D. F. Nachman
Abstract: An ice calorimeter was used to study the reaction of magnesium metal and 1.00M sulfuric acid solution: Mg(s) + H2SO4(aq) →MgSO4(aq) + H2(g). We found the experimental molar enthalpy of reaction to be ΔH = –355 ± 17 kJ/mol at 0°C, 24% lower than the textbook value of ΔH° = –466.9 kJ/mol, reported at 25°C.
Whether a chemical reaction occurs spontaneously or is driven by an outside force, it almost always exchanges energy with the surroundings. Energy exchange can occur as work or as heat flow. When a reaction occurs under constant-pressure conditions, we call the heat portion of the energy exchanged the enthalpy change of the reaction. Because the work can be hard to measure accurately, and because the work often represents only a minor fraction of the total energy change, we often focus only on the heat form of energy exchange. Measuring the enthalpy change of a chemical reaction gives information about the relative bonding energies of the products and reactants. If the product substances experience stronger bonding forces than the reactants do–if the potential energy of the products is lower than that of the reactants–the reaction process will release the energy difference as heat and/or work, and we call it an exothermic reaction.
We used short strips of magnesium ribbon (cleaned with steel wool to remove most of the surface oxide) and 1.00M sulfuric acid solution provided by the stockroom. We weighed the magnesium on the laboratory balance, and measured the acid solution with a Class A 5-mL transfer pipette. To check the temperatures, we used liquid-in-glass emergent stem thermometers with a range of –10°C to +150°C. The wall clock was used to monitor elapsed time. The major piece of equipment used is the ice calorimeter, first described in 1780 by Antoine Lavoisier and Simon de la Place1, and modified for use in this course2. The device captured the heat released by the reaction under study, converting solid ice to liquid water. By measuring the amount of water produced, we calculated the amount of heat captured in the calorimeter. To avoid heat leaks between the reaction system and the surroundings, we surrounded the calorimeter with crushed ice. Contact with the atmosphere could not be completely avoided, because constant pressure was desired even though the reaction produces hydrogen gas. We took care to remove air from the calorimeter chamber, so that the only volume changes would be due to changes in the amounts of solid ice and liquid water. We monitored the changes in volume of the ice/water mixture by observing the level of liquid in the inverted volumetric pipette. To avoid adding heat to the calorimeter during the addition of reactants, both Mg and H2SO4 were chilled to ice temperature before mixing. Despite these precautions, some melting of calorimeter ice did take place, as seen in the slow decrease in pipette reading before the reactants were mixed (see Figure 1. Calorimeter data).
We did not attempt to measure the temperature of the reaction mixture, as inserting a thermometer could have introduced heat to the system. Instead we measured the temperature of the surrounding ice. The time lag between marking each elapsed 30-second interval and reading the water level in the pipette was estimated at not more than 3 seconds. We did not confirm the concentration of the sulfuric acid solution by independent measurements. Data
Table 1. Ice calorimeter data
Before reaction During reaction After reaction
time (s) pipet (mL) time (s) pipet (mL) time (s) pipet (mL)
0 0.900 300 0.870 630 0.381
30 0.899 330 0.800 660 0.369
60 0.895 360 0.720 690 0.358
90 0.894 390 0.650 720 0.347
120 0.892 420 0.590 750 0.338
150 0.890 450 0.540 780 0.329
180 0.890 480 0.498 810 0.320
210 0.889 510 0.462
240 0.888 540 0.439
270 0.887 570 0.418
Linear equations on Figure 1...