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Egg-Detection in Pastries: An Analysis in Heat-Alterated Molecule Identification
I would like to know if anybody here knows of a method to detect the presence of ovalbumin--or any unique, egg-related molecule, in a baked good. Here I am anticipating that the "unique egg-related molecule" might alter under high temperatures, and so the suspect molecule for detection may or may not be any molecule which is present in a fresh, uncooked egg. I'm thinking of something that can be brushed on to a pastry to evolve a color change. Any help is appreciated.

"" in a baked good. "" "Baked good", is that some cookie, bread or what ever?

Any morsel which may be available at any given cafe that has been baked. Particularly detection of whatever unique molecule that is generated in the aftermath of some temperature-induced egg-associated protein transformation. ^_^

There are a number of commercial kits available to determine the presence of egg in commercial products and they are based on a technique called sandwich ELISA (enzyme linked immunosorbent assay).  A recent article in Food Research International by Gomaa and Boye discusses the impact of thermal processing on allergen detection (in this case casein, egg, gluten and soy).  Egg recoveries with the kits are low, ranging from 48% down to 0.  It appears that in some cases detecting egg in thermally processed samples is either very difficult to do or the egg products are transformed significantly that they can no longer be called egg (I did not read the paper in enough detail to determine if the authors speak to this point).
The kits tested in this study were from the following companies: Morinaga, Ridascreen and Ridascreen.  I web search will reveal these products and additional information, but I'll leave that to the viewer as I do not want this question to be an endorsement or critique of any commercial products.
The alternative to these kits appears to be flow cytommetry, which has better recoveries, but does not fit the OPs conditions of economical for a mid-income business without access to a laboratory.

Egg white consists of mostly Ovalbumin. It is a protein, and so will be hard to distinguish from other proteins in a baked product. Egg white changes structure to a solid polymer on cooking. You could dissolve it in a mixture of detergent and hair perm, returning the Ovalbumin to its uncooked state, and then run it through a column to identify it.
There might also be biological assays you could run using enzymes/antibodies which would only interact with Ovalbumin.
I don't think there are any chemical means to identify a particular cooked protein easily.
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What chemicals are in tap water that gives it a slightly basic pH level
In chemistry, we recently did a lab where we observed the different ways of reading pH of several different substances. We used things such as red cabbage juice, litmus paper, pH paper and a pH probe. One of the substances we tested was water. When we used the pH probe, it came back with a reading around 8.0. When we tested the distilled water, it came back with a reading very close to 7.0. So what is different between distilled and tap water that causes the tap water to be basic?

This is a good question, but it is hard to answer, because the pH of tap water strongly varies with location. Moreover, water treatment centers will typically add certain chemicals to control the pH to some extent, but this too depends on the location your water came from.

Following from @michielm's comment, you should be able to call your water provider and ask how they treat your water and get an idea of what goes into it.

This is most likely a problem of masurement. pH with a glass electrode in weakly buffered systems like water mirrors a lot of the pH of the solution the glass electrode was stored before. One must rinse the electrode several times in the water before one con make a meaningful reading.

The tap water is likely "hard," i.e., contains some dissolved mineral salts, most likely (primarily) calcium and magnesium carbonates and bicarbonates. The anions of these salts are slightly basic, consuming $\ce{H^+}$ ions and thereby raising the pH. Carbonates (with the obvious exception of alkali metal salts) are only sparingly soluble in water, while bicarbonates are typically very soluble. A couple of other subtle factors may affect the pH:

Lower atmospheric pressure encourages the evaporation of dissolved carbon dioxide from water. Since dissolved carbon dioxide exists in equilibrium with carbonic acid, it would tend to lower water pH. If you live in an area with comparatively low atmospheric pressure, your tap water might have a lower concentration of $\ce{CO_2}$, which would favor higher pH.
At higher temperatures, bicarbonate decomposes to water, carbonate, and carbon dioxide. Carbonates are relatively insoluble, and the consumption of bicarbonate should drive forward the dissociation of any dissolved carbonic acid, so a higher temperature should have the effect of lowering pH.

So what happens when the same water is distilled?

@CodyGuldner, the evaporated water doesn't carry the mineral salts with it. The water that you collect from the condensing flask should be relatively pure, while the majority of impurities (namely, the mineral salts responsible for the basic pH) will remain behind as a residue in the boiling flask.

Wrong! So called "hard" water, containing calcium/magnesium hydrogen-carbonate is slightly acid! Guess why!

Water treatment plants typically buffer the pH of the water supply to be slightly alakaline, with a pH of ~7.5-8 normally (the standard is that it must be between 6.5 and 9.5) with various chemicals including Sodium Hydroxide (lye) or Sodium Carbonate (soda ash).
Take a look at this photo from Wikipedia which shows all of the chemicals added at this particular water purification plant.  From left to right: sodium hypochlorite for disinfection, zinc orthophosphate as a corrosion inhibitor, sodium hydroxide for pH adjustment, and fluoride for tooth decay prevention.)
For treatment systems which add fluoride to tap water, a buffer is further necessary - the most commonly used additive for fluoride is fluorosilicic acid.
The reason tap water is kept slighly alkaline is because acidic water will leach chemicals from the pipes which it passes through. This can be especially problematic for homes with older plumbing (i.e. lead/copper/zinc). It also increases the corrosion and wear on the pipes/fixtures it travels through.

Here is a story of a town whose Sodium Hydroxide pump (used to buffer the pH) malfunctioned, causing chemical burns in their town: https://www.wired.com/2007/04/glitch_poisons_/
As good doc explaining the interaction of the Fluoride and pH: http://www.cdc.gov/FLUORIDATION/pdf/pollick.pdf

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Making a saturated solution of Na₂CO₃
I need to make a saturated solution of $\ce{Na2CO3}$ and according to the available data I should add $22\ \text{g}$ in $100\ \text{ml}$ of water at $25\ ^\circ\text{C}$.
I wanted to make $300\ \text{ml}$ solution, which would require $66\ \text{g}$. But, even after adding $110\ \text{g}$ of the salt, no turbidity appears which might indicate saturation. So I believe that the solution is still not saturated.
However, when I kept the flask overnight, almost the entire salt precipitated out in large crystalline chunks. I am a bit confused as to what the problem might be. Has anyone faced this kind of trouble?

Can you site your source for the solubility of $Na_2CO_3$ as 22 g/100 mL?  I wonder if there is a better estimate.

@bobthechemist Ullmann's Encyclopedia of Industrial Chemistry (2007) gives a solubility of 31.3 g / 100 g solution at 32 °C. Etacude, an online resource, provides a solubility curve an a table for the temperature depencency. The solubility at 20 °C is given as 21.8 g / 100 g water.

$\ce{Na_2CO_3}$ forms several hydrates with decomposition point close to room temperature (30-40 Celsium). Moreover, the solubility of decahydrate decreases with temperature decrease, and solubility of anhydrous salt almost does not depend on temperature and is much higher than solubility of decahydrate.  The relevant pictures are googleable by words 'sodium carbonate solubility curve', see 'images' google service, not general search.
So, I believe you have prepared a solution of the salt at the higher temperature at day, but the salt mostly precipitated in form of decahydrate overnight, because the solution cooled significantly at night. Bonus points for this version if you have significant changes in temperature at your work during the nychthemeron.
Second version is that you indeed prepared seemingly oversaturated solution because it is not saturated for anhydrous $\ce{Na_2CO_3}$, but overnight it crystallized, because it is saturated for decahydrate. I'm a bit sceptical about such possibility, though, because significantly oversaturated solution crystallize even on slightest imperfections, and the still not-dissolved salt is a great collection of such imperfections.
In addition, if the solution stayed overnight open, it indeed may consumed some amount of atmospheric $\ce{CO_2}$ and partially transform into sodium hydrocarbonate with much lesser solubility.

One possibility that occurs to me is that some proportion of the carbonate anions underwent an acid-base reaction with the water to form bicarbonate, which is substantially less soluble (96g/L for $\ce{NaHCO3}$ vs. 215g/L for $\ce{Na2CO3}$ at 20°C according to the respective Wikipedia pages, meaning the carbonate salt is nearly twice as soluble in terms of molar ratios), and therefore would precipitate out. The acid-base equilibrium would be driven forward by the formation of said precipitate over time.
As for why no precipitate formed initially, even when the solution was supersaturated, I'm not sure. Normally, preparation of a supersaturated solution would require heating to increase solubility. Then, upon cooling, crystals will precipitate out if given nucleation sites around which to form and/or the solution is physically disturbed. The solvation of sodium carbonate is an exothermic process, so the process would raise the temperature of the solution, but I doubt that's sufficient to explain the unexpectedly great solubility that you observed.

The reason it crystallized was probably because the temperature of the solvent decreased, lowering the solubility level. Also solubility of many solutes are not linearly proportional. Lots of these are solubility curves, so adding 66g to 300ml would not be the same as 33g to 100 ml. When displayed in a graph, the amount of a solute dissolved in a given amount of solvent at the give amount of temperatures, is shown as a quadratic equation. It appears as half a parabola.

I think you meant "22g in 100ml" instead of "33g in 100ml".

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What is the difference between a mixture and a composed substance?
In my textbook it says, that a composed substance is a substance made of more than 1 type of atom. And It also says, in another chapter it says that a mixture is the grouping of more than one substance without any chemical reaction.
What is the difference between a mixture and a composed substance?

From the definitions in your textbook, it seems that a mixture is a combination of two or more substances that--though combined--maintain their chemical identifiability. Examples of mixtures might be:

Water/Alcohol Mixtures
Salt/Pepper Mixture
In pharmacy, pills, pastes/creams, and suppositories are mixtures

(there are many example of mixtures)
The consituent compounds might be separated with somewhat mild techniques, including dissolving in certain solvents, distillation, and crystallization.
Composed Substances (according to your textbook) are mixtures of two or more species that under go some chemical reaction. The species involved could lose most of their identifiability, becoming completely different compound(s).  Loosely speaking, chemical reactions can not be reversed by simple separation techniques.
(Of course, there is some arbitrariness here: particularly in what comprises "most" of their identifiability and what defines "simple" separation techniques.  Also, mixtures are actually held together with bonds, just like in composed substances.  They may be weaker, but even "weakness" of the interaction is a subjective concept.)

