Where can I find Equilibrium Constants? When trying to develop mathematical models to show the concentration in a reaction, where can one find a database or good source of known equilibrium constants if it is not necessarily very convenient to find them empirically?

First, a question: what kinds of reactions are we talking about? Is it a more specialized (in this context) field like biochemistry, or more general chemical reactions? In practice, finding data for the latter is probably simpler than finding data for the former without doing any experimental work.

In general, I would use the information about Gibbs free energy of formation that can be found here (among many other places) for the reactants and products, and then calculate the equilibrium constant from

$$ K = \exp \left(\frac{-\Delta G^\circ_{\mathrm{rx}}}{RT}\right) $$

where $\Delta G_{\mathrm{rx}}$ is the change in Gibbs free energy for the reaction. Beware that if your reaction conditions - e.g. temperature or pressure - is different from standard conditions, this should be taken into account.

Hope this helps a bit!


Are allotropes of sulfur two different phases? Monoclinic sulfur and rhombic sulfur are two allotropes of sulfur. In phase equilibria, part of the system that is having distinct boundary, mechanically separable and different chemical or physical properties that other part of the system is considered as a phase. Now I want to ask, do chemical properties of Allotropes differ and why so that we can consider them different phases?

Solid allotropes of pretty much anything are different phases. They have different crystal structures, and you can't have a single phase with two different crystal structures inside it.

As for the other criteria, some of them are sufficient but not necessary, and some are difficult to check. The distinct boundary must be there, but that's a knowledge that comes from them being two phases and not vice versa. Physical properties are not all that different, though you can tell one allotrope from another by density. (We used a solution of $\ce{BaI2}$ to this end.)

Chemical properties are even more similar since the molecules are the same. I doubt that a purely chemical test can be devised. But then again, do we really need it?


Resources for learning Chemistry Based on various other Stack exchange site (Mandarin Chinese, Russian and German), we adapt this project here for chemistry, since it's a great idea to have all kinds of resources in one place.

This is a specifically created Community Wiki which gathers resources for learning Chemistry. The lists are a community maintained project, hence everybody with more than 100 reputation points can contribute edits to the appropriate sections. If you feel something is missing, just fill it in. To avoid multiple answers for similar branches, we decided on a general outline, which is locked, i.e. no new answers can be added.
 If you have concerns or questions, you can discuss this list on its parent meta post. If they are of a more general concern, you can also post a new question on chemistry.meta.se. This procedure is used, so the comment section here does not become too overcrowded. If you do choose the second option, please leave the link to it in the comments to this post.

Organization

Answers have a type of resource each.
If possible, state whether the material is directed towards a beginner, intermediate or an advanced audience.
Do not include links that lead to illegal content or sites that host such content. If you see any, please flag for moderator attention and choose "other" so you can point us to the content. We'll delete it as soon as we see the flag. (You can of course also delete it yourself. If you do, please flag it anyway, so that we are aware of it. In this case it is crucial you fill in the edit summary with something like: Removed link to illegal content.)
Both free and commercial resources are allowed, but make sure to include a note if they are the latter. Remember the rules about self-promotion. Include also if registration is required.
Include links to the sites only, don't post images, they would take too much space.
Add the resources in alphabetical order so they're easier to find.
For the resources, a short summary is very much appreciated.

If you have questions about these guidelines, please head to meta to make yourself heard. The above points are not set in stone and might change in the future.

Alphabetical Index

(Text)Books: All books that teach you chemistry with theory and exercises. The subcategories are:

General Chemistry
Inorganic, Organic and Physical Chemistry
Analytical Chemistry, Biochemistry and Chemical Biology, Chemical Engineering, Computational and Quantum Chemistry, Theoretical Chemistry
You can add any subcategory to this post if it is missing.

Online courses and Websites: Free or paid services online that teach you chemistry through lessons as well as sites that give help for learning chemistry. They give material, tips, hints, and various help for self-learners or regular students.
Software: This can be any software ranging from plugins for the browser over mobile apps up to standalone applications for the computer. Pure 2D or 3D visualization programs as well as quantum chemistry programs might not fit in this category, as they are not primarily focused on teaching chemistry.
Video Resources & TV: Video resources which help learning chemistry.

Currently those are all categories. If you think, that a new one should be added, please submit an answer in the corresponding meta thread. (A comment is probably not sufficient, as it does "bump" the question on the active tab.

#Books about Inorganic Chemistry

General texts

Housecroft, C. E.; Sharpe, A. G. Inorganic Chemistry, 4th ed.; Prentice Hall: Upper Saddle River, NJ, 2012.

Weller, M.; Overton, T.; Rourke, J.; Armstrong, F. Inorganic Chemistry, 6th ed; Oxford UP: Oxford, U.K., 2014.
This is the latest version of the textbook commonly known as Shriver & Atkins.

Miessler, G. L.; Fischer, P. J.; Tarr, D. A. Inorganic Chemistry, 5th ed.; Prentice Hall: Upper Saddle River, NJ, 2014.
A relatively shorter text, but explains chemical bonding extremely well.

Ghosh, A.; Berg, S. Arrow Pushing in Inorganic Chemistry: A Logical Approach to the Chemistry
of the Main-Group Elements, 1st ed.; Wiley: Hoboken, NJ, 2014.
A mechanistic approach to reactions in inorganic chemistry with emphasis on the mechanisms.

Lee, J. D. Concise Inorganic Chemistry, 5th ed; Wiley: Hoboken, NJ, 1999.
A classic introductory inorganic chemistry textbook.

Bauer, G Handbook of Preparative Inorganic Chemistry, 2nd ed; Academic Press: New York, 1963
A two volume compendium about small-scale syntheses of elements, maingroup and transition element inorganic compounds.

#Books about Organic Chemistry

General texts

Carey, F. A.; Sundberg, R. J. Advanced Organic Chemistry, Part A: Structure and Mechanisms, 5th ed.; Springer: New York, 2007.
Carey, F. A.; Sundberg, R. J. Advanced Organic Chemistry, Part B: Reactions and Synthesis, 5th ed.; Springer: New York, 2007.
Carey & Sundberg is a classic two-volume text with an extremely in-depth discussion. Mechanisms are elucidated in great detail (primarily via MO theory, with results from both ab initio and semi-empirical methods), and the exposition of pericyclic reactions is notably excellent. It also contains a great deal of illuminating content on conformational analysis. Definitely not appropriate for a first textbook, but essential reading for advanced undergraduates and above.

Carey, F. A.; Giuliano, R. M. Organic Chemistry, 9th ed.; McGraw-Hill: New York, NY, 2014.

Clayden, J.; Greeves, N.; Warren, S. Organic Chemistry, 2nd ed.; Oxford UP: Oxford, U.K., 2012.
Excellent first textbook for organic with lucid explanations. It's worth noting that the first edition, while slightly more verbose, contains more information than the second, where some chapters and sections were cut.

Smith, M. B. March's Advanced Organic Chemistry: Reactions, Mechanisms, and Structure, 7th ed.; Wiley: Hoboken, NJ, 2013.
A comprehensive book that covers nearly every reaction under the sun, with appropriate references to primary literature. Arguably best used as a reference and not as study material, but it is also surprisingly readable.

Solomons, T. W. Graham; Fryhle, C. B.; Organic Chemistry, 10th ed.; Wiley: Hoboken, NJ, 2011.
Introductory textbook for organic chemistry.

Wade, L. G. Organic Chemistry, 8th ed.; Pearson Education: Glenview, IL, 2013.
A systematic introductory textbook for organic chemistry. Follows the traditional functional group approach.

Organic synthesis

Warren, S.; Wyatt, P. Organic Synthesis: The Disconnection Approach, 2nd ed.; Wiley: Chichester, U.K., 2008.
A step-by-step introduction to organic retrosynthetic analysis and the construction of different relations between functional groups. Do also get the accompanying workbook (the solutions are discussed immediately after problems).

Kürti, L.; Czakó, B. Strategic Applications of Named Reactions in Organic Synthesis; Elsevier: Amsterdam, 2005.
An incredible compilation of 250 named reactions, with discussion of mechanisms and examples of application to total synthesis. Nearly 10,000 references to primary literature.

Mundy, B. P.; Ellerd, M. G.; Favaloro, F. G., Jr. Name Reactions and Reagents in Organic Synthesis, 2nd ed.; Wiley: Hoboken, NJ, 2005.
A reference book for numerous reaction mechanisms and common reagents.

Nicolaou, K. C.; Sorensen, E. J. Classics in Total Synthesis; Wiley: Weinheim, Germany, 1996.
Thorough discussion of the retrosynthesis and forward synthesis of 36 molecules, by one of the most well-known synthetic chemists of recent years and (at the time) his Ph.D. student. Examples are taken from almost the entire history of organic synthesis: from Woodward's 1954 synthesis of strychnine to the author's own 1995 synthesis of brevetoxin B. Also check out the sequels, Classics in Total Synthesis II and Classics in Total Synthesis III.

Wuts, P. G. M. Greene's Protective Groups in Organic Synthesis, 5th ed.; Wiley: Hoboken, NJ, 2014.
Comprehensive listing of protecting groups, protection and deprotection conditions with references to primary literature, and handy reactivity charts which assess the stability of protecting groups towards various reagents and conditions.

Spectroscopy

Pavia, D. L.; Lampman, G. M.; Kriz, G. S.; Vyvyan, J. A. Introduction to Spectroscopy, 5th ed.; Cengage Learning: Stamford, CT, 2015.
Excellent introductory text with good coverage of all the typical structure determination techniques: elemental analysis, NMR, IR, MS, and UV-Vis.

Silverstein, R. M.; Webster, F. X.; Kiemle, D. J.; Bryce, D.L. Spectrometric Identification of Organic Compounds, 8th ed.; Wiley: Hoboken, NJ, 2014.
Perhaps a slightly more in-depth discussion than Pavia, but without sacrificing any clarity. Includes a large number of tables and charts of spectroscopic data, making it also very valuable as a reference.

Claridge, T. D. W. High-Resolution NMR Techniques in Organic Chemistry, 3rd ed.; Elsevier: Amsterdam, 2016.
A much more involved (graduate level) discussion of how NMR experiments are designed and how we extract information from them. Along with the more physchem-oriented NMR books, this is recommended for those who want to know what is actually going on in their NMR machine.

#Books about Physical Chemistry

General texts

Atkins, P.; de Paula, J. Physical Chemistry, 10th ed; Oxford UP: Oxford, U.K., 2014.
Classic physical chemistry textbook, but can sometimes be difficult to follow, especially for first-time students.

Atkins, P.; de Paula, J. Elements of Physical Chemistry, 7th ed.; Oxford UP: Oxford, U.K., 2016.
A less detailed book than Physical Chemistry. Can be used for introductory physical chemistry courses.

Engel, T.; Reid, P. Physical Chemistry, 3rd ed.; Prentice Hall: Upper Saddle River, NJ, 2012.

Levine, I. N. Physical Chemistry, 6th ed.; McGraw-Hill: New York, 2009.

Statistical mechanics

Chandler, D. Introduction to Modern Statistical Mechanics; Oxford UP:Oxford, U.K., 1987
Useful for undergraduate and early graduate study of both classical and quantum statistical mechanics, as well as for reference.

McQuarrie, D. A. Statistical Mechanics; University Science Books: Mill Valley, CA, 2000.

