Enthalpy of Formation of Ammonium Chloride PDF

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Enthalpy of Formation of Ammonium Chloride Chemistry 105 section 2 - instructor Otoikian October 24, 2019 Kierstin De Koeyer Lab partner: Chloe Gould Introduction: The main objective of this experiment was to find the enthalpy of formation of solid ammonium chloride. Heat flow at a constant pressure was directly measured using the constant-pressure calorimeter in two different reactions. Then specific heat capacity, mass, and overall temperature change of each of the solutions was used to calculate the energy exchanges made (Flowers, 2019). If a reaction’s system lost heat to its surroundings, it went through and exothermic process and overall temperature change was increased. If a reaction’s system gained heat from its surroundings, it went through an endothermic process and overall temperature change was decreased (Rice, 2016). Hess’s law was also used to add the changes in enthalpy from multiple chemical equations from the reactions in the experiment. Many of the learned thermodynamic and calorimetry principles were used to see how much heat was transferred from a system to its surroundings, or from the surroundings to its system. In this experiment, this process was done by measuring the transfer of energy in the form of heat from the reaction of aqueous ammonia and hydrochloric acid. Then measuring the same for the reaction of dissolving solid ammonium chloride in aqueous water. With the resulting temperature changes from both reactions, the enthalpy of formation of ammonium chloride could be found (Rice, 2016). Methods and Materials: There were several procedures and trials done to find the enthalpy of formation of ammonium chloride using an insulated vessel called a calorimeter. First, the

calorimeter was set up with a stir rod and temperature probe that was connected to a logger program. For part B, 50 mL of 1 M hydrochloric acid was measured and poured into the calorimeter to sit for several minutes and collect temperatures. Then after 3 minutes, 50 mL of 1 M ammonia was added and stirred continuously for the next 7 minutes until 10 minutes of temperatures were recorded. Then the calorimeter was reset and the same procedures were done for the second trial. After the second trial in part B, the calorimeter was rinsed and set up the same way for part C. Then, 100 mL of distilled water was measured and poured into the calorimeter to collect temperature for 3 minutes. In a weigh boat 2.7 g of N H 4 Cl(s) was measured on a scale. It was quickly added and stirred into the calorimeter at the 3 minute mark. The solution was stirred periodically until the temperature had been recorded for 10 consecutive minutes. This process was repeated for the fourth and final trial of the experiment and all data was added to tables and graphs in the excel template given. The initial and final temperatures from all trials and reactions were then used to calculate the average temperature changes for each of the two reactions. Then energy exchanged and heat flow could be calculated to be used later in finding the enthalpy of formation of ammonium chloride (Rice, 2016). Results: In this experiment 2 trials (½) were done in part B for reactions between 1 M solutions of ammonia and hydrochloric acid. Then in part C, 2 more trials (¾) were done for reactions between solid ammonium chloride dissolving in water. Trial 1 had a temperature change of 5.6 ℃ and trial 2 had a change of 5.5 ℃ , which between the trials averaged to be 5.55 ℃. The qsol was calculated to be - 2.32 kJ and ∆ H 3 was calculated to be − 46.4kJ /mol . Then in part C, trial

3 had a temperature change of -.1℃ and trial 4 had a temperature change of -.5 ℃ . The average change in temperature between the two was -.3 ℃ and used to calculate qsol. The energy of solution was calculated to be .126 kJ because the negative value was switched to a positive value from the endothermic reaction laws. Then ∆ H 4 was calculated to be 2.52 kJ/mol from q of the solution. Finally adding the ∆ H given values 1 and 2 and calculated values 3 and 4, ∆H 5 was -159.42 kJ. Then doing similarly, ∆ H 8 and the enthalpy of formation of ammonium chloride was found to be -297.52 kJ.

Data analysis/Calculations: Calculated grams N H 4 Cl needed for 50 mL of 1 M N H 4 Cl solution: 50ml ×

1L 1000ml

= 0.05mol × 1.0 mol L

0.05NH 4 Cl mol ×

53.46g mol

= 2.673gNH 4Cl

∆T calculation for trial 1 in part B: Part B, Trial 1: T i = 25.5 ℃ and T f = 31.1 ℃ ∆T = 31.1 − 25.5 = 5.6 ℃ Both calculations for qsol (∆H  3 and ∆H4): J qsol= 4.184 g℃ × 100.0g × 5.55℃ = 2.332 × 103J ×

∆ H3 =

−2.32kJ 0.05 mol

1kJ 1000J

= 2.32 kJ

kJ = −46.4 mol

J 1kJ qsol= 4.184 g℃ =− .126 kJ × 100.0g ×− .3℃ =− 1.26 × 102 × 1000J

∆ H4 =

.126kJ 0.05 mol

= 2.52kJ mol

7a: using Hess’s Law: adding  all following ∆H values, including the ones calculated previously to find ∆H5. + N H 3(g) → N H 4(aq)

H Cl(g) → H Cl(aq)

∆H 1 = − 35.4kJ ∆H 2 = − 75.1kJ

N H +4(aq) + H Cl(aq) → N H 4 Cl(aq) ∆ H 3 = − 46.4kJ N H 4 Cl(aq) ← N H 4 Cl (s)

∆ H 4 = − 2.52kJ

NH3(g) + HCl(g) → NH4CL(s)

∆H5= ∆H1 +∆H2 +∆H3+∆H  4

∆H5= ( − 3 5.4kJ) + ( − 7 5.1kJ) + ( − 4 6.4kJ) + ( − 2 .52kJ) = − 159.42 kJ

7b: using Hess’s Law: adding  all following ∆H values to find the enthalpy of formation for solid ammonium chloride.

