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Thermal

Thermal

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  • 1. Thermal Physics Topic 3.1 Thermal Concepts
  • 2. Temperature
    • At a macroscopic level, temperature is the degree of hotness or coldness of a body as measured by a thermometer
    • Temperature is a property that determines the direction of thermal energy transfer between two bodies in contact
    • Temperature is measured in degrees Celsius ( o C) or Kelvin (K)
      • Where Temp in K = Temp in o C + 273(.15)
      • Temp in K is known as the absolute temperature
  • 3. Thermal Equilibrium
    • When 2 bodies are placed in contact
    • Heat will flow from the warmer body to the colder body
    • Until the two objects reach the same temperature
    • They will then be in Thermal Equilibrium
    • This is how a thermometer works
  • 4. Use the idea of thermal equilibrium to explain
    • Why kitchen floors are generally cold to bare feet during winter
    • Why hot water bottles are used in winter
    • Why carpet feels warmer than wood
    • Why do car windows fog in winter
  • 5. Thermometers
    • A temperature scale is constructed by taking two fixed, reproducible temperatures
    • The upper fixed point is the boiling point of pure water at atmospheric pressure
    • The lower fixed point is the melting point of pure ice at atmospheric pressure
  • 6.
    • These were then given the values of 100 o C and 0 o C respectively, and the scale between them was divided by 100 to give individual degrees
    • A thermometer uses the expansion and contraction of a liquid ;
      • usually mercury or alcohol ,
      • in a glass capillary tube with a scale to ,
      • measure the expansion or contraction.
  • 7. Temperature - Microscopic
    • At a microscopic level, temperature is regarded as a measure of the average kinetic energy per molecule associated with its movement in the substance
  • 8. Internal Energy
    • The Internal (thermal) energy of a body is the total energy associated with the thermal motions of the particles
    • It can comprise of both kinetic and potential energies associated with particle motion
    • Kinetic energy arises from the translational and rotational motions
    • Potential energy arises from the forces between the molecules
  • 9. Heat
    • The term heat represents energy transfer due to a temperature difference
    • Occurs from higher to lower temperature regions
  • 10. Heat
    • Most people tend to believe that all matter ;
      • contains heat.
    • All matter contains a number of forms of energy ;
      • but not heat.
    • Heat is the transfer of energy from a body with higher temperature to ;
      • one of lower temperature.
  • 11. Methods of Heat Transfer
    • Heat can be transferred from one body to another by
      • Conduction
      • Convection
      • Radiation
  • 12. Conduction
    • Conduction can take place ;
      • within materials and ,
      • between different materials ,
      • that are in direct contact.
    • It can be explained in terms of collisions ;
      • between atoms or molecules and ,
      • the actions of loosely bound electrons.
    • How does a metal rod get hot when only part is held in a flame?
  • 13. Convection
    • This involves the transfer of energy ;
      • from one molecule to another ,
      • with the molecules moving.
    • Another way in which convection works is ;
      • for the hotter material to move up such as hot air.
    • Convection occurs in all fluids ;
      • w h ether they are liquids or gases.
    • Hot air above a stove becomes less dense ;
      • and moves up toward the ceiling…
    • Explain how it is possible to fly in hot air balloon
  • 14. Radiation
    • All energy, including heat ;
      • which is transferred by radiation,
      • is called radiant energy .
    • Higher temperature objects emit ;
      • shorter wavelengths
      • and vice - versa.
    • The coolest objects that emit light ;
      • we can see (red) ,
      • are at a temperature of 500 o C.
  • 15. Thermal Properties of Gases
    • Investigations involved the measurement of
      • Pressure
      • Volume
      • Temperature
    • These experiments used these macroscopic properties of a gas to formulate a number of gas laws
  • 16. Units
    • Temperature is always measured in K
    • Volume is usually in m 3
    • Pressure can be different units as long as you are consistent
    • But 1 atm = 1.01 x 10 5 Nm -2 = 101.3 kPa = 760 mmHg
  • 17. Gas Laws
    • 1. Boyle’s Law :
    • Variation of volume with pressure (const T).
