CALORIMETRY – EXPERIMENT B

ENTHALPY OF FORMATION OF MAGNESIUM OXIDE

# INTRODUCTION

This experiment has three primary objectives:

1.      Find the heat capacity (Cp) of a calorimeter and contents (calibration).

2.      Determine the DHrxn, the enthalpy of reaction, for several different reactions, including the reaction of an unknown with a solution of HCl.

3.      Calculate the DHf, the enthalpy of formation, of MgO.

We will assume that the energy exchanged between the calorimeter and the surroundings during and following the reactions is small and at a slow, constant rate. You will become familiar with calorimetry as this experiment is completed.

# BACKGROUND

Calorimetry measures the energy that a reaction produces or consumes.  For example, the major difference between gasoline grades is the octane number.  Unleaded gas has an octane of 86, while Super Unleaded gas has a higher octane.  Calorimetry could be used to measure the heat or energy produced when gasoline is burned.  More heat (energy) would be produced by the super unleaded gas so it would have a higher enthalpy compared to just unleaded gas.  Calorimetry could be used to see if a gasoline station is selling the grades of gasoline it advertises.

The calories in food hascalories in food have also been measured by calorimetry (hence the term calories).  Usually this is a measurement of calories (cal) per gram of food.  Remember that calories iscalories are easily convertible to joules (J) and grams can be converted to moles if it is a pure chemical.

Enthalpy, represented by the symbol H, is a property chemists use to describe the heat flow into or out of a system in a constant-pressure process. This is often the case since most processes that are carried out are exposed to the atmosphere as are the reactions carried out in this course. The enthalpy of a reaction, DH, is defined as the difference between the enthalpies of the products and the enthalpies of the reactants.  In other words, it is the change in energy for a given amount of a given reaction. The enthalpy of formation, DHf is defined as the enthalpy or heat change that results when one mole of a compound is formed from its elements. The standard enthalpy of formation is defined as the enthalpy of formation measured at 1 atm such that the elements are in their standard state.

If a reaction is exothermic, heat will be released, and the temperature of the system or reaction mixture will rise. (In this experiment the heat and temperature rapidly increase and then slowly decrease as heat is lost to the surroundings.) For endothermic reactions heat will be absorbed or used and the temperature will decrease. In this experiment we will use the experimentally measured enthalpy of reaction for a series of exothermic reactions and Hess' Law to determine the heat of formation for magnesium oxide (MgO). We will also determine the enthalpy of reaction for an unknown metal oxide with an acid.

The enthalpy of reaction, DH, can be calculated using the equation:

B-1

where n is the moles of limiting reagent, DT is the change in temperature of the calorimeter’s contents, and Cp is the heat capacity of the calorimeter. The value for n can be determined knowing the amounts of starting material. The DT for a reaction can be calculated using the temperatures before and after the reaction or the initial and final temperatures. The heat capacity, Cp, of the calorimeter has to be experimentally determined by doing a reaction where the DH is known. The heat capacity of the calorimeter is primarily due to the solution in the cup.

Heat capacity (Cp) has units of kJ/°C.  Physically, this means that it takes the value of the Cp in energy to raise the calorimeter by 1°C.  For example, if a calorimeter has a Cp of 0.200 kJ/°C, the calorimeter, including its contents, must absorb 0.200 kJ of energy to increase 1°C.  A 20 kJ/°C calorimeter increases 1°C with one hundred times more energy, or 20 kJ.  Cp is a relative term to describe how much energy is needed to change the temperature of a substance or system.  The Cp of an ocean is huge (compared to a drop of water).  The heat needed to change the drop of water by 1°C could not change the ocean by 1°C.  The oceans of the world maintain the earth at temperatures that support life.Figure B-1.  Calorimeter Apparatus

In this experiment, the calorimeter is defined as two nested styrofoam cups, the lid, magnetic stir bar, and the temperature probe tip, plus the 60.0 mLl of the reaction mixture (mainly water). In order for the heat capacity of the calorimeter to remain constant, all of these must be present.

NOTE:  If less than 60 mL of reaction mixture was added, it would take less energy to increase the calorimeter and contents by 1°C.  In other words, the heat capacity would decrease.  If more than 60 mL of the mixture was added, more energy would be needed to increase the calorimeter and contents by 1°C.  The heat capacity is then increasing.

