chapter 2 l.
Download
Skip this Video
Loading SlideShow in 5 Seconds..
CHAPTER 2 PowerPoint Presentation
Download Presentation
CHAPTER 2

Loading in 2 Seconds...

play fullscreen
1 / 114

CHAPTER 2 - PowerPoint PPT Presentation


  • 124 Views
  • Uploaded on

CHAPTER 2. Solutions By Dr. Hisham Ezzat . http://www.staff.zu.edu.eg/ezzat_hisham/browseMyFiles.asp?path=./userdownloads/physical%20chemistry%20for%20clinical%20pharmacy/. The Dissolution Process. Solutions are homogeneous mixtures of two or more substances.

loader
I am the owner, or an agent authorized to act on behalf of the owner, of the copyrighted work described.
capcha
Download Presentation

PowerPoint Slideshow about 'CHAPTER 2' - monet


An Image/Link below is provided (as is) to download presentation

Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author.While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server.


- - - - - - - - - - - - - - - - - - - - - - - - - - E N D - - - - - - - - - - - - - - - - - - - - - - - - - -
Presentation Transcript
chapter 2
CHAPTER 2
  • Solutions

By

Dr. Hisham Ezzat

http://www.staff.zu.edu.eg/ezzat_hisham/browseMyFiles.asp?path=./userdownloads/physical%20chemistry%20for%20clinical%20pharmacy/

the dissolution process
The Dissolution Process
  • Solutions are homogeneous mixtures of two or more substances.
    • Dissolving medium is called the solvent.
    • Dissolved species are called the solute.
  • There are three states of matter (solid, liquid, and gas) which when mixed two at a time gives nine different kinds of mixtures.
    • Seven of the possibilities can be homogeneous.
    • Two of the possibilities must be heterogeneous.
the dissolution process3
The Dissolution Process

Seven Homogeneous Possibilities

SoluteSolventExample

  • Solid Liquid salt water
  • Liquid Liquid mixed drinks
  • Gas Liquid carbonated beverages
  • Liquid Solid dental amalgams
  • Solid Solid alloys
  • Gas Solid metal pipes
  • Gas Gas air

Two Heterogeneous Possibilities

  • Solid Gas dust in air
  • Liquid Gas clouds, fog
ways of expressing concentration
Ways of Expressing Concentration
  • Qualitative Terms:
  • Dilute Solution – A dilute solution has a relatively small concentration of solute.
  • A concentrated solution has a relatively high concentration of solute.
slide5

Quantitative terms

Molarity

Chapter 13

slide6

Molality

  • Molality is a concentration unit based on the number of moles of solute per kilogram of solvent.

Chapter 13

molality and mole fraction
Molality and Mole Fraction
  • Weight percent (wt %)
  • Volume percent or percent by volume (vol %)
molality and mole fraction8
Molality and Mole Fraction
  • Molarity
  • We must introduce two new concentration units in this chapter.
molality and mole fraction9
Molality and Mole Fraction
  • Example 14-1: Calculate the molarity and the molality of an aqueous solution that is 10.0% glucose, C6H12O6. The density of the solution is 1.04 g/mL. 10.0% glucose solution has several medical uses. 1 mol C6H12O6 = 180 g
molality and mole fraction10
Molality and Mole Fraction
  • Molality is a concentration unit based on the number of moles of solute per kilogram of solvent.
molality and mole fraction11
Molality and Mole Fraction
  • Example 14-1: Calculate the molality and the molarity of an aqueous solution that is 10.0% glucose, C6H12O6. The density of the solution is 1.04 g/mL. 10.0% glucose solution has several medical uses. 1 mol C6H12O6 = 180 g

You calculate the molarity!

molality and mole fraction12
Molality and Mole Fraction
  • Example 14-2: Calculate the molality of a solution that contains 7.25 g of benzoic acid C6H5COOH, in 2.00 x 102 mL of benzene, C6H6. The density of benzene is 0.879 g/mL. 1 mol C6H5COOH = 122 g

You do it!

molality and mole fraction13
Molality and Mole Fraction
  • Mole fraction is the number of moles of one component divided by the moles of all the components of the solution
    • Mole fraction is literally a fraction using moles of one component as the numerator and moles of all the components as the denominator.
  • In a two component solution, the mole fraction of one component, A, has the symbol XA.
molality and mole fraction14
Molality and Mole Fraction
  • The mole fraction of component B - XB
molality and mole fraction15
Molality and Mole Fraction
  • Example 14-3: What are the mole fractions of glucose and water in a 10.0% glucose solution (Example 14-1)?

