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Syllabus Chemistry 102 Spring 2011 Sec. 501, 503 (MWF 8-8:50, 12:40-1:30) RM 100 HELD Professor: Dr. Earle G. Stone Office: Room 123E Heldenfels (HELD) Telephone: 845-3010 (no voice mail) or leave message at 845-2356 email:

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Chemistry 102 Spring 2011

Sec. 501, 503 (MWF 8-8:50, 12:40-1:30)


Professor: Dr. Earle G. Stone

Office: Room 123E Heldenfels (HELD)

Telephone: 845-3010 (no voice mail) or leave message at 845-2356


(put CHEM 101-Sec. # + subject in subject line of your email)

Office Hours: HELD 408: Tue. And Thurs. 8:00-10:50 AM

I.A. Esther Ocola

S.I. Leader: Mary Hesse


Kotz and Treichel 7th ed. TEXTBOOKS Averill and Eldridge

  • Hardbound ~$200
  • Solution Manual ~$40
  • Online Tutor ~$45
  • Total ~$285
  • Ebook $45 per semester
  • Includes
    • Text
    • Solution manual
    • Online tutorial


Online Dictionary of Chemistry


As A Second Language

General Chemistry I and

Organic Chemistry I

(There are O-chem II and Physics

books in this series if you find these

useful and will have to take those classes.

Yvette Freeman

Publisher's Representative

Pearson Education

Chang’s Essentials


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How grades are determined

The way the real world works

Individual Mastery compared

to a large population

What you are used to and I will


1) Raw scores are determined. Sum of points assigned to correct responses

2) Individual scores are normalized. A context-free evaluation of relative performance

3) Normalized scores are transformed. An absolute score is assigned to a defined scale

4) Letter grades are assigned>89.501 A

>79.501 B

>69.501 C

>59.501 D

<59.501 F



  • Your grade will be based on
    • Four one-hour examinations (200 raw points – T score 100+)
    • A final examination (400 raw points – 2 x T score 100+)

The mere formulation of a problem is far more often essential than its solution, which may be merely a matter of mathematical or experimental skill. To raise new questions, new possibilities, to regard old problems from a new angle requires creative imagination and marks real advances in science.

  • ~Albert Einstein

Problem - A situation that presents difficulty, uncertainty, or perplexity:

Question - A request for data: inquiry, interrogation, query.

Answer - A spoken or written reply, as to a question.

Solution - Something worked out to explain, resolve, or provide a method for dealing with and settling a problem.


c (ms-1)

l (m)

  • Numbers – Significant Figures, Rounding Rules, Accuracy, Precision, Statistical Treatment of the Data
  • Units – 5 of the 7
    • Time – seconds
    • Length – Meters Density?
    • Mass – grams Molecular Weight (Mass)
    • Amount – Moles Mole Ratio, Molarity, molality
    • Temperature – Kelvins
  • Vocabulary – Approximately 100 new terms or words and applying new or more rigid definitions to words you may already own.
  • Principles (Theories and Laws) – Stoichiometry, Quantum Theory, Bonding, Chemical Periodicity, Solutions, Thermodynamics, Intermolecular Forces, Gas Laws, Collogative Properties, Kinetics, Equilibrium, Electrochemistry
  • cp = q/mDT rate = k[A]m[B]n ∆E = q + w
  • DG = DH – TDS Eocell = Ecathode = Eanode
  • PV = nRT %yield = actual/theoretical * 100% K =
  • DT = Kmi



E = n =

inter molecular forces
Inter-molecular Forces

Last Semester you studied INTRAmolecular forces—the forces holding atoms together to form molecules.

Now turn to forces between molecules —INTERmolecular forces.

Forces between molecules, between ions, or between molecules and ions.

ion ion forces for comparison of magnitude
Ion-Ion Forcesfor comparison of magnitude

Na+—Cl- in salt

These are the strongest forces.

Lead to solids with high melting temperatures.

