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Why water is not used in Electrochemistry
In electro-chemistry, we always use Metal and Metal ion electrolyte. My question is
1. why Water is not used as electrolyte. (because books only refer to Metal/Metal ion type electrods, they dont use water in it)

when magnesium electrods are made, Why Magnesium dont react with water or magnesium solution instead of showing redox reaction?

I am just using sodium just as example to understand the concept.

Your question seems to rely on some very specific premises. Electrochemistry in general is way broader than that. Water is very much used as electrolyte when it is suitable, or not used when it is not. Sodium would react with 1M $\ce{NaOH(aq)}$ just as violently as it does with water.

So This means sodium electrode is not possible? So how magnesium electrods are made. Dont they react with water Or magnesium solution, to release hydrogen, instead of showing redox reaction?

In electro-chemistry, we always use Metal and Metal ion electrolyte.

And, from a comment:

I am just using sodium just as example to understand the concept.

I concur with Ivan Neretin: you seem to be in possession of some bad information. Unless you're referring to some specific electrochemical process requiring the use of metallic sodium, in general water is used extensively as an electrolyte for electrochemistry. At my work, we use aqueous electrolytes almost exclusively in the course of our electrochemical process development.

So It means Sodium cant be used as Metal/ Metal ion electrod? which metals are generally used to make electrods ?

actually in some question, Book was referring to Magnesium electrode. So i wasnt able to grasp the concept (as Elements of 1 and 2 group are reactive with Water) . SO I just went little extreme in my example and took sodium.

@aman No, sodium won't work, it's too reactive with water. While magnesium metal does react with water, it's a pretty slow reaction. I could see it being feasible if it's used as a cathode and kept polarized. Gold, platinum, glassy carbon, and graphite are pretty common academic electrodes; steel (mild or stainless, depending), titanium, aluminum, tantalum, niobium, and brass/bronze are all materials I've heard of being used. Highly application-dependent.


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Validating a GC instrument / column
Are there any generic rules / heuristics to identify a good GC vs a bad one?
The context is that during manufacturing-troubleshooting I often get shown GCs of various fine chemicals and I often wonder whether a problem is truly process-related or merely just an artifact of a bad GC operator / column / instrument etc.
Are there any benchmarks or signs to look for that will give a hint that a particular GC wasn't very reliable etc?
For example I've added one such GC graph below. Are there things too look for? e.g. I suppose non flat baselines or clamped off peaks are obviously bad? Merging peaks too? What else?

I used to intern at a GC laboratory so my knowledge is very limited. Nevertheless, what data is available to you?

@QuantumMOCHACCINO I've added one such graph to my post. Typically this is available & the heights / Areas & respective percentages.

What you need is a standard mixture of compounds that you can store indefinitely and inject every day or every week or every 20 samples (some frequency that you find gives you confidence).  The chromatogram for this sample should have well-spaced peaks for your particular column and temperature gradient program, and it should look the same every single time.   Here's one possibility: http://www.sigmaaldrich.com/catalog/product/sial/94234?lang=en&region=US

I would first take a look at the background scan. Make a run with the GC with the same solvent and conditions previous to the run of the sample. Also, if you happen to possess the pure chemical in the sample, analyze it with the same GC instrument (under same conditions) and compare the retention times. You could also try different detectors. Bear in mind that any GC will have noise. So don't worry too much about small peaks.

@QuantumMOCHACCINO Thanks. Couple of points: These are solvent free GCs. i.e. Liquid products. Injected neat. Problem is, the small peaks are still on the Quality Control specs. e.g. Some specs. on impurity peaks are less than 0.5% by GC. Is the GC the wrong instrument for such accurate specs?

For any instrument that quantifies via integration of a peak (GC, NMR, etc.), sample concentration (peak height) is important especially with low signal/noise ratios. Obviously if the peak isn't in the sweet spot - so to speak - (where its area is significantly above the baseline, but not too big that it saturates the detector) the quantification may be unreliable. Also peak shape is a consideration in integration. If the peak is too broad, much of the area under it can be obscured by the baseline. On GC/MS or GC/FID in particular I used to have problems with peaks 'tailing', that is a broad lingering asymmetry after the peak max which would often overlap with species with greater retention times. No es bueno for quantification. Baking out the column, finding a better heating profile/method usually did the trick.
Optimal methods are going to vary widely based the instrument you use, the solvent, what components you have in your mixture, there relative amounts and concentration etc.. A lot of trial and error. As a general rule though typically when I had 2 components that were close in rt, but wouldn't resolve, I'd move to a method with a much slower ramp for the column temp, upping the carrier gas if possible on your equipment might help, too.

Thanks! Is there some method / heuristic you used for the "finding a better heating profile/method " bit?

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How to compare the oxidizing power of perhalate ions
In my book the oxidizing power in descending order of the following is given as: $\ce{BrO4-} > \ce{IO4-} > \ce{ClO4-}$
My doubt is regarding their order of oxidizing tendency.
I thought that chlorine has the highest electronegativity, so can help oxygen stabilize the $-1$ charge the best and be least reducing (electron donor) and most oxidizing.  Looks like that’s not the case.
Another possibility was that iodine was the largest ion there, it’s far bigger than oxygen, so maybe it produced some instability which made the negative charge unstable, or in other words it became more reducing. In that case it should be least oxidizing (which is also not the case unfortunately).
I couldn't just make out whats so special about $\ce{BrO4-}$ that makes it the most oxidizing ion
Although this question gives a bit of a good hint but doesn't answer my question completely.

I would love it if you can find a duplicate for me because anyways i am not getting an answer here @orthocresol

Sorry Shubham, @orthocresol is right -- his answer to the stability question also answers the oxidizing power question. Bromine is the least stable of the three in the high $+7$ oxidation state... so what is it likely to do to become more stable? Oxidize something!

The question does approach the matter from a different angle, though, and so it may be useful to have it open.

@brian okay i take ur point that it may have some common points regarding bro4-. But what about the order of the other 2 ions? The question doesnt only ask about bro4-

Shubham Good point. I recommend editing your question to add a link to @orthocresol's answer, and then describing why it doesn't fully answer your question.

@orthocresol This is not a dupe, because it is not immediately apparant from the question that these will go in the same direction. For a dupe to be a dupe we need the essentially same answers and the essentially identical question. This one approaches the problem from a different angle.

@orthocresol oh i am sorry... Actually i read the duplicate question and its answer partially in the beginning. Later i found it somewhat useful. Thanks

Between perchlorate and periodate I think it is a bit subtle. There are possibly effects arising from differences in the form of the oxyanion in solution (ClO4- vs H5IO6 according to Housecroft) but a possible crude approach would be to rationalise it based on the weak I-O (or I=O) bonds. Not 100% sure so I will refrain from writing an answer... Perbromate is most oxidising due to the d-block contraction as described in my other answer.

I thought you guys at stack exchange know everything.  This is the first time i am not getting an answer

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References for Transition State Theory and Quantum Transtion State Theory
Having asked recently this, I started to look for good references and notes on Transition State Theory and, also, modern versions of it including things like tunneling (which isn't taken into account in Eyring's equation).
Unfortunately, I haven't been able to find any satisfactory reference. I have downloaded a bunch of lecture notes and gained conceptual understanding of TST, but nowhere do I find a explanation with all the mathematical details.
What could you recommend for a thorough and detailed exposition of the TST?
Also, I would like to find stuff about how the theory has been improved over the years (the theory is mainly classical). I guess there are more accurate quantum mechanical versions of the theory.
Thanks a lot.

I am unaware, that there has been any development for an improvement of TST over the original publication by Eyring. The many flaws of his equation are being ignored to a certain extent and it still works. A short summary of the critical points can be found at Wikipedia or in this question. I think starting with the original publication is your only chance to understand it from a rigorous point of view.

@Martin-マーチン If my memory serves me correctly,  Eyring refered to the coordinate inaccurately  as a fourth translational freedom degree. I am unsure if would be simple to catch his idea without a very solid background. Glasstone's "A Treatise on Physical Chemistry" maybe is more friendly as an starting point.

There are many sources that you can check out, in particular the reviews of Truhlar:

Current Status of Transition-State Theory by Truhlar et al.

This article was used in my undergraduate course of chemical kinetics, so I think that it could be a good source. Also we used:

Variational transition-state theory by Truhlar and Garrett

Maybe they are a bit outdated, but they are good as start points. The abstract of this newer article make it seems to be useful for you

Transition state theory: Variational formulation, dynamical corrections, and error estimates by Vanden-Eijnden and Tal

To deeply understand the theory statistical physics is a must. A good source (that also deals with TST) is the Ralph H. Fowler's book.

Thank you! That last reference is exactly the type of thing I was looking for.  It has the details and it clearly shows, for instance, how the ergodic hypothesis is crucial for TST, which is never clear in the chemistry literature, as far as I've seen. Also Atkin's "Physical chemistry" seems to be great for all of this stuff.

Your welcome. You should search for articles on Google Scholar, Scopus or similar sites.  ;-)


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Does the Hydrogen/Chlorine reaction in light proceed instantaneously?
When a mixture of Hydrogen and Chlorine is exposed to short wavelength light it reacts explosively. Does the reaction happen simultaneously through the bulk of the mixture?

There is no such thing as instantaneous in the real world. But some things can be incredibly fast. The question is how fast and this requires some insight into the mechanism. A good answer will contain this information.

It depends on how much UV is used, and how it's distributed... from a point source, the flame front would spread as if from an automotive spark plug. In fact, lasers are being tested as a means of ignition in Otto-cycle engines because the flame front can be better controlled than from a spark.
If evenly spread over the surface, as in laser fusion attempts, then the wave-front would propagate from the exterior inwards. Theoretically, using fiber optics inside the container to distribute and delay the beam equally, the explosion could propagate as fast as desired, with the caveat that the delay would still be half the largest space between tips.
This means, for example, that a one-cubic-meter container could be made to explode in far less than 1/3*108 seconds, which would be the minimum if it were to be ignited at the edge. Though not truly simultaneous throughout the container, one can approach any finite speed, even, apparently, greater than light. This is similar to other phenomena that exceed C, such as the crossing point of scissors or the phase velocity of a wave.

No amount of UV will make it instantaneous. you need to discuss the mechanism and the limiting speed possible to make a good answer.

The limiting speed is the time light takes to transverse the volume of gas. So, for a litre let's say it is around 300 picoseconds


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Weird Wikipedia Section on Oxidizing Behavior of Nitric and Sulfuric Acids
It has been my understanding that when nitric, perchloric, or sulfuric acids act as oxidizing agents, they must exist as whole $\ce{HNO_3}$ or $\ce{HClO_4}$ or $\ce{H_2SO_4}$ molecules that are still protonated. For example, putting copper in dilute sulfuric acid does nothing, but using concentrated sulfuric acid does cause a redox reaction. This agrees with the fact that only concentrated forms of these acids can oxidize; only concentrated forms have undissociated/whole molecules. A dilute solution would contain mostly hydronium and the conjugate base of the strong acid, thus not be able to oxidize. However, the "General Properties" section of this Wikipedia article flat out contradict that.
It says,
However, even dilute nitric acid can oxidize copper to $\ce{Cu^{2+}}$ ions, with the nitrate ions acting as the effective oxidant. (Emphasis: nitrate ions).
Then below it says that concentration is a factor for sulfuric acid. Is something amiss here?
UPDATE: I now embrace what user @matt_black is saying, and by now
I'm not even surprised that something like this could happen. Chemical guidelines are just guidelines. However, I'm now even more curious as to this departure from the trend. If he or someone else would like to give an explanation of the mechanism of reactions involving oxidation (of copper, for instance) by $\ce{HNO_3}$ and $\ce{H_2SO_4}$ that accounts for this anomaly, that would be great.

