Chem 451
An Investigation of Solvatochromic Behavior
Adapted from Machado, C., Vanderlei, G. M. J. Chem. Ed. 2001, 78, 649-651.
This experiment is intended to demonstrate some of the ways in which the polarity of a solvent
can influence the energy levels, and hence the spectroscopic properties of a solute. The dye used
in this case (Reichardt’s dye), provides a striking example as the compound absorbs in the visible
part of the spectrum, and the range of wavelengths absorbed depends sensitively enough upon
the polarity of the solvent that we are readily able to observe color changes simply by changing
the solvent. Indeed, this dye has been employed as a quantitative measure of the polarity of
various solvents.
The procedure itself is quite simple, and the data analysis is (compared to some of the other
things we do in p chem) not particularly involved, yet there are definitely some insights to be
gained, and the various colors observed for the solutions should leave a lasting impression.
Information regarding the theoretical background and data analysis for this project is given
following the procedure.
Procedure
Reichardt’s dye is available as a powder. The molar mass is 551.69 g/mol, and the structure is
shown below. (I am currently accepting bribes to keep this structure out of the hands of the
organic chemists, who would undoubtedly love to use this in a systematic nomenclature
assignment!)
N
O
You will need to work together to prepare 10 mL portions of solution (in each of the solvents
that are provided) that are ~ 1 x 10-4 M. We are really only concerned with determining max for
each solution, so you don’t need to be overly quantitative in preparing these solutions. (We
simply want the maximum absorbance values for each solution to fall within a reasonable range.)
We are dealing with pretty small amounts of solid here, so we will use a little “trick” to deliver
the solid for each sample.
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Solvatochromic Behavior – page 1
Start by preparing 100 mL of a 1 x 10-3 M solution of Reichardt’s dye in acetone. (Weigh the
solid on the analytical balance and use a volumetric flask.) Once this stock solution has been
prepared, obtain a clean 50 mL beaker for each solution that you will prepare. (One for each
solvent, label the beakers accordingly.) Use a pipet to deliver 1 mL of the stock solution to each
beaker, and allow the acetone to evaporate … this process can be hastened by placing the
beakers near the front of the fume hood. Once the acetone has evaporated, use a graduated
cylinder to add 10 mL of each solvent to the beakers. (Be sure to rinse the graduated cylinder
thoroughly between solvents.) Mix using a glass rod or magnetic stirrer as necessary. Keep the
beakers covered to retard excessive evaporation. (A little evaporation won’t hurt the
experimental results in this case, but we don’t want these solutions to dry up completely, and we
want to avoid getting excessive amounts of solvent vapor in the air.)
Acquire the visible spectrum of each solution in the range 400 – 700 nm (use a step size of 2
nm). Use quartz cuvettes because some of the solvents will “interact” with plastic cuvettes in a
disruptive manner. (Naturally, the blank will be the pure solvent.) The maximum absorbance
should not exceed ~1, and the peak should be reasonably well-defined within this wavelength
range. If spectra do not conform to these guidelines, you should be able to make appropriate
adjustments and repeat the acquisition of the spectrum without too much difficulty. Use HachLink® to save each of your spectra as an Excel® file.
Once you have completed the acquisition of the spectra for each of the solutions, you will
prepare a second set of solutions using acetone mixed with varying amounts of water as the
solvent. Again, you will need to dispense the stock solution in 1 mL portions into small beakers,
and use graduated cylinders to add the amounts of water and acetone indicated in the following
table:
% H2O
(by volume)
H2O (mL)
Acetone (mL)
5
10
20
30
40
50
0.5
9.5
1
9
2
8
3
7
4
6
5
5
Use appropriate care with solvents and solutions. Dispose of all solvents and solutions in the
appropriate waste containers.
