_____University of Puget Sound
Experiment 4
Department of Chemistry
Chem 110
THERMOCHEMISTRY AND HESS’ LAW
LABORATORY OBJECTIVES AND ASSESSMENTS
1. Be familiar with practical aspects of Thermochemistry.
a. Describe how a thermocouple probe works (input, mechanism, output).
b. Demonstrate how to calibrate a thermocouple, and acquire temperature
changes, using LoggerPro.
c. Demonstrate proper use of calorimeter by acquiring thermochemical
data in laboratory.
d. Explain, with an example, why the ∆H of some reactions might be
difficult to measure in the laboratory.
e. Critically assess design elements of a Styrofoam calorimetry experiment
(considering, e.g., the heat capacity of the container, reactant amount,
solvent amount, etc.).
f. Demonstrate proper responses to spills of strong acid or base solutions.
2. Understand theoretical aspects of Thermochemistry.
a. Carry out calculations regarding the heat required to change the
temperature of a sample by a given amount, and vice versa.
b. Carry out calculations regarding the heat required to melt a sample.
c. Convert Cs to Cp for a given substance.
d. Use thermochemical equations to assess the validity of Hess’s Law.
INTRODUCTION
CALORIMETRY
Thermodynamics is the study of the transfer of heat energy, and
Thermochemistry is the branch of thermodynamics that studies the transfer of heat
energy associated with chemical reactions. The main purpose of this lab is for you to
get experience in measuring the heat transfer associated with a series of chemical
reactions, and to use those measurements to learn about a relationship called Hess’s
Law.
Chemists keep track of heat transfer associated with chemical reactions using the
quantity qp. If a reaction gives off heat, we say that qp is negative, and the reaction is
called exothermic. If a reaction consumes heat (gets colder), we say that qp is positive,
and the reaction is called endothermic. The SI unit for qp is the same as for other
forms of energy, i.e., Joules.
When a reaction takes place at a constant pressure (e.g., open to air in a typical
laboratory), it turns out that qp is proportional to a change in a quantity called the
molar enthalpy, H. Chemists are interested in the molar enthalpy because it tells us
about changes in chemical bonds taking place during a reaction. The relationship is
qp = n ∆H
(1)
where “∆H” means “change in enthalpy per mole as a result of the reaction”. The SI
unit for ∆H is Joules/mole, although we often use kJ/mole.
The value of qp is often measured using a device called a calorimeter which in this
experiment is essentially a Styrofoam cup with an aqueous solution in it (Figure 1).
When a reaction takes place inside the calorimeter, a thermocouple measures the
temperature change, and the heat transferred is calculated according to
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When a reaction takes place inside the calorimeter, a thermocouple measures the
temperature change, and the heat transferred is calculated according to
q p = -qcalorimeter
(2)
where qcalorimeter is the amount of heat transferred to the calorimeter. (The negative
sign enters in here because this q is from the point of view of the calorimeter.) For the
solution, we can calculate
!T
qcalorimeter = mass ×! Cs(solution) ×! ∆T
(3)
where “∆T”
“!T” means “change in temperature” resulting from the reaction, Cs (solution) is
called the specific heat capacity of an aqueous solution (~4.18 J/g·K), and the mass
refers to the amount of aqueous solution.
Figure 1. Calorimeter from nested
Styrofoam™ cups and lid with
hole for temperature probe.
C
∆H 1
(4)
D
∆H2
(5)
D
!
→
!
→
!
→
B
∆H3
(6)
A
!
→
B
∆HA→
!B
A
C
SUM
SUM
HESS’S LAW
In 1840 Hess proposed that the
enthalpy change of a reaction is
constant, regardless of whether the
reaction occurs in one or several steps.
