9 Exp. 9: Calorimetry and Hess’ Law

Pre-lab:

  1. When 1.104 grams of iron metal are mixed with 26.023 grams of hydrochloric acid in a coffee cup calorimeter, the temperature rises from 25.2 °C to a maximum of 33.5 °C. The reaction that occurs is given below.

 

2 Fe (s) + 6 HCl (aq) → 2 FeCl3 (aq) + 3 H2 (g)

 

  1. a) Determine the amount of heat (in J) absorbed by the reaction mixture. Assume that the specific heat of the reaction mixture is the same as the specific heat of water. The sum of the metal and solution masses will give you the mass of the reaction solution.

 

  1. b) How much heat (in J) was released by the reaction that occurred?

 

  1. c) Is this reaction exothermic or endothermic? Is ∆Hreaction positive or negative?
  2. d) Under constant pressure conditions (as used in this experiment), the heat released by the reaction equals the reaction enthalpy, qreleased = ∆Hreaction. Determine ∆Hreaction in Joules per gram of metal used (J/g).

 

  1. e) Determine ∆Hreaction in kilojoules per mole of metal used (kJ/mol).

 

  1. f) Determine ∆Hreaction in kilojoules per mole of reaction for the balanced reaction equation provided (kJ/mol reaction).
  2. Consider the following three reactions:

2 Fe (s) + 6 HCl (aq) → 2 FeCl3 (aq) + 3 H2 (g)                                                                   ∆HA

Fe2O3 (s) + 6 HCl (aq) → 2 FeCl3 (aq) + 3 H2O (l)                                                              ∆HB

2 H2 (g) + O2 (g) → 2 H2O (l)                                                                                                ∆HC

Show how these equations must be summed together according to Hess’s Law to determine ∆H for the combustion of iron (target equation shown below). Also show clearly how the ∆H values of each of the three reactions must be manipulated to determine the enthalpy of combustion of iron.

4 Fe (s) + 3 O2 (g) → 2 Fe2O3 (s)                                                                                           ∆H =?

  1. Using tabulated ∆Hf° values in the text appendix two determine the enthalpy change (in kJ) that occurs during the formation of water from its elements:

2 H2 (g) + O2 (g) → 2 H2O (l)                                                                                    ∆H =?

Note that this value (and the equation) will be used in your data analysis for this lab.

Objectives:

 

The objectives of this laboratory are as follows:

  • To experimentally measure the ∆H values of two reactions using the technique of constant pressure calorimetry.
  • To apply these ∆H values in a Hess’s Law calculation to determine the enthalpy of
    combustion of a metal

Introduction:

 

As you observed in lab two, The Formation of an Oxide, the combustion of a metal in oxygen produces the corresponding metal oxide as the only product. Such reactions are exothermic and release heat. For example, the combustion of iron releases 1651 kJ of heat energy for every four moles of iron burned:

(Reaction 1)                   4 Fe (s) + 3 O2 (g) → 2 Fe2O3 (s)          ∆H1 = -1651 kJ

Since it is difficult to measure the enthalpy of combustion of a metal directly, in this lab it will be
determined indirectly by applying Hess’s Law of Heat Summation. Hess’s Law states that the
enthalpy change of an overall process is equal to the sum of the enthalpy changes of its
individual steps.

 

Hess’s Law Example:  Determine ∆H for the target reaction 2 NO2 (g) + ½ O2 (g) → N2O5 (g) given the following information,

Reaction A    N2O5 (g) → 2 NO (g) + ³/2 O2 (g)            ∆HA = 223.7 kJ/molrxn

Reaction B    NO2 (g) → NO (g) + ½ O2 (g)                               ∆HB = -57.1 kJ/molrxn

Solution:  Reactions A and B must be carefully manipulated before they can be summed to produce the target reaction. Reaction A must be reversed, causing a sign change to ∆HA. Reaction B must be multiplied by a factor of 2, causing ∆HB to be multiplied by 2. Only then will they yield the target equation when added together:

2 NO (g) + ³/2 O2 (g) → N2O5 (g)                                ∆H = −(223.7) = -223.7 kJ/molrxn

+          2 NO2 (g) → 2 NO (g) + O2 (g)                                                        ∆H = 2 x (-57.1) = -114.2 kJ/mol rxn

