Chapter 10

1. Chapter 10 Temperature

2. Temperature vs. Heat • Thermal (internal) energy • sum total of the kinetic energies of all molecules of a substance • Units of Joules • Temperature • Average KE per molecule of a substance • Heat • Net energy transferred from one object to another due to a temperature difference

3. Temperature Scales • Fahrenheit • Water freezing point = 32F • Water boiling point = 212F • Celcius (Centigrade) • Water freezing point = 0C • Water boiling point = 100C

4. Temperature Scales • Kelvin • Standard metric unit of temperature • Expressed without degree symbol  • 0C = 273 K • Based on making zero the lowest temperature on the scale • Obtained experimentally using P vs. T graph

5. Temperature measurement • Thermal expansion • Change in the length, area, or volume of an object due to temperature changes • Most materials expand with increasing temperature, decrease with decreasing temperature • Contraction is negative expansion • Expansion rates vary by material

6. Temperature measurement • Liquid Thermometers • Mercury and Alcohol (dyed red) most common • High expansion rates • Remain liquids at normal temperatures • Knowing the expansion rates allows calibration

7. Temperature measurement • Bimetallic strips • Two strips of different metals bonded together • When heated, different materials expand at different rates • Metal strip will bend • Used in radial (dial) thermometers and most household thermostats

8. Thermal Expansion--solids • Linear Expansion •  = coefficient of linear expansion • L0 = original length • Area Expansion • Volume Expansion p. 344 common values of 

9. Gas Laws • Boyle’s Law • Temperature is held constant • As Pressure increases, Volume decreases • Charles’ Law • Pressure is held constant • As Temp. increases, Volume Increases • Absolute Temperature (Kelvins)

10. Gas Laws • Ideal Gas Law (ratio form) • Amount of gas remains constant • Applies to all density gases • Fairly accurate with higher density gases

11. Gas Laws • Ideal Gas Law (Physics form) • N = # of molecules of gas • kB = 1.38 x 10-23 J/K • Boltzmann’s constant

12. Gas Laws • Ideal Gas Law (Chemistry form) • n = # of moles of gas • 1 mol = 6.02 x 1023 molecules • 1 mol = molecular mass expressed in grams • R = 8.31 J/(mol*K) • Universal gas constant

13. Gas Laws • Standard Temperature and Pressure (STP) • T = 0C = 273K • P = 1 atm = 101.3 kPa • At STP, all gases occupy 22.4 L • 1 L = 1 dm3

14. Absolute Zero • Graph Pressure vs. Temperature for ideal gases • Y = pressure (Pa), X = Temp. (°C) • Amount and Volume of Gas kept constant • Measure Pressure at several temperatures • Extend line backwards until P = 0 • x-intercept • When P = 0, Absolute Temp. = 0

15. Absolute Zero • Kelvin Scale • °C + 273 • Standard Metric unit of absolute temperature • Rankine Scale • °F + 460 • Engineering applications in USA

16. Kinetic Theory • Courtesy of Daniel Bernoulli • Relation of Absolute Temperature to molecular Kinetic Energy • Molecules of a gas can be viewed as colliding particles • Ignore Rotational and Vibrational motion • Temperature and Pressure depend only on translational motion

17. Kinetic Theory • Temperature is the average KE of individual molecules of the gas • Higher temperatures correspond to more vigorous collisions • Pressure is caused by collisions with outside of the container • As number and strength of collisions increase, so does pressure

18. Kinetic Theory • vrms = root-mean-square velocity • Average speed of individual molecules • kB = 1.38 x 10-23 J/K • Boltzmann’s constant • T = absolute temperature (K) • Equation is derived from the force on the side of a container • Force caused by impulse during collisions Average speed of He, H2, N2 Applet

19. Kinetic Theory • U = total internal energy of an ideal monatomic gas • Monotomic—molecules consisting of 1 atom • He • Does not apply to larger molecules • Diatomic—molecules consisting of 2 atoms • H2, O2, N2, etc.