Solubilitychapter 13 Brownlemay
Transcript of Solubilitychapter 13 Brownlemay
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Solutions
Chapter 13Properties of Solutions
Adapted by SA Green from:
John D. Bookstaver
St. Charles Community College
St. Peters, MO
2006, Prentice Hall, Inc.
Chemistry, The Central Science, 10th editionTheodore L. Brown; H. Eugene LeMay, Jr.;
and Bruce E. Bursten
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Solutions
Solutions
• Solutions are homogeneous mixtures of two or more pure substances.
• In a solution, the solute is dispersed uniformly throughout the solvent.
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Solutions
Solutions
How does a solid dissolve into a liquid?
What ‘drives’ the dissolution process?
What are the energetics of dissolution?
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How Does a Solution Form?1. Solvent molecules attracted to surface ions.2. Each ion is surrounded by solvent molecules.3. Enthalpy (H) changes with each interaction broken or
formed.
Ionic solid dissolving in water
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Solutions
How Does a Solution Form?1. Solvent molecules attracted to surface ions.2. Each ion is surrounded by solvent molecules.3. Enthalpy (H) changes with each interaction broken or
formed.
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Solutions
How Does a Solution Form
The ions are solvated (surrounded by solvent).
If the solvent is water, the ions are hydrated.
The intermolecular force here is ion-dipole.
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Solutions
Energy Changes in SolutionTo determine the enthalpy
change, we divide the process into 3 steps. 1. Separation of solute
particles.
2. Separation of solvent particles to make ‘holes’.
3. Formation of new interactions between solute and solvent.
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Enthalpy Changes in Solution
The enthalpy change of the overall process depends on H for each of these steps.
Start
End
EndStart
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Solutions
Enthalpy changes during dissolution
The enthalpy of solution, Hsoln, can be either positive or negative.
Hsoln = H1 + H2 + H3
Hsoln (MgSO4)= -91.2 kJ/mol --> exothermic
Hsoln (NH4NO3)= 26.4 kJ/mol --> endothermic
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Solutions
Why do endothermic processes sometimes occur spontaneously?
Some processes, like the dissolution of NH4NO3 in water, are spontaneous at room temperature even though heat is absorbed, not released.
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Solutions
Enthalpy Is Only Part of the Picture
Entropy is a measure of: • Dispersal of energy in
the system.• Number of microstates
(arrangements) in the system.
b. has greater entropy, is the favored state
(more on this in chap 19)
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Solutions
Entropy changes during dissolution
Each step also involves a change in entropy. 1. Separation of solute
particles.
2. Separation of solvent particles to make ‘holes’.
3. Formation of new interactions between solute and solvent.
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SAMPLE EXERCISE 13.1 Assessing Entropy ChangeIn the process illustrated below, water vapor reacts with excess solid sodium sulfate to form the hydrated form of the salt. The chemical reaction is
Does the entropy of the system increase or decrease?
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Solutions
Dissolution vs reaction
• Dissolution is a physical change—you can get back the original solute by evaporating the solvent.
• If you can’t, the substance didn’t dissolve, it reacted.
Ni(s) + HCl(aq) NiCl2(aq) + H2(g) NiCl2(s)dry
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Degree of saturation
• Saturated solutionSolvent holds as much
solute as is possible at that temperature.
Undissolved solid remains in flask.
Dissolved solute is in dynamic equilibrium with solid solute particles.
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Solutions
Degree of saturation
• Unsaturated SolutionLess than the
maximum amount of solute for that temperature is dissolved in the solvent.
No solid remains in flask.
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Degree of saturation
• SupersaturatedSolvent holds more solute than is normally
possible at that temperature.These solutions are unstable; crystallization can
often be stimulated by adding a “seed crystal” or scratching the side of the flask.
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Solutions
Degree of saturation
Unsaturated, Saturated or Supersaturated?
How much solute can be dissolved in a solution?
