Energy. Energy (E) is the ability to do work. Many types, but we can say 3 main types: Radiant Potential Kinetic

Transcription

1 ENERGY

2 Energy n n n Energy (E) is the ability to do work. Many types, but we can say 3 main types: Radiant Potential Kinetic

3 US Energy Consumption by Source

4 Light Energy Visible and Invisible Radiant Energy Travels in waves over distances Electromagnetic waves n Waves that spread out in all directions from the source n Visible light, UV light, Infra Red Radiation, X-rays, microwaves, radio waves

5 Potential Energy (PE) Stored Energy Due to position n Gravitational PE n Elastic PE Chemical bonds n Chemical PE n Nuclear energy n Fuels n Attractions between molecules

6 Energy of motion Atomic vibrations Molecular movement n Vibration n Rotation n Translation Movement of subatomic particles Kinetic Energy (KE)

7 Kinetic Energy Can be calculated:

8 How are each type shown here?

9 How are each type shown here? Radiant Rainbow = visible light Kinetic Windmill moving Potential All molecules store energy n Water in clouds n Air n Materials the windmill is made from, the plants at the bottom

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11 Where is the PE? KE?

12 Temperature Scales: Measuring that Thermal Energy Boiling Freezing Fahrenheit ( o F) Celsius ( o C) Kelvin

13 A Note on the Fahrenheit Scale NEVER use it in this class. Ever. Only Belize and the US use this scale. Gabriel Fahrenheit made great thermometers. His scale was replicated the world over because of this. But if you stop and think about it, does 32 F for freezing make sense, or 212 F for boiling? 180 degrees separates them. 100 degrees, as in the Celcius scale (sometimes called the Centigrade scale) makes much more sense. Fahrenheit based 0 F on the freezing point of water mixed with NH 4 Cl, and 32 F for freezing water, and 96 F for human body temperature (he was off by 2.6 ). Why? Because he felt like it and it was easy to draw lines at those intervals. ( According to a letter Fahrenheit wrote to his friend Herman Boerhaave, [8] his scale was built on the work of Ole Rømer, whom he had met earlier. In Rømer s scale, brine freezes at 0 degrees, ice melts at 7.5 degrees, body temperature is 22.5, and water boils at 60 degrees. Fahrenheit multiplied each value by four in order to eliminate fractions and increase the granularity of the scale. He then re-calibrated his scale using the melting point of ice and normal human body temperature (which were at 30 and 90 degrees); he adjusted the scale so that the melting point of ice would be 32 degrees and body temperature 96 degrees, so that 64 intervals would separate the two, allowing him to mark degree lines on his instruments by simply bisecting the interval six times (since 64 is 2 to the sixth power). I took this from Wikipedia.

14 Kelvin Temperatures Based on absolute zero (0 K, -273 o C) The temperature at which ALL KE stops NO molecular motion. Lowest temperature theoretically possible n Can t really get there in real life n (See 3 rd Law of Thermodynamics in a few slides) K = o C Technically , but we can stop at 3 significant digits

15 Why do we need the Kelvin scale? Two reasons We need a scale that is relative to molecular motion for certain topics n You can t use negative numbers to indicate motion when it IS present n -20 C makes NO sense in light of indicating motion n And 40 C ISN T twice as much motion as 20 C, n (40K IS twice the motion of 20K) Because when working with equations, can t use zero n We get undefined answers if we divide n We get answers of 0 if we multiple n And those answers would NOT make sense if compared to answers calculated with a positive or negative number

16 The 4 Es: Energy, Exergy, Entropy, & Enthalpy Energy (E): The ability to do work Entropy (S): The measure of the disorder of a system Exergy: The energy available to do work No symbol Enthalpy(H): The thermal energy (heat) content of a system There will be more on these!

17 Thermodynamics The study of energy flow inter-relation between heat, work, and energy of a system Summary of the three laws: 1. The energy in the universe is constant 2. Things get more disorganized over time in a system until everything is equal 3. You can t reach absolute zero

18 1 st Law of Thermodynamics The energy in the universe is constant E=mc 2 Law of Conservation Matter n Matter can not be created or destroyed Law of Conservation of Energy n Energy can not be created or destroyed n However, matter and energy can both change forms in chemical reactions n Can also interconvert between matter and energy in NUCLEAR reactions (more on this later this year.) Summed up: You can not win. You can t get something for nothing because energy and matter are conserved.