2 NO2 (g) + ½ O2 (g) → N2O5 (g)                                              Target

Thus, ∆HTarget = -223.7 + (-114.2) = -337.9 kJ/molrxn

 

In order to use Hess’s Law to find the heat of combustion of a metal, it is first necessary to

find reaction enthalpies (∆H values) for equations that can be summed together appropriately. To accomplish this, two reactions will be studied in this lab. In one reaction, a given metal will react with hydrochloric acid producing hydrogen and the metal chloride. In the other reaction, the corresponding metal oxide will react with hydrochloric acid producing water and the metal chloride. For example, the reactions involving iron and iron (III) oxide are as follows:

 

(Reaction 2)                      2 Fe (s) + 6 HCl (aq) → 2 FeCl3 (aq) + 3 H2 (g)                   ∆H2

 

(Reaction 3)                      Fe2O3 (s) + 6 HCl (aq) → 2 FeCl3 (aq) + 3 H2O (l)              ∆H3

Since both reactions are exothermic, the heat released (q) will be absorbed into the surrounding reaction mixture. As long as the reactions are performed in an insulated container (such as a coffee cup calorimeter), there will be negligible heat exchange with the container walls or outside air. By monitoring the temperature of the reaction mixture when specific quantities of reactants are used, the amount of heat (in J) released by these reactions can be determined by applying the equation:

 

heat released by reaction (−qreaction) = heat absorbed by reaction mixture

 

(Equation 1)                       (+qmixture)=  (m x Cs x ∆T)mixture

Here m is the total mass of the reaction mixture (in g), ∆T is the maximum temperature change
that occurs during the reaction (in °C), and Cs is the specific heat of the mixture (in
J/g•°C). Note that since the reactions occur in aqueous solution, it is reasonable to substitute
the specific heat of water (Cs= 4.184 J/g•°C) for the specific heat of the mixture.

Recall that at constant pressure (the conditions of this experiment), the heat released by the reaction equals the reaction enthalpy:

 

               qP = ∆H

Since the heat released by each reaction is proportional to the amount of metal/metal oxide used, ∆H2 and ∆H3 can be easily calculated per gram or mole of metal/metal oxide used.

 

It should be noted that reactions (2) and (3) by themselves still cannot be summed to produce Reaction (1). Another reaction is required:

 

     (Reaction 4)                      2 H2 (g) + O2 (g) → 2 H2O (l)                                                ∆H4

H for this reaction (the formation of water from its elements) must be obtained from tabulated
thermodynamic data, ∆Hf° appendix two in the textbook. Finally, the reactions (2), (3) and (4) and their enthalpies may be summed together according to Hess’s Law to determine the enthalpy of combustion of the given metal (1).

 

Safety:

  • Hydrogen gas will be generated during this experiment. Hydrogen is flammable; keep all heat and flames away from your reaction vessel.
  • Hydrochloric acid (HCl) is extremely caustic. Wear splash goggles, gowns, and gloves. If HCl comes into contact with your skin or eyes wash immediately under running water for at least ten minutes. The sodium bicarbonate solution by the sinks may be used to neutralize and clean up any acid spills.

 

Waste/Housekeeping:

Dispose of all waste in the labeled container located in Hood #1.

Materials and Equipment:

Mg (s), MgO (s), Zn (s), ZnO (s), (Al (s), Al2O3 (s) not used 2020), 1M HCl (aq), 6M HCl (aq), coffee cup calorimeter with lid*, logger pro temperature probe, glass stirring rod, 100-mL graduated cylinder, 50-mL beaker, utility clamp, stand, electronic balance, and wash bottle.

Temperature Reading

Instead of a thermometer, you may use a temperature probe and LoggerPro® software to directly monitor temperature changes over time. Detailed instructions for setting up this system will be provided at the beginning of your lab period. Please note that your experimental procedure will still be the same regardless of the method used to monitor temperature.

 

The Heat of Combustion of a Metal/Metal Oxide:

  1. You will be assigned a specific metal/metal oxide pair to investigate by your instructor.
    Record their identities on your report form. Note that you will perform the following procedure for a total of four times, twice with the metal, then twice with the metal oxide.
  2. The table below indicates the quantities of reactants to be used for each metal/metal oxide
    combination. Note that the reactions involving Zn and Al require the concentrated 6M acid.

