1. Energy Chem Honors Chapter 10

2. 7 Forms of Energy Sound- from vibration of sound waves Chemical- fuel, gas, wood, battery Radiant (light)- electromagnetic energy Electrical energy- electrons moving among atoms- as in the conductive wire of an electrical cord Atomic- nuclear (from nucleus of an atom) Mechanical- walk, run Thermal- heat

4. What is energy? Energy is the ability to do work or produce heat Two types: – Potential energy – Kinetic energy

5. Potential energy- energy due to position or composition Ex: – water behind a dam – Attractive and repulsive forces

6. Kinetic energy – is energy due to the motion of the object and depends on the mass of the object and depends on the mass (m) and velocity (v) KE = ½ mv 2

7. Law of Conservation of Energy States that energy can be converted from one form to another but can never be created or destroyed

8. Energy form going in- electrical Energy form going out- heat and light

9. Chemical energy stored in bonds of gasoline molecules converted to mechanical and heat energy through combustion

10. Electrical energy converted to mechanical energy

11. Hoover Dam- hydropower plant converts mechanical energy (flowing water) to electromagnetic, which is transported to homes, and then converted back into mechanical energy in a blender

12. Work- force acting over a distance

13. State Function Property of a system that does not depend on the pathway to its present state. – Ex: Displacement is a state function: I send 2 students to the cafeteria…the cafeteria is a specific distance from here….even though the 2 students take different routes to get to the cafeteria, they bith end up the same distance away from this room.

14. Energy is a state function : Work and heat are not state functions:

15. Temperature and Heat What is the difference between warm water and cold water?

16. Thermal energy: random motions of components of an object Temperature: measures the random motions of the components of a substance (thermal energy) Heat: flow of energy due to a temperature difference

17. System and Surroundings System- part of the universe that we are focusing our attention on Surroundings- everything else in the universe

18. Exothermic process – a process that results in the evolution of heat- energy flows out of the system Endothermic process- a process that absorbs energy from the surroundings- energy flows into the system

19. Exothermic or endothermic? 1. Your hand gets cold when you touch ice 2. ice melts when you touch it 3. Ice cream melts 4. Propane is burning in a propane torch. 5. Water drops on your skin evaporate after swimming 6. Two chemicals mixing in a beaker give off heat

20. Exothermic Endothermic Exothermic Endthermic exothermic

21. Energy from an exothermic reaction comes from the difference in potential energy between the products and reactants Which has lower energy, the reactants or products?

22. Burning match - Total energy is conserved -exothermic so energy flows from system to surroundings - So the energy gained by the surroundings must be equal to the energy lost by the system

23. Burning match cont. The potential energy lost from the burned match came from the stored energy in the bonds of the reactants In any exothermic reaction some of the potential energy stored in the chemical bonds is converted to thermal energy- which is random kinetic energy- as heat

24. Thermodynamics Thermodynamics- study of energy Law of Conservation of energy is also known as THE FIRST LAW OF THERMODYNAMICS!! – The energy of the universe is constant.

25. 1 st Law- A Quantitative application We can use the first law to analyze energy changes in chemical systems Internal energy (E) of a system is the sum of all of the PE (potential energy) and KE (kinetic energy) of a system E = PE + KE

26. ΔE < 0 ΔE > 0 Internal Energy, E H 2 (g), O 2 (g) H 2 O (l)

27. The internal energy of a system can be changed by a flow of work or heat or both so mathematically we have ΔE = q + w Δ = change in q = heat added or liberated by the system w = work done on the system

28. Thermodynamic quantities have – Number – indicates the magnitude of the change – Sign (+ or -) – indicates the direction of the flow from the SYSTEM’s point of view

29. q and w When heat is added to a system, or work is done on a system, the internal energy increases so – When heat is transferred from the surroundings to the system q is positive – When work is done on a system by the surroundings w is positive

30. q and w Work and heat created by the system transferred to the surroundings lowers the internal energy of the system – When work is done by the system on the surroundings, w is negative – When heat flows from the system to the surroundings, q is negative

31. In an exothermic process, what is the sign of q? In an endothermic process, what is the sign of q?

32. Measuring Energy Changes Common units of energy change – Calorie – amount of energy (heat) required to raise the temperature of one gram of water one degree Celsius – Joule – in terms of a calorie 1 calorie = 4.184 J

33. Measuring Energy Changes cont. Express 60.1 cal of energy as Joules 60.1 cal x 4.184 J = 251 J 1 cal

34. Measuring Energy Changes Cont. The energy (heat) required to change the temperature of a substance depends on – The amount of substance being heated (grams) – The temperature change – Identity of the substance

35. Measuring Energy changes cont Identity of the substance- definition of a calorie is based on water, what about other substances?

36. Measuring Energy Changes cont. Specific heat – the amount of energy required to change the temperature of one gram of any substance by one degree Celsius – Water 4.184 J/g °C – Iron (s) 0.45 J/g °C – Silver (s) 0.24 J/g °C

37. Measuring Energy Changes cont. q = mcΔT q = energy in the form of heat m = mass c = specific heat ΔT = change in temperature in °C

38. Measuring Energy Changes cont. Determine the amount of energy (heat) in Joules required to raise the temperature of 7.40 g of water from 29°C to 46°C

39. Measuring Energy Changes cont. Ex: What quantity of energy (in Joules) is required to heat a piece of iron weighing 1.3 g from 25°C to 46°C ?

