Solutions

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© 2009, Prentice-Hall, Inc.

I. What is a solution?

I. What is a solution?

A. Terms A solution is a homogeneous mixture of two or more substances, consisting ions or molecules.

of

Solutions

I. What is a solution?

A. Terms 1. Solute/solvent

The solute is the substance that gets The solvent does the .

.

Solutions

I. What is a solution?

A. Terms 1. Solute/solvent solute solvent

The solute is the substance that gets dissolved.

The solvent does the dissolving.

Solutions

I. What is a solution?

A. Terms 1. Solute/solvent 2. Soluble/insoluble

A substance that is can be dissolved.

I. What is a solution?

A. Terms 1. Solute/solvent 2. Soluble/insoluble

A substance that is soluble can be dissolved.

I. What is a solution?

A. Terms 1. Solute/solvent 2. Soluble/insoluble 3. Miscible/ immiscible

Miscible substances are will dissolve in each other.

that

Solutions

I. What is a solution?

A. Terms 1. Solute/solvent 2. Soluble/insoluble 3. Miscible/ immiscible

Miscible substances are liquids that will dissolve in each other.

I. What is a solution?

A. Terms 1. Solute/solvent 2. Soluble/insoluble 3. Miscible/ immiscible 4. Electrolytes/non electrolytes

Electrolytes are liquids that are capable of conducting electricity.

© 2009, Prentice-Hall, Inc.

© 2009, Prentice-Hall, Inc.

I. What is a solution?

B. Types of solutions

I. What is a solution?

B. Types of solutions 1. Gaseous solutions

Air and other mixtures of gases

I. What is a solution?

B. Types of solutions 1. Gaseous solutions 2. Liquid solutions

In some cases water is the solvent. These are called ( ).

I. What is a solution?

B. Types of solutions 1. Gaseous solutions 2. Liquid solutions

In some cases water is the solvent. These are called (aqueous).

I. What is a solution?

B. Types of solutions 1. Gaseous solutions 2. Liquid solutions 3. Solid solutions

Combinations of solids. Some are homogenous mixtures of metals ( )

Solutions

I. What is a solution?

B. Types of solutions 1. Gaseous solutions 2. Liquid solutions 3. Solid solutions

Combinations of solids. Some are homogenous mixtures of metals (alloys)

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Types of Solutions

• Solutions may exist as

gases, liquids,

or

solids

. – The

solute

is the dissolved substance. In the case of a solution of a gas or solid in a liquid, it is the gas or solid. Otherwise, it is the component of lesser amount.

– The

solvent

is the dissolving medium. Generally it is the component of greater amount.

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Presentation of Lecture Outlines, 12 –20

Gaseous Solutions

• Nonreactive gases can mix in all proportions to give a

gaseous solution

.

– Fluids that dissolve in each other in all proportions are said to be

miscible fluids

.

– If two fluids do not mix, they are said to be

immiscible.

– For example, air is a solution of oxygen, nitrogen, and smaller amounts of other gases.

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Presentation of Lecture Outlines, 12 –21

Liquid Solutions

Liquid solutions

are the most common types of solutions found in the chemistry lab.

– Many inorganic compounds are soluble in water or other suitable solvents.

– Rates of chemical reactions increase when the likelihood of molecular collisions increases.

– This increase in molecular collisions is enhanced when molecules move freely in solution.

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Presentation of Lecture Outlines, 12 –22

Solid Solutions

Solid solutions

of metals are referred to as alloys.

– Brass is an alloy composed of copper and zinc.

– – Bronze is an alloy of copper and tin.

Pewter is an alloy of zinc and tin.

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Presentation of Lecture Outlines, 12 –23

I. What is a solution?

C. Making a solution

I. What is a solution?

C. Making a solution 1. Solute/solvent particles are separated.

Solutions

I. What is a solution?

C. Making a solution 1. Solute/solvent particles are separated.

2. Solvent particles are attracted to solute particles.

This is called solvation.

With the solvent this is called hydration

I. What is a solution?

C. Making a solution 1. Solute/solvent particles are separated.

2. Solvent particles are attracted to solute particles.

This is called solvation.

With the solvent water this is called hydration

I. What is a solution?

C. Making a solution 1. Solute/solvent particles are separated.

2. Solvent particles are attracted to solute particles.

3. ‘Like dissolves like’

Factors in Explaining Solubility

• In most cases,

“like dissolves like.”