Most importantly the volume of reaction mixture must remain constant because of the large heat capacity of water.  We want the Cp to remain constant because it is the standard by which we can calculate unknown DHrxn values in reactions 1, 2, and 4.  The Cp is determined in reaction 3, using the known enthalpy (energy/mole) for a strong acid/strong base reaction: In order to determine the Cp for the calorimeter we need to do a reaction where the heat of reaction is known:

H+ (aq) + OH- (aq) ® H2O (l);        DH = -55.90 (KJ/mol) at 25°C                   B-2

Taking Rearranging equation B-1 we can use this DH and to solve for Cp to obtain:.

B-3

A value for DT can then be determined for a known amount of moles (n). Once the Cp is known we can use it to calculate DH for other reactions where DT has been experimentally determined. Please look in your textbook under calorimetry or thermodynamics for more information on these concepts.  The explanation to determine DT is in the experimental section below.

EXPERIMENTAL

This experiment essentially has three parts. In the first lab period, the data to determine the enthalpy of reaction for Mg + HCl and MgO + HCl will be collected the first lab of this experiment (one trial on each). During the second lab period, data will be collected to calculate the Cp using the reaction of NaOH with HCl (two trials). Also, the enthalpy of reaction for the exp. B unknown reacting with HCl will be determined (two trials). The trials for the Cp and unknown should be run in the same lab period so if you start them the first week of the experiment make sure you can finish them in that period..

Be sure to label each graph carefully with your name, section, date, reaction & trial, mass of reactant, etc. Note that the DH is not done until after Cp is determined the second week. Do not save space in your notebook for A, instead collect data sequentially and put the location of data in your table of contents. Tape graphs into your notebook as you print them.

EQUIPMENT AND MATERIALS:

1.      Temperature probe connected to computer via analog to digital interface box

2.      Vernier Data Logger software

3.      Calorimeter (two nested styrofoam cups and lid labeled with your bin number)

4.      Thermometer

5.      Stirrer-hot plate and teflon stir bar

6.      Spatula and electronic balance

7.      25 mL graduated cylinder (or 25 mL pump dispenser)

8.      Various sizes of beakers and erlenmeyer flasks

9.      10 mL volumetric pipet or 10 mL pipettor

10.  Wash bottle filled with distilled water

CHEMICALS:

1.      3.0M Hydrochloric acid, HCl ((3033) )- about 200mL needed

2.      ~5.0M Sodium Hydroxide, NaOH ((3034)) - see carboy for exact concentration, about 30 mL needed

3.      Magnesium turnings, Mg ((0220)) - about 0.20g needed

4.      Magnesium oxide, MgO ((1011)) - about 1.0 g needed

5.      Unknown B-xxxx in your unknown packet (1012) - need at least 2-3 grams

SAFETY CONCERNS:

Risk Assessment: Moderate to high due to corrosive liquids.

1.      The HCl and NaOH are corrosive. The unknown and MgO are mildly corrosive and some are powders so avoid contact with solids and dust. Avoid contact, wear eye protection at all times when working with these chemicals, and wash hands after handling them.  Do not rub your eyes when using these chemicals. Any small spills should be cleaned up immediately with a damp sponge.

2.      Any contact with HCl or NaOH should be rinsed for 15 minutes with water.

3.      All solid and liquid chemical waste should be disposed of in the “Corrosive Liquids” container.

EXPERIMENTAL PROCEDURE:

First Lab Period (Second Lab Period for Experiment A) The Determination of DH of Mg with HCl and MgO with HCl

Calibration of the Temperature Probe:

n  Start up procedure

n  Select “Start”, “Programs”, “Chemistry Applications” and then click on “CHM152L-B Calorimetry”.

If these icons are not present, click on the icon, "Vernier Programs" and then “Data Logger”. Change the time to 400 seconds.

n  Calibration procedure

n  Fill one beaker with a slurry of ice.  The beaker need not be clean. (Fill with ice and put the minimum amount of water to make ice slush.)

n  Fill another beaker with hot tap water.  The beaker need not be clean.

n  Put the temperature probe tip and thermometer bulb together so they touch and place them into the hot beaker and let sit for one minute. The temperature can be read below the graph. You do not need to click on the collect button. Record the temperatures. If the temperatures are not within ± 2°C, see your TA.

n  Put the temperature probe and thermometer into the cold beaker and let sit for one minute. You do not need to push the collect button.  Record the temperatures. If the temperatures are not within ± 2°C, see your TA.

n  Check the calibration again at the beginning of the second week labalso.

NOTE:  Use the same calorimeter and computer every weekslab period if possible.

