbooks usually review it in the first one or two chapters. The organic chemistry book I have in front of me at the moment is McMurray's Organic Chemistry. This is discussed in chapter 1 of the 3rd edition of that book (current edition is the 8th).

The first paragraph answers the OP's first question nicely. To answer the second question I'd moved stuff in 2nd and 3rd paragraphs around to explain why electrons of an atom don't collapse into its nucleus. I'd then discussed ionic bonding and then gone into covalent bonding. To me this is a much nicer explanation than Terry's.

@MaxW, think carefully about what Colin's first paragraph really says. First it brings in forces that are not really relevant to question. Very few people reading that excellent question would seriously think the answer involves gravity, so why even bring it up? But worse, the paragraph ends by simply repeating the question as if it was an answer: "It's all electromagnetics." Well... no, it just is not. Even invoking Schrodinger's equation sidesteps the real issue of where the repulsive forces originate. So sorry you don't like Fermi exclusion... but that without it, it's all hand waving.

I understand that covalent bonding is an equilibrium state between attractive and repulsive forces, but which one of fundamental forces actually causes atoms to attract each other?

Thinking about this like a physicist, there are four fundamental forces: the strong nuclear force, the weak nuclear force, the electromagnetic force, and gravity. The strong nuclear force holds the protons and neutrons in the nucleus together. The weak nuclear force causes beta decay. ... Gravity is much too weak to have any effect on chemistry. So that leaves the electromagnetic force to control nearly all of chemistry. [From Colin's answer where he nailed this part of the answer.]
Side Note - There is a overlapping field of physics and chemistry called Nuclear Chemistry which involves working with radioisotopes which require special handling to avoid radioactive contamination. For the most part the effect of different isotopic masses of an element on the element's chemistry can be ignored. Nuclear properties, such as radioactivity and nuclear transmutation, are really more physics than chemistry.

Also, am I right to think that "repulsion occurs when atoms are too close
  together" comes from electrostatic interaction?

Not really. There is a significant twist here. Coulomb's Law predicts an attractive force between an oppositely charged bodies (whatever the bodies may be...), but offers no explanation as to why the two bodies don't combine. The explanation of why a $\ce{Na+}$ cation and a $\ce{Cl^-}$ anion don't collapse first requires an understanding of why a proton and electron don't combine to form a neutron. That explanation requires electrons to form orbitals around the nucleus of an atom. The first explanation of orbitals was the Bohr model which has a number of shortcomings and which has now been replaced by quantum mechanics.
Although quantum mechanics offers a far better analysis of the behavior of a hydrogen atom via the Schrödinger equation, there is a serious limitation. The Schrödinger equation can only be exactly solved for one electron system. (This is much like the three body problem with gravity.) So multielectron atoms are built-up by using combinations of one electron orbitals.
Now that we have an atom we can consider $\ce{Na+}$ cation to be a hard sphere with the electronic configuration of a neon atom which has a closed subshell and a $\ce{Cl^-}$ anion to be a hard sphere with the electronic configuration of a argon atom which also has a closed subshell. The two ions then form a ionic bond. The ionic bond between the two ions can sort be fudged in a solid by the use of the Madelung constant. However an ionic bond doesn't explain the behavior of organic molecules like methane, or even a permanganate anion.
So extending the concept of combining single electron orbitals to form multielectron atoms, we can combine multiple atomic orbitals of different atoms to form molecular orbitals. The basic notion is that when the atomic orbitals from two different atoms combine then a molecular bonding orbital with lower energy is formed as well as an molecular antibonding orbital of higher energy. If each atom only contributes one electron, then the net result is that the bond formation releases energy as heat. If both atoms contribute two electrons then both the bonding and the antibonding orbitals are occupied and there is no net bond strength and no heat is produced. (Combing multiple atomic orbitals on an atom is where the ideas of $sp$-, $sp^2$-, and $sp^3$-hybridizaton comes from.)
Now it seems that we're reached a point to cut-off drilling down into the problem. The gist is in that using a linear combination of atomic orbitals (LCAO) to form a molecular orbital the overlap depends on the distance between the atoms. There is a distance for which the orbitals have a maximum overlap, and hence the strongest bond strength.

Overlap is a terrible indicator for bond strength, bigger overlap does not necessarily conclude in stronger bonding, and it often does not.