19 Energy Time

20 Before the 2 nd Law Entropy (S) is a measure of DISORGANIZATION in a system (this simply put; there is a much more complicated description about the unavailable energy to do work) Anything disorganized has higher entropy than something organized Exergy is the Energy available to do work

21 2 nd Law of Thermodynamics Things get more disorganized over time in a system until everything equilibrium is reached (everything is equal) Heat flows from hot to cold, not the reverse Law of Entropy n By nature, things get more disorganized to spread out energy and matter The quality of the energy (which is exergy) decreases over time Summed up: You can not break even. You can not return to the same energy state because things get more disorganized (gain entropy)

22 Exergy and Energy The energy of the universe is constant, but exergy is constantly consumed. This can be compared with a tooth-paste tube: When you squeeze the tube (= conduct any process) the paste (= exergy) comes out. You can never put the paste back in the tube again (try!), and in the end you have only the tube itself (= low-exergy) left. When you squeeze the tube, the depressions (= entropy) will increase. (The entropy of a system increases when exergy is lost) But you can never take the depressions in the tube and 'un-brush' your teeth. (I.e. entropy is not negative exergy.) When you buy energy from the electricity network, you actually buy exergy. You can find the energy as room temperature heat after some time, but you can not take that room temeperature energy back to the electricity company and ask for money back. They won't accept it.

23 Energy and Matter Gain Entropy Over Time

24 Exergy: The Energy available to do work

25 3 rd Law of Thermodynamics You can t reach absolute zero and expect things to happen At absolute zero, all kinetic motion ceases. And that energy needs to go somewhere. It goes to something else. And gets transferred back until everything is at an equal temperature. Summed up: You can not get out of the game, because absolute zero is unobtainable.

26 Law of Conservation of Energy Energy cannot be created or destroyed but it CAN change forms. Example: Burning wood in a fire The energy in chemical bonds is released as heat (KE and PE), light (RE), sound (KE) These forms of energy are less useful n have less exergy

27 Radiant Energy: EM Waves Potential Energy: Stored Kinetic Energy: Motion

28 The CPE in these items could:

29 Rio Summer Olympics Proposed Solar Waterfall photos/architecture/ solartower.asp

30 Combinations of PE and KE are very common on a large scale KE and PE animation

31 PE and KE

32 When E changes forms The amount of energy one thing loses is gained somewhere else. E lost = E gained (Law of Conservation of Energy) But the E gained is usually not all in one place (2 nd Law of thermodynamics) n It is spread out (more entropy) n Often in the forms of heat and light n Which are less useful (less exergy)

33 Energy Transformations

34 Thermal Energy: KE + PE on the small scale What s up with Temperature vs Heat? Temperature is related to the average kinetic energy of the particles in a substance.

35 Thermal energy relationships à As temperature increases, so does thermal energy (because the energy of the particles increased). à If the temperature stays the same, the thermal energy in a more massive substance is higher (because it is a total measure of energy).

36 Heat The flow of thermal energy from one object to another. Cup gets cooler while hand gets warmer Heat always flows from warmer to cooler objects. Ice gets warmer while hand gets cooler

37 Heat and Temperature Heat: the measure of the flow of RANDOM kinetic energy Temperature: the measure of heat So temperature is a measure of kinetic energy of the particles of a substance * Sometimes heat is radiated as IR (infra-red radiation, a form of radiant energy)

38 Thermal Energy Thermal Energy is the total of all the (kinetic and potential) heat energy of all the particles in a substance. PE from how the molecules are placed relative to each other (attractions) Farther = more PE, just like how something farther off the ground has higher gravitational PE

39 Exothermic and Endothermic Processes Endothermic Energy is being gained/ absorbed by the object or substance (called the system) from the surroundings Have positive change in enthalpy values (+ΔH) Exothermic Energy is lost/ released from the object or substance (called the system) to the surroundings Have negative change in enthalpy values (-ΔH)

40 Exothermic Reactions? I studied them before they were cool.

41 The big picture How do we see this energy cycling in the real world, and not just as a part of Chemistry class? Around the house? In the environment? While thinking about a car?

