Chapter 12 SOLUTIONS

1. Chapter 12SOLUTIONS

2. Chapter 12 • Section 1 – Types of Mixtures • Section 2 – The Solution Process • Section 3 – Concentration of Solutions

3. 12.1 - Types of Mixtures • Distinguish between electrolytes and nonelectrolytes. • List three different solute-solvent combinations. • Compare the properties of suspensions, colloids, and solutions. • Distinguish between electrolytes and nonelectrolytes.

4. Introduction Sugar & Water is a mixture Soil is a mixture • 2 or more different substances • Homogeneous Mixture (uniform properties) • Heterogeneous Mixture (nonuniform properties)

5. Solutions • Solution: homogeneous mixture of two or more substances • Solute: substance dissolved in a solution • Solvent: substance doing the dissolving (usually in a large quantity) • Can be any combination of solid, liquid, or gas • A homogeneous mixture of solids is called an alloy • Ex: brass – “solution” of zinc in copper

6. Examples of Solutions In these examples, the substance with the greater volume is the solvent

7. Particle Models for Gold and a Gold Alloy

8. Solutions • Soluble: capable of being dissolved • Insoluble: not capable of being dissolved What substances are soluble? • Depends on your solute & your solvent (and other factors, too) • In general, follow the rule: “like dissolves like”

9. Polarity & Solubility • Polar substances dissolve in polar substances • Ionic compounds • Polar molecules • Nonpolar substances dissolve in nonpolar substances • Nonpolar molecules • Organic compounds

10. Like Dissolves Like

11. Heterogeneous vs. Homogeneous

12. Suspensions • Suspension: particles in a solvent are so large they settle out unless the mixture is constantly stirred • What is an example of a suspension? • Oil and water • Sand and water

13. Colloids • Colloid: particles are small enough to be suspended throughout the solvent by constant movement of the surrounding molecules (particles that are intermediate in size) • Colloidal particles make up the dispersed phase, and water is the dispersing medium. • What are some examples of Colloids? • Paint (solid in liquid) • Whipped cream (gas in liquid) • fog

14. Emulsions – A Specific Type of Colloid a mixture of two or more immiscible (un-blendable) liquids

15. Tyndall Effect • Tydnall Effect: when light is scattered by colloidal particles dispersed in a transparent medium (like water) • This is used to distinguish between a solution and a colloid

16. Solutions, Colloids, and Suspensions

17. Electrolytes • Electrolyte: a substance that dissolves in water to give a solution that conducts electric current • What would make a good electrolyte? • Ionic compounds – the positive and negative ions separate from each other in solution and are free to move making it possible for an electric current to pass through the solution • Polar covalent molecules

18. Nonelectrolyte • Nonelectrolyte: a substance that dissolves in water to give a solution that does not conduct electricity • Neutral solute molecules do not contain mobile charged particles so they do not conduct electric current • Example: sugars, alcohols, organic hydrocarbons

19. Electrical Conductivity of Solutions

20. Strong and Weak Electrolytes • What is an electrolyte? • The strength of an electrolyte depends on the ability of a substance to form ions, or the degree of ionization or dissociation.

21. Strong Electrolytes • Strong Electrolyte: conducts electricity well due to presence of all or almost all of the dissolved compound as ions • What is an example of a strong electrolyte? • Strong acids (HCl, HNO3) • Strong bases (NaOH) • ionic compounds (assuming their soluble)

22. Weak Electrolytes • Weak Electrolyte: conducts electricity poorly due to presence of a small amount of the dissolved compound as ions • What is an example of a weak electrolyte? • Weak acids and bases • Ex: Acetetic Acid, H3PO4 & NH3 (weak base) • [HF] >> [H+] and [F–]

23. Models for Strong and Weak Electrolytes and Nonelectrolytes

24. 12.2 - The Solution Process Listand explain three factors that affect the rate at which a solid solute dissolves in a liquid solvent. Explainsolution equilibrium, and distinguish among saturated, unsaturated, and supersaturated solutions. Explainthe meaning of “like dissolves like” in terms of polar and nonpolar substances. Listthe three interactions that contribute to the enthalpy of a solution, and explain how they combine to cause dissolution to be exothermic or endothermic. Comparethe effects of temperature and pressure on solubility.

