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18 Cards in this Set
- Front
- Back
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Chemical change |
Change that turns one molecule into another Examples: burning and rusting (reactions w/ oxygen) |
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Physical change |
Change that does not alter a molecule's chemical identity Examples: phase changes, solvation |
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Fusion |
Phase change from solid to liquid (a/k/a melting) Endothermic (requires the input of heat), so delta H (enthalpy) is positive Molecule becomes less ordered so delta S (entropy) is also positive |
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Vaporization |
Phase change from liquid to gas (a/k/a boiling) Endothermic (requires the input of heat), so delta H (enthalpy) is positiveMolecule becomes less ordered so delta S (entropy) is also positive |
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Sublimation |
Phase change from solid to gas Endothermic (requires the input of heat), so delta H (enthalpy) is positive Molecule becomes less ordered so delta S (entropy) is also positive |
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Condensation |
Phase change from gas to liquid Exothermic (releases heat), so delta H (enthalpy) is negative Molecule becomes more ordered so delta S (entropy) is also negative |
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Crystallization |
Phase change from liquid to solid (a/k/a freezing) Exothermic (releases heat), so delta H (enthalpy) is negative Molecule becomes more ordered so delta S (entropy) is also negative |
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Deposition |
Phase change from gas to solid Exothermic (releases heat), so delta H (enthalpy) is negative Molecule becomes more ordered so delta S (entropy) is also negative |
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Amount of heat input or released (q) |
q = m x C x delta T m=mass C=specific heat delta T = change in temperature (same in Celsius and Kelvin) *Pay attention to the units used for q *If temperature increases, delta T is positive -If temperature decreases, delta T is negative |
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Specific heat |
Amount of heat required to raise one gram of a substance by one degree Celsius |
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Amount of energy required to cause a phase change |
q = n x heat of phase change n = number of moles if heat of phase change is given in J/mol n = number of grams if heat of phase change is given in J/gram |
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Specific heat of a mixed substance |
m x C x delta T = -(m x C x delta T) |
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State function |
Initial and final results are the same regardless of the path taken to get there Examples: Enthalpy (delta H) |
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Triple point |
Point at which the solid, liquid, and gas phases are in equilibrium |
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Critical point |
Point beyond which there is no phase change between liquid and gas (liquid and gas are indistinguishable) Substances beyond this point are called supercritical fluids |
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Boiling point |
Point at which atmospheric pressure and vapor pressure are equal -Increasing the temperature of a substance increases its vapor pressure until it reaches atmospheric pressure, at which point the substance undergoes a phase transition from the liquid to the gas phase -When atmospheric pressure is lower (as in high altitudes), boiling point is also lower because vapor pressure reaches atmospheric pressure more quickly |
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Phase change diagram for water |
Abnormal because the line of equilibrium between solid and liquid has a downward slope (rather than an upward one), indicating that the liquid phase has a higher pressure (and is therefore denser) than the solid phase -Explains why ice floats -Explained by hydrogen bonding |
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Phase change diagram for CO2 |
Line of equilibrium between solid and liquid phases occurs above normal atmospheric pressure (1 atm or 760 torr) so at normal atmospheric pressure, CO2 passes directly from solid to gas without entering the liquid phase |
