The Real Reason Why Oil and Water Don't Mix

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DOI: 10.1021/ed075p116
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Abstract
The majority of introductory chemistry and organic chemistry textbooks state that oil and water don't mix because of enthalpic effects. These texts generally make the argument that the mixing process is endothermic, reasoning that the water-water hydrogen bonds that must be broken in order to accommodate the solute are much stronger than the subsequent solvent-solute dipole-induced dipole intermolecular forces that are formed. In fact, in most cases the mixing process is exothermic, so the immiscibility of the two liquids must be explained by a loss of entropy in the system. The widely accepted model explaining the hydrophobic effect invokes the formation of icelike clathrate hydrate "cages" around nonpolar solute molecules. Water molecules at the surface of these relatively rigid clathrate structures are strongly hydrogen-bonded to one another. The formation of these solvent "cages" explains why both Delta H and Delta S are negative for the solution process, and the endergonicity of solvation is thus due to entropy and not enthalpy. Authors should remove from their textbooks the incorrect enthalpic/hydrogen-bond explanation for the hydrophobic effect. Because aspects of the correct entropic/clathrate "cage" explanation lie beyond the scope of introductory or organic chemistry courses, it may be wisest to omit any detailed physical explanation of the "like dissolves like" phenomenon. If the overall format of the text permits, a brief discussion of solvation entropy effects might be included in the section dealing with the immiscibility of oil and water Keywords (Audience): High School / Introductory Chemistry
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116 Journal of Chemical Education • Vol. 75 No. 1 January 1998 • JChemEd.chem.wisc.edu
The Real Reason Why Oil and Water Don’t Mix
Todd P. Silverstein*
Chemistry Department, Willamette University, Salem, OR 97301
Background
Most introductory chemistry textbooks include in their
discussion of solubility and miscibility the famous rule of
thumb, “like dissolves like”. The converse of this rule, that
nonpolar solutes are insoluble in polar solvents, is often re-
ferred to as the hydrophobic effect. This effect forms the basis
for many important chemical phenomena: the cleaning
action of soaps and detergents, the influence of surfactants
on surface tension, the formation of biological membranes,
and the stabilization of protein structure are all based in
large part on the hydrophobicity of nonpolar groups. In their
explanation of the hydrophobic effect, introductory chemistry
textbooks often rely primarily on the concepts of enthalpy and
intermolecular forces. Because the solution process is gen-
erally discussed after enthalpy, but before entropy or free
energy, authors are left with little choice but to emphasize
enthalpy over entropy when explaining the thermodynamic
basis of “like dissolves like” and the hydrophobic effect.
For example, one text states that octane and carbon tet-
rachloride are miscible because the nonpolar molecules are
both “held together in the liquid phase by weak dispersion
forces” and thus “are attracted to one another”. However,
for the very reason that both molecules experience similar
intermolecular forces, H of mixing should be close to zero.
These molecules are thus not “attracted” to one another in the
normal enthalpic sense of the word, but rather in an en-
tropic sense. Spontaneous mixing of the two phases is
driven not by enthalpy, but by entropy.
Regarding the immiscibility of octane and water, an-
other text states that “the value of the overall enthalpy
change, Hsolution, is likely to be positive, reflecting an endo-
thermic process. In large part, it is for this reason that polar
and nonpolar liquids do not mix well.” In the legend of an
accompanying figure, this text offers the most common (in-
correct) explanation as to why oil and water don’t mix: “The
weak intermolecular interactions between the nonpolar
molecules and water cannot overcome the very strong forces
between water molecules and allow them to be miscible
with water.” Another version of this explanation appeared
in a recent overhead projector demonstration in this Journal,
entitled “Why Don’t Water and Oil Mix?” The authors stated
the argument this way: “The water molecules are attracted
to each other through hydrogen bonding…resulting in the
expulsion of the nonpolar oil molecules into a separate
layer”. They go on to describe an overhead projector dem-
onstration using magnetic stir bars to simulate the water–
water attraction and marbles to simulate nonpolar solutes.
Once again, the problem here is the stress these authors
place on enthalpy and hydrogen bonding, while omitting
any reference to the importance of entropic effects.
