Chapter 14 – Mixtures and Solutions 14.1Types of Mixtures 14.2 Solution Concentration 14.3Factors Affecting Solvation 14.4Colligative Properties of Solutions

Chapter 14 – Mixtures and Solutions14.1 Types of Mixtures14.2 Solution Concentration14.3 Factors Affecting Solvation14.4 Colligative Properties of Solutions

Section 14.1 Types of Mixtures

• Compare the properties of suspensions, colloids, and solutions and the methods used to separate out their particles.

• Identify types of colloids (particularly sols, emulsions and both types of aerosols) and types of solutions.

• Describe the electrostatic forces in colloids and their role in colloid stabilization.

• Describe and give examples of shear-thickening and shear-thinning suspensions or colloids.

Mixtures can be either heterogeneous or homogeneous


• Define and give examples of Brownian motion and the Tyndall effect.

• Describe the role of the emulsifying agent in forming an emulsion.

Section 14.1 Types of MixturesKey Concepts

• The individual substances in a heterogeneous mixture remain distinct.

• Two types of heterogeneous mixtures are suspensions and colloids; these differ in the sizes of the particles.

• Brownian motion is the erratic movement of colloid particles due to random impacts of the dispersing medium.

• Colloids and suspensions exhibit the Tyndall effect.

• A solution can exist as a gas, a liquid, or a solid, depending on the solvent.


Key Concepts• Solutes in a solution can be gases, liquids, or solids.

• Salt solutions (electrolytes) can destabilize colloids and cause the particles to aggregate (stick together) and then flocculate (settle out).

• When emulsions are formed, and emulsifying agent can stabilize it because it contains both polar and non-polar regions which allow the agent to collect on the particle surface and interact with the dispersing medium.


Key Concepts

• The rheological (flow) properties of suspensions and colloids are more complicated than that of pure liquids or solutions. This is because the particles interact with each other in a way that can alter the flow.

• Some suspensions and colloids are shear thinning – their viscosity drops as they are stirred or shaken. Categories of shear thinning materials include pseudoplastic, and thixotropic materials.

• Some suspensions and colloids are shear thickening – their viscosity increases as they are stirred or shaken. The major shear thickening category is dilatant materials.

Heterogeneous Mixtures

Non-uniform composition; individual substances remain distinct

Suspension – heterogeneous mixture containing particles that settle out if left undisturbed

Colloid - heterogeneous mixture of intermediate sized particles (between 1 nm and 1000 nm) that do not settle out

A mixture containing particles that settle out upon standing

• Cornstarch in water• Muddy water

Can be filtered out by simple filters

>1000 nm (1 micron, ) diameter • Large compared to solvent molecules• Gravity can act on particles

Suspensions

Particle Model for

Suspension

Suspension – Cornstarch in Water

Settling

Filtration

Filtration

In past, only suspensions could be successfully filtered

No longer true – lots of filtration technology now available

Intermediate size between solution particles (molecular/atomic/ionic) and suspensions (1000 nm and up)1 nm to 1000 nm particle sizeParticles in dispersing medium (not solvent)Milk – particles can’t be filtered by ordinary filtration but can by ultra methods

Heterogeneous Mixtures - Colloids

Properties of Solutions, Colloids, & Suspensions

Property Solution Colloid Suspension

Particle Size 0.1-1.0 nm

1-1000 nm >1000 nm

Settles on Standing?

No No Yes

Filter with Paper? No No Yes

Separate by Dialysis?

No Yes Yes

Homogeneous? Yes Borderline No

Category Medium Particles Example

Solid Aerosol G S Smoke, dust in air

Liquid Aerosol G L Fog, cloud

Foam L G Whipped cream

Emulsion L L Mayonnaise, milk

Sol L S Blood, paint, milk of magnesia, mud

Solid Foam(Aerogel)

S G Marshmallow, exp. polystyrene

Solid Emulsion(Gel)

S L Butter, cheese, jello, jelly

Solid sol S S Colored Gems

Separation

of Solution,

Colloid, and Suspension

Suspension versus Colloid

Suspended SiO2 (sand) settles very quickly.

