Catalyst Education 2020 Enthalpy Of Dissolution And Neutrali ✓ Solved

Catalyst Education 2020 Enthalpy of Dissolution and Neutralization

Use simple calorimetry apparatus to determine the molar enthalpies of three different reactions. Understand that in a closed system, the heat of a reaction is equal in magnitude but opposite in sign to the change in heat of the surroundings. Directly determine the molar enthalpy of dissolution of NaOH(s) in water and use Hess’s Law to indirectly determine the same thermodynamic value from two different reactions. Compare molar enthalpy results with literature values using a percent error calculation and compare the molar enthalpy of dissolution as determined by two different methods using a percent difference calculation. Use mass and molarity as needed to mathematically determine the limiting reactant for each reaction so enthalpy per mole can be determined.

This week the enthalpy due to both physical and chemical transformations will be investigated. The first experiment will involve a physical transformation, the reorganization of molecules in forming a solution. The resulting enthalpy in kJ/mol will be called the molar enthalpy of dissolution, ΔHm,dissolution, since NaOH will be dissolving in water. The second experiment will determine the molar enthalpy from the chemical neutralization of an acid with a base, ΔHm,neutralization, again in kJ/mol.

Paper For Above Instructions

The field of thermodynamics plays a crucial role in understanding energy changes that occur during chemical reactions and physical transformations. The experiments involving the determination of molar enthalpies of dissolution and neutralization provide essential insights into these energy changes. This paper aims to detail the principles and methodologies used to determine the molar enthalpies of NaOH dissolution and the neutralization reaction between NaOH and HCl, alongside an analysis employing Hess’s Law for validation of findings.

Understanding Enthalpy Changes

Enthalpy (ΔH) is a fundamental concept in chemistry that quantifies the heat content of a system. When a solute like sodium hydroxide (NaOH) dissolves in a solvent (water), several energy-related processes occur. The breaking of solute-solute interactions within NaOH and solvent-solvent interactions in water require energy input. Conversely, the formation of solute-solvent interactions between NaOH and water releases energy. The net energy exchange determines whether the dissolution is exothermic or endothermic.

If the process gives off heat (exothermic), the temperature of the surroundings increases. In contrast, an endothermic process absorbs heat, leading to a decrease in the temperature of the surroundings. Understanding this energy exchange is vital for calculating the molar enthalpy of dissolution, defined as the amount of heat absorbed or released per mole of solute dissolved.

Experimental Procedures

The experimental procedure for determining the molar enthalpy of NaOH dissolution involves using a coffee cup calorimeter to isolate the reaction. The steps include measuring the mass of water, adding a precise amount of NaOH, and recording the temperature change of the solution using a thermometer. The heat change, q, can then be calculated using the mass of the solution, its specific heat capacity, and the change in temperature.

For the dissolution of NaOH, the equation can be represented as follows:

q_dissolution = - (mass_sol × c_sol × ΔT_sol)

By dividing this value by the number of moles of NaOH used, we can determine the molar enthalpy of dissolution (ΔH_{m,dissolution}), which can then be compared to literature values.

Neutralization Reaction

The second part of the experiment investigates the neutralization of hydrochloric acid (HCl) with NaOH. The balanced neutralization reaction can be summarized as:

NaOH(aq) + HCl(aq) → H₂O(l) + NaCl(aq)

This reaction also involves enthalpy changes, where breaking and forming bonds result in energy release or absorption. Similar to the dissolution experiment, the calorimeter is used to measure the temperature change upon mixing these two solutions, and the molar enthalpy of neutralization (ΔH_{m,neutralization}) is derived from the changes in temperature, molarity, and amount of reactants.

To establish ΔH_{m,neutralization}, we use:

ΔH_{m,neutralization} = - (mass_sol × c_sol × ΔT_sol)

Application of Hess’s Law

Hess’s Law states that the total enthalpy change of a reaction is independent of the pathway taken, meaning it can be calculated through multiple steps. In this experiment, we can leverage this law to compare the directly measured molar enthalpy of dissolution against its calculated value through the neutralization equation.

The summation of the enthalpies from individual steps will provide a comprehensive understanding of the enthalpy associated with NaOH dissolution. The calculations involving Hess's law can be expressed as:

ΔH_{m,dissolution} = ΔH_m,system - ΔH_{m,neutralization}

This relationship forms a basis for validating experimental results and understanding discrepancies between literature values and experimental data.

Calculating Percent Error and Percent Difference

Finally, comparison between experimental values and literature values through percent error and percent difference calculations enhances the reliability of our results. The percent error can be calculated using:

Percent Error = | Experimental Value - Literature Value | / | Literature Value | × 100%

Similarly, the percent difference between two methods of enthalpy determination can be calculated, reinforcing the importance of precise methodology in obtaining accurate scientific data.

Conclusion

This series of experiments successfully illustrates the determination of molar enthalpies through practical calorimetry, emphasizing the underlying principles of thermodynamics. By measuring the enthalpy changes associated with both dissolution and neutralization processes, students gain insight into the heat transfer dynamics of chemical reactions. Employing Hess’s Law not only provides a robust framework for theoretical validation but also promotes deeper understanding through comparative analyses. Future experiments can build on these principles, exploring other solutes and reactions to further enhance knowledge in thermochemical processes.

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