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Chemistry LibreTexts

13.S: Properties of Solutions (Summary)

  • Page ID
    70553
  • These is a summary of key concepts of the chapter in the Textmap created for "Chemistry: The Central Science" by Brown et al. 

    13.1: The Solution Process

    • interaction between solute and solvent molecules
    • hydration – solvation when solvent is water

    13.1.1 Energy Changes and Solution Formation

    • overall enthalpy change in formation of a solution
      • \(\Delta H_{\mathit{soln}}=\Delta H_1+\Delta H_2+\Delta H_3\)
        • \(\Delta H_1=\) separation of solute molecules
        • \(\Delta H_1=\) separation of solvent molecules
        • \(\Delta H_3=\) formation of solute-solvent interactions 
    • separation of solute particles is endothermic
    • separation of solvent is endothermic
    • third is exothermic
    • formation of solution can be either exothermic or endothermic
    • exothermic processes are spontaneous
    • solution will not form if enthalpy is too endothermic
    • H3 has to be comparable to H1 + H2
      • Ionic substances cannot dissolve in nonpolar liquids
      • Polar liquids do not form solutions with nonpolar liquids

    13.1.2 Solution Formation, Spontaneity, and Disorder

    • two nonpolar substances dissolve in one another
    • attractive forces = London dispersion forces
    • two factors in processes that are spontaneous: energy and disorder
    • processes in which the energy content of the system decreases tend to occur spontaneously
      • exothermic
    • processes in which the disorder of the system increases tend to occur spontaneously
    • solutions will form unless solute-solute or solvent-solvent interactions too strong relative to solute-solvent interactions

    13.1.3 Solution Formation and Chemical Reactions

    • distinguish between physical process of solution formation from chemical process that leads to a solution

    13.2: Saturated Solutions and Solubility

    • crystallization – reverse process of solution
    • dynamic equilibrium – when equilibrium exists between process of solution and crystallization
    • solute said to be saturated
    • solubility – amount of solute needed to saturate a solution
    • unsaturated – when there isn’t enough solute to saturate a solution
    • supersaturated – when there is more solute than needed to saturate a solution
    • for most salts crystallization of excess solute is exothermic

    13.3: Factors Affecting Solubility

    Solute-Solvent Interactions

    • solubility increases with increasing molar mass
    • London dispersion forces increase with increasing size and mass of gas molecules
    • Miscible – pairs of liquids that mix in all proportions
    • Immiscible – opposite of miscible
    • Hydrogen-bonding interactions between solute and solvent leads to high solubility
    • Substances with similar intermolecular attractive forces tend to be soluble in one another
    • "like dissolves like"

    Pressure Effects

    • solubility of a gas in any solvent increases as pressure of gas over solvent increases
    • relationship between pressure and solubility: Henry’s Law:
      • \(C_g = kP_g\)
      • Cg solubility of gas in solution phase (usually expressed as molarity), Pg partial pressure of gas over solution, k is proportionality constant (Henry’s Law constant)
      • Henry’s law constant different fore each solute-solvent pair, and temperature

    Temperature Effects

    • solubility of most solid solutes in water increases as temperature of solution increase
    • solubility of gases in water decreases with increasing temperature
    • decreases solubility of O2 in water as temperature increases in one the effects of thermal pollution

    13.4: Ways of Expressing Concentration 

    • dilute and concentrated used to describe solution qualitatively
    • mass percentage of component in solution:

    \[\displaystyle\textit{mass% of component} = \frac{\textit{mass of component in soln}}{\textit{total mass of soln}}\times 100\]

    • very dilute solutions expressed in parts per million (ppm)

    \[\displaystyle\textit{ppm of component} = \frac{\textit{mass of component in soln}}{\textit{total mass of soln}}\times 10^6\]

    • 1ppm = 1g solute for each (106) grams of solution or 1mg solute per kg solution
    • 1ppm = 1mg solute/L solution
    • 1 ppb = 1g of solute/109 grams of solution or 1 m g solute/ L of solution

    13.4.1 Mole Fraction, Molarity, and Molality

    \[\displaystyle\textit{mole fraction of component} = \frac{\textit{moles of component}}{\textit{total moles of all components}}\]

    • sum of mole fractions of all components of solution must equal one
    • \(\displaystyle\textit{molarity} = \frac{\textit{moles solute}}{\textit{liters soln}}\)
    • \(\displaystyle\textit{molality} = \frac{\textit{moles solute}}{\textit{kilograms of solvent}}\)
    • molality goes not vary with temperature
    • molarity changes with temperature because of expansion and contraction of solution

    13.5: Colligative Properties

    Colligative properties are physical properties that depend on quantity

    Lowering the Vapor Pressure

    • vapor pressure over pure solvent higher than that over solution
    • vapor pressure needed to obtain equilibrium of pure solvent higher than that of solution

    Raoult’s Law

    • Raoult’s law: \(P_A = X_AP_{A}^°\)
      • PA = vapor pressure of solution, XA = mole fraction of solvent, PA° = vapor pressure of the pure solvent
      • Ideal solution – solution that obeys Raoult’s law
      • Solute concentration is low, solute and solvent have similar molecular sizes and similar types of intermolecular attractions

    Boiling-Point Elevation

    • normal boiling point of pure liquid is the temperature at which pressure is 1 atm
    • addition of a nonvolatile solute lowers vapor pressure of solution
    • \(\Delta T_b=K_b m\)
    • Kb = molal boiling-point-elevation constant
      • Depends only on solvent
      • boiling point elevation proportional to number of solute particles present in given quantity of solution

    Freezing-Point Depression

    • freezing point of solution is temperature at which the first crystals of pure solvent form in equilibrium
    • freezing point of solution lower than pure liquid
    • freezing point directly proportional to the molality of the solute:
      • \(\Delta T_f=K_f m\)
      • Kf = molal freezing-point-depression constant

    Osmosis

    • semipermeable – membranes that allow passage of some molecules and not others
    • osmosis – the net movement of solvent molecules from the less concentrated solution to the more concentrated solution
    • net movement of solvent always toward the solution with the higher solute concentration
    • osmotic pressure – pressure needed to prevent osmosis, p
      • \(\pi = \left(\frac{n}{V}\right)RT=MRT\)
      • M = molarity of solution
    • if solutions identical osmosis will not occur and said to be isotonic
    • if one solution lower osmotic pressure = hypotonic, the solution that has higher osmotic pressure = hypertonic
    • crenation = when cells shrivel up from the loss of water
    • hemolysis = when cells rupture due to to much water

    Determination of Molar Mass

    • colligative properties can be used to find molar mass

    13.6: Colloids

    Colloidal dispersions (colloids) are intermediate types of dispersions or suspensions

    • intermediate solutions between solutions and heterogeneous mixtures
    • colloids can be gases, liquids, or solids
    • colloid particles have size between 10 - 2000Å
    • tyndall effect – scattering of light by colloids

    Hydrophilic and Hydrophobic Colloids

    • hydrophilic – colloids in which the dispersion medium is water
    • hydrophobic – colloids not dispersed in water
    • hydrophobic colloids have to be stabilized before being put in water
      • natural lack of affinity for water causes separation
      • can be stabilized by the adsorption of ions on the surface
      • adsorped ions interact with water
      • can also be stabilized by presence of other hydrophilic groups on surface

    Removal of Colloidal Particles

    • coagulation – enlarging colloidal particles
      • heating or adding an electrolyte to mixture
      • heating increases number of collisions and particles stick together increasing their size
      • electrolytes causes neutralization of the surface charges of the particles which remove the electrostatic repulsion
      • dialysis – use of semipermeable membranes to filter out colloidal particles