Each colloidal particle of SiO2 (Ludox®) attains (–) charge, which

repels other colloidal particles

Food EmulsionsOil and water don’t mix!

However, many foods exist that are emulsions made by mixing oil & water: salad dressing, butter, ice cream, & milk

Emulsion of liquid butterfat globules dispersed within a water-based liquid

Emulsions (type of colloid)Liquids that do not normally mix are spread throughout each other

Emulsify by slowly adding one ingredient to another while mixing

Process disperses and suspends tiny droplets of one liquid (dispersed phase) through another (continuous phase)

Oil

Water

ContinuousPhase

DispersedPhase

http://upload.wikimedia.org/wikipedia/commons/c/ca/Emulsions.svg

Emulsions & Emulsifying Agents

Emulsifying agents (emulsifiers) are materials that help stabilize an emulsion

Surfactants (surface active agents) are a class of emulsifying agents

Detergents are a category of surfactants

Surfactants and detergents are amiphiphilic – they are partly hydrophilic (polar) and partly hydrophobic (non-polar)

Emulsifier (Emulsifying Agent)Prevents mixture from separating

Attracted to both oil & water, allowing both to mix

Functions by surrounding oil droplets to form protective coat - holds oil droplets in suspension

One end (polar) of emulsifier molecule soluble in water; other end (nonpolar) soluble in oil

A to C: no emulsifier

D: with emulsifier

Oil

Water

Non-polarendPolar

end

http://upload.wikimedia.org/wikipedia/commons/c/ca/Emulsions.svg

Mayonnaise - Emulsifier

Egg yolk contains lecithin (phosphatidyl choline) - naturally occurring emulsifier

Lecithin functions by surrounding oil droplets to form protective coat which holds oil droplets in suspension

Polar End

Non-Polar End

http://upload.wikimedia.org/wikipedia/commons/c/c0/Phosphatidylcholine.png

Soap and MicellesSoap molecule has hydrophilic (polar/ionic) “head” and hydrophobic (nonpolar) “tail”

Head attracted to water, rest isn’t

Nonpolar tails tend to coagulate together to form spherical structure called a micelle

Some emulsions have dispersed phase in form of micelles

Model of Sodium Lauryl Sulfate

Micelle Formation in Na Stearate

Stearate Micelle in Water Solution

Polar heads of micelles attract them to water, and simultaneously repel other micelles – won’t coalesce and settle out

Soap Action

Cleansing Action of Soap

Cleansing Action of Soap

FoamsColloid in which very tiny particles of gas are dispersed in liquid or solid substance

Examples: ice cream, whipped cream, foamed milk, marshmallows, and beaten egg white

Gels – Liquids Dispersed in Solids

More-or-less rigid systems

Solid particles form network structure which traps liquid; alginate, starch, gelatin & pectin are typical gelling agents

Often formed by proteins of eggs or flour in puddings, batters, and doughs

Gelatin - type of protein found in bone and skin tissue of animals

Colloidal particles have erratic movement caused by collisions of molecules of the dispersing medium

Brownian Motion

Dispersed particles don’t settle due to electrostatic interactions between medium and particle

Can break colloid structure by adding an electrolyte (alters interactions)

Colloid Stabilization

Colloidal Particles with Cation ShellColloid Stabilized by Electrostatic Interaction

Like charges repel – keep particles separated

Unstable colloidal dispersions form flocs as particles aggregate due to interparticle attractions

Flocculation can be induced by removal of electrostatic barrier that prevents particle aggregation by addition of salt (electrolyte) or by changing pH

The addition neutralizes or "screens" surface charge of particle - removes repulsive forces that keep colloidal particles separate and allows for coagulation due to van der Waals forces

