Chapter 13 Notes- Properties of Solutions

13.1 The Solution Process[p.528]

  1. Background
  2. Solutions= homogeneous mixtures
  3. Can be mixtures of solid, liquids, or gases
  4. Examples include saltwater, sterling silver, and air
  5. Component- each substance of the mixtures
  6. Solvent- component present in the greatest amount
  7. Solutes – all lesser quantity components
  8. Aqueous Solutions- solutions in which water is the solvent
  9. 2 general factors determine the ability of substances to form solutions
  10. Intermolecular interactions
  11. Natural Tendency of substances to spread into larger volumes
  12. The Effects of Intermolecular Forces
  13. Must consider all forces involved between
  14. Solute-solute
  15. Solute- solvent
  16. Solvent –solvent
  17. Example is NaCL (aq)
  18. Solute dissolves readily in water due to the ion-dipole attraction
  19. Solute- solvent- water molecules surround ions (solvation-hydration when water is the solvent)
  20. Solvent- solvent hydration process the hydrogen bonding is weaken to make space between water molecules for the ions
  21. Energy Changes and Solution Formation (three aspects to consider)
  1. ΔH soln = ΔH1 +ΔH2+ΔH3
  2. Solute –solute interactions must broken
  3. Requires input of energy (endothermic)
  4. ΔH1> 0
  5. Solvent – solvent interactions must be separated too
  6. Endothermic
  7. ΔH2> 0
  8. Solute – solvent interaction
  9. Exothermic
  10. ΔH3> 0
  1. ΔH soln can be endothermic or exothermic
  2. Depends on the amount of energy absorbed in separating the solute-solute interactions and the solvent-solvent actions along with the quantity of heat released by the formation of the solute-solvent interaction
  3. A solution will not form if ΔH solnis too endothermic. The solute-solvent be great enough to counter the other two interactions
  4. General rule:

Polar substances dissolve in polar substances

Nonpolar in nonpolar

  1. Solution Formation, Spontaneity, and Entropy
  2. Solution Formation and Chemical Reactions
  3. Solutions that form spontaneously areusually processes in which the energy content if the system decreases (exothermic)
  4. Lower energy (enthalpy)
  5. Less energy required to maintain interactions
  6. Principle #2 : Processes occurring at a constant temperature which result in greater randomness or dispersal in space (entropy) are spontaneous
  7. In most cases the formation of solutions is favored by the increase in entropy that accompanies mixing
  8. Solution Formation and Chemical Reactions
  9. Note solutions allow for a recovering of the solute by physical means
  10. Chemical reactions my result in a change of phase but the original reactants cannot be recovered without a chemical change

13.2 Saturated Solutions and Solubility [p.534]

  1. Saturated solution- contains the maximum amount of solute in a solvent based on the temperature (solubility)
  2. Once max. amount is dissolved any extra solute will settle at the bottom of the container
  3. Dynamic equilibrium is established between the dissolving and crystallizing of the solute in the solution
  4. Unsaturated solution contains less than the max amount
  5. Supersaturated contains more than the max.
  6. Most prepare at high temperatures then cool slowly
  7. One crystal dropped on solution will cause the extra to crystallize

13.3Factors Affecting Solubility [p.535]

  1. Solute-Solvent Interactions
  2. Gases larger in size and mass tend to increase in solubility to the increase in the London dispersion forces.
  3. When other factors are comparable, the stronger the attraction between solute and solvent, the greater the solubility.
  4. “Likes dissolve in likes”
  5. Network solids have too great of intermolecular attractions to dissolve.
  6. Pressure Effects
  7. The solubility of the gas increases in direct proportion to its partial pressure above the solution
  8. Henry’s Law state the mathematical relationship

Sg = kPg

Sg= the solubility of the gas usually in molarity

k = Henry’s law constant (unique to each substance and temperature)

Pg = partial pressure of gas above the solution

  1. Temperature Effects
  2. As temperature of the water increases, solubility of a solid usually increases
  3. Solubility of a gas decreases with an increase in water temperature
  4. Thermal pollution results from high water temperatures lowering the solubility of oxygen

