University of Phoenix Measuring Reaction Enthalpy Lab Report

i need a lab report about Molar Volume of a Gas Report

Laboratory Report

Reports consist of the following sections:-Report Sheets – the completed Data and Report Sheets cut out of your lab manual.Calculations – If there are calculations, you must clearly and logically show all of the calculations to getfull credit.Discussion – Explain your results, referring to the chemistry theory where appropriate. Did the dataconform to what was expected? What does it show in relation to the experiment? If the results are notwhat was expected explain why. Possible sources of error (estimate experimental error in numericalresults). Conclusion – briefly state any final result(s), or conclusions that can be drawn from the experiment.References – any sources consulted for the writing of the report. These should be referenced (cited) inthe body of the text.

However, a data sheet screenshot is attached for your reference as a KEY. Part 1 is for measuring Heat Capacity of Master Calorimeter (coffee cup) to use in part 2 to calculate reaction (Mg + HCl >> MgCl2 + H2) enthalpy by plotting  Heat evolved (qsys) Vsmoles of Mg graph’s slope.

MEASURING REACTION ENTHALPY
OBJECTIVES
•
To gain proficiency in using computer controlled
data acquisition hardware and software to
collect and analyze data.
•
To understand how to calculate the heat
produced or absorbed in a process from
temperature changes.
•
To learn how to measure the heat capacity of a
calorimeter.
•
To determine the enthalpy of a reaction from
calorimetry measurements.
BACKGROUND INFORMATION
Definition of the System
In this laboratory handout, we will define the
system as the component inside the calorimeter
which initiates the heat flow. For a chemical
process, the system consists of the chemical
reaction. For a physical process, the system is the
object at a different temperature than the
calorimeter. For all systems, we will label the
associated heat absorbed/emitted as qsys. This
enables us to use one set of equations with uniform
labels for all of our calorimetry experiments. Note
that Tro tends to use more specific labels to
represent qsys such as qAl (Example 6.3) and qrxn
(Examples 6.5 and 6.8).
Heat and Work are Measures of Energy
It is experimentally difficult to measure the
internal energy of a system—the energy inside a
system when the energy is “standing still”. The
easiest way to measure energy is when it is moving
from one system to another. Energy can be
exchanged between two systems in the form of
heat (q) or work (w). To calculate the total change
in energy (∆E) of a system, we simply sum up the
heat and work produced:
D
(1)
As Equation (1) implies you cannot determine the
E=q+w energy change to a system unless you
know both the heat energy and the work energy
which has occurred.
One application of this principle is that the same
reaction can produce different amounts of heat
Dr. H.
depending upon how the experiment is set up. For
example, suppose you burn some sugar in a sealed
container with thick rigid walls (to maintain a
constant volume). Then, you burn the same
amount of sugar in a container open to the
atmosphere (to maintain a constant pressure). If
you carefully measure the heat produced by each
of these processes, you will find that the constant
pressure process always produces less heat than
the constant volume process. So, where is the
“missing heat” in the constant pressure
experiment?
This mystery is solved once we consider work—
the other half of Equation (1). In the constant
volume process, no work occurs because the rigid
walls prevent the system from expanding or
contracting. However, in the constant pressure
process, work does occur as the gaseous reaction
products—carbon dioxide and steam—push out
into the atmosphere. If the system was contained
in a flexible container, you could see it expand as
the reaction proceeded. When this work energy is
measured, we find it to be exactly equal to the
amount of “missing heat”. Thus, the total amount
of energy change (∆E) is the same in both
processes. What is altered by the change in
experimental setup is not the total energy change
but rather the proportion of energy emitted as heat
and the proportion of energy manifested as work.
Enthalpy
Since the amount of heat produced by a
reaction depends upon the experimental setup, we
must specify the conditions when comparing two
thermochemistry measurements.
In order to
distinguish between heat produced at constant
volume from heat produced at constant pressure,
enthalpy (H) is defined as the heat energy
consumed or produced at a constant pressure.
∆H = q (at constant pressure)
(2)
Because most reactions are performed under
constant pressure conditions, ∆H is the most
common
parameter
used
for
reporting
thermochemical data. However, note that ∆H is an
incomplete measure of the total energy change in
the system as it ignores any work energy. If we
need to accurately measure ∆E, we must
supplement ∆H with a measure of the work
involved in the process.
