Thermodynamics. Chapters 21-24

Transcription

1 Thermodynamics Chapters

2 Temperature, Heat, and Expansion Chapter 21 2

3 History of Heat As late as 200 years ago, heat was regarded as a fluid, called caloric. It was believed that this caloric fluid flows from hot objects to cold objects. 3

4 History of Heat The scientific study of heat was motivated by the Industrial Revolution with its use of steam engines and machines. In the 19th century, James Prescott Joule (and others) showed that heat is a form of energy. 4

5 Heat Terminology Thermal energy is only called heat when it is moving between objects. Heat transfers thermal energy from one object to another (like work). Energy inside an object is properly called internal energy - the kinetic and potential energy of its particles. 5

6 Temperature and Kinetic Energy When you add heat energy to an object, its temperature may (or may not) increase. Temperature is proportional to the average KE of molecular translational motion. Temperature is not a measure of the total KE of all the molecules in a substance. 6

7 Temperature and Kinetic Energy There is twice as much KE in 2 Liters of boiling water than in 1 Liter of boiling water They both have the same temperature because they both have the same average KE 7

8 Heat Energy Heat: q or Q Energy in the process of flowing from a warmer object to a cooler object. Calorie: amount of energy needed to raise the temperature of 1 gram H 2 O 1 C cal = 1 kcal 8

9 Energy Ability to do work or produce a change. Potential Energy: Kinetic Energy: Sitting, food, fuel, electric Walking, driving, hair dryer 9

10 Potential Energy Chemical Energy: Energy stored in a substance due to its composition Types of atoms in substance Number and types of bonds Particular way the atoms are arranged 10

11 Kinetic Energy Energy of motion: Temperature measure of average KE of the molecules in a sample. All matter is in constant motion and therefore has kinetic energy. 11

12 Converting Energy Units Energy Unit Conversions 1 Joule (J) = cal 1 calorie (cal) = J 1 kilojoule (KJ) = 1000 J 1 Calorie = 1 kcal (nutritional calorie) 1 kcal = 1000 cal 12

13 Heat Transfer Radiation: Two objects near each other and heat is transferred Sunlight or stand near a fire Conduction: Convection: Two objects touch and heat is transferred Pan on stove or spoon in hot bowl of soup A fluid (gas or liquid) is used to transfer heat. Ocean currents or weather patterns 13

14 s07/sci/phys/energy/heattransfer/index.html 14

15 Specific Heat (C p ) The amount of energy required to raise the temperature of 1 g of a substance 1 ºC Measure of substances resistance to temperature change High C p = does not heat/cool rapidly Low C p = heats/cools very quickly 15

16 Specific Heat (C p ) High C p Water in the gulf Grass Pizza Crust Cloth Low C p Sand on the beach Sidewalk Sauce Leather 16

17 Specific Heat (C p ) Formula: Q = m C p T Q = heat in Joules M = mass in grams C p = specific heat in J/g ºC T = change in temp (T final T initial ) 10.0 g iron changed from 50.4 ºC to 25.0 ºC with a release of 114 J heat. Calculate C p C p = Q / m T 114J / (10.0g 25.4ºC) = J/g ºC 17

18 Specific Heat (C p ) Formula: Q = m C p T Q = heat in Joules m = mass in grams C p = specific heat in J/g ºC T = change in temp (T final T initial ) If the temperature of 25g granite, which has a C p of.803 J/g ºC is heated from 25 ºC to 55ºC, how much heat did the granite absorb? C p = Q / m T Q = m T C p Q = (25.0g)(30ºC) (0.803 J/g ºC) Q = 602 J 18

19 Specific Heat (C p ) Type 1. Heat Transferred (q) is the unknown: Aluminum has a specific heat of J/g x o C. How much heat is lost when a piece of aluminum with a mass of g cools from a temperature of o C to a temperature of 22.0 o C? 19

20 Specific Heat (C p ) Type 2. mass (m) is the unknown: The temperature of a sample of water increases by 69.5 o C when J are applied. The specific heat of liquid water is 4.18 J/g x o C. What is the mass of the sample of water? 20

21 Specific Heat (C p ) Type 3. change in temperature ( T) is the unknown: 850 calories of heat are applied to a 250 g sample of liquid water with an initial temperature of 13.0 o C. Find a) the change in temperature and b) the final temperature. (remember, the specific heat of liquid water, in calories, is 1.00 cal/g x o C.) 21

22 Specific Heat (C p ) Type 4. Specific Heat (Cp) is the unknown: When J of heat are applied to a 350 g sample of an unknown material the temperature rises from 22.0 o C to o C. What must be the specific heat of this material? 22

23 Heat of Reaction: Calorimeter 23

24 Heat of Reaction: Simple Calorimeter 24

25 Nutritional Calorie Energy Unit Conversions 1 Joule (J) = cal 1 calorie (cal) = J 1 kilojoule (KJ) = 1000 J 1 Calorie = 1 kcal (nutritional calorie) 1 kcal = 1000 cal A typical breakfast contains 230 nutritional Calories. 230 Cal x 1000 cal/1 Cal= 230,000 cal 230,000 cal x J/1 cal= 962,320 J 25