NaCl, mp = 800 oC

MgO, mp = 2800 oC

ion permanent dipole attractions












Ion – Permanent DipoleAttractions

Water is highly polar and can interact with positive ions to give hydrated ions in water.

attraction between ions and permanent dipoles
Attraction Between Ions and Permanent Dipoles

Many metal ions are hydrated. This is the reason metal salts dissolve in water.

attraction between ions and permanent dipoles16
Attraction Between Ions and Permanent Dipoles

Attraction between ions and dipole depends on ion charge and ion-dipole distance.

Measured by ∆H for Mn+ + H2O --> [M(H2O)x]n+

-1922 kJ/mol

-405 kJ/mol

-263 kJ/mol

dipole dipole forces
Dipole-Dipole Forces

Such forces bind molecules having permanent dipoles to one another.

  • Influence of dipole-dipole forces is seen in the boiling points of simple molecules.
  • Compd Mol. Wt. Boil Point
  • N2 28 -196 oC
  • CO 28 -192 oC
  • Br2 160 59 oC
  • ICl 162 97 oC
interactions in liquid solutions
Interactions in Liquid Solutions

• Hydrophilic and hydrophobic solutes

– A solute can be classified as hydrophilic, meaning that there is an electrostatic attraction to water, or hydrophobic, meaning that it repels water.

1. Hydrophilic substance is polar and contains O–H or N–H groups that can form hydrogen bonds to water; tend to be very soluble in water and other strongly polar solvents

2. Hydrophobic substance may be polar but usually contains C–H bonds that do not interact favorably with water; essentially insoluble in water and soluble in nonpolarsolvents

– The difference between hydrophilic and hydrophobic substances has substantial consequences in biological systems.

– Vitamins can be classified as either fat soluble or water soluble.

  • Fat-soluble vitamins (Vitamin A) are nonpolar, hydrophobic molecules and tend to be absorbed into fatty tissues and stored there.
  • Water-soluble vitamins (Vitamin C) are polar, hydrophilic molecules that circulate in the blood and intracellular fluids and are excreted from the body and must be replenished in the daily diet.

– Ionic substances are most stable in polar solvents.

– Water is the most common solvent for ionic compounds because of its high polarity.

– A more useful measure of the ability of a solvent to dissolve ionic compounds is its dielectric constant (), which is the ability of a bulk substance to decrease the electrostatic forces between two charged particles.

hydrogen bonding
Hydrogen Bonding

A special form of dipole-dipole attraction, which enhances dipole-dipole attractions.

  • H-bonding is especially strong in water because
  • the O—H bond is very polar
  • there are 2 lone pairs on the O atom
  • Accounts for many of water’s unique properties.

H-bonding is strongest when X and Y are N, O, or F

hydrogen bonding in h 2 o
Hydrogen Bonding in H2O

Ice has open lattice-like structure.

Ice density is < liquid and so solid floats on water.

One of the VERY few substances where solid is LESS DENSE than the liquid.

H bonds  abnormally high specific heat capacity of water (4.184 J/g•K)

This is the reason water is used to put out fires, it is the reason lakes/oceans control climate, and is the reason thunderstorms release huge energy.


Double helix of DNA

Portion of a DNA chain

H-bonding is especially strong in biological systems — such as DNA.

DNA — helical chains of phosphate groups and sugar molecules. Chains are helical because of tetrahedral geometry of P, C, and O.

Chains bind to one another by specific hydrogen bonding between pairs of Lewis bases.

—adenine with thymine

—guanine with cytosine

forces involving induced dipoles

How can non-polar molecules such as O2 and I2 dissolve in water?

The water dipole INDUCES a dipole in the O2 electric cloud.

Dipole-induced dipole

forces involving induced dipoles26







The alcohol temporarily creates or INDUCES a dipole in I2.
















Consider I2 dissolving in ethanol, CH3CH2OH.

Solubility increases with mass the gas

  • Process of inducing a dipole is polarization
  • Degree to which electron cloud of an atom or molecule can be distorted in its polarizability.
  • Non-polar gases are most soluble in non-polar solvents.
forces involving induced dipoles27

Formation of a dipole in two nonpolar I2 molecules.