PAGE 2 of 7
The sign of ∆H indicates the direction in which
the heat is flowing. By convention, the participants
in a chemical reaction are defined as the system.
Thus when a reaction liberates heat (exothermic
reaction), the system is losing energy and ∆H is
negative. Conversely, when a reaction consumes
heat (endothermic reaction), ∆H is positive.
Enthalpy and Quantity of Reactants
The quantity of heat generated or consumed by
a chemical reaction depends upon the quantity of
participating reactants. As you will see in today’s
experiment, the more reactants consumed in the
reaction, the greater the quantity of heat consumed
or produced. In order to compare your results to
those of other experiments, we must define a
standard reactant quantity to be used in the
comparison.
By convention, the standard unit for reporting
∆Hrxn is in terms of heat produced per mole of the
reactant of interest. For example, consider the
quantity of energy generated when 1 mole of
methane (main component of natural gas) is
burned in air. The reaction is:
CH4(g) + 2 O2(g) ® CO2(g) + 2 H2O(g)
(3)
The enthalpy of reaction (or heat of reaction)
ΔHrxn is -890 kJ mol—1. This indicates that 890,000
joules of energy are liberated when 1 mole (16.0 g)
of methane and 2 moles (32.0 g) oxygen are
burned.
Now, let’s suppose that you perform an
experiment where you burn a sample of 6.00 g of
methane in excess oxygen and determine that 331
kJ of heat are produced. How well does this result
agree with the literature value? It’s not easy to
make the comparison until you convert your result
into the standard unit of kJ mol—1:
-331 kJ ´16.0 g =-883
kJ (4) 6.00 g 1 mol mol
As you can see, your result differs somewhat from
the literature value but is in reasonable agreement.
Heat Capacity
All matter at temperatures above absolute zero
(0 K) contains energy in the form of the kinetic
energy of the atoms or molecules that make up the
matter. Temperature is directly proportional to the
average kinetic energy of the molecules. When
energy is added to a material through heating, the
average molecular kinetic energy increases and
the temperature also increases. Basically matter
acts as an energy sponge. The quantity of energy
Dr. H.
per mole of molecules that the sponge can hold
depends upon the composition of the matter.
The heat capacity of a substance is the amount
of heat energy required to increase the temperature
of that substance by one degree Celsius. The
specific heat capacity (usually shortened to
specific heat and given the symbol s) is the
amount of energy required to increase the
temperature of one gram of substance by 1°C. The
specific heat capacities for the substances you will
use in this experiment are given in Table 1.
Table 1: Specific heat capacities of the
solutions used in this experiment at 25°C
Substance
Specific Heat (J g–1
°C–1)
H2O (l)
4.184
Part 2
Reaction
Mixture (aq)
3.75
The relationship between the heat (q), mass (m),
specific heat (Cs), and change in temperature (ΔT)
of a sample is given by.
q = m Cs ∆T
(5)
Specific heat capacity is typically used for
homogeneous materials. In our experiments the
aqueous reaction mixture is homogeneous, so the
heat exchanged with the calorimeter contents is
most conveniently calculated using equation (5).
We will be performing our reaction inside an
apparatus composed of a Styrofoam block
calorimeter, and a metal probe. It is difficult to
calculate the specific heat of the calorimeter and
temperature probe using equation (5) because
each of these materials has a different specific heat
and a different amount of mass of present. In this
instance, it is more useful to treat the apparatus as
a unit and to measure the heat capacity (C) of the
entire ensemble. The heat capacity is defined as
the amount of energy required to raise the entire
calorimeter apparatus by 1°C. In equation form,
q = C ∆T
(6)
In our experiment, the heat exchanged with the
calorimeter is most conveniently calculated using
equation (6). However, the actual heat capacity of
your calorimeter is unknown. Thus, in the first part
of this experiment, you will be determining the heat
capacity of your calorimeter.
Calorimeters: Heat Measuring Devices
The heat energy produced or consumed by a
process can be measured in a calorimeter. There
THERMOCHEMISTRY 1
are several types of calorimeters. One type is a
constant volume calorimeters (also known as bomb
calorimeters). These calorimeters can be operated
either adiabatically (without heat transfer) or
isothermally (at constant temperature).
A good description of a constant volume
calorimeter is given in Section 6.5 of Tro.
