Exp 4 Lab Report Joe CHM 113 (1)

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Aaron Joe CHM 113 Lab Report 2/27/24 Chemical Thermodynamics: Heat of Formation of MgO(s) Part A: Introduction: This experiment aims to determine the heat, or enthalpy, involved in the production of MgO(s). In this experiment, the enthalpy values were determined using a basic calorimeter. To calculate the amount of heat released, utilize calorimetry. Hess's law was utilized to determine the enthalpy of formation using three distinct chemical reactions in order to achieve the enthalpy values. According to Hess's law, the enthalpy of a reaction that is carried out in steps will equal the total of the enthalpy changes for each step. The three distinct equations must be applied since the air contains N2. Mg and N2 would react to generate Mg3N2. This measure could only be implemented in a pure oxygen environment, which is not possible in a laboratory setting. Procedure: 1. Set up the calorimeter. 2. Prepare 250mL of 1M HCl solution. 3. Using a graduated cylinder, transfer exactly 100mL of the dilute HCl solution into the calorimeter. 4. Clean a piece of Mg ribbon by rubbing steel wool to remove the layer of oxide on the surface. Weigh the Mg ribbon to approximately 0.25 grams and record the measurement. 5. Wind the magnesium ribbon around the loop of the stirrer. And record the initial temperature of the HCl solution.

6. Immerse the stirrer in the HCl solution and stir gently to ensure the Mg never comes above the surface of the solution. 7. Monitor the temperature rise and record the highest value reached. 8. Repeat steps 1 to 3. In step 4, use MgO instead and measure the weight to around 0.4 grams. 9. Record the initial temperature of HCl solution. Add the MgO and stir until completely dissolved. 10. Monitor and record the highest temperature reached. Equations: 2 reactions being measured: Mg(s) + 2HCL (aq) → MgCl2(aq) + H2(g) MgO(s) + 2HCL (aq) → MgCl2(aq) + H2)(g) Heat of formation of MgO: Mg(s) + 1/2O2(g) → MgO(s) q=mc(delta T) qsol=-qrxn Observations: Table of Physical Data: Mg Ribbon MgO Ribbon Weight (g) 0.334 0.403 Atomic Weight (g/mol) 24.035g/mol 41.035g/mol Number of Moles (mol) 0.013896 0.010048 Table of Temperature Changes and Heat Released:

For the MgO Ribbon reaction, we may have not given it enough time to heat up, and therefore obtained a very small temperature change. Mg Ribbon MgO Ribbon Initial Temperature (C) 23 C 23 C Final Temperature (C) 34 C 26 C Temperature Change (C) 11 C 3 C Heat given off (kcal) 1.1 kcal .3 kcal Enthalpy of Reaction (kcal/mol) -79.1595 kcal/mol -29.857 kcal/mol Since both ribbons display a rise in temperature, we can conclude that both reactions are exothermic. Calculations: The amount of heat released, q(heat)=mc(delta T), divided by the number of moles, yields the enthalpy of reaction. For Mg Ribbon: q(H2O)=(1.0cal/gC) x (100g) x 11.0 C= 1100 cal qsol=-qrxn 1100cal=-1100cal -1.1kcal Enthalpy= -1.1kcal/0.013896 moles = -79.1595 For MgO, we must utilize Hess’ Law in order to calculate the heat of formation : -79.1595kcal/mol - ( -29.857kcal/mol) - (-68.3kcal/mol) = -117.6025 kcal/mole MgO Enthalpy of Formation in (kJ/mol): -119.6025 kcal/mole x 4.184 kJ= -492.047 kJ/mol

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