– This means that polar solvents dissolve polar (or ionic) solutes and nonpolar solvents dissolve nonpolar solutes.

– The relative

force of attraction

of the solute for the solvent is a major factor in their solubility.

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Presentation of Lecture Outlines, 12 –29

Polar/Non-polar Substances

Polar

Water Methanol Ethanol Acetic acid Acetone Ionic compounds

Non-polar

Carbon tetrachloride Toluene (paint thinner) Hexane, heptane, octane Benzene Oil Solutions

Solutions

Solutions

Molecular Solutions

• Polar molecules interact well with polar solvents such as water. – The dipole-dipole interactions of water with a polar solvent can be easily explained as electrostatic attraction.

H

d +

H

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polar

d -

solute

d +

H

d +

H

Presentation of Lecture Outlines, 12 –33

Molecular Solutions

Nonpolar solutes interact with nonpolar solvents

– primarily due to London forces.

Heptane, C 7 H 16 , and octane, C 8 H 18 , are both nonpolar components of gasoline and are completely miscible liquids.

– However, for water to mix with gasoline, hydrogen bonds must be broken and replaced with weaker London forces between water and the gasoline.

– Therefore gasoline and water are nearly immiscible. Copyright © Houghton Mifflin Company.All rights reserved.

Presentation of Lecture Outlines, 12 –34

Ionic Solutions

d +

H H

Polar solvents,

such as water, also – interact well with

ionic solutes

.

Since ionic compounds are the extreme in polarity, we can illustrate the electrostatic attractions of water for cations and anions.

O

d -

+

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H

d +

H H H

d +

-

d +

H H

Presentation of Lecture Outlines, 12 –35

O

d -

I. What is a solution?

D. Dissolving Ionic Compounds in H 2 O 1. Water is polar because it has two oppositely ends

I. What is a solution?

D. Dissolving Ionic Compounds in H 2 O 1. Water is polar because it has two oppositely charged ends

I. What is a solution?

D. Dissolving Ionic Compounds in H 2 O 1. Water is polar because it has two oppositely charged ends 2. The end of water is attracted to the anion

Anion = negative ion

I. What is a solution?

D. Dissolving Ionic Compounds in H 2 O 1. Water is polar because it has two oppositely charged ends 2. The positive end of water is attracted to the anion

Anion = negative ion

I. What is a solution?

D. Dissolving Ionic Compounds in H 2 O 3. The combined of water molecules overcomes the attractive force of solute.

I. What is a solution?

D. Dissolving Ionic Compounds in H 2 O 3. The combined pull of water molecules overcomes the attractive force of solute.

I. What is a solution?

D. Dissolving Ionic Compounds in H 2 O 3. The combined pull of water molecules overcomes the attractive force of solute.

4. The ionic compound is dissociated into individual ions.

Ex. NaCl (s) Na + + Cl -

I. What is a solution?

E. Dissolving Molecular Compounds in H 2 O.

I. What is a solution?

E. Dissolving Molecular Compounds in H 2 O.

1. Water’s charged ends are attracted to charged portions of solute.

Solutions

I. What is a solution?

E. Dissolving Molecular Compounds in H 2 O.

1. Water’s charged ends are attracted to charged portions of solute.

2. The solute is surrounded by water molecules.

I. What is a solution?

E. Dissolving Molecular Compounds in H 2 O.

1. Water’s charged ends are attracted to charged portions of solute.

2. The solute is surrounded by water molecules.

3. The solute remains intact.

I. What is a solution?

F. Rate of Solvation To speed up solvation. . .

I. What is a solution?

F. Rate of Solvation To speed up solvation. . .

1. Stir/shake

I. What is a solution?

F. Rate of Solvation To speed up solvation. . .

1. Stir/shake 2. Break apart solute

I. What is a solution?

F. Rate of Solvation To speed up solvation. . .

1. Stir/shake 2. Break apart solute 3. Heat the solvent

II. Is it soluble?

II. Is it soluble?

A. Not all ionic compounds dissolve in water.

II. Is it soluble?

A. Not all ionic compounds dissolve in water.

1. Solubility Table

II. Is it soluble?

A. Not all ionic compounds dissolve in water.

1. Solubility Table 2. Writing physical states (in equations)

Insoluble = (s) for solid Soluble = ______________

II. Is it soluble?

A. Not all ionic compounds dissolve in water.

1. Solubility Table 2. Writing physical states (in equations)

Insoluble = (s) for solid Soluble = (aq) for aqueous

II. Is it soluble?

A. Not all ionic compounds dissolve in water.

1. Solubility Table 2. Writing physical states (in equations) 3. Solubility Rules

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Self Check – Ex. 1 Is copper (II) chloride soluble?

