Intermolecular Forces and

Liquids and SolidsChapter 12

Midterm II

• Any conflicts with March 20? If yes, let me know ASAP. The original date was March 22.

Phase Diagram of Water

• Note the high critical temperature and critical pressure:– These are due to the strong van der

Waals forces between water molecules.

• The slope of the solid–liquid line is negative.– This means that increasing the pressure above 1 atm will raise the boiling point and

lower the melting point.– Lower the melting point?

Phase Diagram of Carbon Dioxide

Carbon dioxide cannot exist in the liquid state at pressures below 5.11 atm; CO2 sublimes at normal pressures.

Phase Diagram of Carbon Dioxide

Carbon dioxide cannot exist in the liquid state at pressures below 5.11 atm; CO2 sublimes at normal pressures.

At 1 atm, solid CO2 does not melt at any temperature.Instead, it sublimes to form CO2 vapor. Why might it beuseful as a refrigerant?

Phase Diagram of Carbon Dioxide

Carbon dioxide cannot exist in the liquid state at pressures below 5.11 atm; CO2 sublimes at normal pressures.

If you want to send something frozen across the country,you can pack it in dry ice. It will be frozen when it reaches its destination, and there will be no messy liquid left overlike you would have with normal ice.

The slope of the curve between solid and liquid is positive forCO2 as well as almost all other substances. Why does waterdiffer?

Freeze-drying

• Completely remove water from some material, such as food, while leaving the basic structure and composition of the material intact

• Two reasons– Keeps food from spoiling for a long period of time – Significantly reduces the total weight of the food

• How?– Freeze the material– Lower the pressure (<0.006 atm)– Increase the temperature slightly

Normal (right) and freeze-dried (left) spaghetti

Freeze-drying

• How?– Freeze the material– Lower the pressure– Increase the temperature slightly

Normal (right) and freeze-dried (left) spaghetti

Physical Properties of Solutions

Chapter 13

13.1

A solution is a homogenous mixture of 2 or more substances

The solute is(are) the substance(s) present in the smaller amount(s)

The solvent is the substance present in the larger amount

A saturated solution contains the maximum amount of a solute that will dissolve in a given solvent at a specific temperature.

An unsaturated solution contains less solute than the solvent has the capacity to dissolve at a specific temperature.

A supersaturated solution contains more solute than is present in a saturated solution at a specific temperature.

Sodium acetate crystals rapidly form when a seed crystal isadded to a supersaturated solution of sodium acetate.

13.1

Solutions

The intermolecular forces between solute and solvent particles must be strong enough to compete with those between solute particles and those between solvent particles.

How Does a Solution Form?

As a solution forms, the solvent pulls solute particles apart and surrounds, or solvates, them.

How Does a Solution Form

If an ionic salt is soluble in water, it is because the ion-dipole interactions are strong enough to overcome the lattice energy of the salt crystal.

Energy Changes in Solution

• Simply, three processes affect the energetics of the process:Separation of solute

particlesSeparation of solvent

particlesNew interactions between

solute and solvent

Energy Changes in Solution

The enthalpy change of the overall process depends on H for each of these steps.

13.2

Three types of interactions in the solution process:• solvent-solvent interaction• solute-solute interaction• solvent-solute interaction

Hsoln = H1 + H2 + H3

“like dissolves like”

Two substances with similar intermolecular forces are likely to be soluble in each other.

• non-polar molecules are soluble in non-polar solvents

CCl4 in C6H6

• polar molecules are soluble in polar solvents

C2H5OH in H2O

• ionic compounds are more soluble in polar solvents

NaCl in H2O or NH3 (l)

13.2

Concentration UnitsThe concentration of a solution is the amount of solute present in a given quantity of solvent or solution.

Percent by Mass

% by mass = x 100%mass of solutemass of solute + mass of solvent

= x 100%mass of solutemass of solution

13.3

Mole Fraction (X)

XA = moles of A

sum of moles of all components

Concentration Units Continued

M =moles of solute

liters of solution

Molarity (M)

Molality (m)

m =moles of solute

mass of solvent (kg)

13.3

What is the molality of a 5.86 M ethanol (C2H5OH) solution whose density is 0.927 g/mL?

m =moles of solute

mass of solvent (kg)M =

moles of solute

liters of solution

13.3

Strategy:

Find mass of solventKnow mass of solute + mass of solvent = mass of solutionIf mass of solution and mass of solute known, can calculate mass of solventCan calculate mass of solute from moles of soluteCan calculate mass of solution from density and volume of the solutionSolve

What is the molality of a 5.86 M ethanol (C2H5OH) solution whose density is 0.927 g/mL?

m =moles of solute

mass of solvent (kg)M =

moles of solute

liters of solution

0.586 moles of solute per 1 L of solution:5.86 moles ethanol = 270 g ethanol927 g of solution (1000 mL x 0.927 g/mL)

mass of solvent = mass of solution – mass of solute

= 927 g – 270 g = 657 g = 0.657 kg

m =moles of solute

mass of solvent (kg)=

5.86 moles C2H5OH

0.657 kg solvent= 8.92 m

13.3

Temperature and SolubilitySolid solubility and temperature

solubility increases with increasing temperature

solubility decreases with increasing temperature

13.4

No clear correlation between ΔHsoln and the variation of solubility with temperature

Fractional crystallization is the separation of a mixture of substances into pure components on the basis of their differing solubilities.

