Hello Marcin,

This is an important question. In C2H5-OH, the C2H5 moiety is rather hydrophobic , just think of the solubility of CH3-CH3, essentially nil. The OH group is "polar", that is, it is at the origin of the dipole moment of the molecule. The most important thing, however, is the abilty of the OH group to act as both a hydrogen bond donor and a hydrogen bond acceptor with bulk water. Overall, these interactions overcome the hydrophobicity of the C2H5 group.

As the length of the alcohol chain increases, the hydrophobic effect increases.

I am gving you a small table taken from "solubility of things" on the net.

Alcohol solubility chart

The amounts are in mol/100g of H2O at 1atm and 25oC

Name Formula Solubility

Methanol CH3OH miscible

Ethanol C2H5OH miscible

Propanol C3H7OH miscible

Butanol C4H9OH 0.11

Pentanol C5H11OH 0.030

Hexanol C6H13OH 0.0058

Heptanol C7H15OH 0.0008

Best wishes,

JL

Thanks for explanation. To sum up: one cannot say that compound is hydrophilic based only on miscible solubility. Am I right?

Franks (in book mentioned above http://books.google.pl/books?id=k4bxN1vtgigC&pg=PA64&dq=figure+6.3+excess+partial+molar+volumes,+at&hl=pl&sa=X&ei=PmtrUuHnKejZ4QTdpoCYCQ&ved=0CDIQ6AEwAA#v=onepage&q=figure%206.3%20excess%20partial%20molar%20volumes%2C%20at&f=false) says that checking sign of derivative of excess partial molar volume of ethanol gives the answer if it is hydrophobic or hydrophilic. Negative sign of dV2e/dx2 means it is hydrophobic.

Is it general rule?

The reason of my disccusion is fact that popular sources describe ethanol as hydrophilic which as far as I understood is not true and its characteristics is more complex.

Yes indeed,the situation is complex. I suggest you have a look at the following paper:

Journal home > Archive > Letters to Nature > Abstract

Letters to Nature

Nature 416, 829-832 (25 April 2002) | doi:10.1038/416829a; Received 26 October 2001; Accepted 4 March 2002

Molecular segregation observed in a concentrated alcohol–water solution

S. Dixit1, J. Crain1, W. C. K. Poon1, J. L. Finney2 & A. K. Soper3

1. Department of Physics and Astronomy, The University of Edinburgh, Mayfield Road, Edinburgh EH9 3JZ, UK

2. Department of Physics and Astronomy, University College London, Gower Street, London WC1E 6BE, UK

3. ISIS Facility, Rutherford Appleton Laboratory, Chilton, Didcot, Oxon OX11 OQX, UK

Correspondence to: A. K. Soper3 Correspondence and requests for materials should be addressed to A.K.S. (e-mail: Email: [email protected]).

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When a simple alcohol such as methanol or ethanol is mixed with water1,2, the entropy of the system increases far less than expected for an ideal solution of randomly mixed molecules3. This well-known effect has been attributed to hydrophobic headgroups creating ice-like or clathrate-like structures in the surrounding water4, although experimental support for this hypothesis is scarce5, 6, 7. In fact, an increasing amount of experimental and theoretical work suggests that the hydrophobic headgroups of alcohol molecules in aqueous solution cluster together2, 8,9, 10. However, a consistent description of the details of this self-association is lacking11, 12, 13. Here we use neutron diffraction with isotope substitution to probe the molecular-scale structure of a concentrated alcohol–water mixture (7:3 molar ratio). Our data indicate that most of the water molecules exist as small hydrogen-bonded strings and clusters in a 'fluid' of close-packed methyl groups, with water clusters bridging neighbouring methanol hydroxyl groups throughhydrogen bonding. This behaviour suggests that the anomalous thermodynamics of water–alcohol systems arises from incomplete mixing at the molecular level and from retention of remnants of the three-dimensional hydrogen-bonded network structure of bulk water.

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Personally, I think this sort of "experimental picture" is extremely important.

There are further points of interest:

1. The co-operative character of hydrogen bonding (HB). See, e.g., J. Org. Chem. 1982, 47, 4553-4557.

2. I don't like the "hydrophobicity". Jokingly, molecules or fragments therefrom don't suffer rabies. More seriously, this term hides something of tremendous importance, London's dispersion forces. For reasons "nearly unknown" (too much mathematics?), they are almost ignored in most Chemistry courses. Hydrogen bonding (•••) and "polarity" (+ and -) are easier to catch and are way more popular. You see, CH4 is a "nearly perfect gas". while n-C17H36 is a waxy stuff. This tremendous change mostly originates in the increasing dispersion effects. as n increases in CnH2n+2. So, in the Nature paper you see the contributions of both effects.

Yes, this world is complex.

Regards, JL.

It is not very readable. I am attaching a more reader-friendlt tex

dear mr. steac. Before looking on solubility values, i suggest to look inside solubilty. Thus, solvatation enthalpy is an adecuate parameter. solubity enthalpy messures the interaction between solute and solvent; solute-solute and solvent-solvent. So, DH = (interaction (soluto-Soluto)+interaction(soluto-solvent)+interaction(solvent-solvent). If interaction soluto-solvent its bigger thant other to parameteres, then, the solute its soluble. So, for C6-Cn alcohol, the london forces are bigger than Van der Waals interaction (OH-H2O) then, solubility down.

I agree with Dr. Martinez.

JL

Sorry, I agree with Dr. Marquez.

JL

Thanks for help and valuable hints. Now I'll start familarizing with literature.

I am agree with doctor Marquez.

Nice summary of the discussion. And next proof that "A picture is worth a thousand words".