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Energetics

6. Energetics. 6.1 What is Energetics? 6.2 Enthalpy Changes Related to Breaking and Forming of Bonds 6.3 Standard Enthalpy Changes 6.4 Experimental Determination of Enthalpy Changes by Calorimetry 6.5 Hess’s Law 6.6 Calculations involving Standard Enthalpy Changes of Reactions.

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Energetics

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  1. 6 Energetics 6.1 What is Energetics? 6.2 Enthalpy Changes Related to Breaking and Forming of Bonds 6.3 Standard Enthalpy Changes 6.4 Experimental Determination of Enthalpy Changes by Calorimetry 6.5 Hess’s Law 6.6 Calculations involving Standard Enthalpy Changes of Reactions

  2. 6.1 What is energetics? (SB p.136) What is energetics? Energetics is the study of energy changes associated with chemical reactions. Thermochemistry is the study of heat changes associated with chemical reactions. Some terms Enthalpy(H) = heat content in a substance Enthalpy change(H)= heat content of products - heat content of reactants= Hp - Hr

  3. 6.1 What is energetics? (SB p.136) Law of conservation of energy The law of conservation of energy states that energy can neither be created nor destroyed.

  4. 6.1 What is energetics? (SB p.137) Internal energy and enthalpy e.g. Zn(s) + 2HCl(aq)  ZnCl2(aq) + H2(g)

  5. Heat change atconstant pressure Change in internal energy Work done on the surroundings = + 6.1 What is energetics? (SB p.138) Internal energy and enthalpy Enthalpy change (Heat change at constant volume)

  6. 6.1 What is energetics? (SB p.138) Exothermic and endothermic reactions An exothermic reaction is a reaction that releases heat energy to the surroundings. (H = -ve)

  7. Check Point 6-1 6.1 What is energetics? (SB p.139) Exothermic and endothermic reactions An endothermic reaction is a reaction that absorbs heat energy from the surroundings. (H = +ve)

  8. 6.2 Energy Changes Related to Breaking and Forming of Bonds

  9. 6.2 Enthalpy changes related to breaking and forming of bonds (SB p.140) Enthalpy changes related to breaking and forming of bonds e.g. CH4 + 2O2 CO2 + 2H2O

  10. 6.2 Enthalpy changes related to breaking and forming of bonds (SB p.140) Enthalpy changes related to breaking and forming of bonds In an exothermic reaction, the energy required in breaking the bonds in the reactants is less than the energy released in forming the bonds in the products (products contain stronger bonds).

  11. 6.2 Enthalpy changes related to breaking and forming of bonds (SB p.140) Enthalpy changes related to breaking and forming of bonds

  12. Check Point 6-2 6.2 Enthalpy changes related to breaking and forming of bonds (SB p.140) Enthalpy changes related to breaking and forming of bonds In an endothermic reaction, the energy required in breaking the bonds in the reactants is more than the energy released in forming the bonds in the products (reactants contain stronger bonds).

  13. 6.3 Standard Enthalpy Changes

  14. 6.3 Standard enthalpy changes (SB p.141) Standard enthalpy changes CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) H = -802 kJ mol-1 CH4(g) + 2O2(g)  CO2(g) + 2H2O(l) H = -890 kJ mol-1

  15. Enthalpy change under standard conditions denoted by symbol:H ø 6.3 Standard enthalpy changes (SB p.141) Standard enthalpy changes As enthalpy changes depend on temperature and pressure, it is necessary to definestandard states and conditions: 1. elements or compounds in their normal physical states;2. a pressure of 1 atm (101325 Nm-2); and3. a temperature of 25oC (298 K)

  16. Standard enthalpy change of neutralization(Hneut) is the enthalpy change when one mole of water is formed from the neutralization of an acid by an alkali under standard conditions. ø e.g. H+(aq) + OH-(aq)  H2O(l) Hneut = -57.3 kJ mol-1 ø 6.3 Standard enthalpy changes (SB p.142) Standard enthalpy changes of neutralization

  17. Acid Alkali Hneu ø HCl HCl HCl HF NaOH KOH NH3 NaOH -57.1 -57.2 -52.2 -68.6 6.3 Standard enthalpy changes (SB p.142) Standard enthalpy changes of neutralization H+(aq) + OH-(aq)  H2O(l)