This issue has in fact surfaced in the pages of this
Journal in the past. In a brief note published in 1994, Alger
(2) pointed out that many organic chemistry textbooks offer
the same incorrect explanation for the immiscibility of water
and organic liquids. Huque (3), writing in this Journal in
1989, provided a detailed thermodynamic, statistical-
mechanical analysis of the hydrophobic effect. Somehow the
ideas stressed in these papers and many others have been
largely overlooked in introductory and organic chemistry
textbooks.1 In this paper I will describe the scope of the
problem, present thermodynamic data along with a generally
accepted model that explains the hydrophobic effect, and
recommend how textbook authors should approach the
problem.
Discussion
First consider the discussion of the immiscibility of oc-
tane and water quoted above. The choice of octane as solute
is fortuitous because dissolving octane in water is, as these
authors propose, slightly endothermic. However, this is not
the main reason why octane does not spontaneously dis-
solve in water. In fact, dissolving smaller hydrocarbons such
as ethane, propane, butane, and pentane in water is actu-
ally an exothermic process (see Table 1). Even for hexane,
benzene, toluene, and xylene, Hsolution is close to zero.
What then is the correct explanation for the hydrophobic
effect? Consider hexane and water: they are immiscible, yet
Hsolution is about zero (see Table 1). The reason for the im-
miscibility must therefore be entropic and not enthalpic.
Seidell’s thermodynamic data (4) showing the enthalpy of
solution for most organic liquids in water to be negative2
were already available in 1941. In the 1940s, Butler (5) and
Frank and Evans (6) were the first to point out the impor-
tance of increased order in the aqueous phase in explaining
the hydrophobic effect. Later, Klotz (7), Kauzmann (8),
Nemethy and Scheraga (9), and Tanford (10) developed
further Frank and Evans’s entropic model. Tanford in par-
ticular performed a series of experiments in which he mea-
sured changes in enthalpy, entropy, and free energy upon
transferring nonpolar solutes from organic to aqueous solvent
phases. The data collected in Table 1 clearly show that the
thermodynamic barrier to the solution process is entropic
rather than enthalpic.
How can we explain this? How is it that Hsolution is zero
for dissolving hexane in water (Table 1)? Is it not true that
in order to introduce a hexane molecule into bulk solvent,
water–water hydrogen bonds must first be broken to create
a cavity (11), after which the solute molecule may be in-
serted and allowed to interact with surrounding solvent?
For this simple series of steps the entropy should be posi-
tive (due to phase mixing) and the enthalpy should also be
positive (because the broken hydrogen bonds are much
stronger than the newly formed dipole-induced dipole in-
teractions). However, the process does not stop there.
What happens after the nonpolar solute is placed in the
aqueous solvent cavity is clearly explained in most modern
biochemistry textbooks, especially those devoted to biologi-
cal membranes (12–18). If H is zero for dissolving hexane
in water, then after the solvent cavity forms, strong attractive
interactions must be created to counterbalance the positive
enthalpy of water–water hydrogen bond breaking. Further-
more, if the process is endergonic and H is zero, then S
must be negative. Both of these facts may be explained if
one considers the change in structure that takes place in
the water “cage” surrounding the solute molecule. Water
*Email:
TSilvers@Willamette.edu
; fax: 503/375-5425.
Information • Textbooks • Media • Resources
JChemEd.chem.wisc.edu • Vol. 75 No. 1 January 1998 • Journal of Chemical Education 117
molecules at the surface of the aqueous cavity may compen-
sate for their broken hydrogen bonds by making extra hy-
drogen bonds with nearest neighbors. The resulting struc-
ture is often visualized as an icelike clathrate hydrate “in
which water provides a hydrogen-bonded framework that
contains holes…occupied by solute molecules. The frame-
work resembles ice in the sense that every water molecule
is hydrogen-bonded to four other water molecules, but
the…geometry of ice is somewhat distorted…[which] gives
rise to the types of cavity structures” (3, but see discussion
below for alternative model).
From this description it is easy to see that the newly
made hydrogen bonds in the icelike clathrate structure com-
pensate for the H-bonds that were broken initially to make
the cavity; hence H could be zero or even negative. Fur-
thermore, because of the increase in order and rigidity upon
formation of the icelike clathrate hydrate, S for the pro-
cess could easily be negative. Hildebrand (19) has provided
data that have been used to support this model, showing
that at 25 °C, methane’s diffusion coefficient in water is 40%
less than it is in carbon tetrachloride (DH2O= 1.72 ×10{5cm2/s
vs. DCCl4 = 2.89 ×10{5cm2/s). Presumably the loose clathrate
water “cages” serve to inhibit free diffusion of the nonpolar
solute. From these data it seems that both the nonpolar
solute (in this case, methane) and the aqueous solvent ex-
perience a decrease in entropy upon dissolv-
ing in water.
