Thermal Physics

1. Thermal Physics IB SL-2

2. Thermal Physics • Import to understand both the microscopic and macroscopic view of matter and related characteristics. • Microscopic – at the molecular level understand activity. • Macroscopic – what can be seen, can be described and observed, REASON for properties can be explained by the microscopic details.

3. Matter • Matter – Takes up space (has a volume) and has mass. • Matter is composed of particles. • Molecule – smallest chemical unit of a substance that is capable of stable, independent existence. • Atom – smallest unit of an element that can exist either alone or in combination with other atoms of the same or different elements.

4. Kinetic Energy & Matter • Kinetic Theory of Matter – developed to explain the motions of molecules and the energy molecules possess. • Molecules of a substance are in constant motion. Amount of motion depends on average kinetic energy (KE) of molecules. Energy dependent on Temperature. • Collisions between molecules are perfectly elastic (except when chemical changes or molecular excitation).

5. Kinetic Energy & Matter • Van der Waalis forces - Force that holds molecules together between molecules. • The force between molecules vary based on the distance between molecules.

6. 4 States of Matter • Solid • Liquids • Gases • Plasma • Note: fluids (Latin for “to flow”) refers to both gas, liquids and plasma • Vapor = term used for gases

7. Solid State • Solid – state of matter with a fixed volume and shape. • Two types of solids: crystalline and amorphous. • Crystalline – have a regular arrangement of particles, they are fixed in a crystal lattice. • Amorphous – have random particle arrangement, so no fixed lattice structure. But have definite volume. Ex. Butter, paraffin & glass.

8. Solid State • Most are Crystalline solids, where: • Binding forces hold particles in relatively fixed position in lattice (see your text pg 146, Fig 7-4), but molecules still vibrate about position. • Particles close together • KE varies with temp, higher Temp yield higher energy. Lowest KE of all properties of matter. • Binding force of molecules in a solid are much greater than thermal KE.

9. Liquid State (phase) • Liquid – material with fixed volume, but whose shape depends on container. • Particles are almost as close together as solids. So increased distance of separation. This means that binding energy is less than solid.

10. Liquid State (phase) • Molecules not held in a fixed position. So no definite shape of material. • Vibrational energy greater than solid, • Binding force not as great as in solid. So liquids can be separated into drops.

11. Gas State (phase) • Gas – state of matter that expands to fill container. • Average separation of particles is relatively huge. • Particles not held in fixed position. Molecules spontaneously separate.

12. Gas State (phase) • Gas occupies about 1000 times the volume of liquid for the same mass. • KE is the much greater than solid or liquids. KE enables molecules to stay separated. Easily overcomes binding energy of molecules.

13. Solid, Liquid, Gas • Both Liquids and solids are not easily compressed, due to repelling nature of molecules. • However, gasses can be compressed. • Most of matter on earth is either solid, liquid or gas.

14. Plasma • Plasma – state of matter in which atoms are separated into electrons and positive ions or bare nuclei. • So much KE that collisions between particles cause them to tear each other apart. • Electron pulled off of atoms, producing positively charged ions.

15. Plasma • Temperature is extremely high, so High KE • In fact plasma can be defined as – gas heated to an extreme temperature so that it is capable of conducting an electric charge. • Stars and our sun made of plasma. This means that the universe is primarily composed of Plasma, but very little plasma on earth.

16. Temperature • Temperature – hotness (or coldness) measured on some definite scale like a thermometer. [MACRO View] • Temperature – is proportional to the average kinetic energy of the particles. [MICRO View]

17. Temperature • Thermal equilibrium – the state in which two bodies in physical contact with each other have identical temperatures. • Fahrenheit (TF), Celsius (C) , Kelvin (K) are three temperature scales. • We will work in Celsius unless noted, must convert from fahrenheit.

18. Temperature • TF = (9/5 TC )+ 32.0 • T = TC + 273.15 • TF = Fahrenheit temp in degrees • TC = Celsius temp in degrees • T = Kelvin Temp

19. Internal Energy • Internal energy – the energy of a substance due to the random motions of its component particles and equal to the total energy of those particles. • Internal energy of a substance is proportional to the temperature of the substance. • Internal energy – consists of KE and PE of particles. Energy can be in the form of translation, rotation, and vibration (Holt pg. 359 table 10-1)

20. Heat • Heat – the process by which energy is exchanged between objects because of a difference in their temperatures. • Macroscopic View – energy transferred by heat always moves from an object with high Temp to object with lower Temp. • Consider Case or warm Cokes put into a tub of ice water. Cokes get cold and ice water gets warm.

21. Heat • Microscopic view – explains energy transfer • Coke molecules have a higher KE than ice water. • Energy transferred from coke to can by collisions of atoms. This transfers energy to metal can atoms, and by same method to ice water atoms. • Water molecules increase in energy while coke atoms decrease, until average KE are same!

22. Heat • Heat has a SI Unit of joule (J) the same as unit for energy. • J = 1 kg (m2/s2) • Holt pg 367 give summary of other thermal units.

