chapter 4_3&4
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Prentice Hall 2003 Chapter 13
Chapter 4_3
Physical Transformation of
Pure Substances
David P. White
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Viscosity
Viscosity is the resistance of a liquid to flow.
A liquid flows by sliding molecules over each other.
The stronger the intermolecular forces, the higher the
viscosity.
Example: glycerol C3H8O3(1.49 Ns/m2) and water
(1.01x10-3Ns/m2)
Some Properties of Liquids
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Surface Tension
Surface Tension
Bulk molecules (thosein the liquid) are
equally attracted to
their neighbors.
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Surface Tension
Surface molecules are attracted downwards and sideways by other
molecules but not upwards away from the surface.
Therefore, the surface to tighten like an elastic film Surface tension is the amount of energy required to stretch or
increase the surface area of a liquid.
Liquids with strong intermolecular forces have higher surface
tension Cohesive forcesbind molecules to each other.
Adhesive forcesbind molecules to a surface.
Some Properties of Liquids
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Surface Tension
Meniscusis the shape of the liquid surface.
If adhesive forces are greater than cohesive forces, the liquid
surface is attracted to its container more than the bulkmolecules. Therefore, the meniscus is U-shaped (e.g. water in
glass).
If cohesive forces are greater than adhesive forces, the meniscus
is curved downwards. (e.g mercury) Capil lary Action: When a narrow glass tube is placed in
water, the meniscus pulls the water up the tube.
Some Properties of Liquids
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Transformation from one phase to another,
Occur when energy (usually heat) is added or removed from a
substance Sublimation: solid gas. Hsub> 0 (endothermic).
Vaporization: liquid gas. Hvap> 0 (endothermic).
Melting orfusion: solid liquid. Hfus> 0 (endothermic).
Deposition: gas solid. Hdep< 0 (exothermic).
Condensation: gas liquid. Hcon< 0 (exothermic).
Freezing: liquid solid. Hfre< 0 (exothermic).
Phase Changes
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Phase Changes
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Liquid-vapor equilibrium
Collision rate higher in liquid phase.
Molecules in liquid phase have sufficient energy to escapefrom the surface a phase change occurs
evaporation/vaporization.
Evaporation depends on temperature: higher T, greater kinetic
E, hence more molecules leave the liquid.
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Phase Changes
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Explaining Vapor Pressure on the
Molecular Level
Some of the molecules on the surface of a liquid have
enough energy to escape the attraction of the bulk liquid.
These molecules move into the gas phase.
As the number of molecules in the gas phase increases,
some of the gas phase molecules strike the surface and
return to the liquid. After some time the pressure of the gas will be constant
at the vapor pressure.
Vapor Pressure
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Explaining Vapor Pressure on the
Molecular Level
Vapor Pressure
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Explaining Vapor Pressure on the
Molecular Level
Dynamic Equilibrium: the point when as many molecules
escape the surface as strike the surface.
Vapor pressure is the pressure exerted when the liquidand vapor are in dynamic equilibrium.
Volatility, Vapor Pressure, and Temperature
If equilibrium is never established then the liquidevaporates.
Volatile substances evaporate rapidly.
Vapor Pressure
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Vapor Pressure and Boiling Point
Liquids boil when the external pressure equals the vapor
pressure.
Vapor pressure of a liquid increases with increasingtemperature
Vapor Pressure
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Molar heat of vaporization and boiling point
A measure of strength of intermolecular forces in a liquid is the
molar heat of vaporization Energy required to vaporize 1 mole of a liquid
Strong intermolecular forces in liquid, high molar heat of
vaporization, low vapor pressure
Vapor pressure of a liquid increases with increasing temperature Relationship between vapor pressure of a liquid and the absolute
temperature is given by the Clausius-Clapeyron equation
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Phase Changes
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Clausius-Clapeyron equation
Can determine heat of vaporization by slope.
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Phase Changes
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If we know the values of Hvap and P of a liquid at one
T, can use Clausius-Clapeyron equation to calculate the
vapor pressure of liquid at different T
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Phase Changes
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Example 1
Diethyl ether is a volatile, highly flammable organic
liquid that is used mainly as a solvent. The vapor
pressure of diethyl ether is 401 mmHg at 18o
C. Calculateits vapor pressure at 32oC. Hvap = 26.0 kJ/mol
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Boling point
At boiling point, bubbles form within liquid
When bubble forms the liquid originally occupying that space ispushed aside and level of liquid in container is forced to rise.