I think this question is generally poorly answered in textbooks - It seems in most chemistry undergraduate books this is "glossed over". Explaining it in terms of MO theory is sort of great - the orbitals overlap, electron experiences both potentials, and so bonding orbital is lower in energy etc. But this is ignoring some of the subtlety in covalent bond which is present. Covalent bonding is also clearly a quantum phenomena - describing it in terms of Coulombic interactions, purely classically, cannot lead to any meaningful result.
If covalent bonding is a general principle, then it stands that the principles involved should be exemplified by the simplest bound species - $\ce{H2+}$ Considered from a classical viewpoint purely, the electron is not stable in the inter-atomic space (middle of the atoms). It is trivial to show the potential energy is lower at one of the nuclei, and so all matter should just condense - electron loses all kinetic energy by radiating, and so orbit collapses - this is precisely what drove the founders of quantum mechanics - the simple stability of matter.
Let's examine $\ce{H2+}$. At infinite separation the electron is in one of the potential wells, but due to the infinite separation, cannot tunnel from one potential well to the other potential well provided by the $\ce{H+}$ nuclei. When the distance is decreased to a finite distance, but still very long, the electron is able to tunnel from one potential well to the other. This means the volume of space the electron occupies is larger - the electron's wavefunction is more spread out. This corresponds to a smaller curvature, and thus a smaller Laplacian, leading to a lower expected kinetic energy term $<T>$. However what happens to the potential energy? If the electron is no longer in one potential well, but tunnels to the other potential well does the expected potential energy term, $<V>$ increase or decrease? It rises - the electron's expected distance from the nucleus increases, as wavefunction grows in the internuclear distance, facilitating the tunneling of the electron from one well to the other. You could also see it as:
1)If a wavefunction is highly concentrated, it is highly curved -> Large value of the Laplacian, and so high kinetic energy. By being concentrated around the nucleus it also has a large expectation value of $<1/r>$ and consequently a large $<V>$ term.
2) If wavefunction is diffuse and spread out, it has a small curvature, small Laplacian, and so a small $<T>$ term, but the expectation value $<1/r>$ is small, and so $<V>$ is small.
Note that this process of electron delocalisation brings about a degree of delocalisation of electron density into the internuclear region - IE the charge delocalisation (leading to a decrease in KE) funds the redistribution of electron density away from the nucleus into the internuclear space (resulting in an increase in potential energy)
As the internuclear distance becomes smaller, this trend continues - the smaller distance facilitates enhanced tunneling of the electron from one potential well to the other, and so a decrease in the kinetic energy, but a rise in the potential energy. Overall though, the total energy is negative, with a positive gradient, IE value of $(dE/dR)$ is positive, so as R decreased, E also decreases. However this trend slows - as internuclear distance becomes smaller, tunneling this facilitates is no longer much, and so kinetic energy term decreases more slowly. In the case of $\ce{H2+}$ the minimum of kinetic energy term occurs at around $2r_0$. At this distance we can presume the potential well's are so close that no increase in tunneling can occur, and instead by reducing bond length, the volume over which the electron can "wander" is reduced, and so kinetic energy starts to increase again.
At this stage of bonding, the classical picture starts to kick back in again - the potential energy term, until now has been positive, and so disfavours bonding, now becomes negative. In the previous "stage" up until the KE term's minimum, electron density moves away from nucleus to the internuclear region. This now allows for orbital contraction - the electron's wavefunction now shrinks, and becomes more concentrated around the nuclei. This leads to an decrease in potential energy, a rather sharp decrease, taking $<V>$ term from positive to quite negative. At the same time, the volume over which the electron wanders is reduced, and so the kinetic energy term $<T>$ increases, and becomes positive - IE level of delocalisation is reduced. However the potential energy term wins out. This can be seen from the virial equations:
$<V>+2<T> = 0$
Which applies at equilibrium bond length. This can be rearranged to give:
$<T>/-<V> = 1/2$
Thus at the equilibrium bond distance, the kinetic energy term is 1/2 magnitude of potential energy term, leading to a binding energy:
$ E = <T> + <V> = -<T> = 1/2 <V>$
For interest the amount of electron density donated to the internuclear space is 16%, and so the typical view that the bonding interaction is due to this is shown to be wrong. Also note that the potential energy term has a minima at a much shorter R distance than $r_0$ and is still steeply decreasing at the equilibrium bond distance. This shows that the common fallacy that an equilibrium bond distance is reached due to an equilibrium between nuclear-nuclear repulsion and nuclear-electron attraction is wrong.
Also shown to be wrong is the classical view of covalent bonding - covalent bonding is purely a quantum phenomena - it cannot be explained by classical electrostatics + mechanics. This can be seen simply from considering the potential well created by 2 $\ce{H+}$ nuclei - the maxima of such a system is found at the halfway bond point, and so a classical electron will simply decay it's orbit until on top of one of the nuclei. It is therefore found that the only stable classical arrangement of charges is a singularity. Consider for example a lattice of positive and negative charges. A single perturbation from perfect lattice geometry would see the entire lattice collapse - such as system could only be metastable. Thus the stability of matter owes itself to quantum physics.
While this picture is undoubtedly the most simple system, it also should, and has, exemplified the characteristics of covalent bonding in all systems. Interestingly for some systems - Notably $\ce{F2}$ - this breaks down, due the presence of compact lone pairs. This is the phenomena of charge shift bonding, and there are some excellent answers over at:What is charge shift bonding?
Sometimes you may also come across the idea of covalent bonding as being the resonance energy of two different spin states:
$\ce{A\uparrow B\downarrow <->A\downarrow B\uparrow}$
This is precisely the picture we have - the electrons tunnel from one potential well to the other, and we simply use their spin as "labels" so we can say this electron tunneled to this nuclei, and the other went the other way. Thus both have expanded the "volume" of their wavefunctions etc. as above.

