chapter 6: condensed phases and phase transitions

CHAPTER 6: Condensed Phases and Phase

Transitions• Molecules in solid and liquid phases are

much closer to each other than in gas phase.

• Solid and liquid are called “condensed phases” and are harder to describe by simple mathematical equations.

• Can no longer ignore “intermolecular forces” as is done for gases.

CHEM 1310 A/B Fall 2006

Types of Intermolecular Forces

• Electrostatic: Dipole-dipole, ion-dipole, induced dipole (opposite charges attract)

• Dispersion forces: correlations between electron motions (my area of research!)

• Hydrogen bonding

The strength of these interactions ranges from 0 to 5 kcal/mol - much weaker than

intramolecular covalent bond or ionic bonds!


Intermolecular Forces and Coulomb’s Law

All of the intermolecular forces can be pictured as manifestations of Coulomb’s law.

• Ion-ion interactions lead to ionic bond.

Strong Attraction!!


Intermolecular Forces and Coulomb’s Law continued

• Ion – Dipole: You can have non-bonded ion-dipole attraction; these are weaker than ion-ion interactions.

δ −

δ − δ −

δ −

δ +

δ +

δ +

δ +

δ +

δ +

δ +

δ +



• Ion-Inducded Dipole: Ions can even be attracted to molecules with no dipole moment – they can create or induce a dipole moment

δ − δ +

Li cation “pulls” electrons in Aratom twoards it, inducing a dipole

This is a weaker attraction than an ion-dipole interaction

The strength of the interaction depends on how polariziable(how easily the electron density deforms) the atom or molecule is.



• Dipole-Dipole Interactions – Two molecules with dipole moments line up to create a favorable interaction

δ −

δ +


δ −

δ +

δ +

δ −

Two dipoles will line up such that the opposite charges are closer than the like charges, such that the favorable interactions will outweigh the unfavorable repulsions.


• Dipole-Induced Dipole: These interactions are weaker than ion-induced dipole.

δ − δ +

δ − δ +

The Ar atom deforms slightly as the partial positive charge of the HCl molecule approaches.



• Induced dipole – Induced dipole: Also called London Dispersion forces, these favorable interactions result from temporary fluctuations, which cause fluctuations in other molecules.

δ − δ −δ + δ +

Strength of dispersion interactions: larger atoms are more polarizable

He2 < Ne2 < Ar2 < ……..



• Hydrogen Bonding: Interaction which forms between a lone pair on N,O, or F and a hydrogen bonded to N, O, or F

Water forms an extensive network of hydrogen bonds as the lone pairs on oxygen (red atom) interact with hydrogen of other molecules. A hydrogen bond is an intermolecular interaction; it occurs between atoms in different molecules.

Hydrogen bonding can be very strong in the gas phase, but usually only provides a few kcal/mol stabilization in solution.

Hydrogen Bond


Gas vs. Condensed Phases

• Intermolecular forces make molecules attract each other and stick together in the liquid and solid phases.

• Gas phase results if molecules are moving too fast to get stuck together. (For example, at high temperatures or low pressures.)


Intermolecular Forces and Boiling Point

• The stronger the intermolecular forces, the higher the boiling point of a liquid.

• Which has a higher boiling point, radon or helium? H20 or H2S?

• Usually it is hardest to boil ionic liquids, then polar liquids, then nonpolar liquids.


Phase Transitions


Phase Equilibrium• Even when a substance

is mainly in one phase, a small fraction is sometimes in a different phase. For instance, volatile liquids have some molecules escape as a gas from liquid phase.

• The gas and the liquid molecules are in a dynamic equilibrium.


Vapor Pressure

• The small amount of gas above a liquid has a pressure, P, called the vapor pressure.

• Vapor pressure generally increases with temperature, until the liquid boils.

• The boiling point of a liquid is the point at which vapor pressure = external pressure.


Phase Diagrams

• A phase diagram is a plot of P vs. T which can be used to predict what phase (solid, liquid, or gas) is seen for a given temperature and pressure.


Phase Diagram for CO2

• X is the triple point –solid, liquid, and gas exist simultaneously.

• Z is the critical point –can not maintain a distinction between liquid and gas beyond this point

• Solid/liquid line usually curves to the right (at a given T, an increase in P means the substance freezes); water is an exception.


Colligative Properties

• Colligative Properties describe how the physical properties (like boiling point) of a solvent change when a solute is added

• For dilute solutions, this will depend only on the number of solute particles, not their type.


Descriptions of Solutions

• Mass Percent = (Mass of component / Total Mass of Mixture) X 100

Example: What is the mass percentage of NaI in a solution that is 5 g of NaI per 100g of solution?


Descriptions of Solutions

• Mole Fraction = Number of moles A / Total number of moles

• Molality (m) =moles solute / kilograms solvent

Note that molality is NOT the same as molarity(moles/L). Sometimes they are numerically similar since 1 L of water weights 1 kg at 25 ºC.


Molality Example

• A solution prepared by mixing 20.00 g of CdCl2 with 80.00 g of water has a density of 1.1988 g cm-3 at 20 ºC. Compute the molarity and molality of this solution.


Example Continued


Vapor Pressure and Raoult’s Law

• If a nondissociating, nonvolatile solute is added to a solvent, the solvent’s vapor pressure decreases in an amount proportional to the mole fraction of the solution.

∆Psolvent = -Xsolute Pºsolventor equivalently

Psolvent = Xsolvent Pºsolvent (Raoult’s Law)


Raoult’s law and boiling point

• If a solute lowers the vapor pressure of a solvent, it’s harder (requires a higher temperature) to bring the solvent’s vapor pressure up to Patm to boil.

• Solutes raise the boiling points.• This is termed “boiling point elevation”.


Determining Boiling point Elevation

∆Tb = kbmι

∆Tb = change in the boiling pointkb = solvent constantm = molality of soulteι = Van’t Hoff constant – how many particles

a solute breaks into in solution


Van’t Hoff Constants

What is ι for the following compounds?NaCl?

H2SO4?

C6H12O6?


Boiling point elevation example

• 4.58 g of (nonvolatile) picnic acid (C6H3N3O7) dissolves in 240.0 g of chloroform, which has a kb = 3.63 K kg mol-1. How much does the boiling point increase?


Freezing Point Depression

• How are freezing points affected by solutes?

• Solute particles make it difficult to form an orderly solid. The solution must get colder to freeze

• “Freezing point depression”


Freezing Point Depression

∆Tf = kfmι

Same concept as boiling point elevation but kfis different than kb.

For example, why is salt spread on icy roads?


Mixes of Volatile Compounds

• What happens if we mix two volatile compounds together?

• If Raoult’s law still applied,P1 = X1 P1ºP2 = X2 P2º

If this is true, then the solution is “ideal”.


Non-ideal solutions• For a mix of two volatile compounds, if Raoult’s

law still held, then P1=X1P1o, P2=X2P2

o. If this is true, the solution is “ideal”

• Even if the solution is non-ideal, for small X2, a similar equation holds if we replace P2

o with “kH”: P2 = kH X2 (Henry’s Law)

• The greater the vapor pressure of a gas above a liquid, the more the gas will dissolve in the liquid

• What could you do to keep a Coke from going flat?


chapter 6: condensed phases and phase transitions

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