Thermochemistry

1. Thermochemistry

2. Objective • Introduction to Thermochemistry • Define Temperature • Potential Energy • Enthalpy • Endothermic and Exothermic • Entropy?

3. Warm – up X2 • What is potential Energy? • Energy can not be created or destroyed!! • Where does the energy come from when you burn gasoline? • How about an Explosion??

4. The Nobel Prize (Alfred Nobel) • How Noble is the Nobel Prize?? • Found that when you mixed Nitroglycerin (shock sensitive) with diatomatious earth (inert) it was no longer nearly as shock sensitive. • His fortune was made in Dynamite!

5. What happens after the material explodes? • Is an explosion exothermic or endothermic? • Is that positive or negative? • Does the material contain more or less potential energy after it explodes? • Loosing potential energy in the bonds

6. What is an Thermochemistry (V)?? • The study of the transfer of of energy, as heat that is accompanied by chemical reactions and physical changes.

7. Thermochemical Equation (v) A Chemical equation that includes the enthalpy.2H2(g) + O2(g)  2H2O(g) + 483.6 kJ The 483.6 kJ is the amount of energy released (exothermic) 2H2O(g) + 483.6 kJ  2H2(g) + O2(g) The 483.6 kJ is the amount of energy absorbed (endothermic)

8. Thermochemical Equation The equations can also be written: 2H2(g) + O2(g)  2H2O(g); ΔH= -483.6 kJ 483.6 kJ is negative (or exothermic) What side? 2H2O(g)  2H2(g) + O2(g); ΔH= +483.6 kJ 483.6 kJ is positive (or endothermic) What side does it go on? Note the substances and moles formed

9. More Vocabulary for This Chapter • Calorimeter – a device that is intended to measure energy absorbed or released. • Temperature – A measure of the average kinetic energy of the particles in matter. (always Kelvin). • Joule – The SI unit for energy. (N-m)

10. More Vocabulary for This Chapter • Joule – The SI unit for energy. • So if I said Joule per gram: What does that mean? • If I said Joule per mole: What does that mean?

11. Enthalpy (ΔH) change • Enthalpy has the symbol (H) and is the amount of energy contained within a system (bonds, state, temperature). Potential energy within the bonds lost or gained. • Is it more or less stable after? • If I set off some Trinitrotoluene (TNT) is it more or less stable after I set it off?

12. Vocabulary for This Chapter Enthalpy change – Amount of energy, as heat absorbed or released by a system at constant pressure. ΔHrx = Hproducts – Hreactants ΔHmixing = Hsolution – (Hsolute + Hsolvent) Note Enthalpy is an measure of potential energy or energy contained within the bonds.

13. Enthalpy (ΔH) change • Exothermic – release heat in a reaction! • Endothermic absorb heat in a reaction! • Exothermic ΔH is negative. • Endothermic ΔH is positive.

14. Vocabulary for This Chapter • Enthalpy change – Amount of energy, as heat absorbed or released by a system at constant pressure. • Enthalpy of reaction – is the quantity of energy transferred as heat during a chemical reaction. • Endothermic – When heat is absorbed by a system from the surroundings (+). • Exothermic – When heat is released by the system to the surroundings (-)

15. The reason why exothermic is neg and endothermic is positive. • It is adding or subtracting potential energy from the energy stored in the bonds.

16. Compounds whose enthalpies of formation are highly negative a. Do not exist. b. Are somewhat stable. c. Are very unstable. d. Are very stable. Why??? ✔

17. Entropy S is a measure of disorder • What is more random, has a higher entropy ΔS ???

18. Higher Entropy ΔS???Type/more/less

19. Entropy Measure of the degree of randomness in a system? (V) S • What is more random, Solid, Liquid or Gas?

20. Entropy S • Entropy is a measure of the disorder in a system. If disorder increases the change in Entropy is positive. • Change in Entropy is ΔS! • If entropy increases disorder increases. CO2 (s)  CO2(g) (ΔS positive or negative?)

21. Entropy S • Entropy is a measure of the disorder in a system. If disorder increases the change in Entropy is positive. • Change in Entropy is ΔS! • If entropy increases disorder increases (In a Solution?) NaCl(s)  Na+(aq) + Cl-(aq) (ΔS positive or negative?)

22. Entropy S What is more random Reactant or Product? 2H2(g) + O2(g)  2H2O(l) S <0 Two molecules of gas converting into one liquid.

23. Entropy S What is more random Reactant or Product? 3H2(g) + N2(g)  2NH3(g) S <0 Two molecules of gas converting into one of gas.

24. Entropy S What is more random Reactant or Product? 2C6H6(l) + 15O2(g) 6H2O(g) + 12CO2(g) S > 0 Two molecules of gas converting into one of gas. More or less favored? More!!

25. Entropy S What is more random Reactant or Product? KNO3(s)  K+(aq) + NO3-(aq) S > 0 Two molecules of solid converting into two moles of aqueous ions. More or less favored?