23. Energy Transfer • Total energy is conserved when energy is transferred. • Important to remember this on Micro scale • PE + KE +U = 0 •  = change in • PE = potential energy; KE = kinetic energy • U = change in internal energy

24. Heat Capacity • Heat capacity – the quantity of heat needed to raise a body’s temperature 1o C • Heat capacity = Q/T • Q = quantity of heat needed (J, joules) • T = change in temperature • Heat Capacity SI units J/oC

25. Specific Heat Capacity • Specific Heat Capacity (or just “Specific Heat”) – the quantity of energy needed to raise the temperature of 1 kg of a substance by 1o C at constant pressure. • Specific heat capacity - symbol is (cp) where c stands for capacity and “p” indicates specific heat capacity is measured at constant temperature.

26. Specific Heat Capacity • Formula : • Q = energy transferred • m = mass; T = change in temperature • cp = specific heat capacity

27. Specific Heat Capacity • The specific heat equation works for both substances that absorb energy from surroundings and transfer energy to surroundings. • Specific Heat capacity is a property of a substance, and the same mass of two differences substances will have different heat capacity.

28. Problem • How much heat is required to raise the temperature of an empty 20-kg vat made of iron (specific heat 450 J/kg Co) from 10oC to 90oC? • What if the vat is filled with 20 kg of water?

29. Solution • See page 421 in Giancoli • Read section 14-4 • Use Specific Heats on Table 14-1

30. Specific Heat Capacity • Specific Heat of water (cp,w ) is well known. So provides good way to understand heat capacity of certain materials • cp,w = 4.186 kJ/kg* oC) • Qw = energy absorbed by water • Qx = energy released by substance

31. Specific Heat Capacity • A modification of the specific heat capacity formula will allow any objects specific capacity to be determine if placed in water. • Where: w = water and x = other substance

32. Specific Heat Capacity • The previously mentioned procedure is called Calorimetry. Holt pg. 372, Giancoli pg 423. • Calorimetry – an experimental procedure used to measure the energy transferred by heat from one substance to another. • MUST WORK IN Celsius

33. Specific Heat Capacity - Difference • Why do different substance have different specific heat capacities? Go MICRO! • Different substances have different # of molecules. • So, if same internal energy is added to two different substances of same mass, the internal energy is distributed over different # of molecules.

34. Specific Heat Capacity –Difference Continued • Thus, average energy change is different for molecules of 2 different substances. Since the same amount of energy is spread equally with different number of molecules. • Therefore, the temperature change and specific heat capacity will be different for the two substances.

35. Calorimetry • Remember • Heat lost = Heat gained • So all energy is conserved

36. Example problems • Calorimetry • Page 422 Giancoli • Ex 14-5 & 14-6 be able to do.

37. Assignment • Giancoli page 439 #9-21

38. Specific Heat Capacity • Assignment • Holt pg 373-4 Sample Problem 10C • Practice 10C pg. 374 #1-5

39. Phase Change • Phase change – the physical change of a substance from one state (solid, liquid, or gas) to another at constant temperature and pressure. • Heat is the process by which energy is exchanged, either between two bodies at different temperatures or between two bodies at the same temperature when one of them is undergoing a phase change.

40. Latent Heat • Latent heat – the energy per unit mass that is transferred during a phase change of a substance.

41. Latent Heat • Latent Heat • Q = energy transferred by heat during a phase change • m = mass • L = latent heat

42. Phase Solid/Liquid • Melting – Change in phase from solid to liquid. Temperature at which this occurs is called the melting point. • Freezing point – when substance changes from liquid to solid.

43. Phase Solid/Liquid • For crystalline solids – each solid has a definite melting point and a definite freezing point. • Same Temperature of melting/freezing point – at any given pressure for each particular crystalline solid. • Noncrystalline solids – have no definite melting point. Rather when heated they soften gradually. Think Candle wax! • We will focus our study on crystalline solids.

44. Phase Solid/Liquid • To melt a solid energy must be supplied. • The energy supplied to solid, increases the energy of the solid particles. • As the temperature increases, the vibrations of particles increase in amplitude; so more PE is stored in the average stretching of the bonds between the particles. (Increase in Temp yields increased intermolecular PE)

45. Phase Solid/Liquid • Melting of a solid occurs when temperature is reached, where the bonds between the particles cannot absorb any more energy without breaking. • Freezing occurs when temperature is lowered to the point, where the molecular forces between the particles are now able to draw the particles into fixed positions.

46. Phase Solid/Liquid • At the melting point. • All energy supplied to a substance during melting is used to increase PE of the particles. • KE of the particles does not change, since average KE depends on temperature. The Temp is unchanged during the melting process once temp is reached!!

47. Phase Solid/Liquid • At the freezing point • Since Temp is lowered both PE and KE are lost. • Loss of KE means particles vibration is reduced because at lower temp. • Loss of PE allows particles to be draw together (since now binding forces are now greater than PE). • Water exception: ice expands

48. Gas Phase Changes • Vaporization – production of a vapor or gas from matter in another phase. • Evaporation – when vaporization occurs from a liquid. • Sublimation – direct change (vaporization) from solid to gas.

49. Gas Phase Changes • Vaporization just like melting is a constant- temperature process. • At vaporization point • All energy during vaporization goes to increase PE of particles. • Average KE of particles is unchanged. • KE of particles is increased as temperature increased until Temp is reached for vaporization.

50. Phase Change Overview • Refer to Holt pg. 376 or pg. 190 & 196 in Trinklein for heating curve. • Giancoli Pg 425 • Shows Temperature change of ice (Y axis for Temp) as energy is added (X axis for Heat in J).