Mg Reaction. Determining the Enthalpy for Reaction (DH) of Mg with HCl (H+):

Mg(s) + 2 H+(aq) ®  Mg2+ + H2(g)

NOTE:  Only one trial is needed for  reactionfor reaction 1.  Avoid adding extra heat from hands, hot plate-stirrer; make sure the hot plate is turned off.

n  Clean a telfon stir bar, 50 mLl beaker, 25 and 50 graduated cylinder and wash bottle.

n  Add 25.0 ml mL of 3.0 M HCl (use pump dispenser, ok to check volume with 25mL graduated cylinder) and 35.0 ml mL of distilled water to the calorimeter.

n  Put the temperature probe in through the lid and place the lid on the calorimeter. Secure the probe with a clamp so that its tip is in the water but away from the stir bar and off the bottom of the calorimeter.

n  Gently stirStir the solution with a teflon stir bar. (Do not heat.  Set stirrer so the solution is mixed vigorously but slow enough so that it is not splashed.)

n  Put between 0.15 to 0.20 g (not 1.5) of Mg metal turnings into a clean, dry 50 ml mL beaker. Record both the beaker’s mass and the mass of the beaker and Mg.

n  Click the "Collect" button at the top of your screen to start graphing the temperature.  Do not add Mg turnings yet.

n  After about one minute, add the metal Mg turnings without removing the temperature probe from the solution.  (Crack the lid open, add Mg (s), and then close the lid, it is ok if residual Mg sticks to the inside of the beaker since you will reweigh it later to see how much Mg was added to the calorimeter).

n  After about 10 seconds, briefly swirl the solution (holding the cup in your hand) to dissolve any Mg left on the sides of the calorimeter.  Do this again after 60 seconds.

n  Reweigh and record the mass of the beaker that contained the Mg turnings. Subtract this mass from the mass of the beaker and Mg to get the mass of Mg used. Why should you measure the mass of Mg this way?

n  Continue graphing data until a linear line (part 3 in fig B-2) is made.  (At about 350 to 400 seconds.)

n  Adjust graph scale and then do a linear fit (refer to Figure B-3). Generate line 2 by selecting the linear part of the graph and doing a linear fit. Use the analyze tool to find the starting time of the reaction, that would be x in y=mx+b to determine the final temperature. Using m, b, and x calculate the final temperature (Tf). The next page graphically explains this process.

n  Point the cursor at the flat part of the graph before the reaction starts to determine Ti and subtract this from Tf to determine the change in temperature, Tf - Ti = DT. Calculate these values for every graph done for experiment B.

n  Label the graph by clicking on the graph title.  Edit the title to include your last name, experiment title-reaction Mg with HCl, date, section letter, and exact mass of Mg used.

n  Save the graph on your “Z” drive, my documents, or a thumb drive using a logical file name and print it.

n  Write the values for Tf, Ti, and DT directly on this printed graph.

n  Clean and dry the calorimeter, temperature probe, and the 50 ml beaker.

Explanation of the Graph:  Below is a general temperature vs. time graph representative for all reactions trials done for this experiment.  (Figure B-2)  It is divided into three parts.

Note: the dark band between Part 2 and Part 3 is a transition area that is not usable for data analysis. The linear regression is done on part 3 to the right of this dark area!

Figure B-2.  A general temperature vs. time graph

Part 1.  This is the initial temperature.  Only one reactant is in the solution and so our reaction is not happening.  (For example, in reaction 1, only 25.0 ml of HCl and 35.0 ml of water are in the solution.)  Between part 1 and part 2 the reactants are mixed together.

Part 2.  The temperature is changing rapidly.  Both reactants are now in the solution and are reacting to give off heat.  (For example, in reaction 1, this is because the Mg turnings are added to the solution.)  Somewhere within the blocked out region the reaction stops.

Part 3.  The reaction has already stopped.  Since the calorimeter isn’t a perfect insulator, heat is lost to the environment and, as a result, the temperature decreases.  The temperature should be constantly decreasing.  (For example, in reaction 1, all of the Mg has been converted to Mg2+.)

How do you get DT?  DT is equal to the extrapolated final temperature minus the solutions initial temperature.  So DT = Tf – Ti.  Below is the same general graph from figure 1, but it has been extrapolated to find Tf (Figure 2). Tf can also be determined by doing a linear regression on the linear, right hand side of the curve to determine the slope and y-intercept of line 2 and then solving for the temperature using the time at the start of the reaction (line 1).

