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January 16[edit]

Strength of a wood-screw joint[edit]

Is any data published about how much tension wood-screws in soft wood can bear without failing? The wood is far more likely to break than the metal screw - the screw would be ripped out of the wood rather than the metal snapping. I am interested in old-fashioned wood-screws as well as Phillips and similar screws. Thanks. 89.243.221.49 (talk) 00:18, 16 January 2010 (UTC)[reply]

I've had this link in my Bookmarks for a L-O-N-G time! It's the most comprehensive description of the pull-out and shear strengths of screws, nails, staples and so forth that I've been able to track down: http://www.fpl.fs.fed.us/documnts/fplgtr/fplgtr113/ch07.pdf - brace yourself though - it's a WAY complicated subject. Everything from the angle you drive the screw, the grain direction, the type of wood and (especially) the dryness of the wood matter. But if you're prepared to study - that document is the business! SteveBaker (talk) 02:26, 16 January 2010 (UTC)[reply]

Thanks. Do you know what the name of the book is please? 78.149.116.255 (talk) 12:05, 16 January 2010 (UTC)[reply]

"Wood handbook : wood as an engineering material". Madison, WI : USDA Forest Service, Forest Products Laboratory, 1999.

You can buy it on Amazon for $60 here - but since it's all available online here, you might want to save your money. Aside from the chapter on screws and such - there is all sorts of hard data about the properties of wood as an engineering material. Not exactly bedtime reading though. SteveBaker (talk) 16:43, 16 January 2010 (UTC)[reply]

I scanned the woodworking reference SteveBaker provided and in spite of its @SB note the apostrophe that is not there! copious details I think it missed this variable: how much is the screw tightened? Just until it feels right? Cuddlyable3 (talk) 22:33, 17 January 2010 (UTC)[reply]

are these EWGs?[edit]

My question wasn't answered, so I'll post again. I've posted an image this time! I want to know if these compounds can be useful building blocks. Is d acidic enough to undergo the Knoevenagel condensation? (I guess it polymerises easily?)

Also, what is the name for alkyne groups conjugated with a carbonyl system? Are these stable? Can they act as Michael acceptors and/or electron-withdrawing groups? Why is it so hard to find names for what I thought would be common industrial compounds?

item b is Propiolic acid, item a is Acrylic acid. Both have many reactions that could be used to fabricate other materials. item d with two double bonds could certainly polymerize. Graeme Bartlett (talk) 10:45, 16 January 2010 (UTC)[reply]

Well the issue is that it's an acrylate ... deprotonated acrylates don't really form enolates, do they? It would be a double anion. John Riemann Soong (talk) 11:31, 16 January 2010 (UTC)[reply]

I would have thought the answers to these questions would be in your lecture notes. If a lecturer sets homework problems, they almost always relate directly to the content of the lecture! Failing that, textbooks are good. Or have you tried both these sources and had no luck?

Ben (talk) 15:09, 16 January 2010 (UTC)[reply]

Um I'm currently not in school. School starts in five days. I'm asking because I'm trying to synthesise various products and need to know how these groups work in practice (versus on paper). My lecturer never addressed what happens to the Aldol EWGness of a deprotonated COO- group on a Michael acceptor, okay? Neither do our articles.

Also kindly do not insult me by saying these are homework problems. THEY ARE NOT. I drew these molecules as examples of structural motifs I have been considering. John Riemann Soong (talk) 15:27, 16 January 2010 (UTC)[reply]

Why is it so hard to get answers here? Can't anyone answer questions on elementary first-year chemistry? I passed my last orgo final with flying colours (242/250), in fact, I may not have another orgo class until grad school (unless I take orgo III in my last year). But of course, the more I ponder questions, the more I realise what I do not know, because naturally what you do on an exam doesn't translate to real life knowledge. So I'm trying to ask real life chemists. I ask the reference desk because my profs get tired of being pestered all the time and have stopped answering my emails. =( (They are also on vacation.) SO CAN SOMEONE PLEASE START ANSWERING MY BASIC QUESTIONS??

For example, when I look at the malonic acid syntheses (my lecturer focused on the alpha-carbon and the reaction pathway and not you know, the issue of whether a COO- group was still electron-withdrawing). I have a feeling that malonic acid has two COOH groups so simply because of pkas one will be deprotonated and one will be protonated but this also means that now malonic acid only has ONE EWG, not two? And why would the Knoevenagel condensation work with beta-keto acids at alkaline pH's if the COO- group is deprotonated? (There is only 1 carboxylic acid group, not two so naturally being the only COOH group on the molecule it will lose its proton.) But I see reactions occur all the time with acrylic acid when its carboxylic acid group should supposedly be deprotonated at pH 7 and these reactions seem to be activated by BASE. HELP SOMEBODY????