A somewhat more basic type of calorimeter is a
constant pressure calorimeter. This type of
calorimeter is commonly created from
Styrofoam cups as illustrated in Section 6.7 of
Tro. Adiabatic constant pressure calorimeters
of this type have significant limitations (i.e., they
cannot be used for gas phase reactions, at high
temperatures, or with organic solvents that
dissolve Styrofoam, etc). However, they are
inexpensive and they are well suited for the
aqueous reactions we will study in CHEM 120L.
We will be using a microscale high-tech
version of the coffee cup calorimeter. As
pictured in Figure 1, we will be using a small
Styrofoam block with a small hole in the center
instead of a coffee cup for the calorimeter. The
thick foam walls serve to insulate the contents
from exchanging heat with the surroundings.
Instead of an analog thermometer, we will be
using the Vernier temperature probe to measure
temperature.
This will enable us to
electronically read the temperature to the
nearest ±0.01°C and record two data points per
second.
Calorimetry: The Measurement of Heat
To use a calorimeter, a chemical reaction which
generates or consumes heat is performed inside
the calorimeter while recording the temperature.
Alternatively, an object which is warmer or cooler
than the calorimeter can be placed inside and the
temperature monitored as it equilibrates. For
accurate
results,
careful
temperature
measurements must be made before the heat flow
is initiated and after the heat flow is complete. The
difference in temperature (∆T) before and after is
then used to calculate the quantity of heat
exchanged.
As you may recall from Lab 01, the
measurement of ∆T is not as simple as it sounds.
Most chemical or physical processes take time to
go to completion and the response of the probe
lags behind the true temperature of the calorimeter
contents. During this lag time, heat is lost by the
calorimeter. Thus, by the time the reaction is
complete and our probe catches up to the sample
temperature, heat has been lost and the sample
temperature has dropped slightly.
Dr. H.
PAGE 3 of 7
A typical calorimetry temperature curve is
shown in Figure 2. As we did in Lab 01, we will
determine the true “after” temperature by
extrapolation.
This process is graphically
illustrated in Figure 2. The true “after” temperature
is determined by extrapolating back to the time at
which the reaction is initiated. In this experiment,
you will perform this extrapolation algebraically by
using the Vernier software to calculate the best fit
line and predict the “after” temperature.
Figure 2: Sample calorimeter temperature
curve.
This method of calculating ∆T compensates for
any heat loss. Thus, in our calculations we can
assume that no heat is lost to or gained from the
outside environment (qtotal = 0). The total heat flow
within the calorimeter ensemble can be divided into
three components: system (qsys), calorimeter (qcal),
and calorimeter contents (qcontents):
qtotal = qsys + qcal + qcontents = 0
(7)
In order to study the heat flow between
components of the calorimeter ensemble, equation
(7) is rearranged to:
– qsys = qcal + qcontents
(8)
From equation (8) it can be seen that any heat lost
by the system (qsys) in an exothermic process must
be equal to the total amount of heat gained by the
calorimeter (qcal) and its contents (qcontents). Note
that in an exothermic process, qsys will be a
negative quantity.
In our experiment, qcal and qcontents will be
calculated from experimental measurements using
equations (5) and (6) as described earlier. Once
these two quantities have been found, qsys is
PAGE 4 of 7
calculated from equation (8) and the molar change
of enthalpy (∆Hrxn) is calculated from equation (4).
To determine the reaction enthalpy of the
reaction
Determining the Calorimeter Heat Capacity
Mg( )s +2HCl(aq) ® MgCl2(aq)+H2( )g (12) we
need to use equation (8). In this experiment, the
system will be defined as the chemical reaction
shown above in equation (12). The contents are
defined as everything remaining in the calorimeter
after the reaction has occurred – the reaction
products plus any excess HCl. In each trial, we will
be determining the heat generated by the reaction
(qsys) by calculating the sum of the heat energy that
flows into the reaction mixture (qcontents) and the
heat energy that flows into the calorimeter (qcal).
The heat that flows into the reaction mixture is
calculated as:
To determine the heat capacity of a
calorimeter, we need to inject a known amount of
heat into the calorimeter and then measure the
resulting increase in temperature.
For our
Styrofoam apparatus, this is most easily
accomplished by adding a known amount of warm
water. In order to obtain a stable and accurate
initial temperature reading, we will start with a
known amount of cool water already in the
calorimeter. In this experiment, the system is
represented by the warm water and the calorimeter
contents are represented by the cool water.