Nuclear magnetic resonance

Keeler, J. Understanding NMR Spectroscopy, 2nd ed.; Wiley: Chichester, U.K., 2010.
An extremely accessible and readable introduction to the theory behind NMR experiments. Describes nearly every important aspect of NMR, from basic quantum mechanics to the vector model, product operators, and a range of NMR experiments. It seems to be a conscious choice on the part of the author to not delve too deep into the quantum mechanics, in the interests of clarity and understanding.

Levitt, M.H. Spin Dynamics, 2nd ed.; Wiley: Chichester, U.K., 2008.
An extremely in-depth treatment of the quantum mechanics of NMR, which goes well beyond the level of Keeler's book.

Gunther, H. NMR Spectroscopy, 3rd ed.; Wiley: Weinheim, Germany, 2013.

Computational and Quantum Chemistry

Levine, I. N. Quantum Chemistry, 7th ed.; Pearson: Upper Saddle River, NJ, 2012.
This book acts as a first introduction to quantum chemistry and computational methods for students who do not have prior experience with the field.

Atkins, P. W.; Friedman, R. S. Molecular Quantum Mechanics, 5th ed.; Oxford UP: Oxford, U.K., 2010.
Realizes a more extensive coverage than Levine.

Cohen-Tannoudji, C.; Diu, B.; Laloe, F. Quantum Mechanics; Wiley: New York, 1977.
This book is a very good book on quantum mechanics that is often used by physics majors to learn the subject. Reading it is necessary if a better understanding than that provided by previous books is desired.

McWeeny, R. Symmetry; Dover, 2002.

Parr, R. G.; Yang, W. Density-Functional Theory of Atoms and Molecules; Oxford UP: Oxford, U.K., 1996.
The first book on DFT and by leaders in the field, still unbeatable.

Jensen, F. Introduction to Computational Chemistry; Wiley, 2007
This books treats the computational methods at a beginner level.

Szabo, A.; Ostlund, N. S. Modern Quantum Chemistry: Introduction to Advanced Electronic Structure Theory; Dover, 1996

Helgaker, T.; Jorgensen, P; Olsen, J. Molecular Electronic-Structure Theory.
This is the reference book on wavefunction methods.

Cancès, E; Defranceschi, M; Kutzelnigg, W.; Le Bris, C.; Maday, Y. Computational Quantum Chemistry: A primer.
A sophisticated mathematical analysis of wavefunction methods and density functional theory.

Solid State Chemistry and Crystallography

West, A. R. Solid State Chemistry and Its Applications, 2nd ed.; John Wiley & Sons, Inc: Chichester, West Sussex, 2014.
Simon, S. H. The Oxford Solid State Basics, 1st ed.; Oxford University Press: Oxford, 2013.
Smart, L.; Moore, E. Solid State Chemistry: An Introduction, 4th ed.; CRC Press: Boca Raton, FL, 2012.
Tilley, R. J. D. Understanding Solids: The Science of Materials; J. Wiley: Chichester, West Sussex, England; Hoboken, NJ, USA, 2004.

Hoffmann, R. How Chemistry and Physics Meet in the Solid State; DOI: 10.1002/anie.198708461.
Learning solid state theory is difficult for a chemist in part because it is primarily ridden by a different community (physicists) that have a different language and way to understand bonding. This review helps in bridging the gap.

Ashcroft, N. S.; Mermin, N. Solid State Physics; Cengage Learning, 1976.
This textbook is primary reference to learn solid state physics.

Evarestov, R. A. Quantum Chemistry of Solids: LCAO Treatment of Crystals and Nanostructures, Springer, 2012.

Fundamentals of Crystallography, 3th ed., Oxford University Press, 2011

Müller, Ulrich Symmetry Relationships between Crystal Structures.
Applications of Crystallographic Group Theory in Crystal Chemistry, 2nd ed., Oxford University Press, 2017
A rigorous yet accessible explanation of the foundations of crystallography and symmetry relationships with a large collection of exercises to practice. Good for self-study.

Müller, U. Anorganische Strukturchemie, 6., aktualisierte Aufl., unveränd. Nachdr.; Studium; Vieweg + Teubner: Wiesbaden, 2009. (in German)

Dronskowski, R. Computational Chemistry of Solid State Materials: A Guide for Materials Scientists, Chemists, Physicists and Others; Wiley-VCH: Weinheim, 2005.

Molecular Orbital Theory

Fleming, I. Molecular Orbitals and Organic Chemical Reactions: Student Edition; Wiley: Chichester, 2009.
A readable introduction to molecular orbital theory with an emphasis on its application to simple organic chemical reactions (SN1/SN2 etc). Suitable for undergraduates wanting a more detailed understanding of organic reactivity. Also widely considered to be the book for the molecular orbital treatment of pericyclic reactions.

Fleming, I. Molecular Orbitals and Organic Chemical Reactions: Reference edition; Wiley: Chichester, 2009.
An expanded version of Molecular Orbitals and Organic Chemical Reactions: Student Edition with a full list of references and additional coverage of some material. More suited to graduate students / instructors who might want to read deeper into a topic.

Albright, T. A.; Burdett, J. K.; Whangbo, M-H. Orbital Interactions in Chemistry, 2nd ed.; Wiley: Hoboken, NJ, 2013.
Arguably the best book on the subject, both in terms of depth (some mathematical treatment of the topics) and coverage (organic and inorganic molecules are discussed). Given the size (over 800 pages), Albright is more of a reference text than Fleming which can easily be read cover-to-cover, and the level of discussion is more geared towards graduate students / undergraduates in later years who already have some understanding of molecular orbitals.

Jean, Y. Molecular Orbitals of Transition Metal Complexes; Oxford, 2005.
An undergraduate level textbook covering the fundamentals of molecular orbital theory applied to transition metal complexes. The derivations of the MO diagrams for various coordination geometries is provided along with a chapter on 'applications' of MO theory such as how complexes react.

Electrochemistry

Elgrishi, N. et al. A Practical Beginner’s Guide to Cyclic Voltammetry; J. Chem. Educ. 2018, 95, 197–206. (open access)
A short introduction teaching the very basics about CV's theory, supplemented with tips and warnings for pitfalls recording and interpreting the experimental data.  The supplementary information contains five training modules to walk through the more typical minutiae one encounters.


Hypochlorous acid mechanism When hypochlorous acid is dissolved in water, is it the chlorine atom or the oxygen that serves as the anti microbial agent?

The chemical disinfection (inactivation of bacteria) in fresh water can occur through a number of mechanisms, including oxidation of cell walls, inactivation of key enzymes and disruption of nucleic acids, thereby rendering them non-functional. The precise mechanism of inactivation depends upon the nature of micro-organism (bacteria, spores, viruses).

When chlorine was first used as a disinfectant in the USA in 1908 (slightly earlier in Europe), its germicidal power was commonly attributed to the liberation of 'nascent oxygen' from hypochlorous acid $(HOCl)$.

$HOCl\rightarrow HCl+\frac{1}{2}O_{2}$

Subsequent investigations (Chang, 1944, Green and Stumpf, 1946) debunked this theory and showed that chlorine reacts irreversibly with the enzymatic system of bacteria, thereby killing it.

Chang demonstrated by showing that hydrogen peroxide or potassium permanganate which release considerable amounts of nascent oxygen, nevertheless exhibit only weak germicidal activity. Chang was able to demonstrate that there is no liberation of oxygen involved in the inactivation of bacteria by chlorine.

In any case, oxygen gas $(O_{2})$ is a much weaker oxidizing agent than hypochlorous acid, as shown in the table below. Hypochlorous acid is likely to be effective as both an oxidizing agent as well as direct chlorination of the protoplasm and reaction with lipoproteins to form toxic chloro compounts that interfere with cell division (Chang, 1944).

Hypochlorous acid $(HOCl)$ dissociates in water to $H^{+}$ and hypochlorite ion $(OCl^{-})$.

$HOCl\leftrightarrow H^{+} + OCl^{-}$

Because hypochlorous acid $(HOCl)$ is uncharged, it is better able to penetrate cell walls than other chlorine species and is about 80 times more effective than hypochlorite $(OCl^{-})$ at chlorination disinfection.

The microbial disinfection due to total chlorine ($HOCl + OCl^{-}$) therefore depends upon pH and temperature. Whilst low pH and temperature favour disinfection by chlorine as $HOCl$, rate of diffusion through the cell membrane occurs faster at high temperature. The metabolic activity of the cell,rate of reaction of chlorine with enzymes and hence rate of inactivation of the cell is faster at higher temperature.

Disinfection by redox reaction using ozone gas $(O_{3})$ (a fast and powerful oxisizing agent) is sometimes used as an alternative to chlorination to reduces or eliminate the toxicity caused by residual chlorine, especially where there is concern over organic matter in the water which may produce trichloromethanes.

Sources:

Chang, S.L., Destruction of microorganisms, J. Am. Water Works Assoc., 36 (1944), pp. 1192–1206

Green, D.E. and Stumpf, P.K., The mode of action of chlorine, J. Am. Water Work Assoc., 38 (1946), pp. 1301–1305

Water Quality Control Handbook, Second Edition, by E. Roberts Alley (2007, McGraw Hill).

White's Handbook of Chlorination and Alternative Disinfectants, 5th Edition, by Black & Veatch Corp (2010, John Wiley & Sons)


Vacuum-jacketed distilling column I was looking for some distillation equipment on Sigma Aldrich because I wanted to perform some vacuum distillation, when I came across a vacuum-jacketed distilling column.

Now as far as I understood my high-school chemistry, I can perform vacuum distillation perfectly fine with a 'normal' column as long as there is an entry to which I can connect my vacuum source (in my case a cheap aspirator, but he it does the trick).

Before I'm buying all this pretty expensive stuff, I want to make sure that I buy the right equipment. As far as I can see it, this vacuum-jacketed distilling column has nothing to do with vacuum distillation, it is simply a vacuum jacket around a normal column. But then, where it is used for and why, or am I wrong?

The purpose of the vacuum jacket is to provide insulation of the column from cooling by ambient lab air.  You don't want to have to heat the "pot" (the flask containing the crude material you want to purify) any more than you have to, otherwise heat sensitive material in the pot may degrade and turn to tar.  Imagine if you cooled the exterior of your distillation column with cold water while performing the distillation.  You'd have to heat the pot a LOT more to get the material in the pot to finally reach the top of the column, then if you weren't cooling the column.  Air will have the same effect on your distillation, albeit to a smaller extent, than cold water.  So to avoid putting unnecessary extra heat into the pot, you vacuum jacket your column to preclude undesired cooling by ambient air.

Remember, a distillation column is different than a condensing column.  You don't want to cool a distillation column, you do want to cool a condensing column (or condenser).  Back in the day, very long columns were sometimes needed in the organic lab (5-10 feet) to separate closely boiling materials, they are still used in industrial applications.  If you had a 10-foot column that wasn't vacuum jacketed, you would turn most of your material to tar before it distilled over, because of the extra heat you had to supply to make up for the undesired cooling of the column by ambient air.  Eventually improvements like column packing materials (glass beads for example) and spinning bands (a metal helix rotating in the column) came along to increase column efficiency and allow chemists to be able to get good separation efficiencies (measured as "height equivalent theoretical plates", HETP) without having to resort to long columns.  Still a vacuum jacketed column is preferred to lessen unnecessary heating.