The induced forces between I2 molecules are very weak, so solid I2 sublimes (goes from a solid to gaseous molecules).

Induced dipole-induced dipole

forces involving induced dipoles28

The magnitude of the induced dipole depends on the tendency to be distorted.

Higher molec. weight → larger induced dipoles.

Molecule Boiling Point (oC)

CH4 (methane) - 161.5

C2H6 (ethane) - 88.6

C3H8 (propane) - 42.1

C4H10 (butane) - 0.5

intermolecular forces summary
Intermolecular Forces Summary

– London dispersion forces, dipole-dipole interactions, and hydrogen bonds that hold molecules to other molecules are weak.

– Energy is required to disrupt these interactions, and unless some of that energy is recovered in the formation of new, favorable solute-solvent interactions, the increase in entropy on solution formation is not enough for a solution to form.

vapor pressure
Vapor Pressure
  • When the liquid is heated, its molecules obtain sufficient kinetic energy to overcome the forces holding them in the liquid and they escape into the gaseous phase.
  • The result of this phenomenon is that the molecules from the liquid phase generate a population of molecules in the vapor phase above the liquid that produces a pressure called the vapor pressure of the liquid.
  • Plotting the fraction of molecules with a given KE against their KE gives

the KE distribution of the molecules in the liquid.

  • Increasing the temperature increases both

the average KE of the particles in a liquid

and the range of KE for the molecules.

  • Molecules with KE greater than Eo

can escape from the liquid to enter

the vapor phase

  • Molecules must also be at the surface

where it is physically possible for the

molecule to leave the liquid surface.

equilibrium vapor phase
Equilibrium Vapor Phase
  • Volatile liquids have high vapor pressures and tend to evaporate readily; nonvolatile liquids have low vapor pressures and evaporate more slowly.
  • Substances with vapor pressures higher than that of water are volatile, whereas those with vapor pressures lower than water are nonvolatile.
  • Equilibrium vapor pressure of a substance at a particular temperature is a characteristic of the material.

– Does not depend on the amount of liquid

– Depends strongly on the temperature and intermolecular forces

liquids equilibrium vapor pressure









increasing strength of IM interactions

Liquids Equilibrium Vapor Pressure

VP as a function of T.

1. The curves show all conditions of P and T where LIQ and VAP are in EQUILIBRIUM

2. The VP rises with T.

3. When VP = external P, the liquid boils.

This means that BP’s of liquids change with altitude.

4. If external P = 760 mm Hg, T of boiling is the NORMAL BOILING POINT

5. VP of a given molecule at a given T depends on IM forces. Here the VP’s are in the order











HEAT OF VAPORIZATION is the heat req’d (at constant P) to vaporize the liquid.

LIQ + heat VAP

Compd. ∆Hvap (kJ/mol) IM Force

H2O 40.7 (100 oC) H-bonds

SO2 26.8 (-47 oC) dipole

Xe 12.6 (-107 oC) induced dipole

phase diagrams
Phase Diagrams

Lines connect all conditions of T and P where EQUILIBRIUM exists between the phases on either side of the line. (At equilibrium particles move from liquid to gas as fast as they move from gas to liquid, for example.)

phase equilibria water
Phase Equilibria — Water




triple point water
Triple Point — Water

At the TRIPLE POINTall three phases are in equilibrium.

phases diagrams important points for water
Phases Diagrams—Important Points for Water

T(˚C) P(mmHg)

Normal boil point 100 760

Normal freeze point 0 760

Triple point 0.0098 4.58

critical t and p
Critical T and P

As P and T increase, you finally reach the CRITICAL T and P

Above critical T no liquid exists no matter how high the pressure. The gas and liquid phases are intermixed and indistinguishable.

