CHE 1142 Lab3inclass Enthalpy Neutralization Hess with Data Table 2020 PDF

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CHE 1142 3

Calorimetry: Determining the Enthalpy of a Chemical Reaction from Hess’ Law All chemical reactions involve an exchange of heat energy; therefore, it is tempting to plan to follow a reaction by measuring the enthalpy change (∆H). The heat energy change of the reactants and products (the system) that occurs in aqueous solution can often be measured the from heat change that occurs in the surroundings by monitoring temperature changes using the following equation. q = Cp  m  ∆T The term q represents the heat energy that is gained or lost. Cp is the specific heat of water, m is the mass of water, and ∆T is the temperature change of the reaction mixture. The specific heat and mass of water are used because water will either gain or lose heat energy in a reaction that occurs in aqueous solution. A styrofoam cup nested in a beaker will be used as a crude calorimeter, as shown in Figure 1. For purposes of this experiment, assume that the heat loss to the calorimeter and the surrounding air is negligible.

Figure 1 Hess’s law states that the enthalpy changes of a series of reactions can be combined to calculate the enthalpy change of a reaction that is the sum of the components of the series. In this experiment, the temperature change of two reactions will be measured, and Hess’s law will be used to determine the enthalpy change, ΔH of a third reaction.H3 = H1 - H2 1. 2. 3.

NaOH (aq) + HCl (aq) HCl (aq) + NH4OH (aq) NaOH (aq) + NH4Cl (aq)

→ NaCl (aq) + H2O (l) → NH4Cl (aq) + H2O (l) → NaCl (aq) + NH4OH (aq)

H1 H2 H3

OBJECTIVES Determine enthalpy of reaction from a measured temperature change in a reaction solution Use Hess’s law to determine the enthalpy change of the reaction between aqueous ammonia and aqueous hydrochloric acid.  Compare calculated enthalpy change with the experimental results.

 

Adapted from Advanced Chemistry with Vernier 13

CHE 1142 3 - 1

Determining the Enthalpy of a Chemical Reaction

MATERIALS Vernier computer interface computer Temperature Probe Styrofoam cup 250 mL beaker 50 mL or 100 mL graduated cylinders glass stirring rod

2.0 M hydrochloric acid, HCl, solution 2.0 M sodium hydroxide, NaOH, solution 2.0 M ammonium chloride, NH 4Cl, solution 2.0 M ammonium hydroxide, NH4OH, solution ring stand utility clamp

PROCEDURE 1. Obtain and wear goggles. 2. Connect a Temperature Probe to Channel 1 of the Vernier computer interface. Connect the interface to the computer with the proper cable. Use a utility clamp to suspend the Temperature Probe from a ring stand, as shown in Figure 1. 3. Start the Logger Pro program on the computer. Open the file “13 Enthalpy” from the Advanced Chemistry with Vernier folder. Part I Conduct the Reaction Between Solutions of NaOH and HCl 4. Nest a Styrofoam cup in a beaker (see Figure 1). Measure 50.0 mL of 2.0 M HCl solution into the cup. Lower the tip of the Temperature Probe into the HCl solution. 5. Measure out 50.0 mL of NaOH solution, but do not add it to the HCl solution yet. CAUTION: Handle the sodium hydroxide solution with care. 6. Conduct the reaction. a. Click to begin the data collection and obtain the initial temperature of the HCl solution. b. After three or four readings have been recorded at the same temperature, add the 50.0 mL of NaOH solution to the Styrofoam cup all at once. Stir the mixture gently with a glass rod throughout the reaction. c. Data collection will end after three minutes. If the temperature readings are no longer changing, terminate the trial early by clicking . d. Click the Statistics button, . The minimum and maximum temperatures are listed in the statistics box on the graph. If the lowest temperature is not a suitable initial temperature, examine the graph and determine the initial temperature. e. Record the initial and maximum temperatures in the data table. 7. Rinse and dry the Temperature Probe, Styrofoam cup, and the stirring rod. Dispose of the solution as directed.

CHE 1142 3 - 2

Adapted from Advanced Chemistry with Vernier 13

Determining the Enthalpy of a Chemical Reaction Part II Conduct the Reaction Between Solutions of HCl and NH 3 8. Measure out 50.0 mL of 2.0 M HCl solution into a nested Styrofoam cup (see Figure 1). Lower the tip of the Temperature Probe into the cup of HCl solution. 9. Measure out 50.0 mL of 2.0 M NH 3 solution, but do not add it to the HCl solution yet. 10. Conduct this reaction in a fume hood or in a well-ventilated area. Repeat Step 6 to conduct the reaction and collect temperature data.

DATA ANALYSIS Complete the data table on the next page using the following calculations: 1. Calculate the amount of heat energy, q, produced in each reaction. Use 1.03 g/mL for the density of all solutions. Use the specific heat of water, 4.18 J/(g •°C), for all solutions. q = Cp  m  ∆T 2. Find the number of moles of salt product in each reaction, then calculate the enthalpy change, ∆Hrxn, for each reaction in terms of kJ/mol of reactant. (beware of J to KJ conversion) moles = Molarity x Liters Hrxn = q / mol 3. Use the answers from 2 above and Hess’s law to determine the experimental molar enthalpy for Reaction 3. H3 = H1 - H2 1. NaOH (aq) + HCl (aq) → NaCl (aq) + H2O (l) H1 2. HCl (aq) + NH4OH (aq) → NH4Cl (aq) + H2O (l) H2 3. NaOH (aq) + NH4Cl (aq) → NaCl (aq) + NH4OH (aq) H3 4. Use Hess’ law, and the literature values of ∆H orxn given in the data table to calculate the expected ∆Horxn for Reaction 3. 5. What percentage error does the experimental value have compared to the literature values?

Adapted from Advanced Chemistry with Vernier 13

CHE 1142 3 - 3

DATA TABLE Reaction

Initial Temp

Final Temp

T

q

oC

oC

oC

J

moles product made

Hrxn from experiment

Literature Value Horxn

KJ/mol

KJ/mol

Percent experiment error

1 -55.9

2 -52.2

3

NOTES

A:

B:

C:

D:

E:

F:

Close to room temperature

Maximum or minimum temperature

Temperature difference

q=mCpT

Mreactant x Liters reactant used

(D/1000)/E

B-A

G:

H: (F-G)/G x 100...