Both the attraction and repulsion are the result of the electromagnetic interaction.  At long distances, two atoms attract each other because of induced dipole-induced dipole interactions. When they get close enough together, the  exchange interaction acts on the non-valence electrons of the atoms and forces them into higher energy states. This makes the atoms repel each other at short distances.

This is just wrong. In chemistry an atom doesn't have a dipole but rather molecules. Also Coulombic repulsion of the nuclei isn't the primary reason that keeps atoms in molecules some distance apart.

@MaxW: Atoms do have dipole moments when near other atoms. The other error has been corrected.

An Induced dipole-induced dipole interaction is called the London dispersion force https://en.wikipedia.org/wiki/London_dispersion_force

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


text:
Are there any molecules with delta bonds in their ground states?
When looking at excited states of molecules, $\delta$ bonds are relatively common, but I've never come across a molecule with a $\delta$ bond in its ground state.
Are there molecules with $\delta$ bonds in their ground states?
What about even higher angular momentum states ($\phi$ and $\gamma$ bonds)?

Re: phi/gamma bonds, Wikipedia says There is strong evidence to believe that no two elements in the periodic table can form a bond with greater order than 6 and links to http://onlinelibrary.wiley.com/doi/10.1002/anie.200603600/abstract. Guess it's relevant here. It may not make phi/gamma bonds impossible, but probably makes them  rarer.

Chromium(II) acetate has a Cr-Cr quadruple bond - and hence a $\delta$ bond - in its ground state. This sort of "paddlewheel" configuration, with four carboxylic acids coordinating a pair of metal ions, is reasonably common (see also copper(II) acetate, for instance), although the extent to which the metals can be considered covalently bonded varies.

The copper analog is another example

Metal-metal quadruple bonds have seemingly been reported for several cases, quite often by F. A. Cotton. A short search revealed a more recent report on a species with a Cr-Cr quintuple bond: Synthesis of a Stable Compound with Fivefold Bonding Between Two Chromium(I) Centers.
Admittedly, I didn't bother to read the full article, it's just not the kind of chemistry I'm interested in. However, YMMV and you might find it helpful.

I am a bit late to the party, but I just found this question. When I did research on the φ bond for What would follow in the series sigma, pi and delta bonds? I came across some predictions about possible φ bond in the ground state.
According to Gagliardi and Roos, there should be a φ bond in the ground state of the $\ce{U2}$ molecule. That being said, it is probably incredibly tough to synthesise this molecule.
Metal-metal multiple bonds have become fairly common as my earlier answer demonstrates.
It should be clarified  that a quadruple bond does not necessarily mean that a δ bond is involved. Often enough the combination of the dz² metal orbitals lead to the formation of a second σ bond first.
It should further be explained, that a δ bond can be present, but the overall bond order could be lower than four.
Laura Gagliardi, Björn O. Roos, Nature, 2005, 433, 848-851. Available free at Archive ouverte UNIGE.

I figure you could make $\ce{U2}$ molecules by heating the metal enough so that becomes gaseous. Either find a chamber that resists temperatures of around $\mathrm{4000\ K}$ to get a decently high vapour pressure, or more likely, do some kind of ablation (plasma, laser, etc) on solid uranium, though there would probably be contamination with larger metallic clusters.


text:
How does brewing time and brewing temperature affect the flavonoid concentration of green tea?
I've searched a lot for this, but I still cannot find any good explanations for how these two factors affect flavonoid concentrations.

This looks on topic here (not sure), but if you're interested in techniques for brewing tea/etc, you may like our site for [homebrew.se] :)

By "how", do you mean "what is the mechanism behind..." (in which case a few more details on the observed effects would be useful), or "what are the effects of ..." (in which case it looks offtopic, not too sure)?

@Manishearth Ah, but at its core, brewing tea is really an exercise in extraction techniques.

You might want to have a look at
J. Peterson et al., Major flavonoids in dry tea, Journal of Food Composition and Analysis, 2005, 18, 487--501, http://dx.doi.org/10.1016/j.jfca.2004.05.006
and the references therein, e.g.
Christiane Lakenbrink et al., Flavonoids and Other Polyphenols in Consumer Brews of Tea and Other Caffeinated Beverages, J. Agric. Food Chem., 2000, 48*, 2848--2852, http://dx.doi.org/10.1021/jf9908042
From the abstract of the latter:

[...] At brew times up to 2 min the composition of the brew solids was for each product practically independent of brew time, with flavonoids again comprising the major proportion (86−88%) of the total phenolics. [...]

Longer extraction times (> 10 min) or continuous (Soxhlet) extraction lead to partial decomposition of the flavonoids, most likely due to phenolic oxidation, cf.
Won Jo Cheong et al., Determination of Catechin Compounds in Korean Green Tea Infusions under Various Extraction Conditions by High Performance Liquid Chromatography, Bull. Korean. Chem. Soc., 2005, 26, 747--754, http://dx.doi.org/10.5012/bkcs.2005.26.5.747
I hope that helps a bit.


text:
How to name this cyclic compound with an O in it?

