Are there any hazardous products that would come from mixing glass cleaner with lubricant? Are there any known safety issues if you mix:

Mr.Muscle glass cleaner
GT85 spray lubricant

Nothing compromising safety will happen. Now, these are two completely different cleaning agents based on alcohol/water and organic-fuel respectively. Most likely will turn white  because these substances are immiscible. Therefore, it will not be a good idea to apply a mix for you cleaning tasks.  This will be  specially so  over wood or any rough porous surface likely as they will absorb the substances resulting in unappealing white or smear white spots.  Application over polished metal surfaces will be ok.


Relation between moles, molar mass and mass You are asked to prepare a 0.8500 M solution of aluminum nitrate. You find that you have only 50.00 g of the solid.
What is the maximum volume of solution that you can prepare?

I'm confusing my molar masses with grams. The molar mass of Aluminum nitrate $\ce{Al(NO3)3}$ is 213.0 g. If I have 50.00 g of the entire solid $\ce{Al(NO3)3}$ does that mean I have 50.00 times the amount of 213.0 g or will I end up Multiplying 50.00 g of Aluminum nitrate by 0.8500 M, then dividing that by the MM of Aluminum Nitrate. But then I don't know how to set up the rest of my dimensional analysis to reach Liters.$%edit$

If you have 50 g of a compound that has a molar mass of 213 $\frac{\rm{g}}{\rm mol}$, then you have $\frac{50~\rm{g}}{213 \frac{\rm{ g}}{\rm mol}}=\frac{50}{213} \rm{mol}=0.234~\rm{mol}$ of the compound.

If you want to prepare a 0.8500 molar solution, that means the concentration you want is $0.85~\rm{M}=0.85\frac{\rm{mol}}{\rm{L}}$.

With those two numbers, the final answer should be apparent.  (Just divide them and make sure the units work out!)


Acid catalysed hydration of allene In the acid catalysed hydration of allene, acetone is obtained as the major product. This is because of electrophilic addition of $\ce{H+}$ across a double bond to give a vinyl carbocation, and subsequent addition of water and then tautomerism to yield the ketone. What I'm confused about is that why does not an allyl carbocation form, which would be resonance stabilised?(leading to  allyl alcohol?)

Good question!  As the following figure illustrates, there are 2 ways we can protonate an allene.

(image source)

The top line in the drawing starts by protonating at the central allene carbon.  This generates a very unstable primary carbocation.  Initially, this primary carbocation is not stabilized by the adjacent double bond.  It is not until after this carbocation rotates 90° that its p orbital is lined up with the p orbitals in the double bond and allylic stabilization can occur.  Since there is no immediate allylic stabilization when the carbocation is formed, and since it is a primary carbocation, the pathway leading to it is very high energy.

On the other hand, protonation at the terminal allene carbon (bottom line of figure) produces a relatively stabilized secondary carbocation.  Formation of a secondary carbocation is a much lower energy process.  Therefore, protonation at the terminal allene carbon is lower energy and is what will occur.

Edit: Response to OP's comment

  I always thought that a positive charge appearing on sp3 carbon (which
  then becomes sp2) is more stable than a positive charge appearing on
  sp2 carbon

Generally we think of vinylic carbocations as high energy.  That is true, but what is meant by "vinylic" carbocation is different from what we have here.  Here there is an empty p orbital.  In a vinylic carbocation the empty orbital is not a p orbital, but rather something like $\ce{sp^2}$ or $\ce{sp}$ hybridized.

  is the secondary carbocation more stable because of inductive effect,
  the rotational barrier in the primary carbocation, or both

The secondary carbocation is more stable for the same reasons that we have the  tertiary > secondary > primary carbocation stability order.

Inductive effects, $\ce{sp^3}$ and $\ce{sp^2}$ carbons are electron releasing towards an $\ce{sp}$ carbon due to electronegativity differences
Resonance effects from hyperconjugation involving hydrogens on the carbon next to the carbocation center.


Why isn't calcium carbide more corrosive? Acetylene has a pKa of 26, I think.  While that's pretty acidic ... for a hydrocarbon ... I still expect calcium acetylide (carbide) to be a very strong base indeed.  Yet the NFPA fire diamond health code is just a 1, compared to a 3 for both sodium and calcium hydroxides, even though each mole of the carbide ought to create a mole of calcium hydroxide the moment it touches water.  Little kids play with toy cannons powered by somewhat impure calcium carbide, taking no special lab precautions.  Why is that possible?
Update: I have since noticed that those NFPA diamonds don't all agree. The Santa Rosa Fire Department and Wikipedia say Health rating is 1.  Fisher Scientific says it is 2.  NOAA says it is 3.  I may never look at one of those little diamonds the same way again...

In these Safety Data Sheets for $\ce{Ca(OH)2)}$ and for $\ce{CaC2}$, both compounds are rated at "Causes serious eye damage – Category 1" and "Causes skin irritation – Category 2".
That said, $\ce{CaC2}$ is often sold as pellets or "stones". Because of large particle size, only a small amount of $\ce{Ca(OH)2)}$ is produced at a time from the action of water on $\ce{CaC2}$ since the release of acetylene bubbles slows penetration of water into the pellets.
$\ce{Ca(OH)2)}$, on the other hand, is more often sold as a finely divided powder. That dust is, in my experience more of a respiratory hazard. The "lime", as it's slaked in water, also evolves considerable heat, which exacerbates the caustic effect on skin.
Of course, $\ce{CaC2}$ has it's own hazards... it can cause considerable damage from an acetylene explosion.


Why is it possible to image LUMO if these orbitals are, by definition, unoccupied? LUMO stands for lowest unoccupied molecular orbital. If it's unoccupied, how is it possible to image it?

See for example in this paper ("Pentacene imaged with STM and NC-AFM").

Why not? It's not like we poke them with our finger, anyway. All we can really see are the electron transfers. Now, the transfers from the molecule are as good as the transfers to it. If we adjust our setup so as to look at the latter, then we may see the shape of the orbital which is now empty but ready to accommodate our electron.

In one setup, the electrons tunnel from the HOMO to the tip, thus visualizing the HOMO. In another setup, they tunnel in reverse direction (hence the change in the sign of voltage), from the tip to the LUMO, thus visualizing the LUMO.


Estimation of Sodium Hydroxide Why is estimation of sodium hydroxide more accurate with oxalic acid rather than with sulfuric acid or any other strong acid?

There are a few reasons. Some of the most important are:

-Oxalic acid dihydrate isn't hygroscopic (it doesn't absorb water from the atmosphere, whereas sulfuric acid does). If a substance can absorb water from the atmosphere, then when you weigh it out, you won't be able to accurately calculate the number of moles used because you don't know how much is water and how much is your substance. This won't be the case with oxalic acid dihydrate.

It's chemically stable when stored.
As Waylander said (in his comment elsewhere), it's a solid so is easily weighed out.
It has a moderately high molecular mass of $\pu{126 g/mol}$ (as the dihydrate) so any error in the mass weighed corresponds to a relatively small error in the calculated amount in $\pu{moles}$.

As a result, oxalic acid has become a primary standard: So you can use it to make a solution of accurately known concentration to determine the concentration of a solution of sodium hydroxide (by titration).


Why are some ions mutagenic and others are not? It is my understanding that when ionic species (such as $\ce{He^2+}$, or alpha particles) enter the human body, they tend to react with biomolecules, causing damage by (for example) mutating DNA.

If such ions can cause severe damage to the human body, why are other ions like $\ce{Na+}$ or $\ce{Cl-}$ safe to consume and not prone to causing similar mutagenic effects?

You are confused about the cause of mutagenesis in alpha particles.

The reason why an alpha particle is mutagenic is not that it is an ion, it is that it is an ion carrying a great deal of kinetic energy. Alpha particles are produced in energetic nuclear reactions and carry a great deal of energy away from the nucleus that emits them (a typical alpha particle has about 5MeV of energy, enough to break a typical chemical bond). It is the kinetic energy that matters not what the ion consists of. Heavier ions with that sort of energy would be just as bad for biological systems as are other forms of energetic radiation (beta particles are just electrons but will still do damage, gamma rays are just powerful photons but will dump enough energy into molecules to break DNA).

Non-energetic helium nuclei would pick up electrons extremely rapidly from the first molecules they encountered (probably from air even before reaching your skin) and would probably be as harmless to you as the helium filling a balloon).


Use of Concentration for pH When we find the pH of a solution, do we use normality or molarity? My teacher said that if the compound is present by itself, we can use normality but in all other cases we should use molarity. Can you explain why?

Current definition implies that $\mathrm{pH}$ is a function of relative activity.
Originally, the amount concentration of  $\ce{H+}$ in $\pu{mol L-1}$ was proposed, which is also often used these days as an approximation [1].

$\mathrm{pH}$ was originally defined by Sørensen in 1909 … in terms of the concentration of hydrogen ions (in modern nomenclature) as $\mathrm{pH} = −\lg(c_\ce{H}/c^\circ)$ where $c_\ce{H}$ is the hydrogen ion concentration in $\pu{mol dm–3},$ and $c^\circ = \pu{1 mol dm–3}$ is the standard amount concentration. Subsequently …, it has been accepted that it is more satisfactory to define $\mathrm{pH}$ in terms of the relative activity of hydrogen ions in solution
$$\mathrm{pH} = -\lg a_\ce{H} = -\lg (m_\ce{H}γ_\ce{H}/m^\circ)\tag{1}$$
where $a_\ce{H}$ is the relative (molality basis) activity and $γ_\ce{H}$ is the molal activity coefficient of the hydrogen ion $\ce{H+}$ at the molality $m_\ce{H},$ and $m^\circ$ is the standard molality. The quantity $\mathrm{pH}$ is intended to be a measure of the activity of hydrogen ions in solution.
However, since it is defined in terms of a quantity that cannot be measured by a thermodynamically valid method, eq. 1 can be only a notional definition of $\mathrm{pH}.$

So, if the numerical value for normality coincides with the one for molarity, and one can omit the activity in favor of molarity, i.e. when both equivalence factor $f_\mathrm{eq}$ and activity coefficient $γ$ are approx. 1, then yes, normality can also be used to estimate the $\mathrm{pH}.$
I'm not sure how to interpret "present by itself" part of your question, but if it means there is a single solute, then it seems to be incorrect.
The simplest counter example would be a solution of di-, tri- and polybasic acid/base, say $\ce{H2SO4},$ for which the equivalent concentration is double the value of amount concentration.
References

Buck, R. P.; Rondinini, S.; Covington, A. K.; Baucke, F. G. K.; Brett, C. M. A.; Camoes, M. F.; Milton, M. J. T.; Mussini, T.; Naumann, R.; Pratt, K. W.; et al. Measurement of pH. Definition, Standards, and Procedures (IUPAC Recommendations 2002). Pure and Applied Chemistry 2002, 74 (11), 2169–2200. DOI: 10.1351/pac200274112169. (Free Access)


Is there any aqueous solution (or simple method) to remove trace level of sulfur dioxide/nitrogen monoxide and nitrogen dioxide in a gas stream? I need to generate pure gas from a gas stream which is free of $\ce{SO2}$/$\ce{NO}$/$\ce{NO2}.$ The air is already pretreated by activated carbon and purafil. However, I can't find any solution for removing trace level of this pollutants. I find that many solutions of capturing for large scale like flue gas. However, it is not suitable for trace level (1–10 ppb).

