1. Topic 3 Thermal physics

2. Temperature TEMPERATURE determines the direction of flow of thermal energy between two bodies in thermal contact HOT COLD This is another way of saying that if an object is hotter than another, heat energy will flow from the hotter object to the colder! Obvious, but important to remember!

3. Temperature Temperature is also a measure of the average random kinetic energy of the particles in a substance. Note that they are not all travelling at the same speed.

4. Thermal equilibrium Two bodies in thermal contact will eventually reach the same temperature. The two bodies are now said to be in thermal equilibrium. WARM

5. Kelvin Temperature Kelvin temperature is proportional to the average random kinetic energy of the particles in a substance. Note that they are not all travelling at the same speed.

6. Absolute/Kelvin temperature and Celsius T (in Kelvin) = T (in degrees Celcius) + 273 Just to mention for now that zero Kelvin is the lowest possible temperature.

7. Measuring temperature The thermometer has to be placed in thermal contact with whatever is being measured until the thermometer and object are in thermal equilibrium. 0°C 100°C We’ll measure some temperatures in a moment

8. So what is heat? Heat is the amount of thermal energy. For example, the sparks from a sparkler are at around 800°C but do not burn your skin. However, a hot cup of tea at around 100°C will burn your hand badly. This is because the tea contains more heat energy, even though it is cooler.

9. Internal Energy This is the sum of the kinetic energies and potential energies of the particles in a substance If you imagine the forces between particles as a spring, you can see if the particles are pulled apart or squashed together that energy is stored in the spring. Similarly there is potential energy between the particles in a substance.

10. Heat transfer Conduction, convection and radiation. I’m not going to say anything about these, you should know it already. Can you construct a heat transfer mind-map? Colours Few words Connections Drawings Pages 108 to 113 IGCSE Physics

11. Moles! One mole of a substance contains the same number of molecules/atoms as in 12 grams of carbon-12. This number (of atoms or molecules) is known as the Avogadro constant (N A ) which is equal to 6.02 x 10 23 You need to learn this definition.

12. Moles! It follows therefore that 7g of lithium (atomic mass 7), 20g neon (atomic mass 20) or 39 g potassium (atomic mass 39) all contain the same number of atoms (1 mole or 6.02 x 10 23 atoms)

13. Moles! The number of moles of a substance can thus be found by dividing the mass of substance by its relative atomic or molecular mass n = mass/RAM

14. Example How many moles of sulphur atoms are there in 80g of sulphur? How many grams of carbon would have the same number of atoms?

15. Example How many moles of sulphur atoms are there in 80g of sulphur? How many grams of carbon would have the same number of atoms? N = mass/RAM = 80/32 = 2.5 moles

16. Example How many moles of sulphur atoms are there in 80g of sulphur? How many grams of carbon would have the same number of atoms? N = mass/RAM = 80/32 = 2.5 moles Mass of carbon = RAM x n = 12 x 2.5 = 30 g

17. Heat Capacity The relationship between the amount of heat energy a substance requires to raise its temperature by a given amount is called its thermal capacity.It is measured in J.°C -1 or J.K -1.

18. Definition to learn Thermal capacity is the amount of energy needed to raise the temperature of a substance by 1K.

19. Why are thermal capacities different? When a substance is heated, its internal energy increases (potential and kinetic). The stronger the force between the particles in the substance, the more heat energy goes into potential energy (and less into kinetic), so the temperature rise is less than in substances with little force between particles. Obviously the more partciles there are too, the more heat energy can be absorbed.

