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Calorimetry and Hess’s Law Lab Report, Lab Reports of Chemistry

Santa Fe University of Art & Design (SFUAD)Chemistry

Student will learn to measure the ∆H values of two reactions by Calorimetry and Hess’s Law

Typology: Lab Reports

2020/2021

Uploaded on 05/11/2021

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Santa Monica College Chemistry 11
Calorimetry and Hess’s Law Page 1 of 4
Calorimetry and Hess’s Law
Objectives
The objectives of this laboratory are as follows:
• To experimentally measure the ∆H values of two reactions using the technique of constant
pressure calorimetry.
• To apply these ∆H values in a Hess’s Law calculation to determine the enthalpy of
combustion of a metal.
Background
The combustion of a metal in oxygen produces the corresponding metal oxide as the only
product. Such reactions are exothermic and release heat. For example, the combustion of iron
releases 1651 kJ of heat energy for every four moles of iron burned:
(1) 4 Fe (s) + 3 O2 (g) → 2 Fe2O3 (s)
∆
H1 = -1651 kJ
Since it is difficult to measure the enthalpy of combustion of a metal directly, in this lab it will be
determined indirectly by applying Hess’s Law of Heat Summation. Hess’s Law states that the
enthalpy change of an overall process is equal to the sum of the enthalpy changes of its
individual steps.
Hess’s Law Example: Determine
∆
H for the target reaction 2 NO2 (g) + ½ O2 (g) → N2O5 (g)
given the following information,
Reaction A N2O5 (g) → 2 NO (g) + ³/2 O2 (g)
∆
HA = +223.7 kJ
Reaction B NO2 (g) → NO (g) + ½ O2 (g)
∆
HB = -57.1 kJ
Solution: Reactions A and B have to be carefully manipulated before they can be summed to
produce the target reaction. Reaction A must be reversed, causing a sign change to
∆
HA.
Reaction B must be multiplied by a factor of 2, causing
∆
HB to be multiplied by 2. Only then
will they yield the target equation when added together:
2 NO (g) + ³/2 O2 (g) → N2O5 (g)
∆
H = −(+223.7) = -223.7 kJ
2 NO2 (g) → 2 NO (g) + O2 (g) +
∆
H = 2 x (-57.1) = -114.2 kJ
2 NO2 (g) + ½ O2 (g) → N2O5 (g) Target
Thus,
∆
HTarget = -223.7 + (-114.2) = -337.9 kJ
In order to use Hess’s Law to find the heat of combustion of a metal, it is first necessary to
obtain reaction enthalpies (∆H values) for equations that can be summed together appropriately.
To accomplish this, two reactions will be studied in this lab. In one reaction, a given metal will
react with hydrochloric acid producing hydrogen and the metal chloride. In the other reaction,
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Calorimetry and Hess’s Law

Objectives

The objectives of this laboratory are as follows:

  • To experimentally measure the ∆H values of two reactions using the technique of constant pressure calorimetry.
  • To apply these ∆H values in a Hess’s Law calculation to determine the enthalpy of combustion of a metal.

Background

The combustion of a metal in oxygen produces the corresponding metal oxide as the only product. Such reactions are exothermic and release heat. For example, the combustion of iron releases 1651 kJ of heat energy for every four moles of iron burned:

(1) 4 Fe ( s ) + 3 O 2 ( g ) → 2 Fe 2 O 3 ( s ) ∆ H 1 = -1651 kJ

Since it is difficult to measure the enthalpy of combustion of a metal directly, in this lab it will be determined indirectly by applying Hess’s Law of Heat Summation. Hess’s Law states that the enthalpy change of an overall process is equal to the sum of the enthalpy changes of its individual steps.

Hess’s Law Example : Determine ∆ H for the target reaction 2 NO 2 ( g ) + ½ O 2 ( g ) → N 2 O 5 ( g )

given the following information,

Reaction A N 2 O 5 ( g ) → 2 NO ( g ) + ³ / 2 O 2 ( g ) ∆ HA = +223.7 kJ

Reaction B NO 2 ( g ) → NO ( g ) + ½ O 2 ( g ) ∆ HB = -57.1 kJ

Solution : Reactions A and B have to be carefully manipulated before they can be summed to

produce the target reaction. Reaction A must be reversed, causing a sign change to ∆ HA.