More on this in Chap 17(solubility products, p 739)
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Factors Affecting Solubility
• Chemists use the axiom “like dissolves like”:Polar substances tend to
dissolve in polar solvents.Nonpolar substances tend
to dissolve in nonpolar solvents.
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Solutions
Factors Affecting Solubility
The stronger the intermolecular attractions between solute and solvent, the more likely the solute will dissolve.Example: ethanol in water
Ethanol = CH3CH2OH
Intermolecular forces = H-bonds; dipole-dipole; dispersion
Ions in water also have ion-dipole forces.
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Solutions
Factors Affecting Solubility
Glucose (which has hydrogen bonding) is very soluble in water.
Cyclohexane (which only has dispersion forces) is not water-soluble.
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Factors Affecting Solubility• Vitamin A is soluble in nonpolar compounds
(like fats).• Vitamin C is soluble in water.
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Which vitamin is water-soluble and which is fat-soluble?
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Gases in Solution
• In general, the solubility of gases in water increases with increasing mass.
Why?• Larger molecules
have stronger dispersion forces.
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Gases in Solution
• The solubility of liquids and solids does not change appreciably with pressure.
• But, the solubility of a gas in a liquid is directly proportional to its pressure.
Increasing pressure above solution forces more gas to dissolve.
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Solutions
Henry’s Law
Sg = kPg
where
• Sg is the solubility of the gas;
• k is the Henry’s law constant for that gas in that solvent;
• Pg is the partial pressure of the gas above the liquid.
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Temperature
Generally, the solubility of solid solutes in liquid solvents increases with increasing temperature.
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Temperature• The opposite is true of
gases. Higher temperature drives gases out of solution.
Carbonated soft drinks are more “bubbly” if stored in the refrigerator.
Warm lakes have less O2 dissolved in them than cool lakes.
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Chap 13: Ways of Expressing Ways of Expressing
Concentrations of Concentrations of SolutionsSolutions
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Mass Percentage
Mass % of A =mass of A in solutiontotal mass of solution
100
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Parts per Million andParts per Billion
ppm =mass of A in solutiontotal mass of solution
106
Parts per Million (ppm)
Parts per Billion (ppb)
ppb =mass of A in solutiontotal mass of solution
109
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moles of Atotal moles in solution
XA =
Mole Fraction (X)
• In some applications, one needs the mole fraction of solvent, not solute—make sure you find the quantity you need!
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mol of soluteL of solution
M =
Molarity (M)
• Because volume is temperature dependent, molarity can change with temperature.
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mol of solutekg of solvent
m =
Molality (m)
Because neither moles nor mass change with temperature, molality (unlike molarity) is not temperature dependent.
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Mass/M
ass
Mol
es/M
oles
Moles/Mass
Moles/L
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SAMPLE EXERCISE 13.4 Calculation of Mass-Related Concentrations
(a) A solution is made by dissolving 13.5 g of glucose (C6H12O6) in 0.100 kg of water. What is the mass percentage of solute in this solution? (b) A 2.5-g sample of groundwater was found to contain 5.4g of Zn2+ What is the concentration of Zn2+ in parts per million?PRACTICE EXERCISE(a) Calculate the mass percentage of NaCl in a solution containing 1.50 g of NaCl in 50.0 g of water. (b) A commercial bleaching solution contains 3.62 mass % sodium hypochlorite, NaOCl. What is the mass of NaOCl in a bottle containing 2500 g of bleaching solution?
PRACTICE EXERCISE
A commercial bleach solution contains 3.62 mass % NaOCl in water. Calculate (a) the molality and (b) the mole fraction of NaOCl in the solution.
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Colligative Properties
• Colligative properties depend only on the number of solute particles present, not on the identity of the solute particles.
• Among colligative properties areVapor pressure lowering Boiling point elevationMelting point depressionOsmotic pressure
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Vapor Pressure
As solute molecules are added to a solution, the solvent become less volatile (=decreased vapor pressure).
Solute-solvent interactions contribute to this effect.
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Vapor Pressure
Therefore, the vapor pressure of a solution is lower than that of the pure solvent.