I get the idea that you can kind of help the whole issue by doing it in non-aqueous solvent but this only appears to work so far when you are specifically adding BASE to the reaction. John Riemann Soong (talk) 15:42, 16 January 2010 (UTC)[reply]

Take a deep breath and relax! It's not worth getting stressed about. I wasn't trying to insult you at all, just trying to help, as I didn't know the answer off the top of my head, no-one else is answering, and you wanted to find the answer quickly.

You're trying to synthesise these molecules in the lab. What's the normal method for finding the best way of doing a synthesis? Check the literature, e.g. Organic Syntheses, Web of Knowledge, SciFinder, the library. Get your textbooks out and check if the concepts are covered - your lectures might not have covered a certain topic but your textbook probably will.

If you snap at people who are being nice to you by trying to help, they might not help in future. Say 'thank you' once in a while!

Ben (talk) 15:55, 16 January 2010 (UTC)[reply]

I dunno, being accused of asking a homework question is like a huge insult on the RD. I wouldn't have asked it on here if Google Scholar was giving me relevant results! I have checked the literature. I have checked my notes. I certainly crushed my textbook before. I am not a grad student; I have used literature before to learn new reaction pathways, but I am not like, a literature warrior. Literature is sometimes quite tedious to check because they assume you already know the material and are an experienced chemist. "Oh look! We won't show you why this reaction works because we assume you already know the reasons why!" I AM NOT AN EXPERIENCED CHEMIST. I CAN'T FIND THE INFORMATION ANYWHERE AND I NEED HELP. =( John Riemann Soong (talk) 16:04, 16 January 2010 (UTC

OK, fine, just say "I've checked my notes and textbooks but couldn't find the answer". I know you need help but shouting about it isn't going to get you what you want any faster. Take a break, get some fresh air and come back - you'll be able to think more clearly and will feel less frustrated.

If you're having difficulty with carboxylate groups, why not esterify them or conduct your reaction in acidic conditions?

Ben (talk) 16:20, 16 January 2010 (UTC)[reply]

Side reactions will occur on my electrophile -- it's a condensed amino acid with substituted groups on it. Oh whoops, it's been condensed (self-acylated and cyclised) so in fact it is an "amino ketone" ... with a very sensitive alpha proton between an electronegative nitrogen and an EWG ketone carbonyl. If I use acid, I risk various undesirable side reactions. I think. I am looking at the Knoevenagel condensation because I think decarboxylation will favour my desired nucleophile via Le Chatelier's principle to drive the reaction forward.

I know that two mesomeric EWGs (if I use the ester) are theoretically supposed to be better than 1 carbonyl and an inductive electron-withdrawing nitrogen. But what happens is when the enol on my "amino-ketone" forms, the enol C=C bond gets connected to an entire indole ring... so uhhh.... that puts me in a dilemma. Hence why I want to use decarboxylation... but it's so difficult! John Riemann Soong (talk) 16:34, 16 January 2010 (UTC)[reply]

Just try a few of them for real. See what works - this is why we do experiments!

Ben (talk) 16:40, 16 January 2010 (UTC)[reply]

Electrical circuit loading[edit]

I'm trying to teach myself electronics and don't understand why my circuit is doing something. I have a 400ohm resistor and two LEDs in series, and a 3 volt DC motor in parallel across that. All is connected to a standard 9 volt battery. When I close the switch, the LEDs briefly flash, and the motor starts. I assume the motor is taking all the power. What do I do to spread the load across the LEDs and motor, and why is this happening? Thanks in advance! --Kickstart70TC 03:36, 16 January 2010 (UTC)[reply]

It sounds like your motor is the path of least resistance. It's no wonder, when you have a 400 ohm resistor along the route with your LEDs in it. John Riemann Soong (talk) 03:48, 16 January 2010 (UTC)[reply]

I wondered that, but putting the motor after the resistor means that either, but not both, the motors or the LEDs activate (mostly the motor). The resistor is there to protect the LEDs (max 3.2 volts each), so I can't really do without it. --04:42, 16 January 2010 (UTC)

The LEDs would briefly flash because the inductance of the motor briefly stops the current flowing through it and keeps the voltage up. Dmcq (talk) 09:42, 16 January 2010 (UTC)[reply]

