Chapter 13 Properties of Solutions

Solutions

Chapter 13Properties of Solutions

John D. Bookstaver

St. Charles Community College

St. Peters, MO

2006, Prentice Hall, Inc.

Chemistry, The Central Science, 10th editionTheodore L. Brown; H. Eugene LeMay, Jr.;

and Bruce E. Bursten

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December 10

• The solution process

• Why a solution forms?

• Chapter HW

• 1,3,4,6

• Solution process13, 15

• Saturated solutions – Factors affecting solubility19, 21, 25, 27, 29, 31

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Section 13.1The Solution process

• Energy Changes and Solution Formation

• Solution Formation, Spontaneity and Disorder.

• Solution Formation and Chemical Reactions

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Solutions• Solutions are homogeneous mixtures of two

or more pure substances.• In a solution, the solute (present in smaller

amount) is dispersed uniformly throughout the solvent (present in largest amount).

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

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How Does a Solution Form?

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

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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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Energy Changes and SolutionFormation

Three processes affect the energetic of the process:Separation of solute

particlesH1

Separation of solvent particles H2

New interactions between solute and solvent H3

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Energy Changes and Solution Formation

• We define the enthalpy change in the solution process as

Hsoln = H1 + H2 + H3.

Hsoln can either be positive or negative depending on the intermolecular forces.

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• Breaking attractive intermolecular forces is always endothermic.

• Forming attractive intermolecular forces is always 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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Why Do Endothermic Processes Occur?

Things do not tend to occur spontaneously (i.e., without outside intervention) unless the energy of the system is lowered.

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• To determine whether Hsoln is positive or negative, we consider the strengths of all solute-solute and solute-solvent interactions: H1 and H2 are both positive.

H3 is always negative.

– It is possible to have either H3 > (H1 + H2) or H3 < (H1 + H2).

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• Examples: – NaOH added to water has Hsoln = -44.48 kJ/mol.

– NH4NO3 added to water has Hsoln = + 26.4 kJ/mol.

• “Rule”: LIKE DISSOLVES LIKE!!!• polar solvents dissolve polar solutes. Non-polar

solvents dissolve non-polar solutes. Why?– If Hsoln is too endothermic a solution will not form.

– NaCl in gasoline: the ion-dipole forces are weak because gasoline is non-polar. Therefore, the ion-dipole forces do not compensate for the separation of ions.

– Water in octane: water has strong H-bonds. There are no attractive forces between water and octane to compensate for the H-bonds.

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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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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.

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Entropy- Disorder

So even though enthalpy may increase, the overall energy of the system can still decrease if the system becomes more disordered.

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Solution formation -Spontaneity• Spontaneous change tend to occur if

the process results in • a) lower energy for the whole system• So the changes with H < 0 are

favored.

• b) an increase in the total disorder, randomnes, degree of dispersal or entropy S > 0

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• A solution will form except in the cases that the attraction between solute-solute/solvent-solvent are too strong compared with the solute-solvent attractions.

• Solution formation always increase the Entropy of the system.

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Solution formation and chemical reaction

• 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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Student, Beware!

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

Consider:

Ni(s) + 2HCl(aq) NiCl2(aq) + H2(g).

• Note the chemical form of the substance being dissolved has changed (Ni NiCl2).

• When all the water is removed from the solution, no Ni is found (only NiCl2·6H2O). Therefore, Ni dissolution in HCl is a chemical process.

• NiCl2·6H2O is a hydrate that contains Ni2+ ions.

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• Example:

NaCl(s) + H2O (l) Na+(aq) + Cl-(aq).

• When the water is removed from the solution, NaCl is found. Therefore, NaCl dissolution is a physical process.

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• Section 13.3

• Factors Affecting Solubility

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Types of Solutions

• SaturatedSolvent holds as much

solute as is possible at that temperature.

Dissolved solute is in dynamic equilibrium with solid solute particles.

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Types of Solutions

• UnsaturatedLess than the

maximum amount of solute for that temperature is dissolved in the solvent.

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Types of Solutions

• Supersaturated 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.

NaC2H3O2 Sodium acetate usually forms supersaturated solutions.

