Designing Thermochromic Thermometers

Designing Thermochromic Thermometers

Project 1, Session 2
Pre-Lab Worksheet

All work must be well-written and organized. If your penmanship is poor, you must type all responses. If you need to
organize your thoughts, please use a separate sheet of paper.
Question 1 (5 p) Page 4 of the Project 1 Guide v2.8 gives the Van’t Hoff equation:
ln⁡〖K_C= -(∆H^o)/R〗 1/T+(∆S^o)/R
Sketch a plot of ln(Kc) vs. 1/T for this equation. Title the plot and label the axes correctly. Clearly indicate how
you will find the standard enthalpy (Ho) and standard entropy (So) from this plot. Assume Ho > 0 and So > 0.
What are the units for Ho and So?

Question 2 (5 p) Depending on the sign of the standard enthalpy and entropy (Ho and So) there are four possible
ln(Kc) vs. 1/T plots. Sketch the four possible ln(Kc) vs. 1/T plots below. Label the axes correctly for each plot.
Indicate which plot you expect for Task 3.
YOU MUST READ THIS GUIDE BEFORE ARRIVING IN LAB. The lab will
be run with the assumption that you have completed all the reading.
FAILURE TO READ THE MATERIAL BEFOREHAND WILL RESULT IN
NEEDLESS CONFUSION AND FRUSTRATION IN LAB.
The central goal of this three-session project is to explore and learn
how to control the behavior of a chemical system in order to design
thermochromic thermometers that change color at specific
temperatures. With this type of thermometer, temperature changes
induce chemical processes that lead to new chemical substances that
differ in color. In this particular project, you will work to control the
extent of chemical reactions involving colored cobalt (II) complexes.

P1- 1 Background
A chemical system exhibits thermochromic behavior when its color changes over a temperature range. You may
have observed this phenomenon in products such as mood rings, baby bottles that change color when
the
contents are cool enough to drink, or actual thermometers used to measure water temperature in aquariums or
body temperature by placing them on the forehead.
Most commercial thermochromic systems use chemical substances that exhibit liquid crystal behavior and
reflect light of different wavelengths as molecules adopt different arrangement s when the temperature
chan ges. Some coordination metal complexes also exhibit thermochromism. Such systems can serve as the basis
of relatively simp le , inexpensive thermochromic thermometers.
Coordination complexes are made up of a central atom or ion (usually metallic) with
a surrounding array of bound molecules or ions, known as ligands . The number of
ligands determines the molecular geometry of the com plex. For example CoCl 2(NH
3
)
4

has an octahedral shape, while CoCl 2(NH
3
)
2
has a square planar geometry (see
illustration s to the right). In these examples, Cl
– ions (green) and NH3
molecules
(blue/white) act as ligands to the Co
2+
metal ion (grey) in the center of the
coordination complex.
The types and number of ligands affects the energy states accessible to electrons in
the coordination complex and thus determine the system’s color. For example,
octahedral complexes (6 ligands) of Co
2+
frequently have a pinkish color while
square -planar c omplexes (4 ligands) tend to be bluish.
When cobalt (II) chloride (CoCl 2
) is dissolved in methanol ( methyl alcohol, CH3OH,
MeOH), it forms octahedral complexes with the formula CoCl 2
(MeOH)
4
. However,
when CoCl2
is dissolved in other alcohols (Alc), such as ethanol, the most common
complexes are square -planar: CoCl 2
(Alc)
2
. This suggests that one can control the color
of CoCl2
alcoholic solutions by controlling the extent of chemical reactions of the
type:
CoCl 2
(Alc)
2
+ 4 MeOH ⇄ CoCl 2
(MeOH)
4
+ 2 Alc

Blue

Pink

CoCl 2(NH 3) 2 (Square Planar)

