1 Hesss Law Lab Determine the Heat of Formation of Magnesium by ghkgkyyt

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									      Hess’s Law Lab: Determine the Heat of Formation of Magnesium Oxide

                                        Lab Report

Objective: In this calorimetry experiment we are going to do two separate reactions,
determine the Hrxn experimentally for each reaction and use Hess’s Law to add these
reactions together to find a target reaction, which corresponds to the Hf° of MgO. The
experiment will take two lab periods to complete.


Procedure: Part 1
The first reaction you will perform is: Mg(s) + 2HCl(aq)       MgCl2(aq) + H2(g)

        Accurately weigh 0.10-0.12 g of magnesium metal. Weigh your calorimeter and
then add 40-50 mL of 1 M HCl and reweigh to determine the mass of the acid. Note: we
will be using 1 M HCl in place of the distilled water in the calorimeter for this
experiment. Let the hydrochloric acid equilibrate to room temperature in the calorimeter.
Just before you add the metal, record the temperature of the HCl solution.

This is a chemical reaction, in addition to a heat transfer process. Look at the reaction
and the states of the products: what will happen to the metal in this experiment? What
this means experimentally, is that you can consider magnesium as the limiting reactant,
and that it will have to completely dissolve before equilibration will occur, so this
experiment will take a little more time than the last experiments, which was simply a heat
transfer between the metal and the water.

When you add the metal to the acid, continuously stir the heterogeneous solution to
increase the rate (speed) of the reaction. Record the temperature at 30 second intervals
for five minutes, and, if the temperature has not yet stabilized, continue recording the
temperature for another 5 minutes at 1 minute intervals.

Note: The directions state that you should record your temperature readings for 5
minutes. Do so. Do not stop when the temperature seems to stabilize...record the
temperature for 5 minutes for each trial. If the temperature has not yet stabilized, then
record the temperature for additional, 1 minute intervals until the temperature does
stabilize. We will consider “stable” when the temperature reading is the same (to the
tenths place) for 3 or more consecutive readings for 30 second interval, or 2 or more
consecutive readings at the 1 minute interval.

Calculations:
       The specific heat of 1M aqueous HCl is 4.19 Joules / g °C. The heat that the
excess acidic solution absorbed (this is the surroundings) is equal in magnitude and
opposite in sign to the heat that the products lost while being formed.




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                Qexcess HCl aqueous soln (surroundings) = -Q Mg+2HCl   MgCl2 + H2 (system)


                      Since this is true, the next equation is also true:

                     Qsoln = msolncsoln Tsoln = -Q rxn = - (m rxnc rxn T rxn)

                                                and so

                              msolncsoln Tsoln = - (m rxnc rxn T rxn)

• Calculate the heat absorbed by the surrounding solution in each trial of your
  experiments, and this will also determine the change in enthalpy of the reaction that you
  just completed (see the equations above).
• Then determine the molar heat of reaction: how much heat would be lost if your starting
  limiting reactant were one mole?
• Check to see if your two answers are within five percent of each other. Ask the
  instructor if you do not know how to do this.
• Check your answers with the instructor before doing a third trial to see if you are on the
  right track.



Procedure: Part 2
      The reaction for this part is
                     MgO(s) + 2HCl(aq)                 MgCl2(aq) + H2O(l)

The procedure is the same, but with different quantities of materials.
Accurately weigh 0.3-0.4 g of magnesium oxide metal. Weigh your calorimeter and then
add 30 mL of 1 M HCl and reweigh to determine the mass of the acid. While allowing

 Thought Question: The experiment is the same, but the quantities of materials are different.
 What does this tell you about the magnitude (how large the number is) of the heat lost in
 the first reaction compared to the heat lost in the second reaction?


the solution to equilibrate, use your spatula to pulverize any lumps in the magnesium
oxide powder. Lumps are tightly compacted powder, and this decreases the surface area
of the substance available for reaction.

