Solids , Liquids and Solutions

1. Solids, Liquids and Solutions

2. Intermolecular Forces Forces of attraction between different molecules rather than bonding forces within the same molecule. • Dipole-dipole attraction • Hydrogen bonds • Dispersion forces 1

3. Forces and Phases • Substances with very little intermolecular attraction exist as gases • Substances with strong intermolecular attraction exist as liquids • Substances with very strong intermolecular (or ionic) attraction exist as solids 2

4. Phase Differences Solid – definite volume and shape; particles packed in fixed positions; particles are not free to move Liquid – definite volume but indefinite shape; particles close together but not in fixed positions; particles are free to move Gas – neither definite volume nor definite shape; particles are at great distances from one another; particles are free to move 3

5. 4 Three Phases of Matter

6. Effect of Pressure on Boiling Point 5

7. Types of Solids • Crystalline Solids: highly regular arrangement of their components [table salt (NaCl), pyrite (FeS2)]. 6

8. Representation of Components in a Crystalline Solid Lattice: A 3-dimensional system of points designating the centers of components (atoms, ions, or molecules) that make up the substance. 7

9. Unit Cell The smallest portion of a crystal lattice that shows the three-dimensional pattern of the entire lattice 8

10. Packing in Metals Model: Packing uniform, hard spheres to best use available space. This is called closest packing. Each atom has 12 nearest neighbors. 9

11. Types of Solids • Amorphous solids: considerable disorder in their structures (glass and plastic). 10

12. constant Temperature remains __________ during a phase change. Water phase changes 11

13. Phase Diagram • Represents phases as a function of temperature and pressure. • Critical temperature: temperature above which the vapor can not be liquefied. • Critical pressure: pressure required to liquefy AT the critical temperature. • Critical point: critical temperature and pressure (for water, Tc = 374°C and 218 atm). 12

14. Phase changes by Name 12

15. Water Water 13

16. Carbon dioxide Carbon dioxide 14

17. Carbon Carbon 15

18. Sulfur 16

19. Solutions are homogeneous mixtures Classification of Matter 17

20. Solute A solute is the dissolved substance in a solution. Salt in salt water Sugar in soda drinks Carbon dioxide in soda drinks Solvent A solvent is the dissolving medium in a solution. Water in salt water Water in soda 18

21. Solvents 19 Solvents at the hardware store

22. Dissolution of sodium Chloride 20

23. Concentrated vs. Dilute 21

24. Heat of Solution The Heat of Solution is the amount of heat energy absorbed (endothermic) or released (exothermic) when a specific amount of solute dissolves in a solvent. 22

25. Electrolytes vs. Nonelectrolytes The ammeter measures the flow of electrons (current) through the circuit. If the ammeter measures a current, and the bulb glows, then the solution conducts. If the ammeter fails to measure a current, and the bulb does not glow, the solution is non-conducting. 23

26. Definition of Electrolytes and Nonelectrolytes An electrolyte is: • A substance whose aqueous solution conducts • an electric current. A nonelectrolyte is: • A substance whose aqueous solution does not • conduct an electric current. Try to classify the following substances as electrolytes or nonelectrolytes… 24

27. Electrolytes? • Pure water • Tap water • Sugar solution • Sodium chloride solution • Hydrochloric acid solution • Lactic acid solution • Ethyl alcohol solution • Pure sodium chloride 25

28. Answers to Electrolytes ELECTROLYTES: NONELECTROLYTES: • Tap water (weak) • NaCl solution • HCl solution • Lactate solution (weak) • Pure water • Sugar solution • Ethanol solution • Pure NaCl But why do some compounds conduct electricity in solution while others do not…? 26

29. Suspensions and Colloids Suspensions and colloids are NOT solutions. Suspensions: The particles are so large that they settle out of the solvent if not constantly stirred. Colloids: The particle is intermediate in size between those of a suspension and those of a solution. 27

