Ch 9 enthalpy sec 8-11 CHM1020

Chapter 9: Thermochemistry

Section 9.8-9.11

Enthalpy

  • Definition: Enthalpy, denoted as H, represents the sum of the internal energy of a system.

  • Enthalpy Change ( (\Delta H)):

    • Defined as the heat evolved in a reaction under constant pressure.

    • Formula: (\Delta H_{reaction} = q_{reaction}) at constant pressure.

    • Typically, (\Delta H) and change in energy (\Delta E) are similar; discrepancies are prevalent in reactions involving significant gas quantities.

Hess's Law

  • Properties:

    • The enthalpy change for a response is unique to that specific reaction.

    • It is an extensive property, meaning changes in reaction parameters alter (\Delta H_{rxn}).

    • If a reaction is multiplied by a factor, (\Delta H_{rxn}) is modified correspondingly.

    • Reversing a reaction changes the sign of (\Delta H_{rxn}).

    • If expressed as multiple steps, the total (\Delta H_{rxn}) equals the sum of each step’s enthalpy change.

Relationships Involving (\Delta H_{rxn})

  • From Hess’s Law:

    • An overall reaction's (\Delta H_{rxn}) equals the sum of the enthalpies of the individual steps in the reaction pathway.

    • Example Calculation:

      • Given reaction: (C_3H_6O(l) + 4O_2(g) → 3CO_2(g) + 3H_2O(g))

      • Reactions and respective values:

        1. (3C(s) + 3H_2(g) + ½O_2(g) → C_3H_6O(l), \Delta H_{rxn} = -248.4 kJ)

        2. (C(s) + O_2(g) → CO_2(g), \Delta H_{rxn} = -393.5 kJ)

        3. (H_2(g) + ½O_2(g) → H_2O(g), \Delta H_{rxn} = -241.8 kJ)

Bond Energies

  • To estimate (\Delta H_{reaction}), consider:

    • Energy needed to break existing bonds versus energy released when new bonds form.

    • Bond energy: Energy required to break one mole of a specific bond in a compound (usually measured in gas state).

      • Average bond energies provide a way to estimate (\Delta H_{rxn}) effectively when all substances are in gaseous form.

    • Endothermic process (bond breaking): (\Delta H(breaking) = (+) value).

    • Exothermic process (bond forming): (\Delta H(making) = (−) value).

  • Formula: (\Delta H_{rxn} = ,\sum (\Delta H(bonds broken)) + ,\sum (\Delta H(bonds formed)))

Standard Enthalpy of Formation ((\Delta H_f))

  • Standard state: Defines the conditions for measuring enthalpy changes (e.g. pure gases, solids, typical temperatures).

  • Definition of (\Delta H_f°): Enthalpy change when 1 mole of a substantiated compound is formed from its elements in standard states.

    • The standard enthalpy of formation for pure elements in their standard states is 0 kJ/mol.

Lattice Energy

  • Definition: Stability arising from the formation of a crystal lattice (electrostatic attractions between cations and anions).

  • Measurement: Represents energy released when a solid forms from its gaseous ions and is an exothermic process.

    • It's influenced by charge size (larger charges lead to stronger attractions and larger lattice energies) and inter-ionic distance (smaller distances lead to stronger attractions).

Endothermic and Exothermic Reactions

  • Exothermic reactions: (\Delta H < 0) (heat released, surroundings feel hot).

  • Endothermic reactions: (\Delta H > 0) (heat absorbed, surroundings feel cold).

  • In reactions, chemical potential energy changes leading to observable thermal effects.

Practice Problems and Applications

  • Numerous calculations and formulas are presented:

    • Example problems illustrate the application of Hess’s Law and standard enthalpies of formation.

    • Methodologies show the relation between enthalpy changes in reactions, stoichiometry, and practical situations like combustion energy calculations.

Summary of Key Points

  • Enthalpy changes denote energy shifts during reactions, governed by bond energies and structural changes.

  • Hess’s Law relates summation of steps to overall reaction energy transfers.

  • Understanding lattice energies aids in predicting stability transitions in ionic compounds.