Colloid Destablilization - Flocculation

Formation and Coagulationof ColloidWhen strong electrolyte

added to colloidal iron oxide, charge on surface

of each particle is partially neutralized …

… and colloidal particles coalesce into suspension

that quickly settles

Colloidal particle large enough to scatter light

Suspensions also scatter (tend to be more opaque )

Solution particles are too small to scatter

Tyndall Effect

Tyndall Effect – Light scattering

Solution Colloid

Rheology of Colloids & Suspensions

Because these are heterogeneous mixtures, the flow properties of these materials can be unusual, especially for solid particles dispersed/suspended in a liquid medium

Flow properties are the subject of the field of study called rheology

RheologyStudy of the flow of matter, primarily in liquid state, but also as 'soft solids' or solids under conditions in which they respond with plastic flow rather than deforming elastically in response to an applied force.

Applies to substances which have a complex structure, such as muds, sludges, suspensions, polymers and other glass formers, as well as many foods and additives, bodily fluids (e.g. blood) and other biological materials.

http://en.wikipedia.org/wiki/Rheology

Rheology – Shear StressForce applied to overcome fluid’s resistance to flow, divided by area that force is acting upon (units of pressure but unlike pressure, force is parallel to surface)

A

F

shear stress, N/m2 = PaF force applied, N

A surface area subjected to stress, m2

Rheology – Shear RateRelative velocity of fluid layers divided by their normal separation distance (fluid not moving next to bottom plate)

d

U

shear rate, sec-1 U velocity, m/sec

d plate distance, m

Rheology - ViscosityRatio of shear stress to shear rate; indicates resistance material has to a change in form

Property can be thought of as an internal friction (measure of how difficult it is to slide one layer of fluid past an adjacent layer)

viscosity, Pa s shear stress, Pa shear rate, sec-1

Rheology – Newtonian FluidNewtonian fluid has a fixed viscosity that is independent of time and of shear stress

= shear stress/shear rate

constant slope

Rheology – Non-Newtonian Fluids

For our purposes, will focus on following types of non-Newtonian fluids:

Dilatant: shear-thickening

Pseudoplastic: shear-thinning

Thixotropic: time-dependent pseudoplastic flow; at constant applied shear rate, viscosity gradually decreases

Non-Newtonian

Liquids

Shear-Thickening Behavior

Dilatant mixture – “liquid-like” mixture that solidifies when shaken or stirred

Alternative statement: viscosity increases with rate of shear strain

All are stabilized suspensions with high volume fraction of solid

Examples: Silly Putty, oobleck (corn starch & water), silica in polyethylene glycol (liquid body armor), quicksand

http://en.wikipedia.org/wiki/Dilatant

Shear-Thickening Dilatant

Shear-Thinning Behavior

Psuedoplastic & Thixotropic mixtures –thick or solid-like mixture that flows more easily when stirred or agitated (shear is applied); thixotropic is time-dependence while psuedoplastic is not (won’t worry about distinction for examples)

Examples: toothpaste, certain paints and inks, catsup, clay soils that liquefy in an earthquake, mayonaisse

Shear-Thinning Psuedoplastic

Rheological Properties

Why Shear Thinning Occurs

Random coilpolymers elongate & break

Unsheared

Anisotropic Particles alignwith Flow Streamlines

Sheared

Aggregatesbreak up

Homogeneous Mixture - Solution

Solution Characteristics

Homogeneous mixtures

Solute = substance that dissolves• minor component

Solvent = dissolving medium• major component

By visual examination (even under magnification), not possible to distinguish solute from solvent

Types and Examples of SolutionsPhase Solute Solvent Example

G Gas Gas Air – N2, O2, Ar

L Gas Liquid Soda – CO2 / H2O

L Liquid Liquid Antifreeze Glycol / H2O

L Solid Liquid Salt water

S Gas Solid H2 in Pd or Pt

S Liquid Solid Dental Amalgam (Hg/Ag)

