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Energy is a fundamental topic in the fields of physics and chemistry. In this experiment, we utilized a coffee-cup calorimeter in three distinct reaction phases to determine the enthalpies of various reactions and calculate the heat of solutions. Part One of the experiment yielded an exothermic reaction with a heat of reaction equal to 37.3203 kJ/mol. Parts Two and Three, on the other hand, resulted in endothermic reactions with heat of reactions equaling -57.9185 kJ/mol and -453.9854 kJ/mol, respectively.
Understanding heat energy in chemical reactions is crucial, and physical chemistry serves as a bridge connecting the realms of physics and chemistry. Additionally, ensuring thorough solvent mixing is essential to obtain precise experimental values.
Energy can be defined as the rapid movements of molecules in matter, leading to collisions or vibrations that generate light and/or heat. Energy exists in various forms in the universe, including light, sound, nuclear, and chemical energy. Each type of energy can produce heat through molecular motion (IPAC, n.d.).
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This understanding is vital in chemistry as it explains the heat production or absorption during reactions. Heat energy is not lost or destroyed; instead, it is transferred from one entity to another. The change in heat, known as "enthalpy," allows chemists to comprehend how different elements and compounds interact. Heat can be either absorbed or released in a chemical reaction. An endothermic heat change indicates an exothermic reaction, whereas an exothermic heat change signifies an endothermic reaction (University of Massachusetts Boston, 2019).
The measurement of heats of reactions can be accomplished using a calorimeter.
A calorimeter is a device designed "to measure the quantity of heat flow in a chemical reaction" (Helmenstine, 2019). Many laboratories prefer a cost-effective and user-friendly calorimeter that can be assembled using common household and laboratory items. The coffee-cup calorimeter is easily and inexpensively constructed and is widely used in educational laboratories due to its affordability and ease of maintenance. It consists of two nested Styrofoam cups, with a small square fiberboard piece placed over the cups to prevent rapid heat loss and splashing during stirring or reactions. A hole in the center of the fiberboard accommodates a thermometer for monitoring temperature changes. A small stir bar continuously agitates the solution in the nested cups, facilitating rapid and complete dissolution. The magnetic plate beneath the cups propels the stir bar when activated. Calorimeters can determine the heat capacity of a reaction, with "heat capacity" defined as "the amount of energy required to raise the temperature of 1 g of a substance by 1°C" (University of Massachusetts Boston, 2019).
First, the mass of two nested Styrofoam cups and a small stir bar was measured and recorded on a data sheet using an analytical balance. Approximately 100 milliliters (mL) of deionized water was measured with a 100 mL graduated cylinder and carefully poured into the nested cups along with the stir bar. The combined mass of the water, stir bar, and Styrofoam cups was measured on the same analytical balance and recorded. The mass of the water was determined by subtraction.
Subsequently, the mass of an aluminum dish was measured on an analytical balance and recorded. 2 grams of potassium nitrate (KNO3) were added to the dish, and the mass of the dish and potassium nitrate was measured and recorded. The mass of potassium nitrate was calculated by subtraction. The coffee-cup calorimeter was then assembled following the appropriate configuration as depicted in Figure 1 above. The stirrer was activated by turning on the magnetic plate, ensuring efficient stirring without splashing. The initial temperature (Ti) of the water was measured and recorded using the thermometer.
All of the potassium nitrate was added rapidly to the coffee-cup calorimeter, and the calorimeter was covered with the fiberboard lid, with the thermometer positioned inside. The reaction mixture was stirred continuously for approximately 1 minute until it reached the minimum temperature. The final temperature (Tf) of the reaction was measured and recorded.
The coffee-cup calorimeter was emptied into an appropriate waste container, and all equipment in contact with the solution (inner cup, thermometer, and stir bar) was rinsed thoroughly with deionized (DI) water and dried. The coffee-cup calorimeter was then reassembled, with the stir bar placed back into the nested cups.
About 50 mL of 0.1 M hydrochloric acid (HCl) was measured using a graduated cylinder and recorded on a data sheet, and then it was emptied into the nested cups. Using a separate clean graduated cylinder, 30-40 mL of 1.00 M sodium hydroxide (NaOH) was measured, and the volume was recorded. The stirring motor was activated to achieve rapid stirring without splashing, and the initial temperature (Ti) was measured and recorded.
The sodium hydroxide solution was then carefully poured into the inner cup of the coffee-cup calorimeter containing the hydrochloric acid. The fiberboard cover was placed over the cups, and the solution was stirred for approximately 1 minute. The maximum temperature was recorded on the data sheet as Tf.