    • eg Bike pump .
  • 18. Gas Laws
    • Boyle’s experiment can be repeated ;
      • by using a flexible J-tube containing mercury ,
      • and a barometer to measure atmospheric pressure.
    • Air was trapped in one end ;
      • and varied the pressure on it ,
      • by raising the other end of the tube as shown below:
  • 19. Gas Laws
  • 20. Gas Laws
    • The gas pressure of the enclosed air ;
      • is equal to the atmospheric pressure plus ,
      • the pressure resulting from ,
      • the height difference of Hg in the two arms.
    • Taking a number of pressure and volume readings ;
      • the graph looks like:
  • 21. Gas Laws
  • 22. Gas Laws
    • A more revealing graph is obtained by plotting p vs 1/V
  • 23. Gas Laws
    • As a straight line is obtained ;
      • that passes through the origin,
      • we can say mathematically:
    • p  1/V
  • 24. Gas Laws
    • If the same experiment is performed ;
      • at different temperatures,
      • a curve is still obtained but at different positions.
  • 25. Gas Laws
  • 26. Gas Laws
    • The curves are called isotherms ;
      • as the temp at each point on the curve is the same.
  • 27. Gas Laws
    • We have assumed ;
      • that there is a fixed amount of gas ,
      • maintained at a constant temperature.
    • We can also write, p = k /V or pV = k
  • 28. Gas Laws
    • k is a constant of proportionality ;
      • and is equal to the slope.
    • For different points on the graph ;
      • p 1 V 1 = k and p 2 V 2 = k
    •  p 1 V 1 = p 2 V 2 (for const. temp)
  • 29. Gas Laws
    • Boyle’s law states ;
      • at constant temperature,
      • the volume of a fixed mass of gas ,
      • varies inversely with its pressure.
  • 30. Gas Laws
    • 2. Charles’ Law :
    • Variation of volume with temperature (const p).
    • To verify this ;
      • a capillary tube is sealed at one end ,
      • and a small mercury piston is included,
      • trapping a sample of air.
  • 31. Gas Laws
    • A thermometer is strapped to the tube ;
      • and both are placed in glycerol.
    • The pressure is fixed ;
      • and is equal to atmospheric pressure.
  • 32. Gas Laws
    • The glycerol is heated ;
      • and the length of the air column is noted ,
      • at various temperatures.
    • The length of the column can be read ;
      • as the volume in arbitrary units ,
      • as the cross section is uniform.
  • 33. Gas Laws
  • 34. Gas Laws
    • The results obtained can be graphed as below:
  • 35. Gas Laws
    • This graph does not go through the origin.
    • When extrapolated ;
      • it cuts the temp axis at ,
      • -273 o C
  • 36. Gas Laws
  • 37. Gas Laws
    • This creates a new scale defined as ;
      • The absolute temperature scale, where
      • zero degrees kelvin (0K) = -273 o C.
    • A 1K rise = 1 o C rise.
    • A real gas would liquefy before it reaches 0K ;
      • and so no gas is ideal.
  • 38. Gas Laws
    • Any extrapolation should not continue ;
      • past the liquefaction point.
    • No gas can reach a temperature below 0K ;
      • as no gas could have a negative volume.
    • Graphing for two different pressures:
  • 39. Gas Laws
  • 40. Gas Laws
    • Using the Kelvin scale ;
      • we can say mathematically:
    • V  absolute T
    • V = k T or k = V/T
      • where k is the gradient of the graph.
    • For two points, (V 1 ,T 1 ) and (V 2 ,T 2 )
    • V 1 /T 1 = V 2 /T 2
  • 41. Gas Laws
    • Charles’ law states ;
      • at constant pressure,
      • the volume of a fixed mass of gas ,
      • varies directly with its absolute temperature.
  • 42. Gas Laws
    • 3. Guy-Lussac’s Law : (Law of pressures)  
    • Variation of pressure with temperature (const V).