I know $\ce{O3}$ can react with $\ce{NO}$ to $\ce{NO2}$. However, as I mentioned, I need the gas stream free of $\ce{NO}$ and $\ce{NO2}$ simultaneously. So, converting $\ce{NO}$ to $\ce{NO2}$ is not the answer.

Removal of $\ce{SO2, NO2}$ from solution

As mentioned by Maurice, you can remove $\ce{SO2, NO2}$ by bubbling the mixture bubble through NaOH or KOH solution, or washing it through a column with flowing NaOH or KOH solution at countercurrent.

Removal of $\ce{NO}$ from solution

You can bubble the mixture through $\ce{aq.FeSO4}$ solution.
I remember doing this in eudiometry experiments.


Products of the reaction between chromium(III) chloride, zinc and sulfuric acid The aim of my experiment was to make a Cr(II) aquacomplex in situ, and for that I have the reaction:

$$\ce{CrCl3·6H2O + Zn}$$

and I add concentrated sulfuric acid to it. I know that chromium(III) reduces to chromium(II), but I'm not sure about the reaction. Would it be

$$\ce{CrCl3·6H2O + Zn -> Cr(H2O)6Cl2 + ZnCl + H2}?$$

I know that chromium(III) reduces to chromium(II), but I'm not sure about the reaction. Would it be $$\ce{CrCl3·6H2O + Zn -> Cr(H2O)6Cl2 + ZnCl + H2}?$$

Your concept is right and I am glad nobody taught you the concept of nascent hydrogen. It is an obsolete idea of the 1920-40s. Please don't use it anywhere. Wikipedia has a nice summary https://en.wikipedia.org/wiki/Nascent_hydrogen and of course there are several articles which negate this very concept.

You can just simply write the ionic equation involving Cr(III), Zn, resulting in Cr(II) and Zn(II). Hydrogen is just a side-product, which is not participating in the reaction. Write the conditions over the arrow.

Alternatively, write the aquated forms as a complex. Your current equation has incorrect formula of zinc salt. Check what it should be!


Why Fe 3+ is more common than Fe 2+ I heard that $\ce{Fe}(\mathrm{III})$ is more common than $\ce{Fe}(\mathrm{II})$ but I've not heard a very clear explanation. Could someone please explain this incorporating electron configurations in their answer?

The electron configuration of  $\ce{Fe(II)}$ is $1s^22s^22p^6 3s^23p^6 3d^6$.

The electron configuration of  $\ce{Fe(III)}$ is $1s^22s^22p^6 3s^23p^6 3d^5$.

The half filled  orbital $3d$ is more stable than the same orbital filled with 6 electrons. So,  $\ce{Fe(III)} $ ion is more common than $\ce{Fe(II)}$.


How to predict the solubility of an organic compound in different kinds of solvents? How is it possible to know whether a particular solvent will dissolve in a particular kind of solvent, i.e. organic, polar, protic...?
Are there any indicators to look for in molecules or elements that allow one to predict what kind of solvent should be used to dissolve it?

There are some general guesses one can make from looking at the structure but the Abraham solvation equation is commonly used to estimate the solubility of a compound in a given organic solvent:
$$\log P_s = c + e E + s S + a A + b B + v V$$
This equation relies on a set of descriptors to characterize the properties of solvents and solutes to give the water/solvent partition coefficient $P_s$ (the ratio of the solubilities in water and solvent, under certain assumptions):

$p,e,s,a,b,v$ are coefficients that describe the solvent. Abraham and coworkers have collected coefficients for 85 different solvents.

The other variables are descriptors of the solute:

$E$ is the excess molar refractivity, which is a measure of the polarizability of a molecule.

$S$ is the solute dipolarity/polarizability.

$A$ is the overall (summation) hydrogen bond acidity.

$B$ is the overall (summation) hydrogen bond basicity.

$V$ is the McGowan characteristic volume.

You can look at the paper if you want to know exactly what these descriptors are, but their values can be obtained through various experimental and computational means, and combined with the solubility in water, used to estimate the solubility of a given compound in an organic solvent.

Fortunately, you don't have to do all this by hand as several websites, like this, that offer lookup and computation where you only have to specify a compound and pick a solvent from a list.


How does valency change in a reaction? There is a reaction given in my book:

$\ce{2FeSO_4 +Heat -> Fe_2O_3 + SO_2 + SO_3}$

I just noticed that on the left $\ce{Fe}$ has valency $2$, while it has $3$ on the right. How can it change valency?

Changing the number of electrons is easy, but this only changes the charge, not the valency. Does this mean that the number of protons changes?

Yes, valency can change, just like that. Some elements (in fact, most of them) can manifest different valencies; that's their nature. If you look again at this very reaction, you might notice that Fe is not the only element to change its valency here; S does the same. Reactions like this one are called redox reactions. There is a huge body of knowledge that explains what valencies (or rather, what oxidation states) an element can have, and which of them are more stable; the answer by @Raghav Gupta touches on that in more detail.

As to the number of protons (if you mean the protons in the element's nucleus), it has nothing to do with valency. Really, Fe contains 26 protons; is its valency 26? Absolutely not, it is 2 or 3. Also, the number of protons in an element can never be changed at all. If you change it, you'll get another element; also, that would be a nuclear reaction, a thing which is out of scope of chemistry textbooks.


Concentrating Oxalic Acid by Evaporation I bought Oxalic Acid in an aqueous solution from a hardware store years ago. I need to turn it to powder form for easier storage and other things.

Can I simply evaporate the solution? I seem to notice that evaporating it releases fumes that are similar when one is boiling Acetic Acid. Are there more Oxalic Acid being lost just as it is when boiling Acetic Acid (not distillation)

Depends on how hot the plate is I'd say. Oxalic acid is hygroscopic, so it will be hard to completely dry it by just subjecting it to heat as you might just cause it to break down.

There's a paper and procedure on Organic Synthesis on how to obtain the anhydrous form, so you should check that out to see if that meets your requirements.


What is electron configuration used for? What is the purpose of writing electron configuration vs just using the number of electrons off of the periodic table. I know it can be used to determine the number of electrons in each energy level, and therefore, the number of valence electrons. Is there anything else? The number of valence electrons can be found from the periodic table, so why go through that process?

Well, it's almost always this way.

If you know the atomic electronic configuration, you know how many electrons an element has.
If you know how many electrons it has, you'll know how many it needs to have a stable configuration (usually octet).
If you know how many it electrons it needs, you'll know how it reacts with other species.
If you know how substances react, you can predict the reactions involving this substance.
If you know reactions involving this substance, you can either discover a new reaction (knowledge of new reactions makes synthesis easier) or you could invent an unknown compound with interesting properties, which may have its own special uses.


Why is the initial carbocation not formed adjacent to the phenyl moiety? I know that after rearrangement the carbocation is more stable, but we should protonate the compound first where the carbocation it is more stable, i.e.  adjacent to the phenyl group. Why do they suggest a different pathway? Is the shown correct?

I agree with your sentiment, so I run a very quick calculation with a semi-empirical method (GFN-xTB) to find the protomers. This has been done using water as solvent. Note that I did not choose to include any deuterium.

First I calculated the conformational space of the plain molecule (NIST: 1,1'-(3-methyl-1-propene-1,3-diyl)bis-benzene, InChI=1S/C16H16/c1-14(16-10-6-3-7-11-16)12-13-15-8-4-2-5-9-15/h2-14H,1H3/b13-12+, Google search InChIKey: GNQWHYWLSGTMSL-OUKQBFOZSA-N) to obtain the lowest lying conformation and show you the HOMO:

As you can see the double bond is a good protonation cite, with both carbon atoms having about equal contribution; one would therefore expect that the most stable intermediate will form. Undoubtedly this should be the one where the charge can be delocalised into the aromatic π-system.

The automatic generation of the protomers agrees here, and the lowest lying conformation (with partial π-stacking) is shown below, the protonated carbon indicated in bright white:

There is a tentative pathway via a four-membered (green) transition state to rearrange to the product shown in the exam question, but I do not have the time to investigate. Therefore here is my proposed mechanistic scheme:


Why do people still use the mole (unit) in chemistry? I know that the mole is widely used in chemistry instead of units of mass or volume as a convenient way to express amounts of reactants or of products of chemical reactions.

I'm wondering why people in chemistry still excessively use it for their measurement? To be backward-compatible and consistent with traditional textbooks? Why they don't simply express their quantity measurement per atom, per unit volume, per molecule or etc.?

Simply speaking, because it's an appropriate unit to use.

Let's imagine I wanted to measure the length of a rope. What would be an appropriate length to use? Inches? Centimeters? Feet, maybe? It would really be awkward to express it as 0.000189393 miles, or as 304,800,000 nanometers.

(Note: if you can't see why these units are awkward, take any page discussing things like this (e.g. biomass of certain species) and change all the units so that they're nonsense like this. Then put it away for a week and try to read it later.)

Now let's say I've changed my mind and I'd like to measure all the rope created in the world in a year. Now an appropriate unit is almost certainly miles or kilometers, and not inches or centimeters.

Let's consider something else: so far, I've been using length to measure ropes. Would it make sense to measure their combined mass instead? Maybe not for small amounts, since I think I would throttle you if you told me to cut off half a pound of rope, but for global-scale things, tons or metric tons may make sense.

On the other hand, using measurements like the average density of the rope or the combined diameters of all the ropes really wouldn't be much use at all.

What we've seen here is that when we're measuring things, there are measures that make sense (for rope, length, maybe weight) and some measures that really don't (color, average diameter, etc.). In those measures, there are units that are convenient (inches, feet), and some that aren't (nanometers).

This is exactly the problem with chemical units, but much more magnified. When I measure the energy released on hydrolysis of a sample, I'm not measuring the energy of one bond, or two bonds, or a thousand bonds, or even a million. I'm measuring a collection of so insanely many molecules that there's really no number in common language for it.

What is an appropriate way to measure things in these molecules? If I'm only interested in how much I have and not what's in it, per gram or per volume is often a good way to do it. Since when I'm measuring the temperature change of something, I don't particularly care what it's made of, measures of this (specific heat, for instance) are done in units of something per gram.