Both composed substances and mixtures have more than type of atoms.
However, there are some difference between them both.
1)composed substances=mix chemically, mixture=mix physically
Composed substances are usually mix together through chemical process just like burning, titration and redox reaction.
While mixture are mix together through physical process like dissolving and even putting together.
2)Composed substances and mixture properties
Composed substances have different properties with it original atoms.
i.e iron sulphide do not attarct to magnet although iron is magnetic.
While mixture have the same properties with its original atoms.
i.e. a pile of iron filling and sulphur powder. The iron will still attarct to magnet and the sulphur don't
3)Bond between atoms
In composed substances, the atoms are bonded with either ionic bond or covelent bond. i.e the hydrogen and oxygen atoms in a water molecule have colevent bonds between them.
In mixture, the atoms do not bond together. i.e oxygen gas, nitrogen gas, carbon dioxide gas in the air do not have bond between them as they flows freely in the air.
They are also a lot more differences between these both such as melting point, existing states of matter, conductivity of electricity and heat etc. I hope this helps.


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What is a chemical element?
I would like to know what is the difference between a chemical element and a type of atom? Can  anyone give the definition, please!

From a scientific standpoint, there is no difference.  Elements=Atoms=the 100 or so basic building blocks of chemistry.
This is purely conjecture, but I suspect that the term/concept "elements" pre-date "atoms" (at least atoms as we know them today).  I could see "elements" referring to the macroscopic description, e.g. a jar of carbon (graphite), but it may not have been realized back-in-the-day that these were aggregations of many carbon atoms.
"Atom" strikes me as quantum scale, and "Element" strikes me as something on the macroscopic scale, perhaps hearkening back to the pre-atomic days.  But really, there is no difference.

I agree that the distinction is mostly academic/linguistic. I suppose the one issue to clarify is the matter of isotopes; not every neutral atom of a given element is identical, and grouping by element only takes into account protons, so that probably bears mentioning.

At the chemical level, an "element" is "a type of atom"; the later is literally the simple definition of the former (with "type" being further clarified as a specific number of protons found in the nuclei of all atoms of that element). So for most purposes, the two are one and the same.
There is sometimes a distinction made in chemistry literature between a substance in its "molecular" form as opposed to its "elemental" or "atomic" form. "Molecular nitrogen" refers to the pure form of natural gaseous nitrogen, N2. "Elemental nitrogen" or "Atomic nitrogen" refers not only to the pure substance, but in its unbonded, monoatomic form. Most non-noble gases don't naturally exist this way, and in fact a few, like nitrogen, are far more stable in diatomic gas form than any other. However, when making stoichiometric calculations, it's an important distinction; a mole of "molecular" nitrogen is double the mass of a mole of "atomic" nitrogen, because each nitrogen molecule is two nitrogen atoms.
Further confusing things is that depending on who you talk to, the "elemental" form of a substance could be its "natural pure" form, the form the pure substance exists in at room temperature and pressure, or its "atomic" form as single unbonded atoms. If there is a difference, as there is with gases, then the term "elemental" can be very confusing.
And perhaps most confusing of all is that the definition of "pure" a lab chemist might use is different from the definition that a nuclear physicist would use, because of the existence of isotopes. While atoms of a specific element must have a constant number of protons in order to be that element, many elements can exist for a measurable length of time with varying numbers of neutrons. Some have multiple "stable" isotopes (which will exist indefinitely in their current state without decay); chlorine, for example, is stable with either 35 or 37 total nucleons. Usually (but not always), the chemistry of various isotopes is unaffected; where it matters, in the molar mass of reagents, an average of the "naturally-occurring stable isotopes", weighted by percentage in naturally occurring form, is used and is close enough for chemists' purposes.
However, what a chemist would call "pure chlorine" would be unusable for a nuclear physicist researching molten salt reactors, because of the differing behaviors of Chlorine-35 and Chlorine-37 when bombarded with neutrons. Chlorine-35 can capture a neutron to become Chlorine-36, which then degrades (with a rather long half-life of 300,000 years give or take a few thousand) into Argon-36 and Sulfur-36, neither of which are chlorine and therefore don't produce the halide salt that is necessary for the chemistry of the reactor design, and sulfur being highly corrosive to the alloys used for the reactor vessel and coolant loops.

I do not see how your answer is substantially different than mine, except for 2) your involvement of isotopes--which are never the issue when talking about "atoms" in the context of "elements"--these are solely based on proton count.


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Why can't rust form without water?
Shouldn't iron oxide be able to form without water? It is just iron and oxygen. I don't really understand what the dot followed by the $\ce{H2O}$ means either. I was reading on wikipedia, but I have a rather terrible understanding of chemistry.
What purpose does water serve in rust formation, and what does it mean for a molecule to be hydrated-how is that different then just being surrounded by water molecules?

We could explain it from an electrolytic point of view.
Iron has a certain tendency to dissolve in water according to the following equation:
$\ce{Fe (s) -> Fe^2+ (aq) + 2e^-}$
The release of electrons causes a small current to flow in the metal (Remember that metals conduct electricity). This turns the point where iron dissolves into an anode, and the region around this area becomes rich in electrons, a cathode. Now, we have to remember that water dissociates to small extents according to:
$\ce{H2O (l) <=> H^+ (aq) + OH^- (aq)}$
If it was an acid, the $\ce{H^+}$ would have quickly taken all the electrons liberated from the dissolution of iron earlier, but the concentration of hydrogen ions is not large enough in water, so that we get another reaction taking place at the cathode [The hydrogen ions actually 'gather up' near to the surface but cannot do anything, creating a thin protective layer of hydrogen ions around the cathode, but not strong enough to prevent further reaction]. This reaction involves oxygen and other water molecules:
$\ce{2H2O (l) + O2 (g) + 4e^- -> 4OH^- (aq)}$
Now this is the $\ce{OH^-}$ that reacts with the earlier liberated iron (or ferrous) ions:
$\ce{Fe^2+ (aq) + 2OH^- (aq)->Fe(OH)2 (aq)}$
Now more oxygen will react to oxidise the iron (II) hydroxide...
$\ce{Fe(OH)2 (aq) \xrightarrow[O] Fe(OH)3 (aq) \xrightarrow[O] Fe2O3 (s)}$
You probably can recognise the last product.
As the $\ce{Fe^2+}$ are consumed, more iron will get dissolved (Le Chatelier's Principle) and keep the whole process going.
Maybe a little picture to go along with this :)

Note that water is not required if the temperature of the iron and oxygen are sufficiently high. I do not know the precise temperature, but from experience with welding I can say that noticeable amounts of orange-colored iron oxide begin to form as steel approaches its melting point, which is around 1370 degrees C (2500°F). This high-temperature oxidation in air, without water, is similar to what we usually call charring or burning.

@RalphDratman That sounds accurate. Heat basically prompts iron to react with oxygen, first giving FeO, and with more oxygen in the fray, Fe2O3 (rust). We could say that's the actual way that rust is formed, and water acts as catalyst (makes the same reaction possible with lower energy (heat) requirement.

The dot means that the compound is a hydrate of the parent chemical; the parent compound forms a weakly-bonded structure with one or more water molecules per parent molecule, because either the hydrogen atoms of the water, or the partial positively-charged region of the water molecule caused by its "bent" molecular shape, is attracted to partially negative regions of the parent molecule. It's not a "full" traditional ionic bond but it works in somewhat of a similar way. In solid form, this mixture of the parent compound and water molecules forms a crystal structure like ice (it doesn't have to be a well-ordered crystal like you typically get with salts), with properties that differ from a form of the substance that has more, less or no water in its structure.
An everyday example of the difference between a dehydrated and hydrated compound is Portland cement, used in concrete. In its raw form, it's "dehydrated" calcium carbonate along with a few other components, made by heating limestone until the water is released from the rock's structure. The remaining solid forms a fine powder that can't hold much of a shape. Add the water back in while mixing concrete, and the water molecules are re-incorporated into a solid structure with the calcium carbonate to essentially reform a rock in the desired shape.
As far as why rust only forms in the presence of water, water's "polar" nature is at work here as well. Being a polar solvent, it has some affinity for electrons, which are easily given up by the iron as a transition metal, and are attracted to the hydrogen atoms in the water. If a hydrogen atom successfully "captures" an electron, it balances its own electric charge, and is "liberated" from the bond it has with the oxygen in the water molecule, instead pairing with another liberated hydrogen to form a diatomic gas. Now, this hydrogen gas molecule won't get far, because water usually has some "dissolved" oxygen gas in it, and there's more oxygen just beyond the water, waiting to re-oxidize this hydrogen gas back into water. The net result is a positively-charged iron atom (typically in the +2 or +3 oxidation states, having given up 2 or 3 electrons to the water, respectively), negatively-charged hydroxide ions ($\ce{OH-}$), and water. These form a mixture of iron hydroxides:
$$\ce{2(Fe - 2e^{-}) + 4(H_2O + e^+) + O_2 \\ \to 2Fe^{2+} + 4OH^- + 2H_2 + O_2 \\ \to 2Fe(OH)_2 + 2H_2O}\\$$
$$\ce{4(Fe - 3e^{-}) + 12(H_2O + e^+) + 3O_2 \\ \to 4Fe^{3+} + 12OH^- + 6H_2 + 3O_2 \\ \to 4Fe(OH)_3 + 6H_2O}$$
That second equation typically occurs "stepwise", as you'll notice that's quite a lot of electrons and ions floating around at once. Typically, iron(II) hydroxide is formed first, and the iron will then readily oxidize further to its +3 state, either taking on an additional hydroxide to form iron(III) hydroxide, or liberating a hydrogen that will oxidize to form water, resulting in the monohydrate of iron(III) oxide-hydroxide:
$$\ce{4(Fe(OH)2 - e^{-}) + 2H_2O + O_2 \to 4FeO(OH)*H_2O}$$
Iron(III) hydroxide will readily rearrange itself into much the same state (it's not so much decomposition, as the water stays weakly bonded to the oxide-hydroxide to form the monohydrate):
$$\ce{Fe(OH)3 \to FeO(OH)*H_2O}$$
Finally, two of these iron oxide-hydroxide molecules are generally considered to "share" an oxygen from one of the hydroxide ions between them, with the hydrogen instead bonding to the oxygen of the other hydroxide, reforming water as a hydrate (possibly with some of it being released):
$$\ce{2FeO(OH)*H_2O \to Fe_2O_3*xH_2O + yH_2O}$$
Various hydrate structures of this iron oxide exist, which result in different colors of the compound from dark brown to deep red to red-orange. Other oxides are also possible by more direct action of the oxygen on the iron itself, such as iron(II) oxide ($\ce{FeO}$). Heating rust releases the water and forces the iron oxides to rearrange, which can also produce the third oxide of iron, iron (II,III) oxide (known in mineral form as magnetite).
You'll notice that the water participates in the overall reaction without really being consumed by it to form the product; it's destroyed and then re-formed in equal amounts, creating intermediates in the process. As such, water is a catalyst for these reactions. You'll also hopefully notice that this "ideal" reaction chain is just one of the possibilities; "rust" is an ill-defined term in chemistry, with the everyday compound consisting of a mixture of various hydrates of the three iron oxides, all of which are produced depending on subtle localized differences in the temperature and ratios of water, oxygen and iron.

the initial reaction in water liberates H2 gas. Ferrous ion, hydroxide, oxide is further oxidized by dissolved or interface elemental oxygen to give mixed iron oxides and eventually hydrated ferric oxide. Water prevents the formation of an adherent protective oxide layer.