Suppose you have 90 g KNO3 contaminated with 10 g NaCl.

Fractional crystallization:

1. Dissolve sample in 100 mL of water at 600C

2. Cool solution to 00C

3. All NaCl will stay in solution (s = 34.2g/100g)

4. 78 g of PURE KNO3 will precipitate (s = 12 g/100g). 90 g – 12 g = 78 g

13.4

Temperature and Solubility

Gas solubility and temperature

solubility usually decreases with

increasing temperature

13.4

Pressure and Solubility of Gases

13.5

The solubility of a gas in a liquid is proportional to the pressure of the gas over the solution (Henry’s law).

c = kPc is the concentration (M) of the dissolved gasP is the pressure of the gas over the solutionk is a constant (mol/L•atm) that depends only on temperature

low P

low c

high P

high c

Colligative Properties of Nonelectrolyte Solutions

Colligative properties are properties that depend only on the number of solute particles in solution and not on the nature of the solute particles.

Vapor-Pressure Lowering

Raoult’s law

If the solution contains only one solute:

X1 = 1 – X2

P 10 - P1 = P = X2 P 1

0

P 10 = vapor pressure of pure solvent

X1 = mole fraction of the solvent

X2 = mole fraction of the solute13.6

P1 = X1 P 10

PA = XA P A0

PB = XB P B0

PT = PA + PB

PT = XA P A0 + XB P B

0

Ideal Solution

13.6

PT is greater thanpredicted by Raoults’s law

PT is less thanpredicted by Raoults’s law

ForceA-B

ForceA-A

ForceB-B< &

ForceA-B

ForceA-A

ForceB-B> &

13.6

Fractional Distillation Apparatus

13.6

Boiling-Point Elevation

Tb = Tb – T b0

Tb > T b0 Tb > 0

T b is the boiling point of the pure solvent

0

T b is the boiling point of the solution

Tb = Kb m

m is the molality of the solution

Kb is the molal boiling-point elevation constant (0C/m)

13.6

Freezing-Point Depression

Tf = T f – Tf0

T f > Tf0 Tf > 0

T f is the freezing point of the pure solvent

0

T f is the freezing point of the solution

Tf = Kf m

m is the molality of the solution

Kf is the molal freezing-point depression constant (0C/m)

13.6

13.6

What is the freezing point of a solution containing 478 g of ethylene glycol (antifreeze) in 3202 g of water? The molar mass of ethylene glycol is 62.01 g.

Tf = Kf m

m =moles of solute

mass of solvent (kg)= 2.41 m=

3.202 kg solvent

478 g x 1 mol62.01 g

Kf water = 1.86 0C/m

Tf = Kf m = 1.86 0C/m x 2.41 m = 4.48 0C

Tf = T f – Tf0

Tf = T f – Tf0 = 0.00 0C – 4.48 0C = -4.48 0C

13.6

Osmotic Pressure ()

13.6

Osmosis is the selective passage of solvent molecules through a porous membrane from a dilute solution to a more concentrated one.

A semipermeable membrane allows the passage of solvent molecules but blocks the passage of solute molecules.

Osmotic pressure () is the pressure required to stop osmosis.

dilutemore

concentrated

HighP

LowP

Osmotic Pressure ()

= MRT

M is the molarity of the solution

R is the gas constant

T is the temperature (in K) 13.6

A cell in an:

isotonicsolution

hypotonicsolution

hypertonicsolution

13.6

Colligative Properties of Nonelectrolyte Solutions

Colligative properties are properties that depend only on the number of solute particles in solution and not on the nature of the solute particles.

13.6

Vapor-Pressure Lowering P1 = X1 P 10

Boiling-Point Elevation Tb = Kb m

Freezing-Point Depression Tf = Kf m

Osmotic Pressure () = MRT

Colligative Properties of Electrolyte Solutions

13.7

0.1 m NaCl solution 0.1 m Na+ ions & 0.1 m Cl- ions

Colligative properties are properties that depend only on the number of solute particles in solution and not on the nature of the solute particles.

0.1 m NaCl solution 0.2 m ions in solution

van’t Hoff factor (i) = actual number of particles in soln after dissociation

number of formula units initially dissolved in soln

nonelectrolytesNaCl

CaCl2

i should be

12

3

Boiling-Point Elevation Tb = i Kb m

Freezing-Point Depression Tf = i Kf m

Osmotic Pressure () = iMRT

Colligative Properties of Electrolyte Solutions

13.7

A colloid is a dispersion of particles of one substance throughout a dispersing medium of another substance.

Colloid versus solution

• collodial particles are much larger than solute molecules

• collodial suspension is not as homogeneous as a solution

13.8

The Cleansing Action of Soap

13.8

Chemistry In Action: Desalination