The heat evolved or absorbed in a single
reaction will be identical to the heat
evolved or absorbed by an entire
sequence of reactions. Hess’s law is
sometimes called the “law of
summation of heats of reaction.” By
adding the heats of the two individual
reaction steps, one obtains the same
enthalpy value as for the overall
process.:
process:
⇒
⇒∆H
∆HAA!→BB==∆H
∆H11++∆H
∆H22++∆H
∆H33
(7)
Why is this law important? How is it applied to gain useful or needed
information? Experimentally we can measure the heat changes for hundreds of
reactions. There are however, an even greater number of chemical reactions that
because of their toxic, explosive, reactive or corrosive nature or, simply because they
are incomplete or non-stoichiometric, are extremely difficult or experimentally
impossible to study in a calorimeter. For example, the reaction
2C
+
O22
2 CO
(8)
is impossible to measure experimentally, because some C and some CO22 will be
present at all stages of the oxidation of carbon to CO. Because heats of reaction are
additive, we can determine the heat change for an experimentally uncooperative
reaction if we can find sufficient other reactions for which heat changes are known, so
Thermochemistry and Hess’ Law
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that when added together, they give us the desired reaction.
Thermochemistry and Hess’ Law
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Heats of Solution
In the experiment you initially will observe and record results of the instructor
demonstration measuring heats of solution (∆Hsoln) of various solids. “Heat of
solution” is the heat (q) associated with the process of dissolving a substance; the
solvent is usually water.
Hess's Law
Secondly, you will perform three reactions, one of which is a combination of the
other two. Therefore, according to Hess's Law, the heat of reaction of the overall
reaction should be equal to the sum of the heats of reaction for the appropriate
combination of the other reactions. A primary purpose of this experiment is to explore
the validity of Hess’s law.
You will be studying the following reactions.
#1) Solid sodium hydroxide (NaOH) dissolves in water to form an aqueous
solution of ions:
Na+(aq) +
NaOH(s)
OH-(aq)
∆H1 = ?
(9)
#2) Solid sodium hydroxide reacts with aqueous acid to form water and an
aqueous solution of the salt:
NaOH(s) + H+(aq) + Cl- (aq)
H2O(l) + Na+(aq) +
Cl- (aq)
∆H2 = ? (10)
#3) Solutions of aqueous sodium hydroxide and aqueous hydrochloric acid react
to form water and aqueous salt:
Na+(aq) + OH-(aq) + H+(aq) + Cl- (aq)
H2O + Na+(aq) + Cl- (aq)
∆H3 = ? (11)
Equations (9) and (10) are combined algebraically to obtain (11). This is done by arranging
the two reactions (in reverse if necessary) in a series so that when they are added together,
some of the reactants and products will cancel algebraically, leaving you with only the
reactants and products of the third reaction, exactly as written. Recall that if you reverse a
chemical reaction, the sign of the enthalpy reverses.
Thermochemistry and Hess’ Law
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EXPERIMENTAL PROCEDURE
WORK IN PAIRS
ALL REAGENTS, LIQUID AND SOLIDS, MAY BE DISPOSED IN THE SINK FLUSHING WITH WATER.
PART 1.
HEATS OF SOLUTION
Your instructor will demonstrate the reactions to observe the ∆Hsoln for various
salts. Record the observed temperature changes in the table you prepared.
CAUTION! ! !
You will be working with the strong base sodium hydroxide (NaOH) and strong acid
hydrochloric acid (HCl). These chemicals are extremely corrosive to metals and human flesh.
If small amounts of the chemicals are spilled, immediately wipe up the spill, wash the area
with water and wipe it dry. If the chemicals get on your person, wash the affected area
immediately under a stream of water. If large spills occur, have the instructor handle it.
PART 2.
MEASURING THE HEAT OF REACTION
GENERAL SET-UP
WEAR SAFETY GOGGLES AT ALL TIMES.
1. Fill the large beaker found in your work area with DI water so it can
equilibrate to room temperature. This will be shared with the groups on
your workbench. Place the thermocouple into the beaker to monitor
temperature.
2. Open the file with the title “c110 Hess Law”. You should see the
temperature of the water on the screen.