Figure B-3.  Extrapolation of a Temperature vs. Time Graph to Find DT.

Line 1.  This line represents the time when the reactants were mixed.

Line 2.  This line helps us model what the final temperature would be if the reaction and temperature measurement were instantaneous. It compensates for the heat lost from the calorimeter so that we can estimate the final temperature if the reaction and temperature measurement were instantaneous. This line is important because it compensates for heat lost to the environment while temperature is measured during and after the reaction.

Line 3.  This line is drawn at a right angle to line 1 to intersect the point where lines 1 and 2 meet. It is there to help read the final temperature, Tf, at the y-axis.

Calculations:

IMPORTANT:  For nearly all calculations in this manual the value you are calculating will be in bold print.

In this reaction you were trying to find the DH for reaction of Mg with HCl.  Unfortunately all of the calculations cannot be done until the Cp in equation B-1 is found in reaction 3.  The limiting reagent is Mg, so find the moles of Mg.

B-4

where g Mg is grams of Mg and MM of Mg is the molar mass of Mg (g/mol). Now n is found. Find DT by extrapolation on the printed graph as shown in figure B-3.  Draw all lines with a pencil and a ruler.  Label Ti (initial), Tf (final), and DT on your graph.

MgO Reaction.  Determination of the Enthalpy for Reaction of MgO with HCl:

MgO(s) + 2 H+ (aq) ® Mg2+(aq) + H2O(l)

NOTE:  Only one trial is needed for reaction 2. Never click on “New Graph” to start a new trial, instead click on “Collect” when you want to start a new trial.

n  Add 25.0 ml mL of 3.0 M HCl and 35.0 ml mL of distilled water to the calorimeter.

n  Clean the temperature probe by thoroughly spraying with your wash bottle into a large waste beaker.

n  Put the temperature probe in the calorimeter as was done before and vigorously stir the solution (but don't splash) with a Teflon stir bar.  (Do not heat)

n  Put between 1.0 to 1.2 g of MgO powder into a clean, dry 50 ml mL beaker.  Record the mass of both the beaker and the beaker with MgO.

n  Click the “Collect” button at the top of the screen to start graphing the temperature.  Do not add MgO powder yet.

n  After about one minute, add the white powder MgO without removing the temperature probe from the solution.  (Crack the lid open, add MgO(s), and then close the lid. Its ok if a residual amount of powder remains in the beaker since it will be reweighed later to determine the amount transferred to the calorimeter).

n  After about 10 seconds, swirl the solution to dissolve any MgO left on the sides of the calorimeter holding it with your hands. Do this again after 60 sec.

n  Reweigh the beaker (with traces of MgO powder not transferred) and subtract this from the mass of the beaker and MgO to determine the actual amount of MgO transferred to the calorimeter.

n  Continue graphing data until a linear line (part 3 of figure B-2) is made, then click on Stop at the top of the screen.  (at about 350 to 400 seconds.)

n  Adjust scale of graph, do a linear fit, and find the DT as was done for Mg rxn.

n  Label the graph by clicking on the graph title.  Enter in last name, experiment title-Reaction MgO + HCl, date, section letter, and mass of MgO used.

n  Save the trial on your “Z” drive, my documents, or a thumb drive and print it.

n  Clean and rinse all glassware, the calorimeter and temperature probe.

n  Before leaving, trim graphs to size and tape into the notebook, and have TA sign and date notebook.

Calculations:

These are the same calculations as described for reaction 1.  The calculations are now for the reaction of MgO (s) with HCl (aq). The limiting reagent is MgO so find the moles of MgO.

B-5

Now n (mol of MgO) is found. Extrapolate your printed graph to find DT.

Second Lab Period  – Determination of the Cp and DHrxn of an unknown with HCl

Check the calibration of temperature probe and calibrate if needed.

NaOH + HCl Reaction.  Determination of Cp:

NaOH (aq) + HCl (aq) ® H2O (l) + NaCl (aq)

NOTE:  At least two trials need to be done. Avoid adding extra heat from your hands or hot plate.

n  Clean a 25 mLl graduated cylinder, 10 ml mL volumetric pipet, a spatula, a 50 mL beaker, and a wash bottle.

n  Add 25.0 ml mL distilled water and 10.00 ml mL of NaOH to the calorimeter and measure its temperature.

