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Thermochemistry

This review on enthalpy covers essential concepts in thermochemistry, including temperature as a measure of average kinetic energy, and the distinctions between open, closed, and isolated systems. It explains enthalpy (ΔH) as the heat of a reaction and emphasizes the difference between exothermic and endothermic changes. Additionally, it discusses the importance of potential energy changes during chemical reactions and phase changes, as well as how these changes can be represented through thermochemical equations and enthalpy diagrams.

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Thermochemistry

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  1. Thermochemistry Part II : Enthalpy

  2. Review: Temperature • A measure of the average kinetic energy of the particles that make up a substance or system • Change in temp represents a change in kinetic energy

  3. Systems and surroundings: • System: area where reaction takes place (generally is the reaction) • Surroundings: outside of the system • Closed verses open systems

  4. Open-system • both matter and energy can freely cross from the system to the surroundings and back. • Ex: an open test tube

  5. Closed-System • energy can cross the boundary, but matter cannot. • Ex: a sealed test tube

  6. Isolated-System • neither matter nor energy can cross between the system and the surroundings. • Ex: The universe • there are no surroundings to exchange matter or energy with (as far as we know!)

  7. Enthalpy (∆H) • a.k.a heat of reaction • Chemical rxn (and phase changes) involve changes in potential energy not kinetic • Enthalpy is the potential energy change of a system during process like chem rxn (or phase change) • Units: kJ/mol

  8. Chemical Rxns: • The internal energy of a reactant or product cannot be measured, but their change in enthalpy can • Enthalpy changes represent the diff between PE of products and PE of reactants • PE changes result from bonds being broken and/or formed

  9. Breaking a bond requires energy!

  10. Creating a bond releases energy ! (Strength of bond determines the energy needed to break or create bond)

  11. Exothermic and Endothermic Rxns:

  12. Exo verses Endo: • If energy is released out of the system then it is considered to be an exothermic enthalpy change (exit). • If energy is absorbed into the system then it is considered to be an endothermic enthalpy change (enter).

  13. There are several different enthalpies that we will be looking at : • Enthalpy of Rxn (∆Hrxn) • Standard enthalpy of Rxn (∆H0rxn) -- at standard temp (25°C) and pressure (100 kPa) • Standard molar enthalpy of formation (∆H0f) • Standard molar enthalpy of combustion (∆H0comb) • We use tables to find these VALUES!

  14. Standard enthalpy of Rxn (∆H0rxn) • Exothermic and endothermic • 3 ways to represent: • Thermochemical equation (includes energy in eq) • Chem eq + ∆H0rxn beside • Enthalpy diagrams

  15. Exothermic: • ∆H0is always negative!!!!! • Energy written on product side • H2 + ½ O2 H2O + 285.8 kJ • H2 + ½ O2 H2O ∆H0rxn= -285.8 kJ/mol

  16. Exothermic cont’d: Hint: • Arrow down for neg and up for pos • arrow goes toward the products

  17. Endothermic: • ∆H0is always positive!!!!! • Energy written on reactant side • 117.3 kJ + MgCO3 MgO + CO2 • MgCO3 MgO + CO2∆H0rxn= 117.3 kJ/mol

  18. Endothermic cont’d: Hint: • Arrow down for neg and up for pos • arrow goes toward the products

  19. Compare:

  20. Enthalpy dependent on molar amounts: • Linearly dependent (multiply by the stoich coefficients) • Ex:

  21. Standard molar enthalpy of formation (∆H0f) • Is the quantity of energy that is absorbed (+) or released (-) when one mole of a compound is formed directly from its elements (so this is why you see ½ as coefficient) • Ex:

  22. Standard molar enthalpy of combustion (∆H0comb) • Enthalpy associated with combustion of 1 mol of substance • Ex:

  23. HW  • Pg 643 # 15, 16, 17, 18

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