If only the protonated forms are oxydants why all the perchlorate  salts are strong oxidants/potential explosives?  This chemistry does not add up

@Greg It explains some things, such as how sulfuric acid exhibits its oxidizing powers only when fairly concentrated, and perchloric acid beyond 70%.

A strong acid is more reactive when concentrated than when the same amount of that acid is diluted in water.  The acid is strong precisely because dissociation in water makes it give up free energy.

It is chemistry that is being weird not Wikipedia
The trouble with simple explanations of oxidising acids is that there is no simple explanation that covers all the cases. The specific reactions leading to oxidation are complicated and have a variety of different mechanisms.
So copper in sulphuric acid can be oxidised by a reaction that involves SO2 production. That has a completely different mechanism to the reaction with nitric acid and a different concentration dependence.
And, given the variety of reactions, the need to have protonated molecules of the acid is not a good generalisation: it depends on the specific reaction that is happening.
The oxidising reactions depend on the substrate being oxidised as well as the acid. Which mechanism will dominate isn't easy to determine without trying the reaction. So don't trust simple generalisations.

Which is why we sometimes describe something as acting as a Lewis acid - it is not the "proton donating" characteristic or the hydronium ion that is the mechanism of reaction.


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Why do ionisations from non-bonding orbitals result in little/no vibrational structure in a photoelectron spectrum?
I understand that ionisations from bonding and antibonding orbitals create vibrational structure from the Franck-Condon principle, resulting in a greater overlap integral between different vibrational levels, but surely if a non bonding orbital is ionised the equilibrium bond length will not change (much), resulting in a good overlap integral between the vibrational levels?
If someone could explain the principle to me in this context I would be very grateful.

The idea is that your molecule, prior to ionisation, is nearly always in a vibrational ground state. The intensity of the line is therefore governed by the overlap of the vibrational level of the ionised molecule with the ground-state vibrational level of the unionised molecule. For ionisation from a nonbonding orbital, this is largest for $v' = 0$ i.e. transition to vibrational ground state of ionised molecule; from then on the overlap (and intensity) just decreases

whereas for ionisation from bonding/antibonding the Franck-Condon factor is largest for $v' \neq 0$ - this means that your intensities go up to a certain point, and then only decrease


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How pH paper works
Please tell me how a pH paper works. How does it decide whether the solution is an acid or a base on the basis of ions formed and how it can determine its concentration. How does it indicate an acid or a base?

It's a piece of paper with one or a couple of these coated on it: https://en.wikipedia.org/wiki/PH_indicator
Indicators are dye molecules that change colour depending on the pH of a solution.

Perhaps you could provide a little more detail here about how this works? At the moment this is in danger of being flagged as low quality for being too brief.


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Reaction mechanism for 7-hydroxy-4-methylcoumarin synthesis with an iodine catalyst
7-Hydroxy-4-methylcoumarin can be synthesized from ethyl acetoacetate and resorcinol with iodine as a catalyst. What is the mechanism for this? From what I understand it follows the Pechmann reaction but I'm having trouble to understand how it works in this particular case.

@bon I'm on it ;)

There is a paper about this exact reaction, but in addition to the use of $\ce{I2}$ it mentions the use of microwaves. In the reaction, $\ce{I2}$ acts as a Lewis acid, catalyzing both the transesterificaiton and electrophilic aromatic substitution reactions. It can be seen acting as such in the first step of the mechanism:
$\hspace{2cm}$
Next, a proton transfer occurs:

Next, the resorcinol derivative can act as a nucleophile, and ethanol is ejected, completing the transesterification reaction:

Next, $\ce{I2}$ coordinates to the keto carbonyl, activating the carbon towards electrophilic addition. Because this resorcinol derivative is so activated by the $\ce{-OH}$ and $\ce{-OOCR}$ groups, a weak Lewis acid like $\ce{I2}$ is sufficient to catalyze the reaction:

$\hspace{1.5cm}$
$\hspace{1.8cm}$
$\hspace{1.6cm}$


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Propane Combustion in a Gas Grill with Irregular Flameouts
I'm troubleshooting a problem with a brand-new propane gas grill that is experiencing flameouts and making puffing/popping sounds.  I'm asking here because I want to understand the basic chemistry that would cause this problem.
The problem is demonstrated in this video.
I've already been through all the expected troubleshooting steps:

Replaced the gas regulator
Replaced the burner
Tried three different propane tanks

The grill was purchased about 3 weeks ago. The valve/regulator assembly was replaced last week, and the burner was replaced yesterday, so both parts are essentially new and unused.
It's hard to tell from the video if the puff causes the flameout or the flameout occurs first, followed by accumulation and detonation of a small amount of unburned propane.
I would like to understand the chemistry principles that govern the combustion of propane in a grill, and what design flaw(s) could cause this behavior.

Very important question about that video:  Did you intentionally adjust the color balance in any way?

No.  It's from an iPhone 4s camera and color was not adjusted.  On my photography-calibrated monitor it looks like a faithful representation of what I saw.

It seems to me that the flame is more green than it normally should be.

I would blame that on bad grill design, since you replaced critical parts and the problem persists. It seems like incomplete propane burning.  Unburned gas accumulates below the top lid and when a fresh air blow comes it ignites. The unburned gas might also be accumulating below the lid to a point when it finally "overflows", reaches the air around the grill and sputters. You should perhaps consider getting your money back to purchase another grill model.

Does this behave in the same manner whether the flame is at it's highest or lowest settings?

@airhuff Yes, although it is diminished at lower gas settings.  I gave up on this grill and returned it for a refund.  Then I bought a different brand, which has performed flawlessly, so I'm chalking it up to a design defect.

Good call getting the refund  ;)

It's hard to tell why it got this oscillatory regime, but it isn't unusual.

Did the grill have a side burner? If so, was it upstream from the main burner, and did it exhibit this flameout behavior too?

No side burner.  The grill has two burners in one physical unit and two control knobs.  The issue occurred with both burners lit, and also with only one lit.

The brilliance of the orange in flares suggests suggests incomplete combustion of a carbon residue.  This can happen with a poor oxygen mix, but the blue in the core of the flame indicates that mix is fine.
Also, on the base of the grill there is some charred residue.
Based on the color, and the residues I expect somewhere in the flow you have some carbon deposits that are occasionally getting dislodged, causing the sputtering.

This would be a possibility, but the grill is brand new (used 4 times) and the valve/regulator and burner have respectively 2 and 0 total usage counts.  There is no possibility of carbon deposits anywhere in the gas system.  I'll update the question.

This looks to me like there was a piece of something (plastic film, maybe?) loosely pinned inside the propane feed line, that periodically would flap over and block the flow of gas to the burner.
Reasoning:

The simultaneous disappearance of the flames across the entire width of the burner implies a total cutoff of upstream feed pressure.
The relatively sharp and short aspect of the noise accompanying each flameout supports the supposition of a mechanical blockage upstream, and the "PHWAFFPH" character of the sound implies a soft material (no metallic clank or ping) and is quite reminiscent of an empty plastic grocery bag catching a stiff wind.
The irregular and very long (fluid mechanically speaking) time interval between instances of the flameout is strongly suggestive of some sort of discrete interruption of the gas flow, rather than an oscillatory flow disruption downstream of a static constriction.
The reduced severity of the problem at lower gas flow rates is consistent with this hypothesis, since a lower gas flow rate would result in a lower pressure drop at the flap as it starts to occlude the flow, and thus less of a driving force for pushing the flap fully across the feed line.
Recovery of the gas flow after only a brief flameout is also consistent with a flexible material causing the blockage, since a stiff material would probably not be able to rebound after closing off the flow, but instead be held in place by the pressure coming from the gas cylinder.

Interesting deduction chain but (1) everything in the pipeline from cylinder to burner was replaced with no change, and (2) the same cylinder on a new grill of a different brand exhibited no problems. I still favor a design problem.

@JimGarrison Everything, including the length of line between the regulator and the burner?

Yes, absolutely everything.

@JimGarrison Well then. What were you thinking, going and buying a possessed grill?!

@JimGarrison What about the gas flow control valve/knob, was it replaced too?


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Classical bioisosteres
Classical bioisosteres are functional groups that satisfy with the Grimm's hydride displacement law and Langmuir's definition of isosteres.
My question is why chlorine can be replaced by trifluoromethyl or cyano group and the replacement is categorized in univalent atoms replacement?
Because as seen from the criteria of classical bioisosteres aforementioned, the replacement atoms do not have any hydrogen atom therefore I think Grimm's law is devoid. However both trifluoromethyl and cyano group do not have the same number of atom nor electron as well as electron arrangement as in chlorine.
So what explains that these functional groups are classical bioisosteres?

I think you are right. Grimm's law in its original version (1925) should not apply to the $\ce{-CF3}$ , $\ce{-CN}$ pair. This "law" is very restrictive and has only historical significance. To be honest I could not recall its statement and I have never come across it in the last 10 years or so that I work in the field. The modern definition of classical bioisosteres (as opposed to non-classical bioisosteres like the $\ce{-CO2H}$, tetrazole pair) has been extended exactly for this reason, to include functional group pairs with similar physiochemical properties that do not follow Grimm's law or any other classification of less general scope.


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Dipole and size of a thioether relative to hydrocarbon
I am an enzymologist and I would like to check whether the consensus of how enzymes tell thioethers from hydrocarbons on substrates is a fact or fiction.
So are any of these wrong?

S–π interactions are the sole discriminant (in enzymes at least)
They have the same dipole
They are the same size
A thioether and a hydrocarbon have the same angle

From what I can tell...
Point #1 is not really in question: even if it does not account for much, aromatic residues are often found with thiols and thioethers (1/3 of protein have a methione involved in a S–π interaction).
From what I can tell of point #2, 3.55 vs. 3.58 units in electronegativity (ref. wikipedia: electronegativity) won't make a difference. From point #3, while, yes, sulfide is 360 pm and methane is 380 pm (ref. wikipedia: kinetic diameter), but the former is R-S-R, R-CH2-R, so it might be different.
I am not saying they are very different: thialysine (thioether) is a toxic analog of lysine (alkane) and norleucine (alkane) of methionine (thioether; toxic for SAM usage only) due this reason. But nature has evolved to discern similar compounds.
An example, cystathionine (pictured below) is a thioether of cysteine + homoserine or homocysteine + serine (both routes exist) and one enzyme (cystathionine beta-lyase, CβL) can bind the amine on the shorter side, deprotonate the α-carbon and eliminate the thiol (homocysteine) and release the imine (hydrolysed to pyruvate), while a different enzyme (cystathione gamma-lyase, CγL) acts on the other (cysteine and ketobutyrate). The two enzymes do not catalyse both reacts as it would be wasteful (not much just a ATP>AMP). The evolution from one to the other has not be published, but from what I can tell S–π interaction is the driving force. AFAIK, the closest literature to this on cystathionine elimination is my own paper, so I know there is no answer there.