Aside: [This is also true in baking. Sometimes corn starch or baking powder is so
compacted that most of the substance doesn’t even react, and so you get tiny white
(uncooked) lumps in your cake or pancake or muffins. This is why chefs sift the dry
ingredients, to make sure that all of the ingredients are fine powders, ensuring that they
will react completely.]




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When more surface area is available, the reaction will occur faster. The faster the
reaction occurs, then there is less time for heat to be lost to the environment; thus, your
experiment will be more accurate.

When you add the metal oxide to the acid, continuously stir the heterogeneous solution to
increase the rate (speed) of the reaction. Record the temperature at 30 second intervals
for 10 minutes. Check to see if all of the powder has reacted. If the temperature has not
yet stabilized, or if the powder has not completely reacted, continue recording the
temperature for another 5 minutes at 1 minute intervals until the powder has completely
reacted (limiting reactant, again) and until the temperature has stabilized. Again, all of
this data should be recorded! This will be part of your observations grade.
Repeat this reaction so that you have 3 good trials, or 3 trials that yield a Hrxn that are
within 5% of each other.

Calculations

                Qexcess HCl aqueous soln (surroundings) = -Q MgO+2HCl   MgCl2 + H2O (system)


• Calculate the heat absorbed by the surrounding solution in each trial of your
  experiments, and this will also determine the change in enthalpy of the reaction that you
  just completed (see the equations above).
• Then determine the molar heat of reaction: how much heat would be lost if your starting
  limiting reactant were one mole?


                                             Lab Report

Objective: Write an appropriate objective that mirrors the main conclusion.

Reactions (3)

Data Table: - all data taken, but you do not have to include all of the times-temperature
values. You DO need to show the initial and equilibrium temperatures. You also need to
include the reference “correct” values given in class, as part of your data. Both sets of
data (data from both reactions) should be in this section; DO NOT set up the data as if
this were two separate lab reports. If you include results in your data table, points will be
deducted.

Calculations: One sample for each type of calculation for EACH reaction.
There should be a header for each type of calculation, and then a verbal for each type of
calculation. IF the verbal is the same for BOTH reactions, then you only need to write it
once. We are half way through the semester. You need to follow these instructions
explicitly. If you want to discuss the format with me BEFORE the due date of the report,
you can do that through email or office hours.
DO NOT include more than 2 trials per calculation or points will be deducted. Show the
answers to all 6 trials in the results table which is separate from the data table, and comes


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 after the calculations section. Points will be deducted if you format your report
 differently.

 Calculations include – again one calculation using Trial 1 data in part 1 and one using
 Trial 1 data from part 2:
         •       change in temperatures ( T)
         •       heat of reaction
         •       molar heat of reaction
         •       average molar heat of reaction
         •       average deviation
         •        Hf° of MgCl2, kJ/mole (experimental)
         •       % error of Hf° of MgCl2, kJ/mole from literature value given in class.

        • To find the Hf° of MgO, kJ/mole (experimental), use Hess’s Law. You will
          have gotten your target reaction ( Hf° reaction of MgO) by the end of the lab.
          You will need ONE more reaction, which you will be able to figure out, given
          the other two reactions plus the target reaction. You will use the text book value
          (from the appendix in the text) for this reaction. Using Hess’s Law and your
          AVERAGE numbers from the experimental data for reactions 1 and 2,
          determine the heat of reaction of the target reaction. No credit will be given for
          “products – reactants” method.
        • Show the complete solution for the Hess’s Law calculation.
        • % error of Hf° of MgO, kJ/mole from literature value given in class.

Results Table: Put all of the results (answers to all calculations above) for all 6 trials in
this table. For the Hess’s Law calculation, just enter the one, final Hrxn for the results
table, and then the % error from the literature value.

No Graphs.

Discussion: Discuss any errors or difficulties from this experiment.

Conclusion: Re-state in sentence form, the most important results listed in the results
table. The numbers should be stated as “17 “ not “seventeen”. Also, percent errors
associated with the main results should be incorporated in context. (Do not list % errors
separately, but next to the result that each describes).




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