The P exerted on the bubble is largely atmospheric pressure.
The P inside bubble is due solely to the vapor pressure of liquid.
when vapor pressure = external pressure, bubble rise to surface ofliquid and burst.
If vapor pressure in bubble lower than external pressure bubble
collapse before it could rise
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Phase Changes
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We can conclude that boiling point of liquid depends on
external pressure.
Bcoz boiling point defined in terms of vapor pressure ofliquid, we expect boiling point to be related to the molar
heat of vaporization.
Higher heat of vaporization, higher boiling point
Both determined by strength of intermolecular forces
Ar and methane (weak dispersion forces) low bp and Hvap
Ethanol and water (H-bonding) high bp and HvapPrentice Hall 2003 Chapter 13
Phase Changes
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Molar heat of fusion- energy required to melt 1 mole of a
solid
Molar heat of sublimation- energy required to sublime 1mole of a solid
Hsub= Hfus+ Hvap
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Phase Changes
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Example 2
Calculate the amount of energy needed to heat 346 g of
liquid water from 0oC to 182oC. Assume that the
specific heat of water is 4.184 J/go
C over the entireliquid range and that the specific heat of steam is 1.99 J/goC.
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Phase diagram: plot of pressure vs. Temperature
summarizing all equilibria between phases.
Given a temperature and pressure, phase diagrams tell uswhich phase will exist.
Any temperature and pressure combination not on a
curve represents a single phase.
Phase Diagrams
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Features of a phase diagram:
Triple point: temperature and pressure at which all three phases
are in equilibrium. Vapor-pressure curve: generally as pressure increases,
temperature increases.
Critical point: critical temperature and pressure for the gas.
Melting point curve: as pressure increases, the solid phase isfavored if the solid is more dense than the liquid.
Normal melting point: melting point at 1 atm.
Phase Diagrams
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Critical Temperature and Pressure
Every substance has a:
Critical temperature, Tc:
Above it gas phase cannot be made to liquefy nomatter how great the applied pressure
Highest T at which a substance can exist as a liquid
Above Tc, there is no fundamental distinction betweenliquid and gas
Critical pressure, Pc: minimum pressure required for
liquefaction at Tc
Phase Diagrams
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Phase Diagrams
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The Phase Diagrams of H2O and CO2
Phase Diagrams
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The Phase Diagrams of H2O and CO2
Water:
The melting point curve slopes to the left because ice is less
dense than water. Triple point occurs at 0.0098C and 4.58 mmHg.
Normal melting (freezing) point is 0C.
Normal boiling point is 100C.
Critical point is 374C and 218 atm.
Phase Diagrams
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The Phase Diagrams of H2O and CO2
Carbon Dioxide:
Triple point occurs at -56.4C and 5.11 atm.
Normal sublimation point is -78.5C. (At 1 atm CO2sublimesit does not melt.)
Critical point occurs at 31.1C and 73 atm.
Phase Diagrams
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Chapter 4_4Properties of Solutions
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A solution is a homogeneous mixture of solute (present in
smallest amount) and solvent (present in largest amount).
Solutes and solvent are components of the solution.
In the process of making solutions with condensed
phases, intermolecular forces become rearranged.
The Solution Process
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The Solution Process
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Energy Changes and Solution
Formation
There are three energy steps in forming a solution:
separation of solute molecules (H1),
separation of solvent molecules (H2), formation of solute-solvent interactions (H3).
We define the enthalpy change in the solution process as
Hsoln= H1+ H2+ H3. Hsolncan either be positive or negative depending on the
intermolecular forces.
The Solution Process
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Energy Changes and Solution
Formation
Breaking attractive intermolecular forces is always
endothermic.
Forming attractive intermolecular forces is alwaysexothermic.
The Solution Process
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Energy Changes and Solution
Formation
To determine whether Hsolnis positive or negative, we
consider the strengths of all solute-solute and solute-
solvent interactions: H1and H2are both positive.
