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Heat of Formation of Ammonium Chloride
Objectives:
1. To use calorimetry to measure the heat of neutralization of the reaction of
aqueous NH3 with aqueous HCl to form aqueous ammonium chloride.
2. To use calorimetry to measure the heat of solution of ammonium chloride.
3. To calculate the heat of formation of solid ammonium chloride using these
data and the known heats of formation of NH3 and HCl solutions.
Discussion:
When a chemical reaction takes place there is always an exchange of energy. This
energy can take several forms, such as heat, electrical energy, or work done against
external forces such as pressure. In this experiment the amount of heat released or
absorbed during a chemical reaction will be studied.
A reaction in which heat is lost by the reaction to the surroundings is exothermic. A
reaction in which heat is absorbed is endothermic. In this experiment any heat evolved
or absorbed during the course of the reaction is transferred to, or from, the water in
which the reaction takes place. The water is insulated by the container from the outside
and acts as the surroundings for the reaction. This container is called a calorimeter. In
this reaction the water is already in the acid/base.
The term heat of reaction can apply to several different types of reactions:
1. The Heat of Formation: the amount of heat involved in producing one mole of a
substance from its constituent elements.
2. The Heat of Solution: The amount of heat involved when one mole of a
substance is dissolved in a solvent.
3. The Heat of Neutralization: the amount of heat involved when 1 mole of water
is produced by the reaction of an acid and a base.
4. The Heat of Combustion: the amount of heat given off when one mole of
combustible material is reacted with excess oxygen.
5. The Heat of Vaporization, Fusion, or Sublimation: the amount of heat
involved in changes of state.
The heat evolved or absorbed in a process at constant pressure is H, the
thermodynamic quantity called enthalpy. The term  is used to denote a difference
between the amount of heat in a system in the final state Hf and the amount of heat in a
system in the initial state Hi.
H = Hf – Hi
If heat is given off during a reaction, then there will be less enthalpy in the system in the
final state. This means that Hf will be smaller than Hi and, therefore, H will be
negative. If heat is evolved H is negative (exothermic). If heat is absorbed H is
positive (endothermic).
In this experiment a reaction will take place in water inside a calorimeter. Any heat
released by the reaction will be absorbed by the water and the temperature of the water
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will drop. The mass of water can be calculated from its volume, the temperature
change can be measured and the exact quantity of heat released or absorbed can then
be calculated.
Hess’s law states that H for a reaction is independent of the steps involved in the
reaction. It is only dependent upon the final products and the initial reactants. It is
equal to the sum of the Hs for the steps in the reaction. Read your textbook for more
information on Hess’s Law.
The objective of this experiment is to find the Hf of the following reaction:
½ N2(g) + 2H2(g) + ½ Cl2(g)  NH4Cl(s)
This would be difficult to measure directly because it involves reactants in the gas
phase and a tremendous amount of heat. Instead, the various Hs involved in a multistep process and Hess’s Law will be employed.
Hess’s Law can be used to determine H for the following reaction
NH3(g) + HCl(g)  NH4Cl(s)
Given the following:
NH3(g)  NH3(aq)
H = -35.4 kJ/mole
HCl(g)  HCl(aq)
H = -75.1 kJ/mole
And the values of H for the following two reactions, which are determined in this
experiment:
NH3(aq) + HCl(aq)  NH4Cl(aq)
NH4Cl(s)  NH4Cl(aq)
The four equations above can be arranged so that when added together they will give
the NET equation:
NH3(g) + HCl(g)  NH4Cl(s).
The following heats of formation are given:
½ N2(g) + 1½ H2(g)  NH3(g)
Hf = -45.8 kJ/mole
Hf = -92.3kJ/mole
½ H2(g) + ½ Cl2(g)  HCl(g)
These, together with the heat of reaction which has already been determined for,
NH3(g) + HCl(g)  NH4Cl(s)
Can be used to calculate the heat of formation (Hf) of NH4Cl(s) from its elements.
