Energy and Chemical Reactions

1. Energy and Chemical Reactions Goals: Assess heat transfer associated with changes in temperature and changes of state. Apply the first law of thermodynamics. Define and understand the state functions enthalpy and internal energy. Calculate the energy changes occurring in chemical reactions and learn how these changes are measured.

2. What is the relation of ENERGY and CHEMISTRY? • Energy accompanies both chemical and physical changes. * Evaporation of water: heat is absorbed by the water molecules and the system is cooled down. * Photosynthesis: sun’s energy is stored as chemical energy in carbohydrates and oxygen from carbon dioxide and water. This energy can then be released. 6 CO2 (g) + 6 H2O (g) + energy C6H12O6 (s) + 6 O2(g) C6H12O6 (s) + 6 O2 (g) 6 CO2 (g) + 6 H2O (g) + energy

3. What is Energy? • Energy is ______________________. • Energy is classified as kinetic or potential. Kinetic Energy – is associated with ______ ____________. Potential Energy – associated with ______ ____________. • Energy can be converted between potential and kinetic. Ex. A waterfall, a turbine

4. List forms of Kinetic Energy

5. List forms of Potential Energy:

6. Potential Energy • Positive and negative particles (ions) attract one another. • Two ions can bond. • As the particles attract they have a lower potential energy. NaCl — composed of Na+ and Cl- ions.

7. What is Internal Energy? • Internal energy (E or U) = ______ • Internal Energy of a chemical system depends on • number of particles • type of particles • temperature • The higher the T the higher the internal energy • So, use changes in T (∆T) to monitor changes in E (∆E).

8. Practice • When light shines on the solar panel, a small electric motor propels the car. What types of energy are involved in these setup?

9. What is the Law of Conservation of Energy? • Or The first Law of Thermodynamics states that __________________________________________________. Thus, The total energy of the universe is constant. THERMODYNAMICS is_________ ______________________.

10. Heat and Temperature • Heat is not the same as temperature. • The thermal energy (heat) of a given substance depends not only on temperature but also on the amount of substance. • Heat transfer occurs when two objects _____ ___________________________________ ________________. • Heat transfer always occurs from an object at _____ temperature to an object at a ________temperature = directionality. • Transfer of heat continues until both objects are at _____________________ (the system has reached ______________________).

11. Thermal Equilibrium • No further temperature changes occurs and the temperature throughout the entire systems is the same. • The quantity of heat lost by a hotter object and the quantity of heat gained by the cooler object when they are in contact are numerically equal (Law of Conservation of Energy). DT measures the energy transferred.

12. Directionality of Heat Transfer • Heat always transfer from _________ object to __________ one. Heat energy is associated with ______________.

13. Exothermic and Endothermic Processes • EXOTHERMIC is when ___ ______________________ ______________________. • ENDOTHERMIC is when ______________________________ _________________. T(system) goes down T(surr) goes up T(system) goes up T (surr) goes down

14. What is the relation of Energy and Chemical Reactions? All of thermodynamics depends on the law of CONSERVATION OF ENERGY: The total energy is unchanged in a chemical reaction. • If PE of products is less than reactants, the difference must be released as KE. PE of system dropped. KE increased. Therefore, you often feel a T increase.

15. James Joule 1818-1889 What are the Units of Energy? • 1 calorie = heat required to raise temp. of ___________________ __________________________________. • 1000 cal = 1 kilocalorie = 1 kcal • 1 kcal = 1 Calorie (a food “calorie”) • But we use the unit (SI) called the JOULE; 1 J = 1 kg m2/s2 • As a rough guide, 1 joule is the absolute minimum amount of energy required (on the surface of Earth) to lift a one kilogram object up by a height of 10 centimeters. • 1 cal = _______ joules

16. Which has the larger heat capacity? What is Heat Capacity? It is the heat required to ___ ______________________ ____________________.

