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Thermochemistry

Learn about thermochemistry, energy changes during phase changes and chemical reactions, heat flow, heat capacity, specific heat, and enthalpy changes.

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Thermochemistry

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  1. Thermochemistry

  2. Definitions • Energy – capacity for doing work or supplying heat. • Thermochemistry – study of energy changes that occur during phase changes and chem. rxns. • Chem. Potential Energy – energy stored in chemical bonds.

  3. Lots of energy stored in bonds! Energy difference. 5473 kJ/mol Little energy stored in bonds. Example

  4. Heat • Represented by q. • Energy that transfers from one object to another because of a Temp. difference between them. • Heat flows from warm  cool until the two objects are at the same Temp.

  5. Exothermic vs. Endothermic • In exothermic processes, the system loses heat as its surroundings warm up. • q has a negative value b/c the system is losing heat. • In endothermic processes, the system gains heat as its surroundings cool down. • q has a positive value b/c the system is gaining heat.

  6. Water Potential Energy  Ice Time  Potential Energy Diagram of Ice Melting at 0ºC. Is the melting of ice an endothermic or an exothermic process? How can you tell?

  7. Measuring Heat Flow • SI Unit of heat flow: Joule (J) • Common unit used in chemistry: calorie (cal) • Amt. of heat needed to raise 1 gram of water by 1ºC. • 1 cal = 4.184 J • Food Calorie (capital “C”) = 1000 cal, or 1 kilocalorie = 4184 J

  8. Heat Capacity • Amount of heat needed to raise an object’s temperature by 1°C. • Depends on the chemical composition and the mass of the object. • EXAMPLE: 1 gram of water requires 1 cal to raise its temperature by 1°C. • 100. g of water require 100. cal to raise the temp. by 1°C.

  9. Same temperature change 10 g H2O 1 g H2O Heat Capacity

  10. Specific Heat (c) • Amt. of heat needed to raise 1 gram of a substance’s temperature by 1ºC. • Expressed in J/g ºC, or cal/g ºC • The higher a substance’s specific heat, the more energy it takes to heat it. • Substance’s with low specific heats heat up and cool down quickly (most metals, e.g.)

  11. Some Specific Heats

  12. Specific Heat (c) • c = heat / (mass x change in Temp.) • c = q / (m x ΔT) • q = m x c x ΔT

  13. Example Problem • The temperature of a 95.4-g piece of Cu increases from 25.0ºC to 48.0ºC when the Cu absorbs 849 J of heat. What is the specific heat of Cu? • SOLUTION: q = m x c x ΔT • 849 J = (95.4 g) c (48.0ºC – 25.0ºC) • 849 J = (95.4 g) c (23.0ºC) • 849 J = (2190 gºC) c • c = 0.388 J/gºC • Based on what you know about metals, does this answer make sense?

  14. Example Problem • When 435 J of heat is added to 3.4 g of olive oil at 21ºC, the temperature increases to 85ºC. What is the specific heat of olive oil? • SOLUTION: q = m x c x ΔT • 435 J = (3.4 g) c (85ºC – 21ºC) • 435 J = (3.4 g) c (64ºC) • 435 J = (220 gºC) c • c = 2.0 J/gºC

  15. Example Problem • How much heat is required to raise the temperature of 250.0 g of mercury by 52ºC? The specific heat of mercury is 0.14 J/gºC. • SOLUTION: q = m x c x ΔT • q = (250.0 g)(0.14 J/gºC)(52ºC) • q = 1800 J = 1.8 kJ

  16. Enthalpy Changes • Enthalpy (H) – the heat content of a system at constant pressure. • Enthalpy change (ΔH) – the heat that enters or leaves a system at constant pressure. • q = ΔH • Neg. ΔH = exothermic process • Pos. ΔH = endothermic process

  17. Thermochemical Equations • Enthalpy change can be written as a reactant or a product. • Reactant  endothermic • Product  exothermic • Example: The reaction of calcium oxide with water is exothermic. • It produces 65.2 kJ of heat per mole of CaO reacted. • CaO(s) + H2O(l)  Ca(OH)2(s) + 62.5 kJ

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