3. For reactions you will use water from the beaker when needed.
4. Each trial, called a “run” by Loggerpro, is started by pressing “Return”
or clicking the green button; the trial is stopped by the same action.
When starting a run, IF you haven’t saved the prior run THEN
Loggerpro will show a dialogue box; simply choose the default (Append
and continue) by pressing “Return” again. The second run will overlay
on the same plot, but you can hide it later.
5. Nest two cups as shown in Figure 1 previously, cover with a lid.
The three reactions are performed with the same procedures as above, modified as
necessary depending on the state of the reactants. In Rxn #1, a solid is dissolved into
100 mL water.
In Rxn #2, a solid is dissolved into 100 mL of a hydrochloric acid solution that is
HALF the concentration of the solution used in Rxn #3. In Rxn #3, 50 mL of a sodium
hydroxide solution is mixed with 50 mL of a hydrochloric acid solution.
NOTE:
DO NOT PLACE THE CHEMICALS DIRECTLY ON THE BALANCE PANS; WEIGH THE
CHEMICALS IN BEAKERS. SPILLS MUST BE CLEANED UP IMMEDIATELY!
NOTE: You must record the molarity of all solutions.
Reaction #1
NaOH(s)
1.
Na+(aq)
+
OH-(aq)
Measure 100.0 mL of water into the nested–cup calorimeter.
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2.
3.
4.
5.
6.
7.
8.
9.
Place the thermocouple into the water and wait for the temperature to
stabilize.
Since the solid hydroxides readily picks up moisture from the air, it is
necessary to weigh it and be prepared to proceed to the next step
without delay.
Start a run by pressing “Return”.
As soon as the temperature is constant, weigh out about 2.0 grams of
solid sodium hydroxide pellets (NaOH(s)) and record the actual mass.
(CAUTION: Handle the solid hydroxide and resulting solution with
care.)
Immediately add the solid hydroxide to the nested–cup setup. Using the
temperature probe, stir vigorously and continuously until the
temperature has stabilized for at least 1 minute after reaching a
maximum and the solid is completely dissolved. You may need to
“Extend Collection” to continue recording past the initial time. As soon
as the temperature has begun to drop after reaching a maximum, you
may terminate the run by pressing “Return”.
Remove the thermocouple and touch it to litmus paper. Record the
results of the litmus test of the solution in your notebook. Put the
thermocouple in the cold water beaker.
Using the examine function, record the initial and final temperatures to
0.1°C from the plot in your notebook. Sketch the plot in your notebook.
Dispose of the salt solution, rinse and wipe the calorimeter dry.
ALL REAGENTS, LIQUID AND SOLIDS, MAY BE DISPOSED IN THE SINK FLUSHING WITH WATER.
10.
Repeat steps 1-9 two more times.
Reaction #2
NaOH(s) + H+(aq) + Cl- (aq)
H2O(l) +
Na+(aq) +
Cl- (aq)
CAUTION: Handle the acid solution with care.
1. Follow the same procedure as in Reaction #1, except that 100.0 mL of
0.50 M Hydrochloric acid (HCl) is used in place of water. Use the same
setup for the control as in Reaction #1. NOTE: The 1.0 M acid solution
supplied must be accurately diluted to obtain the concentration needed
in this reaction.
2. Place the 100.0 mL of the 0.50 M acid solution into the calorimeter.
3. Repeat steps 3 to 10 as in reaction #1. Remember to Record the results
of the litmus test.
4. Using the examine function, record the initial and final temperatures
from the plot in your notebook. Sketch the plot in your notebook.
5. Dispose of the salt solution, rinse and wipe the calorimeter dry.
ALL REAGENTS, LIQUID AND SOLIDS, MAY BE DISPOSED IN THE SINK FLUSHING WITH WATER.
6.
Repeat this reaction two more times.
Reaction #3
Na+(aq) + OH-(aq) + H+(aq) + Cl- (aq)
H 2O +
Na+(aq) +
Cl-(aq)
CAUTION: Handle the HCl and NaOH solutions with care.
