Ice caloremetry Lab

1 Thermochemistry: An Ice Calorimeter Determination of Reaction Enthalpy Michael Orozco, Brayden Raynolds Professor Zachary Galasinski February 15, 2024 Abstract: Focusing on the first law of thermodynamics and its application in calorimetry, calorimetry is a vital process in measuring internal energy transfers during chemical reactions, typically monitors temperature changes. However, this study employs an ice calorimeter to track volume changes, termed Isothermal heat transfer. The experiment involves an exothermic reaction between magnesium and sulfuric acid, resulting in ice melting within the calorimeter. The melted ice's volume change is calculated using the densities of water and ice at zero degrees Celsius. Results indicate a significant decrease in volume due to ice melting, and calculations reveal the heat of fusion energy and molar enthalpy of the reaction. The limiting reactant is identified as sulfuric acid, with the experimental value of molar enthalpy showing a 10% deviation from the theoretical value. Possible sources of error, including heat loss and premature reaction termination, are discussed.

2 Introduction A fundamental idea of energy is that it may be transferred as either a form of heat or work between a system and its surroundings. Energy also can neither be created nor destroyed; this is known as the first law of thermodynamics and is an important concept for understanding how calorimetry works. Calorimetry is a common process used to measure how the internal energy transfers in and out of a certain system of interest, with the system usually being a chemical reaction. To observe calorimetry processes, chemists use a calorimeter to measure heat-induced temperature changes caused by the system and its surroundings. In this Lab We utilized an ice calorimeter to measure the change in volume throughout the reaction rather than the temperature. This is known as Isothermal heat transfer. The ice inside the calorimeter will melt because of the exothermic reaction of Mg(s) + H2SO4(aq) → MgSO4(aq) + H2(g). The amount of melted ice is then measured and utilized. Methods The lab manual procedures were followed accordingly, and no changes were made. To start off we added water and ice to a calorimeter and submerged it halfway into an ice bucket. We waited 10-15 minutes for the colorimeter apparatus to reach equilibrium temperature. We then gathered our magnesium strips which we weighed to 0.2073 grams and cleaned them with steel wool to remove the oxidated surface to allow the magnesium metal to react properly. A 1.0057 M sulfuric acid solution(H2SO4) was provided to us. We used a transfer pipette to transfer 5.00 mL of the acid to the reaction test tube and allowed it to cool for ten minutes to reach equilibrium temperature. Next, we added the Magnesium strips to the Acid in the test tube. To monitor temperature change and time as the reaction began, We used my phone as a timer and my lab partner kept an eye on the volume readings. We recorded a temperature reading every 30

4 120s 0.660 mL 150s 0.640 mL 180s 0.620 mL 210s 0.610 mL 240s 0.600 mL 270s 0.590 mL 300s 0.580 mL 330s 0.575 mL 360s 0.570 mL 390s 0.565mL 420s 0.560 mL 450s 0.560 mL 480s 0.550 mL 510s 540 550 0.550 mL 0.550 mL 0.550 mL

5 Figure 1: Calorimeter Data 0 100 200 300 400 500 600 700 800 0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1 f(x) = − 0 x + 0.93 Time (s) Pipette volume readings (mL) The results displayed in figure 1 show us that that the as the ice melted over time due to heat, the volume dropped in the ice calorimeter.

7 Since we figured out how much Ice was actually melted, we were then able to apply the result and figure out the actual heat of fusion energy in the melting of our ice, the enthalpy of fusion for ice melting at zero Celsius is 0.334 J/mol. 4.8 g Ice× 0.334 J / mol 1.000 gIce = 1.6 J To determine what the limiting reactant was in our experiment we compared the amount of magnesium and sulfuric acid: 0.20 gmg× 1 m ol mg 24.3 gmg = 8.2 molesmg 5.00 ml H 2 S 04 × 1.00 mol H 2 SO 4 1.00 ml H 2 SO 4 = 5.00 moles So, the limiting reactant was H2SO4 Sulfuric Acid. Knowing which was the limiting reactant allowed us to know find out the molar enthalpy of the reaction: 0.0016 kj 5.00 moles = 3.2 × 10 − 4 kj / mol A value of ∆ H° = -3.2 × 10 − 4 kJ/mol for the reaction of magnesium and sulfuric acid solution at was achieved in the experiment.

8 based on the experimental value -466.9 KJ/mol and our experimental value of 3.2 × 10 − 4 kJ/mol we found we had a 10% difference. Discussion In hindsight, we had expected there to be a slight difference in our experiment and our assumption was correct as the Delta H value of enthalpy was significantly smaller than expected. This could have been different due system heat loss via the stop cock, heat loss via glassware and wrong use of our timing. It could also have to do with the fact that we stopped the reaction time to early and the reaction would have kept going past the time we stopped. References 1. Gilletti, Paul. CHM152LL Lab Manual Thermochemistry: An Ice Calorimeter Determination