26. Entropy is also additive ΔS = [Sum of Sproducts – Sum of Sreactants] Problem: Al2O3(s) + 3H2(g)2Al(s) + 3H2O(g) ΔS = ? If entropy of Al2O3(s) =51.00kJ/mol·K; Al(s) = 28.32 kJ/mol·K; H2O(g) is 188.7 kJ/mol·K; H2(g) is 130.6 kJ/mol·K .

27. Entropy S • Entropy is a measure of the disorder in a system. If disorder increases the change in Entropy is positive. • Change in Entropy is ΔS! • If entropy increases disorder increases. 2NO(g)  N2(g) + O2(g) (ΔS positive or negative?)

28. Day 2 (Warm-up) • Does Entropy increase or decrease? • Does Enthalpy increase or decrease? • Combustion of Liquid Octane (C8H18)? • Evaporation of Water?

29. The Most Expensive Combustion Reaction!! (The Russians are coming!) Cdiamond + O2 CO2 ΔH = -395.41 kJ/mol

30. Enthalpy (ΔH) change is the • a. pressure change of a system at constant temperature. • b. entropy change of a system at constant pressure. • c. temperature change of a system at constant pressure. • d. amount of energy absorbed or lost by a system as energy is the form of heat.

31. Objective • Gibbs Free Energy (spontaneous reactions) • Introduction to Specific Heat calculations • Energy Curve, Activation Energy • Explain Hess’s Law

32. Spontaneous (v) reaction • A spontaneous process is the time-evolution of a system in which it releases free energy (most often as heat) and moves to a lower, more thermodynamically stable, energy state. • Or the reaction happens…with no help • Free Energy???? • Thermodynamically stable??? • Ahhhhhhhh……………………..

33. G = H - TSIf Negative Spontaneous • At 300 K is this reaction spontaneous, not spontaneous, or can it not be determined? C2H4 (g) + H2(g)  C2H6(g) ΔH = -136.9kJ/mol; ΔS = -0.1207kJ/mol·K G = H -TS G =-136.9kJ/mol – (300K)(-0.121kJ/mol·K) -100.6 kJ/mol

34. Gibbs Free Energy (ΔG) • ΔG = Gibbs free Energy • If negative the reaction is spontaneous • If positive will not happen. ΔG = ΔH - TΔS

35. The Driving Force In the Nature of Reactions Two factors effect whether a reaction will occur spontaneously. • H – Enthalpy of formation if negative will favor Spontaneous Reaction -- WHY??

36. The Driving Force In the Nature of Reactions Factor Number TWO effect whether a reaction will occur spontaneously. G = H - TS • TS – Entropy Term (Entropy times temperature). If positive will favor Spontaneous Reaction -- WHY??

37. Spontaneous reactions are driven by a. decreasing enthalpy and decreasing entropy. b. decreasing enthalpy and increasing entropy. c. increasing enthalpy and decreasing entropy. d. increasing enthalpy and increasing entropy. 

38. Vocabulary for This Chapter • Enthalpy change – Amount of energy, as heat absorbed or released by a system at constant pressure. • Enthalpy of reaction – is the quantity of energy transferred as heat during a chemical reaction. ΔH = Hproducts – Hreactants Note Enthalpy is an measure of potential energy or energy contained within the bonds.

39. Thermochemical Equation • A Chemical equation that includes the enthalpy.2H2(g) + O2(g)  2H2O(g) + 483.6 kJ The 483.6 kJ is the amount of energy released (exothermic) 2H2O(g) + 483.6 kJ  2H2(g) + O2(g) The 483.6 kJ is the amount of energy absorbed (endothermic)

40. Thermochemical Equation Write the Thermal Chemical Equation For: 2H2(g) + O2(g)  2H2O(g); ΔH= -483.6 kJ - 483.6 kJ is negative (or __________) 2H2O(g)  2H2(g) + O2(g); ΔH= +483.6 kJ 483.6 kJ is positive (or ____________) Note the substances and moles formed

41. Warm up Question • Balance and write the thermochemical equation for the combustion of benzene (C6H6) ΔH = - 3476kJ/mol:

42. Objectives • What is Activation Energy? • Why does it exist? • Potential Energy Diagram? • What is it and how to read it.

43. Activation Energy • the minimum energy that must be put into a chemical system, containing potential for reactants to react! In order for the rock to get to the bottom of the hill it has to go over the hump

44. In Chemical reactions, What causes the hump? • Energy at which they collide (not all have the same energy). Why?? What surrounds an atom or molecule? • Orientation (for most reactions to occur they have to have the correct orientation (meet the right way)

45. Collision Theory • When particles of the reactant hit each other, only a certain percentage of the collisions cause a significant chemical change.

46. Potential Energy Diagram

47. Which is exothermic and which is endothermic?.Which curve shows more storage at the end?

48. Catalyst (v) • A substance that is not used up in a reaction that lowers the activation energy!!

49. Extra

50. Other Vocabulary for This Chapter • Specific Heat – amount of energy required to raise the temperature of 1 gram, K (1 °C ). • Heat – Energy transferred between matter due to their temperature differences. • Which way does heat flow.