On the other hand, if I do care what my sample is made of, then I need a better way to measure it. For example, sodium chloride and calcium chloride look similar, but will have very different energies since there's three ions in a $\ce{CaCl2}$ unit and only two in an $\ce{NaCl}$ unit.

Since you can't have half a molecule, the simplest way to do it is to count the darn things. Unfortunately, measuring things per atom is a really awkward way to do things because of the problems I described above. The energy it takes to melt ice is 0.000000000000000000001029 J/molecule.

What we need then, is some count of molecules that's convenient. We could go by 1000 molecules, or a million molecules, but it's a pain in the butt to convert between that and macroscopic units (how much does 1 million calcium atoms weigh?) Now it just so happens that $6.022 \times 10^{23}$ atoms of carbon-12 weigh 12 grams, and a single carbon-12 weighs 12 amu. A single molecule of water weighs roughly 18 amu, $6.022 \times 10^{23}$ molecules of water weigh 18 grams.

Let's take a look at what using $6.022 \times 10^{23}$ molecules offers us over other units:

It is appropriate. $6.022 \times 10^{23}$ molecules will typically
fall in the gram to kilogram range of substance, which can easily be
measured out on a balance. It's also within the range of what you'd
usually use in a lab.
It allows for easy conversion. You can easily tell when you have
$6.022 \times 10^{23}$ molecules because your atomic mass is equal to
your macroscopic mass.
It makes sense. Intuitively, we'd like to measure things per atom or
per molecule, but doing so leads to some ugly units. Instead, if we
measure per $6.022 \times 10^{23}$ molecules, we still are doing a
counting-based measurement, but the numbers themselves are a lot
nicer.

You probably know this already, but we call $6.022 \times 10^{23}$ molecules a mole. These are the advantages of using the mole as a unit, and not for backwards-compatibility. I am of the opinion that if we shed all our units today and started from scratch, we would start off with grams and liters, but we would probably start using the measure of a mole again within a year. It's just a very powerful, useful way to measure things.

P.S. This is much longer than I was initially intending and was not written while I was in the most awake state of mind. Please let me know in the comments if you think this answer is unclear, or stupid, or just plain wrong.


Suitable antifreeze solution If glycerine $\ce{C3H5(OH)3}$and methyl alcohol $\ce{CH3OH}$ are sold at the same price per kilogram, which would be cheaper for producing an antifreeze solution for the radiator of an automobile radiator?

I think that more depression in freezing point will result in a better antifreeze solution. It depends on molal freezing constant. How can I predict this?

You are on the right track when you mention molality.

Because freezing point depression is a colligative property of the solvent, in this case water, all you need to do is to find out which solute has more moles per kg, aka molality.

The equation for freezing point depression for a non-ionic solute is given as:

$$\Delta T_\mathrm{F} = K_\mathrm{F}\times b$$

Where:
$\Delta T_\mathrm{F} =$ freezing point depression
$K_\mathrm{F} =$ cryoscopic constant (dependent on solvent, not solute)
$b =$ molality, the moles of solute per kg of solvent

The key point here is that freezing point depression is linearly proportional to solute concentration (molality).

Since glycerol has a molecular mass of $\pu{91.1 g/mol}$, 1 kg will contain 11.0 moles of solute. Methanol has a molecular mass of $\pu{32.0 g/mol}$, so 1 kg will contain 31.2 moles of solute.

So, for any given price per kg, methanol will give you $31.2/11.0 = 2.9$ times as many moles of solute, resulting in a 2.9 times greater freezing point depression.  Note that you cannot calculate the actual freezing point depression without knowing $K_\mathrm{F}$.


Why does CO2 accept electrons? CO2 is a lewis acid as it can accept electrons. But as the octets of all atoms of CO2 are complete why does it accept electrons? Shouldn't it donate electrons and be a lewis base?

Consider that the carbon in carbon dioxide is bonded to two electronegative oxygen atoms. This will cause shift the electrons towards the oxygens and away from the carbon, making it electron deficient:

$\hspace{7.5cm}$

This makes the central carbon electrophilic, and explains the reactivity of the carbonyl moiety towards nucleophiles. Similarly, $\ce{CO2}$ undergoes addition reactions with Grignard reagents to form carboxylate anions.

$\hspace{4cm}$


How do catalysts provide alternative routes for reactions? According to collision theory, a reaction will take place between, say, two molecules, if the collision between the atoms has a sufficiently high kinetic energy (to meet the activation energy threshold) and have the correct orientation. If these two conditions are not met, the collision will not result in a reaction.

However, there is one way to make a reaction easier, and that's with the use of a catalyst. Catalysts do not lower the activation energy or increase the kinetic energy, as I've just learned trying to figure out this question, but they provide an alternative activation energy, lower than the initial one, that, if met, will allow the reaction to take place.

My question is this: what is the actual chemistry behind how catalysts provide this alternative route? What's happening with the molecules or atoms? How does it work?

For example, if I wrap zinc in oxidized copper wire and mix it with some hydrochloric acid, the oxidized copper will act as a catalyst for the reaction between zinc and hydrochloric acid. I'm not sure how this takes place, though.

Thanks.

So catalysts provide an alternative pathway by allowing a new intermediate to be formed.
For example consider the reaction of $A$ and $B$ to form $C$. this means that
$$\ce{A + B -> C}$$
Say this reaction has  an activation energy of $x$ $kJmol^{-1}$.

This reaction could also be possible with a catalyst $X$. This means that a new possible route for the reaction could be (there a multiple routes the new reaction could take this is just a simple example):
$$\ce{A + X->AX}$$
$$\ce{AX + B->C +X}$$
Now you have ended up in the same position using a catalyst and when not using a catalyst, seen as you add both reaction and cancel out common species then you get back to the equation of:
$$\ce{A + B -> C}$$
These two reactions with the catalyst will have their own activation energies. Say these two reactions have activation energies $\pu{p}$ and $\pu{q}$ $\pu{kJmol^{-1}}$ and that $\pu{p>q}$ (doesn't matter if $\pu{p}$ applies to the first or second reaction) then if $x>p$ a particle needs less energy to proceed via the route with the catalyst compared to the route without the catalyst.

As a particle now needs less energy for the product $\pu{C}$ to be formed, more particles will now be able to take part in the reaction thus increasing the amount of successful collisions per second which is the same in an increase in rate. This can be seen nicely in a general Maxwell-Boltzmann distribution of molecular energies as the new route will mean there's a lower activation energy overall meaning a lot more particles can react.


Which effect does adding hydrogen chloride have when determining the concentration of sodium hypochlorite in bleach via titration have? Many of you might be familiar with this and I do understand this completely, however I do have a question.

If you are familiar with this, then this is the equation and this is the question for it.

$$\ce{ClO^{-} ~(aq) + 2I^{-} ~(aq) + H2O ~(l) -> Cl- ~(aq) + I2 ~(aq) + 2 {}^{-}OH ~ (aq)}$$

So for the equation above I understand everything except that in the titration process my instructions say to add $\ce{HCl}$ into the solution, I do not understand how adding $\ce{HCl}$ will have any impact.

I will provide an accurate reference later.

The reaction is in equilibrium and if you don't add any acid, then some of the $\ce {ClO-}$ does not react. Adding acid will change the equilibrium to the right so most of the $\ce {ClO-}$ reacts.


Why don't all electrodes produce the same results in the electrolysis of potassium iodide? While in lab, I accidentally used two paper clips while trying to electrolyze potassium iodide and all I saw was was phenolphthalein turning pink at the cathode and no brown at the anode.

I asked my instructor to allow me to repeat the experiment with pencil lead, which produced brown and pink, aluminum foil, which only produced pink, silicon which produced brown and pink, and copper, which produced brown and pink.

Then I used one paperclip at the cathode and a piece of pencil lead at the anode and I found that the anode turned pink and the cathode turned brown. Then I flipped the pencil lead and paper clip (anode to cathode and vice versa), and only the cathode (pencil lead) turned pink and there was no brown at the anode (paper clip).

My original hypothesis was that the increased resistance in the silicon and graphite lowered the current since the battery provided a constant $\pu{9V}$, but being able to get brown and pink with copper invalidated that hypothesis.
This has been bugging me all day, and my instructor says she does not know why this is happening.

To distill down your observations: you always saw pink at the cathode, but only some materials produced a brown color when used as the anode. This means that the anodic materials that didn't produce brown didn't oxidize the iodide to iodine.

So, the question is what was happening at the anode when the iodide was not being oxidized. The answer is simply that the anode material itself was easier to oxidize than the iodide in those cases, so the anode itself was oxidized and the iodide was not, so there was no brown color produced.


Am I correctly showing a hydride shift on the alkyl halide? I did this quiz and I forgot the hydride shift.

With the hydride shift, would these be the right answer?

When you solvolyze the starting bromide a secondary carbocation would be generated.  Since a secondary carbocation is not particularly stable, the molecule will explore other pathways such as hydrogen shift, alkyl shift, etc., in order to produce a more stable carbocation.  The figure below, explores these options for your starting bromide.  Realize that although I've drawn a discreet secondary carbocation that then undergoes rearrangement, in reality cation generation and alkyl or hydrogen shift may occur in a more or less concerted fashion.

In this case there are 3 options (actually there are 4 options, but the fourth one would produce a primary carbocation, so I left it out - can you find it?), a, b and c. Pathways a and c produce secondary carbocations; however pathway b, involving hydrogen shift, produces a more stable tertiary carbocation.  Capture of the tertiary carbocation from pathway c by solvent should produce the expected product.


Electronic density distribution in hydrogen bond I'm currently studying about the atomic configuration of electrostatic interactions. Concept is there that in case of interactions like in N-H...N-H (amides), the stability depends on the position of resonating electron between H...N nuclei (THE NATURE OF THE CHEMICAL BOND. II. THE ONE-ELECTRON BOND AND THE THREE-ELECTRON BOND - Linus. Pauling).

Given that hydrogen bonds formed between N-H...N-H or N-H...O-C are partial covalent in nature so the issue of directionality comes in effect.
I'm just trying to understand, from the view-point of the electron density shared by H...N to form the bond, that how (if possible) atomic configurations/geometry affects the electron cloud distribution ?

I'm basically trying to understand the hydrogen bond forming amides (currently interested in) from the electronic density distribution's point of view; to calculate the single-point-energy using dft.

Thank you.

There are a few descriptors in 'conceptual'  DFT just to characterize hydrogen bonds. E.g.: non-covalent interaction (NCI): note that hydrogen bonds are often consider as NON-covalent bonds, Density Overlap Regions Indicator (DORI).

One can also look at the  deformation density contributions originating from natural orbitals for chemical valence within ETS-NOCV theory.


Regioselectivity in opening of cationic three-membered rings The mechanism of electrophilic addition on alkenes often involves an electrophile attacking the double bond and causing an electromeric shift of electrons. The electrophile then forms a bridged cation if possible.