After mixing the warm and cool water inside the
calorimeter, we will extrapolate the temperature
curves to the moment of mixing to obtain the true
temperature of the calorimeter and water samples
at mixing. Since the extrapolation compensates for
heat loss, we know that all the heat lost by the warm
water (qsys) must be equal to the heat gained by the
calorimeter (qcal) and the cool water (qconents) as
shown in equation (8). Since the warm and cool
water samples are homogeneous materials with
masses that change with each experimental trial,
versions of equation (5) will be used to calculate
qsys and qcontents:
qsys = mwarm Cs(H2O) ∆Twarm
qcontents = mcool Cs(H2O) ∆Tcool
(9)
(10)
The masses and temperature changes can be
found from your experimental measurements and
the specific heat of water is given in Table 1.
Once qsys and qcontents are known, qcal can be
found by solving equation (8):
– qsys = qcal + qcontents
(8)
Pay attention to algebraic signs as you do this.
Because the warm water is losing heat, qsys should
have a negative sign which will become positive in
equation (8). Thus, qcal will be the relatively small
difference between qsys and qcontents. Once qcal has
been found, Ccal is calculated from a rearranged
form of equation (6):
qcontents = mcontents Cs(contents) ∆Tcontents (13) where
mcontents and ∆Tcontents are obtained from
experimental measurements and Cs(contents) =
3.75 J·g-1·C-1. The heat energy that flows into the
calorimeter can be calculated as
qcal = Ccal ∆Tcal (14) where Ccal is the heat capacity
of the calorimeter determined in Part 1 and ∆Tcal =
∆Tcontents. Once these two terms have been
determined, qsys can be quickly found from
equation (8). Be sure to pay attention to algebraic
signs as you solve for qsys.
You will find that the amount of heat generated
is different for each trial as the amount of
magnesium (the limiting reagent) is changed. To
demonstrate the relationship between the amount
of magnesium and the amount of heat produced,
we will plot these variables on a graph. This graph
will also enable us to perform a graphical analysis
to determine the heat of reaction per mole of
magnesium.
REFERENCE
For additional information on the topics covered in
the Background Information, consult Sections 6.3
to 6.7 from Chemistry: A Molecular Approach
(4th Ed), Nivaldo Tro, Pearson, Boston, 2017.
EXPERIMENTAL PROCEDURE
(11)
Supplies (check out per pair; keep until next
week)
Master calorimeter
Note that because the calorimeter and cool water
experience the same temperature change,
∆Tcal=∆Tcool.
Supplies (check out per pair; return at end of
lab)
Spare calorimeter
Determining the Enthalpy of the Reaction of
Magnesium and Hydrochloric Acid
Supplies (provided at station)
One Vernier temperature probe
12″ Ruler
Ccal = qcal / ∆Tcal
Dr. H.
THERMOCHEMISTRY 1
Chemicals and Shared Lab Supplies
Magnesium ribbon, 15-16 cm
Sandpaper, scissors
2.0 M HCl(aq) hydrochloric acid
2.0 M NaOH (for waste neutralization)
Universal indicator (for waste neutralization)
Safety Precautions
PAGE 5 of 7
Part 1: Measuring the Calorimeter Heat
Capacity.
The microscale calorimeters are not all the
same, and must therefore be calibrated. You will
keep one of the calorimeters (the master
calorimeter) for next week’s experiment as well.
The other microscale calorimeter (the spare
calorimeter) will simply be used as a warm water
container and will be returned at the end of this part
of the experiment. Label this spare calorimeter
“warm” with a piece of labeling tape so that you will
not get them confused.
•
Exercise common sense and keep liquids away
from the computers and Vernier interface boxes
to avoid electrical shocks.
•
Contact with hydrochloric acid or sodium
hydroxide causes skin and eye burns. Goggles
and aprons must be worn at all times for
personal protection. If any hydrochloric acid or
sodium hydroxide is splashed onto your skin,
wash it off immediately with water.
For your heat capacity measurement, you will
need a beaker of cool water and warm water at 3545°C in a Styrofoam cup. If the warm water cools
too much, you may need to get more warm tap
water.
•
Magnesium is a reactive solid. It reacts slowly
with water and vigorously with acids to produce
hydrogen gas. While hydrogen gas is an
explosion hazard in large quantities, there is
insufficient hydrogen produced in this
microscale experiment to create a hazard. If
heated above 480°C in air, magnesium ignites
and burns in a highly exothermic reaction.