How to name this cyclic compound with an O in it?

I was going to name it cyclobutane, but not all of the atoms making up the ring are carbons, one is an oxygen. So, then I decided to name it butane, but that doesn't make sense because there is a ring.

This is $1,4$-epoxy butane.

Well, the molecule you draw is tetrahydrofuran, aka THF, but of course the only way you can guess that is if you already know the molecule. It's a very common organic solvent.
If you didn't know its common name, you could try to name it according to IUPAC nomenclature. It's an ether, and it's cyclic, so it can be named as oxacyclopentane. But IUPAC has specific names for cyclic ethers, which in this case is oxolane:

(from this organic chemistry book).

authors: ΔΗΜΗΤΡΗΣ ΤΡΑΚΑΔΑΣ


text:
Temperature dependence in absorption spectroscopy
What are some qualitative guidelines one should know regarding the temperature dependencies (of the gas and maybe the radiation, if that makes sense) for an absorption spectroscopy measurement?
How does, for example, a rise in temperature influence the absorption lines?

I am aware of one effect: doppler broadening. When the temperature increases, the speed distribution spreads out. That spreads out the distribution of the velocity component in the direction of your light source. The light the particles see from the light source then get red- or blue-shifted relative to what actually comes out of the source. So the particle population ends up absorbing a broader range of wavelengths than it did at a lower temperature. The effect is  $\Delta\lambda/\lambda_0 \propto \sqrt{T/m}$

See this related question on the shape of the broadened bands: http://chemistry.stackexchange.com/questions/324/when-simulating-spectral-line-broadening-which-convolution-is-preferred

Temperature will change the population of the different states of the molecule, particularly the rotational states (i.e. in temperatures in ranges that are easily reached in the lab/ some factor around 300 K).
This will affect the intensity of the respective absorption lines in rotational (or rot.-vib.) spectra and therefore the hull shape of the P and Q branches will change
I just found example applets: http://rkt.chem.ox.ac.uk/tutorials/rotation/rot_spectra.html


text:
Can I test for lead with household chemicals?
I picked up a metal fence loop cap today and took it home just because I liked the shape.

(^ not mine)
I cleaned it, but some of the paint is chipping off. I'm only vaguely thinking about lead, and I'm not going to chew on it, so I don't want to buy a test kit. Is there a way I could test for the presence of lead ions with common/ideally household chemicals? I'd be willing to wait until I'm at my school lab again if I needed HCl or something.

Since this doesn't answer your question about testing for lead, I'll leave a comment. The sale of lead-based paint is banned in the US and other countries. In the US, it has been banned since 1978, and fell out of common use before that. If you have reason to believe that your fence loop is newer than 1978 (and it probably is), and you found it in the US, then there is very likely any lead in the paint.

Lifting and thinking should be enough. Second thought is about softness of lead, making such a cap made from lead useless. BTW, what is the purpose of the eye?

General remark: get a bit of lead (metal or salt - whatever you look for) to have a positive control for your experiments.
More non-chemical tests for lead (metal):

lead is soft: you can cut/scrape it quite easily with a normal knife.
sheets of lead are easily bent, and they stay so.
For bigger pieces: Hammering on it would leave a rather big impression of the hammer. The sound would be dull.
it usually looks grayish, not silvery-metallic, but the cut will be silvery.

If it is about the paint:

what colour is it?
white lead: $\ce{2 PbCO3 \cdot Pb(OH)2}$ I'm more familiar with that in the context of paintings or painting of wooden things (window frames)
$\ce{PbSO4}$ also white. I think it has also been used for things like window frames.
side thought: lead paints for wood may have been popular due to the fact that the toxicity to microorganisms helps very much with the anti-fouling effect. Anti-corrosion paints for metal need different properties.
lead yellow: $\ce{PbCrO4}$ Wikipedia says that was the original pigment for school buses.
$\ce{2PbO \cdot PbO2}$(Mennige) red anti-corrosion pigment that was at least here in Germany the most important lead based priming paint. Only difficulty with recognizing by color: lots of non-toxic replacements are produced with similar colour to have the familiar look...
Very sensitive but but not very distinctive test$^*$ for $\ce{Pb^2+}$: the sulfide is black and very insoluble. Kitchen chemistry: put some egg on and even the surface of the metal should turn black. Possibly some heat is needed (I didn't test, though). Otherwise, use acidic $\ce{Na2S}$ solution from the school lab. If the $\ce{Na2S}$ solution doesn't give a black precipitate, there could in theory still be lead in the paint which is not accessible by the sulfide due to the binding material of the paint. In that case, however, the toxicity will be less of a problem, too (i.e. unless you dissolve the paint somehow, burn it, ...)
$^*$ OTOH it is highly improbable that there are substantial amounts of silver in the paint...
reductive borax bead should be gray, oxidative colourless.
Isolating Pb is quite easy, maybe you can do that in your school lab: precipitate $\ce{PbCl2}$ with cold HCl, filter. The dissolve the chloride in hot water. You can then precipitate PbS (black) or $\ce{PbCrO4}$ (yellow).  Looking at the crystals under the microscope can help further. Qualitative inorganic textbooks often have microscope photos of different precipitates.

