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Exploring Enthalpies of Dissolution: Calorimetry Insights into Exothermic Reactions

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Procedure:

Part I:

Begin by setting up the calorimeter on a magnetic stirrer. Measure 100.0 mL of water in a graduated cylinder and pour it into the calorimeter.
Measure and record the initial temperature of the water.
Add a stir bar and turn on the stirrer to ensure gentle stirring without splashing.
Weigh 5.00g of Magnesium Sulfate anhydrous solid in a plastic cup. Slowly add it to the calorimeter while monitoring the temperature. Record the highest temperature reached.
Dilute the resulting solution with water and dispose of it.
Repeat the above steps, and average the recorded temperatures.

Part II:

Take a clean, dry 150 mL beaker and place 100.0 mL of water in it.

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Heat the water with occasional stirring until it reaches approximately 50 degrees Celsius.
Remove the beaker from the hot plate and place it on the lab bench. Simultaneously, place 100 mL of cool water in the clean, dry calorimeter.
Measure and record the temperatures of both the hot and cold water. Pour the hot water into the calorimeter immediately, cover it quickly.
After 15 seconds, measure and record the temperature.
Repeat the above steps.

Part III:

Take safety notes on each substance provided.
Rank the substances based on their cost.
Follow the steps in Part I to determine the heat released by the substance.

Prelab Questions:

The sodium and chloride ions undergo separation, with water molecules surrounding the ions and separating from each other.

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Bonds between ions and water molecules break, and new bonds form between cations, anions, and water molecules.
Heat is absorbed in this process.

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The energy required to separate sodium and chloride ions and water molecules surpasses the energy released from ion-water attraction. The difference is absorbed as heat, leading to a decrease in water temperature.
The thermodynamically favorable increase in disorder causes dissolution to occur spontaneously despite the overall endothermic nature of the reaction.
The temperature change during the solution formation depends on the relative energy required to separate ions in the crystalline lattice solid and the energy released during the formation of water-ion attraction. If the energy needed for ion separation is greater, heat is absorbed, resulting in a lower solution temperature, and vice versa.

In this laboratory experiment, we will explore the determination of enthalpies of dissolution for various salts through calorimetry. Enthalpy of dissolution refers to the amount of heat energy absorbed or released when a certain amount of solute dissolves in a solvent to form a solution. This is a crucial parameter in understanding the thermodynamics of solution formation. Calorimetry, the measurement of heat changes, provides a reliable method to determine these enthalpies accurately.

Materials:

Various salts (e.g., NaCl, KNO3, CuSO4)
Distilled water
Calorimeter
Thermometer
Stirring rod
Weighing balance
Stopwatch

Methods:

Preparation of Solutions:
- Weigh a precise amount of each salt separately.
- Dissolve the weighed salt in a known volume of distilled water to prepare a concentrated solution.
Calorimeter Setup:
- Place the calorimeter on a stable surface.
- Add a known volume of distilled water to the calorimeter.
Initial Temperature Measurement:
- Measure and record the initial temperature of the water in the calorimeter.
Salt Addition:
- Add the prepared concentrated salt solution to the water in the calorimeter.
- Stir the solution gently to ensure homogeneity.
Final Temperature Measurement:
- Measure and record the final temperature of the solution in the calorimeter.

Calculations:

Change in Temperature (∆T): initialΔT=Tfinal−Tinitial
Heat Absorbed or Released (q): q=mcΔT where:
- q is the heat absorbed or released.
- m is the mass of the water.
- c is the specific heat capacity of water (4.18 J/g°C).
- ΔT is the change in temperature.
Enthalpy of Dissolution (∆H): ΔH=moles of soluteq
Moles of Solute: Moles=Mass of soluteMolar mass of soluteMoles=Molar mass of soluteMass of solute

Data:

Part I:

Cold and Hot Water Temperatures
	Initial Temperature	Final Temperature	Change
Hot Water	49.1	36.6	-12.5
Cold Water	23.2	36.6	13.4

Part II:

Calorimeter Constant Data
Mass hot water	100.0
Mass cold water	100.0
Initial temperature cold	21.2
Initial temperature hot	53.0
Final temperature of mixture	36.8
Magnesium Sulfate Dissolution Data
Initial temperature	20.0
Final temperature	27.7

Part III:

Investigation Data
Solid	Initial temp	Final temp
NaC₂H₃O₂	20.0	24.8
CaCl₂	20.0	36.0
Na₂CO₃	20.0	25.0
NaCl	20.0	18.6
LiCl	20.0	38.0
NH₄NO₃	20.0	12.5

Part IV:

Calorimeter Constant Calculations
Temperature change of hot water	= 36.8 - 53.0 = -16.2
Enthalpy change of hot water	= (100.0g)(4.184)(-16.2) = -6780
Temperature change of cold water	= 36.8 - 21.2 = 15.6
Enthalpy change of cold water	= (100.0g)(4.184)(15.6) = 6530
Temperature change of calorimeter	= 36.8 - 21.2 = 15.6
Enthalpy change of calorimeter	= 6780 - 6530 = 250
Calorimeter constant	= 250 / 15.6 = 16.0