42 If the cup is the system, it is undergoing an exothermic process because it is losing heat to the surroundings (hand) Cup gets cooler while hand gets warmer If the ice is the system, it is undergoing an endothermic process because it is absorbing heat from the surroundings (hand) Ice gets warmer while hand gets cooler

43 Which process is endothermic? Which is exothermic?

44 Trophic Levels and Energy Consumers are all heterotrophs 3 Consumers: Carnivores and Omnivores 2 Consumers: Carnivores and Omnimores Energy Out; 90% per level 1 Consumers: Herbivores Producers: Autotrophs

45 Trophic Levels and Energy Producers:

46 Trophic Levels and Energy 3 Consumers 2 Consumers 1 Consumers While this shows only a 1 consumer, all animals lose E the same way; high levels lose more to motion than do lower levels

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50 Class Question: How does a car exhibit all types of energy? Explain!

51 Energy Loss in a Car

52 Can the world really run out of Energy? World-Wide Energy Sources, (2007)

53 PHASE CHANGES & ENERGY

54 Phase Diagrams Tell what state of matter a material is in at a given temperature and pressure The triple point is the pressure and temperature when a solid, liquid, and a gas of the same substance exist at equilibrium Equilibrium: When there is no net change n Here referring to changes in state n Can also refer to temperature and chemicals The critical point is the temperature above which a substance will always be a gas, regardless of pressure Fullerton Phase Diagram Explorer Link

55 Phase Diagrams

56 Phase Diagram for Water

57 A few terms Freezing Point - The temperature at which the solid and liquid phases of a substance are in equilibrium at atmospheric pressure. The same temperature as the melting point Boiling Point - The temperature at which the vapor pressure of a liquid is equal to the pressure on the liquid. Vapor Pressure- The pressure at which the vaporization rates are equal to condensation rates

58 Measuring Vapor Pressure

59 Energy and Phase Changes

60 Phase Changes Enthalpy(H): The heat (thermal energy) content of a system

61 States of Matter and Entropy The states are NOT plateaus because entropy is NOT constant. This isn t a phase change diagram.

62 Energy and Matter and Connected Any change in matter ALWAYS is accompanied by a change in energy

63 Phase Changes and Energy

64 Heating Curve Temperature, C Time, min

65 States on a heating curve States of Matter on a Heating Curve GAS Temperature, C LIQUID SOLID Time, minutes

66 Changes of State on a Heating Curve Changes of state on a heating curve Two states of matter exist on the plateaus Temperature, C LIQUID becoming VAPOR SOLID becoming LIQUID Time, minutes

67 Remember: Phase Changes and Energy

68 Why does temperature remains constant when melting or boiling? During melting or boiling, energy is absorbed from the surroundings Due to the increase in the thermal energy of the particles from the increase in PE of the particles n Molecules are n moving apart n breaking attractions which n Absorbs latent (hidden) heat n can not be measured on a thermometer Substance (system) gets warmer

69 The E s and Heating Endothermic process Energy is absorbed from surroundings Entropy increases Enthalpy is positive (+ΔH) since heat added Exergy decreases

70 Why does temperature remains constant when freezing or condensing? During freezing or condensing, energy is released to the surroundings Due to the decrease in the thermal energy from the decrease in PE of the particles n Molecules are n moving closer n forming new attractions that are n Releasing latent (hidden) heat n can not be measured on a thermometer Substance (system) gets colder

71 The E s and Cooling Exothermic process Energy is lost to surroundings Entropy decreases Enthalpy is negative (-ΔH) since heat is lost Exergy increases

72 Processes of a Heating Curve Temperature, C VAPORIZATION GAS FUSION (MELTING) LIQUID SOLID Time, minutes

73 What happens during each segment

74 Cooling Curve: The Reverse of a Heating Curve Temperature, C Time, minutes

75 Cooling Curve: The Reverse of a Heating Curve Cooling curve Temp, rature, C C gas Gas and liquid present condensation liquid Liquid and solid present Freezing solid Time, min

76 Energy and Phase Change in Nature:

77 Energy and Phase Change in Nature:

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79 Measuring the Energy of Phase Changes The math of thermal energy flow

80 REMEMBER: Energy and Matter and Connected Any change in matter ALWAYS is accompanied by a change in energy This includes changes in temperature and/ or phase

81 Specific Heat : c Things heat up or cool down at different rates. Land heats up and cools down faster than water, and aren t we lucky for that!?

