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Chapter 12: Solutions

Chapter 12: Solutions. Solutions are homogeneous mixtures consisting of a solute and solvent. Not all solutions are liquids! A solution can be a solid, liquid, or a gas. . Types of Solutions. Solution Process.

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Chapter 12: Solutions

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  1. Chapter 12: Solutions • Solutions are homogeneous mixtures consisting of a solute and solvent. • Not all solutions are liquids! • A solution can be a solid, liquid, or a gas.

  2. Types of Solutions

  3. Solution Process • The intermolecular forces between solute and solvent particles must be strong enough to compete with those between solute particles and those between solvent particles.

  4. Solution Process • As an ionic compound dissolves in water, the ions are surrounded by six water molecules.

  5. Solution Process

  6. Solution Process • Energy of solution formation for ionic compounds has two factors. • Factor One: Lattice energy of ionic compound. • Factor Two: Ion-dipole force formed between water and ions. • Which one requires energy? • Which one produces energy? • NH4NO3(s) and MgSO4(s)

  7. Ion-Dipole Force • Intermolecular force between water and ion is called the Ion-Dipole force. • Hydration energy = combined interactions of all six waters with the ion.

  8. Energy of Ionic Solutes DHsolution = DHhydration – DHlattice energy

  9. Insoluble Ionic Compounds • Compounds with very LARGE lattice energies are insoluble. • Compounds with smaller lattice energies are highly soluble. • Ex) KBr(s) vs. PbS(s) • Ex) Al2O3(s) vs. KNO3(s)

  10. Energy of Solution Energy to overcome solute-solute interaction (requires energy = endothermic) Energy to overcome solvent-solvent interaction (requires energy = endothermic) Energy of solute-solvent interaction (produces energy = exothermic)

  11. Energy of Solution • If 1 + 2 – 3 > 0, then net process is endothermic. • If 1 + 2 – 3 < 0, then net process is exothermic. • If 1 + 2 – 3 >> 0, then the two substances will NOT mix.

  12. Spontaneity • Enthalpy is not the sole factor in why things dissolve. • Natural tendency of the universe is to go from order to disorder. • A solution is going from order to disorder when mixing occurs. • The amount of disorder is called the entropy of the system.

  13. Spontaneity • Both CCl4 and C6H14 are non-polar molecules with similar boiling points. • Adding the two together will produce a solution increasing the entropy of the system.

  14. Non-polar and Polar • “Like dissolves like” • Polar substances dissolve in polar solvents. • Non-polar substances dissolve in non-polar solvents. • Non-polar substances do NOT dissolve in polar solvents!

  15. Intermolecular Forces

  16. Types of Solvents

  17. Molecular Compounds • Molecular compounds will dissolve in water IF they are polar or have hydrogen bonding. • CH3CH2OH, CH3CHO • Size, though matters! • CH3CH2CH2CH2CH2CH2OH • W14, #1 • Vitamins

  18. Solubility of Alcohols

  19. Saturated Solutions • When more solid solute is added to a liquid and the solid does NOT dissolve, the solution is said to be saturated. • Solubility limit = maximum amount solute that can be dissolved in a given quantity of solvent. • Ex) @25oC, KNO3 = 36g / 100mL. • Super-saturated solution.

  20. Sodium Acetate Solution

  21. Temperature Effects • The solubility limits of most solid solutes INCREASES with increasing temperature.

  22. Temperature Effects • The solubility of ALL gases decreases with an increase in temperature. • Can have an adverse reaction in ponds and lakes in the hot summer months.

  23. Concentrations • There are many methods for expressing concentration. • Percent • mass / mass • volume / volume • mass / volume

  24. Concentrations • ppm (106) or ppb (109) • Mole fraction = moles A / total moles • Molarity (M) – Chapter 4 • Molality (m)

  25. Colligative Properties • Why do we put “salt” on our roads in the winter? • Why do we add “antifreeze” to the radiator in our cars? • Can we reduce the vapor pressure of “volatile” liquids? • Yes!

  26. Colligative Properties • Any property that depends only on the quantity of solute particles and not their identity. • Vapor Pressure lowering • Freezing-pt depression • Boiling-pt elevation • Osmotic Pressure

  27. Addition of a nonvolatile solute reduces the rate of vaporization, decreasing the amount of vapor Vapor Pressure • A non-volatile solute (solid) can be added to a volatile solvent and LOWER its vapor pressure. • Pa = caPao • Mole fraction is for the solvent and Pao is the vapor pressure for the pure solvent (dependent on T).

  28. Vapor Pressure • Two volatile liquids, of similar intermolecular forces, will both contribute a partial pressure. • Raoult’s Law • Ptotal = caPao + cbPbo • Graphical interpretation • Fractional distillation

  29. Ideal vs. Non-ideal Behavior

  30. BP and FP • When a solute is added to a solvent: • the freezing-pt is lowered. • the boiling-pt is raised. • Formulas are similar: • DT = Kf cm • DT = Kb cm

  31. Osmosis • In biological systems, the cellular walls are made up of cellulose. • Cellulose allows water molecules to pass in and out of the cells. • Larger molecules or ions are generally “blocked” from entering. • Note: glucose is transported into the cells via a complex process.

  32. Osmosis • When two solutions of different concentrations are separated by a semi-permeable membrane, solvent molecules will flow from lower solute concentration to higher solute concentration.

  33. Osmosis • This is important in maintaining water in our cells.

  34. Osmotic Pressure • The osmotic pressure is measured in atmospheres. • p = MRT • M is the molarity of the solution • R = 0.08206 L atm/mol K • T = temperature, K

  35. Molar Masses • Many experiments involving colligative properties can find a molar mass of a solute. • Generally, this requires us to work backwards through several formulas.

  36. Ionic Solutes • Because ionic solutes break apart into ions, an adjustment to our formulas must be made. • i = van’t Hoff factor. • Equals the number of ions per formula unit. • True factor depends on the concentration. As the concentration increases, i becomes less than predicted. • Why?

  37. Ionic Solutes

  38. Ionic Solutes • Assume “ideal” values for each electrolyte. • Becomes multiplier in all formulas. • What is i for: • CaCl2 • Na2CO3 • Al(NO3)3 • NH3

  39. Colloids • When larger solute particles remain suspended in a solution, they form a colloid. • Colloids will exhibit the Tyndall Effect due to the larger particles.

  40. Colloids • Tyndall effect in nature.

  41. Colloids • There are several types of colloids. • Aerosol = a liquid or solid in a gas. • fog, smoke • Foam = a gas in a liquid or solid. • whipped cream, marshmallow • Emulsion = a liquid in a liquid or a solid in a solid. • milk, butter, mayo • Sol = a solid in a liquid or solid. • cement, paint, gems

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