S Solid Solid Brass (Cu/Zn alloy)

Solution Terminology

Soluble• Substance dissolves in solvent• If solvent and solute are liquids, also

said to be miscible

Insoluble• Substance doesn’t dissolve• If solvent and solute are liquids, also

said to be immiscible - oil and water

Chapter 14 – Mixtures and Solutions

14.1 Types of Mixtures14.2 Solution Concentration14.3 Factors Affecting Solvation14.4 Colligative Properties of Solutions

Section 14.2 Solution Concentration

• Describe concentration using different units.

• Determine the concentrations of solutions.

• Calculate the molarity of a solution.

Concentration can be expressed in terms of percent or in terms of moles.

Section 14.2 Solution Concentration

Key Concepts• Concentrations can be measured qualitatively and

quantitatively.

• Molarity is the number of moles of solute dissolved per liter of solution.

• Molality is the ratio of the number of moles of solute dissolved in 1 kg of solvent.

• The number of moles of solute does not change during a dilution.

M1V1 = M2V2

Measure of how much solute is dissolved in a specific amount of solvent or solution

General terms• Dilute – small amount of solute• Concentrated – large amount of solute

Solution Concentration

See table 14-3, page 480

Percent by mass

Percent by volume

Molarity

Molality

Mole Fraction

ppm (not in book)

Concentration Units Used

% by mass = 100 mass of solute mass of solution

Example problem 14.1, page 481

Dissolve 3.6 g NaCl in 100.0 g H2O

% by mass = 100 3.6 103.6

= 3.5 %

Percent by Mass

Have 14.0% by mass KI solution in H2ODensity = d = 1.208 g/mLGrams KI in 25.0 mL of solution?Answerd = m/V m = V dm = 25.0 mL 1.208 g solution/mL m = 30.2 g solution0.140 g KI/g solution 30.2 g solution = 4.23 g KI

Percent by Mass Using Density

Practice (% by mass)

Problems 9 – 12, page 481


Problems 1 – 2, 7, page 985

Volumetric Measurements

Variety of laboratory containers available in which volume can be measured

Need to choose one appropriate for task at hand

Beaker

Volume graduations should be used only for "ballpark" estimates

Erlenmeyer flaskShape constructed to facilitate swirling or mixing of reactantsNOT precise-250 mL flask typically holds ~ 270 mL Use only for approximate measurements

Volumetric FlasksUsed to prepare solutions of exact concentrations of solutionsHave precise graduation line in neck of flask

Volumetric & Graduated PipettesUse bulb/pipette pump/pipette filler to fillConvenient way to accurately transfer small volume of liquidVolumetric pipette -used to measure one volume onlyGraduated pipette - lots of lines so you can measure many different amounts

Volumetric Glassware

Accurate/Precise Volumetric Glassware

Burette (or Buret)

Volumetric Flask

Volumetric (a) and Graduated (b) Pipettes (or pipets)

Beaker

Graduated Cylinder

% by volume = 100 volume of solute volume of solutionAlthough book and your own experience tell you that

Vsolution = Vsolute + Vsolvent

Additivity of volumes does not always occur!

For purposes of doing calculations in this course, we will assume that volumes add, but won’t when describe lab methods of preparation of solutions

Percent by Volume

Volumes Don’t Always Add

50 mL of

ethanol …

… and 50 mL of water …

… when mixed,

give less than 100

mL of solution

Preparation of 10% Ethanol Solution by

Volume

10.00 mL C2H5OH

100.0 mL

Volumetric pipette

Volumetric flask

Dispense ethanol into flask, slowly add water with swirling

When well mixed, top off to mark with water

Practice (% by volume)




Molarity = M = moles of solute liters of solution

Sample problem 14.2 page 483

100.5 mL of solution contains 5.10 g glucose (180.16 g/mol) Molarity?