This is very interesting, but I can't find the question in there.

I meant to ask whether any of the four points were wrong. I put an example at the bottom to emphasize that they are different, so I must be wrong somewhere.

I just realised that I got the dipole business utterly wrong as sulfur has two lone pairs.
I checked dimethyl sulfide, propane and dimethyl ether and was surprised how differently they behaved. The former has a total dipole moment of 1.5 D ref, while the second near zero (0.1 D) as expected from a non-polar molecule, while the last has 1.3 D.
Whereas there is only a 0.03 difference in electronegativity between C and S (vs. .89), the geometry is different, which I did not know about. Wikipedia on says:

thioether article: "C–S–C angle approaching 90°. The C–S bonds are about 180 pm."
ether article: "C–O–C linkage defined by a bond angle of about 110° and C–O distances of about 140 pm".
alkane article: "154 [pm] for a C–C bond [and] with an angle of 109.47° between them."

So the side chain of methionine should actually be able to hydrogen bond with stuff (not electrostatically, but dispersively). So I was surprised to find papers I had not found prior (ex1, ex2).
S-π are probably the strongest force nevertheless (the bond has an enthalpy of 2.6, 1.1 or 0.7 kcal mol-1 depending on the angleref, which is comparable to N−H···:O hydrogen bond (1.9 kcal mol-1))


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Cleanliness of Fuel
Is there any unit to measure cleanliness of a fuel?
Sort of like calorific value which measures the fuel efficiency, is there any unit which signifies the cleanliness of a fuel.
I have heard of "methane no", but I am not sure whether it is really a unit to measure cleanliness or whether it can be used for fuels other than hydrocarbons?

What are you talking about?

If we consider the degree to which a fuel produces carbon dioxide on burning as the opposite of "clean", then hydrogen would be a perfectly clean fuel (producing only water on combustion); methane with a 4:1 ratio of hydrogen to carbon atoms would be less clean; and long chain hydrocarbons with almost down to a 2:1 H:C ratio would be less clean still.

The amount of sulfur present in a fossil fuel such as coal or oil is another "clean fuel" factor: a fuel that does not contain sulfur does not produce sulfur dioxide on combustion (leading to acid rain).

Why not purity (of hydrocarbons?) by w/w assay?

Please clarify what you mean by cleanliness of fuel. I can think of at least two ways of defining it.

The trouble with the idea of a "clean" fuel is that the fuel is often not the dominant factor in determining whether burning is clean in an engine.
Usually the stuff we don't want in a combustion reaction in an engine consists of oxides of nitrogen, carbon monoxide, particulates, sulphur oxides and other similar things that contribute to harmful things getting into the atmosphere. Some of these are caused by the content of the fuel, like sulphur oxides, but these are mostly very low in actual modern fuels. Raw coal contains a lot of nitrogen and sulphur which makes it dirty however you burn it. But raw coal is almost never used any more because of that problem.
Most vehicle fuels will now have very low levels of things like sulphur. The problem is that just because the fuel doesn't have nasty components doesn't mean the exhaust won't. The trouble is most fuels are burned in air and the reactions in real engines are hard to control so combustion is perfect. So a range of undesirable compounds can still be formed. Some occur because combustion is incomplete (carbon monoxide and particulate carbon, for example). Others are formed because other components of air react to form undesirables (e.g. nitrogen oxides). The nasties are as much a product of the specific environment in the engine as they are of the content of the fuel itself.
What this implies is that the cleanliness of a fuel can not be defined without some consideration of the engine conditions it is burned in. So there will never be a perfect scale of fuel "cleanliness." Obviously if the fuel contains something like sulphur it will be dirtier than a sulphur-free fuel but even a pure hydrocarbon has ways to produce nasty output.

I do not know of an established  way to do this, but it would certainly be a good indication to measure and publish the following things:
What is the ratio of hydrogen to carbon? The more hydrogen,  the more water is produced. Usually this means more energy density (compare oil to methane). Despite the scare in mass media this is almost  irrelevant regarding both health and environmental concerns however (water vapour and carbon dioxide are very comparable in  regards to greenhouse effects). This can used  to guess the next point:
How much 1) soot 2) carbon monoxide is produced in a particular engine (or oven or whatever). Something like nitrocellulose burns rather clean even if there is not much oxigen present.
Are there other elements present(for example nitrogen, sulfur etc) Those certainly burn much "dirtier", often producing irritating gases (nitrous gases etc.) This is probably the greatest concern regarding health.
As a side note: One should also consider whether there are volatile things which are unhealthy, benzene rings come to mind.

This does not provide an answer to the question. To critique or request clarification from an author, leave a comment below their post. - From Review

Why does this not answer the question? I  hardly think it is possible to have a one dimensional quantity for "cleanliness" as there are several aspects. To my knowledge those  are represented rather well in this answer.  The  units  themselves  should be rather obvious for all points.

I felt that it was rather speculative when I was reviewing it. Other people may feel differently.


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How was cholestyramine discovered?
How were the properties of bile acid sequestrants like cholestyramine discovered? The earliest reference to the chemical I could find was the following article which does not discuss the history of the drug:
Glueck CJ, Fallat R, Tsang R. Pediatric familial type II hyperlipoproteinemia: therapy with diet and cholestyramine resin. Pediatrics. 1973; 52:669-79. [IDIS 37416] [PubMed 4355362]
Does this kind of drug exist in nature or is it purely synthetic?

It is definitely synthetic, a trimethyl-ammonium derivative of polystyrene.
The first to use it as a human medication seems to be Stanley S. Bergen, Jr. et al.
See Effect of an Ion Exchange Resin on Serum Cholesterol in Man Proceedings of the Society for Experimental Biology and Medicine, Volume 102, 1959, Pages 676–679:

It is well known that bile acids are a major
  end product of cholesterol metabolism and
  that cholesterol is probably the only precursor
  of bile acids([references]1-3 ) . Rate of oxidation of cholesterol
  to bile acids is believed to be regulated
  by need for bile acids for digestive purposes,
  cholesterol normally being oxidized at
  a rate sufficient to replace the bile acid that
  escapes reabsorption and is lost in the feces
  ([references] 4 - 5). If an agent could be found that would
  sequester bile acids and thereby promote their
  excretion in the feces, 2 consequences of clinical
  interest might be anticipated: (1) increase
  in rate of oxidative degradation of cholesterol.
  and (2) decrease in serum cholesterol level.
  Tennent ([reference] 6) described lowering of serum cholesterol
  levels in animals fed the chloride salt
  of a basic anion exchange resin. This material
  (MK-135) has marked affinity for bile
  acids in vitro and is believed to exert its serum
  cholesterol-lowering effect by virtue of this
  property( 6 ) . The present report summarizes
  results obtained when MK 135 was administered
  to 26 patients, many with elevated serum
  cholesterol levels.

Reference 6 is Tennent, D. M., Siegel, H., Zanetti, M. E., Kuron, C. W., Ott, W. H., Wolf, F. J., Circulation, 1959, v20, 996.
Even though Bergen's paper only uses the term "MK 135", according to the editor's note on the current online version, this paper is the first use of cholestyramine as a medication in humans.


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Manipulate melting point of frozen water
How to manipulate water in way that ice melting point isn't 0°C, but for example 4°C?

One way of doing so would be to significantly increase the pressure. As you can see from this phase diagram, at a pressure of 1 GPa, water ice will remain frozen at temperatures up to about 30°C.


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Is the variation in carbonation of soft drinks the variable that affects the effectiveness of the mentos reaction?
I went and bought a bottle of Fanta and some Mentos to replicate the famous youtube videos of the big fizzy reaction.
The result was dismal, it weakly foamed up and dribbled over the edge.
I tried again with Pepsi Max and had a spectacular result (that the kids loved).
I wondered if the carbonation of the drinks was what impacted the difference in reaction.
My question is: Is the variation in carbonation of soft drinks the variable that affects the effectiveness of the mentos reaction?

Try diet Pepsi. A 2L one.

Thanks! Pepsi max worked a charm and the kids loved it. Going back to my question...

It's my understanding that the two key variables here are carbonation, as you suggested, and viscosity.
People performing this sort of demonstration usually use drinks completely sweetened with artificial sweeteners like aspartame or Sucralose, because they're much less viscous than their full-sugar equivalents, and especially than the high-fructose corn syrup variants.
Common artificial sweeteners are much sweeter than sucrose and glucose [1], so they're added in much smaller quantities and so affect the properties of the liquid less.
The viscosity affects the ability of the liquid to have the rush of bubbles from rapid nucleation required for the fountain effect, and the carbonation affects how much impetus there is behind it.
For an good comparison, you could try three variants of (popular cola product), including a newer one partially sweetened with steviol (look for the ones with green labels: Pepsi Next and Coca-Cola Life are two examples): the wholly artificially sweetened one should jet best, then the part-steviol, part-sugar one, then the full-sugar one.

"Sucralose is about 320 to 1,000 times sweeter than sucrose, three times as sweet as aspartame and twice as sweet as saccharin." -- Wikipedia (I'm not sure how that's measured, but the sucralose figure comes from the FDA approval document cited on the Wikipedia page.)


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Hybridization of sulfur in sulfur dioxide
One of the canonical structures for sulfur dioxide - $\ce{SO2}$ - has sulfur (with a lone electron pair) double bonded to each oxygen atom to form a total of 4 bonds for sulfur - which can be achieved via valence expansion into empty d-orbitals.
What then is the hybridization of the valence-expanded sulfur?
It is described as sp². But how can that be? This seems unlikely
because d-orbitals are involved since the sulfur underwent valence
expansion.
On might imagine a pair of electrons from the 3s/3p oribital(s) being promoted to an empty d-orbital and then having the 3s and 3p orbitals hybridize in to sp².
If this is true it would mean that the lone electron pair of the
valence-expanded sulfur consists of 2 electrons occupying
and unhybridized d-orbital. But is this correct?

The role of d orbital in sulfur bonding is a controversial issue. It seems to me that the problem is not about if d orbital helps, rather than whether the role of d orbital should be regarded as a polarization function.

The structure of sulfur dioxide ($\ce{SO2}$) is quite complicated.
The image from socratic.org* explains it quite well, please find the $\ce{SO2}$ hybridization diagram via the Internet Archive.
As seen, all the atoms have $\ce{sp^2}$ hybridization.
I'll only focus on the central sulfur atom.

Two $\ce{sp^2}$ orbitals form $\ce{\sigma}$-bonds with the two oxygens.
The other $\ce{sp^2}$ orbital is where the lone pair lives in.

