CHEM 3.2

• Almost all substances dissolve in water under observable temperature and pressure. But some substances do not.
• Oil


  • Forms immiscible bubbles that do not mix with water
• Substances that mix easily

  Possess "compatible" intermolecular forces of attraction
• Explain the effect of temperature on the solubility of a solid and of a gas

  STEM_GC11PP-IIId-f-113
• Explain the effect of pressure on the solubility of a gas

  STEM_GC11PP-IIId-f-114
• Intermolecular Forces of Attraction (IMFAs)


  Particles of substance exhibit intermolecular forces of attraction and repulsion to one another
• London dispersion forces (LDFs)


  Exist in all molecules, polar or nonpolar. Strength depends on size, surface area, and polarizability
• Ion-ion interactions


  Between charged substances. Strength follows Coulomb's law
• Dipole-dipole interactions


  Between polar, uncharged substances. Strength depends on electronegativity of atoms in a bond
• Induced Dipole Interactions


  Ions and dipoles induce formation of temporary dipoles on nonpolar molecules
• Hydrogen Bonding


  Polar molecules that have H atoms bonded to O, N, or F (e.g., H2O, NH3, and HF)
• Solution Process


  1. Solute particles must be dissolved in solvent particles

  2. A change in IMFAs between the particles:

  3. Solvent-solvent interaction

  4. Solute-solute interaction

  5. Solute-solvent interaction
• Molar enthalpies


  Energy changes that accompany each of the interactions
• Solution Process


  1. Solvent molecules are separated

  2. Solute and solvent molecules form a solution
• Heat of Solution, ΔHmixing


  Also referred to as ΔHsoln. Sum of all the enthalpy changes associated with each step
• ΔHmixing = ΔH1 + ΔH2

  Is always positive
• Sign of ΔHmixing


  Depends on the magnitudes of the three enthalpy changes
• ΔHmixing > 0


  Is endothermic. ΔH1 + ΔH2 is greater than ΔH3. Solvent-solvent and solute-solute interactions are stronger than solute-solvent interactions
• ΔHmixing < 0


  Is exothermic. ΔH1 + ΔH2 is less than ΔH3. Solvent-solvent and solute-solute interactions are weaker than solute-solvent interactions
• Entropy (ΔS)


  The inherent tendency toward disorder in highly favorable processes. The chaos when solute and solvent particles mix in comparison to their initial ordered states
• Solubility


  The extent to which a solute dissolves in solvent at a particular temperature. Governed by both ΔHmixing and ΔSmixing
• Solute and solvent with similar IMFAs tend to mix favorably. Solute and solvent with different IMFAs do not usually mix.
• Like dissolves like


  Substances with similar intermolecular forces of attraction dissolve to one another
• Solvation


  The process in which solvent molecules surround an ion or a molecule
• Hydration


  The solvation process when water is the solvent. Ions are attracted to the dipole by ion-dipole interaction
• Dissolution of Ionic Solids in Liquids

  1. NaCl in water:

  2. NaCl: crystal lattice composed of Na+ and Cl-

  3. H2O: polar solvent
• Dissolution of Ionic Solids


  1. Step 1: Separation of ions in the NaCl crystal lattice. Na+ and Cl- ions. Endothermic process; ΔH1 > 0 (crystal lattice energy)

  2. Step 2: Separation of solvent molecules. H2O molecules (H-bonding). Endothermic process; ΔH2 > 0

  3. Step 3: Interaction of solute and solvent particles. Na+ interact with partially negative O atoms. Cl- interact with partially positive H atoms. IMFAs formed between ions and H2O molecules; ΔH3 < 0 (hydration energy)
• Crystal Lattice Energy (CLE)


  The energy released when a mole of formula units of a solid is formed from its constituent ions in the gas phase. Reflects the strength of IMFA present in solid. Higher CLE, strong IMFA holding the particles together. The energy needed to break the solid → ΔH1 = -CLE, always positive
• Solvation Shells


  Form when the solvent molecules surround ions (or atoms or molecules) of the solute in a specific arrangement
• Hydration Shells


  The solvation shell when water is the solvent
• Hydration Energy, ΔH3


  The energy released when one mole of formula units becomes hydrated. Increases with charge density of ion
• Charge densities and hydration energies of some ions
• Heat of Solution, ΔHmixing, of Ionic Solids in Liquids

  • ΔHmixing > 0 is endothermic, |–CLE| > |HEcation + HEanion|
  • ΔHmixing < 0 is exothermic, |–CLE| < |HEcation + HEanion|
• Dissolution of Liquids in Liquids

  Solvation process: highly dependent on IMFAs present in solute and solvent particles
• Benzene (C6H6) and toluene (C7H8)


  • Exhibit only LDFs; similar IMFAs. Miscible liquids
• Immiscible liquids


  Do not mix well and form distinct layers. Exhibit distinct IMFAs
• Ideal solution


  IMFAs of solute and solvent are 'compatible'. ΔHmixing = 0
• Dissolution of Gases in Liquids

  1. Polar gases (e.g., NO and NO2) are soluble in polar solvents (e.g., H2O)

  2. Nonpolar gases (e.g., H2S) are soluble in nonpolar solvent (e.g., hexane)

  3. Some polar gases enhance their solubility by reacting with water

  4. Few nonpolar gases are soluble in water because few can react with water

  5. Gases have very weak IMFAs, energy needed to break IMFAs often negligible

  6. Dissolution of gases in liquid is always exothermic
• The rule of thumb "like dissolves like" means that substances with similar intermolecular forces of attraction are soluble to one another.
• Hydration energy (HE) increases with charge density.