5.10 g gluc. 2.83x10-2 mol gluc.

M = 2.83x10-2 mol/ 0.1005 L = 0.282 M

Molarity

Practice (molarity)




Problems 5, 6, 8, 9, page 985

Don’t Know Volume of

H2O Required

1.0 L water initially available to make 1.0 L of solution

25.0 g (0.100 mol) CuSO45H2O

Water left over

1.0 L of 0.100 M CuSO4

For 1.0 M aqueous solution, cannot add 1.0 mol of compound to 1.0 L of H2O

Final volume of solution not guaranteed to be 1.0 L using this method

Need to use volumetric flask• Weigh out solute• Add to volumetric• Fill volumetric to mark

Preparing Molar Solutions

Preparing Molar Solutions Using Volumetric Flask

Weigh Place in flask Fill to mark

Preparing Solutions of Known Concentration

(Adds Solute in Dry Form to Flask)

Prepare less concentrated solution from more concentrated one by adding solvent

Diluting Solutions

Often want to use concentrated (stock) solution to prepare a precisely known concentration of a dilute solution

When diluting, add only solvent – moles of solute remains constant

mol solute = M1 V1 = M2 V2

1 = concentrated (stock) solution

2 = dilute solution

Diluting Solutions

Making Solution by Dilution

M1 = 12.0 M V1 = ? LM2 = 1.50 M V2 = 5.00 L

M1 V1 = M2 V2

V1 = M2 V2 / M1

= 1.50 M 5.00 L / 12.0 MV1 = 0.625 L

Dilute 0.625 L of 12.0 M solution to 5.00 L

Make 5.00 L of 1.50 M KCl from 12.0 M Stock


Make 0.50 L of 0.300M CaCl2 from 2.00M stock CaCl2. mL stock needed?

M1 V1 = M2 V2

V1 = V2 M2

M1

V1 = 0.50L 0.300M 2.00M

V1 = 0.075 L = 75 mL

Diluting Solutions

Practice (dilutions)




Molality = m = moles of solute kg solventNote: only concentration unit that has solvent rather than solution in denominator; also in kg, not g

Prob. 14.4 4.5 g NaCl + 100.0 g H2OMolality?4.5 g NaCl 0.077 mol NaClm = 0.077 mol NaCl /0.1000 kg = 0.77m

Molality

Practice (molality)




Mole fraction = XA = moles of cpd A total moles in solution

= nA/(nA + nB)

A = solvent or solute cpd = compound

Mole Fraction

XA & XB represent mole fractions of

substances A and B (must add up to 1)

nB

Practice (mole fraction)


Problems 83 – 85, pages 508-9


Not discussed in book but widely used for concentrations in water, especially with regard to environmental issues

ppm = g solute x 106

g solution

ppm, ppb, etc

Parts per million (ppm): often used to express concentration of very dilute aqueous solutions. “Technical" definition is:

ppm = g solute x 106

g solution

Since amount of solute relative to amount of solvent is typically very small, density of solution is ~ same as density of solvent (water). For this reason, parts per million also expressed as:

ppm = mg solute = mg solute kg solution L solution

ppm, ppb, etc



Section 14.3 Factors Affecting Solvation

• Describe how intermolecular forces affect solvation.

• Define solubility and the terms unsaturated, saturated, supersaturated, dilute and concentrated.

• List the factors affect solubility and be able to predict if a certain material will readily dissolve in another.

Factors such as temperature, pressure, and polarity affect the formation of solutions.


• Use Henry’s law to calculate gas solubility.


Key Concepts

• The process of solvation involves solute particles surrounded by solvent particles. Strong forces will occur between solute and solvent when both are of the same “type” (polar or non-polar).

• Solutions can be unsaturated, saturated, or supersaturated.

• Henry’s law states that at a given temperature, the solubility (S) of a gas in a liquid is directly proportional to the pressure (P) of the gas above the liquid.