CoCl 2(NH 3) 4 (Octahedral)
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According to this equation, if you have a blue solution of CoCl 2
(Alc)
2
and add
colorless methanol (MeOH), the equilibrium will shift to the right, forming more
CoCl 2
( MeOH)
4
. Hence, as the amount of MeOH is increased, the solution should
transition from blue to purple and then go pink . Notice , in theory the equilibrium
may be driven back to the left by adding Alc (which is also colorless) so the pink
solution should return to blue .
The extent of this cobalt -alcohol reaction also depends on the system temperature.
For example , if the reaction is exothermic increasing the temperature should favor
formation of the reactant CoCl 2
(Alc)
2
and the chemical system would become more
bluish.
In P roject 1 we will focus on the behavior of CoCl 2
(Alc)
2
complexes in which the
alcohol (Alc) is ethanol or isopropanol. Common abbreviations for these alcohols are
EtOH for ethanol and iPrOH for isopropanol. Hence, CoCl 2
(Alc)
2
solutions of ethanol
and isopropanol will be referred to as CoCl 2
(EtOH)
2
and CoCl 2
(iPrOH)
2, respectively .
To take advantage of reaction s such as:
CoCl 2
(Alc)
2
+ 4 MeOH ⇄ CoCl 2
(MeOH)
4
+ 2 Alc

Blue

Pink

in the design of thermochromic thermometers we need to explore how the chemical nature of the alcohols,
their concentrations, and the temperature of the system a ffects chemical equilibria. Developing this knowledge
and applying it to the design of thermochromic thermometers are the central challenges of Project 1 .
When studying chemical equilibria a good starting point is to determine the equilibrium constant ( K c
). To better
unders tand K c, consider the generic balanced chemical equation:
aA + bB ⇄ cC + dD
Reactants Products
where the capital letters represent chemical formulas of the reactant s and products and the lower case
letters
are the stoichiometric coefficients ( the numbers in front of the chemical formulas) . As indicated, the reactants
are on the left, the products on the right. For this generic balanced chemical equation , the equilibrium
constant
expression is written:
�
�
=
[ � ]
�
[ � ]
�
[ � ]
�
[ � ]
�

Here the brackets signify molar concentrations of the reactant or product enclosed by the bracket.
From
Equation 3 , we see that K c
is the ratio of the product molar concentrations raised to their stoichiometric
coefficients over the reactant molar concentrations raised to their stoichiometric coefficients . So
clearly, if the
reactant molar concentrations are greater at equilibrium than the product molar concentrations , Equation 3
predicts K c
should be smaller than when the product molar concentrations are greater than the reactant molar
concentrations.
Now let’s apply this to the cobalt-alcohol system of Equation 1 . In this case t he equilibrium constant
expression
is :
�
�
=
[ ����
2
( ���� )
4
] [ ���]
2
[ ����
2
( ���)
2
] [ ����]
4

Methanol (CH4
O)
(MeOH)
Ethanol (C2 H6
O)
(EtOH)
Isopropanol (C
3H 8
O)
(iPrOH)
Equation 3
Equation 4
Equation 2
Equation 1
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Take a look at Equation 1. Let’s assume the equilibrium strongly favors the product side. Would you expect K c
to
be large or small? From Equation 1 , if the equilibrium strongly favors the product s , then [ CoCl 2
(MeOH)
4
][Alc]
2
>
[CoCl2
(Alc) 2
][MeOH]
4
, so that Equation 4 predict s a large K c
( K c
> > 1). Notice the reverse would be true if the
equilibrium strongly favored the reactant side – namely K c
would be small ( K c
< < 1). Hence, the magnitude of K c