2. Use your 20 mL plastic syringe to measure out
7 mL of cool water and transfer into your
master calorimeter. Reweigh to obtain the
combined mass of the master calorimeter and
cool water.
General Notes:
1. This experiment should be done in pairs.
2. Use the same balance throughout this
experiment; offset errors may result from a
failure to do so.
3. Do not weigh master calorimeters directly on
the balance.
Instead, place a piece of
weighing paper on the balance pan, tare, and
then place the calorimeter on the paper to
obtain its net weight.
4. Throughout this procedure, you should
minimize holding calorimeters with your hand
as can this can transfer your body heat to the
calorimeter. To hold the calorimeter in a fixed
position during measurements, place it inside
a small ring stand.
5. During
measurements,
gently
and
consistently stir the calorimeter contents with
the thermometer probe until temperature is
equilibrated. Avoid contact between the
probe and the walls and bottom of the
calorimeter when stirring. This can scratch
and damage the calorimeter surface and allow
reaction solution to be absorbed into the
Styrofoam.
Dr. H.
1. Weigh your empty master calorimeter and
record the mass to the nearest 0.001 g.
3. Use your 20 mL plastic syringe to measure out
7 mL of warm water and transfer into your
spare calorimeter.
4. Place the temperature probe in the spare
calorimeter containing warm water. Wait for
15-30 seconds for the probe to equilibrate and
then start the Vernier data collection.
5. At 25 seconds, remove the temperature probe
from the warm water. Dry the temperature
probe carefully with a towel, and place
immediately into the cool water in the master
calorimeter. Do not stop data collection during
this process. Start stirring the cool water gently
and evenly with the temperature probe and
continue to do so until the end of the
experiment. Take care to not scrape the sides
or bottom of the calorimeter with the probe.
6. At 60 seconds, while one team member
continues to stir the sample, the other team
member should pour the warm water from the
spare calorimeter into the master calorimeter.
7. With continued stirring, collect data until the
temperature has remained approximately
constant (or is in a gradual decrease) for at
least 60 seconds. After this is achieved, stop
the experiment.
8. Weigh your calorimeter and combined water
samples to obtain the total mass.
PAGE 6 of 7
9. Empty your calorimeter contents into the sink,
dry the inside with a paper towel and allow to
air dry. Save your data according to the file
name pattern shown below. Do not proceed
until you have saved your raw data!
10. Repeat steps 1-9 twice to obtain three trials.
Part 2: Measuring the Heat of a Reaction
In this procedure, you will be using only the
master calorimeter to measure the enthalpy for
reaction (12):
Mg( )s +2HCl(aq) ® MgCl2(aq)+H2( )g
(12)
1. Obtain a ~13 cm length of magnesium ribbon.
If necessary, use sandpaper to remove any
black oxide until the entire surface is shiny.
Measure the length of the ribbon to the nearest
0.1 mm and its mass to the nearest 0.001 g.
2. Cut the 13 cm ribbon into lengths of
approximately 2.5, 3, 3.5, and 4 cm. (Tip: You
do not need to obtain the exact lengths given.
However, it is important to accurately know the
actual length used.) Measure and record the
lengths of these strips to the nearest 0.1 mm.
If necessary, bend the strips into a curve so
that they will fit inside the calorimeter cavity.
Follow the instruction from instructor.
3. Measure and record the mass of your dry
empty master calorimeter.
4. Using your graduated cylinder, obtain 60 mL of
2.0 M HCl and transfer to a clean beaker.
5. Use your 20 mL plastic syringe to transfer 14
mL of the 2.0 M HCl into your calorimeter.
Reweigh to obtain the combined mass of the
calorimeter and 2.0 M HCl.
6. Place the temperature probe into the HCl and
wait 15-30 seconds for the temperature to
equilibrate. Start collecting data. Continuously
stir the solution in your calorimeter throughout
this procedure, taking care to not scrape the
bottom or sides. At 20 seconds, initiate the
reaction by dropping the 2.5 cm strip into the
acid. The strip will eventually be completely
consumed. If absolutely necessary, use the
probe (not your hand) to move the strip. After
the temperature has steadily decreased for at
least 60 seconds, stop the experiment. Note:
the curve might not be completely smooth
during the reaction itself; this is perfectly
normal and will not affect your results.
