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CHEMISTRY The Central Science 9th Edition. Chapter 13 Properties of Solutions. Text, P. 417, review (Chapter 11). 13.1: The Solution Process. Solutions homogeneous mixtures Solution formation is affected by strength and type of intermolecular forces

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13.1: The Solution Process

  • Solutions
    • homogeneous mixtures
  • Solution formation is affected by
    • strength and type of intermolecular forces
    • forces are between and among the solute and solvent particles

Hydration of solute

  • Attractive forces between solute & solvent particles are comparable in magnitude with those between the solute or solvent particles themselves
  • Note attraction of charges
  • What has to happen to:
    • Water’s H-bonds?
    • NaCl?
  • What intermolecular
  • force is at work in
  • solvation?

Text, P. 486


Energy Changes and Solution Formation

  • There are three energy steps in forming a solution:
  • the enthalpy change in the solution process is
  • Hsoln = H1 + H2 + H3
  • Hsoln can either be + or - depending on the intermolecular forces

Text, P. 487


Text, P. 488

MgSO4 Hot Pack

NH4NO3 Cold Pack


Breaking attractive intermolecular forces is always endothermic

  • Formingattractive intermolecular forces is always exothermic
  • To determine whether Hsoln is positive or negative, consider the strengths of all solute-solute and solute-solvent interactions:
          • H1 and H2 are both positive
    • H3 is always negative

Rule: Polar solvents dissolve polar solutes

  • Non-polar solvents dissolve non-polar solutes
  • (like dissolves like)
  • WHY?
    • If Hsoln is too endothermic a solution will not form
      • NaCl in gasoline: weak ion-dipole forces (gasoline is non-polar)
        • The ion-dipole forces do not compensate for the separation of ions

Solution Formation, Spontaneity, and Disorder

  • A spontaneous process occurs without outside intervention
  • When energy of the system decreases, the process is spontaneous
    • Some spontaneous processes do not involve the system moving to a lower energy state (e.g. an endothermic reaction)
  • If the process leads to a greater state of disorder, then the process is spontaneous
    • Entropy

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

Text, P. 489


Solution Formation and Chemical Reactions

  • Example:
  • Ni(s) + 2HCl(aq)  NiCl2(aq) + H2(g)
  • When all the water is removed from the NiCl2 solution, no Ni is found only NiCl2·6H2O(achemical reaction that results in the formation of a solution)
      • Water molecules fit into the crystal lattice in places not specifically occupied by a cation or an anion
      • Hydrates
      • Water of hydration
  • Think about it: What happens when NaCl is dissolved in water and then heated to dryness?

NaCl(s) + H2O (l)  Na+(aq) + Cl-(aq)

  • When the water is removed from the solution, NaCl is found
    • NaCl dissolution is a physical process

13.2: 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

13.3: Factors Affecting Solubility

  • 1. Solute-Solvent Interaction
  • “Like dissolves like”
  • Miscible liquids: mix in any proportions
  • Immiscible liquids: do not mix


  • Intermolecular forces are important:
    • Water and ethanol are miscible
      • broken hydrogen bonds in both pure liquids are
      • re-established in the mixture
  • The number of carbon atoms in a chain affects solubility: the more C atoms in the chain, the less soluble the substance is in water

Generalizations, continued:

  • The number of -OH groups within a molecule increases solubility in water
  • The more polar bonds in the molecule, the better it dissolves in a polar solvent (like dissolves like)
  • Network solids do not dissolve
    • the strong IMFs in the solid are not re-established in any solution

Read “Chemistry & Life”, P. 494

Fat soluble vitamin

Water soluble vitamin


2. Pressure Effects

  • Solubility of a gas in a liquid is a function of the pressure of the gas

High pressure means

    • More molecules of gas are close to the solvent
    • Greater solution/gas interactions
    • Greater solubility
  • If Sg is the solubility of a gas
  • k is a constant
  • Pg is the partial pressure of a gas
  • then Henry’s Law gives:
  • Carbonated Beverages!