  • 43. Gas Laws
    • A bulb enclosing a fixed volume of gas ;
      • is placed in a beaker of water ,
      • and attached to a mercury manometer ,
      • as shown below:
  • 44. Gas Laws
  • 45. Gas Laws
    • The water is warmed ;
      • the air expands in the bulb and tube ,
      • and causes the Hg ,
      • to be forced up the manometer.
  • 46. Gas Laws
    • One arm is then raised ;
      • to increase pressure ,
      • and compress the air ,
      • back to the original volume.
    • The height difference is recorded.
  • 47. Gas Laws
  • 48. Gas Laws
    • As the graph does not pass through the origin ;
      • it must be extrapolated.
  • 49. Gas Laws
    • We can now mathematically say:
    •   p  absolute T
    •   p = k T or k = p/T;
      • where k is the gradient of the graph.
    • Using two points on the graph;
    • p 1 /T 1 = p 2 /T 2
  • 50. Gas Laws
    • The law of pressures state ;
      • at a constant volume,
      • the pressure of a given mass of gas ,
      • varies directly with the Kelvin temperature.
  • 51. Gas Laws
    • GENERAL GAS EQUATION
    • All three laws studied above ;
      • can be combined into one.
    • From Boyle’s law:
  • 52. Gas Laws
  • 53. Gas Laws
    • The curves plotted from the above situations are shown:
  • 54. Gas Laws
    • Going from (1) to (2)
    • V  by raising T at constant pressure.
    • V 1 /T 1 = V * /T 2 (Charles’ Law) 
    • Going from (2) to (3) temp constant ;
      • piston is withdrawn causing V  & P  .
    • p 1 V * = p 2 V 2 (Boyle’s Law) 
    • Multiplying  and 
  • 55. Gas Laws
    • (V 1 x p 1 V * )/T 1 = (V * x p 2 V 2 )/T 2
    • (p 1 V 1 )/T 1 = (p 2 V 2 )/T 2
    • pV/T = constant
    • This is called the ideal gas law ;
      • and is accurate except ,
      • at low temperatures ,
      • and high pressures.
  • 56. Gas Laws
    • The only gas ;
      • that approximates the ideal gas law ,
      • over a wide range of conditions ,
      • is Helium,
      • to 4.2K.
  • 57. Gas Laws
    • p 1 V 1 /T 1 = C
    • C depends on the number of moles, n , and ;
      • Universal gas constant, R
      • (8.315 J mol -1 K -1 )
    • For one mole of gas, pV = RT
    • For n moles of gas, pV = nRT
  • 58. Gas Laws
    • One mole of a substance contains ;
      • 6.023 x 10 23 particles.
    • This value is called Avogadro’s number ;
      • and is given the symbol N A .
    • To find the moles present ( n ) ;
      • divide the total number of molecules N ,
      • by Avogadro’s number.
  • 59. Gas Laws
    • n = N/N A
    • m grams of a substance contains ;
      • m/m A moles,
      • where the molar mass is expressed in grams.
    • n = m/m A
  • 60. The Mole
    • The mole is the amount of substance which contains the same number of elementary particles as there are in 12 grams of carbon-12
    • Experiments show that this is
    • 6.02 x 10 23 particle s
    • A value denoted by N A and called the Avogadro Constant (units mol -1 )
  • 61. Molar Mass
    • Molar mass is the mass of one mole of the substance
    • SI units are kg mol -1
  • 62. Example
    • Molar mass of Oxygen gas is
    • 32 x10 -3 kg mol -1
    • If I have 20g of Oxygen, how many moles do I have and how many molecules?
    • 20 x 10 -3 kg / 32 x10 -3 kg mol -1
    •  0.625 mol
    •  0.625 mol x 6.02 x 10 23 molecules
    •  3.7625 x 10 23 molecules
  • 63. Use your periodic table to…
    • Find out how many atoms are in
    • 46g of sodium
    • 64mg of copper
    • 180g of water
    • Find out how many moles is
    • 10g of Argon
    • 1kg of Iron
  • 64. What is the molar mass of…
    • Nitrogen
    • Chlorine
    • Neon
    • Water
    • Sulfuric acid (H 2 SO 4 )
    • Calcium carbonate (CaCO 3 )