Phase Changes

Melting Point

0°C

Properties of Phases

Solid: volume = fixed shape = fixed movement of particles = vibrations strength of intermolecular bonds = very strong Liquid: volume = fixed shape = no movement of particles = sliding strength of intermolecular bonds = strong Gas: volume = no shape = no movement of particles = random motion strength of intermolecular bonds = negligible

Equations

q = mcΔT heat energy (J/cal) = mass (g) * heat capacity (J/gK) * (final temp - initial temp) heat capacity = the number of heat units needed to raise the temperature of a body by one degree q = mHf heat energy (J/cal) = mass (g) * heat of fusion (J/g) the heat of fusion = the amount of energy required to change 1 gram of the substance from solid to liquid at its melting point -q (metal) = q (water)

Phase Transition Chart

x-axis = heat energy y-axis = temperature the graph shows time passing as constant heat is supplied to ice at first the temperature of ice is raised to the melting point, then the energy is used to melt the ice (temperature does not change, but energy is still being supplied). The temperature is then raised to the boiling point and the energy is used to change the state from liquid to gas.

Boiling Point

100°C

Law of the Conservation of Energy

energy cannot be created nor destroyed but only changed from one form to another

Endothermic

heat moves from surroundings into a system (or heat moves from environment​ into water) heat goes IN (melting/boiling) melting: H2O (s) --> H2O (l) ΔH = +6.01KJ boiling: H2O (l) --> H2O (g) ΔH = +88.0KJ

Exothermic

heat moves out of a system to surroundings heat goes OUT (condensing/freezing) condensing: H2O (g) --> H2O (l) ΔH = -88.0KJ freezing: H2O (l) --> H2O (s) ΔH = -6.01KJ

Phase Transitions

s --> l = melting/fusion, l --> s = freezing l --> g = boiling, g --> l = condensing s --> g = sublimation, g --> s = deposition