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In non-polar liquids, the intermolecular forces are known as London dispersion forces (or dipole induced dipole interaction). This is the weakest interaction that exists between molecules. The strength of this interaction depends on polarisability of non-polar molecule. As an example, let's take the following examples: $$ \begin{array}{lr|lr|lr|lr} \hline \...


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Latent heat of vaporization/evaporation: It's the amount of heat required for liquid ---> gas phase change of a substance at it's Boiling point. Not exactly. It's the amount of heat required for liquid ---> gas phase change of a substance at the particular temperature, there is usually and implicitly the substance boiling point. But generally, it is ...


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A supercritical fluid is just a state of matter. like solid, liquid or gas. Hexane is no more or less toxic if it has become supercritical and then brought back to standard temperature and pressure, as would water, if it changed from liquid to ice or to supercritical "steam" and back to water. Further, supercritical $\ce{CO2}$ is used for dry-...


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Firstly, sulfur has a larger atomic radius than carbon, so we would expect DMSO (Ddimethyl sulfoxide) to have a much larger and thus polarisable electron cloud than acetone. The London dispersion forces between DMSO are thus stronger, causing DMSO to have a higher boiling point. Secondly, if we compare the dipole moments of the $\ce{S=O}$ and $\ce{C=O}$ ...


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A more plausible explanation is that it is not water that is boiling. The "organic substances" and maybe the "nitrogenous substances" in the juice could include some more volatile components that may pass selectively into the gas phase at a lower temperature than most of the solution. Such a vapor would not be the pure organic or ...


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It does seem that placing the hydroxyl group at the terminal carbon maximises the hydrogen bonding interactions between alcohol molecules and also the van der Waal forces between them. This is best illustrated with a crude diagram I have made: Finer differences due to the position of the hydroxyl group along the chain, such as the differences between the 2- ...


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Distance In vacuum, the calculated (classical) interaction between to aligned dipoles decreases with the square of the distance. If the dipoles can't come very close to each other, the interaction will be weak. If you treat hydrogen bonds as a special case of dipole-dipole interaction, you will find that a N-H ... O=C interaction is much stronger than a C-H ....


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It is not a messed up cycle. Let's consider what happens with an example. If you heat up a mixture containing $10$% ethanol + $90$% water, the mixture boils at $92$°C, producing a vapor containing $50$% ethanol + $50$% water. With this operation the liquid loses more ethanol than water. So its concentration in ethanol decreases, and it is necessary to heat ...


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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 ...


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The boiling temperature always increases when adding salt into water. It looks as if salt was attracting water molecules and preventing them from quitting the liquid. So that it is necessary to overheat the solution to let water molecules get into the vapor phase. I have often made this measurements in class and always found about $\pu{104°C}$.


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I do use my comment above because the meaning of the books is, with little interpretation, clear and the answer is simple and don't require digging in theory of fractional distillation. So it would be pity to let the question unanswered. A mixture as those discussed boils over a range of T in the sense that its boiling point BP depend on its composition. The ...


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The main reason acting here is Dipole Moment. The higher the dipole moment of a molecule, the greater will be the intermolecular attractive forces and the higher will its boiling point be. Dipole moment is the displacement of electron density in a molecule and it is a vector quantity. The net dipole moment of a molecule is the vector summation of all the ...


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As Poutnik said, the melting point is well defined for pure substances. In the case of egg yolk, it is not a pure substance and also when you heat it, some molecules decompose. So we can't measure or even define a melting point for it. Also, the temperature is not constant, when the egg is melting. A substances' phase doesn't get influenced only by ...


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[OP] the average kinetic energy of evaporating water molecules You have to specify whether you are talking about the kinetic energy just before the water molecule breaks the hydrogen bonds to its neighbors or just afterwards. A millisecond before or after the event, of course, the average kinetic energy will be determined by the bulk temperature. One way to ...


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At gas/liquid phase equilibrium, average kinetic energy of evaporating molecules, i.e. those just passed to a gas phase, is equal to average kinetic energy of condensing molecules. The latter is then approximately proportional to $T$. If these average values were not equal, the system would not be in thermal equilibrium. Feedback to comments: @theorist In ...


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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 ...


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When a $\ce{H}$ atom is bound to an electronegative atom like oxygen, its only electron is mostly placed between the two nuclei. Outside of this $\ce{H}$ atom, the $\ce{H}$ nucleus is nearly "naked". Speaking naively, there is no other electron to occupy the space around the nucleus towards outside. It looks as if the nucleus $\ce{H}$ is not ...


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This is a very interesting question! Also, I don't believe it is a duplicate of the other cited questions, since those seem to concern the average kinetic energy of liquid vs. gaseous water, while the OP is asking whether those that evaporate from the liquid are at the tail of the distribution, and whether there's a relationship between the KE of this ...


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Trial and error, but most often the boiling point of the lower-boiling solvent. The absolutely correct answer would be to look at a phase diagram, determine whether you have a positive or negative azeotope (or non at all), find out where on the phase diagram you are and then derive the temperature. That would tell you approximately at what temperature your ...


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I suggest that an entropic term explains the little difference observed. While this aspect is normally important to justify what isomer melts at lower temperature, in principle it can be invoked to justify why an isomer having a longer - not too much - branch boils at a bit higher temperature. In this case, which might be rare as for a balancing between ...


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In addition to the poor polarizability of $\ce{C-F}$ bonds described here, we may consider hydrogen bonding. $\ce{CHCl3}$ has this to a greater extent than most chlorinated organic compounds because the electronegative $\ce{CCl3}$ function draws off the electron density in the bonding orbitals, leaving more of the unoccupied antibonding orbitals on hydrogen. ...


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The reason is that after arsenic, there is shielding. This causes the intermolecular forces to become weak. Thus, the melting point decreases. I hope you know about shielding of d and f electrons


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