Now, we have dealt with 4 electrons, and only have 2 electrons to deal with.
The 2 remaining electrons actually live in the unhybridized $\ce{p}$ orbital.
At this point, the two oxygen atoms will also have 2 electrons left to pair with the sulfur atom.

The two oxygen atoms get to keep their own electron.
Sulfur shares two electrons among itself and the other two oxygen atoms.

Thus, no electron lives in the $\ce{d}$ orbital.

$\color{Red}{\mbox{red}}$ represents number of electrons (I am too lazy to draw the fish-hooks).
$\color{Green}{\mbox{green}}$ represents $\ce{sp^2}$ orbital.
$\color{blue}{\mbox{Blue}}$ represents $\ce{p}$ orbital.

Note that the two $\ce{sp^2}$ orbitals between sulfur and the two oxygen atoms are in $\ce{\sigma}$-bond.

* The website is no longer available; there is no version in the Internet Archive.

We have been taught that $SO_2$ had one $ d\pi - p\pi$ bond and one $p\pi - p\pi$ bond. So is this wrong?

@samjoe $d$ orbital is hardly used, even in $\ce{SO3}$.

@KennyLau Does ozone follow the same analysis? If so, does this mean the central O in ozone allocates electrons differently to the end oxygens? If this is true, are there any significant consequences?

@YushiLi I'm sorry, but this was 5 years ago, and I've forgotten all chemistry since then.

@KennyLau then what is hybridization of S in SO3? If you search on google, I will get more than 10 website saying that both SO2 and SO3 are sp2 and not a single one except this saying something else

@Ritil Really? I can see there clearly just that...

@Mithoron It could be that due to my search history and preferences, Google shows me only such kind of websites which shows sp² hybridisation. I can give examples

@Ritil This answer says sp2 and others too, because it is approximately sp2!

@Mithoron APPROXIMATELY!

@Ritil In SO2 it's approximate, because lone pair isn't oxygen. SO3 on the other hand is more symmetric. BTW I don't get the point of this conversation.

The canonical structure for sulphur dioxide nowadays has charge separation, one oxygen bonded to sulphur in a single bond, the other in a double bond. But not too long ago the canonical structure was indeed what you proposed with one pp-π bond and one dp-π bond.
This was generally explained with one of sulphur’s 3d-orbitals taking part in hybridisation giving rise to a ‘sp²d’ state. Taking Kenny’s scheme, though, sulphur’s 3d-orbitals are a good step above the 3p orbitals in energy. Technically, the 4s orbital should almost come before it. So this hybridisation is highly unlikely.

4s is above 3d on energy only for K and Ca. For all other elements, 3d remains lower than 4s in energy.

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


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Why is the boiling point of pentane (36.1 °C) slightly higher than that of diethyl ether (34.6 °C)?
Both are chains of the same length, except diethyl ether has an oxygen atom in the middle instead of a carbon. Diethyl ether has a slight dipole from the oxygen atom, so shouldn't the intermolecular forces be stronger than for pentane which give it a higher boiling point?

Curious: b.p. of propane: −42.04 °C; b.p. of dimethyl ether: −24 °C

Heptane: 98.7 °C; Dipropyl ether: 90 °C

Nonane: 151.0 °C; Dibutyl ether: 142.4 °C

It seems like it is an actual phenomenon.

It may have to do with branching (contact surface area)...

Maybe this reduces the van der waals/ hyfrophobic interactions? Maybe the oxygen atom is interfering?

The disruption of hydrophobic interactions has probably a greater effect than the establishment of polar interactions.

This is perhaps not a very scientific answer, but consider how branched alkanes have a lower boiling point than straight chain alkanes.  This is due to the inability of the individual molecules to "stack" (if you want to get technical it's to do with surface area).  Now imagine stacking diethyl ether molecules.  The problem arises that two oxygen molecules (2 lone pairs each) on two separate diethyl ether molecules would not want to stack directly on top of eachother as the oxygen atoms have decently high electron densities.  The repulsive force would likely be minimal, as demonstrated by the insignificant difference in boiling points between diethyl ether and n-pentane.  However, this is only an educated guess.
To strengthen my argument, check out Diethyl sulfide, which has a much higher b.p. than both n-pentane and diethyl ether.  This could be due to the larger atomic radius (and therefore larger surface area) of Sulfur in comparison to Oxygen.  The charge has a larger area to be distributed over, resulting in less repulsion between "stacked" sulfur atoms.

None of the answers  explain why alkanes with more than four carbon atoms show the irregularity, while the smaller ones don't.

Even though diethyl ether is a polar molecule but because of steric hindrance,only weak dipole-dipole interactions exist between its molecules.So it have lower boiling point than pentane

Perhaps, I'm missing something here, but how being more polar correlates with a lower b.p.?

First of all this effect is very small. It is on the same order as the b.p. difference of H2O and D2O.
The effect is due to slightly smaller surface area of ether relative to pentane.
Oxygen atom in ether has 2 lone pairs. Lone pairs occupy more volume than hydrogen atoms. As a result the C-O-C angle in ether (110$^o$) is smaller than C-C-C angle in pentane (112$^o$-113$^o$). As a result pentane molecules are more stretched out, have more surface area to interact with neighbors.
For a similar reason neo-pentane is a gas while other pentane isomers are liquids. More "ball-like" structure, less surface area, lower boiling point. For the same reason SF6 and WF6 (bp = 17$^o$C) are two other gases worth mentioning.


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Definition of drift
What does drift actually mean in analytical chemistry? I can't find the exact definition and now am not sure whether my intuitive understanding of it is correct.

According to the International Vocabulary of Metrology (JCGM 200 as well as ISO/IEC Guide 99), instrumental drift is defined as

continuous or incremental change over time in indication, due to changes in metrological properties of a measuring instrument

with the note that

Instrumental drift is related neither to a change in a quantity being measured nor to a change of any recognized influence quantity.

Thus, instrumental drift is caused by a continuous or incremental change in the value of an unmeasured influence quantity during the period of repeated observations.
The consideration of instrumental drift is important when evaluating uncertainty in measurement. According to the Guide to the Expression of Uncertainty in Measurement (JCGM 100 as well as ISO/IEC Guide 98), it must be asked whether all of the influences that are assumed to be random really are random. If there is a sufficient number of observations, the arithmetic means of the results of the first and second halves of the period and their experimental standard deviations may be calculated and the two means compared with each other in order to judge whether the difference between them is statistically significant and thus if there is an effect varying with
time.


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Diffusion Phenomena : Why does diffusion occur?
Why do things diffuse that is why do they go from a place of higher concentration to a place of lower concentration?

Hi and welcome to Chemistry Stack Exchange. It basically boils down to the fact that it is more likely that the particles are going spread apart and become evenly distributed rather than them all accumulating in one spot. So basically, the second law of thermodynamics!

The short answer
They don't.
The long answer
Molecules are very fast. For example, a single molecule in the air around you can be as fast as 500 km/h. The reason that you do not feel them is because they are too small.
Also, you need to understand that there are very, very, very many molecules in a "small" volume.
So, until this point, you have many molecules which travel very fast.
Suppose a gas is divided to high and low concentration areas. The area with high concentration would have more molecules, and let's say 50% of them is travelling towards the area of low concentration (they all move randomly, and in a straight line unless collision happens).
Let's say, 50% of the gas molecules in the area of low concentration is also travelling to the area of high concentration.
Then, you would have a net movement of molecules from high to low concentration.
The keyword here is net. Some molecules do go from low to high concentration. However, they are outnumbered.
After they reach equilibrium, there is still gas molecules travelling between the two regions, just that they are equal in number.
Extra
Of course, because my explanation being not very formal and strict, there would be some errors in my explanation, but I hope that you understand it.

That was very helpful :) @Kenny Lau

As molecules possess kinetic energy of their own, they keep on colliding with each other.
To avoid this collision, they have a tendency to move away from each other, especially the gaseous molecules, as they have the highest kinetic energy.
So they start moving.
But the net movement of molecule is from high concentration to low concentration so that equilibrium can be reached.That's why diffusion occurs.


text:
Determining actual safety from a vague material safety data sheet
My desktop computer system employs a liquid cooling loop featuring a 6-litre external reservoir and radiator combination "cooling tower" unit. It isn't a cooling tower in the strict sense as it does not achieve temperature reduction through any evaporative method - only by radiating heat through the metal structure. However, the reservoir is not sealed and I have noticed a reduction in the level of liquid over the past few months since I started using it - I must assume that I'm breathing-in this liquid in some amount.
The coolant used describes itself as:

XSPC EC6 is a high performance, eco-friendly coolant for PC water-cooling. It’s based on a blend of refined vegetable extracts with non-toxic corrosion inhibitors and non-toxic dyes. It offers superb protection for copper, brass, steel, nickel, aluminium and has been tested with a range of acetal and acrylic plastics

Never one to trust marketing materials, I requested a material safety data sheet from the manufacturers, and details of the ingredients. I only received an MSDS back without an ingredients listing. The MSDS reports the following:

COMPOSITION / INFORMATION ON PRODUCT COMPONENTS
  Hazardous Components for EC
  Low oral toxicity proprietary antifreeze formulation based on refined vegetable extracts and non-toxic corrosion and scale inhibitors.
  Classified as GRAS (Generally Regarded as Safe) as a food additive and not classified as harmful for handling or transport purposes.
  However we recommend good working practices are employed and personal protection gear is worn when handling any chemical.
[...]
EXPOSURE CONTROL/PERSONAL PROTECTION
  Respiratory Protection: Respiratory protection if there is a risk of exposure to high vapour concentrations.
[...]
TOXICOLOGICAL INFORMATION
  Acute Toxicity: LD 50/oral/rat: >20000 mg/kg
[...]
ECOLOGICAL INFORMATION
  Mobility: The product will dissolve rapidly in water.
  Ecotoxicity: The product is rated as:
  Algae EC50 (72h):>100mg/l.
  Fish : LC50 (96h):.100mg/l, Oncorhynchus mykiss
  Bacteria: >1000 mg/l Daphnids (acute) EC50 (48h):>100mg/l

My guess is that it's some glycol-based solution - I'm disappointed they don't reveal the actual ingredients.
Is this substance anything I should be concerned about? Is there any  carcinogenic potential?
Also, how can I determine the specific heat capacity of the coolant? As it is used in a reservoir system instead of an active heat-exchange system the heat-capacity is important - I would like to be able to calculate exactly what TDP (Thermal Design Power) computer components I can use and how long I can run them for at-load without fear of overloading the cooling system.

Unfortunately, it is a rather common practise to keep the actual composition of formulations a secret. However, I doubt that ethylen glycol is a major component in the mixture. I'd rather assume that it contains 1,2-propanediol and 1,2,3-propanetriol, which are both much less toxic that ethylene glycol.

@KlausWarzecha Those substances seem safe - though the docs state it produces  acrolein "when heated" - the temperature of the coolant in operation is about 35-40 degrees Celsius - does that constitute "heating" and is that hot enough to produce acrolein?