Google translate says Trennungsgang is "separation process", so maybe filtration or separation? Thank you for your detailed answer.

If you’re testing for whether it’s entirely made of lead, you can do it with a simple density calculation rather than anything chemical – an easy typical way to do this is to immerse it in a known volume of water and see how much volume the object adds, then divide the weight by that. Lead is the one of the heaviest common metals, with a density of just over 11 g/cm³.
If on the other hand it’s just lead content you’re looking for, I think you can go for, coincidentally enough, HCl.
NB: I have not tested the following procedure(s).
Lead (II) chloride isn’t very soluble (0.9 g/100 ml at 20 degrees Celsius), so if you just add a small quantity of strong HCl to a small sample of the metal and wait a while (leave it overnight or longer if you can, the reaction’s supposedly quite slow with lead: cf this video of trying it quickly and nothing happening, even with strong acid and strong peroxide. You might get a white precipitate of lead chloride forming.
If nothing happens, try adding a little potassium iodide. Lead (II) iodide is over ten times less soluble than the chloride, so you might get a precipitate with that if you don’t with the chloride. Note that adding hydrogen peroxide in the initial step will speed up the reaction of the acid with the metal, but hydrogen peroxide will also decompose dangerously rapidly if you then add iodide later (see the ‘elephant toothpaste’ demonstration, for example). Getting boiling acid everywhere is not something I recommend.
You can also then do a flame test – adding methanol to a lead salt and igniting should give you a somewhat ambiguous blue-ish flame.

Wrt the flame test: as methanol alone burns with blueish colour, you at least need a positive control. Also, a spectroscope to look at the lines would be good, as there are more things that cause blue flames, and AFAIK the color with lead is not very strong. Also, make sure you take $\ce{PbCl2}$, the paint most likely will cause a orangeish colour as organic binding material burns.

All the $\ce{PbCl2}$ I've precipitated was white.

@cbeleites: Ah, yes, oops. Edit time.


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Differences between phenols and alcohols
I know phenols are more acidic as compared to alcohols, but are they considered different from alcohols?
Sure, you can study something as a subset, but are phenols considered a subset of alcohols, or are they considered as completely different from alcohols? My confusion is due to the fact that both of them contain a hydroxyl group.

Then what are the difference between above mentioned ones and carboxylic acid?

Nope. Alcohols consist of an -$\ce{OH}$ group bonded to a saturated carbon ($\mathrm{sp^3}$ hybridized, no multiple bonds).
IUPAC says:

alcohols
Compounds in which a hydroxy group, -$\ce{OH}$, is attached to a saturated carbon atom $\ce{R3COH}$. The term 'hydroxyl' refers to the radical species, $\ce{HO^.}$.

and

phenols
Compounds having one or more hydroxy groups attached to a benzene or other arene ring, e.g., 2-naphthol:

(source: iupac.org)

A phenol consists of an -$\ce{OH}$ bonded to an unsaturated $\mathrm{sp^2}$ carbon. Thus, it does not qualify as an alcohol. One can classify it as an enol, though.

Really, to me, the classification doesn't matter. Classifications are artificial, what is important is, how well the properties fit in the classification.
Many of the alcohol properties depend upon:

Its unsaturated nature: oxidation to ketone/aldehyde/acid
The weaker $\ce{R-O}$ bond and its ability to easily break and form an $\ce{R+}$ cation (this makes it a good participant in $\mathrm{S_N1}$ reactions)

Phenol can obviously not be oxidised at the $\ce{OH}$ to a ketone/acid (though one can do stuff to make it into a quinone). Phenylic carbocations are unstable, thus we don't get any $\mathrm{S_N1}$ reactions, and the $\ce{Ph-O}$ bond stays put.
On the other hand, most of the reactions of phenol depend upon its

Aromatic phenyl ring: All the EAS reactions
Weaker $\ce{O-H}$ bond (i.e., acidic nature): Reimer-Tiemann reaction, etc.

Thus phenols and alcohols don't have too many reactions in common. So, in this case, they have been classified in a sensible manner--if phenols were classified as alcohols, we would basically be clubbing two radically different classes of compounds under one umbrella.

Actually the question in my book was "Which of these cannot be considered as an alcohol", and phenol was an option.

@kartik It is not an alcohol by definition. The second part of my answer deals with why we cannot even consider it as one.

@KartikAnand, one must remember that what is taught in a course is not always necessarily agreed on in the field. The "Central Dogma" comes to mind.

(Re:last paragraph) I agree that classifications are somewhat artificial. But I disagree that phenols and alcohols have nearly no reactions in common. Both alcohols and phenols are capable of acting as weakly acidic species; when deprotonated, their conjugate bases are both strong bases and good nucleophiles; and they can undergo many reactions in common at oxygen, such as alkylation, silylation, acylation, and many others. This is a significant facet of the chemistry of alcohols and should not be overlooked.

