Chapter 5

1. Chapter 5 Energy & Enthalpy Thermodynamics (rev. 0910)

2. Definition • Thermodynamics- is the study of energy transformations.

3. Chemical Reactions • Chemical reactions involve not just the conversion of reactants into products, but also involve an energy change in the form of heat—heat released as the result of a reaction, or heat absorbed as a reaction proceeds. • Energy changes accompany all chemical reactions and are due to rearranging of chemical bonding.

4. Making Bonds • Addition of energy is always a requirement for the breaking of bonds but the breaking of bonds in and of itself does not release energy. • Energy release occurs when new bonds are formed.

5. Bond Energy • If more energy is released when new bonds form than was required to break existing bonds, then the difference will result in an overall release of energy. • If, on the other hand, more energy is required to break existing bonds than is released when new bonds form, the difference will result overall in energy being absorbed.

6. Overall Reaction Energy • Whether or not an overall reaction releases or requires energy depends upon the final balance between the breaking and forming of chemical bonds.

7. Energy is... • the ability to do work or produce heat. • conserved. • made of heat and work. • a state function. (Energy is a property that is determined by specifying the condition or “state” (e.g., temperature, pressure, etc.) of a system or substance.) • independent of the path, or how you get from point A to B.

8. Thermodynamics State functions are properties that are determined by the state of the system, regardless of how that condition was achieved. energy , pressure, volume, temperature Potential energy of hiker 1 and hiker 2 is the same even though they took different paths. 6.7

9. Energy • While the total internal energy of a system (E) cannot be determined, changes in internal energy (E) can be determined. • The change in internal energy will be the amount of energy exchanged between a system and its surroundings during a physical or chemical change. ΔE = E final - E initial

10. Definitions • Work is a force acting over a distance. • Heat is energy transferred between objects because of temperature difference. (Heat is not a property of a system or substance and is not a state function. Heat is a process—the transfer of energy from a warm to a cold object.)

11. System vs Surroundings • The Universe is divided into two halves. • system and the surroundings. • In a chemistry setting, a system includes all substances undergoing a physical or chemical change. • The surroundings would include everything else that is not part of the system.

12. Heat • Most commonly, energy is exchanged between a system and its surroundings in the form of heat. • Heat will be transferred between objects at different temperatures. • Thermochemistry is the study of thermal energy changes.

13. Definition of Heat • Heat energy (or just heat) is a form of energy which transfers among particles in a substance (or system) by means of kinetic energy of those particle. In other words, under kinetic theory, the heat is transfered by particles bouncing into each other. • http://physics.about.com/od/glossary/g/heat.htm

14. Definition of Temperature • Temperature is a measurement of the average kinetic energy of the molecules in an object or system and can be measured with a thermometer or a calorimeter. It is a means of determining the internal energy contained within the system.

15. Temperature = Thermal Energy 900C 400C Energy Changes in Chemical Reactions Heat is the transfer of thermal energy between two bodies that are at different temperatures. Temperature is a measure of the thermal energy. greater thermal energy 6.2

16. Heat vs Temperature • Note that temperature is different from heat, though the two concepts are linked. Temperature is a measure of the internal energy of the system, while heat is a measure of how energy is transferred from one system (or body) to another. The greater the heat absorbed by a material, the more rapidly the atoms within the material begin to move, and thus the greater the rise in temperature. • http://physics.about.com/od/glossary/g/temperature.htm

17. Exo vs Endo • Exothermic reactions release energy to the surroundings. • Endothermic reactions absorb energy from the surroundings.

18. Heat Potential energy

19. Heat Potential energy

20. Three Parts Every energy measurement has three parts: • A unit ( Joules of calories). • A number. • a sign to tell direction. negative - exothermic positive- endothermic

21. Surroundings System Energy DE <0

22. Surroundings System Energy DE >0

23. Same rules for heat and work • Heat given off is “negative”. • Heat absorbed is “positive”. • Work done by the system on the surroundings is negative. • Work done on the system by the surroundings is positive.

24. First Law of Thermodynamics • The energy of the universe is constant. • It is also called the: • Law of conservation of energy. q = heat w = work • In a chemical system, the energy exchanged between a system and its surroundings can be accounted for by heat (q) and work (w). DE = q + w • Take the system’s point of view to decide signs.

25. Thermodynamics DE = q + w DE is the change in internal energy of a system q is the heat exchange between the system and the surroundings w is the work done on (or by) the system w = -PDVwhen a gas expands against a constant external pressure 6.7

26. Enthalpy and the First Law of Thermodynamics DE = q + w At constant pressure, q = DH and w = -PDV DE = DH - PDV DH = DE + PDV 6.7

27. Conservation of Energy • Energy exchanged between a system and its surroundings can be considered to off set one another. • The same amount of energy leaving a system will enter the surroundings (or vice versa), so the total amount of energy remains constant.

