Adapted from a lab used at Roanoke College, Salem, Virginia.
This experiment uses calorimetry to determine enthalpy changes for two reactions that can conveniently be performed in the laboratory. Then it makes use of Hess's Law to calculate the enthalpy change for a reaction which is not directly observed.
An exothermic reaction releases heat, and that heat flows into the surroundings and is usually observed as a temperature increase in the solvent. container, and other immediate surroundings. An insulated container called a bomb calorimeter can be used to limit the heat flow with good accuracy to the solvent and bomb. Thus
The negative sign on the left of the equation says that heat is flowing out of the reaction; so the heat flow out of the reaction is equal to the heat flow into the solution and the calorimeter. It is very important to remember that whenever heat is flowing out of a substance or reaction the value of "q" will be negative. In the case of our experiment, the calorimeter absorbs so little heat that qbomb is nearly zero and thus the equation simplifies to
The temperature change caused by the addition of a given amount of heat will depend on the specific heat, Csp, of the substance. The specific heat of water (and a good approximation for dilute aqueous solutions) is 4.18 J/g oC. Thus, the heat involved can be calculated by the equation
The mass of the substance is m, and the change in temperature is delta T. Change in temperature is always calculated as (final temperature) - (initial temperature).
Our experiment is run at constant pressure, thus the heat changes observed will be equal to the enthalpies of reaction, (delta Hrxn for the reactions considered. We will obtain experimental data to directly calculate the (delta Hrxn for two reactions in this experiment.
The enthalpies of these directly observed reactions and Hess's Law will be used to calculate the enthalpy change for an additional reaction which is not directly observed. Recall that since enthalpy is a state function, the enthalpy change for any process will depend only on the starting and ending points, not on the reaction path followed. The enthalpy, (delta H(A-B), change for the reaction
A ----> B
can be calculated by adding up the enthalpy changes for a series of intermediate steps.
A ----> C ----> D ----> B
Thus, delta H(A-B) = delta H(A-C) + delta H(C-D) + delta H(D-B). This concept is especially useful in calculating the enthalpy change for reactions which are not easy to observe or measure directly.
You will first determine the enthalpies of reaction for the neutralization of HCl with NH3, and the dissolution of NH4Cl. These values in turn will be used to calculate the enthalpy of decomposition of NH4Cl(s). For the purposes of these calculations we will assume that the densities of all aqueous solutions are 1.00 g/mL and that their specific heats are equal to 4.184 J/g oC. Because of these assumptions, it is better to use the mass of water and not the mass of water + mass of NH4Cl in part II. You will do three trials of each measurement.
The diluted hydrochloric acid and aqueous ammonia are corrosive to equipment and skin. They are also volatile, therefore, be very careful to keep their vapors out of the room. The dilute acid, base, and salt solutions may all be flushed directly down the sink with lots of water.
Part I. Determination of the enthalpy of reaction for
Measure 25.0 mL of 2.00 M NH3(aq) using a clean, dry graduated cylinder. Then measure 25.0 mL of 2.00 M HCl(aq) directly into your zeroed calorimeter and record the mass of the acid. Place the thermister in the acid and allow it to equilibrate. Click on "Start", wait until you see 2-3 data points and the add the NH3(aq) to the calorimeter. Stir while watching the graph display on the monitor and stop recording data when the graph levels off. Record the minimum and maximum temperatures reached by the mixture. Be sure to stir continuously.
Here we assume that the NH3 and HCl solutions are initially at the same temperature. This is a good assumption when the solutions have been prepared well ahead of lab time and have been allowed to equilibrate at room temperature.
II. Determination of the enthalpy of dissolution for
NH4Cl(s) ----> NH4+(aq) + Cl-(aq)
Weigh out between 5 and 5.5 g of solid NH4Cl. Thoroughly grind the solid for 1-2 minutes. Reweigh the sample on weighing paper, to the nearest 0.001 g. Place 25.0 mL of distilled water into your zeroed calorimeter. Follow the procedure in Part I for using the ULI and add the NH4Cl solid to the water in the calorimeter and stir to completely dissolve the solid.
Results and Discussion:
Use a spreadsheet such as Excel to do your calculations for this lab. Points will be deducted from your grade if you don't use a spreadsheet. Use spreadsheet formulas to calculate the values for these results tables from the data in your earlier data tables. Print this spreadsheet for inclusion in your lab report. Notice that your spreadsheet doesn't keep track of significant figures. Rewrite the average delta H values for the two reactions with the correct number of significant figures. (These average values are the last values calculated in the two results spreadsheets.)
Finally, use Hess's Law to calculate delta H for the reaction
You will need to use four delta H values:
|your experimentally determined value for||NH3(aq) + HCl(aq) ----> NH4+(aq) + Cl-(aq)|
|your experimentally determined value for||NH4Cl(s) ----> NH4+(aq) + Cl-(aq)|
|given delta H = -34640 J/mole for||NH3(g) ----> NH3(aq)|
|given delta H = -75140 J/mole for||HCl(g) ----> HCl(aq)|
1. Determine the effect that each of the following mistakes would have on: a) the results of the individual sections of the experiment and b) the final calculated value for the enthalpy of decomposition of the ammonium chloride. In each case state whether the value with the mistake would be greater (more positive) than, the same as, or smaller (less positive) than the value calculated if the mistake had not been made.
Be very careful to watch those +/- signs! Explain your answers.
b. A student mis-reads the directions and uses 75.0 mL of each of the solutions in every part of the experiment. The student, however, records 75.0 mL for each volume and uses this value in his calculations.