@Kickstart: I'm assuming that your switch is in series with the battery, so both the LEDs and the motor are energised at the same time. Then you say that the LEDs light up for a moment, and go off as the motor starts turning. Is that right? Then Dmcq's explanation is correct. You also say that you have a 3V motor connected across a 9V battery. This means that the motor will be drawing far more current than it is meant to, so it is not surprising that the voltage drops too low to light the LEDs. You might even damage the motor. You need to arrange for the motor to receive a 3V supply, not 9V. You should be able to do this by adding another resistor in series with the motor. Put a multimeter across the motor and experiment with different resistors until you get 3V across the motor. That should do it. Beware that the resistor might get quite hot, so you may need a high-power wirewound type. --Heron (talk) 11:17, 16 January 2010 (UTC)[reply]

It's all Ohm's law and basic algebra. This is not meant as a flip response; if you want to do any sort of electronics you need to think in Ohm's law almost as a second language. Short Brigade Harvester Boris (talk) 15:47, 17 January 2010 (UTC)[reply]

LEDs don't seem to have received the memo about Ohm's law. (That was a flip response.) Inductance can be handled using complex number algebra. Cuddlyable3 (talk) 22:20, 17 January 2010 (UTC)[reply]

protection of aromatic substitution sites[edit]

Is this possible?

This is a polycyclic compound (tryptophan) -- one site is particularly reactive, and if I "disable it" via a halogen, I think I can then carry on a reaction at another less reactive site (it's intramolecular so it would be kinetically favoured) -- and the halogen would also deactivate the imidazole ring and favour substitution on the benzene ring.

The thing is, is "dehalogenation" practical? I want to do it without affecting any further substituents that I put on. (None of the substituents are aryl, though.) The substituents I'm putting on will be aryl-carbon bonds. John Riemann Soong (talk) 03:47, 16 January 2010 (UTC)[reply]

Toilet bowl - the cleanest place in school?[edit]

As the toilet bowl is always flushing itself after use, is it true that the toilet bowl surface is the cleanest place in school? Why or why not?
If there is any information on the question, can anyone please answer it or upload a reference website address up here? Thank you. Crystal1102 (talk) 05:09, 16 January 2010 (UTC)[reply]

There are undoubtedly things in the chemistry and biology labs, as well as the nurse's office that are kept close to sterile. APL (talk) 05:34, 16 January 2010 (UTC)[reply]

In my year 12 biology class we meaured bacteria counts in different parts of the school. The cleanest spot was the urinal, and the place richest in bacteria was the canteen. Nasties can stick to the toilet bowl. Graeme Bartlett (talk) 06:18, 16 January 2010 (UTC)[reply]

Makes sense. There's a chunk of solid disinfectant in the urinal, but not in the salad :) , usually. --Dr Dima (talk) 08:00, 16 January 2010 (UTC)[reply]

The way I've heard this same story, it was about people's homes, and how the toilet bowl is typically far more sanitary than the kitchen. Wooden butcher's blocks being particularly filthy. Still, there would likely be a difference in the kind of bacteria. It's still a good idea that "employees must wash hands".--Rallette (talk) 11:31, 16 January 2010 (UTC)[reply]

If your urinal has autoflushers why don't people just not touch the restroom or their privates? It's not that hard and saves time. This would probably increase sick days, though, as you still touch nonbathroom things and be reducing your only middle-of-the-day source of hand washing. Sagittarian Milky Way (talk) 14:29, 16 January 2010 (UTC)[reply]

I was waiting around in the toilet at Subway for ages the other week and no-one showed up to wash mine for me. --Kurt Shaped Box (talk) 15:04, 16 January 2010 (UTC)[reply]

In reference to the comment on butcher's blocks above, I'm assuming you mean because of the porosity of the wood and perhaps the deep cuts providing nice niches for bacteria to avoid mechanical debridement during scrubbing, etc. If, for instance, one had a handheld cutting boad small enough to place in the microwave, would zapping it take care of the microbes? Or perhaps putting it in the freezer -- would that kill the bacteria? Perhaps soaking in vinegar...any thoughts? DRosenbach ^{(Talk | Contribs)} 06:08, 17 January 2010 (UTC)[reply]

You also need to bear in mind one of the important principles of kitchen hygiene, don't use the same board for raw meat and other products. Raw meat will (almost) always be cooked before consumption thereby killing any possible harmful bacteria. Problems occur where there is transfer of bacteria from raw meat onto already cooked or raw consumed food. Not that all raw meat is necessarily contaminated, and vinegar is not a reliable disinfectant. Richard Avery (talk) 09:13, 17 January 2010 (UTC)[reply]

