The Enthalpy Change of a Chemical Reaction with Magnesium

References:
Chemistry, 4th ed, by J. McMurry and R. Fay. Prentice Hall, 2004, Sections  4.9, 4.10, 8.8

Purpose

Determine the change in enthalpy ( DH°) for the reaction of magnesium metal with hydrochloric acid, i.e.,

Mg(s) + 2HCl(aq)  ®  Mg+2(aq) + H2(g) + 2Cl-1(aq)

Method:


To determine the enthalpy change for a reaction the heat released or absorbed by the reaction can be measured in a container called a calorimeter.   When an exothermic reaction occurs in a calorimeter, the heat which is released warms up the container and its contents. The contents include the reaction products and the solvent, if it is present. The amount of heat energy released into the calorimeter by the reaction (qp) is related to its change in temperature by the equation:

qp = [Cp(cal) ´ mass(cal) ´ DT] + [Cp(contents) ´ mass(contents) ´ DT]

 (1)


 

where:
 Cp(cal) is the specific heat of the calorimeter (J/g×°C),
Cp(contents) is the specific heat of the contents (J/g×°C), and
DT is the temperature change (°C)


 The experiment is carried out by mixing the reactants (Mg(s) and HCl(aq)) in the calorimeter and measuring the change in the calorimeter temperature. The specific heat and mass of the resulting solution are known. The specific heat of the calorimeter is unique for each calorimeter and must be measured. For convenience the mass and specific heat are combined into one factor (Cp(cal) ´ mass(cal)), the heat capacity. Using the measured heat capacity of the calorimeter, the temperature change, and the values for the specific heat and mass of the final solution, the total heat released can be calculated using the equation above. The heat which is released for one mole of magnesium, at constant pressure, is calculated and represents the enthalpy change, or DH°,for the reaction as it is written.


 Procedure
Note: Please work in pairs for this experiment.

Heat Capacity of Calorimeter  For this procedure, two identical calorimeters are used. Each consists of two nested Styrofoam cups, a cover, and a thermometer as shown in the figure.



 1. Immerse both thermometers into a beaker of water and read the temperature on each to 0.1oC. If they do not agree, the difference in the readings is the correction which must be applied to one of the thermometers each time a reading is recorded. For example, if both thermometers are placed in the same water and thermometer A reads 25.1 oC, and thermometer B reads 25.3oC, then to each reading of thermometer A add 0.2oC as a correction.  Alternatively, you could subtract  0.2oC from each reading of thermometer B.

 2. Using a graduated cylinder, measure 50.0 mL of room temperature tap water into one calorimeter and 50.0 mL of water which has been heated to about 45 oC into the other. At one minute intervals measure and record the temperatures of both calorimeters. At the fourth minute pour the warm water into the calorimeter containing the cold water. Read and record the temperature at minutes 5, 6 and 7. Repeat this entire procedure a second time.

 The heat capacity of the calorimeter is calculated by comparing the heat lost by the warm water to the heat gained by the cold water. Any difference between these values represents heat lost or gained by the calorimeter. To determine the temperature changes, accurately plot the temperatures of the water at each time you recorded it (use Excel) and extrapolate to the mixing time (4.0 minutes).  An example of this computation is shown below using the data from the figure.




From this data:

    heat lost by warm water (7.9°C)(50.0g)(4.18 J/g×°C) = 1651J

    heat gained by cool water (5.0°C)(50.0g)(4.18J/g×°C) = 1045J 

    heat lost to cool water calorimeter = 1651J - 1045J = 606J

    heat capacity of cool water calorimeter (606J/5.0°C) = 121 J/°C


 For every degree increase in temperature of a solution in the cold water calorimeter, 121 joules are absorbed by the calorimeter. 

Calculate the heat capacity of your calorimeter from your two sets of data and determine the average value from your two trials.



Enthalpy of the Reaction - 

Waste disposal: Pour excess HCl solution and completed reaction solutions into the drain. Put the excess magnesium into the trash bins, not in the liquid waste container.

  1. Use a balance to weigh 0.50 g of magnesium turnings (Mg) into your calibrated calorimeter.  Use a spatula to do this so that you do not touch the magnesium. Oily residue from your hands will inhibit the reaction with acid. 
  2. Using a graduated cylinder, measure out 100.0 mL of 1.0 M HCl and measure its temperature. Assume that the HCl solution, the calorimeter and the magnesium metal are all at the same temperature.  This is also the initial temperature of the calorimeter.  Pour the HCl solution into the calorimeter. While stirring gently with your thermometer, measure and record the temperature immediately after mixing and then every 15 seconds thereafter until the temperature reaches a maximum and declines for two consecutive time intervals. Be certain that you record the maximum temperature even if it occurs at mid-interval.

     Clean and dry your apparatus and repeat the experiment a second time.

    For each trial, calculate the total heat released during the reaction of 0.50 g of magnesium using Equation 1. In this equation the value of DT is the difference between the initial and maximum temperatures. The mass of the contents is calculated from the volume of HCl solution (100.0 ml) times its density (1.015 g/ml) plus the mass of the magnesium (0.50 g). The specific heat of the resulting MgCl2 solution is 3.97 J/g×°C. Use the average of your experimental values for the heat capacity of your calorimeter.

    Calculate the average heat released by the reaction (in joules) and the value of DHo for the reaction (in kilojoules per mole of magnesium).


    Sample Data and Calculation Format

    Heat capacity of calorimeter

    a) Thermometer comparison readings

thermometer A _____
thermometer B _____
correction _____

    b) Measured temperature readings for water mixing

                                   Trial 1

time (min) 1.0 2.0 3.0 4.0 5.0 6.0 7.0
warm water ___ ___ ___        
        mix ___ ___ ___
cooler water ___ ___ ___        

       Corrected temperature readings for water mixing

                                   Trial 1

time (min) 1.0 2.0 3.0 4.0 5.0 6.0 7.0
warm water ___ ___ ___        
        mix ___ ___ ___
cooler water ___ ___ ___        


      

    c) Extrapolated temperatures to mixing time (from graph)

warmer water ____
cooler water ____
mixture ____

    d) Temperature change of warm water _____


    e) Heat lost by warm water _____


    f) Temperature change of cool water _____

    g) Heat gained by cooler water _____

    h) heat gained by the calorimeter _____

    i) heat capacity of calorimeter _____


Trial 2 is, of course,  identical.

Average heat capacity of calorimeter _____


 2. Enthalpy of Reaction

 a) mass of Mg(s) ______

b) volume of 1.0 M HCl ______

c) moles of Mg ______

d) moles of HCl ______

e) initial temperature of HCl ______

f) temperatures of reaction solution

time (min)

trial 1

 trial 2
0.25

______

______

0.50

______

______

0.75

Continue taking data until two successive readings show a decrease, extending this table as necessary

______

______

 

trial 1

 trial 2
g) maximum temperature

______

______

h) temperature increase

______

______

i) total mass of solution

______

______
j) heat absorbed by calorimeter  ______ ______
k) heat absorbed by solution  ______ ______
l) total heat released by reaction ______ ______
m) heat released per mole of Mg ______ ______

 

n) average enthalpy change per mol of Mg _______