H3is always negative.
It is possible to have either H3> (H1+ H2) or H3< (H1+
H2).
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Energy Changes and Solution
Formation
Examples:
NaOH added to water has Hsoln= -44.48 kJ/mol.
NH4NO3added to water has Hsoln= + 26.4 kJ/mol.
Rule: polar solvents dissolve polar solutes. Non-polar
solvents dissolve non-polar solutes
If Hsolnis 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.
The Solution Process
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Energy Changes and Solution
Formation
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 solute dissolves in a solvent if attraction for its own
molecules are stronger than the solute-solvent attraction
Solution process governed by energy(exothermic/endothermic) and entropy (tendency
towards disorder)
In pure state, the solvent and solute possess a fair degree
of order (regular arrangement of molecules/ions/atoms)
Order is destroyed when the solute dissolves in the
solvent
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The Solution Process
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Solution Formation, Spontaneity, and
Disorder
A spontaneous process occurs without outside
intervention.
When free energy of the system decreases (e.g. droppinga book and allowing it to fall to a lower potential energy),
the process is spontaneous.
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Solution Formation, Spontaneity, and
Disorder
If the process leads to a greater state of disorder, then the
process is spontaneous.
Example: a mixture of CCl4and C6H14is less orderedthan the two separate liquids. Therefore, they
spontaneously mix even though Hsolnis very close to
zero.
There are solutions that form by physical processes and
those by chemical processes.
The Solution Process
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Solution Formation, Spontaneity, and
Disorder
The Solution Process
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Solution Formation and Chemical
Reactions
Example: a mixture of CCl4and C6H14is less ordered
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 isfound only NiCl26H2O. Therefore, Ni dissolution in
HCl is a chemical process.
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Solution Formation and Chemical
Reactions
Example:
NaCl(s) + H2O (l) Na+(aq) + Cl-(aq).
When the water is removed from the solution, NaCl isfound. Therefore, NaCl dissolution is a physical
process.
The Solution Process
S t t d S l ti d
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Solutions can be characterized:
Saturated solution: contains maximum amount of a solute that
will dissolve in a given solvent at a specific temperature Unsaturated solution: contains less solute than it has the
capacity to dissolve
Supersaturated solution: a solution formed when more solute is
dissolved than in a saturated solution. Not stable and solute willcome out from supersaturated solution as crystals
Crystallization: process in which dissolved solute comes out of
solution and forms crystals.
Saturated Solutions and
Solubility
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F t Aff ti
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Solute-Solvent Interaction
Solubility is a measure of how much solute will dissolve in a
solvent at a specific temperature.
like dissolve like helpful in predicting solubility of substance in agiven solvent
Miscible liquids: mix in any proportions.
Immiscible liquids: do not mix.
Polar liquids tend to dissolve in polar solvents. Intermolecular forces are important: water and ethanol are miscible
because the broken hydrogen bonds in both pure liquids are re-
established in the mixture.
Factors Affecting
Solubility
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Solute-Solvent Interaction
The number of carbon atoms in a chain affect solubility:
the more C atoms the less soluble in water.
The number of -OH groups within a molecule increasessolubility in water.
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.
Factors Affecting
Solubility
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Solute-Solvent Interaction
Factors Affecting
Solubility
Factors Affecting
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Solute-Solvent Interaction
Factors Affecting
Solubility
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Example 3
Predict the relative solubilities in the following cases
Bromine (Br2) in benzene (= 0 D) and in water (= 1.87 D)
KCl in carbon tetrachloride (= 0 D) and in liquid ammonia(= 1.46 D)
Formaldehyde (H2C=O) in carbon disulfide (= 0 D) and in
water
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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. Example: SiO2
Factors Affecting
Solubility
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Pressure Effects
Solubility of a gas in a liquid is a function of the pressure
of the gas.
The higher the pressure, the more molecules of gas areclose 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.
Factors Affecting
Solubility
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Pressure Effects
Factors Affecting
Solubility
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Henrys law states that the solubility of a gas in a liquid
is proportional to the pressure of the gas over the
solution
Where c is the molar concentration (mol/L) of dissolved
gas, k is a constant with the units mol/L .atm, andPis
the pressure of the gas over the solution at equilibrium
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Factors Affecting
Solubility
kPc
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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 CO2decreases and the solubility of CO2decreases.