The carbon on the left side has an alkyl group which shows inductive effect (and hyperconjugation effect) which increases the electron density on the left carbon. The right carbon is not bonded by any such alkyl group. Therefore, the electron density on the left carbon must be higher than that on the right carbon.

Why does the nucleophile $\ce{Nu}$ attack the carbon which has higher electron density?

There are three factors that you should consider in nucleophilic attack.

Steric effects
Electronic effects
Stereoelectronic effects

Two of these are related to lowering the overall energy of the transition state. The third relates to how productive a collision can be.

The one that matters less here (or is already accounted for) is stereoelectronic effects. Notice that you proposed a backside attack for a $\mathrm{S_N2}$ reaction. Therefore, you have considered the proper alignment of orbitals, which is what is meant by stereoelectronic effects.

Steric effects government repulsion of large groups. In your example, attacking the more substituted side would result in greater steric repulsion since the $\ce{R}$ group is larger than $\ce{H}$. Steric arguments therefore argue for attack at the other carbon.

Frequently, the three factors do not say the same thing, and you need to use some intuition and/or experience to figure out which one dominates (in reality, the relative energies will determine the "winner").

In this case, it turns out with a carbocation and a negative nucleophile, perhaps not surprisingly, electronic effects dominate. There are actually two related electronic effects. One is that the partial positive is greater on the more substituted carbon (cf. Markovnikov rule). This carbon is hyperconjugatively stabilized (relative to the other carbon) by greater substitution. In addition, the greater positive charge also implies a weaker $\ce{C-E}$ bond, which further implies a lower energy $\ce{C-E}$ $\sigma^{*}$ orbital. We can then argue that the transition state energy will be lower given that the two interacting orbitals are closer in energy. And indeed, the electronic effects dominate this reaction giving the reactivity shown in your diagram.


When and why do hot material take heat from cold material? This is what I read on Chemwiki:

  Rudolf Clausius was a German physicist, and he developed the Clausius statement, which says "Heat generally cannot flow spontaneously from a material at a lower temperature to a material at a higher temperature." However, there is one exception to this statement and that is sometimes a hot material will take heat from the cold material, making the hot material warmer and the cold material colder.

Chemwiki does not provide a reference for the exception about heat transfer from cold material to hot material. How is this possible? What is an example of this happening?

The statement is wrong and/or poorly formulated. Heat cannot spontaneously flow from cold areas to hot areas, there are no exceptions known. It contradicts Clausius' statement of the second law and hence also the more general observation/statement that the entropy of an isolated system is increasing.

Of course you can drain heat from cold regions and transport it to hotter regions, but this should always go paired with some other (larger) increase in entropy somewhere else, e.g. by doing work like in a Linde refrigeration. An other example is that some salts have a very high entropy of solvation and can cool water to ~ -15 °C in a lab at room temperature. These salts can be/are used in practical chemistry for cooling a reaction mixture.


Query in Chemical Kinetics I have been recently studying chemical kinetics. I've understood what the rate law and the rate constant mean, and also that, kinetics depends completely on experimental observations, due to which we have zero order reactions as well as fractional order reactions.

Take for example, this: $$\text{Rate of a certain chemical reaction} = k [A]^p [B]^q [C]^r $$

Here, $[A]$, $[B]$ and $[C]$ are the molar concentrations of the three reactants involved, while $p$, $q$ and $r$ are the number of molecules of each respectively taking part in the reaction.

So, $\text{the order of the reaction} =p+q+r $.

Now, this $p$, $q$ and $r$ are completely determined experimentally, and cannot be determined theoretically.

How are these values determined experimentally? How does one find out how many molecules of each reactant are taking part in the reaction? Is there any special instrument to do this? If not, then how do scientists find the order of a reaction?

As an attempt to shorten out the answering area, consider this reaction: $$\ce {NO2 +CO -> NO +CO2} $$ Now, experimentally it has been determined that $\text {the rate of this reaction} = k[\ce{NO2}]^2$.  I want to know how have scientists determined that the rate depends on the square of the concentration of $\ce {NO2} $ and not on the other reactant.

N.B.: My question is how it can be experimentally determined, and not why it has to be experimentally determined.

Experimetally :- one reactant is kept constant and concentration of other is changed  and accordingly rate of reaction is measured...then dependence with thta reactant is measured.....and then other one  is kept  constant and ist one is changed then  similarly rate of reaction and thus dependence is measured....... We cant determine the actual dependence of rate of reaction on the concentration of reactants therotically ......we have tk calculate it and then we can formulate the rate law or order of reaction


Does trans-decalin really have a plane of symmetry? My teacher told me that trans-decalin (see below) is achiral due to the presence of both, a centre of symmetry and a plane of symmetry. But I could not spot the plane of symmetry untill now. Can someone point out the plane of symmetry in a diagram?

It's not easy to see from a diagram, because it distorts bonds and angles. I recommend building it with a balls-and-sticks model set. You can also use a molecular viewer to model it; there are a couple of open-source (or at least free) ones out there.

I have calculated the molecule on the DF-BP86/def2-SVP level of theory. The point group of the molecule is C2h. In the following model I have highlighted the plane of symmetry, and the rotational C2-axis. At their intersection is an inversion centre Ci.

(I needed to downscale this a lot, click on the image to get to a high resolution still. Images created with ChemCraft, assembled with GIMP.)


Lattice type from image What type of 2D lattice is depicted in the image shown below? Is it possible to determine without knowing the length of the unit vectors and the angles in between them? I'm confused because it seemingly doesn't fit with either monoclinic or orthorhombic description as $\mathrm{a \ne b}$ and $\mathrm{c\ne d}$

The picture is not a lattice, it is a periodic pattern. In a lattice, all lattice points are related by the translational symmetry given by the unit cell vectors. I can take any two lattice points and take the vector between them. Then, when I add that vector to any other lattice points, I should land on another lattice point.

This is not the case here. If I take the vector between the pair of closest points, and add the vector to the upper right point of the pair, I land on empty space.

So let's paint half of the points a different color. The points of different color are related by two-fold axis (located for instance at the center of two close pairs). The two-fold symmetry applies not only to the two closest points, but to the entire pattern, so it is called a crystallographic symmetry (as opposed to non-crystallographic symmetry).

This crystallographic symmetry imposes no restraints on the dimensions of the lattice. Other symmetry elements would impose restraints, for example in the presence of a four-fold crystallographic symmetry, the lattice would have to be tetragonal.

For the example at hand, the convention is to put the origin on one of the two-fold axes. It looks like you could have one unit cell vector along x and one along y, at approximately right angle (this is not imposed by point group symmetry, so it is "by accident" and you would have to measure carefully to see whether the angle is exactly 90 or maybe 90.0001 degrees).

To summarize, it is monoclinic with an angle that happens to be close to 90 degrees.

For specifics on the lattice types, see e.g. Wikipedia.

Thanks, @Ivan_Neretin, for the helpful comment.


How many electrons are there in the pi-system of cyanogen? I know that there are 18 electrons available for bonding in the entire molecule, and that 6 of these are used for sigma bonds. That leaves 12 left over. Which molecular orbitals best describe the locations of these electrons?

Here is the structure of cyanogen

Each carbon has 4 valence electrons and each nitrogen has 5, for a total of ((4x2)+(5x2)) 18 electrons as you said.  Again, as you noted, 6 electrons are involved in sigma bonds, the 2 C-N sigma bonds and the one C-C sigma bond. Each nitrogen also has a lone pair of electrons and that accounts for (2x2) 4 more electrons.  That leaves 8 electrons to fill the 4 pi bonds.  Just like acetylene with 2 pi bonds, the two carbon-nitrogen triple bonds in cyanogen also have two pi bonds, with each pi bond holding 2 electrons.


Figuring out if a surfactant is present I have recently starting distilling my own essential oils using a still that I built out of an old diesel barrel. Before the first run with the still, I cleaned the inside using soap and water. I distilled until I created 5 gallons of hydrosol. The issue is that it seems that the oil and water are emulsified (hope that is the right term, I don't have a background in chemistry). There appears to be a thin film of oil on the surface, rather than the "globs" I would expect. Another reason I believe this is an emulsion is that no matter how much I skim off the top, that film is always present, making me think that the oil is thoroughly mixed in the solution.I have done some research and it seems to be that there may be a surfactant present that is causing this.

Is there a way that I can test for the presence of a surfactant?

Does it sound like my issue is the presence of a surfactant or something entirely different?

Could the soap I used to clean it be the surfactant?

Thanks!

Update in Response to Answer:

The end goal is to use a few drops of the oil to enhance the smell of a facial product (beard oil), so the end product has to be safe for skin.

I don't have a diagram at the moment, but here is a smaller still that is basically the same as my setup, just much smaller. I am using a propane heat source to control the temperature.

To be honest, I don't know the expected structure of the oils. This is my first attempt at anything like this (I'm a software engineer that recently acquired an interest in chemistry!).

Thanks for the help!

There are multiple issues that need to be addressed before one can give a definite answer to this question. Let's start by what we are certain about:

1- The soap is not the issue (unless you haven't rinsed the barrel thoroughly with water). The remnant soap would be so dilute that should not cause any problem.

2- Since the film keeps forming there seems to be an equilibrium issue for your oil to separate. If you look into partition coefficients of the oils you should be able to estimate how much will remain in the water. Anyway, the less water you have, the more oil you can collect. (concentration factor pushes the oil out to keep the equilibrium) One idea is to add less water to your feed.

3- THEORETICALLY one way to enhance oil extraction is to use a nonpolar solvent. HOWEVER, I STRONGLY ADVISE AGAINST ACTUALLY TRYING IT AT HOME due to the flammable nature of most nonpolar solvents and health issues related to it. If you can find a nonpolar benign solvent (that is a big IF), mix it with your final water oil mixture, and then decant the water from the bottom and then distill the solvent out to get the oil as a residue.

4- Depending on the chemical structure of the oil, some of it will always remain in water. If they are not highly hydrophobic, it might be difficult to separate them from water anyway without using an assistant solvent.


When a candle burns, why does the top of wick glow if bottom of flame is hottest? Totally an elementary question.

Staring at a candle, it appears that the bottom of the wick is dark whereas the top glows. However the bottom of the flame (the blue) is the hottest.

Is the reason for this that the concentration of liquid wax is greater at the bottom, offsetting the greater temperature at the bottom?

The wick temperature does not have to be the same as the flame temperature.The flame is hottest at the bottom, but the wick is hottest at the top.

For a candle, the wick burning isn't the intended purpose of the wick; light comes from burning wax (more generally: fuel), you want to burn the wax not the wick. Rather the purpose of a wick is to help fuel evaporate by soaking up wax and allowing the radiant energy from the flame to heat the wax causing it to evaporate and burn also.