Boiling water is a pretty effective killer of nearly all bacteria (except for the ones adapted to live in thermal vents on the ocean floor, and they are not likely to invade your house), but don't use it on your hands! Dbfirs 20:52, 17 January 2010 (UTC)[reply]

To:Graeme Bartlett Hello can I know what are the processes carried out for this experiment? And can I have a report of it as well? I am also doing a similar experiment and we have no idea how to do it. Thank you. Crystal1102 (talk) 05:54, 19 January 2010 (UTC)[reply]

We had petri dishes with agar gel in the bottom. One was just exposed to the air in the room, and the other one we got a cotton swab, rubbed it on surfaces in the room and rubbed it on the agar in the petri dish. Then they were labeled and stored in the incubator till the next biology lesson a couple of days later. Then we got them out and counted the visible dots due to bacteria cultures. All the students in the class were formed into small groups, with each group having a different room to check out. All my biology notes and writeups were stolen towards the end of the course. But it had no effect on my mark. So I cant show my report, but it would have been two or three pages with aim, method, results and conclusion sections. Don't expect anything extremely rigorous from year 12! I expect we did not know the temperature or duration of the incubation. Graeme Bartlett (talk) 11:03, 19 January 2010 (UTC)[reply]

To: Graeme Bartlett
Thanks alot! Crystal1102 (talk) 12:28, 19 January 2010 (UTC)[reply]

Lorentz transformation[edit]

How would the Lorentz transformation equations change if the two reference frames were moving relative to each other at an angle with respect to the set of axes. Would the gamma factor for x, y, and z just use the component of the velocity in their respective axis for v? —Preceding unsigned comment added by 173.179.59.66 (talk) 06:53, 16 January 2010 (UTC)[reply]

As far as I understand your question, this would only be a rotation to the 3 space coordinates. The transformation would look far more complicated as in the standard case of movement only along one axis. Physically, this wouldn't change anything. 93.132.135.45 (talk) 11:20, 16 January 2010 (UTC)[reply]

From the article Lorentz transformation:

{\begin{bmatrix}c\,t'\\x'\\y'\\z'\end{bmatrix}}={\begin{bmatrix}\gamma &-\beta _{x}\,\gamma &-\beta _{y}\,\gamma &-\beta _{z}\,\gamma \\-\beta _{x}\,\gamma &1+(\gamma -1){\frac {\beta _{x}^{2}}{\beta ^{2}}}&(\gamma -1){\frac {\beta _{x}\beta _{y}}{\beta ^{2}}}&(\gamma -1){\frac {\beta _{x}\beta _{z}}{\beta ^{2}}}\\-\beta _{y}\,\gamma &(\gamma -1){\frac {\beta _{y}\beta _{x}}{\beta ^{2}}}&1+(\gamma -1){\frac {\beta _{y}^{2}}{\beta ^{2}}}&(\gamma -1){\frac {\beta _{y}\beta _{z}}{\beta ^{2}}}\\-\beta _{z}\,\gamma &(\gamma -1){\frac {\beta _{z}\beta _{x}}{\beta ^{2}}}&(\gamma -1){\frac {\beta _{z}\beta _{y}}{\beta ^{2}}}&1+(\gamma -1){\frac {\beta _{z}^{2}}{\beta ^{2}}}\\\end{bmatrix}}{\begin{bmatrix}c\,t\\x\\y\\z\end{bmatrix}}\ ,

Dauto (talk) 04:44, 17 January 2010 (UTC)[reply]

computer symulation for the theory of general relativity of Einstein[edit]

I am a Phyisicist student and am looking for a computer symulation for the theory of general relativity of Einstein, as without such instrument it is very dificult to perform the required calculations. I cannot use the Newton's mechanics as I wish to calculate calculations for bodies near an event horizon of black holes. Do you know where I can find something like this? Thanks in advance 89.139.82.116 (talk) 09:56, 16 January 2010 (UTC)[reply]

What exactly do you want to calculate? Here's a research paper on The general relativistic N-body problem; it is an analytic treatment; you may be interested in our article on numerical relativity, which links to many different research papers. There are probably not a lot of open-source or commercial relativistic numerical simulations, because this is sort of a niche research community, but you can try contacting the corresponding author of one of the papers to see if they have any pointers or code to share. This NASA press release on black hole simulation mentions this PRL paper, Gravitational-Wave Extraction from an Inspiraling Configuration of Merging Black Holes, where black holes are evolved through a computational grid. Nimur (talk) 15:35, 16 January 2010 (UTC)[reply]