3. Temperature Effects

  • As temperature increases
    • Solubility of solids generally increases
    • Solubility of gases decreases
      • Thermal pollution

Text, P. 497


13.4: Ways of Expressing Concentration

  • All methods involve quantifying amount of solute per amount of solvent (or solution)
    • Amounts or measures are masses, moles or liters
    • Qualitatively solutions are dilute or concentrated


  • 3.
  • Recall mass can be converted to moles using the molar mass


  • Converting between molarity (M) and molality (m) requires density
    • Molality doesn’t vary with temperature
      • Mass is constant
    • Molarity changes with temperature
      • Expansion/contraction of solution changes volume

13.5: Colligative Properties

  • Colligative properties depend on quantity of solute particles, not their identity
    • Electrolytes vs. nonelectrolytes
  • 0.15m NaCl  0.15m in Na+ & 0.15m in Cl-  0.30m in particles
  • 0.050m CaCl2  0.050m in Ca+2 & 0.1m in Cl-  0.15m in particles
  • 0.10m HCl  0.10m in H+ & 0.10m in Cl-  0.20m in particles
  • 0.050m HC2H3O2  between 0.050m & 0.10m in particles
  • 0.10m C12H22O11 0.10m in particles
  • Compare physical properties of the solution with those of the pure solvent

1. Lowering Vapor Pressure

  • Non-volatile solutes reduce the ability of the surface solvent molecules to escape the liquid
  • Vapor pressure is lowered
  • Raoult’s Law:
  • PA is the vapor pressure with solute
  • PA is the vapor pressure without solute
  • A is the mole fraction of solvent in solution A

Increase X of solute, decrease vapor pressure above the solution


Ideal solution: one that obeys Raoult’s law

  • Raoult’s law breaks down (Real solutions)
    • Real solutions approximate ideal behavior when
      • solute concentration is low
      • solute and solvent have similar IMFs
    • Assume ideal solutions for problem solving
  • 2. Boiling-Point Elevation
  • The triple point - critical point curve is lowered

At 1 atm (normal BP of pure liquid) there is a lower vapor pressure of the solution

    • A higher temperature is required to reach a vapor pressure of 1 atm for the solution (Tb)
  • Molal boiling-point-elevation constant, Kb, expresses how much Tb changes with molality, m:

3. Freezing Point Depression

  • The solution freezes at a lower temperature (Tf) than the pure solvent
    • lower vapor pressure for the solution
  • Decrease in FP (Tf) is directly proportional to molality (Kfis the molal freezing-point-depression constant):

Text, P. 505

Applications: Antifreeze!


4. 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
    • “Where ions go, water will flow” ~ Mrs. Moss

Osmotic pressure, , is the pressure required to stop osmosis:

  • It is colligative because it depends on the concentration of the solute in the solvent

Isotonic solutions: two solutions with the same  separated by a semipermeable membrane

  • Hypertonic solution: a solution that is more concentrated than a comparable solution
  • Hypotonic solution: a solution of lower  than a hypertonic solution
  • Osmosis is spontaneous
  • Read text, P. 508 – 509 for practical examples
There are differences between expected and observed changes due to colligative properties of strong electrolytes
    • Electrostatic attractions between ions
    • “ion pair” formation temporarily reduces the number of particles in solution
    • van’t Hoff factor (i): measure of the extent of ion dissociation
Ratio of the actual value of a colligative property to the calculated value (assuming it to be a nonelectrolyte)
    • Ideal value for a salt is the # of ions per formula unit
  • Factors that affect i:
  • Dilution
  • Magnitude of charge on ions
    • lower charges, less deviation

11.6: Colloids

  • Read Text, Section 13.6, P. 511 – 515
    • Terms/Processes:
      • Tyndall effect
      • Hydrophilic
      • Hydrophobic
      • Adsorption
      • Coagulation

11.6: Colloids

  • Read Text, Section 13.6, P. 511 – 515
  • Suspensions in which the suspended particles are larger than molecules
    • too small to drop out of the suspension due to gravity
  • Tyndall effect: ability of a colloid to scatter light
    • The beam of light can be seen through the colloid

Hydrophilic and Hydrophobic Colloids

  • “Water loving” colloids: hydrophilic
  • “Water hating” colloids: hydrophobic
    • Molecules arrange themselves so that hydrophobic portions are oriented towards each other

Adsorption: when something sticks to a surface we say that it is adsorbed

    • Ions stick to a colloid (colloids appears hydrophilic)
  • Oil drop and soap (sodium stearate)
    • Sodium stearate has a long hydrophobic tail (Carbons) and a small hydrophilic head (-CO2-Na+)

Removal of Colloidal Particles

  • Coagulation (enlarged) until they can be removed by filtration
  • Methods of coagulation:
    • heating (colloid particles are attracted to each other when they collide)
    • adding an electrolyte (neutralize the surface charges on the colloid particles)