Calorimetry and Hess's Law
Adapted from a lab used at Roanoke College, Salem, Virginia.
Purpose:
This experiment uses calorimetry to determine enthalpy changes for two reactions that
can conveniently be performed in the laboratory. Then it makes use ofHess's Law to
calculate the enthalpy change for a reaction which is not directly observed.
Introduction:
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
-qreaction = qsolution + qbomb -- Eqn. 1
where "q" is the symbol for heat flow which will be measured in Joules.
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
-qreaction = qsolution -- Eqn. 2
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
q = (Csp)(m)(delta T) -- Eqn. 3
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.
Hazards:
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.
Procedure:
Part I. Determination of the enthalpy of reaction for
HCl(aq) + NH3(aq) ----> NH4+(aq) + Cl-(aq)
Measure 25.0 mL of 2.00 M NH3(aq) using a clean, dry graduated cylinder. Record its
temperature. Then measure 25.0 mL of 2.00 M HCl(aq) directly into your calorimeter
Stir while watching the temperature on the thermometer change. Record the maximum
temperature that the mixture reached. 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 a weigh boat. Place 25.0 mL of distilled water into
your calorimeter. Record the waters initial temperature. Add the ammonium chloride
and stir frequently, observing the change in temperature. Record the maximum
temperature reached.
Results and Discussion:
1) Use Hess's Law to calculate delta H for the reaction
NH4Cl(s) ----> NH3(g) + HCl(g)
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)
Be sure that you show how these equations must be combined to get delta H for the
desired equation, NEATLY.
Analysis
2) Desired Use the standard heats of formation for the compounds involved to
determine an accepted delta H for this reaction. Neatly show your work.
3) Determine your percent error.
QUESTIONS:
4. 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.
a. A student mis-reads the final temperature of the mixture in part I, recording a
temperature that is too low.
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.