affords some insight as to where the equilibrium lies.
Can the equilibrium constant ( K c
) be negative? Why or why not?
From Equation 4, you also see that to calculate K c
the equilibrium concentrations of all the species in the system
must be known. In this project you will not have to measure or determine all these concentrations,
as a
spreadsheet has been prepared for you that will calculate K c
once the concentration of CoCl 2
(Alc)
2
is found . So
how do you find [CoCl 2
(Alc)
2
]? Since the cobalt complexes are colored, this task can be accomplished using
visible absorbance spectroscopy. You know from CHEM 151 that abso rbance spectroscopy involves a
mathematical relationship between absorbance and concentration called Beer’s Law, A = εlC , in which ε is the
molar absorptivity and l is the path length. The path length is always 1.00 cm for the cuvettes we use in lab and
the molar absorptivity at λmax = 656 nm of CoCl 2
(EtOH)
2
is 170 cm
– 1
M
– 1
, while that of CoCl 2
(iPrOH)
2
at λ max
= 656
nm is 319 cm
– 1
M
– 1
. So , the molar concentration at equilibrium of the CoCl 2
(Alc)
2
complexes can be found . Let us
call this C eq
. Once this value is known, you can use the Excel file Project 1 Task 2 to find K c
.
Should the equilibrium constant (K c
) be independent of the initial concentr ation of the different species? Why?
P1- 2 Your Challenges
In Project 1 you are expected to explore the behavior of reactions of the type:
CoCl 2
(Alc)
2
+ 4 MeOH ⇄ CoCl 2 (Me OH)
4
+ 2 Alc
w ith the objective of understand ing how to control the chemical equilibrium and how such processes may
be used to design thermochromic thermometers that change color at a specific temperature . To
accomplish this , the following four major tasks must be complete d:
Task 1
Objective : Qualitatively explore the thermochromic behavior of cobalt alcohol solutions (CoCl 2
(Alc)
2
) with
methanol.
Prepare CoCl 2
(EtOH)
2/MeOH in various ratios between 2:1 to 4:1 ( CoCl 2
( EtOH)
2
to MeOH) and
observe the color at room temperature and when placed on ice. Do likewise for
CoCl 2
( iPrOH )
2/MeOH.
Expected outcome : A qualitative feel for the behavior of the cobalt alcohol complexes with methanol at
different temperatures.
Resources:
x 20 mL of 0.0101 M CoCl 2(EtOH) 2 stock solution – use a 20 mL vial & cap to control evaporation ; NO
BEAKERS!
x 20 mL of 0.0103 M CoCl 2(iPrOH)2 stock solution – use a 20 mL vial & cap to control evaporation ; NO
BEAKERS!
x 20 mL of Methanol – use a 20 mL vial and cap to control evaporation and spills; NO BEAKERS!
x 20 – 200 µL micropipette (hanging on the micropipette station attached to the islands)
x 100 – 1000 µL micropipette (hanging on the micropipette station attached to the islands)
x 200 µL pipette tips (for the 20-200 µL micropipette)
x 1000 µL pipette tips (for the 100-1000 µL micropipette)
x Contents of your shared locker (see the locker inventory sheet and photo on D2L )
You will have 30 minutes to complete this task. Use your notebook to clearly record and describe the work
you are performing and the associated results.
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Task 2
Objective : Determine the average value of K C
at room temperature for chemical reactions involving cobalt
alcohol solutions (CoCl 2
(Alc)
2
) with methanol.
a) Prepare CoCl 2
(EtOH)
2
/MeOH in ratio s between 2:1 to 4:1 ( CoCl 2
( EtOH)
2
to MeOH). Do likewise for
CoCl 2
( iPrOH )
2/MeOH. You will have to decide how many ratios to prepare.
b) Using absorbance spectroscopy determine the molar concentration of CoCl 2
(EtOH)
2
or
CoCl 2
(iPrOH)
2
for the ratios prepared above.
c) Using the Excel file Project 1 Task 2 v1.1 to find K c
for each prepared ratio.
Expected outcomes: You should be able to determine the equilibrium constant K c
at room temperature
for CoCl 2
(EtOH)
2
and CoCl 2
(iPrOH)
2
with MeOH at various ratios.
Resources (In addition to items listed in Task 1 ) :
x 5 mL of Ethanol – only for zeroing the spectrometer ; NO BEAKERS!
x 5 mL of Isopropanol – only for zeroing the spectrometer ; NO BEAKERS!
x USB650 absorbance spectrometer with USB cable (one per group)
Detailed operating instructions for the USB650 absorbance spectrometer are given in Absorbance
Spectroscopy with Logger Pro v2.5 technical guide on D2L. This information is not available in lab, so
bring
either a hard or electronic copy .
You will have the rest of the lab session to complete Task 2 . Use your notebook to clearly record and
describe the work you are performing and the associated results.