40. Calorimetry A calorimeter is used to determine the heat associated with a chemical reaction

41. Calorimeter Constant All parts of a calorimeter heat up or cool down as heat is released or absorbed in the chemical process Heat capacity of the calorimeter, C, also known as the calorimeter constant – Defined as the sum of the products of the specific heat and the mass of all components of the calorimeter

42. Calorimeter Constant cont. Formula for the heat energy produced in the calorimeter is q = C calorimeter ΔT

43. Calorimeter Constant Ex A calorimeter has a heat capacity of 1265 J/°C A reaction causes the temperature of the calorimeter to change from 22.34 C to 25.12 C. How many joules of heat were released in this process?

44. 3517 J

45. Calorimetry Can determine the temperature change of the reaction, and use the heat capacity of the calorimeter to determine the ΔH of the reaction What is H?

46. Enthalpy (H) For most chemical reactions chemists are interested in the heat generated at constant pressure, which is denoted as q p = H H is the heat content of a compound ΔH= which is the difference in heat between the reactants and products so ΔH= H products - H reactants

47. Thermochemistry Because chemical reactions occur at constant pressure, ΔH = q p for chemical reaction ΔH is a state function- independent of the path

48. Energy Diagram

49. Energy Diagram with Catalyst

50. Energy Diagram

51. Hess’s Law The enthalpy of a reaction is the sum of the enthalpies of the combined reactions- known as Hess’s Law Hess’s Law states that, whatever mathematical operations are performed on a chemical reaction, the same mathematical operations are applied also to the heat of reaction :

52. Hess’s Law Enthalpies of reaction have been measured and tabulated- which allows us to calculate ΔH of a reaction without making calorimetric measurements for all reactions Because enthalpy is a state function, the change in enthalpy is NOT dependent on the path…..doesn’t matter if the reaction takes place in one step, or multiple steps, change in enthalpy is the same

53. Hess’s Law continued Standard heat of reaction, ΔH° rxn ( ° indicates standard conditions of 25°C and 1 atm) is the heat produced when the specified number of moles in the balanced equation reacts Ex: C 3 H 8(g) + 5O 2(g)  3CO 2(g) + 4H 2 O (g) ΔH° of rxn = -2044 kJ when 1 mol of propane reacts with 5 mol of oxygen

54. Hess’s Law cont. Note that the physical states of the reactants and products are indicated (s,l,g, etc) H 2(g) + ½ O 2(g)  H 2 O (g) ΔH = - 241.8 kJ H 2(g) + ½ O 2(g)  H 2 O (l) ΔH = -285.8 kJ

55. *Hess’s Law* 1. If the coefficients of a chemical reaction are all multiplied by a constant, ΔH° rxn is multiplied by the same constant- once this is done, the reaction is no longer standard, drop the ° Recall combustion of propane….

56. C 3 H 8(g) + 5O 2(g)  3CO 2(g) + 4H 2 O (g) ΔH° rxn = -2044 kJ when 1 mol of propane reacts with 5 mol of oxygen If we multiply the reaction by 2, 2C 3 H 8(g) + 10 O 2(g)  6CO 2(g) + 8H 2 O (g) ΔH rxn = -4088 kJ If we multiply the reaction by ½, ½ C 3 H 8(g) + 5/2 O 2(g)  3/2CO 2(g) + 2H 2 O (g) ΔH rxn = -4088 kJ

57. 2. If two or more reactions are added together to obtain an overall reaction, the heats of these reactions are also added to give the heat of the overall reaction

58. Hess’s Law Ex: N 2(g) + O 2(g) 2NO 2(g) ΔH = 68 kJ We could also break down the rxn into 2 steps….

59. N 2(g) + O 2(g)  2NO (g) ΔH 2 = 180 kJ 2NO (g) + O 2(g)  2NO 2(g) ΔH 3 = -112 kJ ΔH 2 + ΔH 3 = 68 kJ

60. *Hess’s Law* 3. ΔH° forward rxn = - ΔH° reverse rxn Ex: C 3 H 8(g) + 5O 2(g)  3CO 2(g) + 4H 2 O (g) ΔH° rxn = -2044 kJ when 1 mol of propane reacts with 5 mol of oxygen In this ex, burning propane produced a large amount of heat. It would be almost impossible to experimentally calculate the reverse reaction. But the Hess’s Law states hat if we reverse the reaction, we can reverse the sign of ΔH So the energy needed to produce the reverse reaction would be + 2044 kJ

61. Hess’s Law Example Ex: pg 286 Using the enthalpies of combustion for graphite and diamond, calculate the ΔH° rxn for the conversion of graphite to diamond C graphite(s)  C diamond(s)

62. Hess’s Law example cont. The combustion reactions are: C graphite(s) + O 2(g)  CO 2(g) ΔH° = -394 kJ C diamond(s) + O 2(g)  CO 2(g) ΔH° = -396 kJ

63. Hess’s Law Example continued If we reverse the second reaction and add them together the compounds that are on both sides of the reaction cancel out C graphite(s) + O 2(g)  CO 2(g) ΔH = -394kJ + CO 2(g)  C diamond(s) + O 2(g) ΔH = -(-396kJ) C graphite(s)  C diamond(s) ΔH rxn = +2kJ

64. Hess’s Law Another Example Given the following data: 4CuO (s)  2Cu 2 O (s) + O 2(g) ΔH° = 288 kJ Cu 2 O (s)  Cu (s) + CuO (s) ΔH° = 11 kJ Calculate the ΔH rxn ° for the following reaction 2Cu (s) + O 2(g)   2CuO (s)

65. Specific ΔH ΔH° f = heat energy absorbed or released during synthesis of one mole of a compound from its elements at standard conditions ΔH° sol = heat energy absorbed or released when a substance dissolves in a solvent ΔH° comb = heat energy released when a substance reacts with oxygen to form CO 2 and H 2 O

66. Second Law of Thermodynamics States that the entropy of the universe is always increasing – Entropy (S) is a measure of disorder or randomness – Spontaneous process- occurs in nature without outside intervention