  18. 6.3 Standard enthalpy changes (SB p.142) Standard enthalpy changes of neutralization Enthalpy level diagram for the neutralization of a strong acid and a strong alkali

  19. Standard enthalpy change of solution (Hsoln) is the enthalpy change when one mole of a solute is completely dissolved in a sufficiently large volume of solvent to form an infinitely dilute solution under standard conditions. ø 6.3 Standard enthalpy changes (SB p.142) Standard enthalpy change of solution

  20. e.g. NaCl(s) + water  Na+(aq)+Cl-(aq) Hsoln= +3.9 kJ mol-1 ø Enthalpy level diagram for the dissolution of NaCl 6.3 Standard enthalpy changes (SB p.143) Standard enthalpy change of solution

  21. e.g. LiCl(s) + water  Li+(aq) + Cl-(aq) Hsoln= -37.2 kJ mol-1 Enthalpy level diagram for the dissolution of LiCl in water ø 6.3 Standard enthalpy changes (SB p.143) Standard enthalpy change of solution

  22. Salt Hsoln(kJ mol-1) ø NaOH NaCl KOH KBr -44.7 +3.9 -57.8 +20.0 6.3 Standard enthalpy changes (SB p.143) Standard enthalpy change of solution

  23. Standard enthalpy change of formation (Hf ) is the enthalpy change of the reaction when one mole of the compound in its standard state is formed from its constituentelements under standard conditions. ø 6.3 Standard enthalpy changes (SB p.144) Standard enthalpy change of formation

  24. e.g. 2Na(s) + Cl2(g)  2NaCl(s) H = -822 kJ mol-1 ø Na(s) + ½Cl2(g)  NaCl(s) Hf = -411 kJ mol-1 ø Standard enthalpy change of formation of NaCl is -411 kJ mol-1. 6.3 Standard enthalpy changes (SB p.144) Standard enthalpy change of formation

  25. ø What is Hf [N2(g)] ? ø Hf [N2(g)] = 0 6.3 Standard enthalpy changes (SB p.144) Standard enthalpy change of formation N2(g)  N2(g) The enthalpy change of formation of an element is always zero.

  26. 6.3 Standard enthalpy changes (SB p.146) Standard enthalpy change of combustion e.g. C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(l) H1 = -2220 kJ 2C3H8(g) + 10O2(g) 6CO2(g) + 8H2O(l) H2 = ? H2 = -4440 kJ  It is more convenient to report enthalpy changes per mole of the main reactant reacted/product formed.

  27. Standard enthalpy change of combustion (Hc ) of a substance is the enthalpy change when one mole of the substance burns completely under standard conditions. ø e.g. C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(l) 1 mole Hc = -2220 kJ mol-1 ø 6.3 Standard enthalpy changes (SB p.146) Standard enthalpy change of combustion

  28. ø Substance Hc(kJ mol-1) H2(g) C (diamond) C (graphite) CO(g) CH4(g) -285.8 -395.4 -393.5 -283.0 -890.4 Check Point 6-3 6.3 Standard enthalpy changes (SB p.147) Standard enthalpy change of combustion

  29. 6.4 Experimental Determination of Enthalpy Changes by Calorimetry

  30. 6.4 Experimental determination of enthalpy changes by calorimetry (SB p.148) Experimental determination of enthalpy changes by calorimetry Calorimeter = a container used for measuring the temperature change of solution

  31. 6.4 Experimental determination of enthalpy changes by calorimetry (SB p.149) Determination of enthalpy change of neutralization

  32. Example 6-4A 6.4 Experimental determination of enthalpy changes by calorimetry (SB p.149) Heat evolved = (m1c1 + m2c2) ΔT where m1 is the mass of the solution, m2 is the mass of calorimeter, c1 is the specific heat capacity of the solution, c2 is the specific heat capacity of calorimeter, ΔT is the temperature change of the reaction

  33. 6.4 Experimental determination of enthalpy changes by calorimetry (SB p.150) Determination of enthalpy change of combustion The Philip Harris calorimeter used for determining the enthalpy change of combustion of a liquid fuel

  34. 6.4 Experimental determination of enthalpy changes by calorimetry (SB p.151) Determination of enthalpy change of combustion A simple apparatus used to determine the enthalpy change of combustion of ethanol