You do it!

molality and mole fraction16
Molality and Mole Fraction
  • Example 14-3: What are the mole fractions of glucose and water in a 10.0% glucose solution (Example 14-1)?
molality and mole fraction17
Molality and Mole Fraction
  • Now we can calculate the mole fractions.
chapter goals
Chapter Goals

The Dissolution Process

  • Spontaneity of the Dissolution Process
  • Dissolution of Solids in Liquids
  • Dissolution of Liquids in Liquids (Miscibility)
  • Dissolution of Gases in Liquids
  • Rates of Dissolution and Saturation
  • Effect of Temperature on Solubility
  • Effect of Pressure on Solubility
  • Molality and Mole Fraction
chapter goals19
Chapter Goals

Colligative Properties of Solutions

  • Lowering of Vapor Pressure and Raoult’s Law
  • Fractional Distillation
  • Boiling Point Elevation
  • Freezing Point Depression
  • Determination of Molecular Weight by Freezing Point Depression or Boiling Point Elevation
  • Colligative Properties and Dissociation of Electrolytes
  • Osmotic Pressure
chapter goals20
Chapter Goals

Colloids

  • The Tyndall Effect
  • The Adsorption Phenomena
  • Hydrophilic and Hydrophobic Colloids
spontaneity of the dissolution process
Spontaneity of the Dissolution Process
  • As an example of dissolution, let’s assume that the solvent is a liquid.
  • Two major factors affect dissolution of solutes
  • Change of energy content or enthalpy of solution, Hsolution
    • If Hsolution is exothermic (< 0) dissolution is favored.
    • If Hsolution is endothermic (> 0) dissolution is not favored.
spontaneity of the dissolution process22
Spontaneity of the Dissolution Process
  • Change in disorder, or randomness, of the solution Smixing
      • If Smixing increases (> 0) dissolution is favored.
      • If Smixing decreases (< 0) dissolution is not favored.
  • Thus the best conditions for dissolution are:
    • For the solution process to be exothermic.
      • Hsolution < 0
    • For the solution to become more disordered.
      • Smixing > 0
spontaneity of the dissolution process23
Spontaneity of the Dissolution Process
  • Disorder in mixing a solution is very common.
    • Smixing is almost always > 0.
  • What factors affect Hsolution?
    • There is a competition between several different attractions.
  • Solute-solute attractions such as ion-ion attraction, dipole-dipole, etc.
    • Breaking the solute-solute attraction requires an absorption of E.
spontaneity of the dissolution process24
Spontaneity of the Dissolution Process
  • Solvent-solvent attractions such as hydrogen bonding in water.
    • This also requires an absorption of E.
  • Solvent-solute attractions, solvation, releases energy.
    • If solvation energy is greater than the sum of the solute-solute and solvent-solvent attractions, the dissolution is exothermic, Hsolution < 0.
    • If solvation energy is less than the sum of the solute-solute and solvent-solvent attractions, the dissolution is endothermic, Hsolution > 0.
dissolution of solids in liquids
Dissolution of Solids in Liquids
  • The energy released (exothermic) when a mole of formula units of a solid is formed from its constituent ions (molecules or atoms for nonionic solids) in the gas phase is called the crystal lattice energy.
  • The crystal lattice energy is a measure of the attractive forces in a solid.
  • The crystal lattice energy increases as the charge density increases.
dissolution of solids in liquids27
Dissolution of Solids in Liquids
  • Dissolution is a competition between:
    • Solute -solute attractions
      • crystal lattice energy for ionic solids
    • Solvent-solvent attractions
      • H-bonding for water
    • Solute-solvent attractions
      • Solvation or hydration energy
dissolution of solids in liquids28
Dissolution of Solids in Liquids
  • Solvation is directed by the water to ion attractions as shown in these electrostatic potentials.
dissolution of solids in liquids29
Dissolution of Solids in Liquids
  • In an exothermic dissolution, energy is released when solute particles are dissolved.
    • This energy is called the energy of solvation or the hydration energy (if solvent is water).
  • Let’s look at the dissolution of CaCl2.
dissolution of solids in liquids30