Remember:  10.00 mLl is  more precise that 25.0 mLl.  What kind of glassware should you use for this?  ­­­__________________________________.

n  Measure out 25.0 mLl of 3.0 M HCl and adjust its temperature so that it is within ±0.5°C of the NaOH solution. Do not get any of the NaOH solution in the HCl solution. Rinse and dry your thermometer between solutions.

n  Put the temperature probe in the calorimeter and stir the solution with a teflon stir bar.  (Do not heat)

n  Click the Collect button at the top of the screen to start graphing the temperature probe.  Do not add HCl solution yet.

n  After about one minute, add the 3.0 M HCl without removing the temperature probe from the solution.  (Crack the lid open, add HCl(aq), and close the lid.)

n  After about 10 seconds, briefly swirl the solution.

n  Continue graphing data until a linear line (Figure B-2: see part 3 of this figure) is made.  (At about 350 to 400 seconds.)

n  Adjust scale of graph, do a linear fit, and find the DT as was done before.

n  Label the graph by clicking on the graph title as before.  Enter your last name, experiment title, date, section letter, and NaOH with HCl, Trial 1 or 2, and molarity of NaOH.

n  Save the run on your “Z” drive, my documents, or a thumb drive and print it.

n  Clean and dry the calorimeter , and temperature probe, and 50 ml beaker.

n  Repeat this once.

Calculations:

The calibration of the calorimeter is now complete!  Cp can now be calculated and used for all other calculations! First solve for the moles of NaOH (limiting reagent).

(M of NaOH) * (L of NaOH) = mol NaOH = n

Where M is molarity (exact molarity on carboy) and L is liters of NaOH.  (You need to convert from mL). Now n (mol of NaOH) is found.  Plug this into equation B-3. Also the enthalpy of neutralization of a strong base by a strong acid is a constant –55.90 kJ/mol at 25°C.  This is shown by reaction B-2.

H+ (aq) + OH- (aq) ® H2O (l);        DH = -55.90 (kJ/mol) at 25°C                           B-2

With this value, the moles of limiting reactant (n), and after determining DT from your graphs by extrapolation, the equation B-3 becomes a simple plug and chug.

B-3

Repeat this calculation for the second trial and use the average Cp value for the calculation of DH for both Mg and MgO. If the two Cp values differ by more than ±0.02 you may want to run a third trial to determine more a precise value. Calculate the median of the Cp values.

Now the Enthalpy for Reaction (DH) of Mg with HCl can be calculated using theDT and moles of Mg previously calculated and plugging them into the equation below using the Cp determined above.

B-1

Now repeat this calculation using the data for MgO to determine the DH of MgO.

Calculation of the Molar Enthalpy of Formation of Magnesium Oxide, DHf:

Using Hess’ Law to combine the calculated DH values and a known DH (-285.8 kJ/rxn), the

DHf (MgO) can be calculated.

Mg(s) + 2 H+(aq)       ®    Mg2+(aq) + H2 (g)                             DH= ___________

MgO(s) + 2 H+(aq)  ®    Mg2+ (aq) + H2O(l)                          DH= ­­­­­­­­­­­­­­­­­­­___________

H2 (g) + 1/2O2 (g)     ®    H2O (l)                                              DH=­­ -285.8 (kJ/mol)

---------------------------------------------------------------------------------

Mg (s) + 1/2 O2 (g)    ®    MgO (s)                                          DHf= ___________

Example:  Let say the DHrxn for the below equations was determined experimentally.

A + 2B ® C + D                  DHrxn= -2 kJ/mol

2B ® 2C + 3D                     DHrxn = -5 kJ/mol

3A + 4B ® C                        DHf  (C) =  ________

When the first equation is multiplied by three, and the second equation is flipped around, this equation becomes solvable.

3A + 6B ® 3C + 3D           DHrxn= 3(-2 kJ/mol)    = -6 kJ/mol

2C + 3D  ® 2B                    DHrxn = -1(-5 kJ/mol) =  5 kJ/mol

3A + 4B ® C                        DHf  (C) = (-6) + (5)      = -1 kJ/mol

Notice how chemical D and some of B and C cancel out because they are on opposite sides of the first and second chemical equation.

Reaction 4.  Enthalpy for the Reaction of HCl with an Unknown Metal Oxide.