Hildebrand’s data have been interpreted
as supporting the clathrate cage model, but
he himself did not accept this model as a
logical explanation for the hydrophobic ef-
fect. He stressed that if clathrate cages as
such really existed, then methane’s diffusion
coefficient would be much less in water, per-
haps by one or two orders of magnitude, than
it is in CCl4. Although the clathrate cage
model has been widely accepted by most
physical chemists and biochemists (12–18,
28), in fact, experimental evidence for the in-
duction of icelike structure by the addition
of nonpolar solutes to water is equivocal (20,
21), and some researchers have presented
data that argue against the existence of
highly ordered clathrate structures (11, 19–
21). Several authors have developed an alter-
native statistical mechanical approach called
the “scaled particle cavity theory” (22). Simu-
lations using this theory suggest that the
hydrophobic effect stems mostly from the
small size of the water molecule, rather than
from hydrogen-bonded clathrate structures
(23–27). According to these results, the
probability of finding an appropriate void in
the fluid that will accommodate a solute is
low because water is such a small molecule,
not because of its hydrogen-bonding capac-
ity. Entropy loss stems from the exclusion of
water from these relatively large solute cavi-
ties, decreasing rotational and translational
freedom (25–27) of the solvent (and solute).
It is important to point out that both the cav-
ity-based model and the clathrate cage
model rest upon the fact that the hydropho-
bic effect is entropy, not enthalpy driven.
The models differ only in how they explain
the source of the entropy loss.
Finally, recent news reports featuring
“flaming ice” retrieved from the ocean depths
merit a mention here. The arguments discussed above ex-
plain why oil and water don’t mix at standard temperature
and pressure. But they also explain why oil and water do
mix below the surface of the ocean or the arctic permafrost.
In these regions of the earth, temperature is low (often sig-
nificantly less than 0 °C), and pressure is quite high. Both
of these conditions mitigate in favor of the aqueous solubil-
ity of nonpolar solutes. Recall that for most nonpolar solutes,
both S and H are negative; therefore as T decreases, solu-
bility increases. And because the average density of the two
unmixed phases is significantly less than the density of the
aqueous solution, as pressure increases, solubility increases.
Finally, when these solutions freeze under high pressure,
the hydrocarbon solutes remain stably trapped in tight pock-
ets in the frozen hydrate. A recent report (29) stated that fro-
zen clathrate hydrates “trapped beneath oceans and arctic
permafrost may contain up to 53% of the world’s organic
carbon reserves.” Methane is the most common gas found
in frozen hydrates, “and the amount of methane in these
frozen reserves in the U.S. alone may exceed [by more than
30-fold] the estimated total remaining conventional methane
resources in the entire world.” If effectively mined, these
enormous reservoirs of frozen gas hydrates could easily
“transform our energy economy from one based on oil to one
based on gas…”. All this from clathrate hydrates!
tasetuloSralopnoNgnirrefsnarTrofsnoitcnuFcimanydomrehT.1elbaT 52 °retaWottnevloScinagrOmorfC
etuloStnevloS
H
{
T
S
G
a
G
b
fer
HC
4
lCC
4
{5.01 6.22+1.21+
82
HC
4
C
6
H
6
{7.11 6.22+9.01+
82
HC
4
rehte {0.01 7.32+8.31+
3
HC
4
enaxeholcyc {0.01 6.71+6.7+
3
HC
4
enaxoid-4,1 {9.11 9.71+0.6+
3
HC
4
HC
3
HO {0.8 7.41+7.6+
3
HC
4
C
2
H
5
HO {2.8 9.41+7.6+
3
C
2
H
6
lCC
4
{5.7 7.32+2.61+9.51+
61
C
2
H
6
lCC
4
{1.7 4.22+3.51+5.51+
3
C
2
H
6
C
6
H
6
{2.9 9.42+7.51+1.51+
61
C
2
H
6
C
2
H
6
{5.01 2.62+7.51+3.61+
61
C
2
H
4
C
6
H
6
{7.6 8.81+1.21+
82
C
2
H
2
C
6
H
6
{8.0 8.8+0.8+
82
C
3
H
8
C
3
H
8
{1.7 4.72+3.02+5.02+
61
C
3
H
8
C
3
H
8
{5.7 7.82+2.12+
3
C
4
H
01
C
4
H
01
{3.3 7.82+3.52+7.42+
61
C
4
H
01
C
4
H
01
{2.4 7.82+5.42+