Crop yield[edit]

The article describes what it is but are there any kinds of tables to check the yield for different crops/climates/soil? I did a search on the net and it is all widely incoherent. It's measured in different units on each article and I strongly suspect there are people out there who mistook MT for Mt (metric tons for megatons). 93.132.135.45 (talk) 11:15, 16 January 2010 (UTC)[reply]

There are so many variables, in climate and soil conditions alone, that I doubt whether, even if such tables existed, they would be of any practical use. There are no doubt localized figures available for specific crops, so you would do better to use much more specific search terms.--Shantavira|^{feed me} 11:55, 16 January 2010 (UTC)[reply]

I'm looking for figures for use in renewable energies from neutral sources. 93.132.135.45 (talk) 12:04, 16 January 2010 (UTC)[reply]

See the so-called "Billion Ton Study" and references therein. Short Brigade Harvester Boris (talk) 15:49, 17 January 2010 (UTC)[reply]

Does milk spoil or taste old faster the less of it is left in the container?[edit]

If so, does the rate of spoilage become only a little faster or a lot? I have a feeling it at least accelerates somewhat with inverse volume but never empirically found out as there are various reasons we wouldn't get to be affected by this much in normal use. Sagittarian Milky Way (talk) 14:10, 16 January 2010 (UTC)[reply]

It probably depends on conditions. Large volumes of milk have more thermal inertia so they stay cold longer. Realistically, this effect only lasts a few hours (if we're talking about 1-gallon sized quantities), but it might be the difference between rancid and not-rancid if you leave a full gallon vs. a mostly empty gallon on the counter for two hours. Homogenized milk should be, well, homogeneous - so the bacteria and their food source (lactose) should be evenly distributed throughout the milk. Therefore, once the temperature is warm enough to activate them, adding more milk shouldn't change the spoilage rate. But, there may be a surface-area to volume ratio issue. We have an article on soured milk with links to a bunch of links and patents on techniques for intentionally souring milk, e.g., Process for Preparing Gelled Sour Milk. Nimur (talk) 15:11, 16 January 2010 (UTC)[reply]

Perhaps a thought experiment is in order. Consider a container of milk containing one lonely bacterium. Let's assume that refrigeration confines it to one cycle of division per day. On day two, there will be two bacteria; on day three, four bacteria. On day ten, there will be a thousand; on day twenty, a million; on day thirty, one billion (roughly). If each bacterium turns out the same amount of lactic acid each day, the amount of acid generated on day ten will be roughly equal to the amount of acid released in total on days one through nine. Each day's new growth will match the spoilage caused on all the preceding days. This is why the apparent rate of spoilage will tend to increase with time; it doesn't matter how full or empty the container is. (That said, as the container empties, there is more and more room air being exchanged into the jug or carton each time you open it. That room air may contain additional microbial contaminants which increase the load in the container and speed spoilage. Again, the faster spoilage isn't directly the result of a less-full container; there's just been greater exposure of the emptier container to contaminated air, and the emptier container has been opened and closed more often than the full one.) TenOfAllTrades(talk) 16:10, 16 January 2010 (UTC)[reply]

In the absence of bacteria things alter their taste by oxidation, which happens by surface/volume. Take a look at an apple pie. If you have the purse and the taste for it, examine an opened bottle of wine over the course of some days or weeks. 93.132.135.45 (talk) 18:16, 16 January 2010 (UTC)[reply]

How does Parathyroid hormone work on the kidneys?[edit]

Ok so I've done some research and now I'm stuck! As far as I know parathyroid hormone (PTH) has 3 main effects on the kidneys: Increases reabsorption of calcium ions Reduces reabsorption of phosphate ions Causes conversion of calcifediol to calcitriol

What I need to know about is how these effects are brought about. I know there are PTH 1 receptors to be found in the kidneys, does the binding of PTH to them trigger a cAMP dependent pathway activating a protein kinase which would produce a substance of some sort? The wiki article on PTH says PTH inhibits the cAMP dependent pathway but an article I found in the Journal of the American Society of Nephrology 20(8):1693-1704, August 2009 said quite the opposite. Also, I know that PTH increases the uptake of calcium ions by activating Ca2+-ATPase and a Na+-Ca2+ antiporter but can't find any information on how this occurs...

