Hess’s Law Determining Enthalpy of MgO Formation
🔒 LAB RESOURCES:
Student Procedure [PDF] [DOC]
Teacher Annotated Procedure [PDF]
Complete Lab Guide [HERE]
GoogleSheet Data & Analysis [HERE]
SUMMARY/ OVERVIEW:
Students use a calorimeter to measure the enthalpy changes (ΔH) for two simple reactions involving MgO. By applying Hess's Law, which states that the total enthalpy change is independent of the path taken, these two measured ΔH values are algebraically combined to indirectly calculate the difficult-to-measure ΔH for the formation of MgO from its elements.
ESTIMATE TIME ⏰: 40-60 minutes
Preparation and Initial Measurements (20-30 minutes)
Reaction 1 (e.g., $\text{MgO} + 2\text{HCl}$) (30-45 minutes)
Reaction 2 (e.g., $\text{Mg} + 2\text{HCl}$) (30-45 minutes)
SAFETY PRECAUTIONS:
Eye Protection is Mandatory: Always wear approved safety goggles throughout the entire experiment.
Handle Hydrochloric Acid (HCl) with Care: The $\text{HCl}$ solution used is corrosive.
Dispense $\text{HCl}$ carefully. If acid contacts skin, immediately rinse the area with copious amounts of water for several minutes and notify the instructor.
Avoid inhaling the acid fumes; work in a well-ventilated area or under a fume hood if the acid concentration is high.
Weighing Precautions: Ensure the solid magnesium (Mg) and magnesium oxide (MgO) are weighed quickly, as $\text{Mg}$ can slowly react with air and $\text{MgO}$ can absorb moisture, affecting the accuracy of the initial mass.
Cleaning the Calorimeter: Ensure the calorimeter is thoroughly cleaned and dried between the two reactions to prevent contamination, which could affect the $\Delta T$ of the second reaction.
Proper Waste Disposal: The solutions contain products of the acid reactions ($\text{MgCl}_2$ and excess $\text{HCl}$). Dispose of all liquid waste in the designated acidic waste container as instructed by your teacher. Do not pour them down the sink.
Introduction
This lab is designed to illustrate and verify Hess's Law of Constant Heat Summation, a fundamental principle of thermochemistry. The experiment focuses on determining the enthalpy of formation ($\Delta H_f^\circ$) for magnesium oxide ($\text{MgO}$), a reaction that is nearly impossible to measure directly with standard laboratory calorimetry.
The formation of magnesium oxide is:
$$\text{Mg}(\text{s}) + \frac{1}{2}\text{O}_2(\text{g}) \rightarrow \text{MgO}(\text{s}) \quad \Delta H_f^\circ$$
Attempting to measure this reaction directly in a simple calorimeter is impractical because it's a vigorous, highly exothermic reaction ($\text{Mg}$ burning in air) that doesn't occur under controlled, dilute conditions.
Background
Enthalpy Change ($\Delta H$)
Enthalpy is a measure of the heat content of a system at constant pressure. The enthalpy change ($\Delta H$) represents the amount of heat released ($\Delta H < 0$, exothermic) or absorbed ($\Delta H > 0$, endothermic) during a chemical reaction. In this lab, $\Delta H$ is determined indirectly using a calorimeter and the following formula derived from the conservation of energy:
$$q_{\text{reaction}} = - (q_{\text{solution}})$$
The heat absorbed by the solution is $q_{\text{solution}} = m \cdot c \cdot \Delta T$, where $m$ is the mass of the solution, $c$ is the specific heat capacity of the solution (assumed to be that of water, $4.184\ \text{J/g} \cdot ^\circ\text{C}$), and $\Delta T$ is the change in temperature. The molar enthalpy ($\Delta H$) is then found by dividing $q_{\text{reaction}}$ by the moles of the limiting reactant.
Hess's Law
Hess's Law states that if a reaction can be expressed as the algebraic sum of a sequence of other reactions, the enthalpy change for the overall reaction will be the sum of the enthalpy changes for the individual reactions. Enthalpy is a state function, meaning the final result depends only on the initial and final states, not the path taken.
The law allows us to find the unknown $\Delta H_f^\circ$ for $\text{MgO}$ by constructing a hypothetical path using two reactions that can be easily measured in the lab:
Reaction 1: Solid magnesium oxide reacting with hydrochloric acid ($\text{HCl}$).
$$\text{MgO}(\text{s}) + 2\text{HCl}(\text{aq}) \rightarrow \text{MgCl}_2(\text{aq}) + \text{H}_2\text{O}(\text{l}) \quad \Delta H_1$$
Reaction 2: Solid magnesium metal reacting with hydrochloric acid ($\text{HCl}$).
$$\text{Mg}(\text{s}) + 2\text{HCl}(\text{aq}) \rightarrow \text{MgCl}_2(\text{aq}) + \text{H}_2(\text{g}) \quad \Delta H_2$$
By running both reactions in a calorimeter, measuring their respective $\Delta H_1$ and $\Delta H_2$, and combining them algebraically with the known enthalpy of formation of water ($\Delta H_f^\circ$ of $\text{H}_2\text{O}$), we can calculate the $\Delta H_f^\circ$ for $\text{MgO}$.