Enthalpy

1. Enthalpy Chemistry 11 - Energetics Source: Brown and LeMay’s Chemistry the Central Science

2. Objectives • Define the terms exothermicreaction,endothermic reaction, and enthalpy change of reaction (ΔH) • State that combustion is an exothermic reaction. • Apply the relationship between temperature change, enthalpy change, and the classification of a reaction as endothermic and exothermic. • Deduce, from an enthalpy level diagram, the sign of the enthalpy change of the reaction.

3. Energy • What is it? • How / why is it important? • Are there different kinds of energy? • How do we measure it?

4. Chemical reactions • Distinguish between a physical and chemical reaction.

5. Chemical reactions • A physical reaction involves a change of state or phase but the chemical composition of the substance stays the same. • A chemical reaction involves a change in the composition of the substance. • What are some chemical reactions / processes that you can think of? • Apply what we have learned in bonding to explain what a chemical reaction is.

6. Energy and chemical reactions • How do you think energy and chemical reactions might be related?

7. The first law of thermodynamics • There are many types or forms of energy. • Name some types of energy that we have encountered before. • Consider a rock sitting on top of a cliff that is made to fall to ground level. Describe the changes that have occurred in terms of energy.

8. The first law of thermodynamics • Energy can be converted from one form to the other. • It can neither be created nor destroyed, it can only transform from one type to another. • Is the amount of energy in the universe changing? decreasing? increasing?

9. The first law of thermodynamics • The first law of thermodynamics states that the amount of energy in the universe is constant. • So why are we always being told to conserve energy if it’s not going anywhere?

10. The first law of thermodynamics • In this unit, we will often refer to systems, of which there are many different types: • isolated • closed • open

11. The first law of thermodynamics A comparison of types of systems We will often work with isolated systems, because it is the ideal situation. Do we actually work with isolated systems in the “real world”?

12. The first law of thermodynamics • Define the internal energy of a system. • How would we represent a change in internal energy?

13. The first law of thermodynamics The change in internal energy of a system may be expressed as: ΔE = Efinal – Einitial What would a positive ΔE imply? A negative ΔE?

14. The first law of thermodynamics • Remember that the sign of a quantity only signifies direction. • - ΔE: Energy is lost . • + ΔE: Energy is gained.

15. The first law of thermodynamics Changes in internal energy In which process is energy lost? Energy gained? Which process has + ΔE? – ΔE? Internal energy, E final state initial state final state initial state

16. The first law of thermodynamics Energy diagram for the interconversion of H2(g), O2(g), and H2O(l) H2(g), O2(g) Internal energy, E • Which process has a +ΔE? • ΔE? H2O(l)

17. The first law of thermodynamics Energy diagram for the interconversion of H2(g), O2(g), and H2O(l) H2(g), O2(g) What does it mean if there are arrows going both ways? ΔE > 0 ΔE < 0 Internal energy, E H2O(l)

18. The first law of thermodynamics • Recall two terms that we have discussed before: • heat, Q • work, W • What is heat? • What is work? • How do they relate to energy?

19. The first law of thermodynamics • If heat is added to a system, what happens to internal energy? What if heat is taken away? • If work is done on a system, what happens to internal energy? What if work is done by the system?

20. The first law of thermodynamics The effect of heat and work on internal energy

21. The first law of thermodynamics The effect of heat and work on internal energy Derive an expression that relates the change in internal energy ΔE to heat Q and work W.

22. The first law of thermodynamics ΔE = Q + W This is the quantitative definition of the first law of thermodynamics.

23. The first law of thermodynamics Two gases, A(g) and B(g), are confined in a cylinder with a moving piston. Substances A and B react to form a solid product: A(g) + B(g) C(s) As the reaction occurs, the system loses 1150J of heat to the surroundings. The piston moves downward as the gases react to form a solid. This action requires 480 J of work to be done on the system. What is the change in internal energy of the system?

24. The first law of thermodynamics Review of chemical equations: A(g) + B(g) C(s)

25. The first law of thermodynamics Review of chemical equations: 2H2O(l)  2H2(g) + O2(g)

26. The first law of thermodynamics ΔE = Q + W ΔE = -1150 + 480 ΔE = -670 J

27. The first law of thermodynamics Calculate the change in internal energy of the system for a process in which the system absorbs 140 J of heat from the surroundings and does 85 J of work on the surroundings.