Reaction B must be multiplied by a factor of 2, causing ∆ HB to be multiplied by 2. Only then

will they yield the target equation when added together:

2 NO ( g ) + ³ / 2 O 2 ( g ) → N 2 O 5 ( g ) ∆ H = −(+223.7) = -223.7 kJ

2 NO 2 ( g ) → 2 NO ( g ) + O 2 ( g ) + ∆ H = 2 x (-57.1) = -114.2 kJ

2 NO 2 ( g ) + ½ O 2 ( g ) → N 2 O 5 ( g ) Target

Thus, ∆ HTarget = -223.7 + (-114.2) = -337.9 kJ

In order to use Hess’s Law to find the heat of combustion of a metal, it is first necessary to obtain reaction enthalpies (∆H values) for equations that can be summed together appropriately. To accomplish this, two reactions will be studied in this lab. In one reaction, a given metal will react with hydrochloric acid producing hydrogen and the metal chloride. In the other reaction,

the corresponding metal oxide will react with hydrochloric acid producing water and the metal chloride. For example, the reactions involving iron and iron(III) oxide are as follows:

(2) 2 Fe ( s ) + 6 HCl ( aq ) → 2 FeCl 3 ( aq ) + 3 H 2 ( g ) ∆ H 2

(3) Fe 2 O 3 ( s ) + 6 HCl ( aq ) → 2 FeCl 3 ( aq ) + 3 H 2 O ( l ) ∆ H 3

Since both reactions are exothermic, the heat released ( q ) will be absorbed into the surrounding reaction mixture. As long as the reactions are performed in an insulated container (such as a coffee cup calorimeter) there will be negligible heat exchange with the container walls or outside air. By monitoring the temperature of the reaction mixture when specific quantities of reactants are used, the amount of heat (in J) released by these reactions can be determined by applying the equation:

heat released by reaction ( − q reaction ) = heat absorbed by reaction mixture ( +q mixture)

= ( m x c x ∆ T ) mixture

Here m is the total mass of the reaction mixture (in g), ∆ T is the maximum temperature change

that occurs during the reaction (in °C), and c is the specific heat capacity of the mixture (in J/g•°C). Note that since the reactions occur in aqueous solution, it is reasonable to substitute the specific heat capacity of water (= 4.184 J/g•°C) for the specific heat capacity of the mixture.

Recall that at constant pressure (the conditions of this experiment), the heat released by the reaction equals the reaction enthalpy:

q P = ∆ H

Since the heat released by each reaction is proportional to the amount of metal/metal oxide

used, ∆ H 2 and ∆ H 3 can be easily calculated per gram or mole of metal/metal oxide used.

It should be noted that reactions (2) and (3) by themselves still cannot be summed to produce Reaction (1). Another reaction is required:

(4) 2 H 2 ( g ) + O 2 ( g ) → 2 H 2 O ( l ) ∆ H 4

∆ H for this reaction (the formation of water from its elements) must be obtained from tabulated

thermodynamic data in the textbook. Finally, the reactions (2), (3) and (4) and their enthalpies may be summed together according to Hess’s Law to determine the enthalpy of combustion of the given metal (1).

  1. Assemble your equipment as shown in the figure below. The thermometer (or temperature probe) and the stirring rod must be inserted through the holes in the calorimeter lid. The thermometer bulb should be immersed in the acid, but not touch the bottom of the calorimeter. Clamp the thermometer in place using the slotted stopper and utility clamp.
  2. Measure the temperature of the HCl in the calorimeter (while covered with the lid). Once thermal equilibrium is established, record the temperature. Next, carefully add the metal sample to the acid. Quickly replace the lid and monitor the temperature change until the reaction is complete. Stir the mixture continuously with the stirring rod as the reaction occurs. Record the maximum temperature achieved by the mixture. Note that the mixture first warms up as the reaction occurs, but will then gradually cool as heat is lost to the surroundings. However, as Styrofoam is a poor conductor of heat this cooling will occur slowly. Thus it will be very easy for you to identify the maximum temperature.
  3. When finished, dispose of your chemical waste as directed by your instructor. Then rinse the calorimeter, thermometer and stirring rod thoroughly with distilled water, dry, and repeat the experiment again. Once you have completed both trials with the metal, perform your two trials using the metal oxide using the identical procedure.

Note that the thermometer must be clamped in place using the slotted stopper and utility clamp/stand. You may also want to place the nested cups in a medium beaker for extra stability.