1. Chapter 13 Properties of Solutions
CHEMISTRY The Central Science 9th Edition
Chapter 13 Properties of Solutions
David P. White
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Chapter 13

2. The Solution Process
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.
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3. The Solution Process
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4. The Solution Process
Consider NaCl (solute) dissolving in water (solvent):
the water H-bonds have to be interrupted,
NaCl dissociates into Na+ and Cl-,
ion-dipole forces form: Na+ … -OH2 and Cl- … +H2O.
We say the ions are solvated by water.
If water is the solvent, we say the ions are hydrated.
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6. The Solution Process

7. The Solution Process Energy Changes and Solution Formation
There are three energy steps in forming a solution:
separation of solute molecules (H1),
separation of solvent molecules (H2), and formation of solute-solvent interactions (H3).
We define the enthalpy change in the solution process as
Hsoln = H1 + H2 + H3.
Hsoln can either be positive or negative depending on the intermolecular forces.
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9. The Solution Process Energy Changes and Solution Formation
Breaking attractive intermolecular forces is always endothermic.
Forming attractive intermolecular forces is always exothermic.
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10. The Solution Process Energy Changes and Solution Formation
To determine whether Hsoln is positive or negative, we consider the strengths of all solute-solute and solute-solvent interactions:
H1 and H2 are both positive.
H3 is always negative.
It is possible to have either H3 > (H1 + H2) or H3 < (H1 + H2).
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12. The Solution Process Energy Changes and Solution Formation Examples:
NaOH added to water has Hsoln = kJ/mol.
NH4NO3 added to water has Hsoln = kJ/mol.
“Rule”: polar solvents dissolve polar solutes. Non-polar solvents dissolve non-polar solutes. Why?
If Hsoln is too endothermic a solution will not form.
NaCl in gasoline: the ion-dipole forces are weak because gasoline is non-polar. Therefore, the ion-dipole forces do not compensate for the separation of ions.
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13. The Solution Process 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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14. The Solution Process Solution Formation, Spontaneity, and Disorder
A spontaneous process occurs without outside intervention.
When energy of the system decreases (e.g. dropping a book and allowing it to fall to a lower potential energy), the process is spontaneous.
Some spontaneous processes do not involve the system moving to a lower energy state (e.g. an endothermic reaction).
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15. The Solution Process Solution Formation, Spontaneity, and Disorder
If the process leads to a greater state of disorder, then the process is spontaneous.
Example: a mixture of CCl4 and C6H14 is less ordered than the two separate liquids. Therefore, they spontaneously mix even though Hsoln is very close to zero.
There are solutions that form by physical processes and those by chemical processes.
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16. The Solution Process Solution Formation, Spontaneity, and Disorder
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17. Ni(s) + 2HCl(aq)  NiCl2(aq) + H2(g).
The Solution Process
Solution Formation and Chemical Reactions
Example: a mixture of CCl4 and C6H14 is 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 is found only NiCl2·6H2O. Therefore, Ni dissolution in HCl is a chemical process.
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18. NaCl(s) + H2O (l)  Na+(aq) + Cl-(aq).
The Solution Process
Solution Formation and Chemical Reactions
Example:
NaCl(s) + H2O (l)  Na+(aq) + Cl-(aq).
When the water is removed from the solution, NaCl is found. Therefore, NaCl dissolution is a physical process.
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19. Saturated Solutions and Solubility
Dissolve: solute + solvent  solution.
Crystallization: solution  solute + solvent.
Saturation: crystallization and dissolution are in equilibrium.
Solubility: amount of solute required to form a saturated solution.
Supersaturated: a solution formed when more solute is dissolved than in a saturated solution.
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21. Factors Affecting Solubility
Solute-Solvent Interaction
Polar liquids tend to dissolve in polar solvents.
Miscible liquids: mix in any proportions.
Immiscible liquids: do not mix.
Intermolecular forces are important: water and ethanol are miscible because the broken hydrogen bonds in both pure liquids are re-established in the mixture.
The number of carbon atoms in a chain affect solubility: the more C atoms the less soluble in water.