28. The first law of thermodynamics ΔE =+55 J

29. Exothermic vs endothermic • Chemical reactions may be classified as exothermic or endothermic. • One type means it releases heat and the other means it absorbs heat. Which is which?

30. Exothermic vs endothermic • When a process absorbs heat, it is said to be endothermic. • When a process releases heat, it is said to be exothermic. • endo- means “within” • exo- means “outside” • -therm means “heat” Can you think of some processes (physical or chemical) that are endothermic or exothermic?

31. State functions • Consider 50 g of water at 25C in a pot. • That system carries with it a certain amount of internal energy. • What are processes that may have occurred to reach this endpoint?

32. State functions • The pot of 50 g of water at 25C may have been obtained either by: • cooling 50 g of water from 100C to 25C, or • melting 50 g of ice and raising the temperature from 0C to 25C • Does the process by which that system reached the endpoint affect its internal energy?

33. State functions • No, it does not. • The process by which we arrived at that result doesn’t matter. The internal energy of 50 g of water at 25C is the same no matter how it was done. • This is an example of a state function, a quantity for which the magnitude only depends on the present state of the system, not the path the system took to reach that state.

34. State functions You are travelling from Manila to Baguio. Manila is at sea level, 0 m, and Baguio is much higher, 1500 m above sea level. Is the altitude change a state function? Is the distance traveled a state function?

35. State functions There are several routes or paths to get to Baguio from Manila: • via EDSA or C-5 and Katipunan • via NLEX or National Highway • via SCTEX or National Highway • via Kennon Road or Marcos Highway Does the path you take affect the altitude change from Manila to Baguio? Does it affect the distance traveled from Manila to Baguio?

36. State functions Quite obviously, the path taken does not affect the altitude change. Baguio is still 1500 m higher than Manila. Therefore, altitude change is a state function. Again, quite obviously, the path taken will affect the distance you travel. Therefore, distance traveled is not a state function. Can you think of any other state functions?

37. State functions Quite obviously, the path taken does not affect the altitude change. Baguio is still 1000 m higher than Manila. Therefore, altitude change is a state function. Again, quite obviously, the path taken will affect the distance you travel. Therefore, distance traveled is not a state function. Can you think of any other state functions?

38. Enthalpy Let’s go back to the First law of thermodynamics: ΔE = Q + W How would you measure heat? How would you measure work?

39. Enthalpy • Heat can be measured using calorimeters, which we discussed in physics. • What are the main parts of a calorimeter? • How is work defined in physics?

40. Enthalpy • Consider the following reaction: Zn(s) + 2H+(aq) Zn2+(aq) + H2(g) • Work is done by the system in this reaction. How will we use W = Fdto measure work?

41. Enthalpy Zn(s) + 2H+(aq) Zn2+(aq) + H2(g) • The work is done by the expanding hydrogen gas. • Work can be measured if that expanding gas can be harnessed to exert a force on something.

42. Enthalpy W = Fd for gases, force is related to pressure: P = F / A F = PA therefore: W = PAd W = P ΔV where P is the pressure in the system and ΔV is the change in volume.

43. Enthalpy W = P ΔV What does + ΔV imply? – ΔV?

44. Enthalpy Relationship between change in volume and work

45. Enthalpy Relationship between change in volume and work Therefore the more accurate relationship would be: W = - P ΔV

46. Enthalpy • A thermodynamic function called enthalpy, H, (enthalpeinmeans “to warm”) accounts for heat flow in processes at constant pressure, such that: H = E + PV where H is enthalpy, E is the internal energy of the system, P is the pressure of the system and V is the volume of the system. Why is it important to specify “at constant pressure”?

47. Enthalpy H = E + PV Is internal energy a state function? pressure? volume?

48. Enthalpy H = E + PV Internal energy E, pressure P, and volume V are all state functions, therefore enthalpy H is also a state function.

49. Enthalpy H = E + PV • However, in the context of chemical reactions, changes in these quantities are far more relevant. • Instead of H, we are interested in ΔH. • Derive an expression for change in enthalpy ΔH.

50. Enthalpy H = E + PV ΔH = Δ(E + PV) ΔH = ΔE + PΔV Which of these quantities can be measured?