Chapter 6 Thermochemistry

1. Chapter 6Thermochemistry

2. Energy • The ability to do work or transfer heat. • Work: Energy used to cause an object that has mass to move. • Heat: Energy used to cause the temperature of an object to rise.

3. Potential Energy Energy an object possesses by virtue of its position or chemical composition. PE = mgh

4. 1 KE =  mv2 2 Kinetic Energy Energy an object possesses by virtue of its motion.

5. kg m2 1 J = 1  s2 Units of Energy • The SI unit of energy is the joule (J). • An older, non-SI unit is still in widespread use: The calorie (cal). 1 cal = 4.184 J

6. System and Surroundings • The system includes the molecules we want to study (here, the hydrogen and oxygen molecules). • The surroundings are everything else (here, the cylinder and piston).

7. Work • Energy used to move an object over some distance. • w = F d, where w is work, F is the force, and d is the distance over which the force is exerted.

8. Heat • Energy can also be transferred as heat. • Heat flows from warmer objects to cooler objects.

9. Transferal of Energy • The potential energy of this ball of clay is increased when it is moved from the ground to the top of the wall.

10. Transferal of Energy • The potential energy of this ball of clay is increased when it is moved from the ground to the top of the wall. • As the ball falls, its potential energy is converted to kinetic energy.

11. Transferal of Energy • The potential energy of this ball of clay is increased when it is moved from the ground to the top of the wall. • As the ball falls, its potential energy is converted to kinetic energy. • When it hits the ground, its kinetic energy falls to zero (since it is no longer moving); some of the energy does work on the ball, the rest is dissipated as heat.

12. First Law of Thermodynamics • Energy is neither created nor destroyed. • In other words, the total energy of the universe is a constant; if the system loses energy, it must be gained by the surroundings, and vice versa. Use Fig. 5.5

13. Internal Energy The internal energy of a system is the sum of all kinetic and potential energies of all components of the system; we call it E. Use Fig. 5.5

14. Internal Energy By definition, the change in internal energy, E, is the final energy of the system minus the initial energy of the system: E = Efinal−Einitial Use Fig. 5.5

15. Changes in Internal Energy • If E > 0, Efinal > Einitial • Therefore, the system absorbed energy from the surroundings. • This energy change is called endergonic.

16. Changes in Internal Energy • If E < 0, Efinal < Einitial • Therefore, the system released energy to the surroundings. • This energy change is called exergonic.

17. Changes in Internal Energy • When energy is exchanged between the system and the surroundings, it is exchanged as either heat (q) or work (w). • That is, E = q + w.

18. E, q, w, and Their Signs

19. Exchange of Heat between System and Surroundings • When heat is absorbed by the system from the surroundings, the process is endothermic.

20. Exchange of Heat between System and Surroundings • When heat is absorbed by the system from the surroundings, the process is endothermic. • When heat is released by the system to the surroundings, the process is exothermic.

21. State Functions Usually we have no way of knowing the internal energy of a system; finding that value is simply too complex a problem.

22. State Functions • However, we do know that the internal energy of a system is independent of the path by which the system achieved that state. • In the system below, the water could have reached room temperature from either direction.

23. State Functions • Therefore, internal energy is a state function. • It depends only on the present state of the system, not on the path by which the system arrived at that state. • And so, E depends only on Einitial and Efinal.

24. State Functions • However, q and w are not state functions. • Whether the battery is shorted out or is discharged by running the fan, its E is the same. • But q and w are different in the two cases. • E = q + w

25. Work When a process occurs in an open container, commonly the only work done is a change in volume of a gas pushing on the surroundings (or being pushed on by the surroundings).

26. Work We can measure the work done by the gas if the reaction is done in a vessel that has been fitted with a piston. w = −PV

27. Internal Energy Example • Calculate E for a system undergoing an endothermic process in which 15.6 kJ of heat flows & where 1.4 kJ of work is done on the system.

28. PV Work Example • Calculate the work associated with the expansion of a gas from 46 L to 64 L at a constant external pressure of 15 atm.

29. Internal Energy, Heat & Work Example • A balloon is being inflated to its full extent by heating the air inside it. In the final stages of this process, the volume of the balloon changes from 4.00 x 106 L to 4.50 x 106 L by the addition of 1.30 x 108 J of energy as heat. Assuming that the balloon expands against a constant pressure of 1.0 atm, calculate E for the process. To convert between L * atm and J, use 1 L * atm = 101.3 J.

