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Understanding Phase Changes: Energy Transfer in Solids, Liquids, and Gases

Explore the fascinating world of phase changes and how energy transfer influences states of matter. When heat energy is added or removed from solids, liquids, or gases, molecules experience changes in kinetic energy, affecting their movement and temperature. During phase changes, such as melting or freezing, the temperature remains constant despite energy transfer, as it is used to disrupt or restore intermolecular forces. Discover practical applications and the significance of melting and freezing points across pure substances.

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Understanding Phase Changes: Energy Transfer in Solids, Liquids, and Gases

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Presentation Transcript

  1. Hot Topics, Cool Stuff: Phase Changes

  2. When heat energy is added to a solid, liquid, or gas • Molecules gain kinetic energy, move faster, and temperature increases. • When heat energy is removed from a solid, liquid, or gas • Molecules lose kinetic energy, move slower, and temperature decreases

  3. During a Phase Change… • Ex. Solid liquid • attractive forces can no longer hold defined shape of solid due to greater movement of molecules: forces are broken • Heat energy is used to convert solid m.c. to liquid m.c.

  4. During a Phase Change… • Ex. Liquid  solid • attractive forces need to be restored to bring molecules back to closely-packed, defined shape • Heat energy used to convert liquid m.c. to solid m.c.

  5. When heat energy is either added or removed during a phase change: • Heat energy is used to break or restore forces between molecules so that the phase change can occur • This causes the molecules to become either organized or disorganized • Practical Application: As ice in a glass of lemonade melts on a hot day, the temperature inside of the glass will not change.

  6. During a phase change… energy is NOT used to speed up/ slow down the molecules. The temperature DOES NOT change! • Practical Application: As ice in a glass of lemonade is melting, the temperature inside of the glass will not change until ALL the solid ice has changed to liquid water.

  7. Parts of a Phase Diagram

  8. Melting Points and Freezing Points Are… • Equal • melting and freezing occur at the same temperature: reversible process

  9. Melting Points and Freezing Points Are… • Unique for specific, pure substances • nitrogen has a melting point of –209 oC • water has a melting point of 0oC • sodium metal has a melting point of 98 oC

  10. Melting Points and Freezing Points Are… • The same temperature for pure substances regardless of the amount of substance • 10 grams of water has a melting point of 0 oC • 100 grams of water has a melting point of 0 oC, but it will take longer for all of the solid to melt

  11. Heating Curve of 1 ice cube vs. Heating Curve of 8 ice cubes

  12. Phase Change: Liquid to Gas • Boiling occurs when the vapor pressure of the liquid is equal to the atmospheric pressure

  13. Phase Change: Liquid to Gas • In other words, the liquid pushes against the air just as much as the air pushes on the liquid • Vapor pressure (a.k.a- “Equilibrium Pressure”) = pressure that is exerted by liquid molecules and relates to the tendency of molecules and atoms to escape from a liquid • The higher the vapor pressure of a liquid at a given temperature, the lower the normal boiling point

  14. How do we calculate heat? Q = mCΔt Q = Heat (Joules) m = mass (grams) C = Specific Heat Capacity ( J/goC) Δt = change in Temperature (oC) = tf – ti ti= initial temperature tf =final temperature

  15. Specific Heat Capacity • Specific Heat Capacity = the amount of energy required to increase the temperature of one gram of substance by one degree Celsius. • Unique for each substance • The higher the specific heat capacity, the more energy required to change its temperature. • Practical application: This is why our metal ring stand gets hotter than the water inside of a beaker

  16. Specific Heat Capacity

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