Note: At least 2 trials also need to be done for reaction of HCl with the unknown. Avoid adding extra heat from hands, hot plate, etc.

n  Add 25.0 ml mL of 3.0 M HCl and 35.0 mL of distilled water to the calorimeter.

n  Clean the temperature probe by thoroughly spraying with your wash bottle into a 600 mL beaker.

n  Put the temperature probe in the calorimeter and vigorously stir the solution with a teflon stir bar but avoid splashing.  (make sure the heat is off)

n  Before using your unknown, make sure it is a powder. If it is clumpy, grind it up in a clean, dry mortar and pestle. Put between 1.0 to 1.2 g of unknown powder into a clean, dry 50 mL beaker.  Record the mass of the beaker and its contents.

n  Click the "Collect" button at the top of your screen.  Your temperature probe will display the data it is collecting on the graph.  Do not add unknown yet.

n  After about one minute, add the white powder unknown without removing the temperature probe from the solution.  (Crack the lid open, add unknown powder, and then close the lid. It is ok if a residual amount of powder remains in the beaker since it will be reweighed later to determine the amount transferred to the calorimeter)

n  After about 10 seconds, briefly swirl the solution to dissolve any unknown left on the sides of the calorimeter.  Do this again after 60 seconds.

n  Reweigh the beaker with unknown powder not transferred into the calorimeter.  Continue graphing data until a linear line (part 3 in fig B-1) is made and then click on "Stop".  (At about 350 to 400 seconds.)

n  Adjust scale of graph, do a linear fit, and find the DT as was done before.

n  Label the graph by clicking on the graph title.  Enter in last name, experiment title, date, section letter, Unknown with HCl, Trial 1 or 2, & mass unknown.

n  Save the run on your “Z” drive, my documents, or a thumb drive and print it.

n  Do a second trial after cleaning the cup and probe.

n  Clean and rinse all glassware and the temperature probe.

n  Before leaving, trim graphs to size and tape into the notebook, and have TA sign and date notebook.

Calculations:

The unknown metal oxide is assumed to have a formula weight of 120.0 g/mol.  These are the same calculations as described in reaction 1.  The calculations are now for the reaction of unknown (unk(s)) with HCl (aq).The unknown is the limiting reagent.

(g unk) /  (120.0 g/mol unk ) = mol unk = n

Now use the value for n (mol of unk), the DT determined from the graph, the Cp and the equation below to calculate the enthalpy of reaction for the unknown with HCl:

B-1

Calculate this value for each trial and report the median value.

Post Lab Work  (In your lab notebook):

Calculate the range and percent range for your Cp and DHrxn for your unknown. Use the calculation check to make sure there are no errors in your calculations.  When doing the calculation check, be aware that the heat of reaction is the same as enthalpy of reaction and that these values are negative numbers for exothermic reactions. Be sure to get a printout of the calculation check. Fill out the report sheet for experiment B and staple the printout of the calculation check to it. Turn the report sheet in to your TA or the prep stockroom window (216 in bld 17) for grading.  The report sheet for the unknown must be turned in by the next lab period.

Compare your value to the published value for DH°f , the Standard Molar Enthalpy of Formation for MgO(s) and calculate the error and relative error as defined in the introduction. Comment on these values and relate these errors to measurements made while completing the experimental procedure.

CALORIMETRY - EXPERIMENT B

### CHM 152L REPORT SHEET FOR UNKNOWN B-XXXX

STUDENT'S NAME_______________________ ID#_______________ Dana ID____

SECTION LETTER____ LOCKER #_____ DATE_________ UNKNOWN # B-_____

TEACHING ASSISTANT____________________ INSTRUCTOR________________

This report sheet should be turned in at the chemistry stockroom.  Do not write the hazard code, which has the form HC-xxxx, for the unknown number.  The unknown number can be found on the top of vial label containing the unknown in the format B-xxxx. The unknown number can also be found in the section blue book.

I.  Calibration of Calorimeter - Identification # or locker #____

Cp Values   ____________    ______________

Mean or Average Cp   ____________ kJ/oC  (use median if more than two trials)

II.  Enthalpy  or Heat of Reaction:  Mg(s) + 2 H+(aq)  ®  Mg2+(aq) + H2(g)

____________ kJ/mol

III.  Enthalpy or Heat of Reaction: MgO(s) + 2 H+(aq)  ®  Mg2+(aq) + H2O(l)

____________ kJ/mol

IV.   Enthalpy or Heat of Formation MgO __________ kJ/mol

V.    Enthalpy or Heat of Reaction: Unknown Oxide + HCl

Values For Each Trial  _________   ___________

Mean or Average Value ____________kJ/mol  For Unknown # B-_______

(use median if more than two trials)