82 Specific heat is the amount of heat required to raise the temperature of 1 kg of a material by one degree C c water = J / g C the number is high; water holds its heat c sand = J / g C less E than water to change it; it doesn t hold heat as well as water does This is why land heats up quickly during the day and cools quickly at night and why water takes longer.

83 Why does water have such a high specific heat? water metal Water molecules form strong attractions with other water molecules; it takes more heat energy to break those attractions than other materials with weaker forces of attraction between them.

84 Specific Heat Capacities of Selected Substances c water = J / g C c ice = 2.09 J / g C c steam = 1.99 J / g C c sand = J / g C c Al = 0.90 J / g C c Fe = J / g C

85 Heat can be Transferred even if there is No Change in State q = mc T

86 Remember this? Which is process is endothermic? Which is exothermic? Now we care about how much energy is being transferred, and are ready to calculate that change.

87 Calculating Changes in Energy: The Calorimetry Equation q = mcr T q = change in thermal energy (+) value means heat is absorbed (-) value means heat is released m = mass of substance r T = change in temperature (T final T initial ) c = specific heat of substance Each substance has a different c (see CRH, p ) Different states of matter for the same substance may have a different c

88 Notice that the negative sign on q signals heat lost by or transferred OUT of Al. Was this an endothermic or exothermic process? Specific Heat Capacity Problems If 25.0 g of Al cool from 310 o C to 37 o C, how many joules of heat energy are lost by the Al? heat gained or lost = q = mc T where T = T final - T initial q = (25.0g) (0.897 J/g o C)(37-310) o C q = J

89 Specific Heat Capacity Problems If 25.0 g of Al cool from 310 o C to 37 o C, how many joules of heat energy are lost by the Al? heat gained or lost = q = mc T where T = T final - T initial q = (25.0g) (0.897 J/g K)(37-310)K q = J Notice that the negative sign on q signals heat lost by or transferred OUT of Al. Was this an endothermic or exothermic process?

90 Or Heat Transfer can cause a Change of State Changes of state involve energy changes at constant T Ice J/g (heat of fusion) -----> Liquid water Is there an equation? Of course!

91 Or Heat Transfer can cause a Change of State Changes of state involve energy at constant T H 2 0 (s) +333J/g à H 2 0 (l) Ice J/g (heat of fusion) à Liquid water q = mδh fusion m= mass ΔH fusion = the enthalpy of melting the change in thermal energy associated with melting Units are J/g or KJ/Kg

92 q = mδh fusion WHY DO I NEED THIS WHEN I HAVE q = mc T? Well, when a phase changes THERE IS NO change in temperature but there is definitely a change in energy!

93 Sample Problem: How much heat energy is required to melt 25.0g of ice, (assuming constant temperature of O C)? m= 25.0g q = mδh fusion ΔH fusion =333J/g for water q= (25.0g)333J/g =8385J Value is positive, which means heat is absorbed, which makes sense!

94 Latent heat* and the PE of particles molecule Regular arrangement breaks up strong attraction weak attraction *Latent means hidden. Latent heat is the thermal energy (potential energy) associated with the attractions between molecules, and can not be measured with a thermometer.

