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Enthalpy

Enthalpy. Thermodynamics. Enthalpy (H). Measures 2 things in a chemical reaction: Energy change Amount of work Most chemical reactions happen at constant pressure (atmospheric pressure)—open container. Enthalpy (ΔH). 2 types of chemical reactions:

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Enthalpy

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  1. Enthalpy Thermodynamics

  2. Enthalpy (H) • Measures 2 things in a chemical reaction: • Energy change • Amount of work • Most chemical reactions happen at constant pressure (atmospheric pressure)—open container

  3. Enthalpy (ΔH) • 2 types of chemical reactions: • Exothermic—heat released to the surroundings, getting rid of heat, -ΔΗ • Endothermic—heat absorbed from surroundings, bringing heat in, +ΔΗ **Enthalpy of reaction— amount of heat from a chemical reaction which is given off or absorbed, units = kJ/mol

  4. Enthalpy of Reaction DH = Hfinal - HinitialHinitial = reactants Hfinal = products IfHfinal > Hinitial then DH is positive and the process is ENDOTHERMIC IfHfinal < Hinitial then DH is negative and the process is EXOTHERMIC

  5. Enthalpy of Reaction Hfinal < Hinitial and DH is negative

  6. Enthalpy of Reaction Hfinal > Hinitial and DH is positive

  7. Water Formation 2H2 + O2 2H20 ΔΗ = -967.28 kJ/mol

  8. More Enthalpy • The reverse of a chemical reaction will have an EQUAL but OPPOSITE enthalpy change • HgO Hg + ½ O2 ΔH = + 90.83 kJ • Hg + ½ O2HgO ΔH = - 90.83 kJ • SOOO-----total ΔH = 0

  9. Stoichiometry Returns

  10. Example 1: • Calculate the ΔH for the following reaction when 12.8 grams of hydrogen gas combine with excess chlorine gas to produce hydrochloric acid. • H2 + Cl2 2HCl ΔH = -184.6 kJ/mol

  11. Methods for determining ΔH • Calorimetry  • Stoichiometry  • Application of Hess’ Law • Enthalpies of Formation  

  12. Hess’ Law • Enthalpy change for a chemical reaction is the same whether it occurs in multiple steps or one step • ΔHrxn = ΣΔHA+B+C (sum of ΔH for each step) • Break a chemical reaction down into multiple steps • Add the enthalpies (ΔH) of the steps for the enthalpy for the overall chemical reaction

  13. Guidelines for using Hess’ Law • Use data and combine each step to give total reaction • Chemical compounds not in the final reaction should cancel • Reactions CAN be reversed but remember to reverse the SIGN on ΔH

  14. USING ENTHALPY Calculate DH for S(s) + 3/2O2(g)  SO3(g) knowing that S(s) + O2(g)  SO2(g) DH1 = -296.8 kJ SO2(g) + 1/2O2(g)  SO3(g) DH2 = -98.9 kJ The two equations add up to give the desired equation, so DHnet = DH1 + DH2 = -395.7 kJ

  15. Example 3: H2O(l) H2O (g) ΔH° = ? Based on the following: H2 + ½ O2 H2O(l) ΔH° = -285.83 kJ/mol H2 + ½ O2 H2O(g) ΔH° = -241.82 kJ/mol

  16. Example 4: NO(g) + ½ O2 NO2(g) ΔH° = ? Based on the following: ½ N2(g) + ½ O2 NO (g) ΔH° = + 90.29 kJ ½ N2(g) + O2 NO2(g) ΔH° = +33.2 kJ

  17. Methods for determining ΔH • Calorimetry • Application of Hess’ Law • Enthalpies of Formation • Stoichiometry

  18. Enthalpy of Formation (ΔHf°) • Enthalpy for the reaction forming 1 mole of a chemical compound from its elements in a thermodynamically stable state. • A chemical compound is formed from its basic elements present at a standard state (25°C, 1 atm) • Enthalpy change for this reaction = ΔHf° • ΔHf°= 0 for ALL elements in their standard/stable state.

  19. Enthalpy of Formation cont. • DHrxn = Hfinal – Hinitial Really, • ΔHf (products) - ΔHf (reactants) • Calculate ΔHrxn based on enthalpy of formation (ΔHf) • aA + bBcC + dD ΔH° =[c (ΔHf°)C + d(ΔHf°)D] - [a (ΔHf°)A + b (ΔHf°)B ]

  20. Calculate the ΔH° for the reaction 2Mg(s) + O2(g)  2MgO(s) DHof for MgO = -601.6 kJ/mol Recall that DHof for elements in their standard state = 0 kJ/mol DHrxn =∑(DHofproducts)(moles of products) – ∑(DHof reactants)(moles of reactants) = (-601.6kJ/mol)(2) – [(0kJ/mol)(2) + (1)(0kJ/ mol)] = - 1203.2kJ/mol

  21. Enthalpy of Formation Values for #6 and 7

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