Energy

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 mv2

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- electricalEnergy form going out- heat and light

9. Chemical energy stored in bonds of gasoline moleculesconverted 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 • 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. -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 Burning match

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. 1st 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. Internal Energy, E H2(g) , O2(g) H2O (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 x4.184 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 = Ccalorimeter 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 qp= H • H is the heat content of a compound • ΔH= which is the difference in heat between the reactants and products • so ΔH= Hproducts - Hreactants

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

48. Energy Diagram

49. Energy Diagram with Catalyst

50. Energy Diagram