By definition, rust includes water in its structure. Its not rust if it doesn't contain water.
Rust is porous- it has holes so corrosion of iron keeps going a lot longer if it rusts rather it forms a straight oxide.
Iron forms oxides instantly which are more protective than rust- not a lot more. Hot iron in oxygen environments corrodes by forming iron oxide.

Water is the catalyst in the process of rusting, therefore you cannot get rust without water.
Taken from http://www.chemicalformula.org/rust:

The overall chemical equation for the formation of rust is
Iron + water + oxygen --> rust
$$\ce{4 Fe(s) + 6 H2O(l) + 3 O2(g) -> 4 Fe(OH)3(s)}$$
Iron(III) hydroxide, $\ce{Fe(OH)3}$ then dehydrates to produce $\ce{Fe2O3.nH2O(s)}$
  or rust

That doesn't help my understanding much-All you did was introduce a new word.

Let me put it this way - the reaction that causes rust requires iron, oxygen, and water.

The OP already knew that; the question was why.

What purpose does water serve in rust formation?"
First, one should ask what is the temperature of the water and the source of the water? Cold from a faucet implying possible oxygen content along with say free $\ce{Cl2}$ or even $\ce{NH2Cl}$, or natural aerated waters containing transition metal salts (including those of $\ce{Fe, Mn}$, some $\ce{Cu}$, …), along with possibly gases other than $\ce{O2}$ like $\ce{N2O}$ from nitrate (present in well water) decomposition, or perhaps boiled distilled water without any dissolved gases or minerals. The composition of the iron alloy, I will ignore and assume pure $\ce{Fe}$, along with a pH range for the $\ce{H2O}$ of between 6 to 8.
I will attempt a brief depiction of the chemistry which could actually occur in a few scenarios to give a closer perspective of real world chemistry. For natural waters, as a starter, the chemistry described below is certainly relevant. However, with the distilled water in open air contact with submerged powdered iron, I will expect a comparatively much longer induction period till any rust is observed per the reactions outlined above. This is due to the underlying electrochemical nature of the corrosion process and no electrolyte (from dissolved salts, but instead awaiting dust particles) and no dissolved $\ce{O2}$ or an acid (source of $\ce{H+}$ could arrive via air containing $\ce{CO2}$). So iron rusting proceeds much faster in those cases containing dissolved oxygen, an acid source (free chlorine), and high mineral content (electrolyte).
There is also a possible direct action of water on iron via the following path without any oxygen presence:
$$\begin{align}\ce{2 [ H2O &<=>  H+ + OH- ]}\\[1.3em]
\ce{Fe + 2 OH- &-> Fe(OH)2 + 2 e-}\\[1.3em]
\ce{2 [ H+ + e- &<=> H^. ]}\\[1.3em]
\ce{H^. + H^. &-> H2}\end{align}$$
$$\begin{align}\text{Net:} &&\ce{Fe + 2 H2O  -> Fe(OH)2 + H2}&&\end{align}$$
This could then be followed by the so called Schikorr reaction, which details the conversion of iron(II) hydroxide into the mixed oxide iron(II,III), also written as Fe3O4 (see Wikipedia https://en.wikipedia.org/wiki/Schikorr_reaction ).
Interestingly, there is also a possible action on any formed ferrous in water containing $\ce{N2O}$ in the further presence of cupric (see Chemical Reduction of Nitrite and Nitrous Oxide by Ferrous Iron, by J. T. Moraghan and R. J. Buresh,in SSSAJ, Vol. 41 No. 1, p. 47–50, 1976.).
The reason for mentioning an acid source is due to the following reaction converting ferrous to ferric in the presence of oxygen:
$$\ce{4 Fe^{II} (aq) + 4 H+ + O2 -> 4 Fe^{III} + 2 H2O}$$
(see http://corrosion-doctors.org/Experiments/rust-chemistry.htm)
So, depending on various sources of the water, in addition to the mechanism with water and oxygen described above, the chemistry could be a bit more complex.

"Rust" can form without water, if by rust you mean an iron oxide. Fe2 O3 ( the red oxide) can form in air at ambient temperatures. The only example I know is fretting corrosion where the characteristic red powder forms around two rubbing steel/iron surfaces. Of course you can make all sorts of iron oxides at elevated temperatures.

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Chemically Speaking, What's Bad About Inhaling Plastics?
I hear about things like dioxin, but I really don't see what's dangerous about plastics-shouldn't you be ble to just breathe it out, or pass it out of your system? How would it cause cancer?

Check out this information

The primary dangers from plastics themselves come when the monomer, and any compound put in to make the polymer matrix flexible, separate from the polymer and are released (leached) into whatever the plastic is containing. There are several such compounds which are known to be problematic to human biochemistry:

Phthalates - These are aromatic diesters, with a structure incorporating a phenyl ring with two ester linkages, typically to alkanes like ethyl or butyl groups. Phthalates are used in certain plastics to link the monomers, such as in poly-ethyl terephthalate (PET aka polyester), and in others as a plasticizer to make the normally rigid plastic flexible (PVC). Phthalates mimic certain hormones in the human body and so can cause issues with growth and maturation, as well as feminizing characteristics in males (reduced body hair, weight gain, moobs). PET is OK for packaging and storage of cold foods (it is, for instance, the plastic used for soda bottles), but you don't want to heat it because that will degrade the plastic to release the phthalate.
The ones used in PVC are especially problematic because they're not well-bonded; instead of being an integral piece of the polymer chain, they form weak hydrogen bonds between the polymer chains (thus the flexibility, even in relatively thick layers of the plastic), and so these phthalates easily leach out into water or "off-gas" into the air, even without heat. These phthalates are a big part of the "plastic smell" of new plastic products, and of water stored in them; the more flexible, and the longer the plastic has been kept enclosed, the more intense the smell. There was a big stink (NPI) in the early 90s when PVC was being used for water supply lines in home construction, and then they found all these problems with the stuff and a lot of homebuilders and chemical companies went bust paying for all this piping to be replaced with copper.

Bisphenols - These are variations of 2,2-bis-(dihydroxyphenyl)propane (which is the infamous Bisphenol A or BPA). Bisphenol A is the monomer for most of the original formulations of high-impact polycarbonate, such as Lexan. Again, when heated, the polymer breaks down and leaches the water-soluble monomer, which is structurally similar in many ways to phthalate and again has hormone-mimicking, endocrine-disrupting properties, which is why newer polycarbonates intended for food storage are made with other bisphenol monomers (most of which have similar endocrine properties, but because they aren't Bisphenol A they're still legal for now and allow the manufacturer to label the product "BPA-free").

Vinyl Chloride - The monomer of poly-vinyl chloride or PVC, this is typically a hazard in production facilities for PVC and when old, aged PVC begins to break down. The vinyl chloride monomer is a known carcinogen, which is the other reason PVC is no longer used for food packaging or water supply, and is being phased out from wastewater lines in favor of nylon.

As far as answering the question "shouldn't we just be able to pass it out of our system", the same could naively be asked of a whole host of chemicals that are bad for us. The biggest problem with that statement is the First Law of Toxicology; anything, in sufficient amounts, is harmful, even if it's absolutely essential for life in the right amounts. Selenium is an essential trace mineral in human biochemistry, required in certain protein synthesis mechanisms, but the LD50 blood concentration is about 3 mg/kg, and harmful effects such as stunted growth begin to be observed at concentrations within foodstuffs of as low as 4 to 5 ppm. Heck, you can die from drinking too much water, faster than you could from lack of it (it causes a condition called hyponatremia where the added water dilutes electrolytes to the point the body can no longer function).
What causes many chemicals to be harmful is in fact our body's own affinity for them as a beneficial chemical. Either the chemical is beneficial to us, or it looks like one that is, and the chemical is normally so rare in our diet and environment that our bodies grab all of it that they find, without an "off switch" to regulate intake.
Which leads to the second problem with that statement; sometimes our bodies don't efficiently metabolize and dispose of chemicals. Vitamin A, for instance, is an essential nutrient, and our body absorbs all it can, and gets rid of it only very slowly. As a result, it "bioaccumulates", and eventually becomes toxic. That's why beta carotene is the preferred "form" of Vitamin A in supplements; it's in fact an intermediate that our bodies can use to synthesize the vitamin as needed, but being water-soluble unlike actual Vitamin A, any unused excess of beta carotene is simply flushed out via the kidneys.
That process in itself can cause problems; the latest round of drug recalls the lawyers are running commercials on are for drugs that did exactly what they were supposed to, but were then metabolized into carcinogenic compounds that concentrated in the urine and caused various cancers of the excretory system.
And the last problem, tying back into the case in point, is that even if the body selectively absorbs it, efficiently processes it, and it isn't acutely toxic in any form or intermediate along the way, as long as there's an input of the chemical into your system, you'll always have some of it. Plastics are ubiquitous in consumer products, and especially food prep and packaging. As a result, virtually every resident of the industrialized world has a detectable amount of artificial phthalates in their blood, because it is simply impossible not to come into contact with them. Even USDA-certified organic foods have phthalates in them, because the cows eat food that's been stored or served in plastic bins and it shows up in their milk and meat, or the vegetable was fertilized with compost stored in a plastic bin or watered from a PVC irrigation system (even the farmer's market items from someone's backyard were likely grown using water collected in a plastic rain catchment and/or dispensed from a hose or a drip irrigator system that leaches plasticizer). What you store it in or prepare it with once it gets to you is immaterial; the damage is done.


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Specific latent heat of fusion of a noble gas
Could someone explain to me where the additional energy goes when applied to an object in the state of melting (such as water) so that the temperature of the object does not increase, and could someone please explain to me why Helium has a specific latent heat.