OH2

2+

O

H H

OH2

H2O

H

O

H

H

O

H

Ca

Cl-

H2O

OH2

H H

O

OH2

Dissolution of Solids in Liquids
dissolution of solids in liquids31
Dissolution of Solids in Liquids
  • The energy absorbed when one mole of formula units becomes hydrated is the molar energy of hydration.
dissolution of solids in liquids32
Dissolution of Solids in Liquids
  • Hydration energy increases with increasing charge density

IonRadius(Å)Charge/radiusHeat of Hydration

K+ 1.33 0.75 -351 kJ/mol

Ca2+ 0.99 2.02 -1650 kJ/mol

Cu2+0.72 2.78 -2160 kJ/mol

Al3+ 0.50 6.00 -4750 kJ/mol

dissolution of liquids in liquids miscibility
Dissolution of Liquids in Liquids (Miscibility)
  • Most polar liquids are miscible in other polar liquids.
  • In general, liquids obey the “like dissolves like” rule.
    • Polar molecules are soluble in polar solvents.
    • Nonpolar molecules are soluble in nonpolar solvents.
  • For example, methanol, CH3OH, is very soluble in water
dissolution of liquids in liquids miscibility34
Dissolution of Liquids in Liquids (Miscibility)
  • Nonpolar molecules essentially “slide” in between each other.
    • For example, carbon tetrachloride and benzene are very miscible.
dissolution of gases in liquids
Dissolution of Gases in Liquids
  • Polar gases are more soluble in water than nonpolar gases.
    • This is the “like dissolves like” rule in action.
  • Polar gases can hydrogen bond with water
  • Some polar gases enhance their solubility by reacting with water.
dissolution of gases in liquids36
A few nonpolar gases are soluble in water because they react with water.

Because gases have very weak solute-solute interactions, gases dissolve in liquids in exothermic processes.

Dissolution of Gases in Liquids
rates of dissolution and saturation
Finely divided solids dissolve more rapidly than large crystals.

Compare the dissolution of granulated sugar and sugar cubes in cold water.

The reason is simple, look at a single cube of NaCl.

The enormous increase in surface area helps the solid to dissolve faster.

Breaks

many smaller crystals

up

NaCl

Rates of Dissolution and Saturation
rates of dissolution and saturation38
Rates of Dissolution and Saturation
  • Saturated solutions have established an equilbrium between dissolved and undissolved solutes
    • Examples of saturated solutions include:
      • Air that has 100% humidity.
      • Some solids dissolved in liquids.
rates of dissolution and saturation39
Symbolically this equilibrium is written as:

In an equilibrium reaction, the forward rate of reaction is equal to the reverse rate of reaction.

Rates of Dissolution and Saturation
rates of dissolution and saturation40
Rates of Dissolution and Saturation
  • Supersaturated solutions have higher-than-saturated concentrations of dissolved solutes.
effect of temperature on solubility
According to LeChatelier’s Principle when stress is applied to a system at equilibrium, the system responds in a way that best relieves the stress.

Since saturated solutions are at equilibrium, LeChatelier’s principle applies to them.

Possible stresses to chemical systems include:

Heating or cooling the system.

Changing the pressure of the system.

Changing the concentrations of reactants or products.

Effect of Temperature on Solubility
effect of temperature on solubility42
What will be the effect of heating or cooling the water in which we wish to dissolve a solid?

It depends on whether the dissolution is exo- or endothermic.

For an exothermic dissolution, heat can be considered as a product.

Warming the water will decrease solubility and cooling the water will increase the solubility.

Predict the effect on an endothermic dissolution like this one.