How does PTH reduce the reabsorption of phosphate ions? Also through the activation of protein kinase A via a cAMP dependent pathway producing some sort of substance? I've also read that PTH inhibits the sodium phosphate co-transporters in the Proximal convoluted tubule - reducing phosphate reabsorption...

As you can probably tell I'm very confused now. Any help would be massively appreiciated! RichYPE (talk) 14:37, 16 January 2010 (UTC)[reply]

It is certainly not unheard of to be confused about the complexity of calcium metabolism. The "business end" of any hormone is it's receptor, in this case two parathyroid hormone receptors. I wonder if you might have mis-read the part in parathyroid hormone about the cAMP dependent pathway -- that section is talking about regulation of PTH secretion by the calcium-sensing receptor (which inhibits cAMP production). In terms of getting an immediate "boost" to serum calcium, the most important action of PTH is to release calcium from the bone (which is where the vast majority of your body's calcium is stored). With regard to the mechanism of renal ion movements, you could start with the nephron article and the distal convoluted tubule, where regulated calcium reabsorption happens. Neither article really gives much detail about the movements of ions, but just remember that in most physiological systems, whenever there are ion channels in the plasma membrane of the cell, ions will flow along a concentration gradient. In the kidney, there are members of the Transient Receptor Potential Vanilloid (TRP) family (TRPV5 and TRPV6) that are responsive to vitamin-D and mediate calcium entry (see also this reference to get you started). I think the effect of PTH on renal reabsorption may be mostly secondary (via its action on vit-D). With regard to phosphate excretion, this article says that "FGF23 synergizes with PTH to increase renal phosphate excretion by reducing expression of the renal sodium-phosphate cotransporters NaPi-IIa and NaPi-IIc in the proximal tubules" (see phosphate homeostasis for more details). --- Medical geneticist (talk) 15:50, 16 January 2010 (UTC)[reply]

Thanks to medical genticist for the reply. I am aware that the conversion of calcidiol to calcitriol (1,25-dihydroxycholecalciferol)is a large part in how PTH affects the kidneys. At the moment, I'm focussing on how the kidneys act to increase calcium reabsorption and reduce its excretion without the influence of calcitriol. My essay is shaping up ok now, can someone read this and tell me if its correct: 'PTH acts by binding to Calcium sensing receptors (CaSR)in the kidneys activating renal adenyl cyclase to increase intracellular concentrations of cyclic AMP which activates protein kinase A and C.' Although now I've got a bit stuck again, what would activating protein kinase A and C do?

I also think PTH binds to GPCRs causing an increase in intracellular cAMP in the kidney cells which stimulates the production of phospholipase C which produces a subtance called IP3 that opens calcium channels allowing more calcium to be reabsorbed. Again is this correct?

Finally, would I be correct in saying that PTH increases phosphate excretion through the inhibition of Brush border membrane sodium dependent phosphate transporters in the proximal convoluted tubule? Thanks, RichYPE (talk) 20:06, 16 January 2010 (UTC)[reply]

the asymmetry of acetic anhydride[edit]

I've asked this question before, but I really must return to it since a) I am starting a subproject where I post electron density images to articles b) I want a sort of caption to go with each image. So again -- why is acetic anydride asymmetric? I really don't think it's my calculator being stuck in sort of local well. It really seems that there is some sort of repulsion that makes the molecule take that conformation. -- further confirmation of the idea that the ethoxy oxygen in the molecule does not have a pi system connect across it at all! i.e. Acetic anhydride is not planar. One thing that I can think of is that there is intrinsic electronic repulsion, in a way not unlike antiaromaticity and this makes the carbonyl groups move to different planes (to avoid withdrawing from each other and to withdraw more from the methyl groups, maybe?).

Really have to explain why:

the two C=O bonds aren't equal
the ethoxy C-O bonds aren't equal either
Neither are the methyl C-H bonds!
To define a scale, the most negative pixel of the molecule has a density value of -41, while the most positive pixel has value +35.43. Te average value of the one carbonyl carbon (rough approximate) appears to be +33 (on one side), while the other carbonyl carbon (on the most positive side) it has a density value of +30. This seems to be quite a significant difference in the electropositivities of acetic anhydride's carbonyl carbons for what would be a presumably symmetric molecule.
There is asymmetry even within the same carbonyl carbon -- for one particular carbonyl carbon, one "side" of it (where a nucleophile might presumably attack) appears to be about as neutral as an amide carbonyl carbon -- values range from -3 to +8, while the other side is drastically more positive -- +27 to +30. Apparently this is due to the nonplanar way the oxygen atoms are sticking out of the molecule.