20. Calculations using Thermal capacity Energy absorbed = Thermal capacity x Temp rise E = QΔT J J.°C -1 °C

21. Specific heat capacity Specific heat capacity is the amount of energy needed to raise the temperature of unit mass of a substance by 1K Specific heat capacity of water = 4186 J.kg -1.°C -1 Specific heat capacity of kerosene = 2010 J.kg -1.°C -1 Specific heat capacity of mercury = 140 J.kg -1.°C -1

22. Calculations using S.H.C. Energy absorbed = Mass x Specific Heat capacity x Temp rise Q = mcΔT Jkg J.kg -1.°C -1 °C

23. For example 500 g of olive oil is heated until its temperature rises by 120°C. If the specific heat capacity of olive oil is 1970 J.kg -1.°C -1, how much heat energy was used? Energy absorbed = Mass x Specific Heat capacity x Temp rise Energy absorbed = 0.5 x 1970 x 120 Energy absorbed = 118200 J

24. Solids, liquids and gases

25. Melting? Changes in kinetic energy and potential energy?

26. Evaporation? Changes in kinetic energy and potential energy?

27. Condensing? Changes in kinetic energy and potential energy?

28. Freezing? Changes in kinetic energy and potential energy?

29. Evaporation Consider a beaker of water at room temperature

30. Evaporation The molecules of water are moving around at different speeds, some fast, some slow. speed of molecule (m/s) # of molecules at a particular speed Average speed

31. Evaporation If a molecule is at the surface, and moving fast enough, it may escape the liquid. This is called evaporation. Freedom!

32. Evaporation Since the average speed of the remaining molecules must now be lower, the temperature of the liquid drops (since temperature is a measure of the kinetic energy of the molecules). Freedom!

33. Evaporation Evaporation can thus take place at any temperature.

34. Increasing the rate of evaporation Increasing the temperature means that more molecules are moving fast enough to escape.

35. Increasing the rate of evaporation Increasing the surface area means that more molecules are at the surface.

36. Increasing the rate of evaporation Increasing the air flow over the surface so that molecules are carried away before they can fall back into the liquid

37. Increasing the rate of evaporation Decreasing the humidity of the surrounding atmosphere to stop water molecules from the atmosphere entering the liquid.

38. Boiling Boiling occurs when vapour is produced in the body of the liquid. This only happens at the boiling point of the liquid. The bubble contains only water vapour, not air!

39. To summarize: Evaporation takes place only at the surface of the liquid and can take place at any temperature.

40. To summarize: Boiling occurs when vapour is produced in the body of the liquid. This only happens at the boiling point of the liquid. Boiling means bubbles!

41. Latent heat

42. In last year’s experiment, you will have noticed that the temperature of the salol stopped changing as the salol changed from a liquid to a solid. Temp (°C) Time (mins) Melting point

43. Latent heat When the molecules of a substance settle into the regular pattern of a solid, energy is released as bonds are formed. This energy released is called latent heat. This stops the temperature from falling. (“latent” = “hidden”)

44. Latent heat The opposite happens when a solid makes. Heat is needed to break the bonds between the solid particles (increasing their potential energy instead of raising the temperature (kinetic energy)) Temp (°C) Time (mins) Melting point solid liquid

45. Specific Latent heat The specific latent heat of a substance tells us how much energy is needed to change the state of 1 kg of substance at constant temperature. Solid to liquid/liquid to solid or liquid to gas/gas to liquid

46. Specific Latent Heat The specific latent heat of fusion (melting) of ice at 0 ºC, for example, is 334000 J.kg -1. This means that to convert 1 kg of ice at 0 ºC to 1 kg of water at 0 ºC, 334000 J of heat must be absorbed by the ice. 1 kg 334000 J absorbed All at 0°C

47. Specific Latent Heat Conversely, when 1 kg of water at 0 ºC freezes to give 1 kg of ice at 0 ºC, 334000 J of heat will be released to the surroundings. 334000 J released 1 kg All at 0°C

48. Specific Latent Heat of Vaporisation For water at its normal boiling point of 100 ºC, the latent specific latent heat of vaporization is 2260000 J.kg -1. This means that to convert 1 kg of water at 100 ºC to 1 kg of vapour at 100 ºC, 2260000 J of heat must be absorbed by the water. All at 100°C 2260000 J input 1 kg