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Factors Affecting Solubility

• *Solute-Solvent interactions

• *Pressure effects (only for gases)

• *Temperature effects

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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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Solute-Solvent Interaction• Polar substances tend to dissolve in polar solvents.• Miscible liquids: mix in any proportions.• Immiscible liquids: do not mix.• Intermolecular forces are important: water and ethanol

are miscible because the broken hydrogen bonds in both pure liquids are re-established in the mixture.

• The number of carbon atoms in a chain affect solubility: the more C atoms the less soluble in water.

Factors Affecting Factors Affecting SolubilitySolubility

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Solute-Solvent Interaction• The number of -OH groups within a molecule increases

solubility in water.• Generalization: “like dissolves like”.• The more polar bonds in the molecule, the better it

dissolves in a polar solvent.• The less polar the molecule the less it dissolves in a polar

solvent and the better is dissolves in a non-polar solvent.

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Example – Place the following substances in order of increasing solubility in water:

C C C O

H

O

H

H

H

H

H

HH

C C C O

H

H

H

H

H

H

H

H

C C C C

H

H

H

H

H

H

H

H

H

H

C C C O

H

C

C

H

H

H

H

H

H

H

H

H

HH

C C C O

H

C

H

H

H

H

H

H

H

H

H

1

2

3

5

4

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Example – Place the following substances in order of increasing solubility in hexane (C6H14):

C C C O

H

O

H

H

H

H

H

HH

C C C O

H

H

H

H

H

H

H

H

C C C C

H

H

H

H

H

H

H

H

H

H

C C C O

H

C

C

H

H

H

H

H

H

H

H

H

HH

C C C O

H

C

H

H

H

H

H

H

H

H

H

5

3

2

1

4

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Solute-Solvent Interaction

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Solute-Solvent Interaction

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Solute-Solvent Interaction• Network solids do not dissolve because the strong

intermolecular forces in the solid are not re-established in any solution.

Pressure Effects• Solubility of a gas in a liquid is a function of the pressure

of the gas.

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Gases in Solution

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

• Larger molecules have stronger dispersion forces.

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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

Sg = kPg

where

• Sg is the solubility of the gas;

• k is the Henry’s law constant for that gas in that solvent;

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

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Henry’s Law Constant

• Is different for each solute-gas pair.

• Varies with temperature

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• The higher the pressure, the more molecules of gas are close to the solvent and the greater the chance of a gas molecule striking the surface and entering the solution.– Therefore, the higher the pressure, the greater the solubility.

– The lower the pressure, the fewer molecules of gas are close to the solvent and the lower the solubility.

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Example – What is the concentration of CO2 in water in a soda bottled under a pressure of 20.0 atm of CO2?

(k = 3.1 x 10-2 mol L-1 atm-1)

0.62 mol L-1 or 0.62 M

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Pressure Effects• Carbonated beverages are bottled with a partial pressure

of CO2 > 1 atm.

• As the bottle is opened, the partial pressure of CO2 decreases and the solubility of CO2 decreases.

• Therefore, bubbles of CO2 escape from solution.

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Temperature

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

Note the exception

Ce2(SO4)3

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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 O2 dissolved in them than cool lakes.

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Ways of Expressing

Concentrations of Solutions

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Homework

• 33, 35, 37, 41, 45, 47, 49

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Concentrations

* Qualitative Expressions

Dilute – Concentrate

*Quantitative Expressions

• Molarity-Molality

• %by mass, ppm,ppb

• Mole Fraction

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Mass Percentage

Mass % of A =mass of A in solutiontotal mass of solution

100

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• Find the percent of KCl in a solution that contains .005 g of KCl in 50 g of solution

• Express the result in ppm and ppb

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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 (X)

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

• REMEMBER THE SUM OF THE MOLE FRACTIONS OF ALL COMPONENTS OF THE SOLUTION = 1

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• Find the mole fraction of CH3OH in a solution that contains 32 gr of methanol in 36 gr of water.

• What is the mole fraction of water in that solution

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mol of soluteL of solution

M =

Molarity (M)

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

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mol of solutekg of solvent

m =

Molality (m)

Because 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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Example - 1.00 g of NaCl is dissolved in 50.0 g of H2O, to make a solution with a total volume of 50.7 mL. Calculate the molarity, mass percent, mole fraction, molality, ppm, and ppb of NaCl in this solution.