13.4Ways of Expressing Concentrations [p.542]

  1. Quantitative measurements indicate whether the solution is dilute or concentrated
  2. Mass Percentage, ppm, ppb
  3. Mass Percentage
  1. ppm
  1. ppb
  1. Mole Fraction, Molarity, and Molality
  2. Mole fraction (Xsub)
  3. = Moles of component

Total moles of all components

  1. No units
  2. Useful when dealing with gases
  1. Molarity (M)
  2. = Moles solute

Liters soln

  1. Useful for relating the volume of the solution to the amount of solute
  1. Molality (m)
  2. = moles of solute

Kilograms of solvent

  1. Unit = moles/ kg or m
  2. Molality does not change with temperature
  1. Conversion of Concentration Units
  2. May be necessary to move between types of concentrations
  3. Density is frequently necessary in such conversions

13.5Colligative Properties [p.546]

  1. Colligative Properties- properties that are dependent of the concentration of the quantity of the substance not the kind.
  2. Lowering the Vapor Pressure
  3. Past knowledge
  4. Vapor pressure = when equilibrium is reached, the pressure of the vapor above the liquid
  5. Involitale = substances without vapor pressure
  6. Volitale = substances with vapor pressure
  7. The more parts in solution the lower the vapor pressure.
  8. Follows Raoult’s Law

PA = XAP˚A

PA= partial pressure of a solvent’ vapor above the solution

XA= mole fraction of the solvent

P˚A = vapor pressure of pure solvent

  1. Boiling- Point Elevation
  2. Boiling point of a solution is higher than the BP of a pure liquid
  3. Equation use:

ΔTb= Kbm

ΔTb = increase in boiling point compared to pure solvent

Kb = molal boiling-elevation constant

m = Molality of the solution

  1. ΔTbis proportional to the Molality of the solution
  1. Freezing- Point Depression
  2. Freezing point of the solution is lower than that of the pure liquid
  3. Equation use:

ΔTf = Kfm

ΔTf = difference between the freezing point of eth solution and the freezing point of the pure solvent

Kf = molal freezing-point constant

m = Molality of the solution

  1. Osmosis
  2. The net movement of the solvent always goes toward the higher solute concentration. (tries to dilute a concentrated solution)
  3. Osmotic pressure = the pressure needed to prevent osmosis of a pure solvent

M= Molarity

Π = osmotic pressure

R = ideal gas constant

T = temperature

n = number of moles

V= volume

  1. Isotonic = no osmosis occurs (equal osmotic pressure)
  2. Hypotonic = solution with lower osmotic pressure
  3. Hypertonic = more concentrated solution of two
  4. Crenation is the shrinking of a cell by placing a cell in a hypertonic solution
  5. Hemolysis –the rupture of a cell due to placing it in a hypotonic solution
  6. IV solutions use to replenish nutrients must be isotonic to avoid cell damage
  7. Edema = retaining of water
  8. Active transport – transport from low concentration to high, not spontaneous, energy must be expended
  1. van’t Hoff, i is a factor unique to each solute and concentration
  2. represents the ratio of the ΔTf (measured) to the ΔTf ( calculated for nonelectrolyte)
  3. value approaches the ideal value as concentration increases
  1. Determination of Molar Mass- can be found by using any the colligative properties

13.6_Colloids [p.556]

  1. Background
  2. Colloidal dispersions or colloids – contain suspended particles that are larger than a solute but smaller then components of a mixture
  3. Represent the dividing line between solutions and heterogeneous mixtures
  4. Can be gases, liquids, or solids
  5. Tyndall Effect- scattering of light by colloidal particles
  6. Hydrophilic and Hydrophobic Colloids
  7. Hydrophilic (water loving)- colloids that like aqueous medium keeps particles suspended (polar groups)
  8. Hydrophobic (water fearing) – must be folded inside molecule to stay suspended
  9. Emulsifiers – have a hydrophilic and hydrophobic end to keep the particle suspended
  10. Removal of Colloidal Particles
  11. Coagulation – enlarges the size of the colloidal particles so they can be filtered out
  12. Dialysis – uses a semipermeable membrane to remove the colloidal particles

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