When you heat a substance (not at its melting point or boiling point), the energy supplied goes to increase the kinetic energy of the atoms/molecules in the substance.
When the same heat is applied to the substance at its melting point (or boiling point), the heat goes to increase the potential energy of the atoms/molecules in the substance. You could refer to the potential energy as the energy used to break the inter-particle bond. In ice, there are relatively strong forces of intermolecular attraction, which makes the ice hard.
Now, temperature is directly related to the kinetic energy of the atoms/molecules of a substance, and since the kinetic energy of the atoms/molecules does not increase during a change of state, the temperature remains constant.
At least, this is what happens at a macroscopic level and in the long term. But if you heat ice at one particular point (at particle level and for a short period of time), that point will melt faster and the temperature will tend to rise there, but this heat energy will then be distributed to the other parts of the ice still cold for melting.
It is only when all the ice has molten than the heat supplied will contribute once again to the increase in kinetic energy, that is, temperature.
And for the last part, every substance that undergoes a change of state has latent heat.

Thanks for your answer. Where exactly is the potential energy stored?

@PeterMorris Sorry for the late reply. Somehow I didn't get any notification... As mentioned in my answer, it goes into the intermolecular bonds and causes them to break. You need energy to break a bond and making back that bond releases back energy.


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what is a pseudohalogen/pseudohalide?
What do you mean by a pseudohalogen/pseudohalide and how do you tell whether a molecule/ion is a pseudohalogen/pseudohalide ?

see the wiki http://en.wikipedia.org/wiki/Pseudohalogen

Pseudohalides are "fake halides."   Their chemistry bears some resemblance to true halides ($\ce{F-}$, $\ce{Cl-}$, $\ce{Br-}$, $\ce{I-}$).
Examples of pseudohalides include $\ce{CN-}$ and $\ce{N3-}$. These are generally weak Lewis bases that bear a formal negative 1 charge.
In my opinion, it is rather subjective as to what can be labeled as a pseudohalide. In general, it is in the eye of the beholder, and for a certain set of chemical circumstances.
In short, there is no "dividing line" for being a pseudohalide. If we assume that halides undergo certain (characteristic) types of chemistry, and the species X in question also generally does those same types of chemistry with a similar mechanism, then you may be able to label X a "pseudohalide."
(Note that within the true halides themselves, they can have very different characteristics, and so there is a bit of a moving target for what constitutes a pseudohalide.)

But F-,Cl- tend to make low-spin complexes and CN- ion tends to make a high-spin complex. Now considering this difference isn't 'pseudohalogen' a misnomer ?

Yes, it is probably a misnomer in the context of coordination chemistry, that you cited here. But, it may have certain aspects of acid/base chemistry that look like halogens.  BTW, I think it's kind of a silly concept: pseudohalides.  It's just asking for a fight as to whether group X is enough like any/all/some halides in all reactivities/spectroscopies to be called "pseudohalide."  Truth is, X is a halide or it is not.

A pseudohalogen is an anion with similar properties to a halogen, but that isn't a halogen. The true halogens are of course the group one column to the left of the noble gases; fluorine, chlorine, bromine, iodine, and astatine, as well as the theoretical ununseptium.
The pseudohalogens are generally -1 compound anions with stable structure (typically featuring multiple bonds) that form compounds with similar reaction behaviors and properties as halides, and can be used as general substitutes in reactions with similar mechanisms. These properties include being weak bases by themselves that form strong acids when protonated, forming neutral salts in combination with metals, and oxidizing electropositive elements. Common pseudohalogens include cyanide/isocyanide, cyanate/isocyanate, azide, their sulfur analogs such as thiocyanate, and complexes such as tetracarbonyls.
However, pseudohalogens are not always identical. Take for instance sodium chloride, and the pseudohalogen analog sodium azide. While the former is relatively harmless table salt, consumed in tens of grams every day, the latter is so generally toxic to life that biochem labs use it as a reference poison (and to make sure whatever nasties they've been growing in vitro for an experiment are good and dead before disposing of the medium). Similarly, while the human body needs a decent dose of chlorine anion as an electrolyte, and smaller amounts of iodine and even bromine are essential trace elements, most of the pseudohalogens are quite toxic, including cyanide, azide and carbonyl complexes. Thiocyanates and isothiocyanates are essential functional groups and anions in human biochemistry, so not overtly toxic like azide, but - First Lesson in Toxicology - too much of anything is bad.

Can one say that the diatomic interhalogen compounds ,eg. ClF are pseudohaogen compounds ?

I would say not, first because interhalogens are typically not ions (they end up sharing their lone electron forming an electrically neutral compound) and because this structure typically isn't preserved in any reaction; the two halogens would much rather oxidize anything else than each other. However, there are halides of pseudohalogens, such as chlorine and fluorine azides. While similarly reactive, the stable pseudohalogen group is typically preserved in an exchange or oxidation.


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How to detect harmful gas occurence in carboy bottle
To fill with drinkable water, I have lots of empty 19lt carboy bottles. Since they came from different people and companies, I have to check their healthiness to decide wheter they can be sterilise for reuse or not.
I need to detect if a 19lt carboy bottle was used previously for gas, lpg, diesel etc. or is there any other harmful chemical (such as urine) in it.
What could be the cheapest and trustable solution for it? Could I use gas sensors or infrared imaging or something else?
Thanks in advance.

I agree with Jerry. But I also wish you moved that, not closed. It's interesting and very useful in real life.

What is the material of the bottles?

Actually I ain't know what the material exactly is. As far as I know, the material is carbon based. I am sure that you are familier with this kind of bottle. You can find a sample image here: http://cdn.tbplatform.com/company/727/products/4392/600-carboy_drinking_water_19_lt.jpg

some information on detection of organic compounds can be found here http://www.epa.gov/oust/pubs/esa-ch6.pdf includes detailed explanations and cost assessment. note that costs are for large batches. equipment will probably be expensive.

Hi Daniel, That link explains lots of things. Thank you very much.

Gas sensors are probably too specific. Infrared imaging is going to have to take account of the carboy material itself. However, sampling the air and performing a gc/ms analysis could work - basically its headspace analysis.

Thank you very much Epicentre. Since I am not familiar with these techniques and devices, I can't decide it is suitable for us or not. Could you please guide me for these devices and techniques. Especially cost is the main parameter. Thanks again.

@Pertinaks I suggest you contact a chemical analysis laboratory in your general area to find out the cost. It will probably be considerably more than the cost of the bottle itself. http://en.wikipedia.org/wiki/Gas_chromatography and http://en.wikipedia.org/wiki/Headspace_technology may help you understand the technique.


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FCI implementation
Could you suggest, please, a manual/textbook/tutorial for implementing FCI (full configuration interaction) method on computers. I am interesting in algorithms mainly. I would like to write my own code, so I want to study how one can provide convenient indexing for configurations and what tricks are usually used to decrease computation time.

MOLPRO has a full CI module (by Peter Knowles), but it also costs quite a bit of money.  You can also do full CI in GAMESS (I think Ruedenberg's implementation), and perhaps any MCSCF/CASSCF program that is efficient enough to include all orbitals in the active space.
I regret to inform you: you may be in for a lot of work if try to write your own code.  :-(  This is an area of very deep, highly mathematical and numerical concepts.
So as far as a manual/textbook:  You may want to try Molecular-Electronic Structure Theory, especially Chapter 11 and 12.  There is just a crazy amount of work that you will have to do prior to CI, including the coding of the various four-center integrals for electron repulsion, kinetic energy, and overlap (earlier chapters).

Molpro is not nearly as expensive as truly commercial packages like Gaussian.  As one can see on http://www.molpro.net/info/products.php?portal=visitor&choice=Licence+types, a single workstation license costs $350/$600 for the serial/parallel version.  That's about the same price as Microsoft Office was the last time I looked (close to 10 years ago).

I happen to have a simple FCI code with detailed documentation of implementation. You can find it at http://scholar.princeton.edu/boxiao/software/fci-201302

Useful links include:

http://vergil.chemistry.gatech.edu/notes/ci/node23.html
http://www.sciencedirect.com/science/article/pii/0009261489874640
http://scitation.aip.org/content/aip/journal/jcp/98/4/10.1063/1.464087
http://www.sciencedirect.com/science/article/pii/0010465589900337
http://www.sciencedirect.com/science/article/pii/0009261488874128
http://www.sciencedirect.com/science/article/pii/000926149085633N
http://www.sciencedirect.com/science/article/pii/000926148485513X
http://www.sciencedirect.com/science/article/pii/0301010480800450
http://www.tandfonline.com/doi/abs/10.1080/002689798168303#.UmRKqJRgbVk

And there is the more recent Monte Carlo FCI method:

http://scitation.aip.org/content/aip/journal/jcp/132/4/10.1063/1.3302277
and many more courtesy of Google Scholar


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How does the volume of a mole of ideal gas relate to pressure and temperature?
The volume of a mole of gas is:

$\ce{22.4 l}$
directly proportional to pressure and Kelvin temperature.
directly proportional to pressure, inversely proportional to Kelvin temperature.
inversely proportional to pressure, directly proportional to Kelvin temperature.
inversely proportional to pressure and to Kelvin temperature.

I said the answer is 3 because the ideal gas law is $PV=nRT$ and "proportional to" is the same as saying which ones are on opposite sides. So $n$ and $P$ are on opposite sides and temp is on the same side as $n$.  But I'm wrong and the answer is not 3, but I don't know why.

Your starting point of reasoning is good, but you misread the equation. The question is about the volume of 1 mole of gas so let's rewrite the ideal gas law to give this relation:
$$P V = n R T \rightarrow \frac{V}{n}=\frac{R T}{P} $$
Now we have the volume of a mole of gas $\frac{V}{n}$ on the left hand side so we can see the proportionalities on the right hand side.
The volume of a mole of gas is inversely proportional to the pressure and directly proportional to the temperature, i.e. 4 in your list of answers

According to Boyle's law, at constant temperature, pressure is inversely proportional to volume.
According to Charles law, at constant pressure, volume is directly proportional to temperature.
Combining the above and the Avagadro's law, we get the ideal gas equation.
Hence, the correct answer is 4.