@james_ash: (I converted your post to a comment, once you have 50 rep you can comment yourself). Hmm, I sort of disagree with your comments on their strengths as acids/bases/nucleophiles, but that is only because their relative strengths are quite different. I'll edit my post, though--you're absolutely right about the other reactions. Basically I want to say that their nature is different enough for them to be classified differently.


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Is carbon tetrachloride an electrolyte?
Is $\ce{CCl_4}$ an electrolyte all by itself, or does something need to be added to it so that it becomes an electrolyte?
Also, is it weak or strong?

No, $\ce{CCl4}$ is not an electrolye.  An electrolye is a material that causes ions (charged entities) to form in the solvent.  Carbon tetrachloride neither dissociates into ions nor induces ion formation in the solvent.
The reason is that the $\ce{C-Cl}$ bond is rather strong and won't break under
normal solution conditions.  Nor does the molecule have any net dipole moment.

$\ce{CCl4}$ is a non polar covalent compound. It consists of only molecules. Thus, it cannot conduct electricity and is a non-electrolyte.

I'm relatively new to this community, and am wondering WHY this answer has 2 "down" votes, with absolutely NO plausible explanation.


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What is a neutral atom?
I was told that an atom's atomic number is defined as follows:

The number of electrons or protons present in a neutral atom is called atomic number. It is represented by Z.

What does neutral mean here? Why isn't it just "..present in an atom..."?

"Neutral atom" is close to nonsense, at least superfluous. An atom is neutral by definition, if not, it is called an ion. Its another instance of white swan, green grass or black soot.

@Georg There are certainly swans, that are not white ;)

It's a bad definition of atomic number. The electron count is only equal to the atomic number in a neutral atom but the term is meaningless outside that context as there are better definitions of atomic number: just count the protons!

@Georg: So table salt ain't made of atoms?

Electrons and protons are charged particles. The electrons have negative charge, while protons have positive charge. A neutral atom is an atom where the charges of the electrons and the protons balance. Luckily, one electron has the same charge (with opposite sign) as a proton.
Example: Carbon has 6 protons. The neutral Carbon atom has 6 electrons. The atomic number is 6 since there are 6 protons.

Here, a "neutral atom" is simply an atom that has no charge.
See, an atom consists of protons, neutrons, and electrons. Protons are positively charged, electrons are negatively charged (with the same magnitude of charge per particle as a proton). Neutrons have no charge.
Now, in a "neutral atom", the number of protons must be equal to the number of electrons, otherwise it would not be neutral.
Basically, the definition is saying that "in a neutral atom, the atomic number is equal to both the number of protons , and the number of electrons, because, well, they both are the same."
Compare this with an ion

An ion is an atom or molecule in which the total number of electrons is not equal to the total number of protons, giving it a net positive or negative electrical charge

Here, the atomic number cannot be equal to both--because they are unequal. In this case, a better definition would be:

The number of protons present in an atom is it's atomic number. It is represented by $Z$.

I totally agree with George that "neutral atom" is close to nonsense. Indeed, IUPAC recommends the following definition for atom:

atom
Smallest particle still characterizing a chemical element. It
  consists of a nucleus of a positive charge ($Z$ is the proton number and
  $e$ the elementary charge) carrying almost all its mass (more than
  99.9%) and $Z$ electrons determining its size.

As one can see, for an atom the proton number, also known as atomic number, which is the number of protons in the atomic nucleus, is the same as the number of electrons. Consequently, any atom is (electrically) neutral by definition, as George mentioned in his comment.

Because the phrase "hydrogen-like atoms" for the series, H, He+, Li2+, Be3+, etc.  is often used, I don't think "neutral atom" is nonsense.

Because sometimes scientists use the term "atom" when it is not neutral.
This is most often in the phrase "hydrogen like atoms":

A hydrogen like atom consists of one nucleus of charge Ze and a single electron of charge -e.

See also:
Correlation in Helium-Like Atoms (2-electron atoms regardless of charge)
The Spectra of some Lithium-like and Sodium-like Atoms (3 and 11-electron atoms)
Quadruply excited beryllium-like atoms – a semiclassical model (4-electron atoms)
Fine-structure studies of the ground state of boron-like atoms (5-electron atoms)
Relativistic corrections to the ground state energies of the carbon-like atoms (6-electron atoms)
etc.

Neutral atom is an atom with no net charge on it like it is not Na+ but Na or Cl- but Cl.
The person has not referred to atom but neutral atom  probably because charged atoms/ions have extra electron (as Cl-) or lesser number of electrons (like Na+), violating the concept that atomic number is the number of protons/ electrons present in a neutral atom.

Atomic number is the number of protons present in an atom, neutral or not. No rule is violated.

