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7.3.5: Discussion questions

  • Page ID
    424810
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    The previous section (data analysis) describes how to handle your data. You should perform the described calculations and analysis and describe them in your notebook. In addition, the following should be addressed in your notebook.

    1. Perform a full propagation of errors analysis of your results and suggest possible sources of systematic error.
    2. Begin your discussion with a quantitative comparison of your experimentally measured values with the literature values of the enthalpy of solution for \(\ce{CH3COONa}\), \(\ce{CH3COONa*3H2O}\), and \(\ce{NaCl}\) (found in the CRC handbook). Do your results agree with accepted values within your estimated uncertainty? If not, suggest possible reasons for any differences found.
    3. Discuss the significance of the sign of \(\Delta T_c\).
    4. Relate the differences in the enthalpies of solution of the two sodium acetate salts to differences in their respective lattice energies. To do this, you might begin by drawing a Hess’s Law type diagram that relates the enthalpy of solution of the solids to the lattice energies and the hydration energies of the gaseous ions. Look up the lattice energy of anhydrous sodium acetate in the CRC Handbook. Note: Use the \(U^{BFHC}_{pot}\) values)
    5. Use the enthalpies of solution of the two sodium acetate salts to calculate the enthalpy change for the following reaction: \[\ce{CH3COONa_{(s)} + 3 H2O --> CH3COONa*3H2O_{(s)}}\] From your result, calculate the lattice energy of sodium acetate trihydrate. Returning to your Hess’s Law diagram, use your results so far and the known gaseous ion hydration energy of \(\ce{Na^+}\) to calculate the gaseous ion hydration energy of the acetate anion (see your instructors if you need help). Compare this with an accepted value. How well does it agree?
    6. Consider now sodium chloride and potassium chloride. You have an experimental value for the enthalpy of solution of sodium chloride. In addition, you used the enthalpy of solution of potassium chloride (230.98 J/g) in the calibration of the calorimeter. Convert the enthalpy of solution of KCl to kJ/mole and compare it to the enthalpy of solution of NaCl. In discussing these results, you should do a similar calculation to what you did for the acetate salts, and calculate their respective lattice energies from your enthalpies of solution and the known gaseous ion hydration energies. In addition, you should relate the differences in lattice energies to the relative sizes of the sodium and potassium cations. What is the general trend that we would expect to see in lattice energy as it relates to the size of the ions?
    7. Why is it appropriate to approximate the conditions in this experiment to infinite dilution? (see Alberty Section 2.21). Look again at the figure showing Integral Heats of Solution (Fig 6.3.1.1). In your experiments, were the sample masses that you used appropriate for an approximate measurement of the enthalpy of solution at infinite dilution? Considering the technical specifications of the calorimeter and the calorimetric thermometer, given in the Instrumentation Specifications, what is the consequence of using sample sizes small enough to yield enthalpies of solution consistent with infinite dilution?

    This page titled 7.3.5: Discussion questions is shared under a CC BY-NC-SA 4.0 license and was authored, remixed, and/or curated by Kathryn Haas.

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