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!

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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!!

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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 +

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Intermolecular Forces and Coulomb’s Law continued

• 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.

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Intermolecular Forces and Coulomb’s Law continued

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

δ −

δ +

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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.

Intermolecular Forces and Coulomb’s Law continued

• 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.

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Intermolecular Forces and Coulomb’s Law continued

• 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 < ……..

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Intermolecular Forces and Coulomb’s Law continued

• 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

CHEM 1310 A/B Fall 2006

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.)

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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.

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Phase Transitions

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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.

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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.

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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.

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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.

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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.

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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?

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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.

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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.

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Example Continued

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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)

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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”.

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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

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Van’t Hoff Constants

What is ι for the following compounds?NaCl?

H2SO4?

C6H12O6?

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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?

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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”

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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?

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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”.

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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?

CHEM 1310 A/B Fall 2006