Generating neat orbitals/surfaces from molden/wfn-files I usually use Gaussian/GaussView and don't have to bother with other programs as I think that GaussView is capable of producing nice orbital/surface pictures like this one of one of uracil's Kohn-Sham-orbitals.

But when it comes that I have to use other programs than Gaussian - such as Turbomole, ORCA or Molcas - I am confronted with molden- or wfn-files that GaussView cannot handle.

Avogadro 1.1.1 (Lib 1.1.1, OB 2.3.2, Qt 4.8.5)

I quite like Avogadro because it can produce as fancy pictures as GaussView but especially for molden files at least on my computer it seems to fail hard. As you can see the orbital is reproduced perfectly from the file but Avogadro somehow cannot manage to show all atoms.

I hope that I'm not the only one around who can't manage this to work properly.

Thanks to Geoff's comments I learned that wfn-files can be imported through the Extensions $\rightarrow$ QTAIM menu which allows AIM analysis.

gmolden 4.x or 5.x

I have absolutely no idea which version this is.

What seems to be easier to open molden-files with molden? I guess that's something to try out. But for me molden stops being an alternative since I am unable to arrange my molecules in space as much as I want. Also molden produces only those stacked-layer orbitals that are not really up to date. Molden can produce pretty good orbitals after one gets how it works. Thank you hokru!

I am admittedly not a big fan of it and didn't try to get more out of it. I was always searching for better alternatives. I'm still no big fan but as those things that I want are implemented and I just have been unable to click the right buttons molden is getting an appropriate way to visualize those orbitals.

Gabedit 2.4.8

Now Gabedit is fancy ... there are a lot of input- and output-files that are supported - luckily molden files are part of it.

At least I can create orbitals somehow acceptable for me and arrange my molecule in space as much as I want it. There is also the povray-output stuff but I was not very successful with this approach, yet.

Now what stops me from using Gabedit at a regular basis. It is that there is everytime this point-and-click-adventurish menu -> orbitals -> selection -> again orbital selection -> quality settings -> isovalue settings -> wait and look. The settings were wrong? That's bad ... do the most things again and choose better values. Is there a simple way to compare orbitals fast? Not really, because the orbital slide show is a little bit tricky, too.

Did I say something about wfn-files? I think they are not supported.

Jmol 14.2.4_2014.08.03

Jmol is another program that can open molden files and I am again unable to create nice orbital images with this program. Even rising the so called "resolution" up to a reasonable degree only produces hexagonal-shaped orbitals.

Wfn-files? Who needs this? I guess no one except me.

Multiwfn 3.3.7

Now there is this free fancy little program from china. A big 400 pages manual/tutorial-pdf comes with it. With it one can quickly have a look on orbitals with a possibility to quickly change isovalues. It lacks unfortunately my favored free arrangement in space plus there are no proper lighting settings ... and the orbitals and structures are not really that beautiful.

But(!) it can read molden files and even wfn-files. It is also able to convert them to cube files that I afterwards can read in with GaussView to create what I want. It can also create those cube files for nearly all other surfaces that it can calculate. I like it much for this simple ability.

Conclusion

Well ... as I can use Multiwfn to convert the files that GaussView cannot read in by itself, there is always this curious feeling that there has to be a simpler alternative to all this. A program that can produce nice images without all the lackings and misbehaviours.

How do you do this stuff for visualisation of molden-/wfn-files to produce nice images of your orbitals and molecular surfaces?

Often different quantum chemical programs require different analysing tools. Unfortunately there is no universal standard (program), that can do it all. If there were, someone would come along and just invent something new.

I personally prefer ChemCraft for producing publication ready images. It is a commercially distributed program. It mainly supports any Gaussian and GAMESS file formats (and Firefly by extension). It can visualize quite a few more formats now. What I especially like about the program is that it can read the files generated by the NBO programs. It makes visualisation of all kinds of localised and canonical orbitals much simpler.
There are some more features of the program I really got attached to, like symmetrizing molecules or building molecules from scratch. It has the big advantage of writing the symmetry adapted data group for GAMESS for custom basis sets, which I have not been able to find anywhere else. But this is certainly not the target of this post.
 Unfortunately it is not able to read molden or wfn[1] files.

The program comes with a large number of visualisation schemes, which can easily be modified to produce a personalised style. Below is a small demonstration.

Most of the times I would use one of the above, depending on whom I have to impress. Often it is far more beneficial to have monochrome images, like the ones below.

A quite recent addition came with styles that allow you to generate simple 2D like structure. It is still possible to overlay these structures with orbital plots. (Unfortunately the transparency is hard-coded and cannot be set explicitly.)

The caveat is obviously, that it cannot read molden or wfn[1] files, not even Turbomole output.

A feature that I am still missing is the batch production of orbital pictures. You still always have to save them by hand, which can be incredibly tedious for large molecules (and large active spaces). (But I have written a feature request for that.)
Support for the program is quite good and the available help files are even helpful. Most of the features can be understood intuitively, however, as always, it takes time to get used to working with it.
Another caveat is, that the windows version outperforms its little linux brother by a lot. If you already run windows somewhere, that is probably fine for you. If you run linux almost everywhere, then you have to deal with some missing features (or recently it can be ported via wine, making it slow though).

I do use molden (Xwindow) quite a lot for intermediate visualisation, since it is one of the fastest interfaces available. It is very robust and has low graphics; you can easily use it whit a rather slow internet connection. It is free for private and non-commercial usage. Since my first days as a computational chemist involved the use of 56k-modem for checkup routines, this was one of the only go to programs. It will always be the first program I install on a work-dedicated computer. One simply does not forget his or her oldest friend.
A very big plus is the very well structured and written postscript files produced by molden. Some of the purists might even like it. Here are a few converted samples, but you can download the source here (via tinyupload.com).

Such files can become extremely handy, when preparing posters (especially when using LaTeX). I have on occasion seen molecular plots generated by molden in publications; I am even guilty of that. However, nowadays I believe there are more elegant methods.

Visualising Turbomole output has been, is and will probably always be a painful experience. For some of the deeper features you should probably be an expert in gnuplot, too. For the basics of structure and orbitals, the current recommendation is to use TmoleX, the GUI to turbomole. If it is done by the same people who designed define, then it's will be quite painfully to use. When I was a heavy user of turbomole, this did not exist, so I should probably shut up.
Judging from the current (version 7.0) manual the workflow I used is still somewhat recommended. It is to let the native modules create the binary *.plt files, which are gOpenMol input. Unfortunately, what was once a very aspiring program is nowadays not developed any more. You can still obtain an unsupported copy. The source code is available, but the current license state of the program is somewhat unclear. Since VMD (see Brians post) and TmoleX do now support gOpenMol files, the necessity of learning a deprecated program is probably not useful any more. One of the prime features of this program however was that it has a tcl interface and hence was fully scriptable - apparently something that VMD now also has.
It is well capable of producing nice orbital pictures, you might have seen them on this page already. (I tweaked the Gaussian output a bit, since I do not have Turbomole available.)

Although it supports quite a few file formats, molden input and wfn[1] is unsupported.

[1] Let's come to the last part: wfn files. However you produce them, they are probably not up to date. The main (and probably only) purpose for those files is the analysis with AIMPAC -  a program that is no longer developed. They have one very big disadvantage: they cannot deal with functions higher than of f type. If you are using a triple zeta basis set, you probably have those functions. Another disadvantage is, that they don not include information about possible effective core potentials.
For the purpose of analysing the electron density with the quantum theory of atoms in molecules (QTAIM) the much superior version wfn extended (*.wfnx or *.wfx) should be used. I am currently only aware of Gaussian to create such files correctly. The main program which uses these files is AIMAll. It is also a commercially available program. You can learn more about these files on the specification page there.
The only valuable point of creating molecular orbitals with this program is, if you want to overlay it with the bond paths instead of just sticks. Just like in this example:

Apart from this usage, the process is very tiring and needs many steps to be completely performed. Working the program non-local is very, very slow. And to be perfectly honest, what is the purpose of the above other than impressing a few organic chemists?
Another negative aspect of the program is the very high demand of memory and storage, especially for large molecules.

In conclusion: If you are not doing QTAIM analysis you should probably refrain from *.wfx files for visualisation purposes. For every QC program there usually comes a workable interface/visualisation program. With any program there is the activation barrier of learning how to do things to overcome. In any case, I suggest having a few ready to work with, you never know where you go and you might have to learn a new program any time.
I think the biggest loss of any visualisation tool is the neglect of writing good tutorials or even a basic documentation.


How sure can we be about the structure of purine? While answering the question Basicities of nitrogen atoms in purine, I noticed, that there might be a rapid interchange of the hydrogen bonded to the nitrogen atoms.
I have found a crystal structure in D. G. Watson, R. M. Sweet and R. E. Marsh, Acta Cryst. 1965, 19, 573-580. (DOI: 10.1107/S0365110X65003900, CCDCID: 1239746, 1239747), that resolves it as 1. However, at least in solution I would assume that there should be a distribution of structures according to the scheme below. In all cases the aromaticity is retained.

I am not an expert on crystal structures, and I don't know how to read the .cif file correctly. I remember from my course on crystallography, that the positions of hydrogen are notoriously hard to confirm.
Can we therefore really be sure about this structure? Is there evidence in the crystallography data for a large disordering?

To provide a bit more context, I am asking this, because it is likely to change part of the answer I have given to the linked question. If there exists an equilibrium between the above protomers, then that would certainly make it even harder to find the most basic nitrogen in purine.

The article text (Ref.1) itself does go into details. The first thing to note: while aromaticity on the larger scale is retained in your structures 3 and 4, this is only true for the 10-π-electron system as a whole. In many of these types of 10 π systems made up of a six-membered and a five-membered ring, the two rings are not equal with the double bond in the five-membered ring often being much more susceptible to hydration. This prior thought should bias us against structures 3 and 4.
A pretty obvious first thought would be to take a look at the arrangement of individual molecules in the crystal structure. The paper does this in figure 3 which I have reproduced below. Wherever we expect the single proton, it should surely form a hydrogen bond to another purine molecule, meaning that these two nitrogens should be sufficiently close and in a favourable angle arrangement.

As the figure shows, the two nitrogens of the five-membered ring are in an arrangement that would suggest a hydrogen bond between them. The cut-off bits outside the unit cell are fragments of the next purine molecule not completed for brevity. This arrangement rules out 3 and 4 on the basis of no available hydrogen bonds and limits the discussion to 1 and 2.
As for the question whether the hydrogen prefers structure 1 or 2 this is admittedly more complex. As the authors note:

Since there is a possibility of tautomerism of the hydrogen between N(7) and N(9), particular attention was given to this problem. Two least-squares cycles, not quoted in Table 2, were computed in which a half-hydrogen was assigned to each of N(7) and N(9). Refinement of H(10), H(ll), H(12) and the two half-hydrogens resulted in the half-hydrogen attached to N(9) shifting to a position about 1 Å from N(7) of the neighbouring hydrogen-bonded molecule.

