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Understanding Energy: Types, Conservation, and Calculations in Chemistry

This lecture provides an in-depth exploration of energy, defining it as the capacity of a physical system to perform work. It distinguishes between kinetic and potential energy, discusses the units of energy like joules and calories, and explains the First and Second Laws of Thermodynamics. The concept of heat and specific heat capacity is highlighted, with practical examples demonstrating how to calculate heat energy in different materials. The intricate relationships between energy, work, and heat are examined, providing essential insights for students of chemistry and society's energy needs.

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Understanding Energy: Types, Conservation, and Calculations in Chemistry

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  1. Lecture 11 Energy, Chemistry and Society April 27, 2005

  2. What is energy? Energy: the capacity of a physical system to do work Work (W): equal to a force (F) multiplied by a distance (d); force exerted to move an object a certain distance Force (F): the change of velocity for a mass; mass (m) x acceleration (a)

  3. Energy: Two Kinds • Kinetic • Anything that is moving • Heat • Electricity • Potential • Stored energy • Chemical – remember bond = spring model • Hydroelectric • Roller coaster

  4. Energy Units • Joules (J) 1 J = amount of energy needed to lift a 1 kg book, 10 cm against Earth’s gravity = amount of energy in 1 heartbeat • calories (cal) = 4.184 J 1 cal = amount of heat needed to raise the temperature of 1 g of water 1 °C • Calories (food calories) 1 Cal = 1 kilocalorie = 1 kcal • Other energy units are British Thermal Units (btu), ergs, & foot-pounds. • Energy of fuels are usually expressed as kJ/mole or kcal/mole.

  5. Law: Conservation of Energy • Energy is always conserved. • It can change forms, but is never created or destroyed. • Can be mysterious… always consider the dissipation of energy to surrounding environment in the form of sound, heat, etc. • Ultimately, most energy on earth comes from the sun. • Biology: photosynthesis yields fossil fuels • Weather: wind, hydroelectric

  6. First Law of Thermodynamics • Also called, the Law of Conservation of Energy and Mass. • Thermodynamics: Physics that deals with the relationships and conversions between heat and other forms of energy. • Good Jeopardy! question… We’ll call these Ken Jennings points

  7. Converting Heat to Work to Electricity

  8. Efficiency Ex. from book discusses converting electricity into heating a home – this final electric heater has an efficiency of 98% Overall efficiency = 0.60 x 0.90 x 0.75 x 0.95 x 0.90 x 0.98 = 0.34 = 34% Why? Because heat cannot be completely converted into work

  9. Second Law of Thermodynamics • States: it is impossible to completely convert heat into work without further changes to the universe. • Entropy in a closed system cannot decrease.

  10. Definitions Heat or Thermal Energy: • Random motion of molecules Entropy: • a measure of the amount of energy in a physical system that cannot be used to do work • Disorder, randomness • High entropic states are more probable than lower ones

  11. Combustion • What makes a usable fuel? • Most common energy-generating reaction is combustion. • Recall that combustion is the combination of fuel with oxygen. CH4(g) + 2O2(g) CO2(g) + 2H20 (g) + energy Energy on reactant side = Endothermic Energy on this side (products) means the reaction is Exothermic

  12. Exothermic vs. Endothermic

  13. Experimentally determining heat energy

  14. Calculating energy • We’ll mainly consider heat • Historically, heat has been abbreviated as q • Transfers of heat cause changes in temperature – Heat transferred in: temperature increases – Heat transferred out: temperature decreases • q ∝ ΔT

  15. Calculating energy • Amount of heat needed to change temperature directly proportional to mass – More heat needed to raise temperature of 50 g of something than 10 g of something • q ∝ mΔT

  16. Calculating energy • Some things are harder to heat than others – Very easy to heat metals: small amount of heat needed to change temperature by 1 °C – Very hard to heat wood, water: large amount of heat needed to change temperature by 1 °C • Specific heat capacity: amount of heat needed to change temperature of 1 g of a material by 1 °C • q = msΔT

  17. Table of Specific Heat Capacities (s) for various compounds

  18. Example • How much heat is required to heat 100 g of water from 15 to 25 °C? q = msΔT q = 100g x 4.184 J/g•°C x 10 °C q = 4184 J = 4.184 kJ

  19. Example • To what final temperature can 75 g of copper, initially at 80 °C, be raised by the addition of 750 J of heat? q = msΔT final temp = 80 °C + 26 °C = 106 °C q ms 750 J 75 g x 0.387 J/g •°C ΔT = = = 26 °C

  20. Example • How much heat is released when 32 g of aluminum are cooled from 15 to 5 °C? q = msΔT q = 32 g x 0.900 J/g •°C x 10 °C q = 288 J

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