I do not expect that acrolein is formed at temperatures around 35-40 °C. Classical catalytical transformations of  glygerol to acrolein were gas phase reactions using aluminium oxide or silicon oxide catalysts and temperatures in the range of 180-350 °C. Your computer will hopefully never reach these temperatures ;-)

General common sense should always remind you not to drink the liquid stick your finger in it, rub it all over your skin or breathe its vapours for extended periods of time. Paracelsus is best known in Germany for this quote:

All things are poison and nothing is without poison; only the dose allows that a thing is not a poison.

If you are now scared out of your wits, then please relax and consider that preface a general disclaimer. Because all the toxicologic data I can see points towards the liquid being entirely harmless. Here’s why:

Acute Toxicity: LD 50/oral/rat: >20000 mg/kg

That’s more than $20~\mathrm{g/kg}$ — an incredible number for $\mathrm{LD_{50}}$ values. Usually, the highest $\mathrm{LD_{50}}$ values I see are $2~\mathrm{g/kg}$. And the greater than basically means ‘we couldn’t kill the rats with that oral dose so we gave up.’

Low oral toxicity proprietary antifreeze formulation based on refined vegetable extracts and non-toxic corrosion and scale inhibitors. Classified as GRAS (Generally Regarded as Safe) as a food additive and not classified as harmful for handling or transport purposes. However we recommend good working practices are employed and personal protection gear is worn when handling any chemical.

(Emphases mine) As you can see, there are numerous references to ‘nope, this thing really won’t harm you.’ The GRAS classification actually seems to be concerning food additives. No danger here.

Respiratory Protection: Respiratory protection if there is a risk of exposure to high vapour concentrations.

I’ll admit, this sounds a tad unsafe at first glance. But if you know what MSDS write for truly unsafe compounds, you’ll realise that this is saying ‘If we write “safe”, somebody will stick their nose in, suffocate and then sue us, so we have to put some kind of caveat here.’

The only thing that actually hints at a potentially harmful substance is the thing about high vapour concentrations. But remember that this is a cooling liquid; you don’t want a cooling liquid evaporating away, that would reduce the cooling efficacy. The minute amounts you are losing will dilute well in the air in your room. Open the window once a day or every two days and you’ll have removed anything that can potentially harm you. The liquid is as safe as it gets.
However, don’t sue me if you decide to turn off common sense and start drinking the liquid or stop venting your room or whatnot.

It can never be said that you are "safe" from any chemical until you find out what the fluid is.
My money is on it being a variant of propylene glycol with merely some dye and that they're trying to play the "secret unique recipe" card. Its an easy way to up-charge on a super useful and cheap product. You also said "a small container over a few months", correct? If so, I believe at that rate of evaporation, you're probably safe. I'm sure exposure to a fog machine in a "club" would expose you to more glycol.
If your really worried and need to get the info, you can play a slightly deceptive and (some would say) unethical trick. Call they're 800 number on the back of they're packaging and say some of their fluid spilled in (x) family members eye and you're taking them to the hospital. Say you've been flushing the eyes with water and need to know the primary or any toxic chemicals in the reservoir.
--- This situation above is completely hypothetical of course.

Heh. Very nice, but there's no guarantee they'll actually tell you what's in it: They may just give you a spiel of something to tell the doctor and how to flush the eye properly.


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How would using different salts in a salt bridge effect the voltage generated by a voltaic cell?
If I use a range of different nitrate solutions (e.g. calcium nitrate, iron nitrate) for the salt bridge in a voltaic cell, what effect would it have on the flow of ions in the voltaic cells?
Will the different properties of the salts such as ionic radius effect the voltage generated by the voltaic cell?
If a certain property would most likely have an effect, what would it be and why?
(Both the cathode and the anode are kept constant with copper in copper nitrate solution and zinc in zinc nitrate solution)

"If I use a keep changing"?!

How can you do electrolysis if you keep changing the salt?

The salt bridges used by mine (so far) were filled with aqueous electrolytes with solutions of either potassium chloride, potassium nitrate, or sodium chloride.  For an electrochemical cell of $\ce{Zn | Zn^2+||Cu^2+|Cu}$ described by you, i.e. based on
$\ce{Cu^2+ + Zn \rightleftharpoons Cu + Zn^2+}$
these appear to me a resonable choice, because

Their ions (of KCl, KNO3, NaCl) would not interfer with the redox reactions, so the potential between the two electrodes (the "voltage" you mention) were not affected.  Referring to the standard potentials, $\ce{Fe^3+}$ of iron nitrate in contrast, is more likely to interfer with your system than the $\ce{Ca^2+}$ of calcium nitrate, esp. in aqueous solution; but the design of the salt brigde (diaphragm) may reduce this electrochemical interference.
Beside electrochemical inertia in the window of investigation, an additional criterion to select electrolytes for salt bridges is their cations and their anions should possess a similar mobility in the electrolyte of tehe salt bridge (the flow you mention).  The rate of the ions to move not only depends on the viscosity of the medium and the charge of the specis, but on the efficient radius of the ion (the ion plus the hydration sphere), too; yet this should not influence the potential between the electrodes, rather than the current between the electrodes if the set up is intended to transform chemical into electrical energy.

In short: within reason, keep the KCl in the salt bridge.


text:
Does the activity product decrease with increasing ionic strength?
I have a question regarding what happens to the solubility of $\ce{CaSO_4}$  or the rate that $\ce{CaSO_4}$ will dissolve when we add $\ce{NaCl}$ to the solution:
$\ce{CaSO_4}$ will dissolve into $\ce{Ca^{2+}}$ and $\ce{SO_4^{2+}}$:
$$\ce{CaSO_4$ $<=>$ $Ca^{2+}$+$SO_4^{2+}}$$
The ionic strength has the expression:
$$I = \frac{1}{2} \sum (m_i z_i^2)$$
where $m_i$ is the molar concentration of species $i$, and $z_i$ is the charge of species $i$.
If we add $\ce{NaCl}$ to the solution, it follows that we get more ions, and therefore the ionic strength will increase.
The expression for activity coefficients for ionic strength up to about 0.5 is:
$$\log\gamma_{i} = -A z_{i}^{2} \left({\sqrt{I}\over 1 + \sqrt{I}}\right) - 0.3 I$$
where $A$ is a constant, $\gamma_{i}$ is the activity coefficient for species $i$, and $I$ is the ionic strength. If I then use $A = 0.5$, I get:
$\gamma_{Ca_{2+}}$=0.296 at $I=0.5$
$\gamma_{Ca_{2+}}$=0.291 at $I=0.4$
$\gamma_{Ca_{2+}}$=0.297 at $I=0.3$
$\gamma_{Ca_{2+}}$=0.318 at $I=0.2$
Since we increase the ionic strength by adding $\ce{NaCl}$, is it correct that the activity product of calcium and sulfate would decrease (even though I did not get that in the above calculations), and that because of this calcium sulfate would dissolve more?

LaTeX tips: use $\ce{Ca^{2+}}$ for better formatting!

Yes your conclusions are right. Increasing ionic strength tends to favour dissolution


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Edge length unit cells ionic crystals
In ionic crystals there are mainly 3 types of structures: Face-centered cubic and octahedral interstices; face-centered cubic with tetrahedral interstices; and body-centered cubic.
My question is how can I know the edge of the unit cell of this crystals for calculating the theoretical density of a crystal?

Well, you need to know what type of unit cell the lattice is and the ionic radius of the atom.

As you see in my question they are all cubic. I know I need to know the ionic radius I'm looking for a formula. For example I know that for Central Cubic Faces $a= 2 (r_{cat} + r_{an})$

What is the Spanish word you are thinking of? 'corpse' is not the right translation I think.

@bon - "cubic with central corpse" = "body-centered cubic (bcc)"

https://en.wikipedia.org/wiki/X-ray_crystallography


text:
Why is the first ionization enthalpy of zinc higher than expected?
The following sentence is stated in my current chemistry book:

The first ionization enthalpy of zinc is higher because it represents an ionization from the $\rm 4s$ level.

Now zinc does have an ionization enthalpy of $\pu{906 kJ/mol}$, which is higher than copper's $\pu{745 kJ/mol}$ or nickel's $\pu{736 kJ/mol}$ in the same series. But I am not sure that the reason for this increase stated in the book is correct. For example, even copper has an ionization from the $\rm 4s$ level, but it does not have anywhere near as high ionization enthalpy as zinc. Is the book correct, or is there some other reason for this?

The book is incorrect. Cu ionisation is also from 4s. The reason is purely because of the increase in effective nuclear charge in Zn.

Do you know the real reason behind this unexpected increase

I've been wondering about this myself. The sudden jump in Ie from Cu to Zn isn't very well explained by mere d orbital filling because it doesn't seem to be present in p orbital filling. I have a theory (only a theory) that what's really going on is the S orbital is only getting one electron, then the p orbitals are filled, then the S orbital is completed at the end. If that's the case, I have no idea why it would happen that way.

It is because the shell is complete.
Zinc's electronic configuration is $\rm [Ar] 3d^{10} 4s^2$, which is a complete configuration for a transition metal. The next element in the same period is $\ce{Ga}$, which starts filling $\rm 4p$ orbitals ($\rm [Ar] 3d^{10} 4s^2 4p^1$).
The zinc's electronic configuration causes a sort of special stability because of the filled electronic states. Also, the effective nuclear charge is greater than $\ce{Cu}$'s, so it's more difficult to ionize $\ce{Zn}$.


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Is nifedipine a weak base or weak acid?
I was wondering whether someone could explain me if nifedipine is a weak base or a weak acid.
I barely can find any information about nifedipine's physicochemical properties in studies, however some suggested that it is a weak acid. If it is a weak acid, which hydrogen atom is dissociable and what would the pKa of that group be?
Thanks in advanced.
Nifedipine's structure, as shown on wikipedia:

Here is a reference which states that the "strongest basic" pKa = 5.33. As for which proton, I'm not sure.

@ToddMinehardt perhaps the H atom at position 4 of the 1,4-dihydropyridine.

My guess is still the N-hydrogen. However a $pK_a$ of 4 is more acidic than I would expect, e.g. in comparison with phtalimide. But there may be higher order effects that affect the acidity.

Nifedipine is a weak acid.

$pK_b = 10.07$
$pK_a = 3.93$
$pK_a = 4.3$

Removed the statement that the N-hydrogen might be the acidic one.

My speculation is a bit related to the answer provided by @aventurin.  In absence of experimental evidence to support this, I think in a first step the proton in (then) para-postion to the nitrogen is abstracted:

leading to a protonated, pyridinium specis, then in equilibrium to the non-protonated form.  This may be favourable by thecreation of an aromatic system.  Again, provided as a speculation including the notice that the initial deprotonation might be irreversible.

The abstraction of a proton in the first step would result in a nifedipine anion.

Your image shows a 3-nitrophenyl compound; however, nifedipine is a 2-nitrophenyl compound.

That nitrogen has a lone pair, your proposed "deprotonation" would lead to a nitrogen with 10 electrons around it (and as aventurin points out it should be negatively charged not positively). Your reaction is in fact a dehydrogenation and the intermediate results from a loss of hydride, not a deprotonation.