The coffee-cup calorimeter was emptied into the appropriate waste containers, and all equipment in contact with the solution (inner cup, thermometer, stir bar, and graduated cylinder) was rinsed with DI water and thoroughly dried.
Using an analytical balance, the mass of the nested cups and stir bar together was determined and recorded on a data sheet. Subsequently, 75 mL of 1.00 M hydrochloric acid was measured using a clean graduated cylinder and transferred into the nested Styrofoam cups. On the same analytical balance, the mass of the cups, hydrochloric acid, and stir bar was measured and recorded. The mass of hydrochloric acid was calculated by subtraction.
A 15 cm cleaned strip of magnesium was folded to fit into the cup. The coffee-cup calorimeter was reassembled appropriately as shown in Figure 1. The stirring motor was activated to achieve rapid stirring without splashing, and the temperature of the hydrochloric acid was recorded as Ti. The folded magnesium strip was added to the hydrochloric acid solution, and the fiber-board was placed over the cups with the thermometer inside. The solution was stirred until a maximum temperature was reached, and this value was recorded as Tf.
Following the completion of each part of the experiment, the contents of the coffee-cup calorimeter were emptied, and the lab area and equipment were thoroughly rinsed and dried. The apparatus was then reassembled, ready for use by the next lab section (University of Massachusetts Boston, 2019).
| Measurement | Value |
|---|---|
| Mass of water | 97.8218 g |
| Mass of KNO3 | 1.9221 g |
| Mass of Solution | 99.7439 g |
| Heat Capacity of Calorimeter, Ccal | 417.3285 J/˚C |
| Initial temperature, Ti | 21.7˚C |
| Final temperature, Tf | 20˚C |
| Temperature change, ΔT | 1.7˚C |
| Calorimeter's heat change, qcal | 709.4585 J |
| Solution heat change, qsoln | 709.4585 J |
| Heat of Solution per mole KNO3 dissolved | 37.3203 kJ/mol |
| Moles of KNO3 dissolved | 0.01901 mol |
| Endothermic or Exothermic? | endothermic |
| % error | 6.97% |
| Measurement | Value |
|---|---|
| Volume of 0.1 M HCl(aq) | 50 mL |
| Volume of 1.00 M NaOH(aq) | 35 mL |
| Moles of NaOH | 0.035 mol |
| Mass of solution (1 mL=1 g) | 85 g |
| Heat capacity of calorimeter, Ccal | 355.64 J/˚C |
| Initial temperature, Ti | 21.8˚C |
| Final temperature, Tf | 27.5˚C |
| Temperature change, ΔT | 5.7˚C |
| Neutralization heat change | 2027.148 J |
| Moles of H3O+ neutralized with OH- | 0.8750 mol |
| Heat of neutralization per mole of H3O+ | -57.9185 kJ/mol |
| Endothermic or Exothermic? | exothermic |
| % error | 0.97% |
| Measurement | Value |
|---|---|
| Mass of HCl(aq) solution | 73.803 g |
| Mass of Mg(s) strip | 0.2905 g |
| Mass of solution | 74.0935 g |
| Heat capacity of calorimeter, Ccal | 310.0072 J/˚C |
| Initial temperature, Ti | 22.0˚C |
| Final temperature, Tf | 39.5˚C |
| Temperature change, ΔT | 17.5˚C |
| Reaction's heat change | 5425.126 J |
| Moles of Mg(s) reacted | 0.01195 mol |
| Heat of reaction per mole Mg(s) reacted (kJ/mol) | -453.9854 kJ/mol |
| Calculated Heat of reaction per mole Mg(s) reacted | -462.5 kJ/mol |
| % error | 1.84% |
The laboratory experiment was divided into three distinct parts to investigate various types of reactions and to practice calculating enthalpy-related values. In Part One, the mixture of 100 mL of water with 1.92 grams of potassium nitrate resulted in a temperature decrease of approximately 1.7 degrees Celsius. The calculated heat of solution per mole of potassium nitrate dissolved was 37.32 kJ/mol. This indicates that the reaction was endothermic as heat was absorbed. The experimental error for this part of the experiment was calculated to be 6.97%.
In Part Two, the combination of 50 mL of 0.1 M hydrochloric acid with 35 mL of 1.00 M sodium hydroxide in the coffee-cup calorimeter led to a temperature increase of 5.7 degrees Celsius. The calculated heat of neutralization per mole of H3O was -57.92 kJ/mol. This reaction was identified as exothermic because heat was released. The experimental error for this part was found to be 0.97%.
