Energy

1. Energy 11/12 & 11/13

2. Warm Up • Refer to the diagram below: • What is going to happen in this situation?

3. Forms of Energy • Energy- ability to do work and produce heat • W=FX • Work= Force x Distance

4. Measurement of Energy • Measured by amount of work it can do (physics) - OR- amount of heat it can be changed into (chemistry) • SI unit of measure = Joule (J)

5. Forms of Energy A. Mechanical Energy • Energy that can exert a force or produce motion • Potential: energy of position, stored energy • chemical potential energy – energy stored in a substance b/c of its composition • Kinetic: energy of motion B. Chemical Energy • Energy associated with chemical change C. Electrical Energy • Current carries energy to do work D. Electromagnetic radiation E. Magnetic F. Nuclear G. Heat- energy in the process of flowing from warmer to cooler object H. Sound

6. Conservation of Energy • Energy is changed, or converted, from one form to another • Law of Conservation of Energy • Energy cannot be created or destroyed • AKA: 1st Law of Thermodynamics • The total amount of energy remains the same during all energy changes.

7. Temperature of all things flow towards equilibrium • It equals out. • 2nd Law of thermodynamics • Absolute Zero=No movement=No Energy • 3rd Law of thermodynamics

8. Describe the energy conversion in this picture

9. Thermochemistry • Thermochemistry is the study of heat changes that accompany chemical reactions and phase changes. • In thermochemistry, the system is the specific part of the universe that contains the reaction or process you wish to study.

10. Everything in the universe other than the system is considered the surroundings. • Therefore, the universe is defined as the system plus the surroundings. universe = system + surroundings

11. Temperature vs. Heat • Heat • Energy that transfers from one object to another because of a temperature difference between them. • Temperature • A measure of the average kinetic energy of the particles in a sample of matter

12. Heat and temperature • Heat is energy. It can do work. • Temperature is a man-made, arbitrary scale indicating which direction heat is flowing…is heat going into the system, temperature rising or is heat leaving the system, temperature declining. • Heat is measured with an instrument called a calorimeter. • Heat is NOT measured with a thermometer. • Temperature is measured with a thermometer. • Heat is measured in Joules. • Temperature is measured in degrees.

13. How to measure heat and energy… • calorie • Quantity of heat that will increase the temperature of 1 gram of water by 1oC • 1 calorie (cal) = 4.19 Joules (J) • 1 Calorie (food) = 1000 cal (heat) • How would you calculate food Calories to Energy? • How many joules of energy will a bowl of cereal containing 230 Cal produce?

14. What are the Units of Energy? • It all depends… • In the food we eat, the units are Calories • In science, we use the calorie or the Joule • To convert between calories and joules, use the following conversion factor • 1 calorie = 4.19 Joules (J)

15. Example Problems • Covert 465 Joules to calories. • Convert 110 calories to Joules.

16. So when we talk about Calories (or calories) we are talking about energy today and the amount of energy we are taking in to our body. • We must USE that energy we have taken in OR our bodies will convert it to the storage form of energy…FAT. • Fat is simply the body’s way of saying…“don’t want to use that energy now? OK. I will save it for you for later”.

17. Specific Heat • When heated, different substances change temperature at different rates. • Specific Heat • The amount of heat it takes to raise the temperature of 1 g of substance by 1oC. • Specific heat of water = 1 cal/g.oC • Specific heat of Iron = 0.11 cal/g.oC • What are the implications for this?

18. Where would you rather sit on a 115 degree day, a metal bench or a kiddie pool full of water?

19. Calorimetry • Measurement of the amount of heat released during a reaction • Heat measured using a calorimeter (refer to diagram) • Calorimeter = device to measure the transfer of heat to water

20. Test Tube Thermometer Stirring Rod Water Sealed Container Reaction Simple Calorimeter

21. How to calculate heat… • Equation Q=mc∆T • To calculate the calories of heat transferred during a chemical reaction, multiply: • Q=heat absorbed or released • Mass of substance in calorimeter (g)  m • Change in temperature (oC)  T • Tfinal-Tinitial • Specific heat  This value is given to you • 1calorie/g.oCOr you can use 4.19 joules/g.oC

22. When working numerical problems we will quickly become confused if we don’t adopt a universal convention for when we use a positive sign or a negative sign. • Sign Convention for heat, Q If Heat is transferred into the system Q > 0 + absorbing heat, ENDOthermic If Heat is transferred out of the system Q<0 - releasing heat, EXOthermic

23. Examples (cal) • 2000 grams of water has its temperature raised by 3.0 oC. How much heat was produced? (1 cal/g.oC) • 6000 cal • How many calories must be added to 5000 g of water to change its temperature from 20oC to 30 oC? • 50,000 cal • If 500 g of water at 25oC loses 2500 calories, what will be the final temperature? • 20oC

24. More Examples (joules) • The temperature of a sample of iron with a mass of 10g changed from 50.4oC to 25oC with the release of 114J of energy. What is the specific heat of iron? • 0.449 J/goC • The temperature of a sample of water increases from 20oC to 46.6oC as it absorbs 5650 J of heat. What is the mass of the sample? (4.19J/goC) • 50.7 g

