Determining the Enthalpy of Reaction

Determining the Enthalpy of Reaction

In this lab, you will examine 2 reactions. The first reaction is between Magnesium Metal and Hydrochloric Acid, and the second reaction is the dissolution (dissolving) of the compound ammonium chloride. Using a coffee-cup calorimeter, you will conduct the reaction, examine the quantities of the reactants and stoichiometry, and then you will also determine the enthalpy change for the reaction.

Part 1: Pre-Lab Questions

1. For part 1 of your experiment, you will be reacting 1.00g of Mg with 100.0mL of 1.0mol/L of HCl(aq). Calculate the number of moles of each of the reactants used in the experiment.

2. For part 2 of your experiment, you will be dissolving 1.00g of NH4Cl in 25.0mL water. Calculate the number of moles in NH4Cl.

3. Draw a basic diagram of a calorimeter. When conducting chemical reactions involving heat & enthalpy, why is it important that we use a calorimeter to run the reaction? Why can’t we do the reaction in an open beaker?

4. In the reaction of Magnesium + Hydrochloric Acid, explain why magnesium is able to displace Hydrogen in the reaction.

5. In part 2 of this lab, you will dissolve an ionic salt in water. What makes an ionic compound (like ammonium chloride) soluble? Why are some salts insoluble?

Lab Objectives:

· To determine how the initial (given) quantities of the reactants affect the quantities of the products

· To determine the heat transfer involved in the reaction between Mg and HCl.

· To determine the heat transfer involved in the dissolving of NH4Cl

· To covert the heat transfer into an enthalpy of reaction

Materials Required:

Balance Magnesium Metal Strip Stirring Rod

Coffee Cup Calorimeter Hydrochloric Acid (1.0mol/L) Ammonium Chloride Salt

Thermometer Funnel 50mL Graduated Cylinder

Goggles Sand Paper

Procedure: Make a lab-doodle (or flow-chart) to summarize the procedure in your lab report

1. Obtain a strip of Magnesium metal. Take the mass of the magnesium metal and record the exact mass (should be close to 1.00g).

2. Using the sand paper, sand down the Magnesium metal to remove any impurities

3. Using a funnel and graduated cylinder, carefully measure 100.0mL of the 1.0mol/L HCl(aq)

4. Transfer the HCl(aq) into your coffee cup calorimeter.

5. Using the thermometer, take the initial temperature of the HCl(aq) solution. Record the temperature.

6. Add the 1.00g strip of Mg(s) into your coffee cup calorimeter

7. Apply the lid to the calorimeter. Allow the reaction to run and measure the maximum temperature reached by the reaction.

8. When the reaction is complete, pour the contents of the calorimeter into the waste bin and rinse your calorimeter. You will know that the reaction is complete when the Mg completely “dissolves”

9. Repeat the experiment using 1.00g of NH4Cl (measured on the scale) and 25mL water (measured with a graduated cylinder).

10. Use the stirring rod to dissolve the salt if necessary.


Create an observation table (or tables) to record quantitative data you capture in this experiment.


1. Write the balanced chemical equation for Reaction 1. Calculate the Limiting Reactant and Excess Reactant for Reaction 1 (Mg + HCl). Using your calculations, determine the mass of Hydrogen gas produced.

2. Based only on your qualitative observations (ignoring the calculations from question 1), how can you tell which reactant is limiting and which reactant is excess? Why?

3. Why is it not necessary to calculate the limiting reactant/excess reactant in Reaction 2 (Ammonium Chloride)?

4. From your data, calculate the following for each part of the experiment. (make sure it is organized )

a. The temperature changes of the water

b. The mass of the water

c. The quantity of heat absorbed (or given off) by the water during the dissolving, given that specific heat of water is 4.184 J/g oC

d. The number of moles of solid used

e. The quantity of heat involved per mole of solid dissolved (value from part c divided by value from part d). This is called the molar enthalpy of solution ∆Hrxn. Assign a positive or negative sign to ∆Hrxn depending in whether you think the dissolving process was exothermic or endothermic.

5. Write a thermochemical equation for the dissolving process for each solid (include the heat term in each equation).

6. Magnesium oxide, MgO, and magnesium chloride, MgCl2, are very similar, white ionic solids with the following properties:

Compound Melting Point (oC) Solubility
MgO 2800 Insoluble
MgCl2 1412 Very soluble


a. Give the formula of the ions of each compound

b. Account for the drastic difference in physical properties

7. The value of ∆Hrxn for the formation of an acetone – water solution is negative. Explain this in general terms that discuss intermolecular forces of attraction.


1. Write a mathematical equation that shows the relationship between ∆Hrxn, Q (heat) and n (moles).

2. Discuss 2 experimental sources of error that occurred in this experiment.

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