 

Table 1: Recommended Mass of Metal and Metal Oxide
Mg / MgO Zn / ZnO Al / Al2O3 (not used 2020)
0.15 g Mg, 25 mL 1M HCl

0.25 g MgO, 25 mL 1M HCl

0.40 g Zn, 25 mL 6M HCl

0.60 g ZnO, 25 mL 6M HCl

0.15 g Al, 25 mL 6M HCl

0.75 g Al2O3, 25 mL 6M HCl

 

  1. Use an electronic balance to weigh your empty, dry calorimeter (the two nested Styrofoam®
    cups). Remove it from the balance, then pour approximately 25-mL of HCl (aq) into it and
    weigh it again. Record these masses on your report (the difference is the mass of HCl (aq)
    used).
  2. Now weigh the recommended mass of your assigned metal using a tared or recorded beaker or weigh boat. Do not place the metal or oxide directly on the balance.
  3. Assemble your equipment as shown in the figure below. A temperature probe will replace your thermometer. You may use a stir rod, or a swirling method. The temperature probe and the stirring rod must be inserted through the holes in the calorimeter lid. The tip of the temperature probe should be immersed in the acid, but not touch the bottom of the calorimeter. Clamp the thermometer in place using a utility clamp.

 

  1. Logger Pro Data Collection:

 

  1. Connect your temperature probe to the Lab Quest interface and open the Logger Pro program. The program will display a temperature reading.
  2. Under the Experiment tab, choose Data Collection and increase the time of data collection. 350 seconds should suffice. *If time runs out during your experiment, re-push Collect and Append to Latest in order to keep your data.
  3. With the probe tip in the acid solution, push Collect to get a baseline temperature reading of the solution at thermal equilibrium.
  4. Once you have a baseline, add your metal or oxide, and quickly replace the lid. Keep swirling or stirring until the reaction is complete. You will know the reaction is finished when the temperature has reached a maximum and begins to decrease.
  5. When you are sure that the reaction is finished, push Stop.
  6. Under the Analyse drop down, select the Statistics tab. This will give you the minimum (baseline) and maximum temperatures needed for your calculations*.

*Make sure that the min and max in you Statistics data makes sense to you. If it does not, you can manually record temperature data by placing the cursor on the minimum temperature and recording the y-axis value at the bottom left of the graph.

  1. When finished, dispose of your chemical waste in the labeled container in hood one, rinse
    the calorimeter, thermometer and stirring rod thoroughly with distilled water, dry, and repeat
    the experiment again. Once you have completed both trials with the metal, perform your two
    trials using the metal oxide using the identical procedure.

Note*: Nesting your calorimeter in a medium beaker and clamping your temperature probe or thermometer will add stability. The following example data tables may help you collect your data.

Metal + HCl Reaction

 

Experimental Data

Assigned Metal:

 

Data Table 1: Collection Data (Metal)
Trial 1 Trial 2
Mass of dry, empty calorimeter (g)
Mass of calorimeter plus HCl (g)
Mass of HCl used (g)
Mass of metal used (g)
 Total mass of solution (HCl and metal) (g)
Initial (equilibrium) temperature of HCl (ºC)
Final (maximum) temperature of mixture (ºC)

 

Data Analysis:

1)  Write the balanced equation for the reaction between your assigned metal and HCl. All
balancing coefficients should be whole numbers.

 

 

 

 

 

 

 

 

2)  Complete the table below with the results of your calculations.

 

Data Table 2: Calorimetry Calculation Data (Metal)
Trial 1 Trial 2
Total mass of mixture, m
Temperature change of mixture, ∆T
Specific Heat. of mixture, s (use Cs water)

of water)

Heat of mixture, in J (*what sign should this have?)
Heat of reaction, in J (*what sign should this have?)
Hrxn in J/g of metal used
Hrxn in kJ/mol of metal used
Hrxn in kJ/mol(rxn) for the rxn as balanced in Data analysis#1(above)
Average ∆Hrxn in kJ/mol(rxn)

 

3)  Show your work for the following calculations using your Trial 1 data only:

  • Heat absorbed by mixture, in J
  • Heat released by reaction, in J
  • Hrxn in J/g of metal used
  • Hrxn in kJ/mol of metal used
  • Hrxn in kJ for reaction as balanced in (1)

4)  Is this reaction exothermic or endothermic? What is your experimental evidence supporting this? Is ∆Hrxn positive or negative?

