11

It turns out that the opposite of what happens in polar solvents takes place when a non-polar solvent is used. At the same temperature, o-nitrophenol is more soluble in benzene than it's m and p isomers. Sidgwick et al.1 did a study of this and obtained the following results. (Note that the solvent they used was toluene and not benzene, but they are similar ...


10

The boiling point of non-ionic compounds are highly depend on their H-bonding abilities. For example, boiling point of water (molar mass: $\pu{18.02 g/mol}$) is $\pu{100 ^\circ C}$ at $\pu{1 atm}$ while that of ethanol (molar mass: $\pu{46.07 g/mol}$) is $\pu{78.4 ^\circ C}$ at $\pu{1 atm}$, even though ethanol is heavier and have more other intermolecular ...


9

Mathew Mahindaratne has provided analysis based on experimental values of the boiling points of the two compounds. I would like to offer a different view using bonding analysis. Before I begin tackling the question, we shall first clarify the concept of the hydrogen bond. While the popular view of the hydrogen bond is as a particularly strong type of "...


7

The nitrophenols have completely different physical behavior based on the position of nitro group: $$ \begin{array}{c|ccc} \hline \text{Compound} & \text{Melting point} & \text{Boiling point} & \text{Water solubility at } \pu{25 ^\circ C}\\ \hline \text{2-Nitrophenol} & \pu{43-45 ^\circ C} & \pu{215 ^\circ C} & \pu{2 g/L} \\ \text{...


4

I would like to object the statement dicarboxylic acids would not engage hydrogen bonding. Terephthalic acids ($\ce{HO2C-C6H4-CO2H}$) are a prominent example since early attempts to formalize the description of intra- and intermolecular hydrogen bonding in the solid state. Consider the pairing of benzoïc acid and $p$-terephthalic acid, sharing the same ...


3

Hydration of acetone is an equilibrium process. $$\ce{(CH3)2C^16O + H2^18O <=> (CH3)2C^18O + H2^16O }$$ The equilibrium constant of this reaction will depend on how the bond strengths of the $\ce{O-H}$ bond in water and the $\ce{O=C}$ bond in acetone is affected by isotope exchange. The isotope-dependent strength of hydrogen bonds might be a secondary ...


3

The boiling points of non-polar hydrocarbons are determined by the extent of van der Waals forces in between them. More specifically, molecules with a larger surface area have larger van der Waals forces of attraction between the molecules. From the structures, it is obvious that but-1-yne has a larger surface area than isobutane, and hence has the higher ...


3

Whenever a non-polar substance is put into water, water molecules will organise themselves around it in a cage-like formation. The molecules exposed to the non-polar substance at any given time will orient themselves to form as many hydrogen bonds with the rest of the solution as possible. (Image from here) The way I understand it intuitively is that these ...


2

From the abstract of an article titled "The crystal structure of galactaric acid (mucic acid) at −147°: An unusually dense, hydrogen-bonded structure" [...] D$_m$ = [1.790] $\pu{g cm−3}$ [...] The crystal structure has a system of strong, intermolecular hydrogen-bonds, which accounts for the high crystal density and low solubility in water. The ...


2

Perhaps the words in the picture are explained at greater length in the text, or perhaps actual numbers are given for the dissociation constants (https://www.quora.com/Why-is-o-flurophenol-is-more-acidic-than-p-flurophenol). As it turns out, the dissociation constants are 8.7, 9.3. 9.9 and 10.0 for ortho, meta, para and unsubstituted phenol. Ortho-...


1

All the compounds that have a relation $\alpha-\beta$ between the carbonyl groups. The hydrogen between two $\ce{CO}$ is very acid because two electrowithdrawing remove electron density from the $\ce{C-O}$ bond. Therefore for this kind of compound the most stable for is the enol. In Organic Chemistry this is very useful in creating new carbon-carbon bonds in ...


1

When Point mutations occur in one strand of DNA, irregular base-pair emerge. One of the main features of this irregular state is that the width of the double-strand differs from the standard form(usually larger width as the new, irregular base pairings are weak) in that specific area. In other words, the portion of DNA which has the defective base-pair ...


1

These exponents $10$ and $12$ are a little bit arbitrary. They have been chosen so as to obtain an energy curve corresponding to the measured H-bond energies. Other attempts have been published, like the famous Morse curve, which is, for covalent bonds versus HO distance $d$ : $$E(d) = D[1 - e^{-k(d - d_o)}]^2$$ where $D$ is the dissociation energy of the ...


1

There is the general trend in the groups 15, 16, 17 of raising of boiling points for the binary compounds with hydrogen, going down the groups. But the first members of each group - $\ce{NH3, H2O, HF}$ - have anomally with their boiling points being exceptionally high, due hydrogen bonds. The strength of these bonds and the boiling point decreases in the ...


1

"In a way yes". However, rather than starting from nomenclature to get facts, do the vice versa. Otherwise and again "in a way", they could be seen as permanent dipole-dipole interactions, too. Hydrogen bonds have energy of about a tenth as compared to covalent bonds. So they are considerably weaker but still stronger than other ...


1

You don't need a lone pair. Water can hydrogen bond to methane, attacking the back side of one of the carbon-hydrogen bonds in the methane. In a molecular orbital the sigma bonding orbitals in the methane are sufficiently delocalized to serve in place of lone pairs, overlapping the antibonding orbitals in the water [1]: Abstract (from [1]) Quantum chemical ...


1

I would say that your analysis of first, second and fourth one is correct, as H-bonds are stronger than dipole-dipole interactions. However, you made a mistake while filtering compounds having dipole-dipole interaction. The $1$ and $3$, both have dipoles! The $\ce{C=O}$ bonds in $\ce{CO2}$ are polar and hence they possess dipoles, but unlike a water molecule,...


1

In fact, yes. It can be explained by a HOMO-LUMO interaction. I am currently running NBO (natural bond orbital) computations for a theoretical study of hydrogen bonds in some models. And NBO interprets hydrogen bond as the donation of electrons from the lone pair (non-bonding electrons) of the hydrogen acceptor to the (LUMO) antibonding orbital of the ...


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