1.
2.
Measure 50.0 mL of 1.0 M acid (HCl) into the nested–cup calorimeter.
Measure 50.0 mL of 1.0 M base (NaOH) into a clean graduated cylinder.
Thermochemistry and Hess’ Law
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3.
4.
5.
6.
7.
8.
Ensure that both solutions have the same temperature.
Place the thermocouple into the calorimeter, then start a run by pressing
“Return”.
As soon as the temperature stabilizes, add the 1.0 M NaOH solution to
the nested–cups. Using the temperature probe, stir gently and
continuously for about 5 minutes or until the temperature maximizes.
When the temperature has stabilized for at least 1 minute after reaching
a maximum, you may terminate the run.
Remove the thermocouple and touch it to litmus paper. Record the
results of the litmus test of the solution in your notebook.
Using the examine function, record the initial and final temperatures
from the plot. Sketch the plot in your notebook.
Dispose of the salt solution.
ALL REAGENTS, LIQUID AND SOLIDS, MAY BE DISPOSED IN THE SINK FLUSHING WITH WATER.
9.
10.
Repeat the reaction two more times.
Rinse the calorimeter. Clean and put away the apparatus.
ANALYSIS OF RESULTS
Calculation of ∆Hs for Reactions 1, 2 and 3
To calculate the enthalpies for the solutions and the reactions, use the data and
calculated results in Tables 3, 4 and 5 with equations (1), (2) and (3) to find ∆H in
kJ/mol for each run. Be sure to calculate the moles of equations (2) and (3) based on
the moles of NaOH. Be sure to show one example of each calculation.
Determine the median value for your measured enthalpies of each procedure and
enter them into a final summary table (Table 6) in your notebook, along with the
enthalpies you calculate for each of the procedures from the tabulated data in your
text and calculate the error.
Testing Hess’s Law with ∆Hs for Reactions 1, 2 and 3
To test Hess's law, find the calculated heat of reaction, ∆Hcalc, by combining the
heat of reaction per mole for reactions (9) and (10) in such a way that the net reaction
stoichiometry is the same as reaction (11). We'll call this the calculated molar heat of
reaction, ∆Hcalc. According to Hess's law, this value should equal the measured molar
heat of reaction, ∆Hmeasured, for reaction (11). The value of the “calculated molar heat of
reaction” minus “measured molar heat of reaction” is called the error.
HINTS: You want to combine the equations for reactions (9) and (10) to obtain
equation (11). This can be done by adding the two equations so that after you cancel
any species that appear as both reactant and product, the result is the same as
equation (11). It's important to treat different forms as different species -- for example,
NaOH(s) on one side of the arrow can't be canceled out by NaOH(aq) on the other
side. The trick here is that one or both of equations (9) and (10) may have to be
reversed to make this work out.
When you have the equations set up properly, you'll want to calculate the overall
∆H, called the calculated molar heat of reaction, by adding the ∆H’s for reactions #1
and #2. There are two main points to be remembered here. First, the ∆H’s for each
reaction have to be expressed in units per mole. Second, if you reversed either
equation #1 or #2, remember to change the sign of its ∆H also.
Thermochemistry and Hess’ Law
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Questions
On a separate page in your notebook, answer the following questions:
Q.1. A 2.00 g sample of NaOH(s) is dropped into a mixture of 90.0 g of water and
25.0 g of ice (at 0°C.) Will all the ice melt? What will be the final temperature?
(Use your values and ∆Hfusion for water = 6.02 kJ/mol.)
Q.2. Re-write the reactions (9), (10) and (11) to show the Hess’s law summation
with ∆Hs as in example equations on page 3.
Q.3. Using the specific heat capacity of water, Cs = 4.184 J/g·K, find the molar heat
capacity for water (J/mol·K).
Q.4. Predict how the accuracy of your results (in Reaction 1) would be affected if
you used 0.2 grams NaOH(s) instead of 2 grams. Important considerations
include the accuracy of your thermometer and balance, among others.