Well according to the afore mentioned definition when it isor electron it does violate the concept.


text:
Difference between non-thermal and thermal ionized gases
Why or when does non-thermal plasma become thermal? How do we know before hand if the ion temperature will differ from the electron temperature, on what parameters does it depend - where is the quantitative gap between the two options?
(Has it to do with the heat release in ion-recombination reactions?)

Ah, I see :) If there is any way you can (I dunno if there is), could you make it apparent in the question what you are looking for (i.e, the chemistry-ness). You don't have to, but it may make it better.

It depends on the density of the plasma or the mean free path respectively since thermalization takes place via collisions between the different components of the plasma.

A thermal plasma has its electrons and ions at the same temperature.  A raw plasma has hot electrons and cold ions from the vast difference in masses when microwave excited.  One way to thermalize is to run the plasma through a Helmholtz pair of coils with a strong uniform magnetic field parallel to plasma flow.  The low mass electrons tightly spiral around the parallel lines of force (that, of course, are not there), colliding with the nuclei to thermal equilibrium.
(That can sound a lot better than it works.)


text:
Where can I obtain tables of P-V-T/compression factor data?
Some recent research I'm dealing with has required me to look for tables of experimental pressure/volume/temperature values (equivalently, a table of compressibility/compression factors in terms of any two of pressure, volume, or temperature) for various substances. Thus far, I've managed to obtain compression factor tables for a short list of substances from Perry's Chemical Engineers' Handbook.
I am wondering, however, as to where I should be looking for more extensive tables (e.g. tables with finer sampling), as well as tables for substances that were not included in that handbook. What books/websites/journal articles should I be looking at?
My attempts at searching on the Internet have not been very fruitful; I would appreciate knowing what keywords should I be using with the usual search engines to get better results.

This borders on being one of those infernal make-a-list questions, but in its current form, it seems OK to me :)

@Manish, I know, but I wouldn't be asking this question if it wasn't very troublesome to pull up references for this sort of thing... :( as I said in a comment to the answer, tabulating data like this is hard work.

Yep. I doubt there are many lists like this--so it's fine :)

One place to look for simple fluids is http://webbook.nist.gov/chemistry/fluid/.  But there are only a limited number of compounds treated.
Be warned that the data in the NIST Webbook is smoothed.  Collecting data of this sort is very difficult.  Many different sets of observations are combined and fitted to an appropriate equation.  What is at the URL given above is computed from that equation.
If you are looking for overall compression factors, this will do very nicely.  If you want five figure precision in compression factors, well, this won't do.
Another place to check is the  which is a primary source of raw data.

"Collecting data of this sort is very difficult." - yes, I've had to do a simplified version of that experiment myself, so I understand very well how difficult it was to obtain those data. Speaking of which... you speak of raw data. Those would be what I need (and somehow I think asking for five-figure precision here is a bit much to ask, so I didn't say so in the OP). Would you happen to have pointers on where I can obtain the raw data you speak of? (Corollary: I don't want the smoothed data, as I'll be doing comparisons with some equations of state.)

Try the NIST data first. It is most likely the most accurate data available. Individual experiments almost never totally agree for a large number of reasons, which is why raw data may not be a good thing to use. Raw data will have to be found in the literature. The "Journal of Physical and Chemical Reference Data"
(an American Institute of Physics publication) is a place to start.


text:
Calcium carbonate toxicity
I can't make heads or tails of the information on the net — although there seems to be plenty of it, it seems contradictory.
Supposedly, calcium carbonate is used in antacids—yet there are warnings about ingestion and contact with skin.
I need to handle it in pure powder form. I'm mainly concerned about skin contact, should it be avoided? (Am I getting confused between two or more forms of calcium carbonate?

Looking at the Material Safety Data Sheet (MSDS), most of the issues with it seem to stem from the physical consequences of it being in fine dry powder form. This means that if you get it on your skin, it can get into cracks and pores and easily dry out your skin to the point of cracking, and will cause irritation or worse if you get it in your eyes, for example. It's not toxic if you get it on you, though -- rubbing a lump of limestone against your skin isn't going to do anything except abrade it.
So, yes, if you're working with the powder, wear gloves. And safety glasses. And long sleeves. (See also: lab coat.)
If in doubt, follow the MSDS.

It's essentially chalk.  I'm not aware of any longitudinal studies of gymnasts' long term exposure to powdered calcium carbonate, but my sense is that the worst that can happen is your skin can get dry and irritated. If you have a part time job as a hand model, you will want to be especially careful to avoid skin contact.
If you're in a lab, of course, it goes without saying that you should take necessary safety precautions such as proper eyewear and skin protection as well as a lab coat.

