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Why is water more stable as a solid below 0 o C but as a liquid above it?

Why is water more stable as a solid below 0 o C but as a liquid above it?. What controls why only some things happen?. Energy is neither created or destroyed during chemical or physical changes, but it is transformed from one form to another.  E universe = 0. TYPES of ENERGY

deirdre-wilson
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Why is water more stable as a solid below 0 o C but as a liquid above it?

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  1. Why is water more stable as a solid below 0 oC but as a liquid above it?

  2. What controls why only some things happen?

  3. Energy is neither created or destroyed during chemical or physical changes, but it is transformed from one form to another.Euniverse = 0

  4. TYPES of ENERGY Kinetic Potential Mechanical Gravitational Thermal Electrostatic Electrical Chemical Radiant

  5. Dropping a rock? • Gravitational  Mechanical • Gravitational  Thermal • Gravitational  Gravitational

  6. Using a flashlight? • Thermal  Radiant • Chemical  Thermal • Chemical  Radiant • Electrostatic  Radiant

  7. Driving a Car?

  8. One turn of a giant windmill produces 0.26 kWh of electricity. How many kJ is this?

  9. SYSTEM and SURROUNDINGS System: The thing under study Surroundings: Everything else in the universe Energy transfer between system and surroundings: Endo: heat added to system Exo: heat released by system

  10. Dissolution of KNO3 HCl(aq) + NaOH(aq)  H2O(l) + NaCl(aq)

  11. HEAT: What happens to thermal (heat) energy? Three possibilities: • Warms another object • Causes a change of state • Is used in an endothermic reaction

  12. Temperature Changes from Heat Exchange Bill, go to the next slide. You know you want to. Example 1: 5 g wood at 0 oC + 5 g wood at 100 oC Example 2: 10 g wood at 0 oC + 5 g wood at 100 oC Example 3: 5 g copper at 0 oC + 5 g copper at 100 oC Example 4: 5 g wood at 0 oC + 5 g copper at 100 oC Clicker Choices: 1: 0 oC 2: 33 oC 3: 50 oC 4. 67 oC 5: 100 oC 6: other

  13. Temperature Changes from Heat Exchange 5 g wood at 0 oC + 5 g wood at 100 oC • 0 oC • 33 oC • 50 oC • 67 oC • 100 oC • other

  14. Temperature Changes from Heat Exchange 10 g wood at 0 oC + 5 g wood at 100 oC • 0 oC • 33 oC • 50 oC • 67 oC • 100 oC • other

  15. Temperature Changes from Heat Exchange 5 g copper at 0 oC + 5 g copper at 100 oC • 0 oC • 33 oC • 50 oC • 67 oC • 100 oC • other

  16. Temperature Changes from Heat Exchange 5 g wood at 0 oC + 5 g copper at 100 oC • 0 oC • 33 oC • 50 oC • 67 oC • 100 oC • other

  17. Conclusion I: What happens to thermal (heat) energy? When objects of different temperature meet: • Warmer object cools • Cooler object warms • Thermal energy is transferred • qwarmer = -qcooler

  18. Conclusion II: What controls the magnitude of an object’s temperature change?

  19. Quantitative: Calculating Heat Exchange: Specific Heat Capacity

  20. Calculate the specific heat capacity of copper: Calculate the specific heat capacity of wood:

  21. Which has the smallest specific heat capacity? • Wood • Copper • Silver • Water

  22. Specific Heat Capacity The energy required to heat one gram of a substance by 1 oC. Usefulness: #J transferred = S.H. x #g x T How much energy is used to heat 250 g water from 17 oC to 100 oC?

  23. What happens to thermal (heat) energy? When objects of different temperature meet: • Warmer object cools • Cooler object warms • Thermal energy is transferred • qwarmer = -qcooler specific heat x mass x T = specific heat x mass x T warmer object cooler object

  24. Heat transfer between substances:

  25. Conceptually Easy Example with Annoying Algebra: If we mix 250 g H2O at 95 oC with 50 g H2O at 5 oC, what will the final temperature be?

  26. Thermal Energy and Phase Changes First: What happens?

  27. Thermal Energy and Phase Changes First: What happens?

  28. Thermal Energy and Phase Changes First: What happens?

  29. But what’s really happening? • Warming: • Molecules move more rapidly • Kinetic Energy increases • Temperature increases • Melting/Boiling: • Molecules do NOT move more rapidly • Temperature remains constant • Intermolecular bonds are broken • Chemical potential energy (enthalpy) increases

  30. Energy and Phase Changes: Quantitative Treatment Melting: Heat of Fusion (DHfus) for Water: 333 J/g Boiling: Heat of Vaporization(DHvap) for Water: 2256 J/g

  31. Total Quantitative Analysis Convert 40.0 g of ice at –30 oC to steam at 125 oC Warm ice: (Specific heat = 2.06 J/g-oC) Melt ice: Warm water (s.h. = 4.18 J/g-oC)

  32. Total Quantitative Analysis Convert 40.0 g of ice at –30 oC to steam at 125 oC Boil water: Warm steam (s.h. = 1.92 J/g-oC)

  33. Enthalpy Change and Chemical Reactions DH = energy needed to break bonds – energy released forming bondsExample: formation of water: DH = ?

  34. Enthalpy Change and Chemical Reactions DH is usually more complicated, due to solvent and solid interactions. So, we measure DH experimentally. Calorimetry Run reaction in a way that the heat exchanged can be measured. Use a “calorimeter.”

  35. Calorimetry Experiment N2H4 + 3 O2 2 NO2 + 2 H2O Energy released = E absorbed by water + E absorbed by calorimeter Ewater = Ecalorimeter = Total E = H = energy/moles = 0.500 g N2H4 600 g water 420 J/oC

  36. Hess’s Law If reactions can be “added” so can their H values.

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