Model 1: What impact does temperature have on NH4CL solubility in water?

Model 2: How will changing the limiting reactant impact the procedure’s outcome?

Model 3: How does the number of intermediates involved in the reaction affect the accuracy of Hess’s Law’s prediction of the enthalpy change of a chemical reaction?

Introduction

Three different experiments were conducted in our heat lab. We performed our lab to learn more about enthalpy/ heat change. In Model 2 we changed the limiting reactant of the equation in Model 1 to observe the enthalpy changes. Additionally, we tested out Hess’s Law and learned more about it.

Background

In the preceding experiment, measurements of change in H (enthalpy) will be conducted while a known quantity of reactants are made to react in a calorimeter under normal circumstances (room atmosphere and atmospheric pressure). It is considered that the heat capacity and heat gained/lost from the environment are insignificant.

Variables

Independent Variable	Heat Energy
Dependent Variable	Enthalpy
Controlled Variables	Pressure, Temperature, The amount of reactants

Materials

Model 1

• 250 mL beaker
• Styrofoam cup
• Pen
• Paper
• Ammonium chloride
• Calculator
• Distilled water
• Thermometer Model 2
• Hydrochloric acid (aq)
• Graduated cylinder
• Sodium Hydroxide (aq)
• Calculator
• Pen
• Paper
• Thermometer
• Styrofoam cup
• Distilled water
• 250 mL beaker

Model 3

• Hydrochloric acid
• Thermometer
• Pen
• Paper
• Thermometer
• Distilled water
• 250 mL beaker
• Styrofoam cup

• Calculator
• Sodium hydroxide

Procedure

Model 1

1. First, put 2 styrofoam cups in a 250 mL beaker . This makes a calorimeter.
2. Break up the NH4Cl and measure the mass.
3. With deionized water wash the graduated cylinder.
4. Looking at Table 1 on the worksheet, measure the right amount of deionised water.
5. Write down the volume of the water in the data table.
6. Pour the deionised water into the calorimeter made in step 1.
7. Calculate the mass of the solid reactant and write it down.
8. Transfer the solid reactant into the calorimeter once the temperature readings have stabilized.
9. Stir the calorimeter by rotating it.
10. Stop collecting data once the temperature measurements have stabilized again.
11. Wash the cup and graduated cylinder (rinse with deionized water) and properly dispose of the solution.
12. For the second run (to confirm your results), repeat the procedure.
13. Share the information gathered with your classmates.

Model 2

1. Make the calorimeter with the 250 mL beaker and 2 Styrofoam cups (using model 1).
2. Reactant 1 is measured to the specified volume. Pour Reactant 1 into the calorimeter after recording the value in the data table.
3. Flush out the graded cylinder.
4. Reactant 2 is measured, and the volume is recorded in the data table.
5. Reactant 2 should be poured into the calorimeter once the temperature readings have stabilized. Gently twirl the calorimeter to stir.
6. Stop collecting data once the temperature measurements have stabilized once more.

• Clean the cup and graduated cylinder, then rinse them with deionized water before properly disposing of the solution.
• Run the experiment once more to get a second run.
• For each data run, ascertain the starting and maximum temperatures. In the data table, note the values.
• Share the information gathered with your classmates.

Model 3

1. Construct the calorimeter like step 1 of Model 1 .
2. Use deionized water to clean the graduated cylinder.
3. based on the prescribed reaction, obtain reactant 1.
4. Fill the calorimeter with the graduated cylinder’s contents.
5. Pick up reactant 2.
6. Take notes.
7. Reactant 2 should be added to the calorimeter when the temperature has stabilized.
8. Gently stir; when the temperature returns to normal, halt the data collection.
9. Consider the safety precautions and the teachers’ directions while disposing of the solution.
10. Cleanse the graduated cylinder and cup, then give them a last rinse in deionized water.
11. To obtain better data, conduct the experiment a second time.
12. While taking safety precautions into account, clean up any materials.
13. Share the information gathered with your classmates.

Safety and Ethical Issues

• Avoid touching the solid NaOH and use caution when handling the resultant solution.
• NaOH is corrosive and will burn you.
• Rinse with a lot of running water right away if NaOH or HCl solutions get on your skin or in your eyes.
• Be sure to adhere to the established safety guidelines and practices when handling, storing, and discarding chemicals.
• Make sure to wear safety goggles, lab coats, and gloves to minimize injury.
• Only use chemicals in the amounts and for the purposes intended.
  • While doing the lab do not eat or drink anything.

DATA ( model 1)

Reaction	Run	Mass of NH4CL (g) ±0.01	Volume of water (mL) ±0.5	ΔT (℃) ±1.0	Heat q gained” or lost (kJ)	Number of moles of NH4Cl (mol)	ΔH (kJ/mol)	Average ΔH (kJ/mol)
A	Run 1	2.0	50.5	-2.6	0.54935	0.037	14.84	14.93
	Run 2	2.0	49.2	-2.7	0.55581	0.037	15.02
B	Run 1	2.0	100.2	-1.1	-0.461	0.037	12.46	11.88
	Run 2	2.0	100.0	-1.0	-0.418	0.037	11.30
C	Run 1	4.00	50.01	-4.99	-1.044	0.075	-13.92	-13.545
	Run 2	4.00	50.01	-4.72	-0.987	0.075	-13.17
D	Run 1	4.0	100.0	-2.5	-1.088	0.0748	14.54	15.985
	Run 2	4.0	100.0	-3.0	-1.304	0.0748	17.43
E	Run 1	6.0	50.1	-8.8	1.845	0.075	27.54	27.07
	Run 2	6.0	50.1	-8.5	1.782	0.075	26.601
F	Run 1	6.0	100	-5.0	2.090	0.112	18.66	14.8
	Run 2	6.0	100	-3.4	1.255	0.112	10.94