I used 6-31G* for my calculations -- I think that's as rigourous as my program can go. Please answer promptly because I am about to edit the acetic anhydride article! John Riemann Soong (talk) 15:01, 16 January 2010 (UTC)[reply]

As with most things, experimental evidence is important - otherwise you're adding OR to Wikipedia. The structure of acetic anhydride has been probed experimentally, see for example J. Phys. Chem. A (2000) 104 (7), 1576–1587.

I had a look at the crystal structures of some anhydrides and they almost all had the same conformation (Jmol model here, based on Angew. Chem. Int. Ed. (2008) 47, 8030-8032). The exceptions were cyclic anhydrides whose geometry is constrained to nearly planar (e.g. succinic anhydride, maleic anhydride, phthalic anhydride).

Ben (talk) 15:34, 16 January 2010 (UTC)[reply]

So acid anhydrides (that are free to rotate) are CONFIRMED to be asymmetric, correct? My question is, what is the EXPLANATION for this? (Also I don't think adding in silico calculations are OR -- these get done all the time by undergraduates everywhere and I'm sure I must be the millionth person to calculate the mysterious density of AcOAc just that it's never been posted on Wikipedia.) John Riemann Soong (talk) 15:45, 16 January 2010 (UTC)[reply]

Also, why didn't anyone tell me this amazing fact during lecture??? Why is it not mentioned in the article acid anhydride? What is the explanation for it? Why can't I find it? What is wrong with everybody? John Riemann Soong (talk) 15:47, 16 January 2010 (UTC)[reply]

I'm not a grad student. I can't understand that paper, because I can't seem to sort through the alphas and phis and wavefunctions to find out if the methyl C-H bonds are equal and if not, WHY. John Riemann Soong (talk) 15:57, 16 January 2010 (UTC)[reply]

Neither am I, I just look at the diagrams and read the abstract and conclusion and try and understand it as best I can. It basically says the molecule does lose some resonance stabilisation by adopting non-planar conformations, but this is approximately compensated for by loss of steric hindrance and dipole-dipole repulsion (I guess that means δ− carbonyl oxygens close together in space in one of the planar conformers). Overall it says that "the delicate balance of these interactions allows in AA for a wide variation in skeletal torsion and valence angles to occur without much variation in internal energy. The different geometries along the rotation paths can be rationalized in terms of mesomeric and steric effects."

Ben (talk) 16:10, 16 January 2010 (UTC)[reply]

On average, do the C-O bonds (on either side) average out? That is, statistically, are the bond lengths symmetric? I get the idea that at any one time the bond lengths aren't equal. (Not that it makes that much sense anyway since the bonds are constantly vibrating.) Like ... why doesn't the NMR of acetic anhydride show two separate singlets (or worse, a complex mix of chemical shifts?)

Theoretically, does this sort of imply there wasn't much resonance stabilisation anyway? I mean usually resonance stabilisation >> conformation stabilisation (except in extreme cases). I actually thought there was internal electronic repulsion that forced the pi systems to be orthogonal to each other, but I guess not. John Riemann Soong (talk) 16:18, 16 January 2010 (UTC)[reply]

One consequence of a nonplanar conformation is that acetic anhydride appears to be more electrophilic than it would otherwise be in its planar conformation, because basically the electropositiveness of one side of the carbonyl carbon gets transferred to the other, so the side that doesn't have any oxygen dipoles near it is REALLY blue. But this is generally not included in class as a reason for why anhydrides are more reactive than esters! Instead we talk about dipole moments using PLANAR Lewis structures. WTF? Time to edit the textbooks anyone? John Riemann Soong (talk) 16:23, 16 January 2010 (UTC)[reply]

The Jmol model (from a crystal structure) has two different C=O lengths, 0.117 nm and 0.125 nm.

The ¹H NMR spectrum of AA (from SDBS) shows a single peak at 2.219 ppm. The ¹³C spectrum shows one peak at 166.63 pm (carbonyl carbons) and one at 22.07 ppm (methyl carbons). I assume these spectra were recorded at room temperature, so we're probably seeing averaging as the molecule rotates through many conformations, faster than the NMR timescale.

If you cooled the sample down, you'd probably see more peaks, as the rotation stops and an asymmetric conformer predominates.

