BELL RINGER

1. BELL RINGER • What is energy to you? • List as many forms of energy and examples that you can think of: http://highschoolenergy.acs.org/content/hsef/en/what-is-energy/what-is-energy.html

3. Chemistry, The Central Science, 10th edition Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten Chapter 5Thermochemistry

4. Energy • The ability to do work or transfer heat. • Work: Energy used to cause an object that has mass to move. • Heat: Energy transferred from hotter to colder objects

5. 1 KE =  mv2 2 Kinetic Energy Energy of motion

6. Calculate kinetic energy • Determine the kinetic energy of a 625-kg roller coaster car that is moving with a speed of 18.3 m/s. •  KE = 0.5*m*v2 • KE = (0.5) * (625 kg) * (18.3 m/s)2 • KE = 1.05 x105 Joules

7. Calculate potential energy • A cart is loaded with a brick and pulled at constant speed along an inclined plane to the height of a seat-top. If the mass of the loaded cart is 3.0 kg and the height of the seat top is 0.45 meters, then what is the potential energy of the loaded cart at the height of the seat-top? • PE = m*g*h • PE = (3 kg ) * (9.8 m/s/s) * (0.45 m) • PE = 13.2 J

8. Potential Energy Energy an object possesses by virtue of its position or chemical composition. Ep= mgh g = 9.8 m/s2

9. Electrostatic Potential Energy • arises from the interaction between charged particles • Eel = (kQ1Q2) / d • k= 8.99 x 10^9 J-m/C2C=Coulomb, unit of electrical charge • Q1,Q2= electrical charges • Both + = charges repel each other Eel = + • Opposite charges = attract each other Eel = - *The lower the energy of a system= the more stable

10. kg m2 1 J = 1  s2 Units of Energy • The SI unit of energy is the joule (J). • The calorie (cal) is the amount of energy required to raise the temperature of 1 g of water by 1 C. 1 cal = 4.184 J • Calorie (Cal) is energy unit for nutrition. 1 Cal = 1000 cal = 1 kcal

12. Calculate the amount of calories and joules are in 1 serving of Doritos:

13. Calculate the amount of calories and joules are in 1 serving of Coke: • 140 Calories= 140,000 calories= 585,760 J

14. 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).

15. Open System- matter and energy can be exchanged with surroundings • Closed System- energy can be exchanged with surroundings but not matter • Isolated System- neither exchanged

16. 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. http://www.physicsclassroom.com/class/energy/Lesson-1/Definition-and-Mathematics-of-Work

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

18. 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.

19. 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.

20. 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.

21. 5.2 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

22. 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.

23. 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 In chemical reaction initial = reactants; final = products Use Fig. 5.5

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

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

26. 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.

27. E, q, w, and Their Signs

28. The internal energy for Mg(s) and Cl2(g) is greater than that of MgCl2(s). Sketch an energy diagram that represents the reaction MgCl2(s) Mg(s) + Cl2(g) Mg(s) + Cl2(g) MgCl2

29. Exchange of Heat between System and Surroundings • When heat is absorbed by the system from the surroundings, the process is endothermic. (feels cold)

30. Exchange of Heat between System and Surroundings • When heat is released by the system to the surroundings, the process is exothermic. (feels hot)

31. A system can lose energy in two ways:   the system can release energy (q) to the surroundings   the system can do work (w) on the surroundings  A system does 20 J of work on the surroundings. The change in energy of the system, E, is ______ J.  A system has a change in energy of -25 J. No work is done by or on the system, so _______ J of heat must have been _____________________ (absorbed / released).  A system has a change in energy of +45 J. If 30 J of work was done on the system, then _______ J of heat must have been _____________________ (absorbed / released).

32. Determine the change in energy, E, for each system: A system gives off 25.0 kJ of heat and has 15.0 kJ of work done on it. _____ • lose 25.0, gain 15.0 = -10.0 kJ A system takes in 75.0 kJ of heat and has 25.0 kJ of work done on it. _____ • gains 75.0, gain 25.0 = +100.0 kJ • A system does 45.0 kJ of work and loses 80.0 kJ of heat. _____ • lose 45.0, lose 80.0 = -125.0 kJ

33. EXO ENDO EXO ENDO EXO EXO EXO ENDO ENDO ENDO Classify each statement as talking about an [EXO]thermic or [ENDO]thermic reaction: • _____  surroundings get hot • _____  PE diagram is uphill • _____  energy is a product • _____  H is positive • _____  reactants have more energy • _____ H is negative • _____ PE diagram is downhill • _____ surroundings get cold • _____ products have more energy • _____ energy is a reactant

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

35. 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.

36. 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.

37. Many paths, same destination

38. 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.

39. 5.3 Enthalpy • Zn(s) + 2HCl(aq)  ZnCl2(aq) + H2(g) • video clip • Hydrogen gas produced expands against atmosphere which requires the system to do work

40. 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).

41. 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: pressure-volume work w = −PV P=pressure, V = Final – initial volume negative sign necessary to conform to sign conventions from table

42. Enthalpy Greek word enthalpein= to warm • 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

43. 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

44. 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.

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

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

47. Enthalpy is a state functionbc internal energy, pressure, and volume are all state functions.

48. Determine the sign of ΔH Endothermic= process in which heat is absorbed = (+) Exothermic= process in which heat is released = (-) • a. ice cube melts • + ice cube absorbs heat from surroundings= endothermic • b. 1 g butane is combusted into CO2 and H2O • - reaction gives off heat = exothermic

49. 5.4 Enthalpies of Reaction The change in enthalpy, H, is the enthalpy of the products minus the enthalpy of the reactants: H = Hproducts−Hreactants

50. Enthalpies of Reaction This quantity, H, is called the enthalpy of reaction, or the heat of reaction. 2H2(g) + O2(g)  2H2O(g) ΔH = -483.6 kJ Balloon filled with Hydrogen and oxygen gas