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Why derived name of nomenclature is used for Alkanes and alcohols only?
My chemistry textbook says derived system of nomenclature of organic compounds is only limited to alkanes and alcohols!
What is the reason behind it? Is it chemical or like a ritualistic/cultural?

Could you cite that textbook, and, for the good of everyone else not having that arcane folio, explain what "derived system on nomenclature" is?

The derived system is really, really old. Even my textbook from the 20th century has IUPAC naming. My recommendation would be to just find IUPAC naming online, or just get a new textbook.
I can't seem to find much on the derived system, but here's another question about it; LINK
I tried searching about, and there really isn't much about it. I don't see why you'd want to learn about it though, as history of nomenclature isn't very useful or interesting. Besides, if you try to use the derived system, no one else is going to know what compound you're describing.

Bangladeshi textbooks have this system included in the context. I was just interested to know why this system was used for alkanes and alcohols only.


text:
How are the electrons of benzene actually configurated?
I found this diagram:

According to it, it is not true that the six electrons of carbons in the pi-orbitals above the ring are all identical.
But, how can it be the case, seeing that this molecule should have an order-6 rotational symmetry?

The electrons are indistinguishable, you cannot look at the system and say "this electron is in the HOMO whereas that one is not"

But I sure can count the number of electrons in the non-bonding pi-orbital?

Yes but you can't say which electrons they are. There are two of them in each orbital but they are not individually distinguishable.

I don't think I ever said they are?

The $C_6$ rotation can alter the phase of the molecule's electronic wavefunction, which is what is depicted in the above diagram, but it does not affect the electron density $|\psi|^2$ which is the quantity that must have invariance under $C_6$. In this case the wavefunction itself should also actually be invariant under the $C_6$ because it is a closed shell molecule (symmetry $A_{1g}$)

So, as I told you in chat what you have here is the orbitals of something which looks like, but which is not, benzene. It is the orbitals of one the all its two mesomeric form, it is so a static view.
Now let me show you the orbitals which are in the benzene. I only draw with the p-orbitals, I won't do the hybridation with the hydrogen it will be too difficult for me two draw and it will not be more effective.

How to draw a molecular diagram when the molecule has a lot of atoms?
First draw the diagram like if in the molecule there is just its squeleton (atoms which interact to create the shape of the molecule, for example with benzene only for the carbon atoms) and do like if these atoms were hydrogen atoms.
Second draw the correct orbitals with the same symetry of what you got with hydrogen just before.
You are obviously not obliged to do it but it is easier to see which orbitals in the fragments have the same symetry or not.

First step
Here for the benzene I will draw the MO's diagram of benzene with $\ce{H_6}$ first, with a molecule which is a rectangle of hydrogen atoms and the other fragment is a longer dihydrogen molecule than the original one. After seeing which of them have the good symetries I get this

Second step
Here we want the same but with the p-orbital of the carbons so keeping the same symetries or antisymetries you get this

And then here you have the correct orbitals.

Be careful
The color on the orbitals represent the sign of the wave function which has no meaning in the reality. So you can make the black and the white anywhere you want if you obviously respect the molecule symetries.
In this example I draw exactly the same orbital.

I say that because the orbitals I drawn for benzene are not exactly in the same color logic of the MO's diagram of $\ce{H_6}$ but as you can see each one have the same symetries or antisymetries.

And then the answer of your question
On the MO's diagram you show they show the interaction of the p-orbitals in one of the mesomeric structure, like this

And then the symetries are not the same as in the benzene but at one of its mesomeric form which is a static view like this
In fact you need to imagine the same but with an interaction of all p-orbitals and the get a ring orbital for $\psi_1$, and the it explain the real shape of the benzene like this
Remember that orbital diagrams are done to think about what you observe in the nature and not to predict it. Well you can use quantum theory to predict things but it is an other subject.
I hope it can help !

Some questions: 1. The s orbitals don't react, right? 2. The first picture is not really an MO diagram, but the possible configurations, right? 3. How do you know they are the actual mesomeric forms? Sure, they look symmetrical, but...

Some answers: 1. The s orbitals of which atoms? Orbitals of hydrogens obviously interact if not the hydrogens will not make bond with carbons... 2.The first picture really is a MO's diagram, just it's not one of a real molecule. 3. There are because of the experiment you did. As I told you "diagrams are done to think about what you observe in the nature not to predict it" so if you want to make the MO's diagram of benzene you must need to know it is an aromatic coumpound. :-)

How do you go from, you know, "it is aromatic", to the pictures there?

https://en.wikipedia.org/wiki/Aromaticity you should read about that. It is a complex subject.

Your depiction of the orbitals is almost the same as in the question, with some leeway they are both qualitatively in the realm of correct. (Even for a schematic representation, they miss certain key features.) It does, however not answer the question why these reproduce the molecular symmetry. And the part about mesomeric structures is completely wrong. An MO scheme cannot really show such things, and at the same time adhere to symmetry, which the OP's clearly does. The part that you have to know benzene is aromatic to construct the diagram is also completely false.

@Martin-マーチン I am open to learning more on the subject. At the time I gave the answer I was finishing to have an introduction to that subject, even though we came deep into it. Now I am not an expert and nowadays I am doing chemical engineering so I won't be able to correct that by myself.

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


text:
Powdering aluminum chemically?
I need a fine aluminum powder for many reactions/experiments. Could I dissolve the metal in an acid and then drop it out of solution using a base, or vice versa? What would be the best chemicals/process to produce a high yield of aluminum powder?

could you explain the chemicals needed/produced in the reaction? other than aluminum metal of course

Burn glucose in your muscle cells and actuate a file. ;-)

I already have aluminum filings that i want to convert into a finer powder. For example: I tried making thermite with the filings but couldn't get it to oxidize, my guess is the surface area is too low still for the aluminum to react with the finer iron-oxide powder.

The finer the Al powder, the larger the surface exposed to air (and, oxygen), so it will oxidize more rapidly on surface (passivation).  Not tried out by myself, but this sounds worth a try.

There is no realistic way to produce aluminium powder at home. You should stick with buying it. There are some work-arounds for magnalium, but I don't recommend it. Aluminium powder is commercially available. There are several grades with varying grain sizes. I recommend to buy a coarse one (maybe even granules) for chemical experiments and one of the finest for aluminotermia experiments.

"Could I dissolve the metal in an acid ..."? Yes, and produce one liter of hydrogen per gram.

There is no simple and affordable way to obtain fine aluminium powder chemically. As permeakra said, in the absence of a ball mill, it's just easier to buy some.
Note that a finely dispersed aluminum powder obtained commercially is always going to have low level of impurities such as $\ce{Fe, Mn, Si, Mn}$ (from the grinder's balls) and traces of wax/fats on the surface (added by manufacturer on purpose to prevent adhesion and separate particles from each other).
As for the chemical/physicochemical methods, there is a great overview included in the US5885321A patent, among which there are the following:

exploding of aluminum wire in a vacuum by a high electric current (particles' size varies a lot);
condensation of vaporized aluminum in a current of inert gas (very energy-inefficient);
decomposition of aluminium carbonyls (there is no experimental evidence for any of these to exist);
reduction of halides in solution (issue with separation of the particles from the salts);
CVD of aluminium thin films by decomposition of tertiary amine complex of aluminum hydride (complex multistep process with low yield and expensive and volatile precursors);
decomposition of the dialkylaluminum hydride in $\ce{H2}$ atmosphere at elevated temperatures (up to $\pu{260 °C}$) with a catalyst ($\ce{Ti, Zr, Hf, Ln}$) and a tertiary amine (no particle size control, questionable efficiency);
decomposition of diethylhydridoaluminum or diisobutylhydridoaluminum in diisopropyl ether or triethylamine at $\pu{90 °C}$ to $\pu{185 °C}$ with titanium isopropoxide catalyst.

The patent itself proposes

decomposing alane-adducts in organic solvents under an inert atmosphere to provide highly uniform particles selectably sized from about $\pu{65 nm}$ to about $\pu{500 nm}$

where adducts are

trialkyl amines and tetramethylethylenediamine, ethers and other aromatic amines;

and conditions such as

atmospheric pressure and at temperatures as low as $\pu{50 °C}$ with xylene solvent, ... , toluene, dioxane, and tetramethylethylenediamine.

As a catalyst, titanium halide, amide and alkoxide are used. Authors also provide a complete table of the systems tested:

References

Higa, K. T.; Johnson, C. E.; Hollins, R. A. Preparation of Fine Aluminum Powders by Solution Methods. US5885321A, March 23, 1999. (PDF)


text:
Is there a substance I can use to coat a projectile to eliminate friction heat?
Let's say that I was going to shoot a can of beer into space from one of those "vacuum cannons." I would expect that the friction with the air around the projectile would cause the whole thing to just melt and burn (explode).
Is there such a substance that, if used to coat the projectile of my contraption, would eliminate the heat problem by basically "consuming the heat in a reaction" on the way out of the atmosphere?

Daresay if you had such a coating you'd be a multi-quadrijillionaire.

Large bodies such as meteors burn while plunging through the atmosphere because of adiabatic compression of the air in front of them, rather than friction. I'm not sure whether the same remains true for an object as small as a beer can. I suspect the beer can would actually remain quite cool, as the square-cube law suggests heat would be transferred proportionally much more efficiently away from the beer can.

@NicolauSakerNeto Uh. Jinx...ish?

@Brian Pretty darn close, though I had to edit some more stuff in!

@NicolauSakerNeto Yours is a lot more substantial, so all in all I'd say you win.

I guess you're talking about ablative coating, like used for atmospheric reentries - https://en.wikipedia.org/wiki/Atmospheric_entry#Ablative

A quite common technique is to put a small hole in the top of the beer can such that a thin film of beer could flow around the can cooling it.

IF something is travelling really fast through the atmosphere, friction heat may not be the biggest problem. So eliminating friction won't keep the beer cool.

Try using carbon fiber in a uniquely woven shape. Maybe apply a dual layer with the top layer having a unique shape for heat dissipation. There have been a few carbon fiber helix designs that maintained structure and supported weight with a 2000 F acetylene/butane flame directed on it.
Another option is go with layers of powdered graphite. See if you can convert thermochemical properties with blends. Maybe you can add a coolant system for your beer can with the excess energy stored.
If you are really trying to have your beer survive extreme heat,"Thermo-electric Reactions and Currents Between Metals in Fused Salts" by Thomas Andrews is probably a good start for prototype blends. I think google has a free .pdf available through Google books.
Here is a graph that might help with selecting your surface metals, if you ever take this project on seriously:

If nothing works, at least your efforts should land you a pretty hardcore beer-koozie.
Final deciding factors come down to the degree of effort you are wishing to exert and material expense.

Do you want to break up this post into a few more readable paragraphs? Also, you can insert photos using the image button above the text.


text:
Coordination number of carbon in methyllithium tetramer
According to J. D. Lee's Concise Inorganic Chemistry, the coordination number of carbon in $\ce{Li4(CH3)4}$ is $7$. But the structure of the tetramer is
Doesn't carbon have coordination number $6$ here? Am I missing something?

