Heat of Formation and Hess's Law

Welcome and Overview

  • Review of calorimetry problems.

  • Introduction to enthalpy of formation and Hess's law from chapter 6 and select topics from chapter 9.

Change in Enthalpy (ΔH)

  • Absolute value of ΔH cannot be measured.

    • ΔH relates to the flow of thermal energy, which is random molecular motion.

    • A reference point is necessary for defining ΔH.

Reference Points

  • Example of altitude for flying: sea level is a standard reference.

  • In chemistry: Reference point for carbon-12 is 12 grams (Avogadro's number).

Standard Enthalpy of Formation (ΔH_f°)

  • Defined as the heat change when one mole of a compound is formed from its elements in their standard states.

    • Standard conditions: 25°C and 1 atmosphere pressure.

  • Heat of formation for an element in its most stable form is zero.

Example of Standard States

  • Oxygen:

    • O2 = 0 (standard form)

    • O3 ≠ 0

  • Carbon:

    • Graphite = 0 (standard form)

    • Diamond ≠ 0

Calculation of ΔH for Reactions

  • ΔH for a reaction at standard states:

    • Sum of ΔH_f of products - Sum of ΔH_f of reactants.

Generic Equation

  • General expression:

    • A moles of A + B moles of B → C moles of C + D moles of D

    • ΔH = (C × ΔH_f of C + D × ΔH_f of D) - (A × ΔH_f of A + B × ΔH_f of B)

Heats of Formation Table

  • Reference table for ΔH_f values (table 6.4 in Chang's book).

  • Importance of using given ΔH_f values to calculate heat of reactions.

Approximation of ΔH

  • Calculated values can approximate experimental measurements within 10%.

  • Thermodynamic functions may change with temperature; full understanding involves complex calculations beyond basic algebra.

Experimental Determination of ΔH_f

  • Heat of formation determined through calorimetry by synthesizing compounds from elements in their most stable states.

  • Example:

    • Hydrogen + Oxygen → Hydrogen Peroxide

    • ΔH for this reaction can be calculated and added to the ΔH_f table.

Hess's Law

  • States that ΔH for a reaction is equal regardless of whether it occurs in one step or multiple steps.

  • Applies particularly under constant pressure conditions.

Example Reaction:

  • C + 2H2 → CH4 (methane) cannot be measured directly through calorimetry but can be inferred using Hess's law.

Steps to Determine ΔH Using Hess's Law

  • Reverse given reactions and sum them to derive the target reaction.

  • Example with specific reactions leading to the formation of CH4 allows calculation of the ΔH from known reactions.

Sample Test Question

  • Calculating ΔH of hydrogen peroxide synthesis using provided reaction data.

Step-by-Step Method

  1. Ensure the correct positioning of substances in the equations.

  2. If necessary, reverse or adjust reaction equations to match desired moles.

  3. Add ΔH values appropriately based on the steps taken.

Additional Examples for Practice

  • Suggested homework problems to practice Hess's Law calculations with given ΔH.

Conclusion of Chapter 6

  • Overview of ΔH of solution and dissolution in lab work.

Lattice Energy

  • Defined as energy required to vaporize a mole of ionic compound into gaseous phase.

  • Heat of hydration describes the energy change when gaseous ions dissolve in water.

Example: Sodium Chloride Dissolution

  • Discusses the lab measurement of temperature change when dissolving sodium chloride in water.

  • Relation to Hess's Law for understanding dissolution reactions.

Final Notes

  • Mention of the Hindenburg disaster as a historical example linking thermochemistry to real-world events (exothermic reactions).

  • Summary of hydrogen as a potential alternative fuel source.