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Raoult’s Law
PA = XAPAwhere• XA is the mole fraction of compound A
• PA is the normal vapor pressure of A at that temperature
NOTE: This is one of those times when you want to make sure you have the vapor pressure of the solvent.
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SAMPLE EXERCISE 13.8 Calculation of Vapor-Pressure Lowering
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 (Appendix B).
PRACTICE EXERCISEThe 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?
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Boiling Point Elevation and Freezing Point Depression
Solute-solvent interactions also cause solutions to have higher boiling points and lower freezing points than the pure solvent.
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Boiling Point ElevationThe change in boiling point is proportional to the molality of the solution:
Tb = Kb m
where Kb is the molal boiling point elevation constant, a property of the solvent.Tb is added to the normal
boiling point of the solvent.
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Freezing Point Depression• The change in freezing
point can be found similarly:
Tf = Kf m
• Here Kf is the molal freezing point depression constant of the solvent.
Tf is subtracted from the normal freezing point of the solvent.
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Boiling Point Elevation and Freezing Point Depression
In both equations, T does not depend on what the solute is, but only on how many particles are dissolved.
Tb = Kb m
Tf = Kf m
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Colligative Properties of Electrolytes
Because these properties depend on the number of particles dissolved, solutions of electrolytes (which dissociate in solution) show greater changes than those of nonelectrolytes.
e.g. NaCl dissociates to form 2 ion particles; its limiting van’t Hoff factor is 2.
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Colligative Properties of Electrolytes
However, a 1 M solution of NaCl does not show twice the change in freezing point that a 1 M solution of methanol does.
It doesn’t act like there are really 2 particles.
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van’t Hoff Factor
One mole of NaCl in water does not really give rise to two moles of ions.
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van’t Hoff Factor
Some Na+ and Cl− reassociate as hydrated ion pairs, so the true concentration of particles is somewhat less than two times the concentration of NaCl.
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The van’t Hoff Factor
• Reassociation is more likely at higher concentration.
• Therefore, the number of particles present is concentration dependent.
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The van’t Hoff Factor
We modify the previous equations by multiplying by the van’t Hoff factor, i
Tf = Kf m i
i = 1 for non-elecrtolytes
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Osmosis
• Semipermeable membranes allow some particles to pass through while blocking others.
• In biological systems, most semipermeable membranes (such as cell walls) allow water to pass through, but block solutes.
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OsmosisIn osmosis, there is net movement of solvent from the area of higher solvent concentration (lower solute concentration) to the are of lower solvent concentration (higher solute concentration).
Water tries to equalize the concentration on both sides until pressure is too high.
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Osmotic Pressure
• The pressure required to stop osmosis, known as osmotic pressure, , is
nV
= ( )RT = MRT
where M is the molarity of the solution
If the osmotic pressure is the same on both sides of a membrane (i.e., the concentrations are the same), the solutions are isotonic.
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Osmosis in Blood Cells
• If the solute concentration outside the cell is greater than that inside the cell, the solution is hypertonic.
• Water will flow out of the cell, and crenation results.
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Osmosis in Cells
• If the solute concentration outside the cell is less than that inside the cell, the solution is hypotonic.
• Water will flow into the cell, and hemolysis results.
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Molar Mass from Colligative Properties
We can use the effects of a colligative property such as osmotic pressure to determine the molar mass of a compound.
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Colloids:
Suspensions of particles larger than individual ions or molecules, but too small to be settled out by gravity.
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Tyndall Effect
• Colloidal suspensions can scatter rays of light.
• This phenomenon is known as the Tyndall effect.
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Colloids in Biological Systems
Some molecules have a polar, hydrophilic (water-loving) end and a nonpolar, hydrophobic (water-hating) end.
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Colloids in Biological Systems
Sodium stearate is one example of such a molecule.
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Colloids in Biological Systems
These molecules can aid in the emulsification of fats and oils in aqueous solutions.
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END Chap 13