Solution Terminology

Solvation – process of surrounding solute particles with solvent particles

• Attractive intermolecular forces between solute and solvent must exist for this to happen

• In water, solvation process called hydration

Intermolecular Forces in Solution

For a solute to dissolve,

strength of solvent–solvent

forces …

… and solute–solute

forces …

… must be comparable to solute–solvent

forces

Intermolecular InteractionsAll Can Be Involved in Forming Solutions

Ion DipoleHydrogen Bond Dipole-Dipole

Ion – Induced DipoleDipole – Induced

Dipole Dispersion

Solubility Rule

“Like dissolves like” general rule for determining solubility

Polar substances (especially water) dissolve other polar substances and ionic compounds

Nonpolar substances (especially hydrocarbons) dissolve other nonpolar substances

Typical Nonpolar Hydrocarbon (C18)

Immiscible Oil and Water

Nonpolar oil (hydrocarbon)

layer

Polar water layer

Like Dissolves Like

Solubility of Alcohols in H2O & Hexane

0.12

0.12

0.030

0.0058

Aqueous Solutions - Ionic Compounds

Water strongly polar

Water molecules in constant motion (KMT) and collide with surfaces

Ionic solids in water tend to dissolve due to strong ion – dipole interactions

• Must be stronger than attraction among ions on crystal surface for dissolution to occur

Ion–Dipole Forces in Dissolution

Positive ends of dipoles attracted to

anions

Negative ends of dipoles attracted to

cations

Not all ionic compounds dissolve in water

• CaCO3 (calcium carbonate)

• CaSO42H2O (gypsum)

These compound have stronger forces between ions than between ions and the polar water molecules

Aqueous Solutions - Ionic Compounds

Water is a good solvent for molecules with polar groups like OH

(sucrose)

Aqueous SolutionsMolecular Compounds

Dissolving Sugar Cube in Water

Solvation process energy requirements• Solute particles separate (endothermic)• Solvent particles move apart

(endothermic)• Solute-solvent particles interact

(exothermic)

Overall energy change: heat of solution

Can be either endo- or exo-thermic

Heat of Solution

Energetics of Solution Formation

Solvation only occurs at point of contact between solvent and solute

Can accelerate process by• Agitation (stirring)• Increasing surface area of solute• Raising temperature

CaCO3 less soluble in hot water than in cold water but probably still reaches its solubility limit faster than at lower T

Factors That Affect Solvation Rate

Maximum amount of solute that can dissolve in a given amount of solvent at a specified temperature and pressure

One common unit is g solute per 100 g solvent

Refer to table 14.4, page 494 for solubilities of aqueous solutions of ionic compounds, molecular compounds and gases (later slide will show)

Solubility

Solubility limit occurs due to dynamic equilibrium process (similar to vapor pressure)

Solvated solute particles collide with undissolved solute – some stick

Rates of dissolution (solvation) and crystallization become equal at saturation

Maximum Solubility (saturation)

Solution Formation and Equilibrium

Saturated: dynamic

equilibrium

Unsaturated:

All solvation

Unsaturated:

Some crystallization

Unsaturated solution - more solute can be dissolved

Saturated – at solubility limit

Supersaturated – temporarily have more solute dissolved than the known solubility at a given temperature

• Unstable state

Solubility

Solubility –Dynamic Equilibrium

SolvationCrystallization

Many (but not all) substances have increased solublility at higher temperatures

Effect of Temperature on Solubility

Aqueous Solubilities – Table 14.4

Increase with T

Decrease with TDecrease with T - gasesLittle change with T

Solubility Variation with Temperature

Variation of Solubility of

Ionic Compounds

with Temperature(large y scale)

Can form by cooling a saturated solution (if solute has increased solubility with temperature)