Task 3
The equilibrium constant ( K C
) should be independent of the initial concentration of the different species.
However, it may be affected by the temperature of the system. The relationship between K c
and T (in
Kelvins) is given by the Van’t Hoff equation:
ln �
�
= −
∆ �
�
�
1
�
+
∆ �
�
�

� = �� + �
where ‘H
o
and ‘S
o
are the standard enthalpy and entropy change of the reaction, respectively, and R is
the ideal gas constant ( R = 8.314 J/(K ·mol)). Th is relationship suggests that the value of ‘H
o
and ‘S
o
may
be derived experimentally using K c
values at different temperatures T . This can be done by plotting ln(Kc
)
versus 1/T to prepare a Van’t Hoff plot which should result in a straight line with slope m = -‘H
o
/R and y -intercept b = ‘S
o
/R.
Objective: Infer the values of ‘H
o
and ‘S
o
for chemical reactions involving CoCl 2
(EtOH)
2
with MeOH and
CoCl 2
(iPrOH)
2
with MeOH. In particular, f or CoCl 2
(EtOH)
2/MeOH and CoCl 2
( iPrOH )
2/MeOH you are
expected to find values for ‘H
o
and ‘S
o
by determining K c
at different temperatures.
a) Prepare CoCl 2
(EtOH)
2/MeOH in ratios between 2:1 to 4:1 ( CoCl 2
( EtOH)
2
to MeOH) and determine
K c
for each preparation as a function of temperature T . Do likewise for CoCl 2
( iPrOH )
2
/MeOH.
b) Using the Excel file Project 1 Task 3 v1.1 to find K c
as a function of temperature T for each
prepared ratio.

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c) From the Excel spreadsheet output, work-up the K c
versus T data to determine ‘H
o
and ‘S
o
for
CoCl 2
(EtOH)
2
/MeOH and CoCl 2
(iPrOH)
2
/MeOH. (Hint, this will involve constructing a Van’t Hoff
plot. )
Expected outcomes: ‘H
o
and ‘S
o
values for CoCl 2
(EtOH)
2
/MeOH and CoCl 2
(iPrOH)
2
/MeOH ratios that can
be analyzed for trends.
Resources (In addition to items listed in Task 2 ) :
x Ice Melter (used to depress the freezing point of ice; on the reagent bench )
x Plastic transfer pipettes (on the reagent bench)
x Plastic ice bowls (only 1 per student; the ice machine is located in the hallway)
x Vernier surface temperature sensor with Go!Link interface and cap to align sensor tip in a cuvette (digital
te mperature probe for measuring temperature of the solution in a cuvette)
x Go!Temp digital thermometers (for monitoring bath temperatures only)
Measurements of absorbance at different temperatures can be systematized using this procedure:

1) Set-up Logger Pro to collect absorbance and temperature co nc urrently for 100 seconds at 1
second intervals.
2) Load 4 mL of the reaction mixture into a 2-dram (7 mL) screw cap glass vial.
3) Chill in an “Ice Melter” ice bath for 5 -10 minutes.
4) Quickly transfer via disposable pipette enough chilled solution to fill t he cuvette ¾ full ( ≈ 3 mL ).
5) P lace cuvette in spectrometer . Immediately insert surface temperature probe .
6) Monitor temperature display on Logger Pro window – in itiate data collection when the
temperature appears to hit a minimum .
7) Migrate data to Excel for work up.

