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Understanding Chemical Reactivity and Enthalpy in Thermodynamics

Explore the relationship between chemical reactivity and enthalpy in thermodynamics, understanding energy transfer and predicting reactivity using standard enthalpy values and Hess's Law. Learn how to calculate heat of combustion.

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Understanding Chemical Reactivity and Enthalpy in Thermodynamics

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  1. STATEFUNCTIONSandENTHALPY

  2. heat energy transferred work done by the system energy change FIRST LAW OF THERMODYNAMICS ∆E = q + w Energy is conserved!

  3. Chemical Reactivity What drives chemical reactions? How do they occur? The first is answered by THERMODYNAMICS and the second by KINETICS. Have already seen a number of “driving forces” for reactions that are PRODUCT-FAVORED. • formation of a precipitate • gas formation • H2O formation (acid-base reaction) • electron transfer in a battery

  4. Chemical Reactivity Energy transfer allows us to predict reactivity. In general, reactions that transfer energy to their surroundings are product-favored. So, let us consider heat transfer in chemical processes.

  5. ENTHALPY: ∆H ∆H = Hfinal - Hinitial If Hfinal > Hinitial then ∆H is positive Process is ENDOTHERMIC If Hfinal < Hinitial then ∆H is negative Process is EXOTHERMIC

  6. ∆H along one path = ∆H along another path This equation is valid because ∆H is a STATE FUNCTION Depend ONLY on the state of the system NOT how it got there. Can NOT measure absolute H, ONLY measure ∆H.

  7. Standard Enthalpy Values Most ∆H values are labeled ∆Ho Measured under standard conditions P = 1 bar Concentration = 1 mol/L T = usually 25 oC with all species in standard states e.g., C = graphite and O2 = gas

  8. Standard Enthalpy Values ∆Hfo = standard molar enthalpy of formation — the enthalpy change when 1 mol of compound is formed from elements under standard conditions. See Table 6.2 and Appendix L

  9. ∆Hfo, standard molar enthalpy of formation H2(g) + 1/2 O2(g) --> H2O(g) ∆Hfo (H2O, g)= -241.8 kJ/mol By definition, ∆Hfo= 0 for elements in their standard states.

  10. HESS’S LAW In general, when ALL enthalpies of formation are known, Calculate ∆H of reaction? ∆Horxn =  ∆Hfo(products) - ∆Hfo(reactants) Remember that ∆ always = final – initial

  11. Using Standard Enthalpy Values Calculate the heat of combustion of methanol, i.e., ∆Horxn for CH3OH(g) + 3/2 O2(g) --> CO2(g) + 2 H2O(g) ∆Horxn =  ∆Hfo(prod) -  ∆Hfo(react)

  12. Using Standard Enthalpy Values CH3OH(g) + 3/2 O2(g) --> CO2(g) + 2 H2O(g) ∆Horxn =  ∆Hfo(prod) -  ∆Hfo(react) ∆Horxn = ∆Hfo(CO2) + 2 ∆Hfo(H2O) - {3/2 ∆Hfo(O2) + ∆Hfo(CH3OH)} = (-393.5 kJ) + 2 (-241.8 kJ) - {0 + (-201.5 kJ)} ∆Horxn = -675.6 kJ per mol of methanol

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