At P < 4.58 mmHg and T < 0.0098 ˚C solid H2O can go directly to vapor. This process is called SUBLIMATION

This is how a frost-free refrigerator works.

terminology and supercritical fluids
Terminology and Supercritical Fluids

• Critical temperature (Tc)

– Temperature above which the gas can no longer be liquefied, regardless of pressure

– The highest temperature at which a substance can exist as a liquid

– Above the critical temperature, the molecules have too much kinetic energy for the intermolecular attractive forces to hold them together in a separate liquid phase

• Critical pressure (Pc)

– Minimum pressure needed to liquefy a substance at the critical temperature

• Critical point — combination of critical temperature and critical pressure

• As the temperature of a liquid increases, its density decreases.

• As the pressure of a gas increases, its density increases.

• At the critical point, the liquid and gas phases have exactly the same density, and only a single phase exists, called a supercritical fluid, which exhibits many of the properties of a gas but has a density typical of a liquid.

solid liquid equilibria















0 °C


Solid-Liquid Equilibria

Raising the pressure at constant T causes water to melt.

The NEGATIVE SLOPE of the S/L line is unique to H2O. Almost everything else has positive slope.

The behavior of water under pressure is an example of LE CHATELIER’S PRINCIPLE At Solid/Liquid equilibrium, raising P squeezes the solid.

It responds by going to phase with greater density, i.e., the liquid phase.

In any system, if you increase P the DENSITY will go up.

Therefore — as P goes up, equilibrium favors phase with the larger density (or SMALLER volume/gram).

changes of state
Changes of State
  • Changes of state are examples of phase changes, or phase transitions
  • Any of the three forms of matter (gas, liquid, solid) is converted to either of the other two
  • Six most common phase changes:

1. melting solid → liquid

2. freezing liquid → solid

3. vaporization liquid → gas

4. condensation gas → liquid

5. sublimation solid → gas

6. deposition gas → solid

temperature curves
Temperature Curves

Heating Curve – A plot of temperature

versus time where heat is added

  • Sample is ice at -23ºC
  • As heat is added, the temperature of the ice increases linearly with time
  • Slope of the line depends on both the mass of the ice and the specific heat of ice, the number of joules required to raise the temperature of 1 g of ice by 1ºC
  • As the temperature of the ice increases, the water molecules in the ice crystal absorb more and more energy and at the melting point, they have enough kinetic energy to overcome attractive forces and to move
  • As more heat is added, the temperature of the system does not increase further, but remains constant at 0ºC until all the ice has melted
  • Once all the ice has been converted to liquid water, the temperature of the water begins to increase – temperature increases more slowly than before because the heat capacity of water is greater than that of ice
  • At 100ºC, water begins to boil; temperature remains constant until all the water has been converted to steam
  • Temperature again begins to rise, but at a faster rate because the heat capacity of steam is less than that of ice or water
temperature curves43
Temperature Curves
  • Cooling curves – A plot of temperature vs. time when heat is removed

1. As heat is removed from steam at 200ºC, the temperature falls until it reaches 100ºC, where the steam begins to condense to liquid water

2. No further temperature change occurs until all the steam is converted to the liquid; then the temperature again decreases as the water is cooled

3. Temperature drops below the freezing point for some time — region corresponds to an unstable form of the liquid, a supercooled liquid that will convert to a solid

4. As water freezes, temperature increases due to heat evolved and then holds constant at the melting point


Units of Concentration

  • There are several different ways to quantitatively describe the concentration of a solution, which is the amount of solute in a given quantity of solution.

1. Molarity– Useful way to describe solution concentrations for reactions that are carried out in solution or for titrations. Volume of a solution depends on its density, which is a function of temperature

Molarity = moles of solute = mol/L = mmol/mL

liter of solution

2. Molality– Concentration of a solution can also be described by its molality (m), the number of moles of solute per kilogram of solvent. Depends on the masses of the solute and solvent, which are independent of temperature.

Molality = moles of solute

kilogram solvent

    • Mole fraction – Used to describe gas concentrations and determine the vapor pressures of mixtures of similar liquids. Depends on only the masses of the solute and solvent and is temperature independent. Mole fractions sum to one for a given mixture.

