7

The problem is, we want to see a trend in everything. There are various factors that govern melting point and boiling point which is the reason of perceived anomalies. "The equation which has more number of variables, is harder to solve." For melting point, few of them are: Crystal system, size of atom, atom-atom distance, distance between two ...


7

According to Wikipedia, ethanol flash point is $\pu{55 ^\circ F} = \pu{12 ^\circ C}$. Auto-ignition temperature is $\pu{793 ^\circ F} = \pu{423 ^\circ C}$. It means that ethanol vapor cannot be put on fire with a match at a temperature under $\pu{12 ^\circ C}$. But its vapor will not spontaneously ignite at temperature lower than $\pu{423 ^\circ C}$.


7

Chemistry is very complicated and I'm probably missing certain edge cases, but I think the answer is almost certainly 'no'. Salts lower the vapour pressure of water because intermolecular forces attract the salt ions to the water, raising the energy needed for water molecules to escape. In order for water's vapour pressure to rise, the salt would need to (on ...


6

Liquid evaporation occurs at any temperature, just its rate changes with temperature. You must have seen washed clothes getting dry in open air, water evaporating at room temperature. OTOH, I guess you have not observed boiling sea, lakes or rivers to obtain clouds - putting aside geothermal activities. There are few things for you to understand: Molecules ...


4

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 ...


3

If you heat the commercial concentrated ammonia solution ($25$% $\ce{NH3}$) at usual pressure, it will boil at $32$°C and the vapor contains $3$% $\ce{H2O}$ (and of course $97$% $\ce{NH3}$). So the liquid looses much ammonia and nearly no water. Its total volume decreases a bit but the concentration of ammonia decreases more, so that it is necessary to heat ...


3

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 ...


3

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- ...


3

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 ....


2

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 $\pu{104°C}$.


2

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 ...


2

[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 ...


2

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 ...


2

But my source of confusion is that ambient pressure is not the only pressure pressing down on the liquid. Pressure does not press down. When something is under pressure, it exerts it in all directions. Maybe it would help imagining this problem in the absence of gravity (you would have to put the sample in a stretchy balloon to get some pressure while being ...


1

Benzene-1,4-diol has a boiling point of 287°C and Benzene-1,3-diol has a boiling point of 277°C and Benzene-1,2-diol has a boiling point of 245.5°C. This could be attributed to the ease of formation of 2 intermolecular hydrogen bonds with the 2 hydroxy groups of these molecules increasing when the hydroxy groups are placed apart, to minimise steric ...


1

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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