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22. Factors Affecting Solubility
Solute-Solvent Interaction
The number of -OH groups within a molecule increases solubility in water.
Generalization: “like dissolves like”.
The more polar bonds in the molecule, the better it dissolves in a polar solvent.
The less polar the molecule the less it dissolves in a polar solvent and the better is dissolves in a non-polar solvent.
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23. Factors Affecting Solubility
Solute-Solvent Interaction
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24. Factors Affecting Solubility
Solute-Solvent Interaction
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25. Factors Affecting Solubility
Solute-Solvent Interaction
Network solids do not dissolve because the strong intermolecular forces in the solid are not re-established in any solution.
Pressure Effects
Solubility of a gas in a liquid is a function of the pressure of the gas.
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26. Factors Affecting Solubility
Pressure Effects
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27. Factors Affecting Solubility
Pressure Effects
The higher the pressure, the more molecules of gas are close to the solvent and the greater the chance of a gas molecule striking the surface and entering the solution.
Therefore, the higher the pressure, the greater the solubility.
The lower the pressure, the fewer molecules of gas are close to the solvent and the lower the solubility.
If Sg is the solubility of a gas, k is a constant, and Pg is the partial pressure of a gas, then Henry’s Law gives:
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28. Factors Affecting Solubility
Pressure Effects
Carbonated beverages are bottled with a partial pressure of CO2 > 1 atm.
As the bottle is opened, the partial pressure of CO2 decreases and the solubility of CO2 decreases.
Therefore, bubbles of CO2 escape from solution.
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29. Factors Affecting Solubility
Temperature Effects
Experience tells us that sugar dissolves better in warm water than cold.
As temperature increases, solubility of solids generally increases.
Sometimes, solubility decreases as temperature increases (e.g. Ce2(SO4)3).
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31. Factors Affecting Solubility
Temperature Effects
Experience tells us that carbonated beverages go flat as they get warm.
Therefore, gases get less soluble as temperature increases.
Thermal pollution: if lakes get too warm, CO2 and O2 become less soluble and are not available for plants or animals.
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33. Ways of Expressing Concentration
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 or liters.
Qualitatively solutions are dilute or concentrated.
Definitions:
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34. Ways of Expressing Concentration
Mass Percentage, ppm, and ppb
Parts per million (ppm) can be expressed as 1 mg of solute 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.
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35. Ways of Expressing Concentration
Mass Percentage, ppm, and ppb
Mole Fraction, Molarity, and Molality
Recall mass can be converted to moles using the molar mass.
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36. Ways of Expressing Concentration
Mole Fraction, Molarity, and Molality
We define
Converting between molarity (M) and molality (m) requires density.
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38. Colligative Properties
Colligative properties depend on quantity of solute molecules. (E.g. freezing point depression and melting point elevation.)
Lowering Vapor Pressure
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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39. Colligative Properties
Lowering Vapor Pressure
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40. Colligative Properties
Lowering Vapor Pressure
Raoult’s Law: PA is the vapor pressure with solute, PA is the vapor pressure without solvent, and A is the mole fraction of A, then
Recall Dalton’s Law:
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41. Colligative Properties
Lowering Vapor Pressure
Ideal solution: one that obeys Raoult’s law.
Raoult’s law breaks down when the solvent-solvent and solute-solute intermolecular forces are greater than solute-solvent intermolecular forces.
Boiling-Point Elevation
Goal: interpret the phase diagram for a solution.
Non-volatile solute lowers the vapor pressure.
Therefore the triple point - critical point curve is lowered.
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43. Colligative Properties
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 atm for the solution (Tb).
Molal boiling-point-elevation constant, Kb, expresses how much Tb changes with molality, m:
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44. Colligative Properties
Freezing Point Depression
At 1 atm (normal boiling point of pure liquid) there is no depression by definition
When a solution freezes, almost pure solvent is formed first.
Therefore, the sublimation curve for the pure solvent is the same as for the solution.
Therefore, the triple point occurs at a lower temperature because of the lower vapor pressure for the solution.
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45. Colligative Properties
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 the pure solvent.