Concurrent Capture of Absorbance and Temperature v2. 1 gives d etailed operating instructions for the
USB650 absorbance spectrometer with the Vernier surface temperature sensor. The document is on D2L.
This information is not posted in lab, so bring a copy.
You will have one lab session to complete this task. It is critical that you have a plan for acquiring the
data
you need . As you work in the lab, use your notebook to clearly record and describe the work performed
and the associated results.

Task 4
In the final phase of this project you are expected to apply what you have learned in the previous two
sessions to design an inexpensive, self -contained, disposable thermochromic thermometer that change s
color at a specific temperature. In particular, you should design a thermometer that can be used for
quality control of products that should never be exposed to temperatures below freezing (or around 2
o
C,
the temperature of a wat er -ice mixture), such as refrigerated blood samples in blood banks or vegetables
in supermarket storage units .
Resources:
You will have access to the same resources that were available in previous session s to complete Task 4 .
You are expected to explore the properties and behaviors of different reaction mixtures using your
previous experimental results to guide the process. To assist in performing a cost analysis, a table
of
material cost follows:

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P1 v2.8 8 -22-15 MY

Thermochromic Thermomete r Material Costs for Task 4

This information can be used to guide your thermochromic thermometer design to meet the “inexpensive”
criterion.
P1- 3 Preparatory Tasks
This project requires you to work with a partner . You are not expected to personally conduct all the lab work;
rather, you and partner should coordinate efforts to successfully complete the project . The two of you will be
responsible for monitoring time. It is important that you develop a clear plan of action before starting any
experimental work. You should come pre pared with ideas about how to implement the experiments that will
help you meet the tasks given the resources available to you.
In this project you will employ a number of te chniques/procedures. Before arriving to lab it is very important
to review the technique guides . In particular, you should:
x Task 2 : Understand how to use a spectrometer for taking absorbance readings (Absorbance Spectroscopy
with Logger Pro v2.5 technical guide on D2L ).
x Tasks 1 -4 : Learn the proper use of a micropipette (Micropipette v2.5 technical guide on D2L ) .
x Task 3 : Be able to capture temperature and absorbance data simultaneously in a small volume of solution.
( Concurrent Capture of Absorbance and Temperature v2.1 technical guide on D2L ).
Make sure that you:
x Review the technical guides. You are expected to know how to properly use the lab equipment
without too much reliance on your TA.
x Don’t forget to complete the P re -L ab Quizzes in D2L (one every session) .
x Don’t forget to submit (lab D2L Dropbox) the P re -Lab Worksheet s before lab (one every session) .
Things to bring to lab :
x A hard copy or an electronic copy of the:
o Project 1 Guide v2.8 (this documen t)
o Pre-Lab Worksheet (one for each session: P1S1, P1S2, P1S3) – MUST BE COMPLETED BEFORE LAB!
o In -Lab Report (one for each session: P1S1, P1S2, P1S3)
o Absorbance Spectroscopy with Logger Pro v2.5 technical guide (For Task 2)
o Micropipette v2. 5 technical guide (For all Tasks)
o Concurrent Capture of Absorbance and Temperature v2.1 technical guide (For Task 3 )
o Excel spreadsheets: Project 1 Task 2 v1.1 ( for Task 2 only), Project 1 Task 3 v1.1 (for Task 3)
x A laptop computer with Excel (to facilitate the analysis of the data) and Logger Pro (to directly
collect data from a USB650 spectrometer and the Vernier temperature sensor s ). The lab is NOT
equipped with computers for student use.
Material Cost
1- Dram vial + cap $0.29 each
2- Dram vial + cap $0.19 each
20 mL Plastic vial + cap $0.18 each
15 mL Centrifuge tube + cap $0.24 each
0.01 01 M CoCl 2
(EtOH)
2
$58.11 per liter
0.01 03 M CoCl 2
(iPrOH)
2
$36.38 per liter
Methanol $38.48 per liter
Ethanol $57.28 per liter
Isopropanol $35.55 per liter
1” Parafilm square $0.02 each
Tape $0.08 per meter