Enthalpy

1. Enthalpy • Most chemical and physical changes occur under essentially constant pressure (reactors open to the Earth’s atmosphere) • very small amounts of work are performed as system expands or contracts • the change in internal energy occurs primarily, or exclusively, as heat that is gained or lost.

2. Enthalpy • If a process occurs at constant pressure and the only work done is PV work, the heat flow is described by the enthalpy of the system. • Enthalpy (H): • a state function defined by the equation: H = E + PV (Question: Are P and V state functions?)

3. Enthalpy • Although the enthalpy of a system cannot be measured, the change in enthalpy (D H) can. • heat gained or lost by a system when a process occurs at constant pressure D H = Hfinal - Hinitial = qP where qP = heat gained/lost at constant pressure

4. Enthalpy • Two common types of work done by chemical systems: • electrical work • redox reactions incorporated into galvanic cells such as batteries • mechanical work (P-V work) • work done by expanding gases • pressure ~ constant • volume of system increases

5. Work done on surroundings ignition Enthalpy • P-V Work: • expanding gases in a cylinder of car engine • P-V work done on piston • eventually turns car wheels

6. Enthalpy • The amount of P-V work done at a constant pressure can be found: w = - P D V where D V = change in volume P = constant pressure • The negative sign indicates that work is being done by the system.

7. Enthalpy • For a process occurring at constant pressure in which the only work done is PV work, D E = qP + w D H = qP w = -P DV D E = D H - P DV (at constant pressure)

8. Enthalpy Example:If the volume of a cylinder increases from 2.00 L to 3.50 L at a constant pressure of 1.50 atm while it absorbs 2.515 kJ, what is the change in internal energy of the system?

9. Enthalpy

10. Enthalpies of Reaction • The change in enthalpy always compares the final and initial enthalpies of the system: D H = Hfinal - Hinitial • For a chemical reaction: • Hfinal = H products • Hinitial = H reactants • The enthalpy change for a chemical reaction is: D H = Hproducts - Hreactants

11. Enthalpies of Reaction • The enthalpy change that accompanies a chemical reaction is called the enthalpy of reaction • D Hrxn • Also called heat of reaction

12. Enthalpy • For a chemical reaction, D Hrxn = Hproducts - Hreactants • D H < 0 (negative) • heat islostby system • exothermic CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (l)

13. Enthalpy • For a chemical reaction, DHrxn = Hproducts - Hreactants • D H > 0 (positive) • heatgained/absorbed by the system • endothermic CO2 (g) + 2 H2O (l) CH4 (g) + 2 O2 (g)

14. Enthalpies of Reaction P • If D Hrxn = positive • endothermic • heat must be added • If D Hrxn = negative • exothermic • heat is given off H R time R H P time

15. Enthalpies of Reaction • DHrxn is associated with a specific chemical reaction. • extensive property • Depends on the amount of material • Thermochemical equationsare balanced chemical equations that show the associated enthalpy change • balanced equation • enthalpy change (DHrxn)

16. Enthalpies of Reaction • An example of a thermochemical equation: CH4 (g) + 2O2 (g) CO2 (g) + 2H2O (l) DH = -890. kJ • The coefficients in the balanced equation show the # moles of reactants and products that produced the associated DH. • If the number of moles of reactant used or product formed changes, then the DH will change as well.

17. Enthalpies of Reaction • For the following reaction: CH4 (g) + 2O2 (g) CO2 (g) + 2H2O (l) DH = -890. kJ -890. kJ -890. kJ 1 mol CH4 2 mol O2 -890. kJ -890. kJ 1 mol CO2 2 mol H2O

18. Enthalpies of Reaction • Guidelines for Using Thermochemical Equations: • Enthalpy is an extensive property • The magnitude of DH is directly proportional to the amount of reactant consumed or product produced

19. Enthalpies of Reaction • The thermochemical equation for burning 1 mole of CH4 (g): CH4 (g) + 2O2 (g) CO2 (g) + 2H2O (l) DH = -890.kJ • When 1 mole of CH4 is burned, -890. kJ of heat are released. • When 2 moles of CH4 are burned, -1780. kJ of heat are released.