7. Empty the reaction mixture from the
calorimeter into a large beaker labeled
“product mixture”, rinse with deionized water,
Dr. H.
dry the inside with a paper towel, and allow to
air dry. Save your data according to the file
name pattern below. Do not proceed until you
have saved your raw data!
9. After your data has been saved, repeat the
experiment using the 3, 3.5 and 4 cm strips.
DATA ANALYSIS TO PERFORM IN LAB
Part 1: Measuring the Calorimeter Heat
Capacity.
Refer to the Vernier operating instructions
handout provided to you in lab to find the
temperatures of the warm water, cool water, and
mixture at the time of mixing.
After performing the analysis, print out a graph
for your Trial 1 dataset.
Part 2: Measuring the Heat of a Reaction.
Refer to the Vernier operating instructions
handout provided to you in lab to find the
temperatures of the reactants and of the product
mixture at the time of mixing.
After performing the analysis, print out a graph
for your Trial 1 dataset.
DATA ANALYSIS TO PERFORM AT HOME
Note that all specific heat values needed for
your calculations are given in Table 1 of the Prelab
reading.
Part
1:
Calorimeter
Calculation
of
Heat Capacity.
The heat capacity of the calorimeter will be
calculated using equations (8), (9), (10), and (11).
In order to perform this calculation, you will need to
find mcool, mwarm, ΔTcool, and ΔTwarm from your
experimental data. Since the calorimeter starts out
at the same temperature as the cool water,
∆Tcal = ∆Tcool. Calculate the heat capacity of the
calorimeter for each of your trials. If the heat
capacity value is negative or greater than 10 J/°C,
check your work before proceeding – it is very likely
that you made a calculation mistake.
Calculate the average and standard deviation
of your calorimeter heat capacity.
Record the
average heat capacity of your calorimeter in
your Lab xx Laboratory Report NOW. You will
need this value in next week’s experiment.
Part 2: Calculation of Heat of the System.
The heat produced by the system, qsys, will be
calculated using equations (8), (13), and (14). In
order to perform this calculation, you will need to
THERMOCHEMISTRY 1
find mcontents and ΔTcontents from your experimental
data. For the calorimeter, use the value of Ccal
determined in Part 1 and note that ∆Tcal =
∆Tcontents. Calculate the heat produced by the
system, qsys, for each of your trials. Be sure that
your answer has the correct algebraic sign for a
heat producing process.
The heat produced in each of your trials is
different because different amounts of the
magnesium were used. However, reaction (12) is
the same reaction in all four trials. Therefore,
although each trial produces a different qsys, all four
trials should produce the same heat of reaction for
1 mole of magnesium. In this experiment, we will
utilize a graphical analysis to obtain a “best fit”
value. As can be seen equation (15) below, the
heat of the system for a given trial (qsys) is given by
the heat of reaction per mole of magnesium (ΔHrxn)
multiplied by the number of moles of magnesium in
the system (nsys):
D
qsys = Hrxn´nsys (15) Y Slope X
Equation (15) predicts that if the experimental
values of qsys are plotted on the Y axis versus nsys
on the X axis for each trial, the resulting plot should
be a straight line with a slope of ΔHrxn and an
intercept of zero. Thus, an accurate calculate of
the slope of the best fit line will provide the “best fit
value” of ΔHrxn for your data set.
Using your data, prepare a manual plot of qsys
versus nsys on a piece of high quality graph paper.
Computer generated graphs using Excel or other
software will NOT be accepted. Plot each of your
data points using an X with a box around it (e.g.
). After plotting your data points, draw the best
fit straight line through your points while forcing the
intercept to be (0,0). Note that the best fit line may
not pass exactly through any of your experimental
data points.
Select two convenient points on the best fit line
(not experimental data points) which are far apart
from each other. Indicate these points on your
graph by using a dot with a circle around it (e.g. ).
For each point, write down your estimated x and y
values on the graph next to the point. Using these
two points, calculate the slope of the line to obtain
the best fit value of ΔHrxn. Report your final value
of ΔHrxn in units of kJ/mol. Be sure to attach this
graph to your report when you submit it for
grading.
Dr. H.