Mole fraction () = moles of component .

total moles in the solution

4. Mass percentage (%) – The ratio of the mass of the solute to the total mass of the solution. Result can be expressed as mass percentage, parts per million (ppm), or parts per billion (ppb). ppm and ppb are used for highly dilute solutions, and correspond to milligrams (10-3) and micrograms (10-6) of solute per kilogram of solution, respectively

mass percentage = mass of solute 100%

mass of solution

parts per million (ppm)= mass of solute 106

mass of solution

parts per billion (ppb)= mass of solute 109

mass of solution


Units of Concentration

– Mass percentage and parts per million or billion can express the concentrations of substances even if their molecular mass is unknown because these are simply different ways of expressing the ratios of the mass of a solute to the mass of the solution

molality and mole fraction
Molality and Mole Fraction
  • 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 fraction47
Molality and Mole Fraction
  • 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
molality and mole fraction48
Molality and Mole Fraction
  • What are the mole fractions of glucose and water in a 10.0% glucose solution?

Vapor Pressure of Solutions andRaoult’s Law

• The addition of a nonvolatile solute, one whose vapor pressure is too low to measure readily, to a volatile solvent decreases the vapor pressure of the solvent.

• The relationship between solution composition and vapor is known as Raoult’s law,

PA = AP0A,

where PA is the vapor pressure of component A of the solution (the solvent), A is the mole fraction of A in solution, and P0A is the vapor pressure of pure A.

• This equation is used to calculate the actual vapor pressure above a solution of a nonvolatile solute.


Vapor Pressure of Solutions andRaoult’s Law

  • Plots of the vapor pressures of both components versus the mole fractions are straight lines that pass through the origin.
  • A plot of the total vapor pressure of the solution versus the mole fraction is a straight line that represents the sum of the vapor pressures of the pure components.
  • The vapor pressure of the solution is always greater than the vapor pressure of either component.
  • Solutions that obey Raoult’s law are called ideal solutions, in which the intermolecular forces in the two pure liquids are almost identical in both kind and magnitude and the change in enthalpy on solution formation is essentially zero (Hsoln 0).
  • Real solutions exhibit positive or negative deviations from Raoult’s law because the intermolecular interactions between the two components A and B differ.
  • Negative deviation
    • If the A–B interactions are stronger than the A–A and B–B interactions, each component of the solution exhibits a lower vapor pressure than expected for an ideal solution, as does the solution as a whole.
    • The favorable A–B interactions stabilize the solution compared with the vapor.
  • Positive deviation
    • If the A–B interactions are weaker than the A–A and B–B interactions yet the entropy increase is enough to allow the solution to form, both A and B have an increased tendency to escape from the solution into the vapor phase.
    • The result is a higher vapor pressure than expected for an ideal solution.
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.
  • This graph shows how the solution’s vapor pressure is changed by the mole fraction of the solute, which is Raoult’s law.
effect of temperature on the solubility
Effect of Temperature on the Solubility
  • Solubility of a substance generally increases with increasing temperature.
  • No relationship between the structure of a substance and temperature dependence.
  • Solubility may increase or decrease with temperature; the magnitude varies widely among compounds.
  • The variation of solubility with temperature is used to separate the components of a mixture by fractional crystallization, a technique for purifying compounds.
  • For fractional crystallization to be effective the compound must be more soluble at high T than at low T: lowering the temperature causes it to crystallize out of solution.
  • Solubility of gases in liquids decreases with increasing temperature
  • Attractive intermolecular interactions in the gas phase are essentially zero for most substances.
  • When a gas dissolves, its molecules interact with solvent molecules and heat is released when these new attractive interactions form, therefore, dissolving most gases in liquids is an exothermic process (Hsoln < 0)
  • Adding heat to the solution provides thermal energy that overcomes the attractive forces between the gas and the solvent molecules, thereby decreasing the solubility of the gas
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.
pressure and solubility of gases henry s law
Pressure and Solubility of Gases: Henry’s Law
  • External pressure has little effect on the solubility of liquids and solids. The solubility of gases increases as the partial pressure of the gas above a solution increases.
  • The concentration of molecules in the gas phase increases with increasing pressure, as does the concentration of dissolved gas molecules in the solution at equilibrium at higher pressures.
  • Henry’s Law:C = kP, where C is the concentration of dissolved gas at equilibrium; P is the partial pressure; and k is the Henry’s law constant, which is determined experimentally for each combination of gas, solvent, and temperature: units are M/atm.
  • Concentration of a gas in water at a given pressure depends strongly on its physical properties.
  • Gases that react chemically with water do not obey Henry’s law; all of these gases are much more soluble than predicted by Henry’s law
colligative properties of solutions
ColligativeProperties of Solutions
  • The colligative properties of a solution depend primarily on the number of solute particles present rather than the kind of particles.