26 Measuring Heat Exchange Q = m C p T A piece of metal with a mass of 4.68g absorbs 25 J of heat when its temperature increases by 182 C. What is the specific heat of the metal? C p = 25 J / (4.68 g)(182 C) =.029 J/g C 26

27 Measuring Heat Exchange Heat loss = heat gain Q = -Q m C p T = -(m C p T) A 75.0g piece of metal is placed in boiling water until its temperature is C. A calorimeter contains g water at a temperature of 24.2 C. The metal is removed from the boiling water and immediately placed in the water in the calorimeter The final temperature of the metal and water in the calorimeter is 34.9 C. Assuming perfect insulation of the calorimeter, what is C p of the metal? (100 g)(4.18j/g C)(10.7 C)=- (75g)(C p )(-65.1 C) 27

28 Measuring Heat Exchange Heat loss = heat gain Q = -Q m C p T = -(m C p T) (100 g)(4.18j/g C)(10.7 C)=- (75g)(C p )(-65.1 C) (100 g)(4.18j/g C)(10.7 C)=- (75g)(C p )(-65.1 C) (100 g)(4.18j/g C)(10.7 C)= (C p ) (75g) (65.1 C) C p = J/g C 28

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30 Temperature Temperature is a measurement of the thermal energy of a system. Specifically temperature is the measurement of the average kinetic energy of all particles in a system 30

31 Temperature So, the average amount of movement, or kinetic energy of the particles that compose a material determines in what material state it resides. 31

32 Temperature Scales Kelvin: Universal temperature measurement used in science. Celsius: A variant on the Kelvin scale. This is the standard scale for the metric system. The Celsius scale is standardized to coincide with the freezing and boiling points of water. Fahrenheit: Traditional English measurement scale. 32

33 Temperature Scales (cont d) 33

34 Conversion Formula Celsius to Kelvin K = C Kelvin to Celsius C = K Fahrenheit to Celsius C = (F - 32) x 5/9 Celsius to Fahrenheit F = (C x 9/5)

35 Temperature vs. Energy What states of matter do the flat areas correspond to? Where do you find this same type of thing on a phase diagram? Why can matter in the same state nonetheless have considerably higher energy? 35

36 Phase Transitions 36

37 Water CO 2 ml 37

38 Expansion of Water Water has maximum density at 4º C Water crystallizes into an open hexagonal form 38

39 Expansion of Water The hexagonal lattice contains more space than the liquid state. 39

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41 Expansion of Water Water at ordinary temperatures contracts and increases in density as it is cooled. at about 4 C it reaches a maximum density and then decreases in density as it approaches the freezing point. 41

42 Water Water has the highest specific heat of any common substance, 1 calorie/gm C = J/gm C provides stability for the temperature of the human body, makes it an effective cooling agent The high heat of vaporization of water makes it an effective coolant for the human body via evaporation of perspiration, extending the range of temperatures in which humans can exist 42

43 Heat of Fusion Molar: the amount of heat necessary to melt (or freeze) 1.00 mole of a substance at its melting point. Mass: the amount of heat necessary to melt (or freeze) 1.00 gram of a substance at its melting point. 43

44 Heat of Vaporization Molar: the amount of heat necessary to boil (or condense) 1.00 mole of a substance at its boiling point Mass: the amount of heat necessary to boil (or condense) 1.00 gram of a substance at its boiling point 44

45 Phase Transitions 45

46 Sample Problem Heat of fusion of ice : 333 J/g Heat of vaporization of liquid water : 2257 J/g 46

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48 Laws of Thermodynamics Thermodynamics is the study of the inter-relation between heat, work and internal energy of a system Three Laws of Thermodynamics 48

49 1 st Law Energy can be changed from one form to another, but it cannot be created or destroyed. The total amount of energy and matter in the Universe remains constant, merely changing from one form to another. Energy is always conserved. 49

50 2 nd Law the Law of Increased Entropy in all energy exchanges, if no energy enters or leaves the system, the potential energy of the state will always be less than that of the initial state. In an isolated system, a process can occur only if it increases the total entropy of the system. Heat cannot spontaneously flow from a material at lower temperature to a material at higher temperature. 50

51 2 nd Law, cont d It is impossible to convert heat completely into work. In an isolated system, concentrated energy disperses over time, and consequently less concentrated energy is available to do useful work. Work can be obtained from these nonequilibrium differences, but that loss of heat occurs, in the form of entropy, when work is done 51

52 Entropy 52

53 Entropy nature tends from order to disorder in isolated systems nature tends toward maximum entropy for any isolated system 53

54 Entropy it is a measure of the "multiplicity" associated with the state of the objects. If a given state can be accomplished in many more ways, then it is more probable than one which can be accomplished in only a few ways. 54