17. What is Specific Heat Capacity? It is the heat required to ________ __________________________. How much energy is transferred due to T difference? The heat (q) “lost” or “gained” is related to a) sample mass b) change in T and c) specific heat capacity

18. Specific Heat Capacity Substance Spec. Heat (J/g•K) H2O 4.184 Ethylene glycol 2.39 Al 0.897 glass 0.84 Aluminum Practice: Change sp.heat of Al from J/gK to J/molK

19. Practice When a 4.5 g sample of a pure element absorbs 31.5 J of energy, its temperature increases from 20oC to 74.3oC. Identify the element. Specific Heat a. Silver 0.129 J/goC b. Mercury 0.139 J/goC c. Copper 0.385 J/goC d. Iron 0.444 J/goC

20. If 25.0 g of Al cool from 310 oC to 37 oC, how many joules of heat energy are lost by the Al? Notice that the negative sign on q signals heat “lost by” or transferred OUT of Al. Specific heat capacity = heat lost or gained by substance (J) (mass, g) (T change, K) heat gain/lose = q = (sp. ht.)(mass)(∆T)

21. Heat TransferNo Change of State q transferred = (sp. ht.)(mass)(∆T)

22. Heat Transfer • Use heat transfer as a way to find specific heat capacity, Cp • 55.0 g Fe at 99.8 ˚C Initial T = 372.8 K • Drop into 225 g water at 21.0 ˚C Initial T = 294 K • Water and metal come to 23.1 ˚C Final T = 296.1 K • What is the specific heat capacity of the metal?

23. Heat Transfer • Because of conservation of energy, q(Fe) = –q(H2O) (heat out of Fe = heat into H2O) or q(Fe) + q(H2O) = 0

24. Heat TransferChange of State Changes of state involve energy (at constant T) • Ice + 333 J/g (heat of fusion) -----> Liquid water q = mass * heat of fusion

25. Heat Transfer and Changes of State • Requires energy q = mass * heat of vaporization Liquid ---> Vapor + energy

26. Practice The heat of vaporization of benzene, C6H6, is 30.8 kJ/mol at its boiling point of 80.1 °C. How much heat is required to vaporize 128 g benzene at its boiling point? a.4.04 kJ b.18.8 kJ c.19.3 kJ d.50.5 kJ e.4.04 x 103 kJ

27. Heating/Cooling Curve of Water Note that T is constant as ice melts

28. +333 J/g +2260 J/g What quantity of heat is required to melt 500 g of ice and heat the water to steam at 100 oC? Heat of fusion of ice = 333 J/g Specific heat of water = 4.2 J/g•K Heat of vaporization = 2260 J/g

29. What quantity of heat is required to melt 500 g of ice and heat the water to steam at 100 oC? 1. To melt ice 2. To raise water from 0 oC to 100 oC 3. To evaporate water at 100 oC 4. Total heat energy

30. Heat Transfer in a Physical Process CO2 (s, -78 oC) ---> CO2 (g, -78 oC) Heat transfers from surroundings to system in endothermic process.

31. Heat Transfer in a Physical Process • CO2 (s, -78 oC) ---> CO2 (g, -78 oC) • A regular array of molecules in a solid -----> gas phase molecules. • Gas molecules have _______ (higher/lower) kinetic energy than in the solid phase.

32. ∆E = E(final) - E(initial) = E(gas) - E(solid) CO2 solid Energy Diagram: Heat Transfer CO2 gas Gas molecules have higher kinetic energy. Also, WORK is done by the system in pushing aside the atmosphere.

33. heat transfer in (endothermic), +q w transfer in (+w) w transfer out (-w) 1st Law of Thermodynamics: Energy is Conserved heat transfer out (exothermic), -q SYSTEM ∆E = q + w

34. What is Enthalpy? Enthalpy is ________________________. Most chemical reactions occur at constant Pressure (P), so Heat transferred at constant P = qp qp = ∆Hwhere H = enthalpy ∆E = q + w and so ∆E = ∆H + w (and w is usually small) ∆H = heat transferred at constant P ≈ ∆E ∆H = change in heat contentof the system ∆H = Hfinal - Hinitial

35. Enthalpy ∆H = Hfinal - Hinitial If Hfinal > Hinitial then ∆H is positive Process is _________________ If Hfinal < Hinitial then ∆H is negative Process is ________________ • Select between endothermic and exothermic.

