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Thermochemistry Review (Ch.1-7)

Overview of Supplemental Instruction

  • Madison (tutor) explains her role in Chem 1 support: Action Center resources, nightly tutoring, and a new program called supplemental instruction.
  • She will attend the lecture, take notes, and then host two weekly sessions with lesson plans and activities for deeper learning.
  • A survey will be sent next week to schedule session times.
  • Purpose: provide extra learning materials and structure to help students understand the material.

Key Concepts in Thermochemistry (review)

  • Energy and state function
    • Energy is a state function: it depends only on the initial and final state, not on the path taken.
    • Change in enthalpy for a reaction can be calculated from enthalpies of formation when you know the initial and final states.
  • Enthalpy change of reaction (ΔH_rxn)
    • Can be calculated from standard enthalpies of formation (
      \Delta H_f^\circ\nobreak{(substance)}) of the reactants and products.
    • Use formation data for each substance involved in the reaction to obtain ΔH_rxn.

Standard State and Enthalpy of Formation

  • Standard state is a defined set of conditions used for comparison:
    • Gas: pressure = 1 atm.
    • Solid and liquid: pressure = 1 atm as well.
    • Temperature: typically 25 °C unless stated otherwise.
    • Solutions: concentration = 1 M.
  • Examples and conventions discussed:
    • Oxygen (O₂) standard state: gas.
    • Carbon (as element) standard state: solid (graphite).
    • Mercury (Hg) and bromine (Br₂) are liquids at standard state.
    • For elements in their standard state, ΔH_f° = 0.
  • Enthalpy of formation (ΔH_f°):
    • Enthalpy change to form 1 mole of a substance from its constituent elements in their standard states.
    • The symbol ΔH_f° denotes standard enthalpy of formation.
    • Formation reactions must produce 1 mole of the substance on the product side.
    • For formation reactions, the left side shows reactants (elements in standard states).
    • Elements in their standard states have ΔH_f° = 0 (e.g., H₂(g), O₂(g), C(graphite)).
  • Formation reaction conventions and common mistakes (to watch for):
    • Ensure you form 1 mole of the target substance. If a coefficient would produce more than 1 mole, you may use fractional coefficients (e.g., 1/2, 3/2) to satisfy 1 mole of product while keeping elements in their standard states.
    • Don’t use the wrong phase for the standard state (e.g., iodine should be solid at standard state, not gaseous).
    • For elements like Cl₂, the standard state is Cl₂(g). For Fe, standard state is Fe(s).
    • For carbon, the standard state is graphite (not diamond).
  • Writing formation equations (examples):
    • NaCl formation: Na(s) + (1/2) Cl₂(g) → NaCl(s) with ΔH_f°(NaCl) corresponding to 1 mole of NaCl.
    • C₂H₆ formation: Form 1 mole of C₂H₆ from elements in standard states:
    • Left: C(s, graphite) + (3/2) H₂(g)
    • Right: C₂H₆(g)
    • FeCl₃ formation: Fe(s) + (3/2) Cl₂(g) → FeCl₃(s)
    • For each formation reaction, the enthalpy change is ΔH_f° of the product (times any stoichiometric coefficients if not forming 1 mole).
  • Balancing rules and practical notes:
    • If a reaction forms exactly 1 mole of product, ΔH = ΔH_f°(product) if reactants are in their standard states.
    • If you multiply the formation reaction by a factor n, ΔH is multiplied by n as well. If you flip the reaction (for decomposition), ΔH changes sign
    • In a formation equation, you start with elements in their standard states on the left and form the compound on the right.

Calculating ΔH_rxn from Formation Enthalpies

  • General formula:
    \Delta H{\text{rxn}}^\circ = \sum (\Delta Hf^\circ \text{ of products}) .\text{times their coefficients} \, - \, \sum (\Delta H_f^\circ \text{ of reactants}) .\text{times their coefficients}
  • Key points:
    • Use ΔH_f° data for each substance in the balanced equation.
    • For elements in standard states, ΔH_f° = 0.
    • Multiply ΔH_f° by the respective stoichiometric coefficients before summing.
    • This approach assumes formation data are available for all substances involved.
  • Worked example (thermite-like reaction): Aluminum and iron oxide to produce aluminum oxide and iron
    • Reaction: 2 Al + Fe₂O₃ → Al₂O₃ + 2 Fe (a famous exothermic reaction)
    • Use ΔHf° values: Fe₂O₃ and Al₂O₃ (the product ΔHf° values are used; elemental Al and Fe have ΔH_f° = 0).
    • Then ΔHrxn° = [1×ΔHf°(Al₂O₃) + 2×ΔHf°(Fe)] − [2×ΔHf°(Al) + 1×ΔH_f°(Fe₂O₃)]
    • Since ΔHf°(Al) = ΔHf°(Fe) = 0, the calculation reduces to ΔHrxn° = ΔHf°(Al₂O₃) − ΔH_f°(Fe₂O₃).
    • The sign and magnitude depend on the given ΔHf° values; the thermite reaction is highly exothermic (negative ΔHrxn°).
  • Additional notes on coefficients and data use:
    • If the balanced equation has a coefficient other than 1 for a product, multiply its ΔH_f° by that coefficient.
    • If the balanced equation has a coefficient other than 1 for a reactant, multiply its ΔH_f° by that coefficient.
    • If the reaction involves a standard-state element with 0 ΔH_f°, that term contributes nothing to the sum.
  • Example capacities and student practice:
    • Aluminum + iron oxide example demonstrates how to apply ΔH_f° data to obtain reaction enthalpy without needing all reactants' and products' de novo enthalpies.
    • In practice, you may be given ΔHf° for several substances and asked to compute ΔHrxn° using the above summation rule.

Exothermic vs Endothermic Reactions

  • Definitions:
    • Exothermic: energy (heat) flows from the system to the surroundings; ΔH_rxn° < 0; heat is released.
    • Endothermic: energy flows from the surroundings into the system; ΔH_rxn° > 0; heat is absorbed.
  • Conceptual view:
    • In a thermal chemical equation, heat (ΔH) is included as part of the equation to indicate energy change.
    • If heat is a product, it appears on the right; if heat is a reactant, it appears on the left.
  • Practical representation:
    • A standard chemical equation may omit ΔH; a thermal chemical equation includes ΔH at the end (e.g., + or − value).
  • Relationship to real-world measurements:
    • Exothermic reactions often feel hot; endothermic reactions feel cool to the touch as heat is absorbed.

Thermochemical Equations and Manipulations

  • Manipulation rules for balancing energy information with equations: 1) Multiply the entire equation by a factor n: ΔH scales by n
    • If you double the equation, ΔH doubles as well (e.g., Q becomes 2Q).
      2) Flip the equation: ΔH changes sign.
    • If a reaction releases energy (negative ΔH) and you flip it, the reverse reaction absorbs energy (positive ΔH).
    • When flipping, the amount of energy remains the same in magnitude; the sign changes.
      3) Phase considerations: changing phase may involve additional energy (e.g., fusion, vaporization).
    • If you alter the phase (gas↔liquid↔solid) independently of the chemical reaction, energy changes may include phase-change enthalpies separate from chemical ΔH.
  • Important cautions:
    • Do not arbitrarily change phase states when applying formation enthalpies unless the phase change is explicitly part of the reaction data.
    • When using fractional coefficients in formation data (e.g., 3/2 Cl₂), maintain consistency with the equation’s 1 mole product convention.

Energy Stoichiometry (Dimensional Analysis) and Problem-Solving Technique

  • Concept: relate energy to moles of a