Effect of Temperature on Solubility
effect of temperature on solubility43
Effect of Temperature on Solubility
  • For ionic solids that dissolve endothermically dissolution is enhanced by heating.
  • For ionic solids that dissolve exothermically dissolution is enhanced by cooling.
  • Be sure you understand these trends.
effect of pressure on solubility
Effect of Pressure on Solubility
  • Pressure changes have little or no effect on solubility of liquids and solids in liquids.
    • Liquids and solids are not compressible.
  • Pressure changes have large effects on the solubility of gases in liquids.
    • Sudden pressure change is why carbonated drinks fizz when opened.
    • It is also the cause of several scuba diving related problems including the “bends”.
effect of pressure on solubility45
Effect of Pressure on Solubility
  • The effect of pressure on the solubility of gases in liquids is described by Henry’s Law.
colligative properties of solutions
Colligative Properties of Solutions
  • Colligative properties are properties of solutions that depend solely on the number of particles dissolved in the solution.
    • Colligative properties do not depend on the kinds of particles dissolved.
  • Colligative properties are a physical property of solutions.
colligative properties of solutions47
Colligative Properties of Solutions
  • There are four common types of colligative properties:
    • Vapor pressure lowering
    • Freezing point depression
    • Boiling point elevation
    • Osmotic pressure
  • Vapor pressure lowering is the key to all four of the colligative properties.
lowering of vapor pressure and raoult s law
Lowering of Vapor Pressure and Raoult’s Law
  • Addition of a nonvolatile solute to a solution lowers the vapor pressure of the solution.
    • The effect is simply due to fewer solvent molecules at the solution’s surface.
    • The solute molecules occupy some of the spaces that would normally be occupied by solvent.
  • Raoult’s Law models this effect in ideal solutions.
lowering of vapor pressure and raoult s law50
Lowering of Vapor Pressure and Raoult’s Law
  • Lowering of vapor pressure, Psolvent, is defined as:
lowering of vapor pressure and raoult s law51
Lowering of Vapor Pressure and Raoult’s Law
  • Remember that the sum of the mole fractions must equal 1.
  • Thus Xsolvent + Xsolute = 1, which we can substitute into our expression.
lowering of vapor pressure and raoult s law52
Lowering of Vapor Pressure and Raoult’s Law
  • This graph shows how the solution’s vapor pressure is changed by the mole fraction of the solute, which is Raoult’s law.
examples
Examples

The vapor pressure of water is 17.5 torr at 20°C. Imagine holding the temperature constant while adding glucose, C6H12O6, to the water so that the resulting solution has XH2O = 0.80 and XGlu = 0.20. What is , the vapor pressure of water over the solution

= 14 torr

slide54

Glycerin, C3H8O3, is a nonvolatile nonelectrolyte with a density of 1.26 g/mL at 25°C. Calculate the vapor pressure at 25°C of a solution made by adding 50.0 mL of glycerin to 500.0 mL of water. The vapor pressure of pure water at 25°C is 23.8 torr

slide55

The vapor pressure of pure water at 110°C is 1070 torr. A solution of ethylene glycol and water has a vapor pressure of 1.00 atm at 110°C. Assuming that Raoult's law is obeyed, what is the mole fraction of ethylene glycol in the solution? Answer: 0.290

P°H2O =1070 torr

PH2O = 1 Atm = 760 torr

PH2O

P°H2O

760 torr

1070 torr

XH2O = ---------

= ---------

=

XH2O + XEG = 1

0.7103 + XEG = 1

1- 0.7103 = XEG

= 0.290

XEG =

slide56

Many solutions do not obey Raoult's law exactly: They are not ideal solutions.

If the intermolecular forces between solvent and solute are weaker than those between solvent and solvent and between solute and solute, then the solvent vapor pressure tends to be greater than predicted by Raoult's law.

Conversely, when the interactions between solute and solvent are exceptionally strong, as might be the case when hydrogen bonding exists, the solvent vapor pressure is lower than Raoult's law predicts.

Although you should be aware that these departures from ideal solution occur, we will ignore them for the remainder of this chapter.

more examples
More Examples

Sucrose is a nonvolatile, nonionizing solute in water. Determine the vapor pressure lowering, at 27°C, of a solution of 75.0 grams of sucrose, C12H22O11, dissolved in 180. g of water. The vapor pressure of pure water at 27°C is 26.7 torr. Assume the solution is ideal.