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Forms of Henry's law versus ideal gas law
This might be a slightly daft question as I'm a physicist rather than a chemist, but I have a slight problem in using Henry's law which I think I can circumvent using the ideal gas law – I'd be grateful if anyone can confirm or deny my logic!
I have an equation that yields the volume of oxygen gas per unit mass of tissue. We'll call this quantity $C$ with S.I units of $\mathrm{m^3\: kg^{-1}}$. We also have the partial pressure $p$ measured in $\mathrm{mmHg}$ and we wish to relate the two; several texts have suggested using Henry's law of the form $C = kp$ where $k$ is Henry's constant for oxygen. However, this causes a problem, as I can find no form of Henry's law with the correct units to allow this conversions (units of $\text{vol} / ({\text{pressure} \times \text{mass}})$).
My first question is does such a series of constants exist and does anyone know where?
In the interim, I decided to try to circumvent this problem by using the ideal gas law, which I believe holds at the relatively low pressures in very shallow water. Knowing one mole of oxygen gas has a mass of $32\:\mathrm{g}$, and occupies a volume of $22.4\:\mathrm{l}$ at standard temperature and pressure of $760\:\mathrm{mmHg}$, I re-wrote the expression for the mass of $\ce{O_2}$ in a given volume;
$m_{\ce{O_2}} = \frac{0.024\,p}{(0.032)760RT} = \frac{0.7\,p}{760RT}$
Using the ideal gas constant, and a temperature of $310.15\:\mathrm{K}$, I get
$m_{\ce{O_2}} = 3.572 \times 10^{-7} \cdot p$
If I define $M_{\ce{O_2}}$ as the unitless mass fraction of $m_{\ce{O_2}}$ per unit mass, then knowing the density of oxygen gas is around $1.331\:\mathrm{kg/m^3}$, I can divide this into $M_{\ce{O_2}}$ to get the expression
$C = 2.6835 \times 10^{-7} \cdot p$
which has the correct units. My second question is then is this valid, and if not, why not?

I wrote out an answer to this but realized I was wrong. But I've noticed a number of things:
First off, your form of the ideal gas law has the Molar Mass inverted. If you make the substitution that $n=m/M$, then the rearrangement reads $m = MVp/RT$, you have the $M$ on the bottom.
Secondly, You are selecting a Volume of 1 Liter, so your calculation will solve for a relation of mass and pressure @ 1L, which I don't think you want.
Thirdly (?), I'm guessing that the 760 in the denominator is there so you can use a gas constant in terms of atm? Or did I miss something?
So those are my observations about your attempt with the ideal gas law. I also am concerned over the bit with the mass fraction, as I think you converted a tissue mass into an $\ce {O2}$ volume using $\ce {O2}$'s density, which would be wrong if that was the case.
Looking again at Henry's law in the $p=kC$, you can see that in order to use the standard k units available, you would need to convert the $C$ you have from $L/kg$ to $mol/L$. You can convert $L$ to $mol$ using the ratio you used earlier, and you end up with $C$ in terms of $mol/kg_{tissue}$. The trick here is that you need the density of the tissue (not of $\ce {O2}$) to convert this into the volume you're looking for in order to use the equation. I think haha.

Thank you! I think you're right; here's how I did it; I multiplied my C by the density of oxygen gas and density of tissue which I assumed was the same as water. Now C was in the correct units, I could use Henry's law; I corrected the O2 constant to 37 degrees and voila, it worked. Thanks for that!


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What is the substance in the new water cooler filter that give bubbles and taste a bit like soap
Recently the water cooler was replace with a new filter.
However, when I refill my water kettle with the water from the water cooler, there are many bubbles and the taste was a bit like soap.
What chemical cause such effects? (Is it related to the new filter?)

The most likely cause of a "soapy" taste and sudsing from filtered water is a reverse osmosis filter, for which the usual membrane preservative used to keep the filter from degrading before its first use is plain old glycerin, a common ingredient in and byproduct of traditional soap (though most cleansing products nowadays, even those marketed as soaps, are technically anionic surfactant "detergents", based on alkyl-sulfates). The glycerin is hygroscopic, meaning it will attract and hold water on the membrane surface to keep it from drying out.
An activated carbon filter sometimes has a similar preservative to improve the material's shelf life, or else a binder that allows the manufacturer to mold a block of the carbon rather than loading loose powder into filter housings, which reduces carbon dust in the factory (a major fire and health hazard of these types of plants).
In all cases, the manufacturer usually recommends flushing any new filter with between one and five gallons of source water to wash out these substances, along with any loose filter media or other byproducts of manufacturing. This obviously didn't happen in your case. The glycerin in the water won't kill you; it's an FDA-approved food additive used as a sweetener and "plasticizer" (it's used in cake frosting for instance, to avoid it hardening after it sets). In sufficient doses it's an irritant to the mucous membranes and the GI tract, which can cause nausea and diarrhea, but you'd have to eat almost a kilo of the pure stuff for it to become dangerously toxic.


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Parts of a manometer
There is a poor quality photo in the diagram of a manometer below:

The picture is too poor for me to understand what they mean by "glass pointer marks calibrated volume".  I would be interested if anyone (a) knows what this is - preferably with a manufacturer's number so I can inquire further or (b) has a higher quality image of the part.

I can't tell you a manufacturer, but I can tell you the purpose of the pointer.
When the height of the column of mercury changes, it changes the height of the mercury exposed to the atmosphere. Since the ruler is (usually) fixed in place, the level of the mercury needs to be adjusted to the zero on the ruler. The pointer shows you the zero point. You adjust the level of mercury so that it barely touches the point.


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Crystallization from a solution
I was wondering if there is any other way to crystallize materials from a solution other than using supersaturation?
Thanks,

No. Per definition you need to have at least local supersaturation. You can archive it by different means (cooling, evaporating of solvent, mixing with other solvent...) but in the end you need supersaturation. See: http://xray.chem.ufl.edu/growing%20tips.htm or http://web.mit.edu/x-ray/cystallize.html
Of course you can use techniques like sublimation but your question was about solution (btw in case of sublimation you also have supersaturation, only difference is that your solvent is gas).

Not really. By definition, if some solute is in a solution below the saturation point at that temperature, it's going to remain in solution. The only way to get it out is to lower the temperature or evaporate some of the solvent to increase the concentration.
If you don't care especially about the integrity of the chemical, you could always precipitate it with some other reagent, but then of course your crystals would not be of the chemical you started with, but perhaps could be purified through some other means...


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A few questions about the conditions of the Diels Alder Reactions
I'm attempting to extract the diels alder reaction into a generic reaction and am trying to gain a better understanding of it.
Can a conjugated diene react with another conjugated diene via the diels alder reaction? And is the reaction limited to only dienes or can it be expanded to polyenes?
Does anyone happen to have a link as to the details on hetero diels alder reactions and the structures that can undergo such reactions and limitations on chemical species that can react via this reaction.

Yes, conjugated dienes can react with other conjugated dienes in Diels-Alder reactions. A well known example of this is the dimerization of cyclopentadiene. All the standard caveats, restrictions, and other considerations still apply, namely:

The diene must not be rotationally locked in such a way that it cannot adopt an s-cis/synperiplanar conformation. As a logical corollary, cyclic dienes that are already locked into the correct s-cis conformation often have lower activation barriers.
The kinetics of the reaction are accelerated by electron-withdrawing substituents on the dienophile and electron-donating substituents on the diene.
The stereochemistry of the major product is highly substrate-dependent. A preference for the endo isomer exists in cases where the dienophile has unsaturated substituents. In those cases, the transition state is stabilized by an overlap of the $\pi$-electron system between said unsaturated substituents and the $\pi$ orbitals of the newly forming double-bond at the back of the diene. (Just as a side note, Carey & Sundberg refer to this as the "Alder rule" in Advanced Organic Chemistry, but I've never seen that term used anywhere else.)

In principle, there's no reason why two conjugated $\pi$-bonds of a polyene cannot react similarly, provided they satisfy the orbital symmetry and conformational requirements and the reaction kinetics are favorable. Even aromatic molecules can undergo Diels-Alder reactions under the right set of conditions (see, for example, the Diels-Alder reaction between benzene and furan via a benzyne intermediate). There are also numerous instances of well-documented Diels-Alder reactions involving heteroatomic dienophiles, alkynes as dienophiles, polymerization reactions, etc.
If you can get your hands on the aforementioned Carey & Sundberg book, they discuss the Diels-Alder reaction at length (including relatively rigorous explanations of the stereospecifity and reaction kinetics via molecular orbital theory). If you're looking for a comprehensive reference, I doubt there's anything superior to the venerable March's Advanced Organic Chemistry. There's a chapter on Diels-Alder reactions (along with virtually every other organic chemistry reaction known to science), as well as a massive list of accumulated primary literature references if further detail is required.

I'm speaking from memory, so this could be a little loose, but I believe dienes react just as well with other dienes via the diels alder reaction as with simple alkenes (of course, only attacking one of the conjugated double bonds). As for polyenes, I believe the same is true, but I'm not certain if there are added complexities there, and it would probably have to be evaluated on a case by case basis.
The reason for the disclaimer is that it's possible that the configuration of double bonds may or may not lock in the polyene into a planar or partially planar structure, in which case depending on how many other double bonds and their locations in one or both of the reagents could interfere with the traditional diers alder reaction. But as to your question, yes, you can apply the reaction to a general case as far as I know.


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Esterification of a carboxylic acid and the production of water
In an esterification reaction between a carboxylic acid and an alcohol, a carboxylate ester and water is produced.
I found the following diagram online, and my textbook has a similar one as well

Question: In forming the water in an esterification of a carboxylic acid, why does the hydrogen ion come from the alcohol and the hydroxide come from the carboxylic acid and not the other way around?
I think it might be because carbon's bond to the OH- is weaker than the R bond to the H+ ion and requires less energy to remove those two ions from their respective compounds.

The mechanism of the reaction doesn't involve the direct exchange of ions in the way that, e.g., a metathesis/double replacement reaction might. Take a careful look at the reaction mechanism from the Wikipedia page. Notice that one of the intermediates (the product after step three) is a tetrahedral complex with two hydroxyl ($\ce{-OH}$) groups. One of those hydroxyls is indeed the original hydroxyl of the carboxylic acid moiety, but the other one is actually formed from the protonated oxygen of the carbonyl. Either one of them could, with equal likelihood, be protonated in a subsequent step to provide the water molecule that serves as a leaving group. In the image that you linked, it's convenient to circle those specific atoms, but that shouldn't lead you to believe that the mechanism of the reaction is in fact a simple exchange of ions.
The question that you didn't ask directly, but perhaps implied, is: why is protonation of the alcohol by the carboxylic acid, followed by nucleophilic attack by the carboxylate on the carbon of the protonated alcohol, not a viable mechanism? First, a protonated alcohol is more acidic than a carboxylic acid by six or seven orders of magnitude, which means that the initial proton transfer would be slow and unfavorable. Second, the resulting ester is apt to be hydrolyzed eventually (albeit slowly without catalysis, but then the whole reaction proceeds very slowly), yielding the original carboxylic acid and alcohol.

