Solutions © 2009, Prentice-Hall, Inc. Chapter 13 Properties of Solutions Chemistry, The Central Science, 11th edition Theodore L. Brown, H. Eugene LeMay,

Solutions

© 2009, Prentice-Hall, Inc.

Chapter 13Properties of Solutions

Chemistry, The Central Science, 11th editionTheodore L. Brown, H. Eugene LeMay, Jr.,

and Bruce E. Bursten

MC 9 out of 75

FRQ almost every year

Extension of Chapter 4

Thanks to John D. Bookstaver

Solutions

13.1 The Solution Process

• Solutions are homogeneous mixtures of two or more pure substances.– Formed when 1 substance (solute) disperses

uniformly throughout another (solvent).• Solute – present in lesser quantity• Solvent – present in greater quantity

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13.1 The Solution Process

Solutions can be mixes involving any state.

This mix always forms a solution!

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How Does a Solution Form?Solvation

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

If the solvent is water, this process is termed hydration.

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How Does a Solution Form?Dispersion

In the absence of a strong solvent-solute attraction, the particles randomly spread out or Disperse.

This is typical for nonpolar solutions!

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Why do Solutions Form?

The ability of substances to form solutions depends on 2 factors:

1. Type of intermolecular interactions involved

2. Natural tendency of substances to spread into larger volumes when not restrained

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The Effect of Intermolecular Forces

There are 3 forces at work during solvation.1. Break Solvent-Solvent attraction.

2. Break solute-solute attraction.

3. Make solvent-solute attraction

If #3 is comparable to, or greater, than #1 and #2, than a

solutions forms!

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Break Solvent – Solvent Forces

H to N,O, or FHydrogen bonding

Dipole-dipole

London or

Dispersion(induced dipoles)

Polar Solvent

Nonpolar solvent

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Break Solute – Solute Forces

H to N,O, or FHydrogen bonding

Dipole-dipole

London or

Dispersion(induced dipoles)

Polar Molecules

Nonpolar Molecules

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Make Solute – Solvent Forces

Strong

Weak

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NaCl and Water

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

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NaCl and C6H14

Solute-soluteionic

(strong)

Solvent-solventLondon dispersion

(weak)

Solute-solventInduced dipole

(weak)

Insoluble!

The solute-solvent force is too weak to overcome the ionic

bonding!

+ >

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Energy Changes and Solution Formation

• Simply put, three processes affect the energetics of solution:– separation of solute particles,

• endothermic

– separation of solvent particles,• endothermic

– new interactions between solute and solvent.• exothermic

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Energy Changes in Solution

The enthalpy change of the overall

process depends on H for each of these

steps.

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Energy Changes in Solution

Things do not tend to occur spontaneously (i.e., without outside intervention) unless

the energy of the system is lowered.

(overall exothermic)

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Why Do Endothermic Processes Occur?

Yet we know that in some processes,

like the dissolution of NH4NO3 in water, heat is absorbed,

not released.

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Solution Formation, Spontaneity, and Entropy

This is where the second factor effecting solutions is involves!

2. Natural tendency of substances to spread into larger volumes when not restrained

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Solution Formation, Spontaneity, and Entropy

Dissolving, when it occurs, is a spontaneous process!

If you remember from chemistry, reactions that increase the entropy (disorder and randomness) tend to be spontaneous.

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Enthalpy Is Only Part of the Picture

The reason is that increasing the disorder or randomness (known as

entropy) of a system tends to lower the energy

of the system.

Diffusion is therefore a

spontaneous reaction

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Enthalpy Is Only Part of the Picture

So even though enthalpy may increase

(endothermic), the overall energy of the

system can still decrease if the system

becomes more disordered.

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Silver chloride is essentially insoluble in water.

Would you expect a significant change in entropy of the system when 10 g of AgCl is added to 100 g of water?

NO!

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Solution Formation and Chemical Reactions

Just because a substance disappears when it comes in contact with a solvent, it doesn’t

mean the substance dissolved.

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Solution Formation and Chemical Reactions

• Dissolution is a physical change — you can get back the original solute by evaporating the solvent.

• If you can’t, the substance didn’t dissolve, it reacted.

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13.2 Types of Solutions

• Saturated– In a saturated solution,

the solvent holds as much solute as is possible at that temperature.

– Dissolved solute is in dynamic equilibrium with solid solute particles.

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13.2 Types of Solutions

• Unsaturated– If a solution is

unsaturated, less solute than can dissolve in the solvent at that temperature is dissolved in the solvent.

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13.2 Types of Solutions

• Supersaturated– In supersaturated solutions, the solvent holds

more solute than is normally possible at that temperature.