TL; DR: coordination number of carbon atom in methyllithium tetramer is indeed 7, arising from 6 intramolecular interactions, as you suggested, and an extra bond with the lithium atom of the next tetramer.

The key point is that you need to look beyond the tetramer. I'm sorry to inform you that the image from Wikipedia of the standalone cubane $\ce{Li4(CH3)4}$ is useless for the determination of coordination number (C.N.) of carbon. In fact, there is a 3D network of interlinked clusters, and the description given in J. D. Lee's Concise Inorganic Chemistry is in fact very precise [1, pp 304]:

The structure of the $\ce{(LiCH3)4}$ cluster is unusual. The four $\ce{Li}$ atoms occupy the corners of a tetrahedron. Each methyl $\ce{C}$ atom is above a face of the tetrahedron, and forms a triple bridge to the three $\ce{Li}$ atoms that make up the face of the tetrahedron. The intramolecular $\ce{Li-C}$ distance is
  $2.31~\mathrm{Å}$. The $\ce{C}$ is bonded to the three $\ce{H}$ atoms in the methyl group. The $\ce{C}$ is also bonded to a $\ce{Li}$ atom in another tetrahedron (with an intermolecular $\ce{Li-C}$ distance of $2.36~\mathrm{Å}$). The coordination number for the $\ce{C}$ atom is therefore $7$.

The problem I stumbled upon is that I didn't find any ready-to-use CIFs for methyl lithium with 3D coordinates (there are plenty for the derivatives and some powder XRD of questionable quality). I ended up creating my own CIF based on data published by Weiss et al. [2] with VESTA program and visualized the content of the unit cell with OLEX2:

Figure 1. Unit cell of the methyl lithium tetramer with the central cubane core $\ce{Li4(CH3)4}$ linked via four crystallographically equivalent $\ce{C}$ atoms with 4 neighboring cubanes displayed as tetrahedra $\ce{Li4}$ for clarity. Color code: $\color{#EEEEEE}{\Large\bullet}~\ce{H}$; $\color{#CC80FF}{\Large\bullet}~\ce{Li}$; $\color{#909090}{\Large\bullet}~\ce{C}$.

Figure 2. Coordination environment of the carbon atom and interatomic distances (Å).
In addition, I run the test for C.N. with Dirichlet subroutine of the ToposPro, which also suggests the C.N. 7:
Central atom:C1 0.320 0.320 0.320 Rsd:0.900
 D(CP):0.624  ( 0.3698 0.3698 0.3698 )
 D(VDP):0.364  ( 0.3490 0.3490 0.3490 )
Atom:0.980 < r < 2.370  <r>=1.747   Top: 1.229 < R < 1.556  <R>=1.357
CN=7:0:0 NV=10 V=3.052/6.126 S=13.376 Cpac=0.161 Ccov=5.175
G3=0.114818789
Face distribution: {3/1 4/3 5/3 }
Vertex distribution: {3/10 }

  N  Atom   x      y      z     Dist.   D1      D2     SAng.
  1  H_1  0.3510  0.1920  0.3510  0.97957  0.89982  0.07975  27.31224
  2  H_1  0.1920  0.3510  0.3510  0.97957  0.89982  0.07975  27.31224
  3  H_1  0.3510  0.3510  0.1920  0.97957  0.89982  0.07975  27.31224
  4  Li1  0.3690  0.3690  0.6310  2.30686  0.89982  1.40704   5.76860
  5  Li1  0.3690  0.6310  0.3690  2.30686  0.89982  1.40704   5.76860
  6  Li1  0.6310  0.3690  0.3690  2.30686  0.89982  1.40704   5.76860
  7  Li1  0.1310  0.1310  0.1310  2.37007  0.89982  1.47025   0.75748

Figure 3. Voronoi-Dirichlet coordination polyhedron for the carbon atom.
Content of the CIF file:
#======================================================================

# CRYSTAL DATA

#----------------------------------------------------------------------

data_VESTA_phase_1

_chemical_name_common                  'Methyl lithium tetramer'
_cell_length_a                         7.24000
_cell_length_b                         7.24000
_cell_length_c                         7.24000
_cell_angle_alpha                      90
_cell_angle_beta                       90
_cell_angle_gamma                      90
_space_group_name_H-M_alt              'I -4 3 m'
_space_group_IT_number                 217

loop_
_space_group_symop_operation_xyz
   'x, y, z'
   '-x, -y, z'
   '-x, y, -z'
   'x, -y, -z'
   'z, x, y'
   'z, -x, -y'
   '-z, -x, y'
   '-z, x, -y'
   'y, z, x'
   '-y, z, -x'
   'y, -z, -x'
   '-y, -z, x'
   'y, x, z'
   '-y, -x, z'
   'y, -x, -z'
   '-y, x, -z'
   'x, z, y'
   '-x, z, -y'
   '-x, -z, y'
   'x, -z, -y'
   'z, y, x'
   'z, -y, -x'
   '-z, y, -x'
   '-z, -y, x'
   'x+1/2, y+1/2, z+1/2'
   '-x+1/2, -y+1/2, z+1/2'
   '-x+1/2, y+1/2, -z+1/2'
   'x+1/2, -y+1/2, -z+1/2'
   'z+1/2, x+1/2, y+1/2'
   'z+1/2, -x+1/2, -y+1/2'
   '-z+1/2, -x+1/2, y+1/2'
   '-z+1/2, x+1/2, -y+1/2'
   'y+1/2, z+1/2, x+1/2'
   '-y+1/2, z+1/2, -x+1/2'
   'y+1/2, -z+1/2, -x+1/2'
   '-y+1/2, -z+1/2, x+1/2'
   'y+1/2, x+1/2, z+1/2'
   '-y+1/2, -x+1/2, z+1/2'
   'y+1/2, -x+1/2, -z+1/2'
   '-y+1/2, x+1/2, -z+1/2'
   'x+1/2, z+1/2, y+1/2'
   '-x+1/2, z+1/2, -y+1/2'
   '-x+1/2, -z+1/2, y+1/2'
   'x+1/2, -z+1/2, -y+1/2'
   'z+1/2, y+1/2, x+1/2'
   'z+1/2, -y+1/2, -x+1/2'
   '-z+1/2, y+1/2, -x+1/2'
   '-z+1/2, -y+1/2, x+1/2'

loop_
   _atom_site_label
   _atom_site_occupancy
   _atom_site_fract_x
   _atom_site_fract_y
   _atom_site_fract_z
   _atom_site_adp_type
   _atom_site_B_iso_or_equiv
   _atom_site_type_symbol
   Li1        1.0     0.131000      0.131000      0.131000     Biso  1.000000 Li
   C1         1.0     0.320000      0.320000      0.320000     Biso  1.000000 C
   H1         1.0     0.351000      0.351000      0.192000     Biso  1.000000 H

References

Lee, J. D. Concise Inorganic Chemistry, 4th ed.; Chapman & Hall: London ; New York, 1991. ISBN 978-0-412-40290-6.
Weiss, E.; Hencken, G. Über Metall-Alkyl-Verbindungen: XII. Verfeinerung der Kristallstruktur des Methyllithiums. Journal of Organometallic Chemistry 1970, 21 (2), 265–268. DOI: 10.1016/S0022-328X(00)83621-9. (In German)


text:
Cheap metal to replace platinum in hydrogen fuel cell
Platinum is used as a catalyst in hydrogen fuel cells because it increases the efficiency. However, if I just want to make a fuel cell for demonstration purposes, I don't want to use a lot of money to buy platinum wires. What "common" metal (aluminium, iron, etc.) works the best for a hydrogen fuel cell?

Actually, I've seen carbon used, but as far as I know only on the electrolysis end; it isn't that great of a conductor. Most metals are bad because they tend to ionize easily (and thus effectively corrode).

for demonstration purpose it is easier to stick with platinum despite its cost. It is possible to use, say, nickel on hydrogen end and a metal oxide covered  electrode on oxygen end, but the setup is a lot less reliable. Platinum wire is useful for other thing as well (say, for flame test)

Platinum "wires" won't get you very far anyway. An old, proven design is the alkaline fuel cell, which supposedly runs very well with silver or nickel.

I would use alluminum plates. I've already tried it once and it's a great way to demostrate how a fuel cell works but where the bubbler (hydrogen) is you'll notice small signs of corrosion after 2 minutes of use. Mine were 1 1cm thick and they did their jobs perfectly. Alluminum should be your best bet.

thanks, this kind of thing was what i was looking for. the experiment has already passed but I'll keep this in mind for future use.

After I started my own project I've ended up using Graphite instead. They are easier to find and replace since the cost is way less.

We have tinkered with this a little. Currently, a handful of others have been experimenting with certain blends of metals, using lower costs to create equal reactions.
I have heard a lot of different ratios of Ni/Cu and Pd blends. But a good experimental blend (if your making a reasonable size batch), would comprise of a ZrO2–Pd(x)D core, with (x) varying as a slew of other metals.
 This article is regarding the direct deposition of hydrogen on palladium, probably a little more refined towards what you where searching.
While im not sure how intricate your going to get with your experimental cell, here is some good food for thought., If you can create a vessel with proper pressure safety, academic groups have been experimenting with super critical fluids, trying to test metallic blends as catalysts under states past standard classical critical conditions.


text:
After opening, when is a can of Diet Coke most acidic? (Based on a true story)
Situation: Clogged toilet, no plunger.
Goal: Trying to use the acidity from a can of Diet Coke as toilet drain cleaner.
Question: How can I maximize the acidity of a Diet Coke?
Question 1: Should I heat up the Coke (using a microwave)?

Is an acidic solution more acidic when it is hot or cold?

Question 2: Would heating the Coke make the Coke less soluble to Carbon Dioxide (which is acidic)?

Is a solution more soluble to gases when it is hot or cold?

Question 3: What is the loss of carbon dioxide (acidity) during the time taken to microwave the soda?

What exactly do you think the acidity will do to help? Why do you think the acidity of any soft drink will be great enough to achieve that effect?

Stop worrying about that. Vinegar should do the trick :)

Coke (or any carbonated drink) is most acidic when it is just opened. The acidity comes primarily due to phosphoric acid. Citric acid if present also adds some acidity. Carbon dioxide (which is present as dissolved form) will increase acidity though by  small amount. Carbonic acid, formed upon dissolution of $CO_{2}$ is a weak acid. Its contribution to acidity, when phosphoric acid is present would be negligible.
Facts

Solubility of a gas is more in cold liquids.
Carbon dioxide dissolves in water. The dissolved $CO_{2}$ is now in equilibrium with carbonic acid. The so formed carbonic acid is much less in concentration than $CO_{2}$ itself.

Considering small effect on acidity from carbonic acid:
 Just opened coke can will have higher carbonic acid (whatever be its amount), due to higher (relatively) dissolved $CO_{2}$. As temperature increases $CO_{2}$  fizzles out to air, and dissolved $CO_{2}$ is reduced  so is carbonic acid.
PS- Microwaving coke can will be dangerous. There is a hazard of explosion and severely damaging your microwave oven of-course !
I am not sure if cleaning with coke will help in your situation.