It should first be noted that the circle drawn derives from experimental results. If you were to mark specific oxygens, e.g. by introducing radioactive $\ce{^18O}$ in place of normal $\ce{^16O}$, you can predict whether and where the ester product will contain $\ce{^18O}$ or not and whether the water will contain $\ce{^18O}$ or not.
$$\begin{align}\ce{R-CO-OH + HO{*}-R &<=>> R-CO-O{*}-R + H2O}\\[0.5em]
\ce{R-CO{*}-O{*}H + HO-R &<=>> R-CO{*}-O-R + H2O{*}}\end{align} $$
Why this happens is due to the reaction mechanism. The mechanism’s first step is a nucleophilic attack of the alcohol oxygen to the $\ce{C=O}$ double bond to give a tetrahedral intermedate with $\ce{O-}$, $\ce{OH}$ and $\ce{HOR+}$ all bound to a single carbon. This breaks down by the $\ce{C=O}$ bond being formed and one group being expelled. If $\ce{HOR+}$ is expelled, the starting materials are regenerated. If $\ce{OH-}$ is expelled, then after proton transfer the ester product is generated.
Within this mechanism, there is never an attack on the alcoholic $\ce{C-O}$ bond so it cannot be broken. Furthermore, it is quite hard to break a $\ce{C-O}$ bond under displacement of $\ce{OH-}$ anyway. The only reason why it works in ester formation is that at the same time $\ce{C=O}$ is regenerated which provides the required energy.

Esterification involves ANDN mechanism which is basically attack of one nucleophile and the departure of another from the same molecule.here OR- group attack the carboxylic acid at the carbonyl carbon causing OH- group to leave. OH- will combine with the H+ ions present in the system to form water.


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Identifying Reaction Types
I need help on identifying precipitation, acid-base, and redox reactions. I know what each one is in theory:

Precipitation: Creating a solid
Acid-Base: neutralization
Redox: exchange of electrons

But given a chemical reaction like:
HCl (aq) + NaOH (aq) → H2O (l) + NaCl (aq)
I need to understand how to classify them. Are there any tell-tale signs to look out for when determining. For example, does a precipitation reaction ALWAYS produce a (s) as a product and therefore any reaction with a (s) on the product side is a precipitation reaction?

They are fairly simple concepts, check your textbook again, or Wikipedia.

It's true that a precipitation reaction will always yield at least one solid product, but not every reaction that yields a solid is, strictly speaking, properly categorized as a precipitation reaction. Precipitation reactions are ones in which at least one of the reactants is in the aqueous phase (i.e., dissolved in water), and a solid forms on the product side which was not present on the reactant side. If all the reactants are solid, then a solid product forming should not be called a precipitate. To be called a precipitate, an insoluble product must form from within solution, either from a solid and solute interacting, or from an interaction strictly between solutes.
Examples of precipitation reactions:
$\ce{AgNO3_{(aq)} + NaCl_{(aq)} -> AgCl_{(s)} + NaNO3_{(aq)}}$
$\ce{CuCl2_{(aq)} + Zn_{(s)} -> ZnCl2_{(aq)} + Cu_{(s)}}$
Notice that at least one of the reactants was in the aqueous phase, and a new solid chemical species is produced on the product side. $\ce{AgCl_{(s)}}$ and $\ce{Cu_{(s)}}$, respectively, are the precipitates. Incidentally, the second reaction is also an example of a redox reaction, since it involved a change in the oxidation states of the reactants (here, zinc is oxidized while copper is reduced).
Identifying acid-base and redox reactions is a lengthier topic, so I suggest you consult a textbook or any of the numerous online resources.

so a reaction can be more than 1 type? In the 2nd example, it is both a precipitation and redox.

@Richard, yes, categories of chemical reactions are not necessarily mutually exclusive, and there are countless reactions that can be reasonably classified in more than one way. Combustion reactions, for example, are also redox reactions. Combination and decomposition reactions may or may not be redox reactions. Which term is used depends on context.


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What is the condensation point of water?
The boiling point of water is $\pu{100^\circ C}$. However, we find that even at room temperature, water can evaporate into gas.
Therefore, water vapor can exist at temperatures of, say, $\pu{50^\circ C}$.
At what temperature, can you say for certain that all water vapor will turn into liquid?
In other words, what is the condensation point of water?

AFAICT, never: Evaporation occurs at all temperatures.

If I'm not mistaken that is technically not right (although this is more semantics than anything else). Once water has become ice (under normal condition $<0^o$C), the process of ice turning into gas is not evaporation but sublimation. So below freezing there is no evaporation.

I believe the OP is assuming 1atm, otherwise the temperature values he posted are meaningless.

You're right - water vapor can exist at temperatures such as $\pu{50^\circ C}$. This is a phenomenon known as evaporative cooling, where molecules of water with higher kinetic energies tend to "release" themselves from the system, and as a result, less and less water molecules are held within that system.
Temperature is a sort of proxy for kinetic energy, and vice versa. Reference, also, the Maxwell-Boltzmann distribution, which reveals that at $\pu{100^\circ C}$, for example, not all molecules in the system possess the kinetic energy that they would at $\pu{100^\circ C}$, but rather the average exists at that temperature.
To answer your question, never in a real scenario.

From a chemical engineering perspective (where we do a lot with steam), the answer is dependent on if you're in a closed system or not. If I put some water in a closed container, it would evaporate only enough such that the gaseous water would reach the vapor pressure at that specific temperature. All temperatures above $\pu{0 K}$ have a non-zero vapor pressure, so you could say that water evaporates at all temperatures, at least for a while, and then it will be in equilibrium with the liquid (or solid) state. However, if you are not in a closed system (say, outside with a cup of water), then the water will continue to evaporate to reach the vapor pressure, but the gas will continue to escape, so the water will never stop evaporating.
The opposite, and the answer to your question, is also true. If you have circumstance when the surrounding pressure due to water vapor is higher than the vapor pressure of water at that temperature (like a rain cloud cooling down), then the water will start to precipitate. Of course, like all equilibrium, even if the current net phase shift is from gas towards liquid water, some of the liquid will still be evaporating - it's just that more gas will be condensing.

In other words, what is the Condensation Point of water?

The dew point is the point at which liquid water would condense from the atmosphere, and the frost point is the temperature at which ice would form from the atmosphere. Both are functions of temperature, pressure, and the water content of the atmosphere.

At what temperature, can you say for certain that all water vapor will turn into liquid?

At absolute zero there would be no gaseous water.
For temperature above 0 Kelvin and 273.15 Kelvin (0 Celsius) or below then water condensing from the atmosphere would form ice.
For temperature between 0 Celsius and 100 Celsius liquid water would condense.
Above 100 Celsius water wouldn't condense unless the gas phase of the system was under pressure.

I wasn't sure about this myself! But after thinking about it logically, I realized that when a gas condenses into a liquid at around 50 Degrees, I knew that when a liquid condensed, it was one step below, so when it condensed it would become a solid. When it freezes (condenses) into a solid, the temperature is around freezing (0 Degrees Celcius) and (32 Degrees Farenheit)

Could you back this answer up with some sources?


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Why are the halogens good disinfectants?
I've been searching around the internet for a while and I know that Chlorine, Bromine and Iodine are used as disinfectants.
My question is, what is the property of the halogens that make them suitable for killing microbes? Is it just their toxicity?

The halogens, particularly in their diatomic free states and within various oxoacids, are strong oxidizing agents by virtue of their high electronegativities, electron affinities, and reduction potentials. The polarizability of the heavier halogens also makes them almost uniquely versatile as both good leaving groups and strong nucleophiles, depending on conditions. Additionally, the relatively low bond dissociation energies of the halogens contribute to their aptitude to react by free-radical mechanisms (coupled with the exothermicity that usually accompanies the propagation of those reactions).
In their diatomic forms, they readily add across the $\pi$-bonds of alkenes and alkynes (which are ubiquitous and essential structural features of myriad types of biomolecules), and they can oxidize numerous other functional groups under the right conditions and in combination with other substances. Various oxoacids that incorporate halogens are often extremely strong oxidizers as well (household bleach, for example, is an aqueous solution of sodium hypochlorite, which is mild by comparison to certain other halogen-containing oxoacids). In those reactions, they typically serve to add oxygen to molecules in place of, e.g., hydrogen, while themselves serving as leaving groups (or parts of leaving groups).
The antiseptic properties of halogens (and halogen-containing oxoacids) are directly consequent to those chemical properties. At least one source suggests that elemental iodine, for example, reacts by mechanisms similar to the ones described above to disrupt key structures of bacterial enzymes, proteins, and lipid membranes: it oxidizes the thiol groups of cysteine or methionine to disulfides, interfering with protein synthesis; it reacts with unsaturated lipids to alter their structure; etc. These are biological equivalents of the very same reactions that are routinely carried out in organic chemistry labs. If you want more details/examples, they should be easy to find with a web search.
In short, the very same awesome chemical reactivity that makes them so indispensable in organic synthesis, also means that they are highly reactive towards (and hence, usually harmful to) all manner of organisms.


text:
Which one is nitric Acid, HNO3 by itself, or when it is included in H2O?
This is a quote from my Chemistry textbook:

Nitric acid ionises according to the following equation:
  $$\ce{HNO3(l) + H2O(l) -> H3O+(aq) + NO3- (aq)}$$

I know that the formula for nitric acid is $\ce{HNO3}$. But I cannot understand what is the product of this equation. What is $\ce{H3O+(aq) + NO3- (aq)}$? Is it a diluted acid with less concentration?
This is another quote (from my textbook) that confuses me:

Nitric acid and sulfuric acid are not quite as strong as hydrochloric acid. Chemical and spectroscopic studies have shown that concentrated solutions of both these acids do contain some unionised molecules. For our purpose, however, their ionisation is effectively complete. This is particularly true in dilute solutions.

I cannot understand the part that is in the bold. What are these unionised molecules? I know that these acids are considered strong acids which completely ionise in water solution. If that is the case, then why does the text indicate that there are some unionised molecules?

"Both these acids do contain some unionised molecules." When does the collective bargaining agreement expire?