As wax travels up the wick, it evaporates and less wax is in the wick the further up you go. Eventually the wax dries up and the radiant energy is heating a wick without any wax. Eventually the wick gets so hot at the tip, that it will glow due to black-body radiation.

In summary: though the blue is the hottest part of the flame, the wick can evaporate wax to cool towards the bottom. There is no wax at the top and thus as the radiant energy of the flame causes it to get hotter until it starts to glow.

Extra: I will note that the top of the wick does not burn when lit as the gas around the wick has too little oxygen to burn (unless the wick slumps out side the flame). When you blow out the candle the low oxygen environment of a flame is gone and thus the hot wick will smoulder.


Rearrangement of Grignard reagent formed from exocycloproyl vinyl bromide First, I prepared the grignard reagent then performed nucleophilic addition to the carbonyl carbon  and then did hydrolysis.

So basically in my product, $\ce{-Br}$ is replaced by $\ce{-CH2OH}$

But answer given is:

Why is my mechanism wrong?

Grignard reagents are capable of rearranging, especially in this situation where there is significant strain. The specific vinyl bromide you drew here has been studied: J. Org. Chem., 1974, 39 (10), pp 1411–1416. The table of data shown below is hard to read without including more of the text, but it was found that at 66 C in THF, after 5-10 minutes, 41% of the initial Grignard had undergone rearrangement, so this is a fairly likely process.

A potential mechanism for the rearrangement is:


How to interpret intervals on a FT-IR spectrogram of soil samples? I am studying to analyze the IR spectra of soil samples. And I'm interested in the question - how to interpret the intervals in (1750..1820) and (3020 ... 2820) of the two samples on the graph?

The attached file - 1200..600 spectra of soil samples taken on Nikolet 6700 FT-IR Thermo Scientific in DRIFT mode and particular clippings of (1750..1820) and (3020 ... 2820) spectra.

UPDATE
Scan method – FT-IR DRIFT 1 200 cm-1 .. 6 000 cm-1

Red line – paleosol untreated (only dry the samples and sieve it threw a 2 mm mesh)
Blue line – weathering sandy soil untreated (only dry the samples and sieve it threw a 2 mm mesh)

You are almost certainly looking at bands corresponding to functional groups in the humic acid fraction of your sample.

The band at 1750 - 1820 cm$^{-1}$ is likely the $\ce{C=O}$ stretching mode in $\ce{COOH}$ groups in humic acid, but this band is also used to differentiate the fulvic fraction from the humic fraction (see quoted text below). The band at 2820 - 3020 cm$^{-1}$ is likely aliphatic $\ce{C-H}$ stretching.

The composition of humic and fulvic acid fractions in soils varies quite a bit, and how the samples were prepared can significantly alter the composition of the fractions as well. Keep that in mind when interpreting your results. Your bands appear to be shifted from those in the reference I chose to cite, but that is to be expected depending on the composition of your soil and the organic matter contained therein (and can also indicate greater or lesser degrees of $\ce{H}$-bonding, resulting in red/blue shifts). The soil organic fractions referenced in the study below are sandy loams from Ibri, Oman, in an agricultural region. Your samples will differ in composition based on geography, geology, climate, and other factors.

Reference: Helal, Murad, and Helal. Characterization of different humic materials by various analytical techniques. Arabian Journal of Chemistry 4(1), 51–54 (2011) (link to free PDF from ScienceDirect):

  The IR spectra of the three humic fractions are shown in Fig. 2. They have a diversity of bands more or less typical to those distinguishing the humic materials (Stevenson, 1994 and Aiken et al., 1985). Major absorption bands are in the regions of 3400–3300 cm$^{-1}$ ($\ce{H}$-bonded $\ce{OH}$ groups), 2940–2900 cm$^{−1}$ (aliphatic $\ce{C–H}$ stretching), 1750–1720 cm$^{−1}$ ($\ce{C=O}$ stretching of $\ce{COOH}$), 1620 cm$^{−1}$ (aromatic $\ce{C=C}$, $\ce{COO−}$, $\ce{H}$-bonded $\ce{C=O}$), 1280–1230 cm$^{1}$−1 ($\ce{C–O}$ stretching and $\ce{OH}$ deformation of $\ce{COOH}$) and 1040 cm$^{−1}$ ($\ce{C–O}$ stretching of polysaccharide or $\ce{Si–O}$ of silicate impurities). The spectra evidently show predominance of $\ce{OH}$, $\ce{COOH}$ and $\ce{COO−}$ groups which are the most characteristic features of soil humic materials. It is clear from the spectra that fulvic acid is characterized by stronger absorption near 1720 cm$^{−1}$ which implies the high carboxylate capacity, and that of humic acid is stronger than that of humin. This is exactly the case in the results of potentiometric titration (Helal, 2007). The spectrum of fulvic acid is also characterized by the absorption at 1400–1390 ($\ce{OH}$ deformation and $\ce{C–O}$ stretching of phenolic $\ce{OH}$). The results of IR spectra indicate that fulvic acid is more aliphatic, and humic acid and humin are more aromatic. It is obvious that the IR results are in good agreement with the other characterization findings.

Figure 2, from the above-referenced article. FA is fulvic acid and HA is humic acid.


How many Ga atoms are connected to one As atom in solid gallium arsenide? In Boylestad and Nashelsky's Electronic Devices and Circuit Theory, the structure of GaAs is shown as follows:

Arsenic is shown making five covalent bonds to neighbouring gallium atoms.

However, I recently also saw an alternative depiction in Razeghi's Fundamentals of Solid State Engineering:

Here, arsenic is shown making four covalent bonds to neighbouring gallium.

Which source is correct? Please give appropriate references or experimental proof.

It is quite well established that stoichiometric gallium arsenide (CAS 1303-00-0) forms a zincblende analogous structure. It consists of two stacked face-centered cubic lattice systems (Ioffe Physical Technical Institute database). See also Uncle Al's Answer. In other words, the arsenic forms the lattice and the gallium fills half of the tetrahedral positions (and vice versa). Therefore every gallium has four arsenic nearest neighbours.
This does not necessarily mean that these are bonds. In this case however, the bonds are best described as covalent (having a somewhat ionic character).
The Britannica Online Encyclopedia (Academic Edition) (You might need access provided through an educational institution) states:

Besides the elemental semiconductors, such as silicon and germanium, some binary crystals are covalently bonded. Gallium has three electrons in the outer shell, while arsenic lacks three. Gallium arsenide (GaAs) could be formed as an insulator by transferring three electrons from gallium to arsenic; however, this does not occur. Instead, the bonding is more covalent, and gallium arsenide is a covalent semiconductor. The outer shells of the gallium atoms contribute three electrons, and those of the arsenic atoms contribute five, providing the eight electrons needed for four covalent bonds. The centres of the bonds are not at the midpoint between the ions but are shifted slightly toward the arsenic.
Such bonding is typical of the III–V semiconductors — i.e., those consisting of one element from the third column of the periodic table and one from the fifth column. Elements from the third column (boron, aluminum, gallium, and indium) contribute three electrons, while the fifth-column elements (nitrogen, phosphorus, arsenic, and antimony) contribute five electrons. All III–V semiconductors are covalently bonded and typically have the zincblende structure with four neighbours per atom. Most common semiconductors favour this arrangement.

This is still a fluorishing field for researchers in the field of materials design semiconductors, semiconductor lasers, computational physics/forcefield development, and more.
The book "Electronic Devices and Circuit Theory" clearly gets it very wrong. I had a look into it and noticed they also argue with the Bohr model, which might not be a very good starting point at all. However, the focus of the book is clearly not solid state chemistry, and therefore it might fit the purpose.


Fischer Projections confusion A bit of trouble I am having with grasping the concept of fischer projections for long chained molecules. Take any non-cyclic carbohydrate from a tetrose and upwards and make an actual model of it and what we see is that the carbon backbone follows a zigzag pattern. However, according to the fischer projections, if we take any chiral carbon the vertical bonds go into the plane of the paper and the horizontal out of the plane of the paper. With a real-life model which follows a tetrahedral structure at each carbon atom and a zigzag pattern this just doesn't seem to work. What is going on here??

The carbon chain of a fischer projection is drawn vertically.

Yes, the vertical and horizontal lines depict 'into the page' and 'out of the page' respectively.

However, when you have a carbon chain with more than 3 carbons, each down carbon on the backbone which would normally be projected away from you (according to the above rule) has to be rotated. Essentially, this creates the 'zig-zag' you would see in other projections (e.g. Natta Projection).


Can oxygen burn in hydrogen? Most combustion happens involves the fuels burning in an oxidant. But there seems to be no reason that it cannot happen in the opposite way. Are there any experiments on oxygen burning in hydrogen? Or some oxidant in solid state (like $\ce{KMnO4,KClO3,}$...) burning in a reductive gas?

Could you be more clear on what 'burning in' something means? The chemical reactions require both reactants available, so they need to mix - in what sense is Hydrogen mixing and reacting with Oxygen any different that Oxygen mixing and reacting with Hydrogen?
In some sense, we do tend to control these reactions via the fuel, not the oxidant (oxygen free environments are difficult to reproduce on Earth's surface). However there may be some examples of controlling the mixture via restricting the oxidant in a fuel-rich environment... even though I can't even get through the sentence without that sounding wasteful.
Nitrous Oxide
Would you consider the example of Nitrous Oxide injection in an internal combustion engine an example of this? Pre-nitrous, there is compressed atmospheric oxygen and hydrocarbons at a set ratio, and to further increase the burn a nitrous system will inject either NO alone (dry system) or NO and more hydrocarbon fuel (wet system) to add more oxidants that would be available through the 21% Oxygen in the atmosphere.

The Chemical: https://en.wikipedia.org/wiki/Nitrous_oxide
Internal Combustion application: https://en.wikipedia.org/wiki/Nitrous

Candles
How about a candle? There's an abundant fuel source (wax), and we frequently control the shape and characteristics of the burn with shaped glass, not to mention put it out via starving it of oxidant.
In either case, if you were to, say take a lump of solid oxygen (nevermind the practical difficulties of dealing with temperatures below 54 Kelvin) and maintained a burn with Hydrogen, what you'd be likely to find would be combustion only happening at the edges, where the oxygen was evaporating and mixing with the reduction agent, similar to lighting a can of gasoline on fire (only the vapors are flammable).


What stopper will stop HCl fumes? The local hardware stores actually sell chemicals in recycled beer bottles with plastic stoppers as lids. The bottle of hydrochloric acid I bought is clearly losing gas to diffusion through the plastic, or a seal that's not tight enough--it smells bad, and the beer bottle's metallic label is slowly corroding.

How can I best seal this bottle? Some ideas are:

Try to mold a stopper from silicone. Use a (synthetic) cork, whatever's used for homebrewing. Line the stopper with a few layers of thin polyethylene or polypropylene (from consumer materials, such as air-filled plastic bags for packing).

