Laboratory
16.2B
Change in Enthalpy for a Reaction (4)
Lab Report = 131
Introduction: (5)
When a chemical reaction takes place, chemical
bonds in the reactants are broken and new
chemical bonds in the products are formed.
Energy is always absorbed in the breaking of
bonds and always released when bonds are
formed. If the energy required to break old bonds
is less than the energy released in forming new
bonds, the difference in energy is given off, and the
reaction is said to be exothermic. The enthalpy
change for an exothermic reaction is given a
negative sign to indicate that energy flows from the
system.
Changes in enthalpy for reactions can be
calculated by using tables of standard enthalpies of
formation of compounds. The standard enthalpy of
formation of a compound, ∆Hƒ˙, is the energy
change involved in the formation of one mole of the
compound from the elements when all reactants
and the product are in their standard states.
Standard conditions are 298˙K and a pressure of
one atmosphere.
1
1
H2(g) + Cl2(g)  HCl(aq) ∆Hƒ˙= -84.0 kJ/mol
2
2
The equations can be rearranged, which is called
Hess’s Law, to obtain the final equation.
Mg(s) + 2HCl(aq)  H2(g) + MgCl2(aq)
The standard enthalpy of formation (∆Hƒ) for the
final equation can be found by using the enthalpies
from the two equations.
In this experiment, precautions must be taken to
retain the heat energy in such a sway that the
immediate environment retains the energy so that
an accurate accounting can be made. A calorimeter
is used to isolate the reaction from the
surroundings.
In this experiment, you will measure the amount of
energy released by the production of hydrogen gas,
H2, from the reaction between magnesium metal
and hydrochloric acid and compare the
experimentally determined result with the value
calculated using the following standard enthalpies
of formation data. You will also analyze the
experiment to account for any differences.
The heat released in the above reaction is a direct
measure of the heat of formation, ∆Hƒ˙, of the
magnesium ion, Mg2+. This is because the heats of
formation of all the other species are zero. By
definition, the heat of formation for any element in
its most stable state at standard conditions is
defined as being zero. The amount of energy
released in the reaction can be calculated from the
mass of the water rises. The amount of energy
released in the reaction can be calculated from the
mass of the water in the solution, the change in
temperature, and the specific heat of water (4.18
J/g-ºC).
Mg(s) + Cl2(g)  MgCl2(s) ∆Hƒ˙= -640.7 kJ/mol
Energy released = (mass) x (specific heat) x
(change in temperature).
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Objectives: (2)
1. Calculate the theoretical change in enthalpy
for the formation of hydrogen gas.
3. Compare the experimental and theoretical
values.
2. Determine the enthalpy change
experimentally.
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Materials: (5)
Polystyrene foam cup calorimeter with lid
Temperature probe system
50-mL graduated cylinder
laboratory apron
safety goggles
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Reagents:
Magnesium ribbion
2.0 M HCl
Laboratory 16-2
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Procedure:
1. Put on your laboratory apron and safety
goggles.
7. Add the HCl to the calorimeter and record the
mass to the nearest ten-thousandths of a
gram.
2. Open up Logger Pro to get the thermometer
readings, same as Expt 15.2
8. Add the magnesium ribbon to the calorimeter
and quickly attach the lid. Make sure the
temperature probe is through the lid. Gentle
swirl the calorimeter to distribute the heat
3. Assemble the calorimeter. Take two
Styrofoam cup’s and one plastic lid. Insert
one cup inside to the other to insulate.
Carefully push the temperature probe through
the calorimeter lid and position the tip so that
it is about 2 cm from the bottom of the cup
9. Click “Collect” on logger pro. Watch the
temperature constantly, because you want
the highest temperature reading. Record the
highest reading.
3. Mass the entire calorimeter (lid and cups) to
the nearest ten-thousandths of a gram.
10. Dispose of the materials into the acid bottle
by the sink.
4. Take a 10-15 cm length of magnesium ribbon
and clean it with a piece of steel wool if
needed.
11. Thoroughly rinse out the calorimeter. Dry it
thoroughly as well.
5. Mass the cleaned piece of magnesium ribbon
to the nearest ten-thousandths of a gram and
record
12. Repeat the procedure two more times.
9. Before leaving the lab, wash your hands
thoroughly.
6. Measure 45-50-mL of 2.0 M hydrochloric acid
and record the temperature of the HCl.
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(40)Data and Observations:
Trial 1
Trial 2
Trial 3
Mass of Mg (g)
Mass of Calorimeter &
HCl soln. (g)
Mass of empty
calorimeter (g)
Mass of HCl soln. (g)
Final Temp (°C)
Initial Temp (°C)
Change in Temp (°C)
Moles of Mg
Total Joules lost by
reactants
Total joules gained by
calorimeter
kJ per mole Mg (expt)
kJ per mole Mg
(theoretical)
Percent error
Avg Percent Error
(20) Stamp…………………
Laboratory 16-2
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(20) Calculations:
(4) 1. Calculate the number of moles of
magnesium.
(4) 2. How much energy was released into the HCl
solution from production of hydrogen gas?
Assume the density and specific heat of HCl is
similar to water.
(2) 3. How much energy was gained by the
calorimeter? What assumption must be made
for this value to be correct?
(4) 4. Determine the experimental value for the
change in enthalpy per mole of magnesium.
(How much heat is produced per mole?)
(4) 5. Calculate what percent error between the
theoretical (prelab) enthalpy and your
experimental value.
(2) 6. Calculate the average percent errors.
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(10) Analysis:
(2) 1. Why is the temperature of the 2.0 M
hydrochloric acid measured in the
graduated cylinder instead of the cup?
(4) 3. Analyze your results and suggest two
reasons for the difference between your
results and the theoretical answer.
(4) 2. The negative chloride ion is not considered to
be part of this reaction. What factors allow
us to ignore the chloride ion?
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(5) Conclusion:
Analyze your data and explain if the data supports
Hess’s Law. If not, explain what factors caused
the difference.
Laboratory 16-2
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