So basically, during refinement Watson attempted to put the hydrogen on both nitrogens equally. However, the refinement would show that the partial hydrogen on $\ce{N}$9 was moved sequentially along the $\ce{N\bond{...}H\bond{...}N}$ distance towards $\ce{N}$7. Taking the final $\ce{N-H\bond{...}N}$ distance of $\pu{285pm}$ and the $\ce{N-H}$ bond at $\pu{85pm}$, this means that there are approximately $\pu{115pm}$ between the two possible hydrogen positions on $\ce{N}$7 or $\ce{N}$9—and most of that was covered by the moving hydrogen. However, Watson also notes that the electron density corresponding to that $\ce{N-H}$ hydrogen was ‘much more smeared’ that those of the other hydrogens.
Sweet and Marsh arrive at the same conclusion using a different method that I did not fully understand.
All things considered, this gives us a pretty strong case that purine is best represented by structure 1 and that the principal tautomerisation structure would be 2 as that is where the hydrogen bond is directed.

Note: The authors state in the introduction to their paper:

This paper presents the results of two entirely independent investigations of the structure of purine (I).
The duplication of effort was discovered only when both three-dimensional studies were almost complete (Watson, 1964). Since different experimental techniques had been used, it was considered appropriate to publish both sets of results in the same paper and to make the necessary comparison between them. It has been a comfort to both parties that, except for systematic differences in the temperature factors, the two sets of results differ by less than their standard deviations.

This is why I focused so much on Watson and ‘ignored’ the other authors as the principal arguments are from Watson’s part of the paper.

Reference:

D. G. Watson, R. M. Sweet, R. E. Marsh, “The crystal and molecular
structure of purine,” Acta Crystallographica 1965, 19, 573-580
(https://doi.org/10.1107/S0365110X65003900).


High-temperature DFT Why is it so difficult to perform DFT calculations that consider temperature?

I have seen that time-dependent DFT is needed to model systems at high temperature.
Why is this the case? What about finite temperature functionals? Why is it also acceptable to use zero-temperature functional with TD-DFT?

I have also read that thermal fluctuations can result in the occupation of electronic states (excited states) above the ground state.

I have tried to rationalise the fact that the electron density changes with temperature as:

If the ions in the until cell move, their degree of orbital overlap and electron density will change. The degree of orbital overlap influences the band-structure, which then impacts all of the electronic and physical properties.

Is this accurate?

First if all, I appreciate you and thank you for putting up such a good question.

One thing I will request you to note here is the definition of temperature. Temperature, in broad sense, is defined as the average kinetic energy (for equilibrium conditions, I hope you are speaking of that) of a system. The contribution mostly arises from translational kinetic energy, and the fluctuation in temperature is proportional to $\frac{1}{N-1}$. (I don't remember if it has some power or something.) In other words, a particle has infinite temperature fluctuation. So if you want to do finite temperature DFT calculations, you will need an ensemble of particles, and it will only add the cost.

In case you have infinite computational power (or at least enough to support no of molecules in the order of Avogadro number), you can, in principle, go for DFT at a finite temperature.

If you are worried about occupation of the higher states, you can calculate the higher states using Kohn-Sham Hamiltionian (these are just eigenfunctions of the Hamiltionian), and then calculate the occupation at a temperature using the formula for the concerned ensemble. Since the states of the KS hamiltionian are its eigenstates, these should not change if you have a static distribution of electrons. The eigenstates mean that these are decoupled by the same KS Hamiltonian, and in principle, you are not allowed to take an electon from one state to other without external purturbation that couples the states.

However, at a given temperature, everything is dynamic. This invokes that the states have finite lifetime, which from Uncertainty principle, imposes broadening in energy, i.e. it is not an eigenstate anymore. This imply there can be coupling between different states, and hence electrons can jump between states without external perturbation. You can externally add broadening to the states using broadning matrix, that mimics the temperature to some extent.

For TDDFT, you actually introduce an alternate electric field, and then this purturbation is used to calculate coupling between the concerned states. Thus, TDDFT calculates the excitation using light. It has nothing to do with the temperature.

Please let me know in comments if you have further doubts.

Greetings from India. जय हिन्द. Jai Hind.


How to construct NMR spectra from chemical shift tensors? If I know the chemical shift tensor (from ab initio methods), $\sigma$, for each atom in my structure, is it possible to construct an NMR spectra for my system? If yes, how and what mathematical implementations do I need?

Yes, it is possible. It is done often. The solution depends on the dynamics and orientation(s) of the spins in the ensemble (as will be shown bellow). I will assume you want to calculate the spectrum in frequency space.

$$
\omega_{cs} = \sigma_{zz}^{LF}(-\gamma \vec{B}_{0})
$$

We can hide a few variables and rewrite this in a much more useful way as:
$$
\omega_{CS} = (\vec{b}^t_0 \overset {\leftrightarrow}{\omega}\vec{b}_0)^{LF} = (\vec{b}^t_0 \overset {\leftrightarrow}{\omega}\vec{b}_0)^{PAS} =(\vec{b}^t_0 \overset {\leftrightarrow}{\omega}\vec{b}_0)^{RF}
$$

where,

$$
\vec{b_0} = \frac{\vec{B_0}}{B_0}.
$$

The above allows us to calculate the chemical shift in many reference frames. The frame where $\omega$ is diagonal (the PAS) is often convenient.

$$
\omega_{CS} = \omega_{xx}^{PAS}(\vec{b}^{PAS}_{0,x})^2 + \omega_{yy}^{PAS}(\vec{b}^{PAS}_{0,x})^2 + \omega_{zz}^{PAS}(\vec{b}^{PAS}_{0,x})^2
$$

We can define $\theta$ as the angle between $Z^{PAS}$ and $b_0$ and $\phi$ as the angle between $X^{PAS}$ and the x-component of the $b_0$ field. This lets us write:

$$
\omega_{CS} = \omega_{xx}^{PAS} \cos^2\phi \sin^2\theta + \omega_{yy}\sin^2\phi \sin^2\theta+ \omega_{zz}^{PAS}\cos^2\theta
$$

The above is the main equation you will need to calculate a chemical shift spectrum in frequency space. Numerically you will sum over all the angles in your ensemble. If your sample is a powder you could use angles sampled from a Gaussian distribution on a sphere for example. Typically people then apply convolution with a Lorentzian function (to account for relaxation and smooth the numerical artifacts). If your sample is a solution and the molecular reorientation time is fast compared to the inverse of the anisotropy, everything but the isotropic component will be averaged away. To put this in context (and to get to some possibly familiar equations) it will be useful to introduce some convention. In the Haeberlen-convention:

$$|\delta_{zz} - \delta_{iso}| \ge |\delta_{xx} - \delta_{iso}| \ge |\delta_{yy} - \delta_{iso}|$$
$$\delta_{iso} = \frac{1}{3}(\delta_{xx} + \delta_{yy}+ \delta_{zz})$$
$$\delta = \delta_{zz} =\delta_{iso}$$
$$\eta = (\delta_{yy}-\delta_{xx})/\delta$$

Where $\delta$ is called the anisotropy and $\eta$ is called asymmetry. Graphically:

Using this convention and some geometry you can get to:
$$
\omega _{CS} = \frac{\delta}{2} (3\cos ^2\theta -1 + \eta \sin^2 \theta \cos 2\phi) + \delta_{iso}
$$

Also, ab initio chemical shift calculations normally are reported in shielding values, you will probably want to convert to shift values using a sutable reference.

Reference:
Multidimensional Solid-State NMR and Polymers, by Klaus Schmidt-Rohr and Hans Wolfgang Spiess, chapter 2


Law of chemical equivalence In this reaction,
$$\ce{KIO_3 + KI -> I_2}$$
According to law of chemical equivalence, shouldn't $m_{eq}$ of $\ce{KIO_3}$ = $m_{eq}$ of $\ce{KI}$ = $m_{eq}$ of $\ce{I_2}$?
However, the book that I was reading from, said:
Total $m_{eq}$. of $\ce{I_2}$ = $m_{eq}$ of $\ce{I_2}$ (from $\ce{KIO_3}$) + $m_{eq}$ of $\ce{I_2}$ (from $\ce{KI}$).
Why is it so?

This Reaction is a Disproportionation Reaction, that is, a reaction involving both oxidation and reduction of a single species, in this case : Iodine. This is not one reaction but a combination of two reactions:
$$\ce{KIO3 -> I2}$$
And
$$\ce{KI -> I2}$$
Therefore, by applying law of chemical equivalence in both reactions and adding the equations :

$m_{eq}$ of $\ce{I_2}$ (from both reactions) = $m_{eq}$ of $\ce{I_2}$ (from $\ce{KIO3}$) + $m_{eq}$ of $\ce{I_2}$ (from $\ce{KI}$)


Nucleophilicty order of various nucleophiles Which one is more nucleophilic in strength $\ce{Br-}$ or $\ce{OH-}$?
I am confused since $\ce{OH-}$ is more electronegative than $\ce{Br-}$ but also a strong  base which contradicts for a better  nucleophile .

Like most concepts that are not captured by a single number, nucleophilicity is not a simple (or at least, relatively simple) concept like acidity.

Nucleophilicity depends a lot on orbital alignment and interaction, in addition to thinking about charges. It's also confusing because nucleophility is not always married to the stability of the adduct. For example, iodide is a great nucleophile, but alkyl iodides are also highly reactive.

In your specific case, since the two nucleophiles have the same charge, you should look at other factors. Bromide is significantly larger than hydroxide, and it is a much better nucleophile, at least when we're considering carbon electrophiles. Indeed, we see bromide used in nucleophilic catalysis but not hydroxide.


How to obtain the radial probability distribution function from a quantum chemical calculation? In the book "Quantum Chemistry" by Ira Levine I found a plot of the radial probability distribution function of argon:

The figure appears in the context of the following section:

  Electron densities calculated from Hartree–Fock wave functions are quite accurate. Figure 11.1 compares the radial distribution function of argon (found by integrating the electron density over the angles θ and φ and multiplying the result by r2) calculated by the Hartree–Fock method with the experimental radial distribution function found by electron diffraction. (Recall from Section 6.6 that the radial distribution function is proportional to the probability of finding an electron in a thin spherical shell at a distance r from the nucleus.) Note the electronic shell structure in Fig. 11.1. The high nuclear charge in 18Ar makes the average distance of the 1s electrons from the nucleus far less than in H or He. Thus there is only a moderate increase in atomic size as we go down a given group in the periodic table. Calculations show that the radius of a sphere containing 98% of the Hartree–Fock electron probability density gives an atomic radius in good agreement with the empirically determined van der Waals radius. [See C. W. Kammeyer and D. R. Whitman, J. Chem. Phys., 56, 4419 (1972).]

Source: Ira N. Levine. Quantum Chemistry 7th ed.; Pearson, 2014; page 294.

What kind of chemical software allows one to do this?

It is possible to do with Gaussian, and, of course, how is it possible?

This is gladly not a big problem as MultiWFN (open source) can do that for you. I'll guide you through the process from start to finish. My current setup is Gaussian 09 Rev D.01 and MultiWFN 3.3.8; there are newer versions out there with which this should work, too.