What to Do
BRING to lab
From Exp. 3–Spreadsheets and Data Acquisition –
1. THREE Separate Stapled Packets;
(1) A printed Formal Results Section for Exercise 1 and a Formal Results
section for Exercise 2. This means (1) a Word document from Exercise 1
containing a table and a plot and (2) a Word document from Exercise 2
including two LoggerPro™ plots, P vs. V and P vs. 1/V). Use the correct
title format, Results formatting and the Tables and Figures formatting for
each document.
(2) The answers to all the questions compiled together on one or two copy
pages from your notebook.
(3) After the pages containing the questions the remaining notebook “copy”
pages associated with Exp.3 laboratory the printed pages
For Exp. 4–Thermochemistry and Hess’ Law –
2. Print a copy of this experiment, read it and bring it to your laboratory class.
3. Bring your Lab Notebook with the PreLab assignment (details below)
completed. You will not be allowed to do the experiment without the prelab
assignment completed.
Bring a flash drive for obtaining a copy data files and documents generated in
the laboratory
PreLab Assignment –
1. Set up your lab notebook appropriately for this experiment including a title
bar and–
2. Table 1. Reagent Table for Hess’s Law including Waste Disposal instructions.
Reagents to include in the Reagent Table are sodium hydroxide, NaOH,
sodium chloride, NaCl, and hydrochloric acid, HCl.
Thermochemistry and Hess’ Law
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3. Prepare pages in your notebook with tables to record data and calculations.
You will need to prepare the following tables for collecting data and for
calculated results. Draw tables 2, 3, 4 and 5 in your laboratory notebook. All
Tables must be in landscape orientation in your notebook and each a halfpage. Leave ~½ page blank after the Tables for observations and calculations
including formula setups.
Table 2. Observed Results from ∆Hsoln Reactions
Salt
Masssalt
Masswater
Ti
Tf
∆T
Sign of ∆H
Endo or exo
Constants – Density of H2O = 1.00 g/mL
Observations
Table 3.
Trial
Reaction 1 Data and Calculated Results – H2O + NaOH(s)
Mass(H2O), g
Mass(NaOH), g
Mass(Sol’n), g
Ti, °C
Tf, °C
∆T, °C qp , kJ n(NaOH)
∆H
1
2
3
Constants – Density of Solution = 1.00 g/mL
Cs(H2O) = 4.184 J/g·K
Observations and Calculations
Table 4.
Reaction 2 Data and Calculated Results – 0.50M HCl(aq) + NaOH(s)
Trial Mass(HCl), g Mass(NaOH), g
Mass(Sol’n), g
Ti, °C
Tf, °C
∆T, °C qp , kJ n(NaOH)
∆H
1
2
3
Constants – Density of Solutions = 1.00 g/mL
Observations and Calculations
Table 5.
Reaction 3 Data and Calculated Results – 1.00M HCl(aq) +1.00M NaOH(aq)
Trial Mass(HCl), g Mass(NaOH), g Mass(Sol’n), g
Ti, °C
Tf, °C
∆T, °C qp , kJ n(NaOH)
∆H
1
2
3
Constants – Density of Solution = 1.00 g/mL
Observations and Calculations
Thermochemistry and Hess’ Law
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During the Lab: Record your data directly into your prepared data tables.
Sketch the time vs. temperature plots in your notebook. Record any additional
observations as you do the experiment. Show all work for the calculations.
Calculate the ∆H for each run and clearly indicate the median value for the
reaction. Use the median ∆H for the Hess’s Law calculations.
To Be Turned In By Each Student at the beginning of lab next week:
Be sure the assignment is removed from notebook and stapled before
coming to lab.
a) Lab Notebook page(s) with the completed Worksheet.
b) Answers to all questions compiled on separate Lab Notebook page(s) of
experiment.
c) Lab Notebook pages with the Tables 2, 3, 4 and 5 and calculations for the
reactions of Hess’s Law.
Thermochemistry and Hess’ Law
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