Assigned Metal Oxide:

 

Data Table 1: Collection Data (Metal Oxide)
Trial 1 Trial 2
Mass of dry, empty calorimeter (g)
Mass of calorimeter plus HCl (g)
Mass of HCl used (g)
Mass of metal oxide used (g)
  Mass of solution (HCl and metal oxide) (g)
Initial (equilibrium) temperature of HCl (ºC)
Final (maximum) temperature of mixture (ºC)

 

Data Analysis

 

1)  Write the balanced equation for the reaction between your assigned metal oxide and HCl. All balancing coefficients should be whole numbers.

 

 

 

 

 

2)  Complete the table below with the results of your calculations.

 

 

Data Table 2: Calorimetry Calculation Data (Metal Oxide)
Trial 1 Trial 2
Total mass of mixture, m
Temperature change of mixture, ∆T
Specific Heat. of mixture, s (use Cs water)

 

Heat of mixture, in J (*what sign should this have?)
Heat of reaction, in J (*what sign should this have?)
Hrxn in J/g of metal oxide used
Hrxn in kJ/mol of metal oxide used
Hrxn in kJ/mol(rxn) for the rxn as balanced in #1(above)
Average ∆Hrxn in kJ/mol(rxn)

Note: You are not required to show your work for the calculations you performed to complete the above table.

 

 

Post Lab; Enthalpy of Combustion of a Metal:

1)  Write the balanced equation for the combustion (Metal + O2(g) –> Metal oxide) of your assigned metal. All balancing coefficients should be whole numbers.

2)  Using Hess’s Law, determine the enthalpy of combustion of your assigned metal as

balanced above. To do this, you will need the balanced thermochemical equations for the two reactions performed in this lab, i.e., the metal and HCl(aq) and the metal oxide and HCl(aq) plus the balanced thermochemical equation for the formation of water from its elements. Clearly show how these three equations (and their reaction enthalpies) must be combined to give the target combustion reaction. See the introductory examples for help. Pay close attention to the sign of ΔHrxn. *You will only need to look up one literature value in this step; this is the enthalpy of formation of water used in the third equation. Be sure to note the number of moles water produced. Why don’t you need the enthalpy of formation (Δ°Hf ) for H2(g) or O2(g)? Provide a clear calculation.

 

3)  Calculate the theoretical value for the enthalpy of reaction, (Δ°Hrxn) for combustion of your assigned metal using the tabulated ∆Hf° values from the text (Burdge Overby: Atoms First) appendix two. *You will only need to calculate the enthalpy of reaction value for your overall (summed) reaction using Hess’ law. This is your assigned metal combining with oxygen to form a metal oxide, see equation . You will only need to look up one literature value in this step; this is the enthalpy of formation (Δ°Hf )  of the metal oxide produced in the overall or target reaction. Why don’t you need the enthalpy of formation (Δ°Hf ) for the solid metal or oxygen gas? Provide a clear calculation.

4)  Determine the percent error in your experimentally determined value for the enthalpy of combustion. See your student lab appendix for formulae and explanation. Discuss the significance of the error and theorize what may have caused the error.

5)  Provide a reaction and practical application for a chemical process that is:

  1. a) Endothermic
  2. b) Exothermic

 

 

 

Your Report:

  1. Complete the pre-lab and take the pre-lab quiz in Blackboard before attending lab.
  2. Use a lab notebook to take notes and make observations.
  3. Create an informal lab document to turn in at the beginning of the next lab. Your lab report should have your name, data arranged in graphs or tables where applicable, conclusive data such as enthalpies of reaction. Relevant sample calculations, including balanced reactions. You must clearly show heat calculations, enthalpy calculations and Hess’ law manipulations for full credit. These should appear in the results and discussion section. Think about the main point of the lab, or the results that you worked for and be sure to include it or them.
  4. Include any Data Analysis and associated questions.
  5. Numbered responses to any post lab question. You do not need to re-write the questions.
  6. Submit your document before your next lab appointment under the assignment tab on your laboratory Blackboard shell.

 

References:

Adapted from: Chem 11 Experiments, Santa Monica College http://www.smc.edu/AcademicPrograms/PhysicalSciences/Chemistry/Lab-Manual/Pages/Chem-11-Experiments.aspx (accessed Jul 04, 2020).