28. Metric Units • The SI (Metric System) unit for all forms of energy is the joule (J).

29. Heat and Work • DE = q + w • - q is exothermic -q = -∆H • +q is endothermic • -w is done “by” the system • +w is done “on” the system • Note: • ∆H stands for enthalpy which is the heat of reaction

30. Practice Problem • A gas absorbs 28.5 J of heat and then performs 15.2 J of work. The change in internal • energy of the gas is: (a) 13.3 J (b) - 13.3 J (c) 43.7 J (d) - 43.7 J (e) none of the above

31. Answer • (b) E = q + w • 28.5 J - 15.2 J = + 13.3 J

32. Practice Problem Which of the following statements correctly describes the signs of q and wfor the following exothermic process at 1 atmosphere pressure and 370 Kelvin? H2O(g) → H2O(l) (a) q and ware both negative (b) q is positive and wis negative (c) q is negative and wis positive (d) q and ware both positive (e) q and ware both zero

33. Answer • (c). An exothermic indicates q is negative and the gas is condensing to a liquid so it is exerting less pressure on its surroundings indicating wis positive.

34. What is work? • Work is a force acting over a distance. w= F x Dd P = F/ area d = V/area w= (P x area) x D (V/area)= PDV • Work can be calculated by multiplying pressure by the change in volume at constant pressure. • Use units of liter•atm or L•atm

35. Pressure and Volume Work • Work refers to a force that moves an object over a distance. • Only pressure/volume work (i.e., the expansion/contraction of a gas) is of significance in chemical systems and only when there is an increase or decrease in the amount of gas present.

36. Work needs a sign • If the volume of a gas increases, the system has done work on the surroundings. • work is negative • w = - PDV • Expanding work is negative. • Contracting, surroundings do work on the system W is positive. • 1 L•atm = 101.3 J

37. Example • When, in a chemical reaction, there are more moles of product gas compared to reactant gas, the system can be thought of as performing work on its surroundings (making w < 0) because it is “pushing back,” or moving back the atmosphere to make room for the expanding gas. When the reverse is true, w > 0.

38. Compressing and Expanding Gases • Compressing gas • Work on the system is positive • Work is going into the system • Expanding gas • Work on the surroundings is negative • Work is leaving the system

39. Clarification Info • If the reaction is performed in a rigid container, there may be a change in pressure, but if there is no change in volume, the atmosphere outside the container didn’t “move” and without movement, no work is done by or on the system. • If there is no change in volume (V= 0), then no work is done by or on the system (w= 0) and the change in internal energy will be entirely be due to the heat involved ( ΔE = q).

40. Examples • What amount of work is done when 15 L of gas is expanded to 25 L at 2.4 atm pressure? • If 2.36 J of heat are absorbed by the gas above. what is the change in energy? • How much heat would it take to change the gas without changing the internal energy of the gas?

41. Enthalpy • The symbol for Enthalpy is H • H = E + PV (that’s the definition) • DH = DE + PDV (at constant pressure)

42. Enthalpy (H) is used to quantify the heat flow into or out of a system in a process that occurs at constant pressure. DH = H (products) – H (reactants) DH = heat given off or absorbed during a reaction at constant pressure Hproducts < Hreactants Hproducts > Hreactants DH < 0 DH > 0 6.3

43. DH = DE + PDV Using DH = DE + PDV • the heat at constant pressure qp can be calculated from: • DE = qp + w (if w = - PDV then…) • DE = qp – PDV (now rearrange) • qp = DE + P DV = DH

44. Examples of Enthapy Changes • KOH(s) → K(aq) + OH-1 (aq) ΔHsolution= - 57.8 kJ mol1 • C3H8(g) + 5O2(g) → 3CO2(g) + 4H2O(l) ΔHcombustion = -2221kJ/mol • H2O(s) → H2O(l) ΔHfusion = 6.0 kJ/mol • Fe2O3(s) + 2Al(s) → Al2O3(s) + 2Fe(s) ΔHreaction - 852 kJ/mol • Ca(s) + O2(g) H2(g) → Ca(OH)2(s) ΔHformation - 986 kJ/mol1

45. 3 Methods • There are a variety of methods for calculating overall enthalpy changes that you should be familiar with. • The three most common are the: • the use of Heats of Formation • Hess’sLaw • the use of Bond Energies

46. Heat of Reaction • To compare heats of reaction for different reactions, it is necessary to know the temperatures at which heats of reaction are measured and the physical states of the reactants and products. • Look in the Appendix of the textbook to find Standard Enthapy tables.

47. Standard Enthalpy of Formation • Measurements have been made and tables constructed of Standard Enthalpies of Formation with reactants in their “standard states”. • Use the symbol DHºf • Standard state is the most stable physical state of reactants at: 1 atmosphere pressure specified temperature—usually 25 °C 1 M solutions • For solids which exist in more than one allotropic form, a specific allotrope must be specified.

48. Establish an arbitrary scale with the standard enthalpy of formation (DH0) as a reference point for all enthalpy expressions. f Standard enthalpy of formation (DH0) is the heat change that results when one mole of a compound is formed from its elements at a pressure of 1 atm. f DH0 (O2) = 0 DH0 (O3) = 142 kJ/mol DH0 (C, graphite) = 0 DH0 (C, diamond) = 1.90 kJ/mol f f f f Because there is no way to measure the absolute value of the enthalpy of a substance, must I measure the enthalpy change for every reaction of interest? The standard enthalpy of formation of any element in its most stable form is zero. 6.5

49. 6.5

50. ΔH°formation • It is important to recognize that the ΔH°formation(abbreviated as ΔH°f) is really just the heat of reaction for a chemical change involving the formation of a compound from its elements in their standard states.