25. If 335 g water at 65.5 oC loses 9750 J of heat, what is the final temperature of the water? • 58.6oC

26. Temperature Conversions • Celsius Scale • Freezing point of water = 0 oC • Boiling point of water = 100 oCelsius • Interval between them is divided into 100 parts • Kelvin Scale • Absolute zero=NO movement • Lowest temperature theoretically possible • Absolute zero = 0 K = -273 oC • Freezing point of water = 273 K • Boiling point of water = 373 K • Size of degree is same as Celsius • Fahrenheit Scale • Freezing point of water = 32 F • Boiling point of water = 212 F

27. Converting • Kelvin = Celsius + 273 (K = oC + 273) • Celsius = Kelvin – 273 (oC = K -273) • Fahrenheit = (oC × 9/5) + 32 • Celsius = (F-32)(5/9)

28. Examples • Convert from K  oC • 110K • 476K • 295K • 1114K • Convert from oC K • 11 oC • 112 oC • -15 oC

29. Covert from oCF • -111oC • 0oC • 45oC • 323 oC • Convert from F oC • 45 F • 0 F • 115 F

30. Kinetic Theory • Matter has small particles in continuous motion • The faster a particle moves, the greater the kinetic energy • Temperature • Measure of the average kinetic energy of particles in the sample. • At absolute zero, the average kinetic energy is zero • Higher temperatures have a greater average kinetic energy • Samples at the same temperature have the same average kinetic energy

31. What kind of energy is this? • Figure 3.12: Equal masses of hot water and cold water separated by a thin metal wall in an insulated box.

32. The H2O molecules in hot water have much greater random motions than the H2O molecules in cold water.

33. The water samples now have the same temperature (50°C) and have the same random motions

34. Heat Transfer Transfer of heat into or out of a sample • Heat transferred into a sample can be used to increased the average kinetic energy of the particles • This causes an increase in temperature • When a sample cools, the particles lose kinetic energy. Heat is given off.

35. Heat Transfer • Heat can also enter or leave a sample without causing a change in temperature. • During the process of ice melting, the heat absorbed is used to rearrange the particles, not to increase the kinetic energy of the particles • No change in temperature occurs

36. Phases Changes that require energy • What happens to molecules in a solid as it melts? • Melting • The amount of energy (heat of fusion) required to melt one mole of a solid depends on the strength of the forces keeping the particles together (Intermolecular force).

37. When liquid water is heated, some molecules escape from the liquid and enter the gas phase. • If a substance is usually a liquid at room temperature (as water is), the gas phase is called a vapor.

38. Vaporization • Vaporization is the process by which a liquid changes into a gas or vapor. • As temperature increases, water molecules gain kinetic energy • At Boiling point, molecules throughout the liquid have the energy to enter the gas or vapor phase. • The amount of energy required to do this is the heat of vaporization

39. Sublimation • The process by which a solid changes directly into a gas without first becoming a liquid is called Sublimation. • Solid air fresheners and dry ice are examples of solids that sublime.

40. Condensation • Some phase changes release energy into their surroundings. • For example, when a vapor loses energy, it may change into a liquid. • Condensation is the process by which a gas or vapor becomes a liquid. It is the reverse of vaporization.

41. Water vapor undergoes condensation when its molecules lose energy, their velocity decreases. • The freezing point is the temperature at which a liquid becomes a crystalline solid. • When a substance changes from a gas or vapor directly into a solid without first becoming a liquid, the process is called deposition. • Deposition is the reverse of sublimation. Frost is an example of water deposition.

42. Temperature Energy or Time Phase changes Heating Curve

43. Practice Substance Freezing point (oC) Boiling Point (oC) Water 0.0 100.0 Gallium 23.0 89.0 Iron 723.0 2780.0 • At room temperature (27 oC), Iron is a solid, mixture, liquid or gas? • At 800 oC, Iron is a solid, mixture, liquid or gas? • During the process of heating water from 27 to 85 oC : • Did the potential energy change? Kinetic energy? • Is it an endothermic or exothermic reaction?

44. Entropy • Entropy (S) is a measure of the disorder or randomness of the particles that make up a system. • Spontaneous processes always result in an increase in the entropy of the universe. • Several factors affect the change in entropy of a system. • Changes of state • Dissolving of a gas in a solvent • Change in the number of gaseous particles • Dissolving of a solid or liquid to form a solution • Change in temperature

45. Overview of Conduction, Convection & Radiation • http://www.wisc-online.com/objects/index_tj.asp?objid=SCE304

46. Methods of Heat Transfer • Conduction: • Transfer of heat between substances that are in direct contact with each other • Occurs mainly in solid • Better conductor  More rapid heat transfer • Examples of good and poor conductors? • Convection: • Up and down movement (circulation) of gases and liquids caused by heat transfer • Does not occur in solid (molecules not free to move around) • Examples of convection?

47. Radiation: • Electromagnetic waves traveling through space • Does not require a medium to transfer heat • Waves transfer heat to the object • Examples of radiation heat transfer?

48. Practice • Boiling water over a campfire • Melting a tub of ice cream on the kitchen counter • Electric Stove versus Gas Stove • Which stove will boil water faster? Why? • Why is the second floor usually warmer than the first floor? Why?

49. Unit 4 Test • Nuclear Reactions • Alpha, Beta, Gamma, Electron Capture, Positron • Balancing Nuclear Equations • Half Life • Radioactive Decay • Energy • Definitions • Calculations • Energy • Temperature Conversions

50. Phase Changes • Diagram • Understanding how KE & PE relate to the diagram • Definitions • Conduction, Convection, & Radiation • Know examples