Run the calculation, for example ar.com at the HF/STO-3G level of theory, but anything else will also be fine:

%chk=ar.chk
#P HF/STO-3G
gfinput gfoldprint iop(6/7=3)
symmetry(loose)
pop=full pop=nbo6

title card

0 1
Ar    0.0  0.0  0.0

This will generate a checkpoint file ar.chk which needs to be formatted:

formchk -3 ar.chk ar.fchk

Now fire up your MultiWFN program, you should get a prompt similar to the following:

 Multiwfn -- A Multifunctional Wavefunction Analyzer (for Linux)
 Version 3.3.8, release date: 2015-Dec-1
 Project leader: Tian Lu  (Beijing Kein Research Center for Natural Sciences)
 Citation of Multiwfn: Tian Lu, Feiwu Chen, J. Comp. Chem. 33, 580-592 (2012)
 Multiwfn official website: http://Multiwfn.codeplex.com
 Multiwfn official forum (in Chinese): http://bbs.keinsci.com
 Bug reporting, question and suggestion, please contact: Sobereva@sina.com
 ( The number of threads:  2   Current date: 2017-03-09   Time: 14:46:57 )

 Input file path, for example E:\Haiyore!Nyaruko-san\Nyaruko.wfn
 (Supported types: .wfn/.wfx/.fch/.31/.chg/.pdb/.xyz/.cub/.grd/.molden, etc.)
 Hint: To reload the file last time used, simply input the letter "o". Input such as ?miku.fch can open miku.fch in the same folder of the file last time used

If you loaded MultiWFN from the directory you're files are located, simply put in ar.fchk. It will load the file and output some basic information and take you to the "Main function menu":

 Please wait...
 Loading various information of the wavefunction
 Loading basis-set definition...
 Loading orbitals...
 Converting basis function information to GTF information...
 Generating overlap matrix...
 Generating SCF density matrix...

 Total/Alpha/Beta electrons:     18.0000      9.0000      9.0000
 Net charge:     0.00000      Expected multiplicity:    1
 Atoms:      1,  Basis functions:      9,  GTFs:     27
 Total energy:  -521.222880802727 Hartree,   Virial ratio:  2.01704078
 This is a restricted single-determinant wavefunction
 Orbitals from 1 to     9 are occupied
 Title line of this file: title card
 Formula: Ar1
 Molecule weight:        39.94800

 Loaded ar.fchk successfully!

                    ------------ Main function menu ------------

For the sake of sanity I will not copy every menu screen here, just simply what you have to put in. On this screen the radial distribution function is hidden in the "Other functions (Part 2)" and there it is option 5, also see the manual chapter 3.200.5 on page 175 (v.3.3.8). So choose

200 // Other functions (2)

and then

5 // RDF

You should get the following menu:

   ====== Plot radial distribution function for a real space function ======
 -1 Exit
 0 Calculate radial distribution function and its integration curve
 1 Select real space function, current:   1
 2 Set sphere center, current    0.0000    0.0000    0.0000 Ang
 3 Set lower and upper limit of radial distance, current:   0.0000   5.0000 Ang
 4 Set the number of integration point in each shell, current:  2030
 5 Set the number of radial points, current:   500

Here you can select a couple of options. The electron density is already preselected, but you can change it in option 1. You'll get a table with all available functions.
The only thing you might want to adjust are the limits, so choose

3 // Limits
0.0,2.0

and calculate the RDF:

0 // Calculate RDF

Choose 1 to plot and view the function. Or choose

3 // Save graph

or choose

5 // Export to RDF.txt

for post-procesing (with gnuplot).

You can shorten the journey if you prepare a file like this mwfn.in

ar.fchk
200 // Other functions p2
5   // RDF menu
1   // Select real space fucntion
1   // Electron density
2   // Set sphere centre
0.0,0.0,0.0
3   // Set limits
0.0,2.0
0   // Calculate RDF
3   // Save graph
5   // export to text file

and pipe it into the program:

multiwfn < mwfn.in

The native graphing capabilities of MultiWFN don't allow for customisation, so you might want to run it through gnuplot. Here is a file which gives you the graph below, I am using gnuplot 4.4, but later versions are more feature rich.

# General Settings
set terminal svg enhanced font "Arial,12" size 800,600
set tics font "Arial, 9"
set output "ar-rdf.svg"
unset key
#or set key top right Left # etc
set title 'Argon Radial Distribution Function HF/STO-3G'
#
# X-Axis: radius / angstrom
set xlabel 'r / Å'
set xtics in scale 0.6, 0.4 nomirror
set xtics 0,1
set mxtics 2
set xrange [0:2]
# Y-Axis: RDF/a.u.
set ylabel "RDF / a.u." rotate left
set ytics  in  scale 0.6, 0.4 nomirror
set ytics 0,10
set mytics 2
set yrange [0:25]
#
#
# PLOT!
plot \
     "RDF.txt"   using ($2):($3) title "HF/STO-3G"   with lines lt 1 axes x1y1

This is the resulting graph:


Diastereotopic methylene protons in achiral molecules Are the protons in methylene groups of achiral molecules such as propane-1,2,3-triol or 3-hydroxypentanedioic acid diastereotopic?
Consider propane-1,2,3-triol. To me, replacing one of the protons in either -CH2 with an arbitrary X nuclei creates (what I work out as) 2 diastereomers. However, NMR simulations of propane-1,2,3-triol in ChemDraw and on nmrdb.org (https://www.nmrdb.org/new_predictor/) do not flag these hyrdogens as diastereotopic and neither does this "Diastereotopic atoms" computational tool (http://www.cheminfo.org/Spectra/NMR/Tools/Diastereotopic_atoms/index.html).
I have also played around with physical models of each isomer and cannot find a plane of symmetry to make them enantiomers.
What am I missing here?

Well, I can't speak to your NMR simulator and computational tools but from my perspective, and that of ChemDraw, the hydrogens at C1, and C3, in diol 1 are diastereotopic. While diol 1 is meso and achiral, introduction of deuterium in 2a renders both C1 and C2 stereogenic. Consequently, 2a is a diastereomer of 2b where the deuterium has been moved to the geminal C1 position. You may want to look at  this.<b


If a temperature of a reaction is raised from 300 to 320K, what factor will the reaction rate increase by In general, if the temperature of a reaction is raised from 300k to 320k, the reaction rate will increase by how much?

So the first thing I do is try out the Arrhenius equation:

$k=Ae^{\frac{-E_a}{RT}}$

so I ignore the activation energy and A because they will cancel out to get the following:

$k=e^{\frac{-1}{300}}$

$k=e^{\frac{-1}{320}}$

and I divide the two k's to get 1.000208

but that can't be right.

It seems to me that I do not know how to answer this question and would like some assistance

I do not want to know a proof of anything but rather to know how to execute methods in solving this problem.

The amount of increase is dependent on the activation energy. For a very low activation energy, there will be almost no change (this is what you obtained by removing $E_a$ and $T$). For a high activation energy, the change will be quite large. For example, with an activation energy of $110\mathrm{\frac{kJ}{mol}}$, you the rate will increase by about $16$ times. $$\frac{\exp({\frac{-E_a}{R*320}})}{\exp({\frac{-E_a}{R*300}})}=\frac{\exp({\frac{-111000\text{J}}{8.314*320}})}{\exp({\frac{-111000\text{J}}{8.314*300}})}=\frac{7.6\times10^{-19}}{4.7\times10^{-20}}=16.14$$

So there really isn't an "in general" answer. However, many text books give a crude estimate that the rate will double when the temperature is increased by $10$. This isn't in general the case, as it depends on the temperature range (a change from $300$ to $320$ is more significant that a change from $600$ to $620$) and the the value of the activation energy.


Why is butter harder if it once got warm? I realized that butter, when kept cool after buying, can be easily cut and spread when taken out of the fridge. If it gets warm once, however, it will be less soft the next time I take it out of the fridge. If it reaches room temperature once, it cannot even be cut when cooled down again to fridge temperature.

What is causing this phenomenon?

(I refer to real butter from cow milk, not what might be sold as butter in some countries)

Butter is a water in oil emulsion that contains butterfat that crystallizes at room temperature. The continuous phase of butter is a network of fat crystals. This network surrounds water and air droplets. One of the unique properties of fats are the interactions of fat crystals, and these interactions are directly related to product hardness. More fat crystal interactions lead to a harder fat. During storage, fluctuations in temperature from refrigerated to room temperature back to refrigerated temperatures result in postcrystallization. This postcrystallization strengthens the fat crystal network since more crystals are formed, leading to a stronger fat crystal network.


Addition of salt to ice Salt is very commonly added to ice even at room temperature to prevent it from melting. Can anyone explain how does it really work?

Uncle Al is correct.  Colligative properties are properties of solution in which only the ratio between number of particles of solute and solvent are considered (often expressed in different forms, such as moles of solute per kilograms of solvent for molality); the identities of either are irrelevant.  So we could dissolve sodium chloride in water or magnesium bromide, but the general properties will be the same (though note that the magnitude of freezing point depression and similar properties will be slightly different depending on the substance).

In particular for salt water, the point at which the solution will freeze is depressed.  To understand this, we must understand what occurs when pure water freezes.  In general terms, when a liquid freezes, its normally hectic and randomly moving particles form a stable, packed structure that we call a solid; a hexagonal structure in particular for water at standard conditions.  The specifics are more complicated, but this is the main idea.

However, if you introduce a solute into the pure water, such as salt, the solute particles are getting in the way of the formation of the packed structure.  As a result, the temperature must be lowered slightly further in order for the solvent (i.e. water) molecules to properly overcome the solute molecules acting as obstacles for the formation of the solid structure.

In fact, when the solvent freezes, more often than not, the solute particles are often excluded, either pushed outside or trapped inside small pockets.  In the case of salt water, a eutectic system is created, if you're interested in learning of them.  Salt water itself cannot freeze into salty-ice or otherwise.


Possible approximations in complex equilibrium If  I have the following equilibrium:
$$\ce{Cu^2+ + 4NH3 <=> Cu(NH3)4^2+}$$
Given an initial concentration of $\ce{Cu^2+}$ ($C_\ce{Cu}$) and an initial concentration of ammonia $C_{NH_3}$, the equilibrium constant $\beta$ can be expressed:
$$\beta=\frac{x}{(C_\ce{Cu}-x)(C_{NH_3}-x)^4}$$

My question is: is there any approximation I can use in order to solve this expression to know the value of $x$? $\beta$ is known.

Thanks

There is some chemistry behind this which you must understand to recognize the approximation. First the copper ammine complex doesn't really form in the manner indicated rather:
$$\ce{Cu^2+ + 4H2O <=>T[acidic solution] Cu(H2O)4^2+}$$

But in basic solution copper will form a gelatinous precipitate which can perhaps be represented as:

$$\ce{Cu(H2O)4^2+ + 2OH^{-} <=>T[basic solution] Cu(H2O)2(OH)2}$$

With ammonia in the aqueous solution various ammine complexes will be formed such that the soluble tetra-ammine complex is formed in a series of reactions something like:

$$\ce{Cu(H2O)2(OH)2(s) + NH3 <=> Cu(H2O)(OH)2(NH3)} $$

$$\ce{Cu(H2O)(OH)2(NH3) + NH3 <=> Cu(OH)2(NH3)2} $$

$$\ce{Cu(OH)2(NH3)2 + NH3 <=> Cu(OH)(NH3)3^+} $$

$$\ce{Cu(OH)(NH3)3^+ + NH3 <=> Cu(NH3)4^2+(aq)} $$

Ammonia is typically present in excess so the overall reaction is written "in shorthand" as the following equation (I highly doubt that the exact sequence of chemical reactions between the gelatinous precipitate and ammonia is fully understood):

$$\ce{Cu^2+ + 4NH3 <=> Cu(NH3)4^2+} \quad \text{K}_\text{eq}$$

and the equilibrium for this reaction is greatly favored to the right (the ammine complex form).