Background
By this point in the course, you should fully appreciate the relationship between the observed
wavelength of an absorption feature in a spectroscopic measurement and the energy levels within
hc
the molecular system. (This relationship is embodied in the equation E 
.) It is also

important to recognize the role played by the molecular-scale environment in determining the
wavelength observed for the absorption. When we were studying chapters 16 and 17 in Atkins
& de Paula, we focused almost exclusively on the quantized energy levels of an “isolated” (i.e.,
gas phase molecule). It should come as no surprise to you that the observed absorption spectrum
will most likely be modified in the condensed phase. (You certainly encountered this reality in
Thoemke / Spring 2003
Solvatochromic Behavior – page 2
your FTIR studies of assorted organic critters … the spectrum of a given compound will most
likely appear different in a nujol mull rather than as the neat liquid, for example.)
This observation is especially relevant when discussing electronic spectra in solution. Any
electronic transition will result, to some extent, in a redistribution of electron density in the target
species. When the target species is a solute molecule surrounded by solvent, molecular
interactions frequently play a major role in determining the overall energy of the system. In
particular the dipole moment of the solute will participate in an electrostatic interaction with the
surrounding solvent molecules. The details of these interactions depend sensitively on the
charge distribution (i.e., dipole moment, ) and molecular structures of the solute and solvent.
Upon electronic excitation, the redistribution of electron density within the solute will change the
solute dipole moment and consequently, the interactions with the solvent … thus changing the
energetics of the system. Further, the energetics of the system will also depend on the polarity of
the solvent. In sum, then, we would expect that the observed absorbance spectrum would
depend, to some extent, upon the details of the solute – solvent interactions. The diagram below
“generically” depicts this scenario.

This picture shows the energy levels of the
same solute molecule under varying

solvation conditions. The levels shown

correspond to a solute molecule for which
the dipole moment of the ground state is
greater than the dipole moment of the
Gas
excited state (i.e., g > e). Note that
phase
solvation lowers the energy of (stabilizes)
Less polar
the solute, but this stabilization occurs to a
solvent
greater extent in the ground state, with its
larger dipole moment. Remember that this
is intended merely to be a “schematic”
More polar representation; don’t give this picture too
solvent
much quantitative significance.
The key thing to note is that the energetic separation between the ground and excited states will
increase with increasing solvent polarity, and thus the transition will shift to shorter wavelengths.
(If the order of dipole moments is reversed, i.e., if g < e the opposite shift will occur…
increasing solvent polarity will shift the transition to lower energy / longer wavelength.)
For the compound of interest in this experiment, Reichardt’s dye, the electronic transition that
occurs upon absorbing visible light is shown below. Note that the charge imbalance is “less
pronounced” in the right hand structure … i.e., all of the atoms have zero net formal charge.
Thus, the dipole moment is smaller in the excited state, so we should anticipate that the transition
should shift to shorter wavelengths with increasing solvent polarity.
Thoemke / Spring 2003
Solvatochromic Behavior – page 3
N
N
O
O
This dye compound is in fact used as a measure of solvent polarity, i.e., solvents can be ranked in
order of polarity based on the color taken on by solutions containing this dye. Reichardt (the
namesake of this compound) is the author of the best-seller Solvents and Solvent Effects in
Organic Chemistry, and has done much to advance our understanding of solvent behavior.
While your previous work in chemistry has given you a reasonable intuitive sense about what we
mean by “polar” solvents as opposed to “nonpolar” solvents, it turns out to be non-trivial to
quantify solvent behavior. Numerous scientists have proposed various ways of considering the
issue. The behavior of a solute in solution depends on the interplay of a number of factors. (If
you think back to thermodynamics, you’ll remember that mixing may or may not be
spontaneous, depending on the details of the enthalpy and entropy changes associated with the
mixing process.) From the molecular standpoint, “solvation” effects depend on the molecular
structure (of both solvent and solute) and molecular quantities such as dipole moment and
polarizability (a measure of the ease with which the electron density of a molecule is distorted by
the presence of a neighboring dipole.) It possible to use quantum mechanical computational
methods to determine these molecular properties, and there are standard experimental measures
as well. This is not the entire story, however, as the bulk properties of the solvent play a role as
well, since in dilute solutions there are far more solvent molecules than solute molecules, and the
collective behavior of the solvent molecules is also relevant. In short, the properties of solvents
and the behavior of a solute in a given solvent depends on both bulk and molecular properties.