36. Energy Transfer and 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

37. Energy Transfer and Chemical Reactivity But energy transfer also allows us to predict reactivity. In general, reactions that transfer energy to their surroundings are product-favored.

38. Using Enthalpy Consider the formation of water H2(g) + 1/2 O2(g) --> H2O(g) + 241.8 kJ Exothermic reaction — heat is a “product” and ∆H = – 241.8 kJ –> product favored

39. Using Enthalpy Making liquid H2O from H2 + O2 involves two exothermic steps. H2 + O2 gas H2O vapor Liquid H2O

40. Using Enthalpy Making H2O from H2 involves two steps. H2(g) + 1/2 O2(g) ---> H2O(g) + 242 kJ H2O(g) ---> H2O(liq) + 44 kJ ------------------------------------------------------------------ H2(g) + 1/2 O2(g) --> H2O(liq) + 286 kJ Example of HESS’S LAW— If a rxn. is the sum of 2 or more others, the net ∆H is ___________________ _______________________________.

41. What is the Hess’s Law? Hess’s Law: Forming H2O can occur in a single step or in a two steps. ∆Htotal is the same no matter which path is followed.

42. ∆H along one path = ∆H along another path ∆H is a State Function • This equation is valid because ∆H is a STATE FUNCTION • STATE FUNCTIONs depend only on ______ ___________________________________. • V, T, P, energy — • Unlike V, T, and P, one cannot measure absolute H. Can only measure ∆H.

43. What are Standard Enthalpy Values? Most ∆H values are labeled ∆Ho Measured under standard conditions: Pressure = ____________ Concentration = ________ Temperature = _________ with all species in standard states e.g., C = graphite and O2 = gas

44. Enthalpy Values Depend on how the reaction is written and on phases of reactants and products: H2(g) + 1/2 O2(g) --> H2O(g) ∆H˚ = -242 kJ 2 H2(g) + O2(g) --> 2 H2O(g) ∆H˚ = -484 kJ H2O(g) ---> H2(g) + 1/2 O2(g) ∆H˚ = ______ H2(g) + 1/2 O2(g) --> H2O(liquid) ∆H˚ = -286 kJ

45. Standard Enthalpy Values NIST (Nat’l Institute for Standards and Technology) gives values of ∆Hfo = standard molar enthalpy of formation It is the enthalpy change when _______ ____________is formed from ______ under _________________. See Table 6.2 and Appendix L

46. 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.

47. Which of the following chemical equations corresponds to the standard molar enthalpy of formation of N2O? NO(g) + 1/2 N2(g)  N2O(g) 2 N(g) + O(g)  N2O(g) N2(g) + 1/2 O2(g)  N2O(g) N2(g) + O(g)  N2O(g) 2 N2(g) + O2(g)  2 N2O(g)

48. Use ∆H˚’s to calculate enthalpy change for H2O(g) + C(graphite) --> H2(g) + CO(g) From reference books we find: H2(g) + 1/2 O2(g) --> H2O(g) ∆Hf˚ of H2O vapor = - 242 kJ/mol C(s) + 1/2 O2(g) --> CO(g) ∆Hf˚ of CO = - 111 kJ/mol

49. Use ∆H˚’s to calculate enthalpy change for H2O(g) + C(graphite) --> H2(g) + CO(g)

50. Using Standard Enthalpy Values In general, when ALLenthalpies of formation are known, ∆Horxn =  ∆Hfo(products) -  ∆Hfo(reactants) Remember that ∆ always = final – initial