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CHEMISTRYThe Molecular Nature of Matter and Change

Third Edition

Chapter 13

The Properties of Mixtures:

Solutions and Colloids

Definitions

Solutions – Homogeneous Mixtures

Particles are individual atoms, ions, or

small molecules.

Colloids – Heterogeneous Mixtures

Particles are either macromolecules or

aggregations of small molecules that are

not large enough to settle out.

Solution TermsSolute vs. Solvent

Concepts to use for ‘Solvent’

Most abundant component

Physical state matches solution physical state

Miscible

Mix in any proportion

Solubility

Maximum amount of solute dissolved in a fixed

amount of solvent at a specified temperature, given that

excess solute is present

Solubilities

S, (NaCl) = 39.12 g / 100 mL water @ 100oC

S, (AgCl) = 0.0021 g / 100 mL water @ 100oC

S, {(NH4)2SO4} = 931 g / 100 mL water @ 100oC

Table 13.1

What makes different

substances mix?

“Like Dissolves Like”

Intermolecular Forces must be

Similar!

Intermolecular

Interactions in Mixtures

Attraction Energies in

kJ/mol

Fig. 13.1b

Intermolecular Interactions in

Mixtures

[Part 2]

Attraction Energies in

kJ/mol

Solutions:

Solvent vs.

Solute

“Like Dissolves Like”

Intermolecular Interactions

Table 13.2

As the

percentage of

the molecule

which doesn’t

interact

favorably with

water

increases, the

solubility

decreases.

Prob. 13.1

CH3CH2CH2CH2OH H C

H

C

H

C

H

C

H

H H

H

O

H

H

HOCH2CH2CH2CH2OH O C

H

C

H

C

H

C

H

H H

H

O

H

HH

Fig B13.1

Amphipathic Molecules: Surfactants {surface active agents}

Soaps

Types of Solutions

o Solid in Liquid

o Liquid in Liquid

o Solid in Solid

o Gas in Liquid

o Gas in Gas

o Gas in Solid

Solid Solutions: Alloys

Table 13.3

Gas in Liquid

Solutions

Enthalpy Changes in Solution Processes

Separation of solute

Separation of solvent

Mixing of solute and solvent

Lattice energy Hlatt:

Hsoln = Hlatt + Hhydr

Lattice energy is energy gained when ions form a solid structure

From the gas phase.

This causes Hsolute to be positive and equal in magnitude to Hlatt

Heats of hydration are always negative, so a dissolution can be

Exothermic or endothermic

Components of Solvation Enthalpies

exothermic

endothermic

M(g) M(aq)H2O

H <0 {always}

X(g) X(aq)H2O

H <0 {always}

H > 0 {always}

MX(S) X(aq)H2O

M(aq)

Charge density (ratio of charge to

size)

2) 2+ ions attract more than +1 of

same size

3) Small 1+ attracts more than

large +2 ion

Fig. 13.7

Heats of Solution

Why do endothermic solution processes occur at all?

Systems that increase in the degrees of freedom of constituents

are favorable.ENTROPY – measure of a system’s degrees of freedom:

Generally a larger number of pieces produced is entropically

favored.

Fig. 13.8

Dissolving Substances in Water: An EQUILIBRIUM Process

Fig. 13.9

Crystallization

Saturated

Supersaturated

Fig. 13.10

Solubility and

Temperature

Effects

Generally,

increasing

temperature

increases

solubility.

Gas solubility

always

decreases

with

temperature.

Fig. 13.12

Gases Dissolved in Liquids

Henry’s Law: Solubility of a gas is directly

proportional to the partial pressure of the gas above

the liquid.

Prob. 13.2

22 N)(NH,2 P k ][N

atm 0.78 atm 1 0.78 P 2N

mol/L10 x 5 atm 0.78 atm)mol/L10 x 7( ][N -4-42

Henry’s Law

S(gas) = kgas x Pgas

Page 500

Solubility of a

Gas:

[Gas] = kH PGas

Table 13.5

Solute Concentrations are Quantitative.

Concentration Terms

Other Related

Concentration Terms

Parts per Thousand (ppt)

Parts per Million (ppm)

Parts per Billion (ppb)

Mass Percent = masssolute/masssolution x 100%

Volume Percent = volumesolute/volumesolution x 100%

Mole Percent = molesolute/(molesolute + molesolvent) x 100%

Once we are able to quantify

solution composition, we can now

predict solution properties.

Prob. 13.51a & 13.3

13.51(a) Calculate the molarity of a solution of 0.82 g of ethanol

(C2H5OH) in 10.5 mL of solution.

C: 2 x 12

H: 6 x 1.0

O: 1 x 16

= 46 g/mol

M = molsolute/Lsolution

molethanol = 0.82 g / 46 g/mol = 1.8 x 10-2mol

Lsolution = 10.5 mL x (1L/1000mL) = 1.05 x 10-2L

M = 1.8 x 10-2mol/1.05 x 10-2L = 1.7 M

m = molsolute/kgsolvent = g/kg 1000

g 563

mol

g180

gglucose

= 2.40 x 10-2 m

gglucose = 2.43 g

Prob. 13.4

mass % = (masssolute / masssolution) x 100%

mass % = [35.0 g / (35.0 g + 150. g)] x 100%

mass % = 18.9%

Mole fraction = molsolute / molsolution = Xsolute

XPrOH = molPrOH / (molPrOH + molEtOH)

Prob. 13.5

M = molsolute/Lsolution

mass % = (masssolute / masssolution)) x 100%

m = molsolute / kgsolvent

X = molsolute / (molsolute + molsolvent)

Fig. 13.14

Strong Electrolyte Weak Electrolyte Non-electrolyte

Electrical Conductivity

Fig. 13.15

Vapor Pressure of Solutions

Raoult’s Law

PSolvent = Xsolvent PoSolvent

Fig. 13.16

Colligative Properties:

1. Vapor Pressure

Lowering

2. Freezing Point

Depression

3. Boiling Point Elevation

Colligative Properties

• Vapor Pressure Lowering

∆PSolvent = Xsolute PoSolvent

PSolvent = Xsolvent PoSolvent

• Boiling Point Elevation

∆Tb = kb msolute

• Freezing Point Depression

∆Tf = kf msolute

Prob. 13.6

∆PSolvent = Xsolute PoSolvent

solvent solute

solute solute

molmol

mol X

molsolute = 2.00 g / 180.15 g/mol & molsolvent = 50.0 g / 32.0 g/mol

Xsolute= 7.06 x 10-3

∆PSolvent = 7.06 x 10-3 101 torr

∆PSolvent = 0.713 torr

Table 13.6

Examples of BP Elevation & FP Depression Constants

Prob. 13.7

Freezing Point Depression: Tf = msolute • kf, H2O

msolute = Tf / kf, H2O

Tf = 32.0oF – 0.00oF = 32.0oF • (5oC/9oF) = 17.8oC

msolute = 17.8oC / (1.86oC/m)

msolute = 9.56 m

Fig. 13.17

Osmotic Pressure: Π = MRT {M = nsolute/Vsolvent}

Prob. 13.8b

Osmotic Pressure: Π = MRT

Π = 0.30 M • 0.0821 (L•atm)/(K•mol) • (37 + 273)K

Π = 7.6 atm

Fig. 13.21

Colloids

Particles >> wavelength of light scatters the light

True

Solution

Colloidal

Dispersion

Tyndall Effect

Brownian Motion

Table 13.7

Colloidal Dispersions

Will not settle out by gravity. Remain suspended.

Fig. 13.18