Entropy, Enthalpy, and Free Energy

1. Entropy, Enthalpy, and Free Energy

2. Heat of Reaction The heat of reaction is the “quantity of energy released or absorbed as heat during a chemical reaction.” To show the heat of reaction, you can add an energy value as a product or as a reactant to the equation. Ex: 2H2 (g) + O2 (g) 2H2O(g) + 483.6 kJ (remember that joules are a unit of heat energy)

3. Heat of Reaction If we produced half as much water as a product, we would release only half as much energy. If we produced three times as much water, we would release three times as much energy.

4. Exothermic reactions release energy A + B C + energy (warm) Endothermic reactions absorb energy A + energy B + C (cold)

5. Heat of Reaction If a reaction is exothermic, the heat of reaction is added as a product. If the reaction is endothermic, the heat of reaction is added as a reactant. Remember that bond forming is exothermic, and bond breaking is endothermic.

6. Heat of Reaction Any equation that includes the quantity of energy released as a reactant or product is a thermochemical equation. The states of matter (s, l, g ect) are always included on a thermochemical equation because they influence the total amount of energy gained or released. Remember it takes extra energy to change phases from solid to liquid to gas.

7. Enthalpy H is the symbol for a quantity called enthalpy. Enthalpy is the amount of heat energy that a substance has. It is only possible to talk about the heat energy of a system, the enthalpy, when looking at the change in heat between the reactants and the products. This is represented ∆H.

8. Enthalpy ∆H stands for change in enthalpy. The enthalpy change is the amount of energy absorbed or lost by a system as heat during a reaction at constant pressure. (∆H and ∆E mean the same thing, and can be used interchangeably on an energy diagram) Enthalpy change is calculated: ∆H = H products – H reactants

9. Change in enthalpy is written with a positive or negative sign if a thermochemical equation is not used (remember that thermochemical equations list energy as a product or reactant).

10. ∆H

11. In an endothermic reaction, where energy is absorbed, the energy of the reactants is higher than the energy of the products. This gives a positive value for ∆H.

13. In an exothermic reaction, where energy is released, the energy of the reactants is lower than the energy of the products. This gives a negative value for ∆H.

14. ∆H > 0 endothermic ∆H < 0 exothermic

15. ∆H and catalysts Remember that a catalyst is anything that lowers the activation energy. Change in enthalpy just looks at the difference between the energy of the final products and the final reactants.

16. ∆H and catalysts Because catalysts do not change the amount of energy of the reactants or the products, a catalyst will not change the ∆H of a reaction

17. Heat of Formation The molar heat of formation is the energy released or absorbed as heat when one mole of a compound is formed by combination of its elements. Molar heat of formation is abbreviated ∆Hf. This indicates the change in enthalpy required to form one mole of a compound. The unit for heat of formation is kJ/mol.

18. Heat of Formation Elements themselves (at their standard states) do not have a ∆Hf – there is not a chemical reaction to form an element. Remember that ∆H is the change in energy between the reactants and products.

19. Heat of Formation Heats of formation are given for the standard states of reactants and products – the states found at atmospheric pressure (1 atm) and room temperature (25º C).

20. Heat of Formation To show that a value is measured in a substance’s standard state, a 0 is added to the ∆H symbol, giving ∆H0 for the standard heat of a reaction. Adding a f indicates the standard heat of formation, so ∆H0f shows that the change in enthalpy value for the heat of formation in standard state.

21. Heat of Formation • ∆Hf0 of several compounds Lead (II) oxide -217.3 kJ/mol Write formula Endo or exo

22. Heat of Formation ∆H0f of dinitrogen monoxide + 90.29 kJ/mol Write formula Endo or exo?

23. The molar heat of formation is the energy released or absorbed as heat when one mole of a compound is formed by combination of its elements. Whenever you are working with a chemical equation that shows multiple moles of a compound, you multiply the heat of formation for the compound by the number of moles you have.