Part V:

Investigation Calculations Table (Based on Data Above)
Solid	Temperature change	Thermal energy change in calorimeter contents	Thermal energy change of calorimeter	Internal energy change	Molar mass	Moles used	Enthalpy of dissolution
NaC₂H₃O₂	4.8	1000	77	-1080	82.03	0.061	-17.7
CaCl₂	16.0	3350	256	-3600	111.10	0.0450	-80.1
Na₂CO₃	5.0	1050	80	-1130	106.00	0.047	-23.9
NaCl	-1.4	-290	-22	310	58.45	0.086	3.7
LiCl	18.0	3770	288	-4050	42.39	0.118	-34.4
NH₄NO₃	-7.5	-1570	-120	1690	80.04	0.062	27.0

In the realm of thermodynamics, understanding the enthalpy of dissolution for various salts through calorimetry experiments provides valuable insights into the energy changes associated with chemical reactions. This laboratory delves into the exploration of exothermic reactions involving salts such as NaCl, NH4NO3, NaC2H3O2, Na2CO3, LiCl, and CaCl2. By examining the enthalpies of dissolution, we gain a deeper understanding of the heat transfer associated with the breaking of bonds in these salts when mixed with water.

Calorimetry and Enthalpy:

Calorimetry, the measurement of heat changes in a system, serves as the cornerstone of this experiment. The enthalpy change ( $Δ H$ ) for a reaction is determined by measuring the heat absorbed or released during the process. This is represented by the equation:

$q = m c Δ T$

where:

$q$ is the heat absorbed or released.
$m$ is the mass of the system.
$c$ is the specific heat capacity of the substance.
$Δ T$ is the change in temperature.

The enthalpy change ( $Δ H$ ) is then calculated using the equation:

$Δ H = moles of solute q$

Exothermic Reactions:

The reactions under investigation are deemed exothermic because they release heat into the system. The enthalpy values associated with each reaction indicate the amount of heat generated per mole of solute. Exothermic reactions are characterized by a negative enthalpy change, as they release energy to the surroundings.

NaCl (s) ⟶ Na⁺ (aq) + Cl⁻ (aq): $Δ H = + 3.90 kJ/mol$
NH₄NO₃ (s) ⟶ NH₄⁺ (aq) + NO₃⁻ (aq): $Δ H = + 25.48 kJ/mol$
NaC₂H₃O₂ (s) ⟶ Na⁺ (aq) + C₂H₃O₂⁻ (aq): $Δ H = - 17.32 kJ/mol$
Na₂CO₃ (s) ⟶ 2Na⁺ (aq) + CO₃²⁻ (aq): $Δ H = - 24.80 kJ/mol$
LiCl (s) ⟶ Li⁺ (aq) + Cl⁻ (aq): $Δ H = - 37.10 kJ/mol$
CaCl₂ (s) ⟶ Ca²⁺ (aq) + 2Cl⁻ (aq): $Δ H = - 81.50 kJ/mol$

The observed changes in these reactions are unequivocally chemical in nature. When salts dissolve in water, they undergo ionization, breaking apart into cations and anions. This process necessitates a chemical reaction to sever the existing bonds between the elements within the compound. The dissolution of the salts is not merely a physical change but a fundamental transformation at the molecular level.

Unfortunately, a class set of data is missing, rendering a comprehensive analysis of percent error or averages unattainable. The absence of this data inhibits the ability to gauge the experimental accuracy and precision across multiple trials. Future iterations of this experiment could benefit from a collaborative effort to pool class data for a more robust analysis.

Several sources of error may have influenced the outcomes of the experiment, impacting the accuracy and reliability of the enthalpy values obtained. A significant potential source of error lies within the calorimeter itself. If the calorimeter was not calibrated accurately, the entire set of calculations could be skewed. Additionally, the calorimeter's inability to maintain a tightly closed system might have allowed heat to escape, compromising the accuracy of temperature change measurements and, consequently, enthalpy calculations.

In conclusion, this laboratory experiment offers a comprehensive exploration of the enthalpies of dissolution for various salts through calorimetry. The calculated enthalpy values provide crucial information about the energy changes associated with these exothermic reactions. The chemical nature of the changes observed underscores the importance of understanding the thermodynamics of solution formation.

While the absence of class data limits the extent of our analysis, the discussion on potential sources of error highlights the need for meticulous experimental design and execution. Through continued refinement and collaboration, future iterations of this experiment can strive for enhanced precision and reliability. This laboratory not only contributes to the foundational knowledge of thermodynamics but also encourages a critical evaluation of experimental methodologies for continuous improvement in scientific practices.

Exploring Enthalpies of Dissolution: Calorimetry Insights into Exothermic Reactions. (2024, Feb 26). Retrieved from https://studymoose.com/document/exploring-enthalpies-of-dissolution-calorimetry-insights-into-exothermic-reactions