95 Latent heat and the PE or particles Energy has to be supplied to oppose the attractive force of the particles. PE ê as molecules separate PE related to the forces of attraction between the particles solid è liquid or liquid è gas average potential energy ê

96 Latent heat and PE The transfer of energy does not change the KE. Temperature does not change. latent heat = change in PE between molecules during change of state Video and song: v=jaagqui9nvy

97 Remember Energy changes accompany changes in state; either: Energy is added (endothermic) Gain thermal energy Molecules Move more (gain KE) Separate (gain PE from broken attractions between molecules) Have a higher entropy Are more disorganized Or Energy is removed (exothermic) Molecules move less Lose thermal energy Move less (lose KE) Move closer (lose PE from new attractions between molecules) Have lower entropy Get more organized

98 Latent Heats You have a certain energy change associated with changing state. These values are usually reported for fusion and vaporization as: ΔH fusion = (latent) Heat of fusion (melting) Δ H vaporization = (latent) Heat of vaporization Δ H sublimation =(latent) Heat of sublimation Different materials have different values for each

99 What about freezing and condensation? Values for freezing and condensation are not typically listed, but are the negative values of those for fusion and vaporization because the energy transferred is the same, but in the opposite direction (latent) Heat of freezing= -ΔH fusion (latent) Heat of condensation= -Δ H vaporization

100 Enthalpy changes with phase changes

101 Enthalpy values for H 2 O H fusion = 334 J/g H vaporization = 2259 J/g H sublimation = J/g

102 From: / hyper physic s.phyastr.gs u.edu/ hbase / tables / phase.html #c1

103 hyperphy sics.phyastr.gsu.e du/ hbase/ tables/ phase.htm l#c2

104 Summing it all up: How do you know what to do to calculate energy changes? Check to see if there is a temperature change. If yes, use q=mcδt. Also, check to see if there is a phase change. If yes, you need to use q= Δ H fusion mass or q= Δ H vaporization mass depending on which one applies* or both if there are two phase changes *If the material freezes or condenses. You can use the negative value Δ H fusion or Δ H vaporization

105 How much energy is required to change 0.5 kg of water at 0 C to ice? Things you know: m = 0.5 kg= 500.g There is no temperature change, and there is change of state (freezing) The water is going to freeze So.. this all tells you to use ΔH fusion (negative of melting value) in q= ΔH fusion m q= (-334J/g)(500.g)= -1.67E5J (The negative value makes sense since you are cooling the water, so energy leaves)

106 How much energy is required to melt 0.5 kg of ice at 0 C temperature raised to 80 C? Total energy required = latent heat (ice at 0 C water at 0 C) + energy (water: 0 C 80 C) = mδh f + mc ΔT = (500g)(334 J/ g) + (500.g x 4.184J/g C x 80 C) = ( J) + (167360J) = =334360J= 3.34E5 J (Value is positive, makes sense since you add E to heat water)

107 Heat & Changes of State What quantity of heat is required to melt 500. g of ice and heat the water to steam at 100. o C? Heat of fusion of ice = 334 J/g Specific heat of water = J/g C Heat of vaporization = 2259 J/g +334 J/g J/g

108 Putting it all together So if I want the total heat to take ice and turn it to steam I need to add values from 3 steps 1. To melt the ice I need to multiply the heat of fusion with the mass q = H fusion m 2. Then, there is moving the temperature from 0 C to 100 C. For this there is a change in temperature so we use q= mc T 3. That just takes us to 100 C, what about vaporizing the molecules? We need q= H vaporization m Add up all the values, and you have it. (However, if you are taking it from below the freezing point to above 100 C, you need to add in the changes with q=mc T there, too!)

109 And now More! Heat & Changes of State How much heat is required to melt 500. g of ice and heat the water to steam at 100 o C? 1. To melt ice : q = m H fusion q = (500.g)(334 J/g) = 1.67 x 10 5 J 2. To raise water from 0 o C to 100 o C : q = mc T q = (500. g)(4.184 J/g o C)(100-0 o C) = 2.1 x 10 5 J 3. To evaporate water at 100 o C: q = m H vaporization q =(500.g)(2259 J/g) = 1.13 x 10 6 J 4. Total heat energy = 1.51 x 10 6 J = 1510 kj

110 Maybe a picture can help.

111 Putting it all together: How are matter and energy related? What influences does energy have on matter? What does this tell us about the world as we know it?

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113 Making Pizza: Changing Matter Describe the pizza making process in terms of: Matter States (s, l, g) Elements, compounds, mixtures n Homogeneous and heterogeneous mixtures Properties and changes Both n chemical and physical n Intrinsic (intensive) and extrinsic (extensive) Energy