In your first reaction, $\ce{H2O}$ acts as a base to abstract an $\ce{H+}$ from the nitric acid. The resulting $\ce{H3O+}$, called hydronium, is the conjugate acid, while the $\ce{NO3-}$, called nitrate, is the conjugate base (this is the molecule of nitric acid, but stripped of its $\ce{H+}$). You could view this as a dilution, in the sense that the resulting solution is less acidic than highly concentrated nitric acid would be, but the process of dilution of acid by mixing with water inevitably involves an acid-base reaction.
As to your second question, the unionized molecules are the neutral water and nitric acid on the reactant side, while the ions are the products resulting from the acid-base reaction. Nitric acid is relatively strong, meaning it tends to ionize almost completely in aqueous solution. Additionally, the less acidic the solution already is prior to addition of the acid, the more extensively the added acid will dissociate into ions (to understand this, please read up on acid-base equilibria and Le Chatelier's principle, as the topics are too lengthy and complex to cover adequately here). This is why the author emphasizes that the dissociation is more complete in dilute solutions. Statistically, however, some proportion of the nitric acid molecules will inevitably not dissociate into ions. The stronger an acid, the greater the extent of dissociation, but there is always some statistical chance that a proportion of the molecules will remain unreacted. For strong acids, however, the overwhelming preponderance of the molecules will dissociate, while only a tiny fraction remain unreacted. Hence, for many (if not most) purposes, it's reasonable when dealing with strong acids to proceed as if the dissociation is truly complete, since the unreacted portion is typically negligible in quantity.

Different compounds and mixtures are described as nitric acid
Depending on context and use several chemicals are described as "nitric acid".
The common laboratory chemical usually called concentrated nitric acid is a 68% solution of nitric acid in water (the composition is set as this is an azeotropic mixture that is easier to manufacture).
But other reagents consist of mostly nitric acid with little water (as lower water levels are often useful in rocket fuels or as nitrating or oxidising agents). White fuming nitric acid is mainly pure HNO3 with less than 2% water. Red fuming nitric acid is basically relatively pure HNO3 but will a large proportion of NO2 in the mix (which might be thought to reduce the water content as it reacts with water to give nitric acid). But those fuming variants are more complicated than they look as complex equilibria exist with various nitrogen oxides and some of these equilibria create water even when the initial mixture doesn't contain any. So it is impossible to get bulk nitric acid that is free of any water.
So the question of which reagent is nitric acid is moot as several reagents are called nitric acid and even the "pure" substance isn't completely HNO3 because of a range of side reactions that can't be prevented.

The product is simply the ionized form of $\text{HNO}_3$; if you prefer it is equally correct to write $$\text{HNO}_3(aq)\rightarrow \text{H}^+(aq)+\text{NO}_3^-(aq)$$
Note that technically nitric acid is $\text{HNO}_3(aq)$and not $\text{HNO}_3(l)$ (which would presumably be the salt hydrogen nitrate in liquid form) As hydrogen nitrate will simply decompose under standard conditions if not dissolved so $\text{HNO}_3$ is sufficient to refer to nitric acid; however, it is understood to refer to $\text{HNO}_3$ dissolved in water, not just $\text{HNO}_3$ alone.

You can get nearly pure nitric acid - White_fuming_nitric_acid - it's not a "salt".


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Cathodic protection of aluminium
I am planning an experiment for a chemistry class about corrosion. I am going to show how different prevention methods affects the corrosion of iron and aluminium parts using bleach as the oxidizer.
One of the methods I plan on is protecting the parts using cathodic protection, the iron nail will be protected by aluminium but I don't know a good, easily available metal to protect the  aluminium stud. Any ideas?

You might want to look at this link which has a nice general discussion about cathodic protection, with an abbreviated activity series for the metals of interest in cathodic protection.  Typically, aluminum, magnesium and zinc are used as sacrificial anodes for cathodic protection.
Strictly speaking from a thermodynamic standpoint, one would suspect that magnesium would be needed to protect aluminum.  However aluminum boat manufacturers state that zinc can be used for protecting aluminum boats in salt water.
I'm completely unfamiliar with the die casting industry, but they use Al/Zn alloys which, if affordable, may provide an interesting pedagogical component to your corrosion activity (namely, what do scientists do when the properties of the elements on the periodic table don't quite fit what we have in mind).

Thanks for your answer! I will see if I can borrow some Mg from the lab.

Glad I could help - if you find the answer useful, you should also consider voting it up, as that helps the Chemistry.SE grow.

Of the metals lower in the Standard Electrode Potential list, Magnesium is probably easiest to come by and least problematic to handle.


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Why does the unstabilised Wittig reaction selectively form cis alkenes?
This question is meant for a simple unstabilised ylide.
The mechanism of the Wittig reaction, as given on ChemTube3d, involves a concerted formation of the oxaphosphetane (this is generally favoured over the traditional stepwise mechanism with betaine formation):

Here the puckered transition state has been clearly shown but involves the direct formation of the oxaphosphetane (without any betaine intermediate). In this transition state, the two possibilities are for the methyl groups to be either cis or trans to each other. Surely the transition state with the methyl groups trans to each other will be more stable. Doesn't this support formation of the trans product?

Related: Which is the currently accepted mechanism of a Wittig reaction? and Why do stabilised ylids lead to trans alkenes in the Wittig reaction?

To my knowledge the mechanism of the Wittig reaction isn't fully resolved yet. But maybe I can give you some ideas about why the Wittig reaction with unstabilized ylides is (Z)-selective (well, with the exception of the Schlosser modification) instead of (E)-selective.
In the excellent book by Clayden, Warren, and Greeves, there is a section beginning p. 690 that describes the oxaphosphetane formation as a concerted, antarafacial [2+2]-cycloaddition reaction (similar to your "puckered transition state").
Here, the phosphonium ylide and the carbonyl compound approach each other at right angles. The substituents are arranged in order to minimise steric repulsions in the transition state. In particular, the phenyl group (Ph) points away from the PPh3 group, and then the R group on the ylide points away from the Ph.

The formation of the oxaphosphetane is irreversible and kinetically controlled. Hence, even though the trans oxaphosphetane is more stable, it is not formed.
(By the way, the book also has some justification why the Wittig reaction of stabilized ylides is E-selective.)

Edit:
Since the OP asked for it I will try to give some justification on why the ylide and the carbonyl group approach one another perpendicularly. For a proper description one would need the frontier orbitals of both compounds and then reason about the orbital interactions that bring about the [2+2] cycloaddition. Since I don't have those I will try to give some oversimplified description that should convey the general principle at work here. In the approximate orbital pictures I'm going to use I'll leave out the substituents for clarity.
To begin with, if one imagines the orbital interactions between the ylide and the carbonyl group, it is clear that the most important (primary) interaction for the reaction is the one between the HOMO of the ylide (nucleophile) and the LUMO of the carbonyl group (electrophile). What happens when both groups approach one another head-on or at right angles is shown in the following picture (for the perpendicular approach a top-view is used).

In both cases there are as many bonding interactions (black) as there are antibonding interactions (red). These situations are overall non-bonding and cannot lead to the desired product. This is the reason why [2+2] cycloadditions are usually thermally forbidden (by symmetry).
However, some compounds have additional orbitals that: (1) are similar in energy to the frontier orbitals drawn in the picture, and (2) possess the right symmetry to interact with the other orbitals. These "secondary orbital interactions" are weaker than the HOMO-LUMO interaction, but might provide enough bonding to tip the balance from an overall non-bonding to a bonding situation.
So, how does this translate to the situation of the Wittig reaction? Since O is more electronegative than C, the HOMO of the carbonyl group will have a larger orbital coefficient at O, while the LUMO will have a larger orbital coefficient at C. The situation is similar with the ylide. Since C is more electronegative than P, the HOMO of the ylide group will have a larger orbital coefficient at C. The ylide group is the nucleophile so the primary interaction will be between its HOMO and the LUMO of the electrophilic carbonyl group. Since the orbitals interact best where the coefficients are large, this primary interaction will be biased towards C–C bonding.
The question is now, is there a suitable secondary interaction? One can argue that P has a lot of orbitals (valence shell extension) and that the LUMO of the ylide might be some empty phosphorus orbital (or at least an orbital with a big coefficient on $\ce{P}$) - this could be a p-orbital perpendicular to the ylide's C=P π-orbitals, or a d-orbital, or something like that. This phosphorus orbital can then interact with the carbonyl group's HOMO, which has its biggest coefficient on O. So this bonding interaction will be biased towards P–O bonding. The situation is shown in the following picture:

The biases in the primary and secondary orbital interactions will lead to some distortions from the pure perpendicular approach which will become more pronounced the nearer the two species get to each other. For example, the ylide might rotate a bit so that its C and P atoms are closer to the carbonyl's C and O atoms respectively, in order to maximize the bonding overlap. And of course the perpendicular approach of ylide and carbonyl group will proceed in such a way that their respective biggest substituents are as far away from each other as possible (see diagram above).
After this quite lengthy answer I want to make clear that this description of mine is only a guess. The real reaction path is still a matter of debate and there are some compounds that rather react via a radical or ionic pathway. But I think it gives a good explanation for most of the observed behaviour and hope it helped the understanding. A more exact frontier orbital description of another thermal [2+2] cycloaddition, namely that of ketene with ethene, can be found in the textbook by Brueckner on p 653.


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How to chemically postpone a chemical reaction
This question is about delaying/postponing chemical reactions in general, but I will explain my point based on an example. Most of you will know the famous vinegar and baking soda experiment which allows you to make a volcano or a rocket. What I have noticed when doing this experiment (in particular the rocket experiment) is that if you just put soda in powder form in the bottle and pour in vinegar, a lot of the formed gas escapes before you can put the cork on thus resulting in a less spectacular launch.
In the video I linked above they wrap the soda in paper to allow some time between the addition of vinegar and the start of the reaction. This is a nice trick, but in many cases you don't want any paper inside your 'reaction' vessel. My question is: what are my options to chemically postpone the start of the reaction?

Does the use of vinegar make it organic?

I'm not sure, I had a hard time picking the tags on this question and I could imagine that people working in organic chemistry/synthesis might be more knowledgeable about the topic. Please feel free to retag if you have suggestions!

In general the answer to your question is no, you cannot postpone the reaction once the reactants are actually in contact.
Reaction rate usually decreases with temperature.  Thus it is possible, but not certain, that if you freeze a reaction mixture down to liquid helium temperatures the reaction would stop.  This isn't really practical in general.

From what I remember, the powder was kept separate.

The powder was put on a piece of tissue paper.
Put the vinegar in the body of the rocket.
Put the cork in with the paper wrapped powder.
On putting it in place the vinegar gets to the paper/powder.

This brings me back...
You could work on the cork to put the powder in it without paper or something.
Also, you could make a paste and put it in a hole on the cork.