– These solutions are unstable; crystallization can usually be stimulated by adding a “seed crystal” or scratching the side of the flask.

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13.3 Factors Affecting Solubility

• Chemists use the axiom “like dissolves like."– Polar substances tend to dissolve in polar solvents.– Nonpolar substances tend to dissolve in nonpolar

solvents.

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Solute-Solvent Interactions

The more similar the intermolecular attractions, the more likely one substance is to be soluble in another.

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Solute-Solvent Interactions

Glucose (which has hydrogen bonding) is very soluble in water, while cyclohexane (which only has dispersion forces) is not.

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What would happen to the solubility of glucose if the hydroxyl groups (-OH)

were replaced with methyl groups (-CH3)?

Decrease, -CH3 is nonpolar

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Solute-Solvent Interactions

• Vitamin A is soluble in nonpolar compounds (like fats).

• Vitamin C is soluble in water.

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Pressure Effects

• In general, the solubility of gases in water increases with increasing mass.

• Larger molecules have stronger dispersion forces.

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Pressure Effects

• The solubility of liquids and solids does not change appreciably with pressure.

• The solubility of a gas in a liquid is directly proportional to its pressure.

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Henry’s Law

Sg = kPg

where• Sg is the solubility of

the gas,• k is the Henry’s Law

constant for that gas in that solvent, and

• Pg is the partial pressure of the gas above the liquid.

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Henry’s Law

2

2

1

1

P

S

P

S

How can this relationship be expressed?

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Henry’s Law Problem

Calculate the concentration of CO2 in a soft drink that is bottled with a partial pressure of CO2 at 4.0 atm over the liquid at 25 ºC if the solubility at 1.0 atm is 0.030 M?

atm 4.0

x

atm .01

0.030

M

x = 0.12 M

x = (4.0 atm)(0.030 M)/(1.0 atm)

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Temperature Effects

Generally, the solubility of solid solutes in liquid solvents increases with increasing temperature.

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Why do bubbles form on the inside wall of a cooking pot when the water is heated on the stove, even though the temperature is well below the boiling point of water?

Increasing the temperature of the water decreases the solubility of the dissolved gasses. These released gasses deposit

on the side of the container.

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Temperature Effects

• The opposite is true of gases.– Carbonated soft

drinks are more “bubbly” if stored in the refrigerator.

– Warm lakes have less O2 dissolved in them than cool lakes.

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13.4 Ways of Expressing Concentrations

• Qualitative Concentrations– Dilute

• Relatively small concentration– Concentrated

• Relatively large concentration

• Quantitative Concentration– Mass percent– ppm and ppb– Mole fraction– Molarity– Molality

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Mass Percentage

mass of A in solutiontotal mass of solution

100

What does a 4.0 % NaCl solution mean?

100solution g 100.0

NaCl g 4.0 % 4.0

4.0 g NaCl for every 96.0 g of water!

% mass is also known as “pph” (parts per hundred)

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Parts per Million andParts per Billion

ppm =mass of A in solutiontotal mass of solution

106

Parts per Million (ppm)

Parts per Billion (ppb)

ppb =mass of A in solutiontotal mass of solution

109

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moles of Atotal moles in solution

XA =

Mole Fraction (χ)

• In some applications, one needs the mole fraction of solvent, not solute — make sure you find the quantity you need!

• Xsolute + Xsolvent = 1

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mol of soluteL of solution

M =

Molarity (M)

• You will recall this concentration measure from Chapter 4.

• Since volume is temperature-dependent, molarity can change with temperature.

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mol of solutekg of solvent

m =

Molality (m)

Since both moles and mass do not change with temperature, molality (unlike molarity) is not temperature-dependent.

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Changing Molarity to Molality

If we know the density of the solution, we can calculate the molality from the molarity and vice versa. (p545)

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13.5 Colligative Properties

• Changes in colligative properties depend only on the number of solute particles present, not on the identity of the solute particles.

• Among colligative properties are– Vapor pressure lowering – Boiling point elevation– Melting point depression– Osmotic pressure

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Vapor Pressure

Because of solute-solvent intermolecular attraction, higher concentrations of nonvolatile solutes make it harder for solvent to escape to the vapor phase.

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Vapor Pressure

Therefore, the vapor pressure of a solution is lower than that of the pure solvent.

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Raoult’s Law

PA = APA

where– A is the mole fraction of compound A, and

– PA is the normal vapor pressure of A at that temperature.

NOTE: This is one of those times when you want to make sure you have the vapor pressure of the solvent.

