Use Hess's Law To Find Δh For The Reactions

Use Hess's Law To Findδh For The Following Reactionsexpress Your An

Use Hess's Law to find ΔH for the following reactions: Express your answers in kJ/mol of the first reactant on the left in each equation. The task involves calculating the enthalpy change for given chemical reactions using Hess's Law, which states that the total enthalpy change for a reaction is the same, no matter how it occurs, provided the initial and final conditions are the same. This typically involves combining known thermodynamic data for related reactions to determine the enthalpy change for the target reaction.

In addition to these calculations, a laboratory-style hypothetical experiment is described, focusing on the combustion of magnesium, which includes applying Hess’s Law with experimental data, standard enthalpies of formation, and calorimetric measurements. The experiment demonstrates how to determine ΔH for reactions that are difficult to measure directly by manipulating related reactions with known enthalpy changes, particularly through calorimetric methods involving magnesium oxidation and combustion.

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Hess’s Law is a fundamental principle in thermodynamics, enabling chemists to determine the enthalpy change of complex reactions by summing the enthalpy changes of simpler, related reactions. Its application is especially useful for reactions where direct measurement is impractical or impossible, such as combustion or formation reactions under standard conditions. The law is mathematically expressed as:

ΔH_reaction = Σ ΔH_steps

where the sum of the enthalpy changes of the known steps equals the overall enthalpy change for the reaction. This principle allows the calculation of ΔH for reactions by manipulating equations with known enthalpy values, often derived from experimental data or thermodynamic tables.

In the context of the laboratory experiment involving magnesium, Hess’s Law exemplifies how thermodynamic principles are utilized in practical settings. The experiment involves using calorimetric measurements to determine the heat released during magnesium oxidation or combustion, then applying Hess’s Law to relate this data to theoretical enthalpy values obtained from standard tables.

The process begins with reactions involving magnesium oxide and magnesium metal reacting with hydrochloric acid, where temperature changes in calorimeters provide data for calculating the heat released or absorbed. These experimental heats (ΔH) are computed using the relationship:

Q = mcΔT

where m represents the mass of the solution, c is the specific heat capacity (approximated as that of water), and ΔT is the temperature change. Converting these heats from joules to kilojoules, and normalizing per mole of reactant, yields the enthalpy change per mole, which can then be related to the overall reaction via Hess’s Law.

Furthermore, the experiment utilizes standard enthalpies of formation from thermodynamic tables to calculate the theoretical heat of combustion or formation for magnesium. The standard enthalpy of formation for magnesium oxide, for example, is a key value used in combining reactions to derive the enthalpy change of the overall process. Comparing experimental data with theoretical values allows for evaluation of experimental accuracy and calculation of percentage errors, enhancing understanding of thermodynamic principles in real-world systems.

In conclusion, applying Hess’s Law to calorimetric data and standard thermodynamic values provides a comprehensive approach to understanding reaction energetics. This method is invaluable in chemical thermodynamics, especially for reactions where direct measurements are challenging, and underscores the importance of thermodynamic data in predicting reaction behavior and energy changes in chemical processes.

References

• Atkins, P., & de Paula, J. (2010). Physical Chemistry (9th ed.). Oxford University Press.
• Zumdahl, S. S., & Zumdahl, S. A. (2014). Chemistry: An Atoms First Approach. Cengage Learning.
• Atkins, P., & de Paula, J. (2014). Physical Chemistry (10th ed.). Oxford University Press.
• Skoog, D. A., West, D. M., Holler, F. J., & Crouch, S. R. (2014). Fundamentals of Analytical Chemistry. Brooks Cole.
• McQuarrie, D. A., & Simon, J. D. (1997). Physical Chemistry: A Molecular Approach. University Science Books.
• Laidler, K. J., Meiser, J. H., & Sanctuary, J. M. (1999). Physical Chemistry (3rd ed.). Houghton Mifflin.
• Chang, R., & Goldsby, K. A. (2016). Chemistry (12th ed.). McGraw-Hill Education.
• Chenoweth, C. K. (2015). Principles of Thermodynamics. Dover Publications.
• Rutherford, D., & Lindner, A. (2020). Thermodynamics in Practice: A Laboratory Manual. Springer.
• Standard Enthalpies of Formation of Compatible Substances, NIST Chemistry WebBook, National Institute of Standards and Technology, https://webbook.nist.gov/chemistry/.