Vapor Pressure Lowered = 26.7-26.1= 0.6

slide58

solution is made by mixing 52.1 g of propyl chloride, C3H8Cl, and 38.4 g of propyl bromide, C3H8Br. What is the vapor pressure of propyl chloride in the solution at 25°C? The vapor pressure of pure propyl chloride is 347 torr at 25°C and that of pure propyl bromide is 133 torr at 25°C. Assume that the solution is an ideal solution.

slide59

. At 25°C a solution consists of 0.450 mole of pentane, C5H12, and 0.250 mole of cyclopentane, C5H10. What is the mole fraction of cyclopentane in the vapor that is in equilibrium with this solution? The vapor pressure of the pure liquids at 25°C are 451 torr for pentane and 321 torr for cyclopentane. Assume that the solution is an ideal solution.

fractional distillation
Fractional Distillation
  • Distillation is a technique used to separate solutions that have two or more volatile components with differing boiling points.
  • A simple distillation has a single distilling column.
    • Simple distillations give reasonable separations.
  • A fractional distillation gives increased separations because of the increased surface area.
    • Commonly, glass beads or steel wool are inserted into the distilling column.
boiling point elevation
Boiling Point Elevation
  • Addition of a nonvolatile solute to a solution raises the boiling point of the solution above that of the pure solvent.
    • This effect is because the solution’s vapor pressure is lowered as described by Raoult’s law.
    • The solution’s temperature must be raised to make the solution’s vapor pressure equal to the atmospheric pressure.
  • The amount that the temperature is elevated is determined by the number of moles of solute dissolved in the solution.
boiling point elevation62
Boiling Point Elevation
  • Boiling point elevation relationship is:
boiling point elevation63
Boiling Point Elevation
  • Example 14-4: What is the normal boiling point of a 2.50 m glucose, C6H12O6, solution?
boiling point elevation64
Boiling-Point Elevation

The addition of a nonvolatile solute lowers the vapor pressure of the solution.

At any given temperature,

the vapor pressure of the solution is lower than that of the pure liquid

slide65

The increase in boiling point relative to that of the pure solvent, DTb, is directly proportional to the number of solute particles per mole of solvent molecules.

Molality expresses the number of moles of solute per 1000 g of solvent, which represents a fixed number of moles of solvent

slide66

Automotive antifreeze consists of ethylene glycol, C2H6O2, a nonvolatile nonelectrolyte. Calculate the boiling point of a 25.0 mass percent solution of ethylene glycol in water.

slide67

Calculate the boiling point of a solution of 2.0 molal of NaCl. Kb, water= 0.52 °C /mola.

Dt = Kbm

NaCl(aq)  Na+ + Cl-

2.0 m 2.0 m 2.0 m

2.0 m + 2.0 m = 4.0m

Dt = (0.52 °C/molal)(4.0 molal) =2.08 °C

BP = NBP +Dt = 100.00°C +2.08 °C = 102.08° C

freezing point depression
Freezing Point Depression
  • Addition of a nonvolatile solute to a solution lowers the freezing point of the solution relative to the pure solvent.
  • See table 14-2 for a compilation of boiling point and freezing point elevation constants.
freezing point depression69
Freezing Point Depression
  • Relationship for freezing point depression is:
freezing point depression70
Fundamentally, freezing point depression and boiling point elevation are the same phenomenon.

The only differences are the size of the effect which is reflected in the sizes of the constants, Kf & Kb.

This is easily seen on a phase diagram for a solution.

Freezing Point Depression
  • Notice the similarity of the two relationships for freezing point depression and boiling point elevation.
freezing point depression72
Freezing Point Depression
  • Example 14-5: Calculate the freezing point of a 2.50 m aqueous glucose solution.
freezing point depression73
Freezing Point Depression
  • Example 14-6: Calculate the freezing point of a solution that contains 8.50 g of benzoic acid (C6H5COOH, MW = 122) in 75.0 g of benzene, C6H6.