• Not all solutions with the same molarity contain the same concentration of solute particles.

• The sum of the concentrations of the various dissolved solute particles dictates the physical properties of a solution.

  • Colligative properties do not depend on the kinds of particles dissolved.
  • Colligative properties are a physical property of solutions.
  • There are four colligative properties:

1. Boiling-point elevation

2. Freezing-Point depression

3. Vapor pressure

4. Osmotic pressure

  • Vapor pressure lowering is the key to all four of the colligative properties.
boiling point elevation
Boiling-Point Elevation
  • The normal boiling point of a substance is the temperature at which the vapor pressure equals 1 atm.
  • Because the vapor pressure of the solution at a given temperature is lower than the vapor pressure of the pure solvent, achieving a vapor pressure of 1 atm for the solution requires a higher temperature than the normal boiling point of the solvent.
  • The boiling point of a solution is always higher than that of the pure solvent.
  • The magnitude of the increase in the boiling point is related to the magnitude of the decrease in the vapor pressure; the decrease in the vapor pressure is proportional to the concentration of the solute in solution.
boiling point elevation57
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 elevation relationship is:
freezing point depression
Freezing Point Depression
  • Dissolving a nonvolatile solute in water not only raises the boiling point of the water but also lowers its freezing point.
  • Freezing-point depression depends on the total number of dissolved nonvolatile solute particles.
  • The molar mass of an unknown compound can be determined by measuring the freezing point of a solution that contains a known mass of solute, which is accurate for dilute solutions.
  • Changes in the boiling point are smaller than changes in the freezing point for a given solvent, so boiling point elevations are difficult to measure and are not used to determine molar mass.
  • The freezing-point depression (Tf) is defined as the difference between the freezing point of the pure solvent and the freezing point of the solution:
  • Relationship for freezing point depression is:
colligative properties of electrolyte solutions
Colligative Properties of Electrolyte Solutions
  • The relationship between the actual number of moles of solute added to form a solution and the apparent number as determined by colligative properties is called the van’tHoff factor (i) and is defined as

i = apparent number of particles in solution

number of moles of solute dissolved

  • The van’t Hoff factors for ionic compounds are lower than expected, their solutions apparently contain fewer particles than predicted by the number of ions per formula unit.
  • Deviation from the expected value increases as the concentration of the solute increases because ionic compounds do not totally dissociate in aqueous solution.
  • Some of the ions exist as ion pairs, for a brief time are associated with each.
  • The actual number of solvated ions present in a solution can be determined by measuring a colligative property at several solute concentrations.
boiling point elevation freezing point depression
Fundamentally, freezing point depression and boiling point elevation are the same phenomenon.

