1. SOLUTIONS

2. RECALL TYPES OF MIXTURES: SUSPENSIONS COLLOIDS SOLUTIONS All mixtures are physically combined and can be physically separated.

3. DEFINITION A solution is a homogeneous mixture of two or more substance in a single physical state

4. Parts of a solution SOLUTE – the substance that is dissolved SOLVENT- the substance that does the dissolving

5. Definitions Solute Solute - KMnO 4 Solvent Solvent - H 2 O

6. TYPES OF SOLUTIONS SOLUTESOLVENTEXAMPLE GAS Air GASLIQUIDSeltzer (CO 2 ) LIQUID Antifreeze (ethyl glycol in water) SOLIDLIQUIDSea water ( salt in water) GASSOLIDCharcoal filter (poisonous gases in carbon) LIQUIDSOLIDDental filling (mercury in silver) SOLID Sterling silver (copper in silver)

7. SOLID SOLUTION Contain two or more metals called alloys Formed by melting the components and mixing them together and allowing them to cool Properties of alloys are different from the original component metals

8. TYPES OF ALLOYS ALLOYCOMPONENTUSES BabbittTin, antimony, copperBearings Bell metalCopper, tinBells Coinage metalsCopper, tin, zincCoins 16 karat goldGold, copper, silverJewelry SterlingSilver, copperJewelry, flatware NichromeNickel, iron, chromium, manganese Heating elements

9. GASEOUS SOLUTIONS All mixture of gases Properties depend on the properties of its components Example: Nitrogen in air serves as a gas that dilutes pure oxygen which is toxic to people and animals, and is very combustible.

10. LIQUID SOLUTIONS Most familiar type of solution The solvent and the solution are liquids Solute may be a gas, a solid, or a liquid It is proper to describe liquids that are soluble to each other as MISCIBLE or can mix. And insoluble liquids as IMMISCIBLE. Or cannot mix. Example: alcohol is miscible in water while oil is immiscible in water.

11. Important terminologies: Soluble – substance that dissolves another substance Insoluble – substance that does not dissolve another substance Miscible – liquids that are completely soluble in each other or can mix Immiscible – liquids that are not soluble in each other or cannot mix

12. AQUEOUS SOLUTIONS Solutions with water as the solvent Aqueus, means like or containing water. Substances that dissolve in water are classified according to whether they produce ions or molecules in solution. Solutions that conduct electricity are called ELECTROLYTES.

13. SOLUBILITY

14. 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

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

16. FACTORS AFFECTING SOLUBILITY

17. Polarity Factors Affecting Solubility Temperature Pressure

18. Polarity Factors Affecting Solubility Temperature Surface Area Stirring

19. “LIKE DISSOLVES LIKE” There is a general saying that "like dissolves like" (but this is a very general rule, with many exceptions).  Refers to the type of bonding between separate molecules  If the bonding between separate molecules of the substance to be dissolved is similar to the type of bonding between solvent molecules, there is a good chance that the substance will dissolve. NATURE OF SOLUTE AND SOLVENT

20. Intramolecular Bonding Intramolecular bonding refers to the chemical bonding that holds atoms together within a molecule of a compound Covalent bonding and ionic bonding are the two main types of intramolecular bonding

21.  Covalent bonding involves the sharing of valence electrons between two atoms. Eg.covalent bonding holds hydrogen and oxygen atoms together to form a water molecule, H 2 O. POLAR- unequal sharing of electrons NON POLAR – equal sharing of electrons

22. COVALENT BONDING

23. Ionic bonding involves the transfer of valence electrons from one atom to another. The electrostatic attraction between opposite charged ions holds the molecule (formula unit) together. An example is sodium chloride, NaCl, which involves the attraction between Na + and Cl - ions.

24. IONIC BONDING

25. Intermolecular Bonding  Intermolecular bonding, on the other hand, is what holds two or more separate molecules together in the solid and liquid phases. What type of intermolecular bonding is involved largely depends on two main factors: 1.whether the bonds within a single molecular are polar or not (an unequal distribution of charge between two atoms involved in a chemical bond due to an unequal sharing of electrons), and 2.the overall shape of the molecule (it's molecular architecture - tetrahedral, linear, bent, etc.).

26. POLAR MOLECULE A polar molecule will have one end of the molecule bearing a partial positive charge while another end carries a partial negative charge. Polar molecules must contain polar bonds.

27. The oxygen end of the molecule has a partial negative charge (δ-), while the hydrogen end is partially positive (δ+).