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Boiling Point Elevation and Freezing Point Depression

Nonvolatile solute-solvent interactions also cause solutions to have higher boiling points and lower freezing points than the pure solvent.

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Boiling Point Elevation• The change in boiling

point is proportional to the molality of the solution:

Tb = Kb m

where Kb is the molal boiling point elevation constant, a property of the solvent.Tb is added to the boiling

point of the pure solvent.

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Freezing Point Depression• The change in freezing

point can be found similarly:

Tf = Kf m

• Here Kf is the molal freezing point depression constant of the solvent.

Tf is subtracted from the boiling point of the pure solvent.

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Boiling Point Elevation and Freezing Point Depression

Note that in both equations, T does not depend on what the solute is, but only on how many particles are dissolved.

Tb = Kb m

Tf = Kf m

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Colligative Properties of Electrolytes

Since these properties depend on the number of particles dissolved, solutions of electrolytes (which dissociate in solution) should show greater changes than those of nonelectrolytes.

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van’t Hoff Factor

• We use the van’t Hoff factor, i, to represent the number of particles we get when a substance dissolves.– Nonelectrolytes (molecules)

• i = 1

– Electrolytes (ionic compounds)• i = # of ions

– NaCl → i = 2– K2SO4 → i = 3

– MgSO4 → i = 2

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van’t Hoff Factor

• We modify the previous equations by multiplying by the van’t Hoff factor, i.

Tf = Kf m i

Tb = Kb m i

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van’t Hoff Factor

However, a 1M solution of NaCl does not show twice the change in freezing point that a 1M solution of methanol does.

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van’t Hoff Factor

One mole of NaCl in water does not really give rise to two moles of ions.

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van’t Hoff Factor

Some Na+ and Cl- reassociate for a short time, so the true concentration of particles is somewhat less than two times the concentration of NaCl.

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© 2009, Prentice-Hall, Inc.

van’t Hoff Factor

• Reassociation is more likely at higher concentration and this effectively lowers the i value.

• Therefore, the number of particles present is concentration-dependent.

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© 2009, Prentice-Hall, Inc.

van’t Hoff Factor

• In addition, higher charge values have greater attractive forces (coulomb's law!) and cause greater reassociation and this also lowers the actual value of i.

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Osmosis

• Some substances form semipermeable membranes, allowing some smaller particles to pass through, but blocking other larger particles.

• In biological systems, most semipermeable membranes allow water to pass through, but solutes are not free to do so.

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Osmosis

In osmosis, there is net movement of solvent from the area of higher solvent concentration (lower solute concentration) to the are of lower solvent concentration (higher solute concentration).

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Osmotic Pressure

The pressure required to stop osmosis, known as osmotic pressure, , is

nV

= ( )RT = MRT

where M is the molarity of the solution.

If the osmotic pressure is the same on both sides of a membrane (i.e., the concentrations are the same), the solutions are isotonic.

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Osmosis in Blood Cells

• If the solute concentration outside the cell is greater than that inside the cell, the solution is hypertonic.

• Water will flow out of the cell, and crenation results.

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Osmosis in Cells

• If the solute concentration outside the cell is less than that inside the cell, the solution is hypotonic.

• Water will flow into the cell, and hemolysis results.

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13.6 Colloids

Suspensions of particles larger than individual ions or molecules, but too small to be settled out by gravity are called colloids.

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Tyndall Effect

• Colloidal suspensions can scatter rays of light.

• This phenomenon is known as the Tyndall effect.

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Mixtures Review

Settle out?

Particle size

Tyndall Effect

Type of mixture

Suspension

Colloid

Solution

yes

no

no

large

intermediate

small

yes

yes

no

heterogeneous

heterogeneous

homogeneous

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Colloids in Biological Systems

Some molecules have a polar, hydrophilic (water-loving) end and a non-polar, hydrophobic (water-hating) end.

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Hydrophilic Colloid

• In this molecules the hydrophilic parts are on the surface end and the hydrophobic parts are folded in to make a hydrophilic colloid.

• Hydrophilic parts become attracted to water and partially hydrated.

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Hydrophilic Colloid

• In this molecules the hydrophilic parts are on the surface end and the hydrophobic parts are folded in to make a hydrophilic colloid.

• Hydrophilic parts become attracted to water and partially hydrated.

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How do hydrophilic colloid form?

1. If ions get adsorbed (stick to surface), it causes them to be attracted by water.

– Ions stabilize the colloid and allow limited hydration and prevents clumping nonpolar colloid by ion to ion repulsion.

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How do hydrophilic colloid form?

2. They can also be stabilized as a colloid by emulsifying agents (molecule with polar and nonpolar portions like soap)