Thermodynamics- Enthalpy of Reaction and Hess’s Law
December 5, 2011
Kylie Case, Emma McKee, Rebecca Smith
Purpose: In this lab, the purpose was to verify Hess’s Law.
Theory:
Four main topics were covered during this experiment including enthalpy of reaction,
heat of formation, Hess’s Law, and calorimetry.
The first being enthalpy of reaction, ΔHrxn, which is the heat or enthalpy change for a
chemical reaction. The energy change is equal to the amount of heat transferred at a constant
pressure in the reaction. The change represents the difference in enthalpy of the products and the
reactants and is independent of the steps in going from reactants from products.
The second, is the heat of formation (ΔH°f), which is also known as standard enthalpy of
formation, and is defined as the change in enthalpy that accompanies the formation of one mole
of a compound from its elements with all substances in their standard states. The enthalpy
change for a given reaction can be calculated by subtracting the enthalpies of formation of the
reactants from the enthalpies of formation of the products. This can be illustrated in this
equation: 𝛥𝐻°𝑟𝑥𝑛 = Σ𝑛𝑝 𝛥𝐻°𝑓 (𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠) − Σ𝑛𝑟 𝛥𝐻°𝑓 (𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠).
During this experiment, Hess’s Law was a concept that the lab concentrated on. Hess’s
Law states that if a reaction can be carried out in a series of steps, the sum of the enthalpies for
each step equals the enthalpy change for the overall reaction. For example, the three chemical
equations used throughout the experiment were:
Equation 1: NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(l)
Equation 2: NH4Cl(aq) + NaOH(aq) → NH3(aq) +NaCl(aq) + H2O(l)
Equation 3: NH3(aq) + HCl(aq) → NH4Cl(aq)
Resulting in Equation 1 plus Equation 2(reversed) will equal Equation 3. In this case, Hess’s
Law gave the ΔH for Equation 3.
The last important concept covered in this lab is calorimetry. Calorimetry is the science
of measuring heat, and is based on observing the temperature change when a body absorbs or
discharges energy as heat. A calorimeter is the device used experimentally to determine the heat
associated with a chemical reaction. During this experiment, a coffee cup calorimeter was used.
The equation 𝑞 = 𝑚𝑐𝛥𝑇 is associated with a calorimeter because it discovers the amount of heat
produced by multiplying the mass of the substance by the specific heat and by the change in
temperature. Since the coffee cup calorimeter will absorb heat, the data must be corrected for
that by solving for qcal and ccal. The value of qcal will be the opposite of 𝑞𝐻2 𝑂 , and in order to find
the value of ccal, the equation 𝑞𝑐𝑎𝑙 = 𝛥𝑇𝑐𝑐𝑎𝑙 must be used.
Procedure:
For Part 1, 50.0 mL of distilled water at room temperature needs to be measured, its
initial temperature recorded, and placed into the calorimeter. Then, 50.0 mL of heated water will
need to be measured, its initial temperature record, and placed into the calorimeter with the room
temperature water. With the motor on and the magnet spinning, immediately insert the
thermometer probe and collect the data. For Part 2, 50.0 mL of HCl and NaOH needs to be
obtained and their initial temperatures recorded. Then add them to the calorimeter, cover the
calorimeter, insert the temperature probe, and start collecting the data. Do this procedure for
Reaction 2 and Reaction 3. Then, the heat change and enthalpy of reaction are to be calculated.
The value of ΔH is to be found two ways, through experimental data and using Hess’s Law.
These values are then compared to find the percent error of the experiment.