Or could I limit the diffusion by putting a plastic bag around it? This is tough to answer experimentally (without a lab) because resisting physical degradation isn't the same as resisting diffusion.

it smells bad, and the beer bottle's metallic label is slowly corroding.

That's, unfortunately, very common. At one of my former labs 'acid' cabinet  was corroded from within. For this reason it is recommended to keep such bottles under fume hood.

AFAIK, common stoppers are made from polyethylene or polypropylene. At one of my former labs concentrated hydrochloric acid was stored in stock plastic canister, similar to that used to store petrol (on the other hand, both nitric and sulfuric acids were stored in 20l glass bottles, even though the stoppers were made from same polymer).  I HIGHLY recommend to by (or order by mail) dedicated chemical bottles with caps, though. They are usually fairly cheap and pretty universal.


Writing the chemical equation given molarity If I am given the molarity of two substances and I know how they react, how can I write the balanced chemical reaction? i.e. experimentally I found:

$\pu{1.0 M}$ $\ce{NH_3(aq)}$ mixed with $\pu{0.1 M}$ $\ce{BaCl_2(aq)}$  yielded no precipitate when $\pu{1ml}$ of each were mixed together. My unbalanced chemical equation so far looks like:

$$\ce{N^{-3}H3+(aq) + (1/10)Ba^{+2}Cl2-(aq) -> N^{-3}Ba^{+2} + Cl^{-2}H+}$$

My logic being: since the solutions are administered in equal volumes, their molarity determines the molar coefficients. Right off the bat I can tell that $\ce{NH}$ is not soluble, but the other compound I think is, which would contradict my observations. Either way, I am incredibly lost.

Have you considered the possibility that your reaction is wrong.  where, for example, is the barium on the product side.  I'll give you one hint here:  barium chloride dissociates in water to barium and chloride ions.  Ammonia reacts with water to produce some ammonium ion and the same amount of hydroxide ion.  However, barium hydroxide is totally ionized.

Can you take it from there?


Is this stoichiometric question answerable? (unknown variable) This question kind of borders on the fields of both semantics and math, as well as chemistry, but I still believe this is the right place to ask.
On a test I took, a question was posed as such:

Via titration we are adding a solution of 25 ml, with 0.5 M NaOH, to 25 ml of HCl. What will the concentration of the acid be?

Now, this question seemed a bit strange to be, and it was definitely new. I'm used to calculating concentrations after dilution and other stoichiometric calculations, but this question seemed like a misworded version of those other questions I've grown familiar to.
So, I asked my teacher if she could clarify. First I asked whether the 25 ml of HCl was pure HCl, or a water solution with HCl in it. She said it was pure HCl. Then I asked if she was looking for the concentration of the acid before ANY reaction between the NaOH and the HCl had occured. To this, she also said yes.
The options were:

0.25 M
0.5 M
0.75 M
1 M

I answered 1 M, my reasoning being that 25 ml of HCl is the same volume as 25 ml of water with NaOH. Obviously, this reasoning is quite flawed, but I didn't have anything better. The answer was wrong. However, when my teacher told me what was meant by the question afterwards, her explanation did not align with the question, nor her previous clarifications. She said that this was a simple case of:

$$C_1 \times V_1 = C_2 \times V_2$$

Where $C_2$ was the unknown, since the concentration of HCl was not specified. I don't see how this makes sense. All the question is saying, is that there is a 25 ml container, with HCl in it. It doesn't even say that this 25 ml consists of a water solution with HCl in it, but a more knowledgeable chemist would've known that HCl cannot exist outside of a water solution (that's what my teacher said at least).
Then the the question says that 25 ml of water, with a concentration of 0.5 M of NaOH is added. And from those two values, I'm supposed to find the concentration of HCl. That basically means:

$0.025$ l of H$_2$O$ \times 0.5$ M NaOH(aq) $= 0.025$ l of water $\times x$ M HCl (aq)

Which would then make the answer 0.5 M. Maybe she wanted me two add the two volumes, which would make sense:

$0.025$ l of H$_2$O$ \times 0.5$ M NaOH(aq) $= 0.05$ l of water $\times x$ M HCl (aq)

Which would then make the answer 0.025, which wasn't even an option.
But even if it was, it still wouldn't make sense to me. Just because you have a solution with a given volume and an unknown concentration, adding another solution with a known volume AND concentration won't just magically let you know what the unknown concentration in the original solution is, now in this new solution. Or am I wrong?

With a lot of assumptions, 0.5 M is a reasonable answer.
Here are the holes you have to fill in.

If it is a titration of NaOH and HCl, it is probably an acid/base titration where we stop the titration when the pH is neutral.

Titrations are often performed to figure out the concentration of a solution. You would titrate a smaller sample of the solution to figure out the concentration. Technically, the concentration of HCl after the titration is zero because all of it is supposed to react as you reach the endpoint of the titration. So it makes sense to ask about the concentration in the original solution.

The stoichiometry of the neutralization of HCl and NaOH is one-to-one:
$$\ce{NaOH(aq) + HCl(aq) -> H2O(l) + Na+(aq) + Cl-(aq)}$$

Of course, it is silly not to give any context when posing such a question on a quiz. I won't comment on the statements you attribute to your teacher other than to say that hydrochloric acid always contains water (pure HCl is a gas under ambient conditions).


Make a silver zinc battery I have a shaver that runs off a rechargeable battery that is dying. Would it be feasible to make a silver zinc battery and use it to replace the existing battery? Cost is not an object, anything less than $2500 I would consider doable. I have a machine shop and a small chemistry laboratory with the standard equipment and glassware, including high vacuum capability, a centrifuge and simple glass blowing capability.

I found a book on silver-oxide zinc chemistry and battery design, but it is $500 and I don't want to spend that if the information can be obtained just as easily elsewhere.

I have tried to find commercial options, but had no luck. Most silver zinc batteries seem to be just for large (multi-million dollar) military or satellite applications. Sony makes a line of silver oxide primary cells for hearing aids, but these are not rechargeable. There is a company called Ultralife that makes medical and military batteries and might have something viable, but before I call them I wanted to check out the opinion of the experts here.

This turns out to be very difficult to do for two reasons. One problem is that the voltage of a silver cell is different from that of Ni-Cd cell, so it would require a specialized, multi-cell configuration to emulate the voltage characteristics of the Ni-Cd rechargeables.

The other problem is that silver cells generally have a sophisticated frame inside of them that is produced by an intricate high-temperature welding process. To duplicate this process and produce a suitable frame would require a significant amount of experimentation and work.


How to measure the percent weight of NaCl in a potato chip I want to determine the weight percentage of sodium chloride in a potato chip.  I can see from this answer that atomic absorption spectroscopy (AAS) can be used to quantify sodium chloride in a solution.  I propose the following measurement scheme:

Dissolve the potato chip in a sample of water with known mass
Homogenize the solution until it has a uniform composition
Draw 10 aliquots from this solution
Use AAS to measure the weight percentage of salt in each aliquot
Use the mean and standard deviation of the weight percentages from
the 10 aliquots to generate an estimate of the true weight
percentage with a confidence interval

My questions:

1) What is the best way to get the mass of the aliquot that belongs to the potato chip?  I want to minimize measurement uncertainty.

2) There could be compounds other than sodium chloride that contains sodium or chloride in a potato chip.  I fear that the AAS won't detect sodium chloride, but sodium and chloride ions separately.  What is the best way to measure the concentration or mass of sodium chloride with AAS?  If AAS is not the best way, can you name an alternative method?

I'm somewhat skeptical that atomic absorption spectroscopy (AAS) would work as you outlined the procedure. The "solution" would have starches and fats floating around in it.

(1) I'm not sure that you could get a homogeneous solution.

As far as the homogeneous solution, think of some parts of the solution settling. So what happens if the solutions sit for an hour before going into the AAS?

You could test this by homogenizing a sample and analyzing it immediately, then analyze it after an hour. Do you get the same result?

(2) Even if you did I'm not sure that all the gunk would burn correctly (uniformly) in the AA flame.

Using a plasma spectrometer would solve this problem.

(3) Depending on what sort of accuracy you want there is also the problem of moisture in the chip. How "dry" is a potato chip?

(4) There is a difference between salt added to the chip and $[\ce{Na^+}]$ and  $[\ce{Cl-}]$ in the chip. I'd guess that $[\ce{Na^+}] \approx [\ce{Cl-}]$ in a biological system, but I'm sure that it isn't exactly $[\ce{Na^+}] = [\ce{Cl-}]$.

Such are the joys of method development in analytical chemistry.


How do energy changes in reactions avoid violating the conservation of momentum? As mentioned in this question, the idea of chemical potential energy explains how the thermal/kinetic energy of particles can change in a reaction; energy is taken in / released by the bonds broken / formed.

However, consider the reaction A -> B + C occurring in a closed system of just these single particles.
A reaction such as this is often endothermic. Therefore, the kinetic energy of the products must decrease. However, it is impossible to remove kinetic energy from the system without changing momentum (unless B and C are travelling in opposing directions to begin with, which is not the case. If they split in opposite directions, that will have already resulted in a KE gain).

How can this happen without violating the principle of conservation of momentum?

EDIT: Here is an explanation showing that it does not work if B and C travel in opposite directions:

How do you measure temperature? The thermometer has to be in thermal equilibrium, so you need to send the thermometer with the object. But that means that temperature, for our purposes, is going to be invariant between different reference frames. No matter how you're moving, the reading on the thermometer is the same.

This is clearly not the case with translational kinetic energy of the bulk material. In the rest frame of the object moving with constant velocity, the kinetic energy is zero.

Therefore, you should not mix up the thermal kinetic energy of the system with the translational kinetic energy of the system. They can't be the same thing. In other words, when the exothermic system loses energy, you don't get to subtract that from the translational kinetic energy.

See also these discussions: here and here.

I would also like to point out that you've just described a rocket. And rocket science, while complicated, is still obeying conservation of energy and momentum.


Which relatively simple molecules violate the octet rule? I'm learning how to draw Lewis diagrams. Everything I've read emphasizes the octet rule. However, to the best of my knowledge, that rule only applies to elements in the first three periods.

On our exam, I expect our professor will ask us to draw some Lewis diagrams of molecules whose atoms do not comport with the octet rule. However, I haven't seen any diagrams of these; it'd be useful to get a sense of what they look like so that my gut doesn't disapprove of what will otherwise be strange looking diagrams.

What are some relatively common molecules that comprise atoms that comprise a great number of valence electrons?

There are 3 types of octet rule "violations" or exceptions

molecules with an odd number of electrons, such as nitric oxide

(image source)

molecules with less than 8 electrons around an atom, $\ce{BeCl2}$ and $\ce{BH3}$ serve as examples

(image source)

molecules with more than 8 electrons around an atom, such as $\ce{PCl5}$ or $\ce{SF6}$

Take a look at the image source links for other examples as well as some practice problems.