½ N2(g) + 2H2(g) + ½ Cl2(g)  NH4Cl(s)
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H is measured in each reaction by measuring the temperature change of the water in
which the reaction is taking place. By knowing the specific heat of water and the
amount of water involved, the heat involved in the reaction can be calculated. The
solution that is used is very dilute and can be considered to be water.
For example, if heat is evolved during the reaction, then
Heat lost by reaction = Heat gained by water.
Heat gained by water = Mass of water x specific heat x temperature change.
The specific heat of water is 4.184J/g K
The density of water is 1.00 g/mL.
The calorimeter is not a perfect insulator, so some heat will be lost to the surroundings.
Temperature readings should be taken over a period of several minutes and a graph
drawn to plot temperature against time. Heat will be lost or gained at a constant rate
and the graph can be extrapolated back to zero time (the time when the reactants were
mixed). The exact temperature at the time of mixing can therefore be found. The
difference between the temperature of the reactants before they were mixed and the
temperature at time zero is T.
Procedure
A. HEAT OF NEUTRALIZATION
Set up the Styrofoam cup calorimeter (use a 250 mL beaker to stabilize!), making sure
that the thermometer does not touch the bottom or sides of the cup. Obtain 50mL of 1.0
M HCl in a graduated cylinder (record to the nearest 0.1 mL) and pour HCl solution into
the calorimeter. Be sure to cover the calorimeter! Heat can easily be exchanged
through the top. Measure and record the temperature of the HCl to the nearest 0.1˚C.
After one minute, measure and record the temperature of the HCl solution again. After
two minutes, measure and record the temperature. After three minutes, measure and
record the temperature again. All three of these temperature measurements should be
constant. This is time = 0 min If the temperature is not constant, continue measuring
at one-minute intervals until it remains constant.
While one lab partner is recording the temperature of the HCl, the other partner should
be cleaning the graduated cylinder. After the cylinder is clean, obtain 50 mL of 1.0 M
NH4OH (again, record to the nearest 0.1 mL). Set aside.
When the temperature of the HCl is constant, quickly pour the ammonia solution into the
calorimeter to react with the HCl and replace the calorimeter cover. Stir. One minute
after mixing read the thermometer and record the temperature. Always record the
elapsed time along with the temperature. Continue measuring the temperature at oneminute intervals, stirring the mixture, for about 5 more minutes. If you are using a
LabQuest to graphically collect the data, you can extrapolate the constant “before
mixing” line forward to the “time of mixing”. Extrapolate the “after mixing” line back to
the “time of mixing”. Find the value of T at the instant the solutions were mixed.
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B. HEAT OF SOLUTION OF NH4Cl
Using the same Styrofoam cup calorimeter as before, measure out 100 mL of deionized
water using a graduated cylinder (record to the nearest 0.1 mL) and transfer to the cup.
Be sure to cover the calorimeter! Heat can easily be exchanged through the top.
Measure and record the temperature of the water to the nearest 0.1˚C. After one
minute, measure and record the temperature of the water again. After two minutes,
measure and record the temperature. After three minutes, measure and record the
temperature again. All three of these temperature measurements should be constant.
This is time = 0 min If the temperature is not constant, continue measuring at oneminute intervals until it remains constant. Mass out approximately 2.00 g of NH4Cl and
record the exact mass in your lab notebook. Quickly transfer the solid to the calorimeter
and record the temperature changes as before. Find the value of T at the instant the
water and ammonium chloride was mixed.
Calculations & Data
1. Calculate the mass of water involved in each reaction and from this and your T
(“change in temperature”) in each case, calculate the amount of heat evolved or
absorbed in each reaction.
2. Calculate the number of moles of ammonium chloride involved in each reaction,
and from this and the amount of heat involved, calculate H for each reaction in
kJ/mole.
3. Construct a graph of temperature change versus time.
4. Using Hess’s Law and the six reactions, calculate the H for reaction below:
½ N2(g) + 2H2(g) + ½ Cl2(g)  NH4Cl(s)
Discussion
Discuss the temperature change of the water when a reaction in solution is exothermic
or endothermic. How does this affect the sign of H?
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