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Solutions

Solutions. Chapter 12. Types of Solutions. Solution  homogeneous mixture of two or more substances in any phase. Solute  component present in the smaller amount Is dissolved in the solvent Solvent  component present in the greater amount what is the solute and solvent? Salt water

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Solutions

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  1. Solutions Chapter 12

  2. Types of Solutions • Solution homogeneous mixture of two or more substances in any phase. • Solute  component present in the smaller amount • Is dissolved in the solvent • Solvent  component present in the greater amount • what is the solute and solvent? • Salt water • 80g of Cr and 5g of Mo • Immiscible  fluids (gas/liquid) that don’t mix. They form two layers • as in oil and water. 

  3. Solubility and the Solution Process Solubility  amount of substance that will dissolve in a solvent at a specified temp. Saturated solution  no more solute can dissolve (is in equilibrium) Unsaturated solution  the solvent can dissolve more solute (not in equilibrium) Supersaturated solution  contains more dissolved solute than a saturated solution (not in equilibrium) 

  4. Factors in Explaining Solubility • “like dissolves like”  similar substances are miscible with each other • Oil will mix with gasoline • Both are hydrocarbons • Oil won’t mix with water • H2O is polar and oil is not • Solubility of a solute in a solvent depends on a balance between • Natural tendency for the species to mix and • The tendency for a system to have the lowest energy possible 

  5. Factors in Explaining Solubility • Example: • Decide whether liquid hexane (C6H14) or liquid methanol (CH3OH) is the more appropriate solvent for the substances grease (C20H42) and potassium iodide (KI). 

  6. Factors in Explaining Solubility • Molecular solutions: “like dissolves like” • Gases are miscible • Hydrocarbons are miscible with hydrocarbons • Polar are miscible with polar • Ionic solutions: • Hydration  the attraction of ions for water molecules • Ionic substances differ in their solubilities in water • Smaller ions have a more concentrated electric field and are more attracted to water increasing their hydration energy • The size of the charge also effects the hydration energy • Ex: which ion has the larger hydration energy, • Na+ or K+ • Na+ or Mg2+ 

  7. Effects of Temperature & Pressure on Solubility • Temperature  solubility often increases with temperature • some ionic compounds may go down, gas solubility goes down (warm pop) • The dissolving process may be exothermic or endothermic. • Pressure  only affects the solubility of a gas • When the partial pressure of the gas over the liquid is increased, the more soluble the gas. CO2 (g) CO2 (aq) • When pressure is released, gas is less soluble, and bubbles out. 

  8. Homework • Pg 516: • 1, 2, 4, 6, 8, 11, 14, 19, 20, 34, 36, 41,44, 46, 49, 52, 53, 55, 63, 65, 67, 73, 83, 91, 96

  9. Effects of Temperature & Pressure on Solubility • Le Chatelier’s principle  when a system in equilibrium is disturbed by a change of temperature , pressure or concentration variable, the system shifts to counteract the change in variable • Henry’s Law  the solubility of a gas is directly proportional to the partial pressure of the gas above the solution • S2 = P2 S1 P1 

  10. Example 27g of acetylene, C2H2, dissolves in 1L of acetone at 1.0 atm pressure. If the partial pressure of acetylene is increases to 12 atm, what is the solubility in acetone? 

  11. Colligative Properties • Colligative properties  depend on the concentration of solute ions or molecules, not the identity of the solute • Concentration  amount of solute dissolved in a solution. • Can be expressed by volume, mass, or moles. 

  12. Ways of Expressing Concentration Molarity (M) = moles of solute liters of solution Mass percentage = mass of solute mass of solution x 100 Molality (m) =moles of solute kg of solvent Mole fraction = moles of A total moles of solution

  13. Examples An experiment calls for 36.0 g of a 5.00 % aqueous solution of potassium bromide. Describe how you would make up such a solution.

  14. Examples What is the molality of I2 in a solution of 5.00 g of I2, in 30.0 g of methylene dichloride, CH2Cl2?

  15. Examples What is the mole fraction of a 3.6 m solution of calcium chloride?

  16. Examples An aqueous solution of citric acid, HC6H7O7 is 2.331 m. What is the molarity of citric acid in this solution? The density of the solution is 1.1346g/mL.

  17. Homework • Pg 516: • 1, 2, 4, 6, 8, 11, 14, 19, 20, 34, 36, 41, 44, 46, 49, 52, 53, 55, 63, 65, 67, 73, 83, 91, 96

  18. Colligative properties • Vapor pressure lowering  when a nonvolatile solute is added to volatile solvent • P = (partial pressure of solvent)(mole fraction of solute) • Distillation  a typical lab technique to separate a more volatile component from a liquid mixture.

  19. Colligative properties • Boiling point elevation: addition of a solute will increase the b.p. of a solvent (b/c of vapor pressure lowering). This effect is proportional to the concentration of solute. • Tb = (Kb)(molality), • Kb = boiling-point-elevation constant • the Kb for water is 0.512C/m

  20. Colligative properties • Freezing point depression  addition of a solute will decrease the f.p. of a solvent. • Tf = (Kf)(molality) • Kf= freezing-point-elevation constant • the Kf for water is 1.85 C/m

  21. Example • An aqueous solution is 0.0222 m glucose. What are the boiling and freezing point of this solution? • Kb = 0.512oC/m Kf = 1.86oC/m

  22. Example A solution of 58.1 mg anethole in 5.00 g of benzene is determined by f.p. depression to have a molality of 0.0784 m. What is the molecular weight of anethole?

  23. Colloids • Colloid  a dispersion of particles of one substance throughout another substance or solution • Ex: Fog… consists of small water droplets and air • Tyndall Effect  The scattering of light by colloidal-size particles • Ex: air appears to be a clear gas but sun rays show up dust and other small particles

  24. Types of Colloids • Colloids are characterized by the state of the dispersed particles • Aerosols  liquid droplets or solid particles dispersed throughout a gas • Fog or smoke • Emulsion  liquid droplets dispersed through another liquid • Butterfat dispersed though homogenized milk • Sol  solid particles dispersed in a liquid

  25. Types of Colloids • Hydrophilic colloid  a strong attraction between the dispersed phase and the continuous phase (water) • Hydrophobic  a lack of attraction between the dispersed phase and the continuous phase (water) • Are basically unstable and given enough time the dispersed phase aggregates into larger particles

  26. Coagulation Coagulation  the process by which the dispersed phase of a colloid is made to aggregate and separate from the continuous phase

  27. Homework • Pg 516: • 1, 2, 4, 6, 8, 11, 14, 19, 20, 34, 36, 41, 44, 46, 49, 52, 53, 55, 63, 65, 67, 73, 78, 83, 91, 96

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