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Key Concepts of Thermodynamics: Energy, Heat, and Reactions

This comprehensive overview covers essential concepts of thermodynamics, including energy definitions, heat transfer, temperature measurement, specific heat calculations, and phase changes. It highlights the principles of exothermic and endothermic reactions, measures of disorder (entropy), and Gibbs' free energy. Practical problems and relevant formulas are provided for deeper understanding. This guide serves as a solid foundation for students studying thermal dynamics and the behavior of energy in physical and chemical processes.

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Key Concepts of Thermodynamics: Energy, Heat, and Reactions

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  1. Ch. 10.4 and 11Thermodynamics You must turn in your notes

  2. Energy • Definition: the capacity to do work, or to produce heat; • SI (metric) unit is Joule • Non-SI unit is calorie • 1 cal = 4.18 J • 1 calorie = the amount of energy needed to raise 1 gram of water 1° C • Heat - energy that is transferredfrom a warm object to acooler object; represented by “q” • Temperature - a measure of the average kinetic energy of an object

  3. Specific heat/Change of Temp • Quantity of heat needed to raise 1 gram of a substance 1° C; unit is J/g°C • FORMULA: q= c x m x ΔT • q= heat (J) • c = specific heat (J/g°C) • m = mass (g) • T = change in temp. (°C)

  4. Specific Heat Practice Problem #1 • Calculate the amount of heat in joules needed to warm 250. g of water from 25.0°C to 95.0°C. (c=4.184 J/°C g)

  5. Practice Problem #2 How much heat is lost when 50.0 grams of Al is cooled from 130.0 °C to 62.0 °C? The specific heat of Al is 0.897 J/g°C

  6. Change of state and heat • heat of fusion- amount of heat needed to melt 1 gram of a substance at its melting point • Hf copper = 205 J/g • Hf water = 80 cal/g= 334 J/g • q = mHf • heat of vaporization- amount of heat needed to boil 1 gram of a substance at its boiling point • Hv water = 540 cal/g = 2260 J/g • q= mHv

  7. Heating/Cooling Curves Change of Temperature Change of Phase Know which formula to use when!

  8. Change of state Practice Problems • Calculate the amount of heat, in Joules, needed to melt 70.0g of copper at its melting point. • Calculate the heat required, in calories, to change 250g of water at 100°C to steam at 100°C. • Calculate the amount of heat needed to change 20g of ice at -10.0°C to water at 80.0°C.

  9. Enthalpy (H) • a measure of heat content of a system • H = change in heat content that accompanies a process • Hrxn = Hfinal - Hinitial • Hrxn = Hproducts - Hreactants • ** ΔHrxn can also be written as ΔHf, for heat of formation** • **Chemical systems in the world tend to achieve the lowest possible energy. Would this occur in an exothermic or an endothermic reaction?

  10. Exothermic reactions • chemicals react and give off heat (feel hot); H is negative; products are more stable • 4Fe + 3O2 2Fe2O3 + 1625 kJ reactants products

  11. Endothermic reactions • chemicals need to absorb energy in order for the reaction to take place (feel cool); H is positive; reactants are more stable • 27 kJ + NH4NO3 NH4+ + NO3- products reactants

  12. Practice problems #1 • CO (g) + NO (g)  CO2 (g) + N2 (g)

  13. Practice #2 • CH4 (g)+ O2 (g)  CO2 (g)+ H2O (g)

  14. Practice #3 • N2 (g) + O2 (g)  NO2 (g)

  15. Entropy (S) • Measure of disorder or chaos in a system • Law of disorder – states that things move spontaneously in the direction of maximum chaos or disorder • ΔSsystem = SP - SR • If ΔS is +, there is an increase in entropy. • If ΔS is -, there is a decrease in entropy

  16. Rules for disorder 1. Entropy increases as particles move apart. • Gas > Liquid > Solid • I2 (s)  I2 (g) 2. Entropy increases when you divide a substance into parts (when the total number of products > the total number of reactants) • 2H2O  2H2 + O2 3. Entropy increases as temp increases b/c particles move faster. (unit for entropy is J/K mole) 4. Entropy increases when you dissolve a solid into a liquid. Entropy decreases when you dissolve a gas into a liquid.

  17. Entropy Examples • Water (liquid) - water (solid) • KCl(s)  KCl (l) • C(s) + O2 (g)  CO2 (g)

  18. Gibbs’ Free Energy (G) • Energy available to do work • Relates enthalpy (H) and entropy (S), using the equation: • ΔG = ΔH – TΔS • ΔG = GP - GR

  19. Still confused? • Try these online notes • http://www.sciencegeek.net/Chemistry/Powerpoint/Unit7/Unit7_files/frame.htm

  20. Spontaneous Reactions • If ΔG is negative, the reaction will occur spontaneously. • If ΔG is positive, the reaction will NOT occur spontaneously.

  21. Quick Spontaneous Chart

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