The bonding around electron deficient atoms such as the boron in $\ce{BH3}$ is best explained in terms of 3-center 2-electron bonds (reference; SE Chem example_1, SE Chem example_2)).

Sometimes, arguments based on d-orbital involvement are advanced to explain the bonding in atoms surrounded by more than 8 electrons.  However d-orbital involvement in non-transition metals is now generally considered to be unlikely.  Instead hypercoordinated (or hypervalent) bonding in which 3-center 4-electron bonds are formed has become a more generally accepted explanation.  A number of examples have previously been discussed here on SE Chem (for example, see here).


How metal-plating on plastics is performed? Metal coated plastics are commercially abundant now-a-days. I want to know how they're plated.

Before posting, I did a quick search and found a nice link. It refers to this process as metallization, although I'm not convinced by the term. Is it right ?

That site already provided the basic information. However, it would be great if someone can mention a few examples(which metal/plastic used) for the various methods deployed.

I couldn't comment on the name.
I am familiar with vacuum deposition, but i have only come across deposition on silicon, not plastic. the substrate ad metal are put in a vacuum chamber, the vacuum means that now the metal can be evaporated at a much lower temperature, and is then deposited onto the substrate.(see link)

I did an experiment of electroless plating of plastic with nickel. I can't say what plastic it was, probably PS or PP. usually electro plating requires electrodes and current to flow between them. electroless plating uses a electrochemical reaction in which a metal precipitates. Due to the treatment of the plastic the metal prefers to precipitate onto the plastic (this is a very simplistic information see liks for more information if you're interested)

the method I used in my experiment:
the plastic is cleaned by diping it in a base bath (sodium carbonate, trisodium phosphate, sodium hydroxide), this is mostly to remove oily debris .

Etching: ...etches the plastic to create microscopic pores that improve
adhesion of metal and following steps
Sensibilization: plastic is
dipped in ti solution, keeps the palladium at its reduced state (next
step)
Activation: plastic is dipped in a palladium solution, Palladium
catalyses the formation of nickel particles.
Plating: since there is no current flowing another particle has to be oxidized for the nickel to be reduced. in this case the reaction used is oxidation of  hypophosphite to phosphite

The result was a very nice metallic finish.

more information:

Electroless nickel plating on Wikipedia
plating kit
plenty of information on electroless deposition
evaporation-deposition


When will organometallics act as base or nucleophile? I would like to ask if the hydrogen next to carbonyl group is acidic,and organometallics are strong base, when will it act as a base and extracts the acidic proton and when will it attack the carbonyl carbon?

Also I would also like to ask, for bases when will it attack the acidic proton to form a enolate and when will it attack the carbonyl carbon directly to form an hemiacetal/hemiketals?

Grignard is a base as well as an nucleophile. When there's a strong electron deficient group it'll attack there, when there' a very acidic group it'll attack there. For other bases it depends upon their relative basicity and nucleophilicity. BAsicity is thermodynamically related while nucleophiicity is kinetically decide. They're silent about each other. You need to practice and gather examples and an experience too :D


Does the salinity in seawater affect the process of alkaline hydrolysis? I am an architecture student doing a project on crematorium spaces. I am looking at resomation or alkaline hydrolysis as an alternative method of disposition of a body. From what I gather from research, the funeral parlours offering the resomation service uses 80 gallons/ 300 L of water for each body. I am just wondering if in theory such an enormous amount of water could be replaced with seawater? Or does the salinity or pH of seawater inhibit the process?

I believe that desalination would be necessary as the salinity of the seawater would interfere with the earth metal salts already utilized to perform the process of alkaline hydrolysis.


Why are snowflakes shaped as they are? (It was the Winterbash icon that got me thinking on this.)

Snow flakes take on a variety of intricate shapes, but what processes are responsible for this?

(Got the image off of Wikipedia)

I mean, I wouldn't have raised this question if those water droplets, descending towards earth, assumed a... less fancy shape. Not only are these minuscule structures very intricate, they also display a high degree of symmetry (with most having multiple planes of symmetry).

Now, as I understand it, all processes proceed so as to maximize the "randomness" of its constituent particles. (Oversimplified version of the Second Law of Thermodynamics, yes, I know... just don't chew me out in the comments section...)

This Law can easily be observed in, and verified by, natural processes.

Now the formation of snow is a natural process, agreed? The way my brain sees it, is that water droplets ought to freeze into random, and by virtue of its "randomness", highly unsymmetrical shapes. But this is not the case here!

For some confounded reason, the shapes assumed by snowflakes are eerily symmetrical. So eerie, that I'm surprised it hasn't attracted the attention of conspiracy theorists yet, as in the case of crop circles. Now I'm not saying you don't find symmetry in nature, there're plenty of examples out there. But what I am trying to point out here, is that you don't typically see symmetry to such a high extent as in the case of snowflakes.

There are innumerable natural processes that give rise to weird shaped stuff, but I'm not really sure what makes snowflakes so different. However, I did notice a few points of difference between the formation of snowflakes and other natural processes:

1) As water droplets/snow descends, it's subjected to significant temperature changes as it passes through different layers of the atmosphere.

2) As water droplets/snow descends, it's subject to a significant increase in pressure as it passes through different layers of the atmosphere.

So, I reiterate my question.

Why are snowflakes shaped as they are?

Spin-off question:

I heard (not sure where) that each snowflake assumes a unique shape. How true is this?

Now, as I understand it, all processes proceed so as to maximize the "randomness" of its constituent particles. (Oversimplified version of the Second Law of Thermodynamics, yes, I know... just don't chew me out in the comments section...)

Yes. True.

  This Law can easily be observed in, and verified by, natural processes.

Sure. Still with you.

  Now the formation of snow is a natural process, agreed? The way my brain sees it, is that water droplets ought to freeze into random, and by virtue of its "randomness", highly unsymmetrical shapes. But this is not the case here!

Nope. Let's consider a counter argument: chemical bond formation between atoms is a natural process, and natural processes seek to maximise entropy. Thus, all molecules should take on highly unsymmetrical shapes, shouldn't they?  The answer is obviously no. Same applies for snow flakes. First of all, "order" and "symmetry" aren't precluded by thermodynamics. Molecular and crystal structure emerges within the realms of laws of thermodynamics and quantum mechanics.

Like I said in the comments, the precise discussion of how a snow flake's shape merges will be quite involved. There will be a lot of variables that affect the growth of snow crystals: temperature, vapour pressure humidity, wind speed and direction etcetera.

The final shape of the snow flake would depend on the precise path it followed, and the conditions it experienced as it descended, and is impossible to predict/calculate.

Of course, the underlying six-fold symmetry comes from the arrangement of water molecules in the ice crystal lattice.

Also, there are no natural processes coordinating the growth of the arms to be perfectly symmetrical; their growth is governed by simple statistics.

The arms of a snow crystal all grow independently, and grow under randomly fluctuating conditions; however, the effects of these fluctuations are somewhat averaged over the entire snow flake as it tumbles through the sky, and you end up with a somewhat regular shape.

There is in fact, asymmetry in snow flakes. Just take a close, hard look at a photograph

Sure there is some symmetry: the snow flake has 6-major arms, and then each arm has sub-branches, and each sub-branch branches out further. The overall branching pattern is the same, because of random averaging--the snow flake randomly rotates and reorients as it falls, each arm feels (on average) the same conditions. But are each of the sub-branches or the sub-sub-branches identical in shape, number, size etcetera? No.


What are the requirements for FRET? I have a question about fluorescence resonance energy transfer (FRET), and the requirements needed for this technique.

Q: Which of the following need to be met for FRET?

A) Strong overlap between the donor emission spectrum and the acceptor absorption spectrum
B) Molecules must be $<10\,\mathrm{nm}$ apart
C) Molecules must be $>10\,\mathrm{nm}$ apart
D) Low quantum yield of the donor
E) Small extinction coefficient of the acceptor

I'm aware that A and B are true, and C is not a requirement. But what about a low quantum yield of the donor or a small extinction coefficient of the acceptor? Do the size of these numbers really matter?

A low quantum yield would indicate that there are efficient intra-molecular non-radiative decay pathways, likely internal conversion or spin-orbit coupling. These decay pathways may be faster that FRET, rendering FRET essentially a non-competetive process.

A small extinction coefficient would indicate a small transition dipole moment. If the transition dipole moment is small it is difficult for the FRET partner to transfer energy since it is a weak transition resulting in a small coupling  between the two chromophores.


Does exothermic solvation mean solute is more soluble at low temp? (I'm surprised this question or a very similar one hasn't been asked, so I wouldn't be surprised if I just failed to find a question with my search that does answer this question.)

I dissolved some $\ce{K3PO4}$ in water a while back, and the solution became very hot. I commented on this to my labmate, and he said "So that means it's more soluble at low temperature, right?" I evaded response because I disagreed but didn't have a thought-out response yet.

I think he thought of the reaction as:
$$\ce{K3PO4 -> 3K^+ + PO4^{3-} + heat}$$

But I think that's missing an important step. I think it's more like this:
$$\ce{K3PO4 -> 3K^+ + PO4^{3-}}$$
$$\ce{PO4^{3-} + H2O <=> HPO4 + OH- + heat}$$

Phosphoric acid's third pKa is pretty high, so the second reaction is pretty significant. This, I would say, is why this particular reaction is so exothermic. So, in this case, I don't really think the heat plays into the solubility of potassium phosphate.

Sometimes, though, solvation comes with a positive change in temperature. For instance, the enthalpy of solvation of sodium iodide is $\mathrm{-7.53\ kJ/mol}$. Sodium iodide, though, is still more soluble at high temperatures. I could be missing something here, though.

In the cases where it is specifically the solvation reaction that causes a temperature change, is it the true that lower temperatures always increase solubility of salts with exothermic solvation reactions (and vice-versa)?

I'd say that the answer is no because the solvation should be dependent on $\Delta H - T\Delta S$ and, since the enthalpy change half of that formula is temperature independent, as $T$ goes up, solvation should be favored by any solvation that has a positive change in entropy. However, I can't really think of any solvation that has a negative change in entropy.

Think of this in terms of equilibrium constant.

$$K=e^{-{\Delta G\over RT}}=e^{-{\Delta H-T\Delta S\over RT}}=e^{-{\Delta H\over RT}}\cdot e^{{\Delta S\over R}}$$

Now it is pretty clear that exothermic reaction means negative $\Delta H$, which in turn means positive coefficient in $-{\Delta H\over RT}$, which consequently is a decreasing function, as is the whole exponential expression. The entropy change (or indeed the sign of it) does not matter.