The paste sounds like a very nice technique. Problem is that these methods are still based on physical separation of the vinegar and the soda. What I am looking for is a chemical way to postpone/delay the reaction (not even specifically this reaction, but I thought this was a nice example case)

The closest I get is catalysis. For reactions that need a catalyst you can do it. For the rest tough luck.

You can delay a reaction- chemical poisons do that. They act in various ways but one way is to react with reactive intermediaries (the initial chemicals need to break some bonds to form the intermediaries) and produce inert products which cannot further react. It may even kill a reaction in process. I don't whether there is one for that chemical reaction.

There are clock reactions that depend on reagents being used up, along with thresholds of certain properties e.g pH or concentraition.

One class of example is the iodine clock reactions, in which an iodine
species is mixed with redox reagents in the presence of starch. After
a delay, a dark blue color suddenly appears due to the formation of a
triiodide-starch complex.

https://en.wikipedia.org/wiki/Chemical_clock


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Does single-layered Aluminium Phosphide exist?
I'm currently working with the band structure of graphene and graphene-like compounds.
α-Boron Nitride (hexagonal, single layer) has many interesting properties. I was wondering if the next period analogue, a hexagonal single-layered Aluminium Phosphide existed. Wikipedia lists this compound, but it seems to only be found with a multilayered sphalerite structure.
Does an allotrope of AlP exist (in a stable state at normal conditions) where the structure is similar to graphene or silicene? If not, why?

(I'm pretty confident that it doesn't exist, so the why is what I'm actually interested in)

I suspect that it does not.  As far as I know, all trivalent phosphorous is highly pyramidal--or at least not planar. This would rule out any graphene-like structure.
The reasoning that I have seen is based on hybridization. In the second row (third period) the 3s and 3p atomic orbitals are well separated, implying that s and p can not mix and that the outer valence is unhybridized. If bonding is with unhybridized p's, then bonding would occur at right angles.
(Kutzelnigg, W. Angew. Chem., Int. Ed. Engl. 1984, 23, 272.)
(Brown, E. Organometallics 2000, 19, 2208-2214.)


text:
Why is the mole a unit of measurement?
I really can't understand why is the mole a unit of measurement and not just a constant. I mean, we use units to express things that we can measure but that are not countable, that are continuous - and not discrete. However, the number of atoms (or whatever) of something is clearly countable, discrete. It is also a huge number, which would request for a constant, a number to multiply to: Avogadro's constant is that number, like mega or kilo in SI for other units. We could still write 14 mole, where mole is the Avogadro constant (exactly the way it is), but every book I read so far states that the mole is a unit of measurement, official in the S.I., and so does Wikipedia. It just makes no sense to me that it is considered that way, but there is obviously a reason why... What is it?

I've been thinking about this question too! Why is "mol" a "Base SI unit" when "radian" is not?

Good question. It got even more confusing since: "The definition of mole was adopted in November 2018 as one of the seven SI base units" :-/

I think you are confusing unit of measurement with fundamental unit.  A unit of measurement is a definite magnitude of a physical quantity.  Consider a very common grouping unit, the dozen which is conceptually the same as a mole and is a little easier for us to grasp since it is a commonly used small number.  As far as I know, the only requirements for a unit of measurement to be valid is general consensus.  If you and I agree that we will measure distance in units of 'Chemistry Textbook Thickness', then it is a valid unit of measurement.
A fundamendal unit is a part of a particular set of units upon which all other units can be derived.  In this case, the mole is related to the fundamental unit of mass and is defined as an amount of items equal to the amount of carbon atoms in 12 grams of $^{12}C$.

Indeed, I think I meant fudamental unit. It just seems to me that it makes sense that the meter is a fundamental unit: you can't just say the distance between you am me is 20. It requires a unit. The mole, on the other hand, is exactly the opposite: it seems to be a unit to count things that do not require units do be counted, like atoms or potatoes.

Yet this does not answer the main question: why not simply consider the mole as a constant? A mole is certainly not a dimensional unit such such as one representing mass or time... It is just a number, somewhat like $pi$ or $10^{24}$.

Like @Oliver, I'm not sure this answer addresses the actual question. Nobody disputes that you could call a mole a unit, and indeed the SI system does, but the question is why does it? "Chemistry Textbook Thickness" is not comparable to moles or dozens, because that's defining a unit of measurement for a concept (length) that we have no other way of defining. When it comes to a simple discrete number of objects, though, the thing looks to me like it's unitless. It's simply a number, and multiplying it by 6E+23 and saying that's a unit doesn't seem logical.

"The definition of mole was adopted in November 2018 as one of the seven SI base units". So now, the mole is a fundamental unit, and does not depend on mass anymore. Which means the above question still stands, and your answer doesn't apply anymore.

Let me add to bobthechemist's answer.  The mole counts quantity.  The fundamental unit involved is Avogadro's Number, approximately $6\times 10^{23}$.  So $6\times 10^{23}$ units of anything is a mole of that thing, whether it is atoms of hydrogen or sand particles.
With this in mind reread bobthechemist's answer.

When my high school teacher wrote at the end of the exam '1 mole of kisses to you', not only she was very ignorant but she made a large disservice to me and to my class.
And when my prep school teacher compared a mole to a dozen, as if they were conceptually the same, he made the same disservice, because mole is not just a number. E.g.
1 dozen eggs = 12 eggs
1 dozen atoms of Fe = 12 atoms of Fe
1 dozen particles = 12 particles
1 mole of Fe = {Nᴀ} particles of Fe
1 mole of some substance = {Nᴀ} particles of that substance
So, if you were to follow the syntax of the definition of a mole, then a mole of kisses is NOT {Nᴀ} kisses. It is actually {Nᴀ} particles of kisses, which makes no sense.
You see – the mole has the same physical dimension of the particle – and is not a dimensionless quantity like the dozen.
Note: I am using {Nᴀ} as the numerical part of the Avogadro constant, Nᴀ = {Nᴀ}[Nᴀ] = {Nᴀ} mol⁻¹

Note that there is Avogadro number and Avogadro constant, where the former is dimensionless like dozen.

I'm not sure I see the distinction your are trying to make. A mole–like a dozen–is a number of countable things and meaningless if the thing is not specified. It does not specify the nature of the thing, just that it should be countable.

@matt_black Like the 1st level ever of math abstraction - the concept of natural numbers without explicit countable physical objects.

@matt_black (i)My point is that we usually don't say a 'A dozen of iron' like we say 'A dozen of eggs'. Because what is 'A dozen of iron'? Is it 'A dozen of atoms of Fe'?
On the other hand, we usually do say 'One mole of Fe'. Is that just a figure of speech – an ellipsis – for 'One mole [of atoms] of Fe'? I don’t think so.
(ii)

@matt_black (ii)My reasoning behind all of this was to try to understand why the amount of substance is usually regarded as a dimensional quantity with the mole being its SI unit – a dimensional unit – unlike the dozen which is regarded as dimensionless.
And that was my answer – if mole is dimensional then it cannot be just "a number to multiply to", quoting the owner of the question. This makes it conceptually very different from the dozen even though both are used for countable things.

Maybe the mole is used to bridge the amount of substance –  uncountable – and the number  of particles – countable – unlike the dozen that bridges countable things with countable things.

In this regard saying 'One mole of atoms of Fe' would probably be incorrect.

@AlanPommer I think that, in most circumstances, chemists should not say a mole of iron. They should specify the countable thing (though a mole of "Fe" implies the thing is iron atoms). This is important to avoid ambiguity. What, for example, does a mole of oxygen mean? It could be atoms or molecules: it should be specified.


text:
is it possible to store formic acid longer if frozen?
I would like to know if formic acid (concentrated to 85%) is stable (no chemical reaction) when frozen.
Is it a good solution to store it for many years?

Could you elaborate on the types of chemical reactions you want to make it stable against? And what you want to do with the formic acid after storage? That will probably help in determining whether freezing is a good idea or not.

Hi Michiel, I'm using formic acid as varroa treatment for bees. I have two bee hives and I can only get 1 or 4L of formic acid. This quantity may last many years so I have to store it in an efficient way. I have read that the formic acid may slowly decompose to carbon monoxide and water, so I would like to know if by freezing it the decomposition will slow down or stop.

Formic acid is thermodynamically unstable to decomposition to water and carbon monoxide: HCOOH ==> CO + H2O. I assume this is what you are referring to.
A common use is acidification of silage to prevent smells. The fact that it decomposes makes it more environmentally friendly than alternatives. I was not aware that it was used in bee keeping to control varroa parasitic mites.
Lowering the temperature would reduce the decomposition rate from both a thermodynamic and kinetic point of view.
The MSDS I have looked at indicate that the bottle should be vented periodically to release the pressure built up from decomposition. However this one also says that formic acid should not be stored below its freezing point of 8C. http://ionics.ca/wp-content/uploads/MSDS-Formic-Acid.pdf Whether this is just for convenience of use or for reasons or safety I do not know.
To me, the build up of pressure seems far more of an issue than any loss of strength in the acid. In order to lose 50% of its strength, a 1L bottle of formic acid (1200g/46=26mol) would have to lose 13mol (about 312L) of carbon monoxide. That's a lot of venting! I would think a bottle should have a shelf life of 2 or 3 years.
I did my university chemical engineering project on a formic acid plant. It is made as follows:
1. CO + Methanol (20 bar, 80C, potassium methoxide catalyst) == Methyl Formate
2. Methyl Formate + H2O (acidic autocatylisis)== Formic acid + Methanol

From this you might conclude that formic acid is stable to decomposition at 20 bar, but that is not the case, because excess reactants and careful use of distillation are used to push both reactions in the desired direction.
That was a long time ago, but I will see if I can find any relevant data in the decomposition reaction. However you would probably be best off just following the manufacturer's recommendations.
Oh, and keep it in the original glass bottle. Many metals catalyse the decomposition.

Large (>250 mL) bottles of formic acid I have bought commercially came with vented caps.


text:
Geometrical isomerism in cycloalkenes
Why doesn't cyclohexene have a cis-trans isomer whereas cyclodecene has one?
Is strain a factor?

You simply can't make a transoid molecule of cyclohexene because there is not enough "length" in the carbon chain.
I don't know if it is "strain" per se--but a putative transoid cyclohexene would have extraordinarily exotic bonding arrangement--thus being very high in energy.

Yes, in cycloalkenes with of 7 or less carbon atoms, the high ring tension is a prohibitive factor, but in higher cycloalkenes it's not that great so they can show the geometrical isomerism.