24. For example, ∆H0f of NO gas is 90.37 kJ. If you had an 3NO, the ∆H0f for those three moles of gas would be 271.11 kJ.

25. Example 1 What is the heat of reaction of the equation below? 2CO (g) + O2 (g) 2CO2 (g) ∆H0f CO (g) = -110.5 ∆H0f CO2 = -393.5

26. Example 1 Change in enthaply, ∆H, is ∆Hf products- ∆Hf reactants 2(-393.5) – 2(-110.5) = -566kJ

27. Hess’s Law Hess’s Law states that the overall enthalpy change in a reaction is equal to the sum of the enthalpy changes for the individual steps in the process. So the enthalpy change for a reaction equals the sum of the steps of the reaction mechanism.

28. Each step of the reaction mechanism has a ∆H, and to determine the overall enthalpy change, you need to have the steps of the reaction mechanism ordered so you get the correct final products and reactants. Remember that the key to a reaction mechanism is to cancel out the intermediates.

29. Hess’s Law Example 1 Calculate the ∆H for the conversion of graphite to diamond for the following reaction. Cgraphite (s) C diamond (s) Use the following equations: Cgraphite (s) + O2 (g) CO2 (g) ∆H = - 394 Cdiamond (s) + O2 (g) CO2 (g) ∆H = -396

30. Hess’s Law Example Cgraphite + O2 CO2∆H = -394 kJ CO2 C diamond + O2∆H = 396kJ C graphite C diamond ∆H =2 kJ

31. Enthalpy and Reaction Tendency Most chemical reactions are exothermic. As the reactions proceed, energy is released and the products have less energy than the reactants.

32. Enthalpy and Reaction Tendency In an exothermic reaction, the products are more resistant to change, more stable, than the original reactants. The tendency throughout nature is for a reaction to proceed in a direction that leads to a lower energy state.

33. Enthalpy and Reaction Tendency It would be logical to think that endothermic reactions, where you go from a lower energy state to a higher energy state, can not happen spontaneously – that you must add heat continuously to get the reaction to happen.

34. Enthalpy and Reaction Tendency Chemical reactions can happen spontaneously, without the addition of any input of energy (heating, catalyst, ect). A spontaneous chemical reaction happens without any input to the reaction system.

35. Enthalpy and Reaction Tendency However, some endothermic reactions do happen spontaneously – something besides enthalpy must determine if a reaction will be spontaneous. An additional factor that determines if a reaction will happen spontaneously is entropy.

36. Entropy Entropy is a measure of the disorder, or randomness, of a system. Entropy is abbreviated “S.” In general, the entropy of a gas is greater than a liquid, which is greater than a solid. Entropy gas>liquid>solid

37. Entropy The entropy of a system increases when a substance is divided into parts, or when a solute dissolves into solution. Entropy also increases as temperature increases. (The particles are moving around more, more disorder!)

39. Entropy ∆S is the change in entropy of a system. ∆S = S products – S reactants When ∆S >0, there is more disorder When ∆S < 0, there is less disorder, more order

40. Free Energy Processes in nature are driven in two directions – toward lowest enthalpy (energy) and toward highest entropy (disorder) To predict which factor will dominate a chemical process, an equation is available that relates enthalpy and entropy at a given temperature.

41. Free Energy Gibbs free energy, abbreviated G, is the combined value of enthalpy and entropy – it is the amount of energy available to do useful work. Again, ∆G, the change in free energy in a reaction is measured. The equation for ∆G is:

42. Free Energy ∆G = ∆H - T ∆S ∆H is change in enthalpy T is temperature in Kelvin ∆S is change in entropy

43. Free Energy When the value for ∆G > 0 you have a non-spontaneous reaction. So positive ∆G, and you will NOT have a spontaneous reaction. When the value for ∆G< 0, you have a spontaneous reaction. So negative ∆G values cause SPONTANEOUS reactions!