Therefore, bubbles of CO2escape from solution.
Factors Affecting
Solubility
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Example 4
The solubility of nitrogen gas at 25oC and 1 atm is
6.8x10-4mol/L. What is the concentration in molarity of
nitrogen dissolved in water under atmospheric
conditions? The partial pressure of nitrogen gas in the
atmosphere is 0.78 atm
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Factors Affecting
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Temperature Effects
Experience tells us that sugar dissolves better in warm
water than cold.
As temperature increases, solubility of solids generallyincreases.
Sometimes, solubility decreases as temperature increases
(e.g. Ce2
(SO4
)3
).
Factors Affecting
Solubility
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F t Aff ti
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Fractional crystallization: separation of a mixture of substances
into pure components on the basis of their different solubilities
Example: 90 g of KNO3is contaminated with 10g of NaCl.
To purify KNO3, dissolve mixture in 100ml of water at 60oC then
cool to 0oC.
Solubility KNO3and NaCl are 12.1g/100g H2O and 34.2g/100g
H2O
Therefore, 78 g of KNO3will crystallize out of the solution
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Factors Affecting
Solubility
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Temperature Effects
Experience tells us that carbonated beverages go flat as
they get warm.
Therefore, gases get less soluble as temperatureincreases.
Thermal pollution: if lakes get too warm, CO2and O2
become less soluble and are not available for plants or
animals.
Factors Affecting
Solubility
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Ways of Expressing
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Mass Percentage, ppm, and ppb
All methods involve quantifying amount of solute per
amount of solvent (or solution).
Generally amounts or measures are masses, moles orliters.
Qualitatively solutions are dilute or concentrated.
Definitions:
Ways of Expressing
Concentration
100solutionofmasstotal
solutionincomponentofmasscomponentof%mass
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Mass Percentage, ppm, and ppb
Parts per million (ppm) can be expressed as 1 mg ofsolute per kilogram of solution.
If the density of the solution is 1g/mL, then 1 ppm = 1 mg
solute per liter of solution.
Parts per billion (ppb) are 1 g of solute per kilogram of
solution.
Ways of Expressing
Concentration
610solutionofmasstotal
solutionincomponentofmasscomponentofppm
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Mass Percentage, ppm, and ppb
Mole Fraction, Molarity, and Molality
Recall mass can be converted to moles using the molar
mass.
Ways of Expressing
Concentration
910solutionofmasstotal
solutionincomponentofmasscomponentofppb
solutionofmolestotal
solutionincomponentofmolescomponentoffractionMole
solutionofliters
solutemolesMolarity
Ways of Expressing
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Mole Fraction, Molarity, and Molality
We define
Converting between molarity (M) and molality (m)
requires density.
Ways of Expressing
Concentration
solventofkg
solutemolesMolality, m
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Example 5
The density of a 2.45M aqueous solution of methanol
(CH3OH) is 0.976 g/ml. What is the molality of the
solution? Molar mass of methanol is 32.04g/mol
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Ways of Expressing
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The choice of concentration unit is based on the purpose
of the experiment
Advantage of molarity is easier to measure the volume ofsolution than to weigh the solvent. Thus molrity is
usually preferred.
However, molality is independent of temperature.
Volume of solution increases with increasing T so that asolution that is 1.0M at 25oC may become 0.97M at 45oC
Can affect accuracy of experiment
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Ways of Expressing
Concentration
C lli ti P ti
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Colligative properties are properties that depend only on
the number of solute particles in solution and not on the
nature of the solute particles. depend on number of solute particles present.
Vapor-pressure lowering, boiling point elevation, freezing
point depression and osmotic pressure
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Lowering Vapor Pressure
If a solute is nonvolatile the vapour pressure of its solution is
always less than that of the pure solvent.
Thus the relationship between solution and solvent vapor pressure
depends on the concentration of the solute in the solution
Non-volatile solvents reduce the ability of the surface solvent
molecules to escape the liquid.
Therefore, vapor pressure is lowered.
The amount of vapor pressure lowering depends on the amount of
solute.