If you let $[\ce{Cu2+}]_i$ be the initial concentration of $\ce{Cu^{2+}}$, let $[\ce{NH3}]_i$ be the initial concentration of $\ce{NH3}$ and the final concentration of the ammine complex be $[\ce{Cu(NH3)4^2+}]_f$ then in excess ammonia a simplfying assumption would be:
$$ [\ce{Cu(NH3)4^2+}]_f \approx [\ce{Cu2+}]_i$$
$$[\ce{NH3}]_f = [\ce{NH3}]_i - 4[\ce{Cu(NH3)4^2+}]_f \approx [\ce{NH3}]_i - 4[\ce{Cu2+}]_i$$

the equations could of course be solved exactly, but you end up with a quadratic expression which really isn't a difficult mathematical problem to solve. It just takes a wee bit more calculating.


Reaction mechanism of glycerol with hydrogen iodide acid I was watching a chemistry video lecture today and I came across this:-

Basically the lecturer was illustrating the mechanism of reaction glycerol with $\ce{HI}$.See from right to left(follow the arrows :-P) in the screen shot.I have a doubt in the fourth step where it is shown by an arrow that the $\ce{I^-}$ ion shifts from terminal position to the central position therby forming a positive charge at the terminal position where another $\ce{I-}$ attacks.He says something like "The $\ce{I-}$ shifts to central position due to areal resonance(or something like that..hope I heard that word correctly) which increases stability".

Now that's my doubt!Why should the iodine ion shift its position at all?Isn't the positive charge more stabilized by hyper conjugation in the central position?And what on earth is areal resonance(I searched google too but no use)?
Can someone help me out ?Thanks in advance!

The iodine isn't changing positions, it's donating a pair of electrons to stabilize the charge by forming an iodonium ion similar to the bromonium ion formed during the halogenation of an alkene shown below. You can see that the bonded pair of electrons was not affected by the step. It looks like they confused the reaction by short-cutting the reaction step.


Homemade Rocket Fuel and Engine I want to put an action figure in orbit but I need to know how to make a secure yet not too  expensive rocket fuel engine. The design of the engine must very primitive and not too complicated.

I've read some information about hydrazine and hydrogenperoxide but it is too dangerous to use and make at a noncommercial scale. Although I studied some chemistry in highschool, I don't have so much knowledge about this matter. So far my best bet is blackpowder based fuel engine. What do you suggest?

Note: I'm not a modeler myself, but tested several recipes of rocket candy and read rocket modeler forum for a while.

First of all, you have to think if you will be satisfied with solid propellant or want a liquid fuel engine.

Building home-made liquid fuel rocket engine is very hard. One need to find pumps, make combustion chamber from something very heat-proof and reasonably stable chemically. Anyway, some successful reports may be found on the internetz. Note, however, that if you can choose safe liquid fuel, like kerosene or some alcohol, there is no safe oxidant for liquid-fueled rocket engine. The closest thing to safe oxidant I can imagine is nitric acid, that require work on fresh air or in specially designed room and tends to form explosives on contact with many things, like milk protein (yep!), some of which are volatile

Solid-fuel model propellants are much easier to prepare and use, there are numerous articles in the internetz about such fuels. The most common one is rocket candy, prepared from potassium nitrate and sugar. I tried it to see how it works (I didn't build any rockets), but instead of sugar I used sorbitol. Rocket candy based on sorbitol can be melted and then can be casted in any form. wiki article contains both info and links to sites with well-testes recipes by successful modelers.

Another possible option is making missile model based on some classical jet. There is quite simple pulse jet and there are some books about home-made turbine jet designs.

Anyway, I recommend to search nearby rocket modeling club and find supervisor.


Comparison of organic and inorganic acid strengths So I know how different factors affect acid/base strength, but I am not sure when they are all in the same question, which affects acid/base strength more.

So, for example this question:

How do I know that $\ce{HBr}$ is the strongest acid...because I also have the inductive electron withdrawal from Fluorine in answer choice B.

You will be likely be expected to memorize approximate pKa values for different functional groups. Molecules containing protons with lower pKa values are more acidic.

In this case, these are the functional groups you have:

a. alcohol (pKa ~16)
b. alkyl halide and carboxylic acid (pKa ~4-5)
c. hydrobromic acid (pKa ~-9)
d. carboxylic acid (pKa ~4-5)
e. ammonia (pKa ~35)

I would have to suspect that b is slightly more acidic than d, but they are going to be in similar ball parks. What you're looking for is the most and the least acidic compounds, and the pKas for those compounds are distinct in this question. Since the acidity of a group is inversely proportional to its pKa, the compound with the lowest pKa among its groups is the most acidic and the compound with the highest pKa among its groups is the least acidic. Things get a bit more complicated than that when multiple groups are involved, but this rule usually holds for simple cases.

Hydrobromic acid has the lowest pKa, so it is the most acidic. Ammonia has the highest pKa, so it is the least acidic.


What are some books, websites, etc. where I can find the detailed protocol for standardization of sodium thiosulfate with dichromate? I am looking for a book or website such as "Vogel's Textbook of Quantitative Chemical Analysis" where I can find methods for analyzing/standardizing a particular species easily.

I am not looking for an exact sub area such as water analysis or drug analysis; rather I want a resource like Vogel's which indexes generic analytical methods, both classical and modern.

Edit:
Actually I am looking for the method of standardizing a sodium thiosulfate solution using a potassium dichromate solution. I can't find the method in "Vogel's Textbook of Quantitative Chemical Analysis 6th Edition".

My go-to for this is typically Kolthoff & Sandell's Textbook of Quantitative Inorganic Analysis.  They unfortunately do not report a detailed protocol for dichromate standardization of thiosulfate solutions therein, but instead cite Volumetric Analysis, Vol III by Kolthoff and Stenger (Amazon and Google Books).  I don't have access to this volume; WorldCat shows numerous copies in university libraries in my area of the U.S.; more than likely there's a library where you are that has a copy.


Freezing Point Depression and Temperature Decrease I don't have a problem understanding the chemical mechanisms underlying freezing point depression or colligative properties in general, but what I cannot seem to grasp is how if you add salt to ice you end up with a salt-ice-water solution that is actually colder than the ice you started with.

Two other places on this site go into detail about this matter but I feel like it is either not totally correct or just not fully explained. I assume I must not fully understand in the first place, also. The two links are posted immediately below pertaining to the subject.

Where does the energy come from to lower the temperature of a brine solution?

Why does ice water get colder when salt is added?

How can the ice be cooling as it melts, as these links seem to claim?  I understand that it would be cooling its surroundings by needing to absorb heat to melt, but it's a phase transition so the temperature of the ice should remain constant. Furthermore, if the temperature remains constant, and the ice continues to absorb heat from the water surrounding it in order to melt since it's above its depressed freezing point, wouldn't you have the impossible feat of heat flowing from cold to hot as the water gets continuously colder, as alleged?

Energy is being used in the phase change from solid ice to liquid water. Heat is being absorbed by the ice. Yes, the ice freezes at a much lower temperature and so it will be liquid at a certain range of higher temperatures than it otherwise would be without salt. As it melts it tries to approach its new
ice-water-salt equilibrium which will be at a depressed freezing point. Why does the temperature fall? Especially if phase changes occur at constant temperature?

The energy must be coming from somewhere and going to somewhere, but where? And how can a hot-to-cold temperature difference for heat transfer actually be maintained if the salt solution is getting colder but the phase change (melting ice) is constant temperature?

I'd really appreciate some sort of insight to this.

Phase change is not at constant temperature, phase equilibrium is

  [OP] it's a phase transition so the temperature of the ice should remain constant

Here is a counter example: If you add ice cubes to hot water, the ice will melt, cooling down the hot water. In this system, there is no thermal equilibrium, so not all of the components are at the melting point of water. If you wait a while, you might see the following: All the ice is gone, or some ice and some water remains, or (if the ice was really cold) all the water is frozen. Only if there is some ice and some water left do you know the equilibrium temperature - equal to the melting point.

A system can gain thermal energy from other types of energy

If you stir a viscous solution, it will get measurably warmer. (This is a classic experiment of thermodynamics.) There is no heat transfer - nothing gets colder. For melting ice, it is a conversion of thermal energy into chemical energy (breaking hydrogen bonds).

  How can the ice be cooling as it melts?

It can't, at least not for pure water (see below for adding water). The temperature of the ice is below the melting point (or at it), and the temperature of the water is above the melting point (or at it, or a bit lower if it is super-cooled). The thermal energy will come from the liquid water, cooling it. If everything is at the melting point, melting and freezing will happen at the same rate, and nothing changes temperature. In a real system where surrounding is warmer, thermal energy from the surrounding flows into the system, melting the ice while the system stays at the melting point (there will be small fluctuations so that the thermal energy reaches the ice).

Scenario of adding salt

Starting with an ice/water equilibrium (at melting temperature), we can add salt (at the same temperature, and let's say there is no enthalpy of dissolution). After adding the salt, the temperature is still the same, but the melting temperature is lower. With the temperature higher than the melting temperature, water molecules break away from the ice, which requires energy. This energy comes from thermal energy in the liquid and in the remaining ice. As a result, the entire system cools down.

Will the temperature change?

  In your example where ice cubes were added to hot water, wouldn't the pure water ice cube remain at 32 F until it fully melted? How is this any different than having an ice cube on a hot plate and melting it? The temperature during the heating of the cube doesn't increase because the absorbed heat is being used to break the chemical bonds of the solid ice.

If the system is not at thermal equilibrium (different temperatures within the system), some parts of the system have to change temperature before equilibrium is reached. The hot water will certainly decrease in temperature. If the ice cubes are colder than the melting point at the start, they will warm up until they have reached the melting point. (Ice can have any temperature below the melting point. For example, ice from a freezer is at about $\pu{-20^\circ C}$.)

If the system is at thermal equilibrium from the start but not at phase equilibrium (or not a chemical equilibrium for a reaction that is endothermic or exothermic), the temperature will also change. In my example of adding salt to a water: ice mixture, the process of melting is endothermic, so the temperature at equilibrium will be colder (and equal to the freezing point of the salt solution). As ice melts, the temperature decreases and the freezing point increases (because the pure melt-water dilutes the salt solution). Once the two temperatures match, the system is at equilibrium (thermal equilibrium and phase equilibrium).


Old sample of sodium tetraphenylborate—why it is red? I have a really old sample of sodium tetraphenylborate. It should be white, but it is red. Does anyone know what was formed and how it influences precipitation of cations in water?

According to Decomposition of Sodium Tetraphenylborate, which relates to the Savanna River site for nuclear weapons material production,

  Stable aqueous NaTPB solutions are a transparent,
  flesh-toned color. During the initial stages of decomposition (or
  instability),the solution turns a golden color. As the quantity
  of TPB- that has decomposed increases, the solution turns dark
  red, then brown, and eventually black.

In the presence of oxygen "phenol and phenylboronic acid
are produced as well as benzene, biphenyl, and other benzene ring
derivatives".


