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    Lecture17222 Lecture17222 Presentation Transcript

    • The Properties of Mixtures: the Solution Process Lecture 17
    • Similia similibus solvuntur
    • Macroscopic rule “like dissolves like” is based on microscopic interactions. How do enthalpy and entropy change in solute-solvent interaction?
    • Three events in the process of solution:
      • Solute particles separate from each other (some energy must be absorbed);
      • Some solvent particles separate to make room for the solute particles;
      • Solute and solvent particles mix together (some energy must be released).
      • There must be change in enthalpy!
    • Solution: separating particles
    • Solute particles separate from each other:
      • Solute (aggregated) + heat  solute (separated)
      • To overcome intermolecular attractions, energy is needed.
      • So the process is endothermic.
      • ∆ H solute > 0
    • Solvent particles separate from each other:
      • Solvent (aggregated) + heat  solvent (separated)
      • To overcome intermolecular attractions, energy is needed.
      • So the process is endothermic.
      • ∆ H solvent > 0
    • Solute and solvent particles mix:
      • Solute (separated) + solvent (separated)  solution + heat
      • The particles attract each other, energy is released.
      • So the process is exothermic.
      • ∆ H mix < 0
    • The three events in solution
    • Heat of solution (∆H soln ) is the total enthalpy change that occurs when a solution forms from solute and solvent. May be both exothermic and endothermic.
    • Thermochemical solution cycle: ∆ H soln = ∆H solute + ∆H solvent + ∆H mix Resembles Hess’s law and Born-Haber cycle.
    • Enthalpy components of the heat of solution
    • Solution implies solvation. Solvation is a process of surrounding a solute particle with solvent particles. Hydration is a process of surrounding a solute particle with water molecules.
    • Heat of hydration: ∆ H soln = ∆H solute + (∆H solvent + ∆H mix ) ∆ H hydr = ∆H solvent + ∆H mix ∆ H soln = ∆H solute + ∆H hydr
    • Heat of hydration
      • NaCl (g)  Na + (g) + Cl - (g)
      • Na + (g) + 6H 2 O (l)  [Na(H 2 O) 6 ] + (aq)
      • Cl - (g) + 6H 2 O (l)  [Cl(H 2 O) 6 ] - (aq)
      • -------------------------------------------
      • NaCl (s) + 6H 2 O (l)  [Na(H 2 O) 6 ] + (aq) +[Cl(H 2 O) 6 ] - (aq)
      • M + (g) [or X - (g) ] + H 2 O  M + (aq) [or X - (aq) ]
      • ∆ H hydr of the ion < 0, always
    • Charge density of an ion is the ratio of the ion’s charge to its volume. In general, the higher the charge density is, the more negative  H hydr is.
    • Coulomb’s law
      • A 2+ ion attracts H 2 O molecules more strongly than a 1+ ion of similar size;
      • A small 1+ ion attracts H 2 O molecules more strongly than a large 1+ ion.
    • Charge densities and heats of hydration
      • decrease down a group of ions (Li + —Na + —K + —Rb + —Cs + —Fr + ) - 1A
      • (F - —Cl - —Br - —I - ) - 7A group
      • increase across a period of ions (Na + —Mg 2+ —Al 3+ ) - 3rd period
    • The heat of solution for ionic compounds in water:  H soln =  H lattice +  H hydration of the ions  H lattice is always positive  H hydration is always negative
    • Dissolving ionic compounds in water
    • Hot (CaCl 2 ) and cold (NH 4 NO 3 ) packs
    • The heat of solution  H soln is only one of two factors determining whether a solute dissolves in a solvent. The other factor is entropy S.
    •  
    •  
    • Entropy is directly related to the number of ways that a system can distribute its energy. It is closely related to the freedom of motion of the particles and the number of ways they can be arranged.
      • Ludwig Eduard Boltzmann (1844–1906), Austrian scientist
    • Freedom of particle motion and entropy
      • S liquid > S solid ; ∆S melting > 0
      • S gas > S liquid ; ∆S vaporization > 0
      • S solid > S gas ; ∆S sublimation > 0
    • Solid state: minimum entropy
    • A solution usually has higher entropy than the pure solute and pure solvent: S soln > (S solute + S solvent ) ∆ S soln > 0
    • Systems tend toward a state of lower enthalpy and higher entropy.
    • Entropy is higher when mixed
    • THE END