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Ch. 10.4 and 11 Thermodynamics

Ch. 10.4 and 11 Thermodynamics. You must turn in your notes. Energy. Definition: the capacity to do work, or to produce heat; SI (metric) unit is J oule Non-SI unit is cal orie 1 cal = 4.18 J 1 calorie = the amount of energy needed to raise 1 gram of water 1° C

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Ch. 10.4 and 11 Thermodynamics

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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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