(MODEL 2) TABLE 4A & 4B

Reaction	Run	1.0 M HCL (aq) +-0.5	1.0 M NaOH (aq) +-0.5	Initial temperature +-0.2	Final temperature +-0.1
1	Run 1	10.1	50.0	19.5	23.0
	Run 2	10.5	50.0	18.1	21.0
2	Run 1	20.0	40.0	20.0	23.0
	Run 2	20.0	40.0	20.0	22.0
3	Run 1	30.0	30.0	20.0	25.0
	Run 2	30.0	30.0	19.0	24.5
4	Run 1	39.8	19.7	20	24
	Run 2	40	19.9	19.7	23.7
5	Run 1	50.0	10.0	19.0	21.1
	Run 2	50.0	10.0	19.0	20.9

Ratio	Run	Temperature change	Heat q (kJ)	Mol HCL	Mol NaOH	Heat of reduction (delta H) (kJ/mol)	Average heat of reaction (delta H) (kJ/mol)
1	Run 1	3.5	0.147	0.277	1.25	0.533	0.523
	Run 2	2.9	0.127	0.288	1.25	0.513
2	Run 1	3.0	752.4	0.02	0.04
	Run 2	2	501.6	0.02	0.04
3	Run 1	5
	Run 2	5.5
4	Run 1	4	0.99	1.09	0.49	2.02	2.01
	Run 2	4	1	1.10	0.50	2
5	Run 1	2.1

Run 2

1.9

TABLE 3

Parameters	Reaction 1		Reaction 2		Reaction 3
	Run 1	Run 2	Run 1	Run 2	Run 1	Run 2
Initial temp	19.3	18.2	20	20	19.5	20
Final temp	32.0	28.1	24.5	26	31.5	37.5
Mass of solid NaOH (g)/ (mL)/(g) (1/2/3)	4.02	4.03	100	100	4.05	4.03
Volume of water (mL)/ HCl/HCl (1/2/3)	100.0	100.0	200	200	100	100
Mass of solution (g)	100.0	100.0	4.5		104.05	104.03
Change of temp	12.7	9.9	4.5	6.0	11.0	17.5
Heat q (kJ)	5.3086	4.1382
Number of moles of NaOH	0.1005	0.10075
Number of moles of HCl (not for reaction 1)	X	X
Molar heat of reaction (change in H) (kJ/mol)	52.818	41.074
Average molar heat of reaction (kJ/mol)	46.964

Conclusion

The three experiments done offer important new understandings into the kinetics and thermodynamics of chemical processes.

Understanding the energy changes that take place during a chemical reaction requires doing the first experiment, which involves measuring the heat change or enthalpy of a process. The relative stabilities of reactants and products can be ascertained using the enthalpy change, which can also be used to forecast whether a reaction will be exothermic or endothermic.

The second experiment, which involves modifying the limiting reactant, serves as an example of how stoichiometry affects reaction outcomes. The experiment demonstrates how the limiting reactant can impact the yield and amount of product obtained by adjusting the number of reactants added. This knowledge is critical for identifying the ideal circumstances for a reaction and enhancing its effectiveness.

The third experiment illustrates the value of thermodynamic concepts in estimating the enthalpy change of a reaction by using Hess’s Law. Hess’s Law can be used to determine the overall enthalpy change by dissecting the reaction into smaller phases.

Analysis

Typical sources of inaccuracy in heat labs include:

• Complete combustion: If the combustion reaction is not complete, the enthalpy value may be lower than it actually is.
  • Heat loss: During the experiment, heat could be lost to the environment, lowering the enthalpy value. By utilizing well-insulated containers and carrying out the experiment in a thermally stable setting, this can be reduced.
  • The computation of enthalpy change makes the assumption that the pressure stays constant throughout the reaction. However, the enthalpy value may alter if pressure changes take place during the process.
  • Inaccurate measurements: Inaccurate enthalpy readings might result from miscalculating the mass of the reactants or products, the temperature change throughout the reaction, or both.
  • Insufficient data: During this lab, the majority of the groups failed to contribute their data to the data table, which had an impact on some of the calculations and analyses we were able to make simply by looking at the data table.

Some actions that can be performed to increase the accuracy of heat labs including enthalpy include:

• Carrying out several trials: the influence of random errors can be lessened and a more accurate average value can be provided by doing many trials.
  • Using more accurate measurement tools: Using more accurate tools for mass and temperature measurements helps lower measurement errors.

• Environmental management Heat loss to the environment can be minimized by carrying out the experiment in a controlled setting with small temperature variations.

Cited Sources

“Laboratory Safety.” EHS, www.ehs.washington.edu/research-lab/laboratory-safety#:~:text=Laboratory%20safety%20practi ces%20include%20appropriate,practices%20and%20safe%20working%20conditions. Accessed 15 May 2023.

Lab Session 09 – University of Louisiana at Monroe, www.ulm.edu/chemistry/courses/manuals/chem1009/session_09.pdf. Accessed 15 May 2023.

Keith E-Books, www.keithebooks.com/sites/default/files/ebooks/Chemistry%20IB%20Diploma.pdf. Accessed 15 May 2023.

“Enthalpy.” Encyclopædia Britannica, www.britannica.com/science/enthalpy. Accessed 15 May 2023.