Lecture'Notes'1:'Physical'Equilibria'–'Vapor'Pressure'!Our!first!exploration!of!equilibria!will!examine!physical!equilibria!(no!chemical!changes)!in!which!the!only!changes!occurring!are!matter!changes!phases.!!As!is!the!case!with!all!equilibria!we!will!see!that!the!stable!state!depends!on!the!free!energy!of!the!system!and!that!at!equilibria!all!species!that!are!present!have!the!same!free!energy!!Free!energy!is!the!balance!between!energy!(on!our!case!we!will!specifically!look!at !t he!ent ha lpy)!and!entropy.!!The!balance!between!these!two!varies!with!the!temperature.!!Areas!of!physical!equilibria!we!will!examine!are!!•!Phase!transitions.!!Changes!between!the!solid,!liquid,!and!gaseous!forms!of!a!substance.!!For!example:!why!does!diethyl!ether!boil!at!a!lower!temperature!than!ethanol?!!•!!Mixtures.!!How!do!energy!and!entropy!govern!the!mixing!of!two!substances?!!!For!example:!why!do!oil!and!water!not!mix.!!•!Colligative!properties.!!How!does!making!a!mixture!affect!the!phase!transitions!of!the!system?!!For!example:!why!does!salt!water!freeze!at!a!lower!temperature!than!pure!water?!!To!understand!this,!we!will!draw!on!two!large!areas!from!CH301!!! Thermodynamics! and!! Intermolecular!Forces!!REVIEW!these!two!areas.!!!!The first topic we will examine is vapor pressure. You should have covered this topic briefly during your discussion of intermolecular forces. Definition of vapor pressure: the partial pressure of a substance in equilibrium with its condensed phase (liquid or solid). The pressure of the vapor over a liquid (solid) at equilibrium.The stability of any given state is given by its free energy. At constant temperature and pressure we utilize the Gibb’s Free Energy G = H – TS G is the free energy. H is the enthalpy. T is the temperature. S is the entropy. The lower the free energy, the more stable the state. Thus lower enthalpy is more stable and higher entropy is more stable at a given T. For phases we have a competition. Intermolecular forces lead to attractions between the molecules. These means the molecules have a lower enthalpy (energy) when they are close together. This favors the solid state. In contrast, from an entropy standpoint, the higher the entropy the lower the free energy. This favors the gas phase as it has the highest entropy. The favored state then depends on the temperature. When the temperature is low, the energy dominates and substances are solid. When the temperature is high, the entropy dominates and substances are gaseous. The exact temperature of a phase change depends on the substance (intermolecular forces). Entropy The highest entropy state for any substance at a given temperature would be if the atoms/molecules were scattered out as far as possible. Thus from an entropy standpoint everything should be a gas exploring the far reaches of the universe. However, all substances have intermolecular forces (interatomic forces) (IMF) that lead to attractions between the molecules. (it might be good to rattle off a few of the IMFs here, just in case IMF term doesn’t ring a bell w/ students—they’re likely to remember ion-ion, dipole, London, etc..) This means that they have a lower energy when they are close together as compared to far apart. Thus the IMF are “holding” the molecules close together. The stronger these forces are, the more difficult it is for the molecules to wander away from each other. Conversely, the lower the IMF the easier ’the wandering’ is. However, no matter the magnitude of the IMF there can always be a few molecules that escape. The number of these molecules and thus the partial pressure of the substance varies with the IMF. Thus at a given temperature, if the IMF are strong, the vapor pressure is low as only a few molecules can overcome the attractions of the molecules for each other. In contrast, a liquid with weak IMF will have many molecules in the gas phase and thus a higher vapor pressure.In CH301 you likely ranked the boiling point of molecules based on their IMF. Boiling point and vapor pressure are the same concept. The boiling point is the temperature at which the vapor pressure equals the total pressure. When comparing vapor pressures we need to be discussing the same temperature. Thus at room temperature, the substance with the lowest boiling point will have the highest vapor pressure (easiest to get into the gas phase). The highest boiling point will have the lowest vapor pressure. In general the trend is to rank things according to the accepted magnitude (or strength) of IMF. Generally dispersion (instantaneous dipoles) forces are the weakest, dipole-dipole interactions stronger, H-bonding next, followed by ionic forces (which are more bonding than IMF). However, as you may or may not have discovered, these trends are not set in stone. Everything has dispersion forces and they can be very large. In addition, some molecular dipoles can be weak. Thus it is very possible for completely non-polar molecules to have stronger IMF than similar polar molecules. For example, the boiling point of CCl4 (no dipole) is higher than that of CHCl3 (polar). What will always correlate is the vapor pressure and the IMF. Larger IMF = lower vapor
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