Data:
Part 1: Determination of the Heat Capacity of the Calorimeter
Initial Temperature
50.0 mL H2O—room temperature:
22.7 °C
50.0 mL H2O—heated:
62.6 °C
Mixing Data
Time (seconds) Temperature (°C)
20
44.4
40
44.4
60
44.3
80
44.2
100
44.1
120
44.0
140
43.9
160
43.8
180
43.7
Calculated Values
Tmix = 44.5 °C
Tavg = 42.7 °C
qcal = +752.41 J
Ccal = 34.51 J/°C
Part 2: Determination of Heats of Reaction
Reaction 1
Initial Temperature
50.0 mL 2.0 M HCl
24.5 °C
50.0 mL 2.0 M NaOH
23.9 °C
Mixing Data
Time (seconds) Temperature (°C)
20
36.7
40
36.7
60
36.6
80
36.6
100
36.5
120
36.4
140
36.4
160
36.3
180
36.3
Calculated values
Tmix = 36.8 °C
qrxn = -5720.115 J
Reaction 2
Initial Temperature
50.0 mL 2.0 M NH4Cl
23.2 °C
50.0 mL 2.0 M NaOH
23.9 °C
Mixing Data
Time (seconds) Temperature (°C)
20
24.7
40
24.7
60
24.6
80
24.6
100
24.6
120
24.6
140
24.6
160
24.6
180
24.6
Calculated Values
ΔH = -57.6 kJ/mol
Tmix = 24.7°C
qrxn = -1860.2 J
ΔH = -18.60 kJ/mol
Reaction 3
Initial Temperature
50.0 mL 2.0 M NH3
22.4 °C
50.0 mL 2.0 M HCl
21.9 °C
Mixing Data
Time (seconds) Temperature (°C)
20
26.4
40
26.4
60
26.4
80
26.4
100
26.3
120
26.3
140
26.3
160
26.3
180
26.3
Calculated Values
Tmix = 26.4 °C
qrxn = -697.575 J
ΔH = -6.97 kJ/mol
Error Analysis:
% 𝑒𝑟𝑟𝑜𝑟 =
|𝐴𝑐𝑡𝑢𝑎𝑙 − 𝑇ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙|
|−26 − (−18.6)|
𝑥 100 =
𝑥 100 = −39.78 %
𝑇ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙
−18.6
The first error that occurred in the experiment was that the group had to assume the
densities of the solutions. In the calculation selection of the lab handout, it stated to assume the
density of the solutions is 1.03 g/mL. Instead of assuming this value, the group could have
calculated it to have a more accurate answer. Like the error before, it would have a small effect
on the data that was produced.
The second error was that the initial temperature measurements for all solutions were not the
same. For example, in the first reaction of Part 2, the initial temperature of hydrochloric acid
was 24.5 °C and the initial temperature of sodium hydroxide was 23.9 °C. In order to find the
initial temperature, the average of these two numbers had to be taken, which could have caused
the calculations to be slightly higher or slightly lower than what they should have been.
However, this error would have a minor effect on the data since there is not a large difference
between the numbers.
The third error that occurred in the experiment was that too much or too little water could
have been added to the solution in the process of mixing the molar solutions. This could have
increased or decreased the molarity of the solution trying to be made. This would affect the
results but only in a minor way.
Conclusion:
The purpose of this experiment was met because the group was able to verify Hess’s Law. From
the calculations, ΔH3 was found to be -26 kJ/mol, and when Hess’s Law was used, ΔH3 was
found to be -18.6 kJ/mol. The percent error calculated from these values ended up being 39.78%. Since the calculated answers were close, the purpose was met.
Questions
1. The direct method worked better to find ΔH3 than the indirect method because in order
to use Hess’s Law, the data had to be nearly error free. If the group does have an error in
the calculations, than the Hess’s Law answer will not match the answer the group
calculated. The Hess’s Law answer had a -61% error compared to the calculated answer.
Therefore, the direct method worked better than the indirect method.
2. Hess’s Law states that if a reaction is carried out in a number of steps, ΔH for the overall
reaction is equal to the sum of the ΔH’s from each individual step.
3. ΔH means the heat or enthalpy change for a chemical reaction. This energy change is
equal to the amount of heat transferred, at constant pressure, in the reaction. This change
represents the difference in enthalpy of the products and the reactants and is independent
of the steps in going from reactants to products.
4. The true initial temperature is found by using a modified linear fit because when
excluding any points that are skewed, it is giving the results a more accurate reading.
The first couple points that were skewed are due to incomplete mixing and lack of
equilibrium with the thermometer.
5. All of the solutions did not have the same initial temperature because they were not
contained in the same area of the room. These areas could have a slightly different
temperature causing the different initial temperatures. Also, transfer of heat from holding
the graduated cylinder could have increased the initial temperature, and depending on
how long the graduated cylinder was held, could have determined how high the initial
temperature rose.