What do you mean by internal electronic repulsion? The δ− carbonyl oxygens repelling each other electrostatically? Or destabilising orbital overlap?

As far as textbooks are concerned, there's a balance to be had between being pedantically correct and being intelligible. Got to walk before you can run.

Ben (talk) 16:36, 16 January 2010 (UTC)[reply]

I could never make sense of my lecturer's "dipole explanation" with planar Lewis structures. What dipole advantage? The molecule is withdrawing in both directions! It makes so much sense now to find out that the asymmetry is one of the big factors.

I don't know what I was exactly thinking with intrinsic repulsion -- I was seeing the pi electrons on each of the carbonyl systems repel each other somehow (hence why they are not on the same plane) -- maybe because they are both EWGs withdrawing from each other, and this is a mutually repulsive interaction. Like the same reason why a Cl-Cl bond is weak. I was thinking that putting them on a different plane would break this interacton and allow the carbonyls to withdraw more from the methyl groups instead. An interesting case I think, is when acetic anhydride gets protonated.

Is protonated acetic anhydride still aplanar? (Now you can hydrogen-bond!) If not ... that certainly hints at "intrinsic" repulsion. Although I get the feeling that the asymmetry encourages protonation of the C=O oxygen. John Riemann Soong (talk) 16:46, 16 January 2010 (UTC)[reply]

Anyway thanks. I guess I'm going to make an article edit. I'll use the paper you found as a source. John Riemann Soong (talk) 16:54, 16 January 2010 (UTC)[reply]

Not sure about protonated AA - try calculating the energies of the planar and nonplanar conformers of it in Spartan and compare. As for your Cl-Cl lone pair repulsion thing, the same argument could be applied to 1,3-butadiene but that's planar. Think about both those cases from an orbital point of view: Cl₂ doesn't have the vacant π MOs that butadiene does.

I'm going to eat now but I will draw some diagrams in a bit. Good luck.

Ben (talk) 17:02, 16 January 2010 (UTC)[reply]

Alkenes aren't really good EWGs though. I mean, the electron density is likely to be be more in the middle of the molecule (intrinsically a bonding interaction) than the ends (which is how it is in AA). John Riemann Soong (talk) 17:07, 16 January 2010 (UTC)[reply]

Howdy, I'm a computational chemist, perhaps I can be of some help. I pulled the full paper and gave it a quick read. The paper reports energies for various conformations of acetic anhydride. The lowest energy structure is definitely the one you would expect, planar acetic anhyrdide with some low energy conformations, some of which are non-planar but are roughly 1 kcal/mol above the planar ones. The issue with your calculation, and the reason you're getting an asymmetric electron density, is because Hartree-Fock (the calculation method you cite in the image information) isn't equipped to handle molecular orbitals with multiple resonance forms, which is precisely the reason that acetic anhydride is planar. The paper reports finding the planar minimum using Hartree-Fock however, and they may have manually forced the geometry to achieve that. I would strongly suggest upgrading to a perturbation theory based method such as MP2. 24.177.124.41 (talk) 17:21, 16 January 2010 (UTC)[reply]

Ah good, someone who actually knows what they're talking about! Thanks for stepping in.

I just had a look in the Spartan database. MP2/6-31G* gives the [sp,sp] planar geometry as the global minimum but MP2/6-311+G** says nonplanar, like this.

Maybe I was reading out of context, but I searched the paper for "minimum" and found the phrase "none of the planar forms (Figure 1) is a minimum energy form".

Ben (talk) 17:38, 16 January 2010 (UTC)[reply]

Oops, that'd be my bad. I misread the point groups. As I look at this paper more, I'm not sure I trust their computational method. B3LYP is a density functional theory and isn't particularly good at measuring weak interactions like you might find between the two carbonyls (depending on the relative electron populations which they don't report). All in all, I'd recommend finding an experimental work instead of a theoretical one for this answer and base your image on that. 24.177.124.41 (talk) 20:00, 16 January 2010 (UTC)[reply]

Hey thanks for stepping in! Do you mean interactions through space (dipole-dipole), or bonding interactions? Both would seem to favour aplanarity. But even if the planar form is the most stable form (crystal structure apparently shows different C=O bond lengths anyway? is it experimental?), the low barrier to rotation suggests very weak resonance stabilisation, which agrees with classical theory of course.John Riemann Soong (talk) 06:02, 17 January 2010 (UTC)[reply]

The crystal structure is experimental (almost all are, although see crystal structure prediction).

Ben (talk) 15:05, 17 January 2010 (UTC)[reply]