The example of $\ce{CaCl2}$ is irrelevant, for anhydrous $\ce{CaCl2}$ is never the phase in equilibrium with the solution. Roughly speaking, one simply cannot just up and dissolve $\ce{CaCl2}$. In effect, there would be two reactions: first the hydration, which produces $\ce{CaCl2.6H2O}$ (that's the exothermic part), and then the dissolution of that.

The implicit approximation we relied upon is that our $\Delta H$ and $\Delta S$ won't change much throughout the dissolution process. This fails when it comes to insanely well soluble things (think of $\ce{NaOH},\;\ce{KOH}$) where by the time we reach equilibrium, we might even have more solute than water. Naturally, the properties of such solution would be quite different from those of pure water.


XRF is "non-destructive," but what actually happens to the sample? Editing for clarity and more details.

My colleague has been using XRF to quantitatively analyze 47-mm diameter Teflon membrane filters for metals, minerals, and rare earths. Much of their analysis has been focused on quantifying metals like iron, steel, and copper. It seems they are looking into replacing their large floor-standing WD-XRF with a benchtop ED-XRF or another floor-standing WD-XRF.

They are also planning to attempt quantification of trace metals that others report to be in the ppt range, so they are considering adding an ICP-MS downstream of the XRF. However, ICP-MS is completely destructive (the membrane filters go through solvent extraction).

My colleague and I discussed that if they were going to put an ICP-MS downstream anyway for trace metals analysis, perhaps it would be better to budget for a benchtop ED-XRF upstream and a more expensive ICP-MS downstream, so that the XRF acts more as a selection mechanism for which filters should be extracted for ICP-MS.

One important question we have is on our assumption that XRF is "non-destructive" to the sample. The membrane filter will certainly be intact after XRF, but will there be some amount of loss of particulate matter that will be significant for trace metals being analyzed in the ICP-MS step of this hypothetical measurement setup?

Thanks for your time and help.

The title of the question is about sample heating by ED-XRF instrument. Yet, OP is wondering its limitations and like to understand what kind of compounds analysts might lose to some sample heating if they switch to ED-XRFs (bench-top?). I hope this may be addressed by MaxW, when OP has said whay kind of samples they are analyzing. However, general knowledge on these instruments would predict both WD- and ED-XRF instruments would heat samples regardless of what kind of instrument is in use, but I'd suggest ED-XRF would heat less than that of WD-XRFs since they are low energy instruments.

Actually, to me, the bigger question is which kind is more sensitive to particular elemental analysis, regardless of sample kind (e.g., environmental or archaeological). My literature survey on the subject indicate that ED-XRF instruments is more qualitative than quantitative toward analysis, specifically towards analysis on low-$Z$ and medium-$Z$ elements (References 1 & 2).

The authors of ref.2 pointed out that the energy dispersive X-ray fluorescence analysis of low-Z-elements can generally be seen as a challenge. They have adjusted the instrument in several ways to get better data. The abstract of the article has well reflected the point:

  Photoninduced energy dispersive X-ray fluorescence analysis in total reflection geometry has been applied to low-$Z$ elements. To avoid the absorption in the $\ce{Be}$-window a windowless $\ce{Si(Li)}$-detector has been used. Icing of the detector crystal is prevented by a thin $\ce{C}$-foil. Physical reasons of the low energy background are reduced using a high energy cut-off reflector. Measurements of low-$Z$ elements from $\ce{O}$ ($Z= 8$) onwards are possible and detection limits in the $\pu{ng}$ region are obtained.

Anybody interested in these adjustments should get the article, which has explained them with some depicted diagrams.

On low-$Z$ element analysis, authors of Ref.2 claimed that:

  WD-XRF is the best XRF approach for measuring and quantifying low-$Z$ elements. ED-XRF spectrometers are subject to limitations, such as Bremsstrahlung radiation and escape peaks, which affect detection and quantification of low-$Z$ elements in particular. Although we will demonstrate in Section 2.3 that, under the right conditions, pXRF spectrometers can simulate the performance of bench-top ED-XRF instruments, there are two notable exceptions: $\ce{P}$ and $\ce{Na}$ (Note: the notation pXRF spectrometers are for portable-ED-XRFs).

The following diagram from Ref.2 well explained point, which shows exceptional correlation ($R^2=0.995$) of expected to measured concentrations of $\ce{P2O5}$ when WD-XRF is used:

References:

Alice M. W. Hunt, Robert J. Speakman, “Portable XRF analysis of archaeological sediments and ceramics,” Journal of Archaeological Science 2015, 53, 626-638 (https://doi.org/10.1016/j.jas.2014.11.031).
C. Streli, H. Aiginoer, P. Wobrauschek, “Total reflection X-ray fluorescence analysis of low-Z elements,” Spectrochimica Acta Part B: Atomic Spectroscopy 1989, 44(5), 491-497 (https://doi.org/10.1016/0584-8547(89)80055-2).


Nomenclature question about diols I'm struggling to understand why the following organic compound:

is named systemically as 2-ethylhexane-1,3-diol (I think its common name is etohexadiol), if the longest carbon chain in the molecule is actually 7 carbons long, as labelled here:
Wouldn't the compound have a hydroxymethyl group on the 3rd carbon and a hydroxy group on the 4th? Or do I just have some big misconception about naming alcohols (very possible).
Thanks for any help anyone can give!

I'll be using the brief guide to the nomenclature of Organic Chemistry as a reference, because I could not gain access to the blue book. This source suffices for this nomenclature example.

Seniority of parent compounds
The systematic name is based on the name of the senior parent compound, which is chosen by applying the following criteria in the order described below and shown in Fig. 1, until a decision is reached. For a complete set of criteria see Ref. [8].

In the given example, both the chains contain the senior most group, which is the hydroxyl (alcohol) group. The smaller chain, however, contains 2 hydroxyl groups as compared to the larger chain, which contains one hydroxyl and one alkylhydroxyl group (which is a lower priority). Since this is the first point of difference, The smaller chain would be the parent compound.
After identifying the parent chain, it is easy to name the locants such that the seniormost groups get lower locants; As you have written, the name would finally be 2-ethylhexane-1,3-diol.
This image is possibly more detailed than this question requires, but the image and the linked reference are good resources, should you need more information on naming tricky compounds.


Is KHF2 an ionic compound or a covalent compound? The statement below is an excerpt from my school textbook:-

  Because of the tendency of fluorine to form hydrogen bond, metal
  fluorides are solvated by $\ce{HF}$ giving species of the type $\ce{HF2-}$ like in $\ce{KHF2(KF.HF)}$. Also,  because of hydrogen bonding, $\ce{HF}$ dimerises to $\ce{H2F2}$ and exhibits dibasic nature and forms acid salt like $\ce{KHF2}$ on reacting with $\ce{KOH}$ whereas $\ce{KHCl2}$ does not exist because chlorine does not form hydrogen bond.

But since $\ce{K+}$ is a cation and  $\ce{HF2-}$ is an anion, it can easily form the ionic bond $\ce{K+HF2-}$. So, is KHF2  an ionic compound $\ce{K+HF2-}$ or a covalent compound $\ce{KHF2(KF.HF)}$?

Yes, $\ce{KHF2}$ is an ionic compound and a covalent compound. There is nothing exceptional about it; in fact, most compounds that we call ionic (not $\ce{NaCl}$, though) have covalent bonds in them. Consider $\ce{MgSO4}$ with ionic bonds between $\ce{Mg^2+}$ and $\ce{SO4^2-}$, and covalent $\ce{S-O}$ bonds within $\ce{SO4^2-}$. Consider $\ce{Na2CO3}$ with similar $\ce{C-O}$ bonds. These are the ways of the world.


Major product for aldol condensation between acetophenone and acetaldehyde Crossed aldol condensation is done between acetaldehyde and acetophenone by using NaOH as the base. So, here we can get 4 possible products. But which one would be the major?

I know that acetaldehyde will act as electrophile to which addition will occur. But what will act as nucleophile to give major product: acetaldehyde or acetophenone?

I don't have any solid reasoning whether acetaldehyde enolate or acetophenone enolate will give major product with acetaldehyde.

I would expect the aldehyde to be the major electrophile in this reaction. The reason is that the steric demands for a nuelophile coming close to the carbonyl carbon are lower than for the same nucelophile coming close to a ketone.

I would expect that the enolate formed from the reaction of the methyl phenyl ketone (acetopheone) to be slightly more stable than the one formed from ethanal. This is because the enolate can be delocalised into the benzene ring.

If we were to form the enolate from ethanal then if this was to attack the carbonyl carbon of acetophenone then we would get a tertary alcohol as the product. The major aldo product in this reaction will be a secondary alcohol formed by the action of the enolate from acetopheone on the ethanal.


Pressure Vs Volume plot for real and ideal gasses $$\pmb {\underline {\text {Background}}}$$

Recently I have been studying about the states of matter and came to the topic of ideal gases and real gases and the laws related to them. While studying it from my textbook I saw the plot of pressure Vs volume for real and ideal gas. The problem is that the interpretation that I am getting from the graph is quite contrary to what I have learnt till now (this might be because of some conceptual misunderstanding of mine. Sorry if the question looks foolish.)

$$\pmb {\underline {\text {Question Explained}}}$$

In the above picture  there is a point where real gas behaves the same as ideal gas (where the curves intersect). Now as we decrease the volume in the graph we find that pressure exerted by the real gas is greater than that of ideal gas for a given volume. This seems to be contrary to what I have learnt which is as the volume of a real gas decreases the pressure exerted by a real gas becomes lesser than ideal gas at the same volume (due to increased intermolecular attraction). (In reality even this isn't the case as decreasing the volume much further causes the intermolecular repulsion to overcome the attraction at some point and hence increase the pressure). But as far as I know the graph has been made with the use of the van die Waals equation (I checked it myself via plotting the graph in Demos and it looked the same as shown above). But the assumption with which the equation (and hence the graph) were derived say something different. The equation I'm talking about is $P_{ideal}=P_{real}+\frac {an^2}{V^2}$ and as far as I can comprehend it says pressure by an ideal gas is greater than pressure by a real gas at a given volume. Hence the question

Why does the equation ($P_{ideal}=P_{real}+\frac {an^2}{V^2}$) and graph(shown below) say totally different thing about pressure at a given volume?

Please do point out if I am going wrong anywhere.

Thanks in advance!

Here in thisCrash Course video the same logic as mine is given for decrease in pressure for real gas.

When the pressure tends to infinity, the volume of the ideal gas tends to zero, and the volume of the real gas does not. It tends to the proper volume of the molecules, which is greater than zero. This is exactly what you see on the upper part of the picture.

And when the pressure is extremely low, the volume increases but not as much as calculated with the ideal law, because there is some small attraction between the molecules. This is exactly what you see on your picture if you draw a horizontal line at its lower part