Self Check – Ex. 2 Is zinc carbonate soluble?

Self Check – Ex. 3 Is calcium hydroxide soluble?

II. Is it soluble?

A. Not all ionic compounds dissolve in water.

B. In reality, ionic compounds all dissolve in varying degrees.

Some compounds are very soluble while others are slightly soluble

II. Is it soluble?

A. Not all ionic compounds dissolve in water.

B. In reality, ionic compounds all dissolve in varying degrees.

1. Solubility depends on ___________

Solubility Curve

II. Is it soluble?

A. Not all ionic compounds dissolve in water.

B. In reality, ionic compounds all dissolve in varying degrees.

1. Solubility depends on temperature 2. Solubility depends on physical state

II. Is it soluble?

A. Not all ionic compounds dissolve in water.

B. In reality, ionic compounds all dissolve in varying degrees.

1. Solubility depends on temperature 2. Solubility depends on physical state 3. Solubility of gas depends on pressure

III. Heat of solution

III. Heat of solution

A. Many reactions warm up when you mix substances

Exothermic – heat given off

III. Heat of solution

A. Many reactions warm up when you mix substances B. Some reactions get colder when you mix substances

Endothermic – heat absorbed

IV. Concentration

IV. Concentration

A. Terms

IV. Concentration

A. Terms 1. Concentrated

A solution that has a lot of per solvent

IV. Concentration

A. Terms 1. Concentrated

A solution that has a lot of solute per solvent

IV. Concentration

A. Terms 1. Concentrated 2. Dilute

A solution that has a lot of .

IV. Concentration

A. Terms 1. Concentrated 2. Dilute

A solution that has a lot of solvent.

IV. Concentration

A. Terms 1. Concentrated 2. Dilute 3. Saturated

A solution with solute.

IV. Concentration

A. Terms 1. Concentrated 2. Dilute 3. Saturated

A solution with the maximum amount of solute.

IV. Concentration

A. Terms 1. Concentrated 2. Dilute 3. Saturated 4. Supersaturated

A solution with .

IV. Concentration

A. Terms 1. Concentrated 2. Dilute 3. Saturated 4. Supersaturated

A solution with more solute than it can hold.

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Solubility and the Solution Process

• The amount of a substance that will dissolve in a solvent is referred to as its

solubility

.

– Many factors affect solubility, such as

temperature

and, in some cases,

pressure

.

– There is a limit as to how much of a given solute will dissolve at a given temperature.

– A

saturated solution

is one holding as much solute as is allowed at a stated temperature. Copyright © Houghton Mifflin Company.All rights reserved.

Presentation of Lecture Outlines, 12 –80

Solubility: Saturated Solutions

• Sometimes it is possible to obtain a

supersaturated solution

, that is, one that contains more solute than is allowed at a given temperature.

– Supersaturated solutions are unstable.

– If a small crystal of the solute is added to a supersaturated solution, the excess immediately crystallizes out.

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Presentation of Lecture Outlines, 12 –81

IV. Concentration

A. Terms B. Measurements 1. Percentage 2. Molarity 3. Molality 4. Mole fraction

Colligative Properties • Consider three beakers: – 50.0 g of ice – 50.0 g of ice + 0.15 moles NaCl – 50.0 g of ice + 0.15 moles sugar (sucrose) • What will the temperature of each beaker be?

– Beaker 1: – Beaker 2: – Beaker 3:

Colligative Properties • The reduction of the freezing point of a substance is an example of a colligative property : – A property of a solvent that depends on the total number of solute particles present • There are four colligative properties to consider: – Vapor pressure lowering (Raoult ’ s Law) – Freezing point depression – Boiling point elevation – Osmotic pressure

Colligative: particles are particles • Colligative comes from colligate – to tie together • Colligative properties have common origin • Colligative properties depend on

amount

solute but do

not

depend on its chemical of identity • Solute particles exert their effect merely by being rather than doing • The effect is the same for all solutes

Colligative properties for nonvolatile solutes: • Vapour pressure is

always

lower • Boiling point is

always

higher • Freezing point is

always

lower • Osmotic pressure drives solvent from lower concentration to higher concentration

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.

