2. CHAPTER 2  Heat  Temperature and Conversions  Specific Heat

3. What is Energy?  The ability to move or change matter. (Units: Joules)  All physical and chemical changes involve energy!

4. Examples of Energy  Kinetic – energy of motion  KE = ½ mv 2  Potential – stored energy/energy of position  Light  Sound  Electricity  Heat (Thermal)  Chemical

5. Law of conservation of energy:  Energy cannot be created or destroyed during any chemical or physical change. Energy may be transferred between the system and surroundings Energy may change forms.

6. Energy and mass are related  Einstein derived an equation to show this relationship in 1905.  Nuclear reactions can create energy from mass.

7. Energy is transferred during physical and chemical changes:  Endothermic – energy is absorbed by the system +  Exothermic – energy is released into the surroundings -

8. What is Heat?  The transfer of energy between the particles of two objects due to a temperature difference between the two objects.  Heat always flows from hot to cold.  Measured in a calorimeter.  Units: Joules, Calories, or calories.

9. TEMPERATURE What is temperature?  Temperature is the measure of the average kinetic energy of all the particles within an object.  Measured with a thermometer.

10. Heat and temperature  The transfer of heat does not always result in a temperature increase. During phase changes, energy goes directly to changing the phase, not into increasing the kinetic energy of the particles.  EX. The heating curve for water.

11. The heating curve for water shows that temperature does NOT change during a phase change.

12. Heating curve points and definitions:  Melting point/ freezing point of water: 0º C  Boiling point of water: 100 º C  Heat of fusion – the amount of energy required to melt a solid  Heat of crystallization – the amount of energy released when a solid forms from a liquid  Heat of vaporization – the amount of energy required to change a liquid into a gas.

13. Scales to Measure Temperature  Fahrenheit Scale (U.S.A.)  Celsius Scale (everyone else)  Kelvin Scale (scientists)

14. How do Thermometers Work?  Usually contain alcohol or mercury.  Temperature increase (particles move faster), liquids expand  Temperature decreases (particles move slower), liquids contract

15. Absolute Zero  The lowest possible temperature  All motion STOPS.  Energy is minimal/absent.  In September 2003, MIT announced a record cold temperature of 450 pK, or 4.5 × 10-10 K in a Bose-Einstein condensate of sodium atoms. This was performed by Wolfgang Ketterle and colleagues at MIT.September 2003MITWolfgang Ketterle

16. SPECIFIC HEAT CAPACITY  Transfer of heat affects substances differently.  Measuring heat transferred to and absorbed by a substance under conditions of constant pressure yields specific heat capacity.

17. SPECIFIC HEAT CAPACITY Specific heat is defined as: The quantity of heat required to raise 1 gram of a substance 1°C or 1 K. Symbol: Cp The p symbolizes that the measurements were taken under constant pressure. Units = Joules/ gram °C or J/gK J/g°C

18. Sample Cp values  Metals have low specific heat values which allows them to heat up with little added energy.  Iron 0.449 J/g°C  Copper 0.385 J/g°C  Platinum 0.133 J/g°C  Water has a relatively high specific heat 4.184 J/g °C

19. Questions:  Which would heat up faster, 5.00 grams of iron or 5.00 grams of water?  Which would cool down faster, 5.00 grams of iron or 5.00 grams of water?  Which is a better thermal conductor?  Which is a better insulator?

20. MEASURING HEAT and SPECIFIC HEAT Must use a calorimeter. Find the change in temperature:  T = (delta T) change in temperature in °C  T = T final – T initial

21. SPECIFIC HEAT CALCULATIONS q =m x Cp x  T

22. Rearrange the formula: m= q/Cp  T Cp = q/ m  T  T = q/ m Cp