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

The effects are easily seen in a phase diagram

Boiling Point Elevation Freezing Point Depression
  • Notice the similarity of the two relationships for freezing point depression and boiling point elevation.
boiling point elevation62
Boiling Point Elevation
  • What is the normal boiling point of a 2.50 m glucose, C6H12O6, solution?
freezing point depression63
Freezing Point Depression
  • 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.
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 depression65
Determination of Molecular Weight by Freezing Point Depression
  • 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?
osmotic pressure
Osmotic Pressure
  • When a solution and a pure solvent are separated by a semipermeable membrane, a barrier that allows solvent molecules but not solute molecules to pass through, the flow of solvent in opposing directions is unequal and produces an osmotic pressure, which is the difference in pressure between the two sides of the membrane.
  • Because of the large magnitude of osmotic pressures, osmosis is important in biochemistry, biology, and medicine; every barrier that separates an organism or cell from its environment acts like a semipermeable membrane, permitting the flow of water but not solutes.
  • Examples of semipermeable membranes include:
  • cellophane and saran wrap
  • skin
  • cell membranes
  • Dialysis uses a semipermeable membrane.
  • Preserving fruits and meats employing cell membranes and using salt prevents bacterial growth.
  • Reverse Osmosis can be used to produce pure water from seawater.
osmotic pressure67
Osmotic Pressure
  • Osmosis is the net flow of solvent through a membrane due to different solute concentrations; the direction of net solvent flow is always from higher solvent concentration to lower solvent concentration or from the side with the lower concentration of solute to the side with the higher solute concentration.
  • Osmosis is a rate controlled phenomenon. The flow from high to low solvent is faster than low to high solvent concentration until equilibrium is reached.
  • Osmotic pressure () of a solution depends on the concentration of dissolved solute particles:
osmotic pressure68
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 pressure69
Osmotic Pressure
  • 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.
solutions of solids
Solutions of Solids

• Solutions are not limited to gases and liquids; solid solutions also exist.

• Amalgams, which are usually solids, are solutions of metals in liquid mercury.

• Network solids are insoluble in all solvents with which they do not react chemically; covalent bonds that hold the network together are too strong to be broken and are much stronger than any combination of intermolecular interactions that might occur in solution.

• Most metals are insoluble in all solvents but do react with solutions such as aqueous acid or base to produce a solution; in these cases the metal undergoes a chemical transformation that cannot be reversed by removing the solvent.

aggregate particles in aqueous solution
Aggregate Particles in Aqueous Solution

• A mixture of gases is the only combination of substances that cannot produce a suspension or colloid because their particles are small and form true solutions

  • A colloid can be classified as:
    • 1. a sol or gel, a dispersion of solid particles in a liquid or solid in which all the solvent has been absorbed by the solid particles, thus preventing the mixture from flowing readily;
    • 2. An aerosol, a dispersion of solid or liquid particles in a gas;
    • 3. An emulsion, a dispersion of one liquid phase in another liquid with which it is immiscible.
interactions in liquid solutions72
Interactions in Liquid Solutions

– All common organic liquids, whether polar or not, are miscible; the strengths of the intermolecular attractions are comparable, the enthalpy of solution is small, and the increase in entropy drives the formation of a solution.

– If predominant intermolecular interactions in two liquids are very different from one another, they may be immiscible, and when shaken with water, they form separate phases or layers separated by an interface.

– Only the three lightest alcohols are completely miscible with water; as the molecular mass of the alcohol increases, so does the proportion of hydrocarbon in the molecule, which leads to fewer favorable electrostatic interactions with water



• Emulsions are colloids formed by the dispersion of a hydrophobic liquid in water, thereby bringing two mutually insoluble liquids in close contact.

• Various agents have been developed to stabilize emulsions, the most successful being molecules that combine a relatively long hydrophobic “tail” with a hydrophilic “head.” Examples of emulsifying agents include soaps and detergents.


  • • Micelles
    • – In the absenceof a dispersed hydrophobic liquid phase, solutions of detergents in water form organized spherical or cylindrical aggregates called micelles, which minimize contact between the hydrophobic tails and water.
    • – In a micelle, only the hydrophilic heads are in direct contact with water, and the hydrophobic tails are in the interior of the aggregate.
  • • Phospholipids – a large class of biological molecules that consist of detergent-like molecules that contain a hydrophilic head and two hydrophobic tails; additional tail results in a cylindrical shape that prevents phospholipids from forming a spherical micelle
  • • Cells
    • – are collections of molecules that are surrounded by a phospholipidbilayer called a cell membrane and are able to reproduce themselves