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Enthalpy (H)

Enthalpy (H). Amount of energy a sample contains at a certain pressure & temperature (heat content) Sometimes used in place of q. Enthalpy change ( Δ H). Enthalpy cannot be measured Change in Enthalpy is measured during a chemical reaction ΔH. Heat of Reaction. Δ H rxn

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Enthalpy (H)

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  1. Enthalpy (H) • Amount of energy a sample contains at a certain pressure & temperature • (heat content) • Sometimes used in place of q

  2. Enthalpy change (ΔH) • Enthalpy cannot be measured • Change in Enthalpy is measured during a chemical reaction • ΔH

  3. Heat of Reaction • ΔHrxn • Heat released or absorbed per one mole of reactant or product • Release of Heat  Exothermic  -ΔH • Absorption of Heat  Endothermic  +ΔH

  4. Finding ΔHrxn • ΔHrxn = H products _ Hreactants • Since H is not measured, calorimetry experiments are done

  5. Example • Example 2 H2 + O2 2 H2O ΔHrxn = -5741.6 kJ • Which contains more energy, the products or the reactants? • Is this reaction exo- or endo- thermic?

  6. Another Notation • Another Notation Sometimes the enthalpy is indicated as part of the chemical equation. • C + O2 CO2 -393.5 kJ • Energy is a product (exorxn) ΔH is negative. • N2 + O2 + 90 kJ  NO • Energy is a reactant (endorxn) ΔH is positive.

  7. Heat of Combustion • Amount of heat released in a combustion reaction. • Combustion reactions can be recognized by the O2 reactant and the CO2 and H2O products. • ΔHc is always negative.

  8. Entropy (S) • The measure of disorder or randomness of a system. • The more disorder there is the greater the entropy. • In general changes tend to occur so the highest entropy is reached.

  9. Spontaneity • Changes that happen on their own. • Depends on enthalpy, entropy, and temperature. • ΔG = ΔH – ΔST • When ΔG is negative, a reaction will be spontaneous.

  10. Spontaneity • An exothermic reaction is more likely to be spontaneous than an endothermic one. • A change that results in more entropy is more likely than one that results in less.

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