Which is the least stable trihalide of nitrogen? Which is the least stable trihalide of nitrogen and why?
I have two conflicting theories, one is that fluorine-fluorine electron repulsion will make $\ce{NF3}$ least stable and another is steric reasons of iodine which give $\ce{NI3}$ as the answer.

Which is correct?

Nitrogen triiodide $\ce{NI3}$ is an extremely sensitive explosive that explodes with the slightest touch when dry.

Nitrogen trifluoride $\ce{NF3}$ is so stable that it is a greenhouse gas global warming concern, with lifetime in the atmosphere of hundreds of years.

$\ce{NI3}$ is unstable due to steric strain as you are saying.


Proton transfer equilibrium in bisulphite adducts In the addition of $\ce{NaHSO3}$ to aldehydes and methyl ketones (higher ketones do not respond well to this reaction), crystalline addition products are formed. An $\ce{-SO3H}$ and an $\ce{-OH}$ group are attached to the same carbon.

$$\ce{R2C=O + NaHSO3 -> R2C(OH)(SO3^{-}Na^+) <=> R2C(O^{-}Na^+)(SO3H)}$$

The line I read online, bothering me is:-

  "The proton transfer equilibrium lies to the right for most ketones
  and to the left for most aldehydes."

What is the reason for this (assuming that it is correct in the first place)? Since the sulphonic acid group is more acidic than the hydroxyl group, should not the proton tranfer equilibrium lie towards the left for both the cases?

I think your confusion arose because of an error in the translation of the text you linked (the discussion on p. 359) and that this text was unclear about which equilibrium it was referring to. In that figure, the equilibrium that the authors are presumably referring to is the (planar) carbonyl- (tetrahedral) hemiacetal equilibrium (as described on page 358). So the equilibrium that the text refers to isn't the "proton transfer equilibrium," but the planar-tetrahedral equilibrium.

In contrast, the scheme you posted with your question assumed that the formation of a hemiacetal (or hemiketal) is irreversible, as indicated by a unidirectional arrow. This reaction should be treated with an bidirectional equilibrium arrow as shown in p. 359 of the document that you linked.

In this context, the steric argument is reasonable and the reaction as a whole analogous to the more well-known relative hydration equilibria of aldehydes and ketones (which is related the general notion of aldehydes being more electrophilic than ketones).


Does an atom want to give or take electrons when it has a shell half full? For example, carbon has 4 electrons in the outer shell and so its half full. Does it want to give electrons in a bond or take electrons, and if any atom has a half full electron shell does it want to take or give electrons for a covalent bond?
I'm new to chemistry so this question may be simple.

When it comes to elements like carbon and silicon, which have 4 valence shells, the maximum possibility is sharing of electrons to form covalent bonds, like in $\ce{CO2}$ or $\ce{CH4}$. However in carbides, like calcium carbide or tungsten carbide, it may gain electrons.


Condition For NaHCO3 to react to form CO2 with organic acids My professor told us the following:

  $\ce{NaHCO3}$ will react to evolve
  $\ce{CO2}$ whenever the acidic nature of the reactant exceeds the
  acidic nature of benzoic acid.

I want to know why this is. And is the given statement valid? He didn't give us the reason.  And in that case should $\ce{NaHCO3}$ react with 1,2-dihydroxy
cyclobutanedione to evolve carbon dioxide?

$\ce{NaHCO3}$ is the basic sodium salt of carbonic acid, a diprotic acid which exists only in small amounts in aqueous solution, because the equilibrium

$$\ce{CO2 + H2O <=> H2CO3}$$

lies mainly on the educt side (this is also the reason why $\ce{CO2}$ evolution is observed). When this fact is taken into account, a $\ce{pK_{a}}$ value of 6.3 at 25 °C is obtained for the first protolysis step (source). Benzoic acid has a lower $\ce{pK_{a}}$ value of 4.19 and is therefore able to liberate carbonic acid from its salts, like any other acid that is significantly stronger than $\ce{H2CO3}$.

$$\ce{C6H5CO2H + NaHCO3 -> C6H5CO2Na + CO2 + H2O}$$

Just in case that you mean 3,4-Dihydroxycyclobut-3-ene-1,2-dione, which is also known as squaric acid, then a reaction with $\ce{CO2}$ evolution is likely. Squaric acid is a fairly strong diprotic acid, with $\ce{pK_{a}}$ values of 1.5 and 3.4, respectively. The acidity can be attributed to the stabilization of the dianion by resonance:

If carbon dioxide evolution with 1,2-dihydroxy cyclobutanedione is possible depends on the $\ce{pK_{a}}$ value of this compound. However, I would expect it to be less acidic than squaric acid.


How to increase the concentration of HCl Today I have purchased some $\ce{HCl(aq)}$ on my local market with a concentration of 23%.
However, I need to obtain a concentration of 37% in order to perform an experiment. Can I use distillation process to increase the acid concentration?
If yes, can I use standard chemistry glassware to process it? (In an isolated system to prevent the $\ce{HCl(g)}$ from escaping of course.)

Distillation unit

Add Calcium Chloride to destroy the azeotrope.  Then distill and capture the vapor phase.

See the second method "Extractive Distillation" here:

http://www.jacobs.com/uploadedFiles/wwwjacobscom/20_Learn_About_Us/25_Products/252_Chemetics/Hydrochloric%20Acid%20Concentration.pdf

(Obviously, proper equipment and safety precautions are essential.)


What is the state of aggregation of iodine formed during electrolysis of molten nickel(II) iodide? Assume the electrolysis of molten $\ce{NiI2(l)}$ (inert electrodes) in a cell. The cell with a heating element is in the room; the room is at $\pu{25 °C}$ and $\pu{101.3 kPa}.$

The half-reaction for oxidation at the anode would be

$$\ce{2 I^-(l) -> I2(\text{state}) + 2 e^-},$$

whereas the global reaction will be

$$\ce{NiI2(l) -> Ni(s) + I2(\text{state})}.$$

Given that iodine is solid at room temperature, should it be marked as a solid in the equations? However, given that the $\ce{NiI2}$ is liquid, I'd think that the newly synthesized iodine would be a gas for a while?

Also, what observations could one make during the electrolysis? I.e. will the iodine accumulate on the positive electrode as a solid, or will it exit the cell in the form of a gas?

Also, is the reduction potential for

$$\ce{I2(s) + 2 e^- -> 2 I^-(\color{red}{aq})}$$

the same as for

$$\ce{I2(s) + 2 e^- -> 2 I^-(\color{red}{l}) }?$$

Basically, during the electrolysis process of molten nickel(II) iodide ($\ce{NiI2}$), $\ce{I– (l)}$ ions would be oxidized to $\ce{I2 (g)}$ at the anode, and $\ce{Ni^2+ (l)}$ ions would be reduced to $\ce{Ni (s)}$ at the cathode. My state assignments were based on following facts assuming the temperature of electrolytic cell is kept below $\pu{900 ^\circ C}$:

Melting point of $\ce{NiI2}$ is about $\pu{797 ^\circ C}$. Therefore, electrolytic cell must be kept above $\pu{800 ^\circ C}$. Thus following equlibrium will be maintained during electrolysis: $$\ce{NiI2 (l) <=> Ni^2+ (l) + 2I- (l)}$$
Boiling point of $\ce{I2}$ is about $\pu{134 ^\circ C}$. Hence, $\ce{I2}$ would be released as a gas at anode.
Melting point of $\ce{Ni^\circ}$ metal is about $\pu{1455 ^\circ C}$. Hence, $\ce{Ni}$ would be deposited as a solid at cathode. If you want to keep nickel metal in liquid form, your electrolytic cell must be kept above $\pu{1500 ^\circ C}$.

I think I have given enough information you to figure out the rest.


Why do branched chain compounds have lower boiling points than the corresponding straight chain isomers? The branched chain compounds have lower boiling points than the corresponding straight chain isomers. For example,

  $\ce{CH_3CH_2CH_2CH_2CH_3}$ -
  No branching-Pentane (n-pentane) ($\mathrm{b.p.}=309~\mathrm{K}$)
  $\ce{CH_3CH(CH_3)CH_2CH_3}$ -
  One branching-2-Methylbutane (Iso-pentane) ($\mathrm{b.p.}=301~\mathrm{K}$)
  $\ce{C(CH_3)_4}$ - Two branches-2,2-Dimethylpropane (Neo-pentane) ($\mathrm{b.p.}=282.5\mathrm{K~}$)

  This is due to the fact that branching of the chain makes the molecule more compact and thereby decreases the surface area. Therefore, the intermolecular attractive forces which depend upon the surface area, also become small in magnitude on account of branching. Consequently, the boiling points of the branched chain alkanes are less than the straight chain isomers.

The above extract from my book, mentions clearly that branching makes the molecule more compact and thereby decreases the surface area. Even for isomeric alkyl halides, the boiling points decrease with branching. The reason is said to be because of the decrease in surface area, same as explained before. How does decrease in the surface area make the intermolecular forces small in magnitude?

The simple explanation is that weak intermolecular forces (the forces that make something condense to a liquid when things are cold enough) depend on the surface area (as well as many other things). But in the case of relatively similar non-polar isomers (where weak intermolecular forces are the dominant forces), the larger surface area will lead to the large force and hence the highest boilling point.

2,2 dimethyl propane is a compact, almost spherical molecule. Pentane is long and "floppier" so will experience more forces between molecules.

The underlying, but slightly simplified, explanation of this is that intermolecular forces (often called van der Waal's forces) depend on attractions caused by quantum fluctuations in the surface electrons of the molecule. These lead to short-lived dipole moments that can also induce dipoles in neighboring molecules to which the original dipole is attracted. The larger the surface area, the more opportunity for such dipoles to exist and therefore a stronger force.


How does impurity of distilled water vary with time in air? If I place distilled water in a glass beaker, how long is it expected to stay distilled?
What are the causes of it's impurity through air or other sources?

Gases like oxygen and nitrogen from air are absorbed quite easily. As you pour the water to the beaker it is most likely already saturated with those gases. Other air components and contaminants (carbon dioxide, sulfur dioxide, nitrogen oxides, ammonia) will also get quickly absorbed. How fast they will contaminate your sample depends on their concentration in the air, temperature of the water, size of the container and so on. Other possible contaminants are dust particles, impurities present inside the beaker, and even glass itself (depending of its type) is slightly soluble in water. The concentration of solid impurities will usually not change very quickly. Of course, as the water evaporates, solid impurities will concentrate.


Intermolecular forces between carbon dioxide and water I'm wondering which intermolecular forces will exist between carbon dioxide and water. Instinctively, I would expect there to be only dispersion forces between the 2 molecules, because carbon dioxide is non-polar due to its linear shape. However, even though carbon dioxide is non-polar, it still contains dipoles. I'm curious if its possible that these dipoles in carbon dioxide will allow it to form dipole-dipole forces (or even H-bonds) with water. Could the partially negative oxygen atomspotentially accept a H-bond from water, even if carbon dioxide is non-polar?

According to the this PubChem data table for the physical and chemical properties of carbon dioxide, it contains two hydrogen bond sites (each of the oxygens). This largely accounts for the high solubility of $\ce{CO2}$ compared to that of non-hydrogen bonding compounds like methane, which has about $\mathrm{1/25^{th}}$ of the Henry's law solubility in water as does $\ce{CO2}$.


Why is this intermediate borane-hydroperoxide adduct unstable? Boron has an empty p-orbital and is highly electron deficient because of an incomplete octet.

