Chapter 12 Solutions

1. Chapter 12 Solutions • Soluble – capable of being dissolved • Homogeneous - solution • Heterogeneous • Suspensions – particles can settle, large particles • Colloid – doesn’t settle, cloudy, particles are smaller than those in suspensions. Particle size 1-1000 nm.

2. Tyndall Effect • Light is scattered by colloidal particles dispersed in a transparent medium. • Distinguishes between a solution and a colloid.

3. Solutes • Electrolyte – substance that dissolves in water to give a solution that conducts electric current. Breaks up into ions. • Ex. Na+ Cl-

5. Solutes • Nonelectrolytes – substance that dissolves in water to give a solution that does not conduct an electric current.

6. The Solution Process • What affects the rate of dissolution? • Surface area • Agitation • Heat

7. Solubility • Solution equilibrium – physical state in which the opposing processes of dissolution and crystallization of a solute occur at equal rates. • Saturated – solution contain max amt. of dissolved solute. • Unsaturated – less than max amt. • Supersaturated – contains more dissolved solute than a saturated solution contains under the same conditions.

9. Solute-Solvent Interactions • “like dissolves like” • What makes substances similar depends on the type of bonding, the polarity or non-polarity of molecules, and the intermolecular forces between the solute and solvent.

10. Dissolving Ionic Compounds in Aqueous Solution • Water is polar and attracts to the ions in the solution. • Hydration – solution process with water as the solvent. • Ions are said to be hydrated CuSO4 : 5H2O hydrates(retain specific ratio of water molecules)

11. Nonpolar solvents • Nonpolar solvent molecules do not attract the ions of the solute enough to overcome the forces holding the crystal together. • Immiscible – liquids that are not soluble in each other. • Miscible – liquids that dissolve freely in one another in any proportion/

12. Effects of Pressure on Solubility • Gas above a liquid. • Equilibrium – go back and forth between liquid and gas. • Increase in pressure causes more gas molecules to collide with the liquid surface. • Increase in gas pressure causes the equilibrium to shift so that more molecules are in the liquid phase. • Ex. Pop bottle

13. Henry’s Law • Solubility of a gas in a liquid is directly proportional to the partial pressure of that gas on the surface of the liquid. • Pressure ∞ Solubility • Effervescence – rapid escape of a gas from a liquid in which it is dissolved.

14. Effects of Temperature on Solubility • Increase in temp increases solubility but not always at the same rate between solutes.

15. Enthalpies of Solution • The formation of a solution is accompanied by an energy change. • Absorb energy • Release energy • Energy is required to separate solute molecules and solvent molecules from their neighbors. • Solvated – solute particle that is surrounded by solvent molecules.

16. Enthalpy of solution – net amount of energy absorbed as heat by the solution when a specific amount of solute dissolves in a solvent.

17. Molarity Molarity – number of moles per liter of solution.

22. Molality

27. https://www.youtube.com/watch?v=9h2f1Bjr0p4

28. Dilutions • You may need to dilute solutions and change the molarity. • When you do this you will not change the amount of moles, just the molarity and volume, so… since M = moles/L moles = ML M1L1=M2L2

29. Example • You have a 56mL solution that is 3M HCl. You want to dilute this solution to a 2M. How much water should you add to the solution? M1L1=M2L2

30. You have in the stock room 12M HCl but you only need 300mL of 4M HCl. How much water should you add and how much 12M HCl should you mix? M1L1=M2L2

31. You want to make a 4M solution of NaOH, the lab calls for 400mL. How many grams of NaOH should you dissolve?