You do it!

determination of molecular weight by freezing point depression
Determination of Molecular Weight by Freezing Point Depression
  • The size of the freezing point depression depends on two things:
    • The size of the Kf for a given solvent, which are well known.
    • And the molal concentration of the solution which depends on the number of moles of solute and the kg of solvent.
  • If Kf and kg of solvent are known, as is often the case in an experiment, then we can determine # of moles of solute and use it to determine the molecular weight.
determination of molecular weight by freezing point depression76
Determination of Molecular Weight by Freezing Point Depression
  • Example 14-7: A 37.0 g sample of a new covalent compound, a nonelectrolyte, was dissolved in 2.00 x 102 g of water. The resulting solution froze at -5.58oC. What is the molecular weight of the compound?
colligative properties and dissociation of electrolytes
Colligative Properties and Dissociation of Electrolytes
  • Electrolytes have larger effects on boiling point elevation and freezing point depression than nonelectrolytes.
    • This is because the number of particles released in solution is greater for electrolytes
  • One mole of sugar dissolves in water to produce one mole of aqueous sugar molecules.
  • One mole of NaCl dissolves in water to produce two moles of aqueous ions:
    • 1 mole of Na+ and 1 mole of Cl- ions
colligative properties and dissociation of electrolytes79
Colligative Properties and Dissociation of Electrolytes
  • Remember colligative properties depend on the number of dissolved particles.
    • Since NaCl has twice the number of particles we can expect twice the effect for NaCl than for sugar.
  • The table of observed freezing point depressions in the lecture outline shows this effect.
colligative properties and dissociation of electrolytes80
Colligative Properties and Dissociation of Electrolytes
  • Ion pairingor association of ions prevents the effect from being exactly equal to the number of dissociated ions
colligative properties and dissociation of electrolytes81
Colligative Properties and Dissociation of Electrolytes
  • The van’t Hoff factor, symbol i, is used to introduce this effect into the calculations.
  • i is a measure of the extent of ionization or dissociation of the electrolyte in the solution.
colligative properties and dissociation of electrolytes82
i has an ideal value of 2 for 1:1 electrolytes like NaCl, KI, LiBr, etc.

i has an ideal value of 3 for 2:1 electrolytes like K2SO4, CaCl2, SrI2, etc.

Colligative Properties and Dissociation of Electrolytes
colligative properties and dissociation of electrolytes83
Colligative Properties and Dissociation of Electrolytes
  • Example 14-8: The freezing point of 0.0100 m NaCl solution is -0.0360oC. Calculate the van’t Hoff factor and apparent percent dissociation of NaCl in this aqueous solution.
  • meffective = total number of moles of solute particles/kg solvent
  • First let’s calculate the i factor.
colligative properties and dissociation of electrolytes85
Colligative Properties and Dissociation of Electrolytes
  • Next, we will calculate the apparent percent dissociation.
  • Let x = mNaCl that is apparently dissociated.
colligative properties and dissociation of electrolytes89
Colligative Properties and Dissociation of Electrolytes
  • Example 14-9: A 0.0500 m acetic acid solution freezes at -0.0948oC. Calculate the percent ionization of CH3COOH in this solution.

You do it!

osmotic pressure
Osmotic Pressure
  • Osmosis is the net flow of a solvent between two solutions separated by a semipermeable membrane.
    • The solvent passes from the lower concentration solution into the higher concentration solution.
  • Examples of semipermeable membranes include:
    • cellophane and saran wrap
    • skin
    • cell membranes
osmotic pressure92
Osmotic Pressure

semipermeable membrane

H2O

2O

H2O

H2O

sugar dissolved

in water

H2O

H2O

net solvent flow

H2O

H2O

osmotic pressure94
Osmotic Pressure
  • Osmosis is a rate controlled phenomenon.
    • The solvent is passing from the dilute solution into the concentrated solution at a faster rate than in opposite direction, i.e. establishing an equilibrium.
  • The osmotic pressure is the pressure exerted by a column of the solvent in an osmosis experiment.
osmotic pressure95
Osmotic Pressure
  • For very dilute aqueous solutions, molarity and molality are nearly equal.
    • M m
osmotic pressure96
Osmotic Pressure
  • Osmotic pressures can be very large.
    • For example, a 1 M sugar solution has an osmotic pressure of 22.4 atm or 330 p.s.i.
  • Since this is a large effect, the osmotic pressure measurements can be used to determine the molar masses of very large molecules such as:
    • Polymers
    • Biomolecules like
      • proteins
      • ribonucleotides
osmotic pressure97
Osmotic Pressure
  • Example 14-18: A 1.00 g sample of a biological material was dissolved in enough water to give 1.00 x 102 mL of solution. The osmotic pressure of the solution was 2.80 torr at 25oC. Calculate the molarity and approximate molecular weight of the material.