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Physical vs Chemical

• Mixing is physical process; chemical properties don ’ t change • Properties of solutions are similar to those of the pure substances • Addition of a foreign substance to water alters the properties slightly

Energy Changes in Solution

• Simply put, three processes affect the energetics of solution: – separation of solute particles, – separation of solvent particles, – new interactions between solute and solvent.

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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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Endothermic Processes?

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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Factors Affecting Solubility

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

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Factors Affecting Solubility

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

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Factors Affecting Solubility

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

• Vitamin C is soluble in water.

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Gases in Solution

• 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

S g

=

kP g

• • • where

S g

is the solubility of the gas,

k

is the Henry ’ s Law constant for that gas in that solvent, and

P g

is the partial pressure of the gas above the liquid.

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Temperature

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

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Temperature

• The opposite is true of gases.

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

– Warm lakes have less O 2 dissolved in them than cool lakes.

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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 © 2009, Prentice-Hall, Inc.

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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Colligative Properties – Pressure Vapor • A solvent in a closed container reaches a state of dynamic equilibrium.

• The pressure exerted by the vapor in the headspace is referred to as the vapor pressure of the solvent.

• The addition of any nonvolatile solute (one with no measurable vapor pressure) to any solvent reduces the vapor pressure of the solvent.

Colligative Properties – Pressure Vapor • Nonvolatile solutes reduce the ability of the surface solvent molecules to escape the liquid.

– Vapor pressure is reduced .

• The extent of vapor pressure lowering depends on the amount of solute.

– Raoult ’ s Law quantifies the amount of vapor pressure lowering that is observed.

Non-volatile solutes and Raoult ’ s law • Vapor pressure of solvent in solution containing non volatile solute is

always

lower than vapor pressure of pure solvent at same T – At equilibrium rate of vaporization = rate of condensation – Solute particles occupy volume reducing rate of evaporationthe number of solvent molecules at the surface – The rate of evaporation decreases and so the vapor pressure above the solution must decrease to recover the equilibrium

Raoult

s Law

P A

=

X A P

A

where – –

X A P

A

is the mole fraction of compound A, and 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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Colligative Properties – Pressure Vapor Example: The vapor pressure of water at 20 o C is 17.5 torr. Calculate the vapor pressure of an aqueous solution prepared by adding 36.0 g of glucose (C 6 H 12 O 6 ) to 14.4 g of water.

Colligative Properties – Pressure Vapor Answer: 14.0 torr

Colligative Properties – Pressure Vapor Example: The vapor pressure of pure water at 110 o C is 1070 torr. A solution of ethylene glycol and water has a vapor pressure of 1.10 atm at the same temperature. What is the mole fraction of ethylene glycol in the solution?

Both ethylene glycol and water are liquids. How do you know which one is the solvent and which one is the solute?

Colligative Properties – Pressure Vapor Answer: X EG = 0.20

Colligative Properties – Pressure Vapor • Ideal solutions Raoult ’ s Law.

are those that obey • Real solutions show approximately ideal behavior when: – The solution concentration is low – The solute and solvent have similarly sized molecules – The solute and solvent have similar types of intermolecular forces.

Colligative Properties – Pressure Vapor • Raoult ’ s Law breaks down when solvent-solvent and solute-solute intermolecular forces of attraction are much stronger or weaker than solute-solvent intermolecular forces.

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: 

T b

=

K b

m

T b

is

added to

the normal boiling point of the solvent.

where

K b

is the molal boiling point elevation constant, a property of the solvent.

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Colligative Properties – BP have points than the pure solvent.

Elevation • The addition of a nonvolatile solute causes solutions to higher boiling – Vapor pressure decreases with addition of non volatile solute.

Higher temperature is needed in order for vapor pressure to equal 1 atm.

Molecular view of Raoult ’ s law: Boiling point elevation • In solution vapor pressure is reduced compared to pure solvent • Liquid boils when vapor pressure = atmospheric pressure • Must increase T to make vapor pressure = atmospheric

Freezing Point Depression

• The change in freezing point can be found similarly: 

T f

=

K f

m

• Here

K f

is the molal freezing point depression constant of the solvent.

T f

is

subtracted from

the normal boiling point of the solvent.

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Colligative Properties Freezing Pt Depression • The addition of a nonvolatile solute causes solutions to have lower freezing points than the pure solvent.

• Solid-liquid equilibrium line rises ~ vertically from the triple point, which is lower than that of pure solvent.

Freezing point of the solution is lower than that of the pure solvent.