25. Factors Affecting the Rate of Dissolution 1. Increase the surface area -- because dissolution occurs at the surface 2. Stirring or shaking -- increases contact between solvent and solute 3. Increase the temperature -- increases collisions between solute and solvent and are of higher energy

26. Dissolving Process Video

27. Solubility • If you add spoonful after spoonful of sugar to tea, eventually no more sugar will dissolve. • There is a limit to the amount of solute that can dissolve in a solvent.

28. Solubility Values • Solubility: measure of how much solute can be dissolved in a specific amount of solvent at specific temperature • Specifically, how much is required to form a saturated solution • Example: solubility of sugar is 204 g per 100 g of water at 20OC • Solubilities vary widely and must be determined experimentally

29. Solutions • Saturated Solution: contains the maximum amount of dissolved solute • How does this happen? • When a solute is added, the molecules leave the solid surface and move about at random. • As more solute is added, more collisions occur between dissolved solute particles. Some of the solute molecules return to the crystal. • When maximum solubility is reached molecules are returning to the solid form at the same rate they are going into solution. • This is called solution equilibrium

30. Solutions • Unsaturated Solution: contains less solute than a saturated solution under the same condition • Supersaturated Solution: contains more dissolved solute than a saturated solution • Usually created by cooling a saturated solution slowly. • If disturbed, a supersaturated solution will precipitate out crystals of excess solute

31. Saturated Solution and Temperature

32. Mass of Solute Added Versus Mass of Solute Dissolved

33. Solubility of Common Compounds

34. Solubility of Common Compounds

35. Dissolving Ionic Compounds in Aqueous Solutions • What type of forces are present in water? • These forces attract the ions in ionic compounds and surround them, separating them from the crystal surface and drawing them into solution. • Hydration: when water is the solvent • Ions are said to be hydrated • A solute particle that is surrounded by solvent molecules is solvated. (like hydrated, but not water)

36. Hydration of Lithium Chloride

37. Liquid Solutes and Solvents • Oil and water do not mix because oil is nonpolar and water is polar. The hydrogen bonding squeezes out whatever oil molecules may come between them. • Two polar substances or two nonpolar substances easily form solutions. • Immiscible: liquids that are not soluble in each other • Miscible: liquids that dissolve freely in one another in proportion

38. Comparing Miscible and Immiscible Liquids Video

39. Temperature on Solubility • In general, an increases in temperature usually increases the solubility of solids in liquids • A few solid solutes are actually less soluble at higher temperatures.

40. Pressure and Solubility • Increases in pressure, increase gas solubilities in liquids • An equilibrium is established between a gas above a liquid solvent and the gas dissolved in a liquid

41. Pressure and Solubility Cont. • Increasing the pressure… causes gas particles to collide with the liquid surface and forces more gas into the solution • Decreasing the pressure… allows more dissolved gas to escape from solution

42. Soda Carbonation Video

43. Henry’s Law • 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 • In carbonated beverages, the solubility of carbon dioxide is increased by increasing the pressure. The sealed containers contain CO2 at high pressure, which keeps the CO2 dissolved in the beverage, above the liquid. • When the beverage container is opened, the pressure above the solution is reduced, and CO2 begins to escape from the solution. • Effervescence: the rapid escape of a gas from a liquid

44. Henry’s Law Video

45. Effervescence

46. Temperature and Solubility of Gas • Increasing the temperature… Increases the average kinetic energy and allows more solute to escape the attraction of the solvent and escape to the gas phase • At higher temperatures, equilibrium is reached with fewer gas molecules in solution

47. Enthalpies of Solution • The formation of a solution is accompanied by an energy change. • The formation of a solid-liquid solution can either absorb energy (KI in water) or release energy as heat (NaOH in water)

48. Enthalpies of Solution Cont • Before dissolving begins, solute particles are held together by intermolecular forces. Solvent particles are also held together by intermolecular forces. • Energy changes occur during solution formation because energy is required to separate solute molecules and solvent molecules from their neighbors.

49. Enthalpy changes during the formation of a solution

50. Enthalpies of Solution Cont. • Enthalpy of Solution: net amount of energy absorbed as heat by the solution when a specific amount of solute dissolves in a solvent • Negative – energy is released (exo) • (Steps 1/2 attractions is less than Step 3) • Positive – energy is absorbed (endo) • (Steps 1/2 attractions is greater than Step3)