  35. Example 6-4B 6.4 Experimental determination of enthalpy changes by calorimetry (SB p.151) Heat evolved = (m1c1 + m2c2) ΔT Where m1 is the mass of water in the calorimeter, m2 is the mass of the calorimeter, c1 is the specific heat capacity of the water, c2 is the specific heat capacity of calorimeter, ΔT is the temperature change of the reaction

  36. Example 6-4C 6.4 Experimental determination of enthalpy changes by calorimetry (SB p.152) Determination of enthalpy change of solution • By measuring the temperature change when a known mass of solute is added to a known volume of solvent in a calorimeter • Heat change = (m1c1 + m2c2) T

  37. Check Point 6-4 6.4 Experimental determination of enthalpy changes by calorimetry (SB p.153) Determination of enthalpy change of formation • The enthalpy change of formation of a substance can be quite high • Found out by applying Hess’s law of constant heat summation

  38. 6.5 Hess’s Law

  39. Route 1 H1 H2 H3 E Route 2 6.5 Hess’s law (SB p.153) Hess’s Law A + B C + D H1 = H2+H3 Hess’s law of constant heat summation states that the total enthalpy change accompanying a chemical reaction is independent ofthe route by which the chemical reaction takes place.

  40. Enthalpy level diagram for the oxidation of C(graphite) to CO2(g) 6.5 Hess’s law (SB p.154) Enthalpy level diagram • Relate substances together in terms of enthalpy changes of reactions

  41. Enthalpy cycle for the oxidation of C(graphite) to CO2(g) 6.5 Hess’s law (SB p.155) Enthalpy cycle (Born-Haber cycle) • Relate the various equations involved in a reaction

  42. 6.5 Hess’s law (SB p.155) Importance of Hess’s law The enthalpy change of some chemical reactions cannot be determined directly because: • the reactions cannot be performed in the laboratory • the reaction rates are too slow • the reactions may involve the formation of side products But the enthalpy change of such reactions can be determined indirectly by applying Hess’s Law.

  43. Given: Hf [CO2(g)] = -393.5 kJ mol-1; Hc [CO(g)] = -283.0 kJ mol-1 ø ø Hf [CO(g)] ø + ½O2(g) + ½O2(g) C(graphite) + ½O2(g) CO(g) H1 H2 CO2(g) Hf [CO(g)] +H2=H1 ø Hf [CO(g)] =H1 -H2 ø 6.5 Hess’s law (SB p.153) Enthalpy change of formation of CO(g) = -393.5 -(-283.0 ) = -110.5 kJ mol-1

  44. Ca(s) + C(graphite) + O2 CaCO3(s) H2 H1 CaO(s) + CO2(g) ø Hf [CaCO3(s)] = H1 + H2 = -1028.5 kJ mol-1 + (-178.0) kJ mol-1 = -1206.5 kJ mol-1 ø Hf [CaCO3(s)] 6.5 Hess’s law (SB p.153) Enthalpy change of formation of CaCO3(s)

  45. ΔH ø MgSO4(s) + 7H2O(l) MgSO4·7H2O(s) aq ΔH1 ΔH2 Mg2+(aq) + SO42-(aq) + 7H2O(l) ø ΔH = enthalpy of hydration of MgSO4(s) ΔH1 = molar enthalpy change of solution of anhydrous magnesium sulphate(VI) ΔH2 = molar enthalpy change of solution of magnesium sulphate(VI)-7-water ΔH = ΔH1 - ΔH2 Check Point 6-5 ø 6.5 Hess’s law (SB p.153) Enthalpy change of hydration of MgSO4(s) aq

  46. 6.6 Calculations involving Standard Enthalpy Changes of Reactions

  47. ø Hreaction =  Hf [products] -  Hf [reactants] ø 6.6 Calculations involving standard enthalpy changes of reactions (SB p.159) Calculation of standard enthalpy change of reaction from standard enthalpy changes of formation

  48. Example 6-6A Example 6-6B Example 6-6C Example 6-6D 6.6 Calculations involving standard enthalpy changes of reactions (SB p.159)

  49. Hf =  Hc [products] -  Hc [reactants] ø ø ø 6.6 Calculations involving standard enthalpy changes of reactions (SB p.162) Calculation of standard enthalpy change of formation from standard enthalpy changes of combustion

  50. Example 6-6E Example 6-6F Check Point 6-6 6.6 Calculations involving standard enthalpy changes of reactions (SB p.162)

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