Decrease in freezing point (Tf) is directly proportional to molality (Kf is the molal freezing-point-depression constant):
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46. Colligative Properties
Freezing Point Depression
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48. Colligative Properties
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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49. Colligative Properties
Osmosis
Eventually the pressure difference between the arms stops osmosis.
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50. Colligative Properties
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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51. Colligative Properties
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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52. Colligative Properties
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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53. Colligative Properties
Osmosis
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54. Colligative Properties
Osmosis
Hemolysis:
red blood cells placed in a hypotonic solution;
there is a higher solute concentration in the cell;
osmosis occurs and water moves into the cell.
The cell bursts.
To prevent crenation or hemolysis, IV (intravenous) solutions must be isotonic.
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55. Colligative Properties
Osmosis
Cucumber placed in NaCl solution loses water to shrivel up and become a pickle.
Limp carrot placed in water becomes firm because water enters via osmosis.
Salty food causes retention of water and swelling of tissues (edema).
Water moves into plants through osmosis.
Salt added to meat or sugar to fruit prevents bacterial infection (a bacterium placed on the salt will lose water through osmosis and die).
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56. Colligative Properties
Osmosis
Active transport is the movement of nutrients and waste material through a biological system.
Active transport is not spontaneous.
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57. Colloids
Colloids are suspensions in which the suspended particles are larger than molecules but too small to drop out of the suspension due to gravity.
Particle size: 10 to 2000 Å.
There are several types of colloid:
aerosol (gas + liquid or solid, e.g. fog and smoke),
foam (liquid + gas, e.g. whipped cream),
emulsion (liquid + liquid, e.g. milk),
sol (liquid + solid, e.g. paint),
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58. Colloids
solid foam (solid + gas, e.g. marshmallow),
solid emulsion (solid + liquid, e.g. butter),
solid sol (solid + solid, e.g. ruby glass).
Tyndall effect: ability of a Colloid to scatter light. The beam of light can be seen through the colloid.
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59. Colloids
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60. Colloids Hydrophilic and Hydrophobic Colloids
Focus on colloids in water.
“Water loving” colloids: hydrophilic.
“Water hating” colloids: hydrophobic.
Molecules arrange themselves so that hydrophobic portions are oriented towards each other.
If a large hydrophobic macromolecule (giant molecule) needs to exist in water (e.g. in a cell), hydrophobic molecules embed themselves into the macromolecule leaving the hydrophilic ends to interact with water.
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62. Colloids Hydrophilic and Hydrophobic Colloids
Typical hydrophilic groups are polar (containing C-O, O-H, N-H bonds) or charged.
Hydrophobic colloids need to be stabilized in water.
Adsorption: when something sticks to a surface we say that it is adsorbed.
If ions are adsorbed onto the surface of a colloid, the colloids appears hydrophilic and is stabilized in water.
Consider a small drop of oil in water.
Add to the water sodium stearate.
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63. Colloids Hydrophilic and Hydrophobic Colloids Prentice Hall © 2003
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64. Colloids Hydrophilic and Hydrophobic Colloids
Sodium stearate has a long hydrophobic tail (CH3(CH2)16-) and a small hydrophobic head (-CO2-Na+).
The hydrophobic tail can be absorbed into the oil drop, leaving the hydrophilic head on the surface.
The hydrophilic heads then interact with the water and the oil drop is stabilized in water.
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65. Colloids

66. Colloids Hydrophilic and Hydrophobic Colloids
Most dirt stains on people and clothing are oil-based. Soaps are molecules with long hydrophobic tails and hydrophilic heads that remove dirt by stabilizing the colloid in water.
Bile excretes substances like sodium stereate that forms an emulsion with fats in our small intestine.
Emulsifying agents help form an emulsion.
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67. Colloids Removal of Colloidal Particles
Colloid particles are too small to be separated by physical means (e.g. filtration).
Colloid particles are coagulated (enlarged) until they can be removed by filtration.
Methods of coagulation:
heating (colloid particles move and are attracted to each other when they collide);
adding an electrolyte (neutralize the surface charges on the colloid particles).
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68. Colloids Removal of Colloidal Particles
Dialysis: using a semipermeable membranes separate ions from colloidal particles.
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69. End of Chapter 13 Properties of Solutions
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