In the third and final part of the experiment, 75 mL of hydrochloric acid reacted with a 0.2905-gram strip of solid magnesium. The dissolution of the magnesium strip in hydrochloric acid caused a temperature increase of 17.5 degrees Celsius. The calculated heat of reaction per mole of magnesium reacted was -453.98 kJ/mol. The percent error was determined to be 1.84%. This reaction was also identified as exothermic since heat was released.
Throughout each part of the experiment, various sources of error were encountered. To evaluate the experimental results, it is essential to compare them to literature values. The expected temperature change in a given reaction can be estimated by conducting the following calculations:
For Part One:
Expected temperature change = (moles of KNO3 dissolved) × (ΔHsolution / Ccal)
Expected temperature change = (0.01901 mol) × (37.32 kJ/mol / 417.3285 J/˚C) = 1.678˚C
For Part Two:
Expected temperature change = (moles of H3O+ neutralized with OH-) × (ΔHneutralization / Ccal)
Expected temperature change = (0.8750 mol) × (-57.92 kJ/mol / 355.64 J/˚C) = -14.079˚C
For Part Three:
Expected temperature change = (moles of Mg(s) reacted) × (ΔHreaction / Ccal)
Expected temperature change = (0.01195 mol) × (-453.98 kJ/mol / 310.0072 J/˚C) = -17.543˚C
These calculated expected temperature changes provide a basis for assessing the experimental results. The discrepancies between the expected and observed temperature changes highlight the presence of errors or factors not accounted for in the experiment. Such discrepancies may include heat loss to the surroundings, incomplete reactions, or variations in the heat capacity of the calorimeter.
Despite these challenges, the experimental values for the enthalpies of the reactions closely align with the expected trends. The endothermic and exothermic nature of the reactions was successfully determined through the measured temperature changes, consistent with theoretical expectations.
Overall, this laboratory experiment served as a valuable learning experience in the field of thermodynamics and enthalpy calculations, providing insight into the energy changes associated with different chemical reactions.
Part One: Heat of Solution
The observed temperature change in Part One differed from the expected value of -1.5895°C based on the literature value for potassium nitrate. One possible source of error in this experiment is the incomplete transfer of potassium nitrate from the aluminum dish to the calorimeter. Some of the solid may have adhered to the dish's base, resulting in less potassium nitrate reacting with the water. To address this issue, it is crucial to ensure that all crystalline solids are effectively transferred into the calorimeter by gently shaking or tapping the dish.
Part Two: Heat of Neutralization
In Part Two, the observed temperature change did not precisely match the expected value of +5.6457°C based on the literature value for H3O+. A potential source of error in this reaction is the incomplete transfer of hydrochloric acid and sodium hydroxide from the graduated cylinders to the calorimeter. Residual droplets on the cylinder walls may have resulted in a smaller quantity of solutions reacting, causing a slight deviation from the expected temperature change.
Part Three: Heat of Reaction of Magnesium and Hydrochloric Acid
In Part Three, the observed temperature change did not align with the expected value of +17.8229°C based on the calculated literature value for solid magnesium. One likely source of error in this portion of the experiment is the discrepancy in the mass of the magnesium strip. The measured mass of 0.2905 grams differed from the intended 0.1 grams, leading to variations in the temperature change. Ensuring a more precise measurement of the magnesium strip's mass, possibly by adjusting its length, could result in a closer match between the observed and expected temperature changes.
In summary, this experiment focused on studying heat energy transfer by combining different compounds to determine the heat of reactions for various chemical processes. The calculated heat of reactions for Parts One, Two, and Three were found to be 37.3203 kJ/mol, -57.9185 kJ/mol, and -462.9854 kJ/mol, respectively. The percent errors for these parts were 6.9656%, 0.9737%, and 1.841%, respectively. Part One was identified as an endothermic reaction since it absorbed heat, while Parts Two and Three were both exothermic reactions, as they released heat.
To improve the accuracy of these experiments, it is essential to ensure the complete transfer of solutes into the solvent to achieve more precise values. The determination of enthalpies for various reactions enhances our understanding of heat energy interactions when different compounds are combined, contributing to the broader knowledge of thermochemistry.
Thermochemistry Experiment: Heat of Reactions and Enthalpy Changes. (2024, Jan 16). Retrieved from https://studymoose.com/document/thermochemistry-experiment-heat-of-reactions-and-enthalpy-changes
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