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Colligative Properties
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Lowering Vapor Pressure
Colligative Properties
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Lowering Vapor Pressure
Raoults Law: vapor pressure of a solvent over a solution,
PAis given by the vapor pressure of the pure solvent, Po
A,
times the mole fraction of the solvent in the solution, XA
In a solution containing only one solute, X1= 1-X2, where
X2 is the mole fraction of the solute.
Colligative Properties
AAA PP
ABAAA PPPP ?
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Lowering Vapor Pressure
Ideal solution: one that obeys Raoults law.
Raoults law breaks down when the solvent-solvent and
solute-solute intermolecular forces are greater thansolute-solvent intermolecular forces.
If both components of a solution are volatile(measureable
vapor pressure) the vapor pressure of the solution is the
sum of the individual partial pressures
PT= PA+ PB
Colligative Properties
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Example 6
Calculate the vapor pressure of a solution made by
dissolving 218 g of glucose (molar mass = 180.2 g/mol)
in 460ml of water at 30oC. What is the vapor pressure
lowering? The vapor pressure of pure water at 30oC is
31.82 mmHg.Assume the density of the solution is
1.00g/ml
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Boiling-Point Elevation
Boiling point is the T at which its vapor pressure equals
the external atmospheric pressure Non-volatile solute lowers the vapor pressure. Therefore,
boiling point of the solution must be affected.
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Colligative Properties
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Colligative Properties
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Boiling-Point Elevation
At 1 atm (normal boiling point of pure liquid) there is a
lower vapor pressure of the solution. Therefore, a higher
temperature is required to teach a vapor pressure of 1 atmfor the solution (Tb).
Molal boiling-point-elevation constant,Kb, (oC/m)
expresses how much Tbchanges with molality, m:
Colligative Properties
mKT bb
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Freezing Point Depression
When a solution freezes, almost pure solvent is formed
first.
Therefore, the sublimation curve for the pure solvent is thesame as for the solution.
Therefore, the triple point occurs at a lower temperature
because of the lower vapor pressure for the solution.
Colligative Properties
Colligative Properties
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Freezing Point Depression
The melting-point (freezing-point) curve is a vertical line
from the triple point.
The solution freezes at a lower temperature (Tf) than thepure solvent.
Decrease in freezing point (Tf) is directly proportional
to molality (Kfis the molal freezing-point-depression
constant):
Colligative Properties
mKT ff
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Freezing involves transition from disordered to ordered
state
Therefore energy must be removed from system Solution has greater disorder than pure solvent, more
energy needs to be removed from it to create order than
pure solvent
Therefore the solution has a lower freezing point than itssolvent
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Example 7
Ethylene glycol is a common automobile antifreeze. It is
water soluble and fairly nonvolatile (b.p 197 oC). Calculate
the freezing point of a solution containing 651 g of this
substance in 2505 g of water. Would you keep this
substance in your car radiator during the summer? The
molar mass of ethylene glycol is 62.01 g/mol.
Kf = 1.86o
C/m and Kb = 0.52o
C/m
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Freezing Point Depression
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Osmosis
Semipermeable membrane: permits passage of some
components of a solution. Example: cell membranes and
cellophane. Osmosis: the movement of a solvent from low solute
concentration to high solute concentration.
There is movement in both directions across a
semipermeable membrane.
As solvent moves across the membrane, the fluid levels
in the arms becomes uneven.
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Eventually thepressure difference between the arms
stops osmosis.
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Osmosis
Osmotic pressure, , is the pressure required to stop
osmosis:
Isotonic solutions: two solutions with the same
separated by a semipermeable membrane.
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MRT
RTV
n
nRTV
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Osmosis
Hypotonic solutions: a solution of lower than a
hypertonic solution.
Osmosis is spontaneous. Red blood cells are surrounded by semipermeable
membranes.
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Osmosis
Crenation:
red blood cells placed in hypertonic solution (relative to
intracellular solution);
there is a lower solute concentration in the cell than the
surrounding tissue;
osmosis occurs and water passes through the membrane out of
the cell. The cell shrivels up.
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Osmosis
Active transport is the movement of nutrients and waste
material through a biological system.
Active transport is not spontaneous.
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