You do it!

osmotic pressure100
Osmotic Pressure

Water Purification by Reverse Osmosis

  • If we apply enough external pressure to an osmotic system to overcome the osmotic pressure, the semipermeable membrane becomes an efficient filter for salt and other dissolved solutes.
    • Ft. Myers, FL gets it drinking water from the Gulf of Mexico using reverse osmosis.
    • US Navy submarines do as well.
    • Dialysis is another example of this phenomenon.
colloids
Colloids
  • Colloids are an intermediate type of mixture that has a particle size between those of true solutions and suspensions.
    • The particles do not settle out of the solution but they make the solution cloudy or opaque.
  • Examples of colloids include:
    • Fog
    • Smoke
    • Paint
    • Milk
    • Mayonnaise
    • Shaving cream
    • Clouds
the tyndall effect
The Tyndall Effect
  • Colloids scatter light when it is shined upon them.
    • Why they appear cloudy or opaque.
    • This is also why we use low beams on cars when driving in fog.
      • See Figure 14-18 in Textbook.
the adsorption phenomenon
The Adsorption Phenomenon
  • Colloids have very large surface areas.
    • They interact strongly with substances near their surfaces.
  • One of the reasons why rivers can carry so much suspended silt in the water.
hydrophilic and hydrophobic colloids
Hydrophilic and Hydrophobic Colloids
  • Hydrophilic colloids like water and are water soluble.
    • Examples include many biological proteins like blood plasma.
  • Hydrophobic colloids dislike water and are water insoluble.
    • Hydrophobic colloids require emulsifying agents to stabilize in water.
  • Homogenized milk is a hydrophobic colloid.
    • Milk is an emulsion of butterfat and protein particles dispersed in water
    • The protein casein is the emulsifying agent.
hydrophilic and hydrophobic colloids105
Hydrophilic and Hydrophobic Colloids
  • Mayonnaise is also a hydrophobic colloid.
    • Mayonnaise is vegetable oil and eggs in a colloidal suspension with water.
    • The protein lecithin from egg yolk is the emulsifying agent.
  • Soaps and detergents are excellent emulsifying agents.
    • Soaps are the Na or K salts of long chain fatty acids.
    • Sodium stearate is an example of a typical soap.
hydrophilic and hydrophobic colloids108
Hydrophilic and Hydrophobic Colloids
  • So called “hard water” contains Fe3+, Ca2+, and/or Mg2+ ions
    • These ions come primarily from minerals that are dissolved in the water.
  • These metal ions react with soap anions and precipitate forming bathtub scum and ring around the collar.
hydrophilic and hydrophobic colloids109
Hydrophilic and Hydrophobic Colloids
  • Synthetic detergents were designed as soap substitutes that do not precipitate in hard water.
    • Detergents are good emulsifying agents.
    • Chemically, we can replace COO- on soaps with sulfonate or sulfate groups
hydrophilic and hydrophobic colloids110
Hydrophilic and Hydrophobic Colloids
  • Linear alkylbenzenesulfonates are good detergents.
synthesis question
Synthesis Question
  • The world’s record for altitude in flying gliders was 60,000 feet for many years. It was set by a pilot in Texas who flew into an updraft in front of an approaching storm. The pilot had to fly out of the updraft and head home not because he was out of air, there was still plenty in the bottle of compressed air on board, but because he did not have a pressurized suit on. What would have happened to this pilot’s blood if he had continued to fly higher?
synthesis question112
Synthesis Question
  • As the pilot flew higher, the atmospheric pressure became less and less. With the lower atmospheric pressure, eventually the blood in the pilot’s veins would have begun to boil. This is a deadly phenomenon which the pilot wisely recognized.
group question
Group Question
  • Medicines that are injected into humans, intravenous fluids and/or shots, must be at the same concentration as the existing chemical compounds in blood. For example, if the medicine contains potassium ions, they must be at the same concentration as the potassium ions in our blood. Such solutions are called isotonic. Why must medicines be formulated in this fashion?
end of chapter 14
End of Chapter 14
  • Human Beings are solution chemistry in action!