49. Latent heat Conversely, when 1 kg of steam at 100 ºC condenses to give 1 kg of water at 100 ºC, 2260 kJ of heat will be released to the surroundings. All at 100°C 2260000 J released 1 kg

50. Another formula! Energy = mass x specific latent heat Q = mL

51. An example calculation Calculate the amount of heat required to completely convert 50 g of ice at 0 ºC to steam at 100 ºC. The specific heat capacity of water is 4.18 kJ.kg -1.°C -1. The specific latent heat of fusion of ice is 334 kJ.kg -1, and the specific heat of vaporization of water is 2260 kJ.kg -1. 50g 0°C 100°C

52. An example calculation Heat is taken up in three stages: 1. The melting of the ice. 2. The heating of the water. 3. The vaporization of the water. 0°C 100°C

53. Stage 1 1. Heat taken up for converting ice at 0ºC to water at 0ºC mass of water x latent heat of fusion = 0.050 (kg) x 334 (kJ.kg -1 ) = 16.7 kJ 0°C

54. Stage 2 2. Heat taken up heating the water from 0 ºC to the boiling point, 100 ºC mass of water x specific heat capacity x temperature change = 0.05 (kg) x 4.18 (kJ.kg -1.°C -1 )x 100 (ºC) = 20.9 kJ 0°C 100°C

55. Stage 3 3. Heat taken up vaporising the water mass of water x latent heat of vaporization 0.05 (kg) x 2260 kJ.kg -1 = 113 kJ 100°C

56. The answer The sum of these is 16.7 + 20.9 + 113 = 150.6 kJ (151 kJ)

57. Formulae you will probably need Q = mcΔT (warming/cooling – change in temp) Q = mL (change of state at constant temp) Power = Energy/time Intensity(W.m -2 ) = Power(W)/area(m 2 ) (waves topic) Many questions will involve sun shining on ice, probably linked with questions about global warming.

58. Kinetic theory/ideal gas We can understand the behaviour of gases using a very simple model, that of an “ideal” gas. The model makes a few simple assumptions;

59. Ideal gas assumptions The particles of gas (atoms or molecules) obey Newton’s laws of motion. You should know these by now!

60. Ideal gas assumptions The particles in a gas move with a range of speeds

61. Ideal gas assumptions The volume of the individual gas particles is very small compared to the volume of the gas

62. Ideal gas assumptions The collisions between the particles and the walls of the container and between the particles themselves are elastic (no kinetic energy lost)

63. Ideal gas assumptions There are no forces between the particles (except when colliding). This means that the particles only have kinetic energy (no potential) Do you remember what internal energy is?

64. Ideal gas assumptions The duration of a collision is small compared to the time between collisions.

65. Pressure – A reminder Pressure is defined as the normal (perpendiculr) force per unit area P = F/A It is measured in Pascals, Pa (N.m -2 )

66. Pressure – A reminder Collisions of the gas particles with the side of a container give rise to a force, which averaged of billions of collisions per second macroscopically is measured as the pressure of the gas Change of momentum

67. Explaining the behaviour of gases When we heat a gas at constant volume, the pressure increases. Why? Increased average kinetic energy of the particles means there are more collisions with the container walls in a period of time and the collisions involve a greater change in momentum.

68. Explaing the behaviour of gases When we heat a gas a constant pressure, the volume increases. Why? Increasing the volume reduces the chance of particles colliding with the container walls, opposing the effect of the particles increased kinetic energy.

69. Explaing the behaviour of gases When we compress (reduce the volume) a gas at constant temperature, the pressure increases. Why? A smaller volume increases the likelihood of a particle colliding with the container walls.

70. Explaing the behaviour of gases In this way we are explaining the macroscopic behaviour of a gas (the quantities that can be measured like temperature, pressure and volume) by looking at its microscopic behaviour (how the individual particles move)