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Chapter 1.4.2: Temperature in Thermal Systems

Chapter 1.4.2: Temperature in Thermal Systems. Objectives: Define specific heat, heat of fusion, and heat of vaporization Use specific heat, heat of fusion, and heat of vaporization to solve problems involving heat transfer. Specific Heat. Units for Thermal Energy and Heat Joule (J)

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Chapter 1.4.2: Temperature in Thermal Systems

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  1. Chapter 1.4.2: Temperature in Thermal Systems • Objectives: • Define specific heat, heat of fusion, and heat of vaporization • Use specific heat, heat of fusion, and heat of vaporization to solve problems involving heat transfer.

  2. Specific Heat • Units for Thermal Energy and Heat • Joule (J) • calorie (cal) • British thermal unit (Btu) • Specific heat, C – amount of energy required to raise the temperature of a unit mass of a substance one temperature unit. • Possible units? • J/g·oC cal/g·oC kJ/kg·oC

  3. C is an intrinsic physical property, like  does not depend on amount but only on the substance itself.

  4. Table 1.7 Specific Heat of Common Substances

  5. Sample Problem • How much energy must be absorbed by 20.0 g of water to increase its temperature from 83.0 °C to 94.0 °C?

  6. Sample Problem • The specific heat of iron is 0.16 cal/g·oC. Convert this value to specific heat in J/g·oC.

  7. Example 1.15, p. 72 • A teakettle hold 0.5 liters of water. How much heat is needed to increase the temperature from 20oC to 100oC. Hints: water = 1.00 g/ml; and 0.5L = 500 mL

  8. Change of State • Linear relationship between heat transfer and temperature does NOT hold during a change of state • Temperature stays constant during the phase change, i.e. T = 0 during the phase change

  9. Boiling point Melting point/freezing point Example for Water Steam Water and Steam Water Ice and water Ice

  10. Melting point (m.p.) – temperature at which a substance melts (or freezes if losing energy). • Boiling point (b.p.) – temperature at which a substance turns to gas (or condenses if losing energy). • Heat of fusion, Hf – amount of energy required to melt one gram of solid. • Heat of vaporization, Hv – amount of energy required to vaporize one gram of a liquid.

  11. Table 1.8 Heat of Fusion (Hf) and Vaporization (Hv) of Selected Substances Back to Problem

  12. Sample Problems • Convert the heat of fusion of iron to heat of fusion in kJ/kg. • How much heat is required to vaporize 0.5 kg of gold.

  13. Example 1.17 Melting Ice and Warming Water • A 10.0 g ice cube has a temperature of -5.0 oC. How much heat is needed to melt the ice cube and warm the resulting water to room temperature (20 oC)?

  14. Links • Virtual experiment: Heating Curves • http://www.harcourtschool.com/activity/hotplate/index.html • Heating Curve Tutorial • http://www.kentchemistry.com/links/Matter/HeatingCurve.htm • Heat Problems • http://dbhs.wvusd.k12.ca.us/webdocs/Thermochem/Thermochem-WS1.html

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