PAGE 7 of 7
WASTE DISPOSAL AND CLEANUP
Pour any excess 2.0 M HCl into the reaction
product mixture waste beaker. Add a few drops of
universal indicator. Use a graduated cylinder to
measure out about 55 mL of 2 M NaOH and slowly
add it to the reaction mixture with stirring with a stir
rod. Observe the indicator color and assess the
acidity or basicity of the solution using the color
chart below.
pH
Acidity/Basicity
Color
0-3
Strongly Acidic
Red
3-6
Mildly Acidic
Orange/Yellow
7
Neutral
Green
8-11
Mildly Basic
Blue
11-14
Strongly Basic
Purple
Ideally, we would like to neutralize the solution so
that its final color is green (neutral) but yellow
(mildly acidic) or blue (mildly basic) colored
solutions are acceptable. Slowly add NaOH with
continuous stirring until an acceptable pH is
obtained. Flush the neutralized mixture down the
sink drain with large volumes of excess water.
Return your spare calorimeter and precision
thermometer to the storeroom. Do NOT return
your master calorimeter‼! Store the master
calorimeter in your personal locker. You will need
this calorimeter with its known heat capacity for
next week’s experiment. Unplug the Vernier probe
from your computer. Return your Styrofoam cup,
lid and temperature probe to the supplies tray next
to your bench.
BEFORE LEAVING LAB…
•
Police and clean up your area – you will not be
“signed out” until your station is clean.
•
Be sure to obtain printouts of your Part 1 Trial 1
and Part 2 Trial 1 graphs with linear regression
and interpolation calculator analyses.
•
Obtain a sheet of high quality graph paper (see
instructor’s bench) for the Part 2 manual graph.
•
Have the experimental section of your
laboratory report “signed out” by your instructor
or teaching assistant.
•
Work on completing the calculations with any
remaining time in the lab period.
Take
advantage of the opportunity to work on these
calculations while your instructor and laboratory
TA’s are available to answer questions.
Name:
Instructor sign:
Be sure to observe proper significant figures and include appropriate units in all entries.
PART 1: THE CALORIMETER HEAT CAPACITY – EXPERIMENTAL DATA
Record in ink
Calorimeter Heat Capacity Data
Trial 1
Trial 2
Trial 3
Mass of empty, dry master calorimeter
Mass of cool water and master calorimeter
Mass of combined water and master
calorimeter
Perform temperature curve analysis as noted below
Temperature of cool water at mixing
(from Vernier analysis)
Temperature of warm water at mixing
(from Vernier analysis)
Temperature of combined water mixture
(from Vernier analysis)
Temperature Curve Analysis: perform this part in the laboratory before leaving.
Using the Vernier software, perform the graphical analysis tasks described in Part
1 of the Data Analysis to Perform in Lab section. After determining the temperatures
of the cool water, warm water, and mixture by extrapolation, record these values in the
data table above. Print out a copy of the Trial 1 temperature curve with the
extrapolation lines. On your printout, indicate the time of mixing by drawing a vertical
line on the graph. Next label ∆Twarm and ∆Tcool on your graph. Be sure to attach this
graph to your report when you submit it for grading.
PART 2: MEASURING THE HEAT OF A REACTION – EXPERIMENTAL DATA Record in ink
Magnesium Strip Data
Length of Mg Ribbon Strip
Heat of Reaction Data
Mass of Magnesium Ribbon Sample
Dr. H.
Mass of Strip Mg Ribbon
Trial 1
Trial 2
Trial 3
Trial 4
THERMOCHEMISTRY 1
PAGE 2 OF 6
Mass of empty, dry master
calorimeter
Mass of 2 M HCl and master
calorimeter
Perform temperature curve analysis as noted below
Temperature of HCl at reaction start
(from Vernier analysis)
Temperature of mixture after reaction
(from Vernier analysis)
Temperature Curve Analysis: perform this part in the laboratory before leaving.
Using the Vernier software, perform the graphical analysis tasks described in Part 2 of the Data
Analysis to Perform in Lab section. After using extrapolation to determine the temperatures of the
HCl at reaction initiation and of the reaction mixture after reaction, record these values in the data
table above. Print out a copy of the Trial 1 temperature curve with the extrapolation lines. On your
printout, indicate the time of mixing by drawing a vertical line on the graph. Next label ∆T of the
reaction on your graph. Be sure to attach this graph to your report when you submit it for
grading.
THERMOCHEMISTRY 1
PAGE 3 OF 6
PART 1: THE CALORIMETER HEAT CAPACITY – CALCULATIONS
Complete In Pencil. Record the results in the table below.