N H 3(g) + H Cl(g) → N H 4 Cl(s) 1N 2 2(g)

+ 32 H 2(g) → N H3(g)

1 2 H2(g)

+ 12 Cl(g) → H Cl(g)

1⁄2 N2(g) + 2H2(g) + 1⁄2 Cl2(g) →NH4Cl(S)

∆H5=  -159.42 kJ

∆H6=-45.8kJ

∆H7=-92.3kJ

∆H8=

∆H5+∆H6+∆H7

∆H8= ( − 1 59.42kJ) + ( − 4 5.8kJ) + (− 92.3kJ) = − 297.52 kJ

Discussion: From part B, the enthalpy change in the reaction of 1 M aqueous solutions of ammonia and hydrochloric acid was calculated to be -46.4 kJ/mol. This reaction was exothermic because ∆H is negative. Then from part C, the enthalpy change in the dissolving of ammonium chloride in water was calculated to be 2.52 kJ/mol. This reaction was endothermic because ∆H is positive. These values could be slightly inaccurate because of experimental error, such as not pushing the temperature probe all the way into the calorimeter. The experiment assumes no heat escaped from the insulated container, which could be false from experimental error. As the enthalpy of formation of ammonium chloride was calculated to be -297.52 kJ using Hess’s law. The literature known enthalpy of formation of ammonium chloride is − 3 14.892 kJ /mol (Rusic, 2005). This poses error happened in the experiment as well, the temperature probe and calorimeter could be the cause of this as well. Inaccurate readings, technological error, or heat escaping could all be tested again through another experiment, either using the same or different equipment as this experiment used. Conclusion: From this experiment the enthalpy of formation of ammonium chloride was calculated. This was done from reactions of two different types to find the enthalpies and further the final enthalpy of formation. Specifically, thermodynamics and calorimetry were used to measure energy and heat transfers in solutions between their systems and surroundings. Then using average temperature changes, mass of the solution, and specific heat capacity of water, the q of the solutions made were found. These values were used to calculate ∆H values to come up with our enthalpy of formation of ammonium chloride being -297.52 kJ. References:

Flowers, P., Theopold , K., & Langley, R. (2019, August 14). 12.3: Heat Capacity, Enthalpy, and Calorimetry. Retrieved October 25, 2019, from https://chem.libretexts.org/Bookshelves/General_Chemistry/Map:_Principles_of_Modern_Chem istry_(Oxtoby_et_al.)/UNIT_4:_EQUILIBRIUM_IN_CHEMICAL_REACTIONS/12:_Thermod ynamic_Processes_and_Thermochemistry/12.3:_Heat_Capacity,_Enthalpy,_and_Calorimetry. Rice, R. Scheiffer, C. Sakamoto, M. Finnegan, M. Wherland, S. (2016). General Chemistry Laboratory Manual: Principles of Chemistry 1 . Belmont, CA: Star Publishing.

Ruscic, B., & Bross, D. H. (2005). Selected ATcT [1, 2] enthalpy of formation based on version 1.118 of the Thermochemical Network [3]. Retrieved October 25, 2019, from https://atct.anl.gov/Thermochemical Data/version 1.118/species/?species_number=656. POST-LAB: RESULTS AND DISCUSSION

1. Write a brief summary of your experiment. As part of this, compare the effects of exothermic and endothermic reactions on the temperature change of the water. In this experiment we used an insulated calorimeter, temperature probe, stir rod, aqueous 1 M solutions of each ammonia and hydrochloric acid in part B. We used the same materials for part C, but dissolved ammonium chloride in 100 mL of water and using both reactions, their enthalpies, and associated enthalpy values to together determine the enthalpy of ammonium chloride formation (Rice, 2016). In part B, the trendline saw an overall temperature change that increased, indicating an exothermic reaction. In part C, the trendline saw an overall temperature change that decreased, indicating an endothermic reaction. Therefore, the reaction between aqueous ammonia and hydrochloric chloric acid was an exothermic process because the system

lost energy in the form of heat to heat up the solution (water). Then, the reaction between ammonium chloride and water was an endothermic process because the system gained energy in the form of heat from the surroundings, which caused an overall decrease in temperature. 2. Give the balanced reaction (include phases) and report your experimentally determined ΔH values for the following: a. The neutralization of aqueous hydrochloric acid with aqueous ammonia. N H +4(aq) + H Cl(aq) → N H 4 Cl(aq) ∆ H 3 = − 46.4kJ b. Dissolving solid ammonium chloride in water. N H 4 Cl(aq) ← N H 4 Cl (s)

∆ H 4 = − 2.52kJ

c. The formation of solid ammonium chloride from its elements. 1⁄2 N2(g) +  2H2(g) +  1⁄2 Cl2(g) →NH4Cl(S)

∆H8=  − 297.52 kJ

3. Lookup the enthalpy for the formation of solid NH4Cl. Reference your source and indicate the quality of it; more points will be given for a higher quality source. The enthalpy of formation of solid ammonium chloride was found to be − 314.892 kJ /mol (Rusic, 2005). 4. Compare the literature value and your experimental value for the formation of NH4Cl. The literature value is − 3 14.892 kJ /mol (Rusic, 2005), while my experimental value was found to be − 2 97.52 kJ /mol (see calculations/results). 5. If your experimental value for the formation of solid NH4Cl is different from the literature value, what factors, either experimental or experimenter, could have caused your values to be too high or too low? Discuss what you could have done to avoid these errors.

The value found in the experiment is a little less than the value found in the literature source. Possible errors could have been made in the experiment due to not stirring the solutions often enough, or having a gap where the probe went in could give inaccurate temperatures. It could be possible with working with technology and the lab equipment that there could have been problems with accurate readings as well. For getting a lower value, this means there had to have been small errors made, like the ones mentioned and leading to lower overall temperatures being recorded....