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Thermochemistry

Thermochemistry. 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. Lots of energy stored in bonds!. Energy difference.

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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. What do Calories mean in food? • 10 grams of sugar has 41 Calories. • When 10 grams of sugar are burned, 41 kcal (170 kJ) of energy are released. • Your body “burns” food for energy. • In order to use the energy available in 10 grams of sugar, you must do 41 kcal worth of work.

  9. 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.

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

  11. 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.)

  12. Some Specific Heats

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

  14. 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?

  15. 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

  16. 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

  17. 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

  18. 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

  19. An Exothermic Reaction • CaO(s) + H2O(l)  Ca(OH)2(s) + 62.5 kJ CaO(s) + H2O(l) ΔH = -65.2 kJ Ca(OH)2(s)

  20. Na2CO3(s) + H2O(g) + CO2(g) ΔH = +129 kJ 2NaHCO3(s) Thermochemical Equations 2NaHCO3(s) + 129 kJ  Na2CO3(s) + H2O(g) + CO2(g)

  21. Thermochemical Equations and Stoichiometry • You can use thermochemical equations in stoichiometry. • How much heat energy is produced when 55.0 grams of ethanol is burned completely? • C2H5OH(l) + 3O2(g)  2CO2(g) + 3H2O(g) + 1300. kJ • Given: 55.0 g C2H5OH • Want: kJ • Conversion factors: • 1 mol C2H5OH produces 1300. kJ when burned • 1 mol C2H5OH = 46.07 g/mol 55.0 g C2H5OH = 1550 kJ

  22. Thermochemical Equations and Stoichiometry • 0.500 grams of methane gas are burned completely beneath a container that holds 100. grams of water, originally at 20.0º. If all of the heat from the combustion reaction goes into the water, what will the water’s final temperature be? • CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) + 803 kJ • First find out how much total heat is released. • Given: 0.500 g CH4(g) • Want: kJ • Conversion factors: • 1 mol CH4 = 16.05 g CH4 • 1 mol CH4 produces 803 kJ when completely burned 0.500 g CH4 = 25.0 kJ

  23. Thermochemical Equations and Stoichiometry • The combustion of 5.00 grams of methane releases 250. kJ of heat. • Now we’ll calculate how hot the water in the container will get if it absorbs all of the heat. • First convert 25.0 kJ to J • q = m x c x T • 2.50x104 J = (100. g) (4.18 J/gºC) T • 2.50x104 J = (418 J/ºC) T • T = 59.8ºC • The water will get 59.8ºC warmer. • The final temperature will be 20.0ºC + 59.8ºC = 79.8ºC. 25.0 kJ = 2.50x104 J

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