So far the skin questions have been answered nicely. Here's the ingestion:

used in antacids—yet there are warnings about ingestion

half of the problem is the acid neutralizing capacity: your stomach may like to keep up its acidic pH.
The other half is of course how that neutralization happens:
$\ce{CaCO3 + 2 H+ \rightarrow Ca^2+ + H2O + CO2 \uparrow}$
You'll produce lots of carbondioxide gas in your stomach

Think of the base as being weakly caustic. This applies not only to antacids based on calcium carbonate but also those based on magnesium oxide, another popular antacid agent.
Calcium carbonate is a compound that stands between full-fledged alkalies and weak  metal-hydroxide bases, much like magnesium oxide discussed here. Such compounds are not very soluble in water as such, but can react in the presence of weak acids such as the fatty acids in your skin. Such reactions, which are possible with magnesia-based antacids as well, may lead to the quoted dryness and irritation.


text:
Why do aromatic hydrogen atoms have a higher chemical shift value compared to aliphatic hydrogen atoms?
In Nuclear magnetic resonance (H-NMR) spectroscopy, the chemical shift of aliphatic hydrogen atoms are very much closer to the 1.0 ppm than that of aromatic hydrogen atoms.
Aromatic hydrogen atoms have a chemical shift value of about 7.0-9.0 ppm whereas the chemical shift value of aliphatic hydrogen atoms ranges between 2.0-3.5 ppm.
Why is this so?

The chemical shift gives you information about how well shielded the nuclei are from the magnetic field. A proton at higher chemical shift values is deshielded, so the aromatic protons are obviously less shielded than aliphatic protons.
One effect that causes deshielding is the presense of electronegative atoms that draw electrons away from other atoms and thereby deshield them. But that is not the cause of the aromatic chemical shift.
The reason for that one is the ring current that is induced in a magnetic field. The induced current creates a local magnetic field that has the same direction as the external field outside the aromatic ring where the attached protons are, leading to the deshielding of the nuclei.

+1.. Does the delocalised electrons in the benzene ring has any effect on this?

They are the cause for the ring current


text:
Adulteration of turmeric with Metanil yellow
I read in the local paper that common food products like turmeric can be adulterated. It continued to specify the adulterant which they called metanil yellow. On searching on the Net, I was able to find that metanil yellow is prepared using some basic raw materials like metanilic acid and diphenylamine. The paper continued to lay out procedures on how to find out whether the turmeric you buy from the local market is adulterated. To do this, one must first dissolve half a spoon full of turmeric powder in 20ml of lukewarm water. It then told to add a few drops of hydrochloric acid. Upon doing this, if the aqueous solution turned pink, violet or purple, it indicated the presence of metanil yellow. It continued to state the risks of consuming turmeric adulterated with metanil yellow. One of these risks was that it is a carcinogen. I'm an Indian, so turmeric is included in nearly all of the dishes I eat in a day, so I became very curious. My question is this:

1.) What is the chemical formula for metanil yellow, how is it prepared (in layman's terms), and how does it react with hydrochloric acid to turn pink, purple or violet (a chemical equation would be helpful)?
2.)What are the risks in consuming metanil yellow?

If You look for "metanil yellow" in Google, You find ample information on formula, MSDSes and- an article from two Indian scientists on riscs from elongated consumption. All other things You ask for wre either silly (eg synthesis in laymans terms) or I start to think for what You need this information. In general: what is Your issue?

The thing with azo dyes like metanil yellow is that they are broken down within the body to the corresponding aromatic amines used to make them. Thus, any danger one might encounter from internal consumption of an azo dye would depend on the aryl amine... and yes, a good number of aryl amines are potentially carcinogenic.

As for why that test with hydrochloric acid works: metanil yellow is in fact an acid-base indicator.

Hmm, I see a few issues with your question: For one, it looks like you haven't properly researched the issue on your own, a Google search would answer the first part of your first question. Secondly, you have asked too many questions in one post--I suggest you restrict the post to either the $\ce{HCl}$ reaction or the toxicity. (Make two posts if you want).

I hope to keep this short.
Metanil yellow is an azo dye synthesized from the coupling of metanilic acid and diphenylamine, as described here. Here's a sketch:

Now, the thing with most azo dyes like metanil yellow is that when they are taken internally, liver enzymes or intestinal flora will usually reduce the azo dye to the components that were used in the diazo coupling. (See this article for more details.) For metanil yellow, the metanilic acid is poorly absorbed from the intestinal tract (and is thus easily excreted), and your remaining worry is the diphenylamine, which is a suspected mutagen and carcinogen (apart from the usual toxicity associated with arylamines). (In short, the metanil yellow itself is not your worry, but the metabolite diphenylamine sure is.)
By way of contrast, the natural coloring agent in turmeric is the phenolic compound curcumin:

As it turns out, both compounds are in fact acid-base indicators, though they have different transition ranges; for metanil yellow, the pH transition range is 1.2-2.3 (red to yellow), while for curcumin, the transition range is 7.8-9.2 (yellow to red-brown). (Data taken from the Handbook of Acid-Base Indicators.) I'm not quite sure where the purple/violet color would come from, but the key point is that natural turmeric would be reddish in an alkaline environment, while metanil yellow will be reddish in an acidic environment; however, vinegar (pH ~ 2.4) is not sufficiently acidic to display the color change for metanil yellow, and thus a stronger acid like hydrochloric acid is needed to display the color change.

It is possible the pink-purple color is a result of a reaction between metanil yellow and curcumin or another component of turmeric.