Your reasoning about the carbonic acid is wrong. The presence of the other acids drive its equilibrium even further towards the dissolved dioxide, I'm sure it does not measureably influence the pH at all. Heating the solution however will increase its pH, and increase whatever little effect it could have on the clogged toilet. But you're right, it probably won't help.

You are right about carbonic acid. I agree. I'll improve my answer. Thanks for your comment.

You can upvote helpful comments, btw. :-)


text:
shifting of equilibrium
Does it mean that if one side of a reaction is favoured, the reaction will shift to the other side? As I was reading up on the van't Hoff equation, it says that an increase in temperature of an endothermic reaction will favour the products. It also showed a graph where lnK increases with temperature (T). But I thought what would happen in this case is the reactants will be favoured and then the equilibrium will shift to the right to de-stress the equilibrium?
So in summary, I'm just quite confused about what it means by favouring a side of an equilibrium.

Doesn't it mean that the favoured side is the side to where the equilibrium shifts?

Favoring side of equilibrium is nothing but the side whose concentration or pressure in the mixture will be higher on reaching equilibrium. Increasing temperature of endothermic reaction leads to the increase in yield. This can be understood from Le Chatelier's prinicple, which states that the reaction will proceed in a direction which will oppose the change enforced on the system.

Thermodynamics and kinetics play a huge role in deciding the favored side.

In general, a equilibrium will shift in a direction that opposes the stress that was placed on the system. For example, if you add heat to an endothermic reaction, you will receive more products, as this will lower the temperature of the solution. When you have the dissociation of an acid, the acid will dissociate more (act like an acid) when you make the solution more basic. An equilibrium will shift to minimize any change in a system, so when evaluating what will happen after a system is stressed, look to see which side of the equilibrium relives that stress, and that side will be favored.


text:
Can hydrofluoric acid etch glass at a pH of 7.0?
Will it still etch glass when very dilute? Will the reaction still proceed slowly? At what pH will the reaction stop completely?

From Wikipedia on hydrogen fluoride: "Although hydrofluoric acid is regarded as a weak acid, it is very corrosive, even attacking glass when hydrated."
On hydrofuoric acid: It etches glass "by reacting with silicon dioxide to form gaseous or water-soluble silicon fluorides. "

It seems fishtank glass can be etched even when the water is kept at a pH of 7.0.
"You might try a razor scraper, to see if that will remove any of it. If not, there is a very high chance they are due to etching. Etching is caused by the action of water, especially acidic water, over a long period of time. "

Your reference [3] is talking nonsense. HF is the only acid that will ever etch glass. And highly diluted, you won't see it in this century or the next.

Karl, HF is one of the compounds formed in fluoridated water.

In extremely low amounts, that are not able to etch glass. The subject of fluorinated water is full of conspiracy therories, you fell for one of them.

In $\ce{SiO2}$ etching, there are two active fluorine-containing species, $\ce{HF}$ and $\ce{HF2-}$.
The rate of etching at 25 degrees C in Angstroms per second is:
$$9.66[\ce{HF2-}] + 2.50 [\ce{HF}] -0.14$$
according to A Study of the Dissolution of SiO2 in Acidic Fluoride Solutions J. Electrochem. Soc. vol. 118, pp. 1772-1775.
The relavent equilibrium constants at 25 degrees C and ionic strength of 1M are
$$\frac{[\ce{H+}][\ce{F-}]}{[\ce{HF}]} = 0.0013M $$
$$\frac{[\ce{HF}][\ce{F-}]}{[\ce{HF2-}]} = 0.104M $$
Because of the -0.14 term in the rate equation, only at enormously high F- concentration would there be etching at pH 7 (such as in this situation).
This study was concerning high (~1M) concentrations of fluorine species, so it might not be completely relevant for low concentrations or very long time periods.

Thanks, but there was no mention of whether the reaction can continue at a pH of 7.0, as in the case of the fishtank.
Does the following from the abstract mean that as long as the temp is above 23 degrees C, the reaction will still proceed, though at a slow rate?
"the rate of dissolution is linearly dependent on the concentration of Formula and HF in relatively dilute fluoride solutions. Both reactions have a similar temperature dependence characterized by an activation energy of 9.1 and 8.1 kcal/mole, respectively."

What units is the rate of etching expressed in? m/s? mm/year?

It appears to be in Å/s.

DavePhD - What would be the rate given a fluoride concentration of 1ppm? How does HF keep reacting with SiO2 at a pH of 7.0? Is this a unique property of HF, or do all acids keep reacting at a pH of 7.0?

@user1491819 Accord to the article, the rate would be zero if the concentration is that low, although they aren't really studying concentrations that low.

@DavePhD can you think of any other plausible ways the fishtank glass could have been etched - (if indeed it actually is)?

What you observe is, unless you kept your fish at pH > 10 for a prolonged period, widely known as "limescale".

Certain acids are corrosive at a pH of 7.0. These compounds do not involve H or OH, so are not neutralised when diluted in water to a pH of 7.0.
"SiFW (silicofluoride treated water) is associated with serious corrosion of lead-bearing brass plumbing, producing elevated water lead (PbW) at the faucet."
"Silicofluorides (SiFs), fluosilicic acid (FSA) and sodium fluosilicate (NaFSA), are used to fluoridate over 90% of US fluoridated municipal water supplies. ...New data refute the long-prevailing belief that PbW contributes little to children's blood lead (PbB), it is likely to contribute 50% or more."[1]
[1]: Journal of Neurotoxicology(2007): Confirmation of and explanations for elevated blood lead and other disorders in children exposed to water disinfection and fluoridation chemicals http://www.ncbi.nlm.nih.gov/pubmed/17420053


text:
For alkali and lanthanide metals, is there a better liquid barrier to oxidation than mineral oil?
Most often mineral oil is used to store, protect reactive metals from oxidation. By far the best choice is to hermetically seal the metal in pure argon or some other inert gas, since, although mineral oil tends to slow down the oxidation process, it's still soluble to oxygen and slowly but surely the metal will begin to oxidize.
While argon is a sure thing it requires special equipment, and once sealed you have to break the container to access the metal.
My question - among liquids is there a better choice than mineral oil? A liquid that will not react with the metals, and provide a better barrier to migration of oxygen?

Not really as all liquids will, to some extent, dissolve a small amount of oxygen or water. The only liquid that comes to mind that will prevent either is mercury, but you probably don't want your stuff contaminated with mercury. If you really wanted to prevent water contamination, adding molecular sieves might help, but other than that, no not really.


text:
Is ammonium a strong or weak conjugate acid?
Ammonium, $\ce{NH_4^+}$, is the conjugate acid of ammonia, $\ce{NH_3}$. I have searched the Internet and so many different answers pop up. The rule of thumb that I read was that a strong acid has a weak conjugate counterpart.
For example, a strong acid has a weak conjugate base. Is ammonium a strong or weak acid? Is ammonia a strong or weak base? Is there something to the rule of thumb that I mentioned?

Related: Strong and Weak Acids/Bases

Both ammonia is a weak base and ammonium ion is a weak acid.  Many, even most, acid/base conjugate pairs are like that.  We should be using comparative instead of absolute adjectives in the rule about conjugate acid-base strengths:
A weaker acid has a stronger conjugate base, not necessarily a totally strong one.
A weaker base has a stronger conjugate acid, not necessarily a totally strong one.

Stronger in respect to what...the weaker acid we have?

@Hydrous Caperilla, yes.

@HydrousCaperilla actually if I understand your question, the phrase "stronger conjugate base" does not actually mean stronger than the weaker acid.  What the first statement means is: If you have acid A1 and its conjugate base B1, and you have acid A2 and its conjugate base B2, then if A1 is a weaker acid than A2, the conjugate B1 will be a stronger base than B2.

Consider these reactions:
\begin{equation}
\ce{HA + H2O <=> A- + H3O+}\tag{1}
\end{equation}
\begin{equation}
\ce{A- + H2O <=> HA + OH-}\tag{2}
\end{equation}
The equilibrium constant for reaction (1) is called "acid dissociation constant", $K_{\rm a}$, and the second is "base association constant", $K_{\rm b}$. (Note that the names are there for historical reasons, they're not correct, strictly speaking) We can factor out the water's concentration to a good approximation in dilute solutions. Consider that
\begin{equation}K_{\rm a} \times K_{\rm b} = \frac{\ce{[A- ][H3O+ ]}}{\ce{[HA]}} \times \frac{\ce{[HA][OH- ]}}{\ce{[A- ]}} = \ce{[H3O+ ][OH- ]} = K_{\rm w}\tag{3}\end{equation}
$K_{\rm w}$ is constant in constant temperature, hence $K_{\rm a}$ and $K_{\rm b}$  are inversely proportional to each other. Again, to a good approximation, we have
\begin{equation}-\log{K_{\rm a}} + (-\log{K_\rm b})=-\log{K_\rm w}\: \Longrightarrow {\rm p}K_{\rm a} + {\rm p}K_{\rm b} = 14\tag{4}\end{equation}
p$K_{\rm a}$s between -1.7 ($\ce{H3O+}$/$\ce{H2O}$) and 15.7 ($\ce{H2O}$/$\ce{OH-}$) are said to be those of weak acids in water. Acids with $K_{\rm a}$s in this range do not dissociate fully in water, and named weak acids (and weak bases, for that matter1) in water.
Now, think of it like this: $x+y=14$. If $x$ is 5, $y$ is 9, and if $x$ is -5, $y$ is 19. (4) is where your rule of thumb originates from. So

A weak acid like acetic acid with a p$K_{\rm a}$ of 4.76 will have a weak conjugate base with a p$K_{\rm b}$ of $14-4.76=9.24$.
A strong acid like $\ce{HCl}$ with a p$K_{\rm a}$ of -7 will have a conjugate base with a p$K_{\rm b}$ of $14-(-7)=21$.

Thus as Oscar's answer points out, it's inaccurate the way your guideline is phrased, but it's naturally following if we use comparatives for phrasing. I'd dare say most conjugate bases of weak acids mentioned in textbooks are weak ones.

1: As the comment points out, the weak base range is p$K_{\rm b}$>1. That slipped.

The definition of "strong" as having $pK$ below -1.7 may not be entirely true.  In https://en.wikipedia.org/wiki/Base_(chemistry)#Strong_bases a "strong base" is considered one whose conjugate acid has $pK_a>13$, so the strong base needs merely $pK_b<1$ not $<-1.7$.

@Oscar Noted; that somehow slipped through.


text:
How to separate benzyl alcohol and meta-cresol
Benzyl alcohol and meta-cresol boiling point's are not so different, the fractional distillation will be very difficult, what can i do for separate this two composites ?

I guess extraction with NaOH in water would be OK

Your observation on the boiling points is correct.
Take a look at the melting points instead: the m.p. of benzyl alcohol is -15 °C, while the m.p. of m-cresol is 11 °C. You might want to cool down the mixture and see whether you can see some enrichment.