1. Calculate the change in temperature of the warm water (∆Twarm) and the cool water and calorimeter
(∆Tcool) for each of your three trials. Show your work for Trial 1 in the space below. Record your answers
for all trials in the calculated values table on the next page.
2. Calculate the mass of cool water (mcool) and warm water (mwarm) for each of your three trials. Show your
work for Trial 1 in the space below. Record your answers for all trials in the calculated values table on
the next page.
3. Calculate the heat capacity of your master calorimeter for each of your three trials. Show your work for
Trial 1. Record your answers for all trials in the table below.
Dr. H.
4. Calculate the average and standard deviation of your calorimeter heat capacity values.
Calculated Values for Calorimeter
Calibration
Trial 1
Trial 2
Trial 3
Mass of cool water (mcool)
Mass of warm water (mwarm)
Change in temperature of cool water (∆Tcool)
Change in temperature of warm water
(∆Twarm)
Heat Capacity of Master Calorimeter
(Ccalorimeter)
*
Average Heat Capacity of Master Calorimeter
Standard Deviation of Heat Capacity of Master
Calorimeter
±
* Write the average heat capacity of your calorimeter in your Lab 10 Laboratory
Report NOW!
PART 2: MEASURING THE HEAT OF A REACTION – CALCULATIONS
Complete In Pencil. Record the results in the table at the top of the next page.
1. Calculate the mass of magnesium per unit length of your magnesium strip. Show your calculation and
clearly circle your answer.
2. Calculate the mass of magnesium used in each trial. Show your work for Trial 1. Record your answers
for all four trials in the calculated values table on the next page.
3. Calculate the mass of 2 M HCl used in each trial. Show your work for Trial 1. Record your answers for
all four trials in the calculated values table on the next page.
4. Calculate the mass of the calorimeter contents used in each trial. Show your work for Trial 1.
Dr. H.
THERMOCHEMISTRY 1
PAGE 4 OF 6
5. Calculate the change in solution temperature (∆T) for each of your four trials. Show your work for Trial
1. Record your answers for all four trials in the calculated values table below.
6. For each of your trials, calculate the heat evolved in the reaction (qsys). Show your work for Trial 1.
Record your answers for all four trials in the calculated values table below.
7. Calculate the moles of magnesium used in each trial. Show your work for Trial 1. Record your answers
for all four trials in the calculated values table below.
THERMOCHEMISTRY 1
PAGE 5 OF 6
PART 2: MEASURING THE HEAT OF A REACTION – CALCULATIONS (CONTINUED)
Calculated Values Table
Mass of magnesium used (mMg)
Mass of 2 M HCl (mHCl)
Mass of calorimeter contents (mcontents)
Change in temperature of contents
(ΔTcontents)
Heat evolved in trial (qsys)
Moles of magnesium used (nsys)
Dr. H.
Trial 1
Trial 2
Trial 3
Trial 4
PART 2: MEASURING THE HEAT OF A REACTION – DATA ANALYSIS
1. On a piece of high quality graph paper, manually plot the heat evolved by the system on the Y-axis against
the number of moles of magnesium used on the X-axis for each trial. Choose your axes scales so that:
(a) you utilize as much of the graph paper as possible; and (b) the major and minor graph markings are at
reasonable increments. Draw a best fit line on this graph. When you draw your best fit line, force it through
the (0,0) origin as there is no heat generated when no magnesium is present. Be sure to attach this
graph to your report when you submit it for grading.
2. Examine your graph. How would you describe the relationship between the amount of heat generated by
the system and the moles of magnesium? Are there any points that do not appear to fit the expected
trend?
3. Calculate the enthalpy of reaction (ΔHrxn) from your data set by graphical analysis. On your plot, select
two points on your best fit line that: (a) are very far apart from each other; and (b) will be relatively easy to
accurately estimate and label their x and y values. Mark these two points on your line by using a dot with
a circle around it. For each point, write down your estimated x and y values on the graph next to the point.
Use the x and y values of these two points to calculate the slope of the best fit line. Show your calculation
work below. Finally, calculate the enthalpy of reaction in units of kJ/mol and make sure it has the proper
algebraic sign. Show your work and record your answer in the blank below.
Enthalpy of Reaction (ΔHrxn) = _____________________
Dr. H.