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Thermochemistry

This comprehensive overview of thermochemistry explores the concepts of heat transfer during chemical reactions. We define energy in terms of its various types, including mechanical and chemical potential energy. The key distinction between exothermic and endothermic processes is explained, showcasing how heat flows out of or into systems. Additionally, we discuss the law of conservation of energy with relevant examples and specific heat calculations. Practical applications include understanding how specific heat capacities affect climate buffering.

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Thermochemistry

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  1. Chemistry Ms. Piela Thermochemistry

  2. Heat transfers in chemical reactions • Energy • Definition: The capacity to do work or supply heat • Basic Types: Mechanical, Electrical, Potential, Kinetic Thermochemistry –

  3. Energy that is stored within chemical bonds • Determined by the arrangement and bonding of atoms • Examples: Gasoline, Food, etc Chemical Potential Energy

  4. Exothermic processes (‘exits’) • Heat flows out of a system Exothermic & Endothermic System Surroundings

  5. System loses heat (negative sign) Examples: Explosions!, Sun, Fire Exothermic and Endothermic

  6. Endothermic processes (‘enter’) • Heat is absorbed from surroundings Exothermic and Endothermic System Surroundings

  7. System gains heat (positive sign), Heat change is > 0 Examples: Boiling water, ice melting Exothermic and endothermic

  8. Energy transfers due to temperature differences between objects Flows from hot to cold until temperature equilibrium is reached Heat –

  9. Energy cannot be created or destroyed, only transferred • Can be transformed • Example: Law of Conservation of Energy –

  10. Calorie (cal) – the amount of heat needed to raise 1 gram of water 1 °C • Heat calorie is not to be confused with dietary calories • 1 dietary Calorie = 1 kcal = 1000 cal • Joule (J) – metric unit of heat and energy Units of Heat Conversion Factor! Woo! 1 cal = 4.184 J

  11. It takes 50.2 J to raise the temperature of a 100.0 g piece of glass. How many calories is this? Conversion Practice

  12. A small chocolate bar has about 210,000 calories. How many Joules is this? Conversation Practice

  13. The amount of heat required to raise 1 gram of a substance 1 °C Heat capacity is the amount of heat required to raise a substance 1 °C Specific Heat – VS.

  14. q = heat (J or cal) m = mass (g) C = specific heat ∆T = Change in Temp. (°C) Specific Heat Equation q = mC∆T

  15. Specific heat plays a role in buffering climates • Specific heat of land = 0.8 J/g ° C • Specific heat of water = 4.184 J/g °C • How does this help buffer climates? • Water heats up slower and cools down slowly, releasing heat gradually Specific Heat Application: Land vs. Water

  16. On the notes, you’ll see a table of specific heat capacities for common metals. We’ll be using these throughout all the problems in this section, so keep it handy! Table of Specific Heat Capacities

  17. A 1.55 g piece of stainless steel absorbs 141 J of heat when its temperature increases by 178 °C. What is the specific heat of stainless steel? Specific Heat Example 1 .511 J/g °C

  18. How much heat is required to raise the temperature of 250.0 g of mercury from 52 °C to 60 °C? Specific Heat Example 2 280 J

  19. If the initial temperature of a 10.0 g substance is 20.0 °C and 100.0 J of heat are absorbed, what is the final temperature? The specific heat of the substance is 1.15 J/g °C. Specific Heat Example 3 28.7 °C

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