0.338 M 1.96 %

X = 0.00613 0.342 m

19600 ppm 19600000 ppb

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December 12

• Colligative properties.• HW 55, 57, 59, 61, 63, 65• Pre lab for experiment 11 due

wednesday in a separate paper. • Review questions from Pearson site

due friday – before test• Test on solutions on friday 16

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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 areVapor pressure lowering Boiling point elevationMelting point depressionOsmotic 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.

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

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Raoult’s Law

PA = XAPAwhere

PA is the vapor pressure of the solution• XA is the mole fraction of compound A• 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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Lowering Vapor Pressure

• Raoult’s Law: PA is the vapor pressure of the solution, PA is the vapor pressure of pure solvent, and A is the mole fraction of A, then

• Recall Dalton’s Law:

• Ptotal = P1 + P2 = 1Po1 + 2Po

2

AAA PP

totalPP AA

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• Ideal solution: one that obeys Raoult’s law.• Raoult’s law breaks down when the solvent-solvent and

solute-solute intermolecular forces are greater than solute-solvent intermolecular forces.

Example - Calculate the Vapor Pressure of a solution prepared by dissolving 50.0 g of sugar (C12H22O11) in 200. g of H2O. (Vapor Pressure of pure H2O = 23.76 torr)

23.55 torr

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Example

• Calculate the vapor pressure of a solution of 25.0 g of H2O and 30.0 g of C2H5OH at 25o C.

Vapor pressures:

H2O = 23.76 torr

C2H5OH = 64.84 torr

36.9 torr

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Example

What is the composition of a pentane-hexane solution that has a vapor pressure of 350 torr at 25ºC?

The vapor pressures at 25ºC are– pentane 511 torr– hexane 150 torr

What is the composition of the vapor?

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FREEZING POINT DEPRESION

• When a solute is dissolved in a solvent, the solvent particles form a shell around the solute particles. When the temperature decreases, the solute particles disrupt the crystal formation of the solvent because of the shells of hydration, therefore more kinetic energy must be removed from the solution in order for the solvent to solidify, causing the freezing point to depress.

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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 normal freezing point of the solvent.

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Boiling Point ElevationThe 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 normal

boiling point 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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Freezing Point Depression• At 1 atm (normal boiling point of pure liquid) there is no

depression by definition• When a solution freezes, almost pure solvent is formed

first.– Therefore, the sublimation curve for the pure solvent is the

same as for the solution.

– Therefore, the triple point occurs at a lower temperature because of the lower vapor pressure for the solution.

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• The melting-point (freezing-point) curve is a near-vertical line from the triple point.

• The solution freezes at a lower temperature (Tf) than the pure solvent.

• Decrease in freezing point (Tf) is directly proportional to molality (Kf is the molal freezing-point-depression constant): miKT ff

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December 14

Colligative properties

Review for lab 11

Will do the lab Friday. Prelab due tomorrow

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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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Colligative Properties of Electrolytes

However, a 1 M solution of NaCl does not show twice the change in freezing point that a 1 M 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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The van’t Hoff Factor

• Reassociation is more likely at higher concentration.

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

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i = vant Hoff factor

= number of particles produced when a substance dissolves

For nonionic substances, i=1

For ionic substances, i is the number of ions produced per formula unit that dissolves.

NaCl, i = 2 Na3PO4 i = 4

Na3PO4 3 Na+ + PO43-

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Explain

• Sodium chloride may be spread on an icy sidewalk in order to melt the ice; equimolar amounts of calcium chloride are even more effective. (10 pts)

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Deviations from Raoult’s Law• If Solvent has a strong affinity for solute (H

bonding).– Lowers solvent’s ability to escape.– Lower vapor pressure than expected.– Negative deviation from Raoult’s law.

Hsoln is large and negative exothermic.

• Endothermic Hsoln indicates positive deviation.

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The van’t Hoff Factor

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

Tf = Kf m 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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• Hemolysis: – red blood cells placed in a hypotonic solution;

– there is a higher solute concentration in the cell;

– osmosis occurs and water moves into the cell.

– The cell bursts.

• To prevent crenation or hemolysis, IV (intravenous) solutions must be isotonic.

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– Cucumber placed in NaCl solution loses water to shrivel up and become a pickle.