So shouldn't this intermediate be more stable than $\ce{BR3}$ (whose electron deficiency is only somewhat fulfilled through hyperconjugation) since boron's electron deficiency is satisfied by the increased negative charge (although the majority of negative charge is still on oxygen due to its high electronegativity).

Basically an intermediate form, that fulfills the octet rule, would be more stable than the trialkylborane, but only if we suppose that there are no other effects! And if there were no other ffects, it wouldn't be an intermediate form but an isolable compound.
In your reaction scheme you see that $\ce{OH-}$ will leave. You maybe know that $\ce{OH-}$ is - under basic conditions - a good leaving group. Moreover a peroxide is mostly pretty unstable, so it's happy to react and let $\ce{OH-}$ leave. The hydroboration–oxidation reaction gives you the tetrahydroxyborane-anion ($\ce{B(OH)4-}$) and three alcohols ($\ce{ROH}$) as final product. So you'll get an anion that fulfills the octet rule and is more stable than your reactant.

Edit: More details on the hydroxide as leaving group
The reaction finds place under strongly basic conditions. Usually you take $\ce{NaOH}$ to deprotonate the $\ce{H2O2}$ and form $\ce{OOH-}$, so in your solution there are plenty of $\ce{OH-}$ ions. The intermediate also splits off an hydroxide ion which is a good leaving group under this conditions. This means that the ion is stabilized in solution and is not less stable as the intermediate. Even if you can't really compare the stability of those two things because they are not the same... So let's only say, that the hydroxide ion is pretty stable in a basic solution.
As I also said in my original answer above, peroxides are pretty unstable. This is because the two oxygens are right next to each other. As you may know, oxygen is one of the most electronegative elements of the periodic table, so it "fights to get electrons". If you connect two oxygens like in a peroxide, they both fight to get the electron of the other one, so it resembles rope pulling. This isn't the best case scenario for a stable bond.
If you look at hydrogen peroxide, especially in concentrated solutions, it tends to decompose without any external influence.
$\ce{2H2O2 -> 2H2O + O2}$
This reaction releases 98,2 kJ/mol [1] of energy. This means that the products (oxygen and water) are way more stable than the hydrogen peroxide itself. This is mainly due to the very unstable rope pulling-bond between the two oxygens in $\ce{H2O2}$: $\ce{RO-OR}$.
This should explain why the intermediate group is so unstable.


Relation between ΔH and ΔE for the reaction in gaseous phase at constant temperature and pressure Which of the following statements is correct for the reaction at constant temperature and pressure:
$$\ce{CO(g) + \frac{1}{2}O_2(g) -> CO2(g)}$$

$\Delta H = \Delta E$
$\Delta H > \Delta E$
$\Delta H < \Delta E$
None of these.

Since this occurs at constant temperature, $\Delta E = 0$ and $W < 0$. So, $\Delta H$ turns out to be less than $\Delta E$, but this is not the answer. Why?

By definition, $$\Delta H=\Delta E+\Delta(PV)$$For a reacting ideal gas mixture at constant temperature, $$\Delta(PV)=(\Delta n)RT$$where $\Delta n$ is the change in the number of moles between reactants and products.  Therefore, $$\Delta H=\Delta E+(\Delta n)RT$$For the reaction under consideration, $$\Delta n=-\frac{1}{2}$$


Why does potassium have the electronic configuration of 2.8.8.1? We all know that potassium, the 19th element on the periodic table has the electronic configuration of 2.8.8.1.  However, why not 2.8.9?  In the element scandium, it has the electronic configuration of 2.8.9.2.  I would like to know why potassium is not 2.8.9.

You could actually have the 2.8.9 configuration, putting the last electron in a 3d rather than 4s orbital.  But the 2.8.8.1 one is lower in energy and thus more stable.  How electrons fill subshells in neutral atoms to make the lowest energy configuration can be seen here.


Which rings of this heterocyclic compound are aromatic? I am wondering which of the inner two rings of this compound are aromatic?

I know the rules for aromaticity. But this case is difficult for me.

Regarding the smaller ring, I guess the sulfur contributes one lone pair of its valence electrons into the system. This means from the nitrogen (in the 5-ring) only one electron is needed for the π-system (with 6 electrons). But this nitrogen atom has only one lone electron pair, thus one of the electrons would not be part of the  π-system of the smaller ring. However, this nitrogen atom in the 5 membered ring is also part of the 6-ring and there it could share both of its free electrons to this larger ring. Then the second nitrogen which is only in the 6-membered ring could contribute one electron only to the aromatic system, and thus having a happy free lone pair left.

What do you think of that? Which rules, logic or reasoning can confirm this? And isn't it a problem that the first nitrogen shares one electron with the 5-ring and the two with the 6-ring?

Also, is there software which can predict this reliably? I haven't found anything so far.

I think that all the four rings are aromatic.

If the top nitrogen donates it's pair of electrons to the ring, and oxygen takes it's pair, the three fused rings would have 14 electrons in the rings' peripheral, causing it to be aromatic:

Even the 3D drawings of this molecule confirm it's planarity of the 3 fused rings:

I think we can at least suspect all the rings in the molecule to be aromatic.


How to test for Chromium(III) Oxide I got some chromium(III) oxide ($\ce{Cr2O3}$) sold as pigment, called "Chrome Green". I've been having a hard time trying to react it with $\ce{HCl}$ to make $\ce{CrCl3}$. I thought it was just hard to break the $\ce{Cr-O}$ bond.

Now I'm just afraid that the pigment that I bought was actually cut with something else, like a silicate-type of pigment. How do I test for chrome(III) oxide, if I can't get it to become a salt in the first place?

The test for $\ce{Cr(III)}$ can be performed in the solid state using an oxidative melt. You will need

chromium(III) oxide (obviously)
sodium carbonate
potassium nitrate
a Bunsen or Teclu burner
a magnesia trough or a crucible
suitable pliers to hold the trough

Mix the green pigment with up to the threefold amount of sodium carbonate and potassium nitrate, place the mixture on the magnesia trough and melt it.

A yellow melt indicates the formation of sodium chromate according to

$\ce{Cr2O3 + 3KNO3 + 2Na2CO3 -> 2Na2CrO4 + 3KNO2 + 2CO2 ^}$


Oxidation - Reduction Ionic Half Reaction Trying to figure out how to balance the following molecular equation:

$$\ce{KMnO4 + Na2C2O4 + H2SO4 -> MnSO4 + CO2 + Na2SO4 + H2O}$$

I have compiled all of their oxidation states

$$\ce{KMnO4 + Na2C2O4 + H2SO4 -> MnSO4 + CO2 + Na2SO4 + H2O}$$

$$\ce{(+1)(+7)(-2 * 4) (+1 * 2)(+3 * 2)(-2 * 4) -> (+2)(-6)(-2 * 4) + (+4)(-2 * 2) + (+1 * 2)(+6)(-2 * 4) + (+1 * 2)(-2)}$$

Now I know that theoretically Mn changes from +7 to +2, so it gains 5 electrons. I also know Carbon in this situation goes from a state of +3 to +4, so it loses one electron.

However, after this I am drawing blanks on how to setup my ionic half reactions with these polyatomic ions, and I am suspicious that I may have set this up incorrectly. Any help at all would be appreciated, I can not find much to explain this type of equation with this particular molecular formula.

The idea of Redox reactions make sense when talking about individual atoms or more clearly depicted changes, but I am having trouble figuring out what is even changing into what in this example.

Credits to my awesome teacher for this method.
(I had to add a potassium ion to the right hand side, there was no other way)


Why are the two bonds in sulphur dioxide identical? Triatomic molecules are either linear or bent. If we analyse the $\ce{SO2}$ molecule,  it turns out that it is bent.
I know that $\pi$ bonds do not alter shape, but merely the bond lengths.
Now since one of the bonds between S and O is a p$\pi$-p$\pi$ bond and the other bond is a d$\pi$-p$\pi$ bond, the two $\pi$ bonds are not same, so the bond length should not be the same, but my book says the bond lengths are the same which is quite puzzling. I thought maybe its due to resonance but there's no way resonance takes place here.  I read this question,but it don't answer many questions as:
Why d-orbitals of sulphur cannot be used  ?
And  in answer by @Kenny Lau ,
what's specify reason for 'Sulfur sharing two electrons among itself and the other two oxygen atoms.'
So, that makes me think then, what could  be the reason for same bond length in $\ce{SO2}$?
It is not clear as to why d-orbitals of sulphur cannot be used and in the answer to this question @Kenny Lau does not specify reason for 'Sulfur sharing two electrons among itself and the other two oxygen atoms.'

There are two ways to look at the $\ce{SO2}$ molecule. The valence bond way is this:

$\ce{SO2}$ is a bent molecule with C2v symmetry point
group. A valence bond theory approach considering just s and p
orbitals describes the bonding in terms of resonance between two
resonance structures:

The sulfur–oxygen bond has a bond order of 1.5 and does not invoke d
orbital participation. In terms of electron-counting formalism, the
sulfur atom has an oxidation state of +4 and a formal charge of +1.
(Wikipedia)

However, there is a molecular orbital approach which seems to give a deeper appreciation of the bonding. The picture is this:

(Picture source with detail explanation)
This is my simple explanation:

The sulfur is hybridized sp$^2$-p. The bond angle is 119$^o$ - you
can hardly get closer to 120$^o$!

We will assume that the oxygens are also hybridized sp$^2$-p.

Then the sigma bonding is designated in the diagram by the numbers 1 and 2 for the sigma bonds between sulfur and oxygen. Lone pairs are designated by numbers 3 (for S), and 4, 5, 6 and 7 (for O).
The pi system has three molecular orbitals, one from each atom and four electrons: one from each oxygen, two from sulfur. The lobes in these orbitals are drawn at the lower right side of the picture. Orbital 8 has the lowest energy, with all lobes bonding. The non-bonding pi orbital (#8) has reversed signs for the lobes of oxygens and therefore a node at sulfur. The antibonding pi orbital (#10) has two nodes and is unoccupied.
So the final count is two filled sigma bonds and one filled pi bond; five lone pairs and one filled non-bonding pi orbital. The formal bond order is 1.5, same as in the valence bond explanation with resonance.
The only thing I might change in the diagram is the relative energies of the sulfur orbitals and the oxygen orbitals: I would think they might be closed to the same energy. The diagram might be drawn as it is in order to make it clearer and give more room for all the numbers.
We just don't need d orbitals to explain anything, but their involvement cannot be entirely eliminated.


How to mix small amount of alcohol into coconut oil I want to mix some alcohol (say rum) into a coconut oil base.
At the minute when I try this the alcohol seperates and sits on the bottom as a liquid. Is there any way I can mix these together so the alcohol remains suspended into the coconut oil.

When you consider that rum and most other hard liquors contain at least 50% water, it is not surprising that the result is a two-phase system. If you began with absolute or 95% ethanol, or perhaps even grain neutral spirits, which vary but often come as 180 proof (90% ethanol) you might get the homogeneous solution you want. Another approach would be to add an edible emulsifier such as lecithin, which will help the ingredients to form a stable emulsion.

The earlier suggestion to hydrogenate the oil is not likely to work, as hydrogenation does not make oil able to mix with water. Coconut oil already has a high saturated fat content, as can be seen by the fact that it is often a solid at room temperature.