3
C
5
H
21
C
5
H
21
{1.2 2.13+1.92+7.82+
61
C
6
H
41
C
6
H
41
0.04.82+4.82+4.82+
61
C
6
H
6
C
6
H
6
1.2+5.71+6.91+3.91+
61
C
6
H
6
C
6
H
6
0.02.71+2.71+
82
C
6
H
6
C
6
H
6
0.00.71+0.71+
3
C
6
H
5
HC
3
C
6
H
5
HC
3
7.1+2.12+9.22+6.22+
61
C
6
H
5
HC
3
C
6
H
5
HC
3
0.00.02+0.02+
82
C
6
H
5
HC
3
C
6
H
5
HC
3
0.05.91+5.91+
3
C
6
H
5
C
2
H
5
C
6
H
5
C
2
H
5
0.2+7.32+7.52+9.52+
61
C
6
H
5
C
2
H
5
C
6
H
5
C
2
H
5
0.01.32+1.32+
3
C
6
H
5
C
3
H
7
C
6
H
5
C
3
H
7
3.2+2.62+5.82+9.82+
61
m
ro-
p
enelyx-enelyx0.03.42+3.42+
3
Note: All thermodynamic data are in units of kJ/mol.
a
Calculated from experimental values of
H
and
S
.
b
Measured experimentally.
Information • Textbooks • Media • Resources
118 Journal of Chemical Education • Vol. 75 No. 1 January 1998 • JChemEd.chem.wisc.edu
Conclusion
Although data and theory have been available for more
than half a century to explain correctly why oil and water
don’t mix, this knowledge has not yet found its way into
many introductory and organic chemistry textbooks in the
United States. It seems curious that after such a long period
of time this error remains uncorrected in introductory texts;
apparently, the simplicity of the hydrogen-bonding expla-
nation makes it a powerful and attractive model. Interest-
ingly enough, some of the blame for the widespread accep-
tance of this explanation may lie with pioneering membrane
biochemists who, even as they developed a robust new entropic
model, confusingly held onto aspects of the old enthalpic
explanation. Several examples of this are evident in the lit-
erature—for instance, in his introduction, Tanford (16, p 3)
refers to “the dominant role of water self-attraction in the
hydrocarbon-water system” and states that “the cause of the
[immiscibility] effect must lie solely in a lack of affinity be-
tween hydrocarbon and water”. By using terms like “water
self-attraction” and “lack of affinity”, Tanford seems to retain
vestiges of the very enthalpic model that he proceeds to
demolish several pages later.
The final question that must be addressed here is what
type of physical explanation for the hydrophobic effect is
appropriate at the beginning undergraduate level. Should
textbook authors continue to offer the age-old simple, famil-
iar, but incorrect enthalpic/hydrogen-bonding explanation?
Most emphatically not. Should authors offer in its place de-
tails of the correct entropic explanation? Perhaps not. This
explanation may lie beyond the scope of either first-year
introductory chemistry or second-year organic chemistry.
There may not be any need for these students to know about
clathrate hydrate “cages” or scaled particle cavity theory. A
discussion of the importance of solvation entropy effects
could fit nicely in the section on the hydrophobic effect, but
only if the format of the text allows such a discussion (i.e.,
if entropy and enthalpy have both already been introduced).
Otherwise, the simple rule “like dissolves like”, in the con-
text of a discussion of different types of intermolecular
forces, seems quite sufficient.
Acknowledgments
I wish to acknowledge here helpful discussions with my
colleagues Frances Chapple, Norm Hudak, Joyce White-
head, and Anders Nilsson.
Notes
1. Interestingly, my brief and anecdotal survey of European
(British, French, German, and Swedish) introductory chemistry
texts showed that these authors do stress the importance of en-
tropy in explaining the hydrophobic effect. Perhaps this is because
European introductory chemistry courses generally introduce the
concepts of enthalpy, entropy, and free energy all together, in
advance of their discussion of solubility.
2. Seidell’s monograph tabulated solubilities of organic liquids
in water at various temperatures. From the decrease in solubility
with temperature it was clear that the solution process was exo-
thermic. By plotting {ln
K
eq vs. 1/
T
it was possible to determine
H
solution from these data.
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