30. Enthalpy • If a process takes place at constant pressure (as the majority of processes we study do) and the only work done is this pressure-volume work, we can account for heat flow during the process by measuring the enthalpy of the system. • Enthalpy is the internal energy plus the product of pressure and volume: H = E + PV

31. Enthalpy • When the system changes at constant pressure, the change in enthalpy, H, is H = (E + PV) • This can be written H = E + PV

32. Enthalpy • Since E = q + w and w = −PV, we can substitute these into the enthalpy expression: H = E + PV H = (q+w) −w H = q • So, at constant pressure the change in enthalpy is the heat gained or lost.

33. Endothermicity and Exothermicity • A process is endothermic, then, when H is positive.

34. Endothermicity and Exothermicity • A process is endothermic when H is positive. • A process is exothermic when H is negative.

35. Enthalpies of Reaction The change in enthalpy, H, is the enthalpy of the products minus the enthalpy of the reactants: H = Hproducts−Hreactants

36. Enthalpies of Reaction This quantity, H, is called the enthalpy of reaction, or the heat of reaction.

37. The Truth about Enthalpy • Enthalpy is an extensive property. • H for a reaction in the forward direction is equal in size, but opposite in sign, to H for the reverse reaction. • H for a reaction depends on the state of the products and the state of the reactants.

38. Calorimetry Since we cannot know the exact enthalpy of the reactants and products, we measure H through calorimetry, the measurement of heat flow.

39. Heat Capacity and Specific Heat • The amount of energy required to raise the temperature of a substance by 1 K (1C) is its heat capacity. • We define specific heat capacity (or simply specific heat) as the amount of energy required to raise the temperature of 1 g of a substance by 1 K.

40. heat transferred Specific heat = mass  temperature change q s = m T Heat Capacity and Specific Heat Specific heat, then, is

41. Constant Pressure Calorimetry By carrying out a reaction in aqueous solution in a simple calorimeter such as this one, one can indirectly measure the heat change for the system by measuring the heat change for the water in the calorimeter.

42. Constant Pressure Calorimetry Because the specific heat for water is well known (4.184 J/mol-K), we can measure H for the reaction with this equation: q = m  s  T

43. Calorimetry Problem • When a 3.25g sample of solid sodium hydroxide was dissolved in a calorimeter in 100.0 g of water, the temperature rose from 23.9 C to 32.0 C. Calculate H (in kJ/mol NaOH) for the solution process: NaOH (s) ----> Na+(aq) + OH (aq)Assume it's a perfect calorimeter and that the specific heat of the solution is the same as that of pure water.

44. Enthalpy Example • When 1 mole of methane is burned at constant pressure, 890 kJ of energy is released as heat. Calculate H for a process in which a 5.8 g sample of methane is burned at constant pressure.

45. Constant Pressure Calorimetry When 1.00 L of 1.00 M Ba(NO3)2 solution at 25.00C is mixed with 1.00 L of 1.00 M Na2SO4 solution at 25.00C in a calorimeter, the white solid BaSO4 forms, and the temperature of the mixture increases to 28.10C. Assuming that the calorimeter absorbs only a negligible amount of heat, that the specific heat capacity of the solution is 4.18 J/0C*g, & the density of the final solution is 1.0 g/mL, calculate the enthalpy change per mole of BaSO4 formed.

46. Bomb Calorimetry Reactions can be carried out in a sealed “bomb,” such as this one, and measure the heat absorbed by the water.

47. Bomb Calorimetry • Because the volume in the bomb calorimeter is constant, what is measured is really the change in internal energy, E, not H. • For most reactions, the difference is very small.

48. Bomb Calorimetry

49. Constant Volume or Bomb Calorimetry It has been suggested that H gas obtained by the decomposition of water might be a substitute for natural gas (principally methane). To compare the energies of combustion of these fuels, the following experiment was carried out using a bomb calorimeter with a heat capacity of 11.3 kJ/0C. When a 1.50 g sample of methane gas was burned with excess oxygen in the calorimeter, the temperature increased by 7.30C. When a 1.15 g sample of hydrogen gas was burned with excess oxygen, the temperature increase was 14.30C. Compare the energies of combustion per gram for H & CH4

50. Calorimetry Continued • A 110. g sample of copper (specific heat capacity = 0.20 J/0C*g) is heated to 82.40C & then placed in a container of water at 22.30C. The final temperature of the water & the copper is 29.40C. What is the mass of the water assuming that all the heat lost by the copper is gained by the water?