30. Types of Colloids 28

31. The Tyndall Effect Colloids scatter light, making a beam visible. Solutions do not scatter light. Which glass containsa colloid? colloid solution 29

32. Factors Effecting Solubility • The solubility of MOST solids increases with temperature. • The rate at which solids dissolve increases with increasing surface area of the solid. • The solubility of gases decreases with increases in temperature. • The solubility of gases increases with the pressure above the solution. 30

33. Therefore… Solids tend to dissolve best when: • Heated • Stirred • Ground into small particles Liquids tend to dissolve best when: • The solution is cold • Pressure is high 31

34. Saturation of Solutions • A solution that contains the maximum amount of solute that may be dissolved under existing conditions is saturated. • A solution that contains less solute than a saturated solution under existing conditions is unsaturated. • A solution that contains more dissolved solute than a saturated solution under the same conditions is supersaturated. 32

35. Solubility Chart 33

36. Calculations of Solution Concentration Concentration - A measure of the amount of solute in a given amount of solvent or solution Grams per liter - the mass of solute divided by the volume of solution, in liters Molarity - moles of solute divided by the volume of solution in liters Parts per million – the ratio of parts (mass) of solute to one million parts (mass) of solution Percent composition - the ratio of one part of solute to one hundred parts of solution, expressed as a percent 34

37. Colligative Properties Colligative properties are those that depend on the concentration of particles in a solution, not upon the identity of those properties. • Boiling Point Elevation • Freezing Point Depression • Osmotic Pressure 35

38. Freezing Point Depression Each mole of solute particles lowers the freezing point of 1 kilogram of water by 1.86 degrees Celsius. Kf = 1.86 C  kilogram/mol 36

39. Boiling Point Elevation Each mole of nonvolatile solute particles raises the boiling point of 1 kilogram of water by 0.51 degrees Celsius. Kb = 0.51 C  kilogram/mol 37

40. Freezing Point Depression and Boiling Point Elevation Constants 38

41. Properties of Acids • Acids taste sour • Acids effect indicators • Blue litmus turns red • Methyl orange turns red • Acids have a pH lower than 7 • Acids are proton (hydrogen ion, H+) donors • Acids react with active metals, produce H2 • Acids react with carbonates • Acids neutralize bases 39

42. Acids you SHOULD know: Strong Acids Weak Acids Sulfuric acid, H2SO4 Phosphoric acid, H3PO4 Hydrochloric acid, HCl Acetic acid, HC2H3O2 Nitric acid, HNO3 40

43. Sulfuric Acid • Highest volume production of any chemical in the U.S. • Used in the production of paper • Used in production of fertilizers • Used in petroleum refining Thick clouds of sulfuric acid are a feature of the atmosphere of Venus. (image provided by NASA) 41

44. Nitric Acid • Used in the production of fertilizers • Used in the production of explosives • Nitric acid is a volatile acid – its reactive components evaporate easily • Stains proteins (including skin!) 42

45. Hydrochloric Acid • Used in the pickling of steel • Used to purify magnesium from sea water • Part of gastric juice, it aids in the digestion of protein • Sold commercially as “Muriatic acid” 43

46. Phosphoric Acid • A flavoring agent in sodas • Used in the manufacture of detergents • Used in the manufacture of fertilizers • Not a common laboratory reagent 44

47. Acetic Acid • Used in the manufacture of plastics • Used in making pharmaceuticals • Acetic acid is the acid present in vinegar 45

48. Acids are Proton Donors Monoprotic acids Diprotic acids Triprotic acids H3PO4 HCl H2SO4 HC2H3O2 H2CO3 HNO3 46

49. Strong Acids vs. Weak Acids Strong acids are assumed to be 100% ionized in solution (good proton donors). HCl H2SO4 HNO3 Weak acids are usually less than 5% ionized in solution (poor proton donors). H3PO4 HC2H3O2 Organic acids 47

50. Strong Acid Dissociation 48