Seed crystal, shock, rough surface (scratch) can initiate rapid crystallization

AgI crystals used to trigger rain in air supersaturated with water vapor

Supersaturated Solutions

Supersaturated Sodium Acetate

Gas Solubilities In Water

Note that these generally decrease with increasing temperature

(opposite in organic solvents)

Pressure and Solubility of Gases

… thus more frequent collisions of gas molecules with surface …

… giving higher concentration of dissolved gas

Higher partial pressure means more molecules of gas per unit volume …

Gases more soluble at higher pressure

Henry’s Law

S1/P1 = S2/P2

Si = solubility of gas

Pi = partial pressure of gas above solution

S2 = S1 (P2/P1)


Solubility of Gases with Pressure


Lowering pressure by unscrewing bottle cap lowers CO2 solubility and excess escapes

Warming also lowers solubility


Increased pressure underwater causes N2 to dissolve in higher concentration in blood – can cause bends when diver surfaces


0.85 g of gas dissolves at 4.0 atm in 1.0 L at 25 CHow much dissolves in 1.0 L at 1.0 atm at same T?

S1 = 0.85 g/L P1 = 4.0 atm P2 = 1.0 atm

S1/P1 = S2/P2 S2 = S1 (P2/P1)

S2 = 0.85 g/L (1.0 atm / 4.0 atm)

S2 = 0.21 g/L

Pressure and Solubility – Henry’s Law

Practice (Henry’s Law)


Problems 89, 90, 92 – 94, page 509




Section 14.4 Colligative Properties of Solutions

• Describe colligative properties.

• Identify four colligative properties of solutions.

• Calculate the boiling point elevation and freezing point depression of a solution.

• Explain the meaning of the terms osmotic pressure, dialysis and reverse osmosis and be able to give an example.

Colligative properties depend on the number of solute particles in a solution.

Key Concepts

• Nonvolatile solutes lower the vapor pressure of a solution.

• Boiling point elevation is directly related to the solution’s molality.

∆Tb = Kbm• A solution’s freezing point depression is always lower

than that of the pure solvent.

∆Tf = Kfm• Osmotic pressure depends on the number of solute

particles in a given volume.


Key Concepts

• Dialysis is a purification process that employs the same type of semipermeable membrane that is involved in the phenomenon of osmotic pressure. Artificial kidney dialysis is one example of its use.

• Applying pressure to the side opposite the pure solvent side of a semipermeable membrane is a purification process called reverse osmosis. It has many uses, among them being the desalinization of seawater to generate fresh water for drinking.


Physical properties of solutions that depend upon number of particles and not identity of dissolved solute particles

Colligative Properties

Electrolytes completely ionize

NaCl(s) Na+(aq) + Cl-(aq)

1m solution yields 2m ions

Nonelectrolytes dissolve but don’t ionize

Sucrose (table sugar)

1m solution yields 1m dissolved particles

Can distinguish based on conductivity

Electrolytes & Nonelectrolytes

Electrolytes & Nonelectrolytes

Nonelectrolyte Electrolyte

Solution Conductivity Tests

Adding nonvolatile solute lowers VP of solvent

Mixture of solvent and solute occupies surface; fewer solvent particles enter vapor phase

The greater the number of solute particles, the lower the VP

Vapor Pressure (VP) Lowering

VP lowering due to number of solute particles and is a colligative property


H2OPure waterSucroseSugar & water

Water VP lowered

1 mole of nonelectrolyte lowers VP by given amount1 mole of electrolyte lowers VP by n , where n = # of particles produced

H3PO4 3H+ + PO43-

4 particles, 4 effect relative to nonelectrolyte


When nonvolatile solute lowers VP of solvent, BP also affectedMore heat needed to supply additional kinetic energy to raise VP to atmospheric pressure

Boiling Point (BP) Elevation

Vapor pressure effect directly related to BP

Tb = Kbm

Kb = BP elevation constant (units C/m)Difference in BP between pure solvent and 1m solution of nonelectrolyte