20. Enthalpies of Reaction Example: How much heat is gained or lost when 10.0 g of butane, C4H10 (l) are burned at constant pressure? 2 C4H10 (l) + 13 O2 (g) 8 CO2 (g) + 10 H2O (g) DH = -5271 kJ

21. Enthalpies of Reaction

22. Enthalpies of Reaction Example: How much heat is gained or lost when 10.0 g of water vapor are formed at constant pressure in the following reaction? 2 C4H10 (l) + 13 O2 (g) 8 CO2 (g) + 10 H2O (g) DH = -5271 kJ

23. Enthalpies of Reaction

24. Enthalpies of Reaction • Guidelines for Using Thermochemical Equations (cont). • The enthalpy change for a reaction is equal in magnitude but opposite in sign to the DH for the reverse reaction. 2 H2O2 (l) 2 H2O (l) + O2(g) DH = -196 kJ 2 H2O (l) + O2(g) 2 H2O2 (l) DH = +196 kJ

25. Enthalpies of Reaction • Guidelines for Using Thermochemical Equations (cont): • The enthalpy change for a reaction depends on the physical state of the reactants and products. CH4 (g) + 2O2 (g) CO2 (g) + 2H2O (l) DH = -890. kJ CH4 (g) + 2O2 (g) CO2 (g) + 2H2O (g) DH = -802 kJ

26. Enthalpies of Reaction • Why does the DHrxn depend on the physical state of the reactants and products? • Energy is either absorbed or released when chemicals change from one physical state to another. H2O (l) H2O (g) DH = + 44 kJ

27. Calorimetry • The enthalpy change associated with a chemical reaction or process can be determined experimentally. • measure the heat gained or lost during a reaction (or process) at constant P • Measure the change in temperature

28. Calorimetry • Calorimetry: • the experimental measurement of heat gained or lost during a chemical reaction or process • Calorimeter • an instrument used to measure the heat gained or lost during a chemical reaction or process.

29. Calorimetry • Calorimetry is used to experimentally determine: • Heat capacity or specific heat • DHrxn • Enthaply change for a reaction • DHsoln • Enthalpy change when a substance is dissolved in a solvent (dissolution) • DHcombustion • Enthalpy change when a substance is burned in the presence of oxygen

30. Calorimetry If you leave your keys and your chemistry book sitting in the sun on a hot summer day, which one is hotter? Why is there a difference in temperature between the two objects?

31. Calorimetry • The temperature increase observed when an object absorbs a certain quantity of energy is determined by its heat capacity (C). • Amount of heat required to raise the temperature of an object 1oC (or 1 K) • As heat capacity increases, more heat must be added to produce a specific temperature increase.

32. Calorimetry • For pure substances, heat capacity is usually given for a specified amount of substance: • Molar heat capacity: • amount of heat required to raise the temperature of 1 mole of a substance by 1oC • Specific heat: • amount of heat required to raise the temperature of 1 g of a substance by 1oC

33. Calorimetry • Specific Heat = quantity of heat transferred mass x temp change = q mass x DT • Molar Heat = quantity of heat transferred Capacity moles x temp change = q mol x DT

34. Calorimetry Example:If 418 J is required to increase the temperature of 50.0 g of water from 24.0oC to 26.0oC, what is the specific heat of water?

35. Common units for specific heat are: • J cal • g.K g.oC Calorimetry

36. Calorimetry Example:What is the molar heat capacity of aluminum if it takes 9.00 J to raise the temperature of 5.00 g of aluminum from 298.0 K to 300.0 K?

37. Common units for molar heat capacity are: • J cal • mol.K mol.oC Calorimetry

38. Calorimetry Example:If the specific heat of Al (s) is 0.90 J/g.K, how much heat is required to raise the temperature of 10.0 kg of Al from 25.0oCto 30.0oC?

39. Calorimetry

40. Calorimetry Example:If the specific heat of Fe(s) is 0.45 J/g.K, what change in temperature would be observed when 1.0 kJ of heat is added to 45 g of Fe(s)?

41. Calorimetry