– Limp carrot placed in water becomes firm because water enters via osmosis.

– Salty food causes retention of water and swelling of tissues (edema).

– Water moves into plants through osmosis.

– Salt added to meat or sugar to fruit prevents bacterial infection (a bacterium placed on the salt will lose water through osmosis and die).

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• Active transport is the movement of nutrients and waste material through a biological system.

• Active transport is not spontaneous.

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Example

What is the osmotic pressure of a solution of 7.95 g of NaCl in 50.0 mL of an aqueous solution at 75˚C?

155 atm

118,000 mm Hg

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Molar Mass from Colligative Properties

We can use the effects of a colligative property such as osmotic pressure to determine the molar mass of a compound.

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One major application of vapor pressure lowering and colligative properties is in molar mass problems

1. An aqueous solution contains 1.00 g/L of a detergent. The osmotic pressure of this solution at 25˚C is 17.8 torr. What is the molar mass of the detergent?

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2. 1.008 g of a compound was dissolved in 11.38 mL of benzene (d=0.879 g/mL) and the solution froze at 4.37˚C. What is the molar mass of the compound?

Tf(benzene) = 5.48˚C

Kf(benzene) = 5.12˚C/molal

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3. 3.101 g of a nonvolatile nonelectrolyte were dissolved in 100. g of CCl4. The vapor pressure of CCl4 was lowered by 1.85%. What is the molar mass of the solute?

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• Colloids are suspensions in which the suspended particles are larger than molecules but too small to drop out of the suspension due to gravity.

• Particle size: 10 to 2000 Å.• There are several types of colloid:

– aerosol (gas + liquid or solid, e.g. fog and smoke),

– foam (liquid + gas, e.g. whipped cream),

– emulsion (liquid + liquid, e.g. milk),

– sol (liquid + solid, e.g. paint),

ColloidsColloids

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– solid foam (solid + gas, e.g. marshmallow),

– solid emulsion (solid + liquid, e.g. butter),

– solid sol (solid + solid, e.g. ruby glass).

• Tyndall effect: ability of a Colloid to scatter light. The beam of light can be seen through the colloid.

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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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Colloids in Biological Systems

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

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Colloids in Biological Systems

Sodium stearate is one example of such a molecule.

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• Sodium stearate has a long hydrophobic tail (CH3(CH2)16-) and a small hydrophobic head (-CO2

-Na+).

• The hydrophobic tail can be absorbed into the oil drop, leaving the hydrophilic head on the surface.

• The hydrophilic heads then interact with the water and the oil drop is stabilized in water.

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Colloids in Biological Systems

These molecules can aid in the emulsification of fats and oils in aqueous solutions.

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Hydrophilic and Hydrophobic Colloids• Focus on colloids in water.• “Water loving” colloids: hydrophilic.• “Water hating” colloids: hydrophobic.• Molecules arrange themselves so that hydrophobic

portions are oriented towards each other.• If a large hydrophobic macromolecule (giant molecule)

needs to exist in water (e.g. in a cell), hydrophobic molecules embed themselves into the macromolecule leaving the hydrophilic ends to interact with water.

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• Typical hydrophilic groups are polar (containing C-O, O-H, N-H bonds) or charged.

• Hydrophobic colloids need to be stabilized in water.• Adsorption: when something sticks to a surface we say

that it is adsorbed.• If ions are adsorbed onto the surface of a colloid, the

colloids appears hydrophilic and is stabilized in water.• Consider a small drop of oil in water.• Add to the water sodium stearate.

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ColloidsColloids

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• Most dirt stains on people and clothing are oil-based. Soaps are molecules with long hydrophobic tails and hydrophilic heads that remove dirt by stabilizing the colloid in water.

• Bile excretes substances like sodium stereate that forms an emulsion with fats in our small intestine.

• Emulsifying agents help form an emulsion.

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Removal of Colloidal Particles• Colloid particles are too small to be separated by physical

means (e.g. filtration).• Colloid particles are coagulated (enlarged) until they can

be removed by filtration.• Methods of coagulation:

– heating (colloid particles move and are attracted to each other when they collide);

– adding an electrolyte (neutralize the surface charges on the colloid particles).

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• Dialysis: using a semipermeable membranes separate ions from colloidal particles.