While we have not been able to address the electrical properties of molecules in class, these ideas
are treated in chapter 21 of Atkins & de Paula (see especially pp. 686 – 693). We are
particularly interested in the molecular dipole moment, and the bulk dielectric constant. It is
important to recognize that the dipole moment is a property of an individual molecule, while the
dielectric constant is a consequence of the collective behavior of the molecules in a bulk sample.
The dielectric constant can be related to the dipole moment and the polarizability of the
molecules via the Debye equation.
Thoemke / Spring 2003
Solvatochromic Behavior – page 4
For our purposes, we do not need to get into all the details of these two quantities; rather, we will
simply consider how the characteristics of the absorption spectrum of Reichardt’s dye can be
related to the solvent properties.
Analysis
Start with the spectra that you’ve acquired for all of the various solvents (for now, you will not
use the acetone – water mixtures, these will be considered later in the analysis). These should be
gathered collectively into a single spectrum, and printed out in an appropriately labeled graph.
Prepare a table that lists the solvent, the perceived color of the solution, the color (or colors)
absorbed by the solution, and the value of max. (You will need to consult the handout from early
in the semester, or table 17.1 in Atkins & de Paula to correlate the wavelength absorbed with the
color absorbed.)
Next, determine the transition energy (in kJ mol-1) that corresponds to each of your max values.
Thinking back to the “energy level diagram” given earlier in the handout, it makes sense to graph
the transition energy as a function of dipole moment and in another graph as a function of
dielectric constant. These values are given in the table on the next page.
Now, perform the first part of the analysis on the acetone : water mixtures. (Plot out spectra,
generate table of colors and max values, determine corresponding transition energies.) Prepare a
graph of transition energy versus vol. % H2O for the set of acetone:water mixtures.
All of the foregoing graphs and tables should be presented in a clear, organized fashion. You
should finish your report with a 2-3 paragraph discussion of the following issues. Study your
graphs of transition energy versus solvent dipole moment (a molecular property) and versus
solvent dielectric constant (a bulk property), and think about these graphs within the context of
the energy level diagram given above. Based on your data, which quantity offers the strongest
trend? Which of these quantitative measures would be most useful for predicting the value of
max of Reichardt’s dye in a different solvent (one that we did not use.)? What do these
observations tell you about the nature of solvent – solute interactions … do these interactions
depend more strongly on molecular solvent properties or bulk solvent properties? Now, think
about these same issues with respect to the acetone:water solvent mixture. Do your results
suggest any “abrupt” change in solvation behavior as more water is added to the system, or is
there simply a gradual transition from “pure acetone behavior” to “pure water behavior”? You
probably found that some of the solvents listed did not readily dissolve the Reichardt’s dye. Can
you make any correlation between this observation and the solvent parameter data given in the
table below? (In other words, does this observed solubility behavior relate to any specific value
of dielectric constant or dipole moment?)
Note that these questions are meant to be speculative. There is not necessarily a “single correct”
answer. I am looking for a thoughtful discussion of these issues that is developed using
scientific reasoning based on your experimental data.
Thoemke / Spring 2003
Solvatochromic Behavior – page 5
Solvent Data
Values from Gordon, A.J., Ford, R.A., The Chemist’s Companion: A Handbook of Practical Data, Techniques, and
References Wiley, 1972
Solvent
Acetone (2-propanone)
Acetonitrile
1-butanol
Cyclohexane
Dimethyl Sulfoxide (DMSO)
Ethanol
Ethyl Acetate
Isopropyl Alcohol (2-propanol)
Methanol
1-propanol
Water
Thoemke / Spring 2003
Dielectric Constant
r, as defined in A & dP,
Chap. 21
Dipole Moment
 / (10-30 C m)
(at 298 K, unless noted otherwise)
9.61
13.1
5.00
0
13.2
5.64
5.94
5.54
5.67
5.60
6.14
20.7
36.2
7.33
2.02 (@ 293 K)
49
24.3
6.02
18.3
32.6
20.1
78.5
Solvatochromic Behavior – page 6