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I. The Nature of Solutions

Solutions. I. The Nature of Solutions. A. Definitions. Solution - homogeneous mixture Soluble – the ability to be dissolved. Solute - substance being dissolved. Solvent - present in greater amount. B. Solvation. Solvation – the process of dissolving.

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I. The Nature of Solutions

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  1. Solutions I. The Nature of Solutions

  2. A. Definitions • Solution - homogeneous mixture • Soluble – the ability to be dissolved Solute - substance being dissolved Solvent - present in greater amount

  3. B. Solvation • Solvation – the process of dissolving solute particles are surrounded by solvent particles First... solute particles are separated and pulled into solution Then...

  4. Solutions may exist as….gases, liquids, or solids.

  5. Label the solute & solvent in each of the following: (a) 14-karat gold - solute: Ag, solvent: Au (b) Water vapor in air - solute: H2O, solvent: air mixture (c) Carbonated water - solute: CO2 , solvent: water (d) Hot tea - solute: tea bag, solvent: water

  6. In addition to solutions, there are also suspensions and colloids.

  7. In summary: • Asolutionis always transparent, light passes through with no scattering from solute particles which are molecule in size. The solution is homogeneous and does not settle out. A solution cannot be filtered but can be separated using the process of distillation.

  8. A suspension is cloudy and heterogeneous. The particles are larger than 10,000 Angstroms which allows them to be filtered. If a suspension is allowed to stand the particles will separate out.

  9. A colloid is intermediate between a solution and a suspension. While a suspension will separate out a colloid will not. Colloids can be distinguished from solutions using the Tyndall effect. Light passing through a colloidal dispersion, such as smoky or foggy air, will be reflected by the larger particles and the light beam will be visible.

  10. The Solution Process There are 3 factors that can affect the rate of dissolution:  • Increasing the surface area of the solute •  Agitating a solution • Heating a solvent

  11. Solution Equilibrium is the physical state in which the dissolution and crystallization of a solute occur at equal rates. • This point is difficult to predict precisely because it depends on: • Nature of solute • Nature of solvent • The temperature

  12. C. Solubility • Solubility • maximum grams of solute that will dissolve in 100 g of solvent at a given temperature • varies with temp • based on a saturated solution

  13. UNSATURATED SOLUTION more solute dissolves SATURATED SOLUTION no more solute dissolves SUPERSATURATED SOLUTION becomes unstable, crystals form C. Solubility concentration

  14. C. Solubility • Solubility Curve • shows the dependence of solubility on temperature

  15. NONPOLAR NONPOLAR POLAR POLAR B. Solvation “Like Dissolves Like”

  16. B. Solvation • Soap/Detergent • polar “head” with long nonpolar “tail” • dissolves nonpolar grease in polar water

  17. Why are ionic compounds generally not soluble in nonpolar solvents? • The nonpolar solvent molecules do NOT attract the ions of the crystal strongly enough. • The strong ionic bond CANNOT be broken!

  18. Immiscible vs. Miscible LIquids • Immiscible – liquid solutes & solvents that are NOT soluble in each other (oil & water) • Miscible – liquids that dissolve freely in one another

  19. Effect of pressure & temperature on solubility • Changes in pressure have very little effect in liquid & solid solubility BUT – increases in pressure increase gas solubilities of liquids !!! WHY? - Remember gases are made up of mostly empty space, meaning that they can be compressed - The solute, then, can literally be “compressed” into the solvent.

  20. HENRY’S LAW • The solubility of a gas in a liquid is directly proportional to the pressure of that gas on the surface of the liquid • Examples: • The “bends” • Soda – EFFERVESCENCE

  21. C. Solubility • Solids are more soluble at... • high temperatures. • Gases are more soluble at... • low temperatures & • high pressures (Henry’s Law).

  22. Formations of solutions are accompanied by energy changes. • The formation of a solid-liquid solution can either absorb heat (ENDOTHERMIC) or release heat (EXOTHERMIC). • The heat of solution is the amount of heat energy absorbed or released when a specific amount of solute dissolves in a solvent.

  23. ENDO VS EXOTHERMIC RXN VIDEOS • https://www.youtube.com/watch?v=GQkJI-Nq3Os • https://www.youtube.com/watch?v=Yex063_Fblk

  24. Solutions II. Concentration

  25. Concentration • The amount of solute in a solution. • Describing Concentration • % by mass - medicated creams • % by volume - rubbing alcohol • ppm, ppb - water contaminants • molarity - used by chemists • molality - used by chemists

  26. Percent Concentration • Percent by mass and volume is a ratio of the solute amount to the solution amount, expressed as a percent.

  27. mass of solvent only 1 kg water = 1 L water Molality

  28. substance being dissolved total combined volume Molarity • Concentration of a solution.

  29. Molarity 2M HCl What does this mean?

  30. 500 mL of 1.54M NaCl 500 mLwater 500 mL volumetric flask 500 mL mark 45.0 gNaCl Preparing Solutions • 1.54m NaCl in 0.500 kg of water • mass 45.0 g of NaCl • add water until total volume is 500 mL • mass 45.0 g of NaCl • add 0.500 kg of water

  31. Dilution • Preparation of a desired solution by adding water to a concentrate. • Moles of solute remain the same.

  32. Solutions III. Colligative Properties

  33. A. Definition • Colligative Property • property that depends on the concentration of solute particles, not their identity

  34. Why is vapor pressure affected? • Vapor pressure is the tendency for molecules to escape from a liquid. • If a NON-VOLATILE solute is added, it will lower the tendency of solution to change into a gas.

  35. B. Types • Freezing Point Depression (tf) • f.p. of a solution is lower than f.p. of the pure solvent • Boiling Point Elevation (tb) • b.p. of a solution is higher than b.p. of the pure solvent

  36. B. Types Freezing Point Depression

  37. B. Types Boiling Point Elevation Solute particles weaken IMF in the solvent.

  38. B. Types • Applications • salting icy roads • antifreeze • cars (-64°C to 136°C) • fish & insects

  39. Solutes are classified according to if they yield IONS or MOLECULES in solution. • Electrolyte: a substance that dissolves in water to give a solution that conducts electrical current • Formed from any soluble IONIC compound • Nonelectrolyte: a substance that creates a solution that does NOT conduct electricity

  40. - + - - + + acetic acid salt sugar B. Solvation Non- Electrolyte Weak Electrolyte Strong Electrolyte solute exists as ions and molecules solute exists as ions only solute exists as molecules only DISSOCIATION IONIZATION

  41. DISSOCIATION – the separation of ions that occurs when an ionic compound dissolves

  42. C. Calculations t: change in temperature (°C) k: constant based on the solvent (°C·kg/mol) m: molality (m) n: # of particles t = k · m · n

  43. C. Calculations • # of Particles • Nonelectrolytes (covalent) • remain intact when dissolved • 1 particle • Electrolytes (ionic) • dissociate into ions when dissolved • 2 or more particles

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