Molecular view of Raoult ’ s law: Freezing point depression • Depends on the solute only being in the liquid phase – Fewer water molecules at surface: rate of freezing drops – Ice turns into liquid – Lower temperature to regain balance – Depression of freezing point

Colligative Properties Freezing Pt Depression • Reminder : You should be able to do the following as well: – Calculate the freezing point of any solution given enough information to calculate the molality of the solution and the value of K f – Calculate the molar mass of a solution given the value of K f and the freezing point depression (or the freezing points of the solution and the pure solvent).

Deviations from ideal

• Real solutions can deviate from the ideal: – Positive (P vap > ideal) solute-solvent interactions weaker – Negative (P vap < ideal) solute-solvent interactions stronger

VI. Colligative Properties

A. Vapor Pressure: it’s lowered B. Boiling Point: it’s elevated C. Freezing Point: it’s lowered D. Osmotic Pressure: it’s lowered

Colligative Properties Osmosis • Some substances form semipermeable membranes particles.

, allowing some smaller particles to pass through, but blocking other larger • In biological systems, most semipermeable membranes allow water to pass through, but solutes are not free to do so.

• If two solutions with identical concentration ( isotonic solutions ) are separated by a semipermeable membrane, no net movement of solvent occurs.

Colligative Properties Osmosis • Osmosis : the net movement of a solvent through a semipermeable membrane toward the solution with greater solute concentration.

• In osmosis, there is net movement of solvent from the area of

lower solute

concentration to the area of

higher solute

concentration.

– Movement of solvent from solvent concentration to concentration high low solvent

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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Osmosis: molecular discrimination • A

semi-permeable

membrane discriminates on the basis of molecular type – Solvent molecules pass through – Large molecules or ions are blocked • Solvent molecules will pass from a place of lower solute concentration to higher concentration to achieve equilibrium

Osmotic pressure

• Solvent passes into more conc solution increasing its volume • The passage of the solvent can be prevented by application of a pressure • The pressure to prevent transport is the

osmotic pressure

Calculating osmotic pressure

• The ideal gas law states

PV

• But n/V = M and so 

nRT

 

MRT

• Where M is the molar concentration of particles and

Π

is the osmotic pressure • Note:

molarity

is used not

molality

Colligative Properties in living systems: Osmosis • Osmosis plays an important role – Membranes of red blood cells are semipermeable.

• Placing a red blood cell in a hypertonic solution crenation .

( solute concentration outside the cell is greater than inside the cell cell in a process called ) causes water to flow out of the

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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Colligative Properties • Placing a red blood cell in a hypotonic solution cell ( solute concentration outside the cell is less than that inside the ) causes water to flow into the cell.

– The cell ruptures in a process called hemolysis .

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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Colligative Properties Osmosis • Other everyday examples of osmosis: – A cucumber placed in brine solution loses water and becomes a pickle.

– A limp carrot placed in water becomes firm because water enters by osmosis.

– Eating large quantities of salty food causes retention of water and swelling of tissues (edema).

Osmotic pressure and molecular mass • Molar mass can be computed from any of the colligative properties • Osmotic pressure provides the most accurate determination because of the magnitude of

Π

– 0.0200 M solution of glucose exerts an osmotic pressure of 374.2 mm Hg but a freezing point depression of only 0.02ºC

Osmotic Pressure

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

( )

RT =

M

RT V 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 .

© 2009, Prentice-Hall, Inc.

Determining molar mass

• A solution contains 20.0 mg insulin in 5.00 ml develops an osmotic pressure of 12.5 mm Hg at 300 K

M

 

RT M

 12 .

5

mmHg

0 .

0821

L

 1 760

mmHg atm mol

K

300

K

 6 .

68  10 4

M

Mass Percentage

Mass % of A = mass of A in solution total mass of solution  100 © 2009, Prentice-Hall, Inc.

Parts per Million and Parts per Billion

Parts per Million (ppm) ppm = mass of A in solution total mass of solution  10 6 Parts per Billion (ppb) ppb = mass of A in solution total mass of solution  10 9 © 2009, Prentice-Hall, Inc.

Mole Fraction (X)

X A

= moles of A total moles in solution • In some applications, one needs the mole fraction of

solvent

, not solute — make sure you find the quantity you need!

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Molarity (M)

M = mol of solute L of solution • You will recall this concentration measure from Chapter 4.

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

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Molality (m)

m = mol of solute kg of solvent 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.

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