Except for number of particles, Tb does not depend on solute

Boiling Point (BP) Elevation

Boiling Point Elevation

Solvent Molecule

Nonvolatile Solute Molecule

Boiling Point Elevation

Solvent BP (C) Kb (C/m)

Water 100.0 0.512

Benzene 80.1 2.53

CCl4 76.7 5.03

Ethanol 78.5 1.22

Chloroform 61.7 3.63

Solute disrupts organization of crystal

Tf = Kfm

Kf = FP constant (units C/m)

Difference in FP between pure solvent and 1m solution of nonelectrolyte

1.86 C/m for water

Use of road salt a common application

Freezing Point (FP) Depression

Freezing Point Depression

Solvent FP (C) Kf (C/m)

Water 0.0 1.86

Benzene 5.5 5.12

CCl4 -23 29.8

Ethanol -114.1 1.99

Chloroform -63.5 4.68

FP & BP Changes – Effective m

Must use effective molality or particle molality in equations

1m NaCl solution = 2m effective

For 1m CaCl2 solution in water

BP elevation?

Effective molality = 3m

BP elevation = 3m 0.512 C/m

= 1.54 C

FP & BP Changes


BP and FP of 0.029 m aqueous soln. NaCl?

m = 0.029 x 2 = 0.058 m (2 particles)

Kb = 0.512 C/m Kf = 1.86 C/m

Tb=Kb m= 0.512 C/m 0.058 m = 0.030C

Tf = Kf m = 1.86 C/m 0.058 m = 0.11C

BP = 100.000 C + 0.030C = 100.030 C FP = 0.00 C 0.11C = 0.11C

Practice (FP & BP Changes)




FP & BP Changes – Phase Diagram

Osmosis – diffusion of solvent particles across a semipermeable membrane from region of low to region of high solute concentration

Membrane allows solvent and smaller size particles to pass, blocks larger ones

Membrane pore radius typically 0.1 to 100 nm (bond lengths 0.1 to 0.25 nm)

Osmosis

Semipermeable membrane separating water and aqueous

solution of glucose

Movement of solvent (water) from dilute to concentrated solution across membrane

Osmosis

If membrane blocks solute from passing, pressure difference develops

Pressure = osmotic pressure

Also a colligative property - depends on # particles

Osmotic pressure MM = solution molarity of solution on one side of membrane (pure solvent on other side)

Osmotic Pressure

Osmosis & Osmotic Pressure

Pure solvent

Pure solvent Solution

Solution

Semipermeable membrane

Osmotic pressure

Membrane allows solvent to pass but not solute

Osmosis and Osmotic Pressure

Net flow of water from

outside (pure H2O) to solution

Solution increases in

volume until … … height of solution exerts

osmotic pressure (π) of solution

Drinking Seawater will Cause Dehydration of Body Tissues

Water leaves cells through osmosis to dilute seawater in digestive tract

Dialysis - process that relies on concentration-driven transport (osmosis) for separation/purification

Common application: artificial kidney dialysis

Osmosis / Dialysis

Renal Dialysis – Artificial Kidney

Hemodialysis attempts to mimic the action of the nephron (kidney) -separates low MW solutes (e.g., urea, creatinine) from the blood of patients with chronic uremia

Renal Dialysis – Artificial KidneyKidney dialysis filters and removes waste products and excess water from blood. Inventor (Kollf, 1943, Holland) originally used cellophane as the membrane.

Principle of Dialysis

If apply pressure opposite that of direction of osmotic pressure, can purify solution

Reverse Osmosis

Reverse Osmosis

Pure solvent Solution

Semipermeable membrane

Pressure applied > osmotic pressure

Reverse Osmosis Water Purification

>30

END OF CHAPTER

Documents

Chapter 14 – Mixtures and Solutions 14.1Types of Mixtures 14.2 Solution Concentration 14.3Factors Affecting Solvation 14.4Colligative Properties of Solutions