28.  Water is an example of a highly polar molecule. Not only are the individual H—O bonds very polar (the shared electrons sit much closer to oxygen than to hydrogen, because oxygen has a higher electronegativity), but because the molecule has a bent shape the molecule itself is also polar. The oxygen end of the molecule has a partial negative charge (δ-), while the hydrogen end is partially positive (δ+).  What holds one water molecule tightly to the next is the strong attraction between the δ+ hydrogen end of one water molecule and the δ- end of a different molecule.

29. NON POLAR MOLECULE Nonpolar molecules either have no positive and negative ends, because the bonds making up the molecule are nonpolar, or because the entire outer "edge" is negative while the core of the molecule is positive (or vice versa), thus having no oppositely charged ends.

30. SOLUTEPOLAR SOLVENT NONPOLAR SOLVENT IonicSolubleInsoluble PolarSolubleInsoluble NonpolarInsolublesoluble

31. The Effect of Temperature on Solubility Generally, increasing the temperature will increase solubility of solids and liquids. But increasing temperature will lower the solubility of gases (the gas will escape from solution, going back to the gas phase).

32. The Effect of Pressure on the Solubility of Gases Pressure has no effect in the solubility of solids and liquids but has a strong effect on the solubility of gases. The solubility of gases increases when the pressure above the gas is increased. In other words, more gas will dissolve when pressure is increased. This is known as HENRY’S LAW (William Henry, English chemist).

33. Dissolution- the rate at which a substance dissolves

34. 1.Particle size – area of solute particles exposed to the action of the solvent particles. Increase in surface area of the solute particles, solubility increases Example: fine table salt dissolves faster than rock salt FACTORS AFFECTING DISSOLUTION:

35. 2.Stirring or Agitation increases the solubility of solid solute particles in a solvent. Because it hastens the contact between the surface of the solute and the solvent particles

36. 3.Application of heat - solvent molecules move faster and come in contact frequently with the solute particles, increasing solubility. Except for other solutes where solubility hastens with decrease in the temperature of the solvent. Example: sodium hydroxide pellets dissolves slowly in hot water than in cold water.

37. 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

38. Solubility Solids are more soluble at... Solids are more soluble at... – high temperatures.  Gases are more soluble at... low temperatures & high pressures (Henry’s Law). EX: nitrogen narcosis, the “bends,” soda

39. Solubility Table LeMay Jr, Beall, Robblee, Brower, Chemistry Connections to Our Changing World, 1996, page 517 0 10 20 30 40 50 60 70 80 90 100 Solubility vs. Temperature for Solids Solubility (grams of solute/100 g H 2 O) KI KCl 20 10 30 40 50 60 70 80 90 110 120 130 140 100 NaNO 3 KNO 3 HClNH 4 Cl NH 3 NaCl KClO 3 SO 2 shows the dependence of solubility on temperature gases solids

40. How to determine the solubility of a given substance? Find out the mass of solute needed to make a saturated solution in 100 cm 3 of water for a specific temperature(referred to as the solubility). This is repeated for each of the temperatures from 0ºC to 100ºC. The data is then plotted on a temperature/solubility graph,and the points are connected. These connected points are called a solubility curve.

42. How to use a solubility graph? A.IDENTIFYING A SUBSTANCE ( given the solubility in g/100 cm 3 of water and the temperature) Look for the intersection of the solubility and temperature.

43. Example:What substance has a solubility of 90 g/100 cm 3 of water at a temperature of 25ºC ?

45. Example: What substance has a solubility of 200 g/100 cm 3 of water at a temperature of 90ºC ?

47. B.Look for the temperature or solubility Locate the solubility curve needed and see for a given temperature, which solubility it lines up with and visa versa.

48. What is the solubility of potassium nitrate at 80ºC ?

49. At what temperature will sodium nitrate have a solubility of 95 g/100 cm 3 ?

50. At what temperature will potassium iodide have a solubility of 230 g/100 cm 3 ?

51. What is the solubility of sodium chloride at 25ºC in 150 cm 3 of water ? From the solubility graph we see that sodium chlorides solubility is 36 g.

52. Solubility in grams=unknown solubility in grams 100 cm 3 of waterother volume of water ___36 grams____=unknown solubility in grams 100 cm 3 of water150 cm 3 water Place this in the proportion below and solve for the unknown solubility. Solve for the unknown quantity by cross multiplying. The unknown solubility is 54 grams. You can use this proportion to solve for the other volume of water if you're given the other solubility.

53. C.Determine if a solution is saturated, unsaturated,or supersaturated. If the solubility for a given substance places it anywhere on it's solubility curve it is saturated. If it lies above the solubility curve, then it's supersaturated, If it lies below the solubility curve it's an unsaturated solution. Remember though, if the volume of water isn't 100 cm 3 to use a proportion first as shown above.

54. Temp. ( o C) Solubility (g/100 g H 2 O) KNO 3 (s) KCl (s) HCl (g) SOLUBILITY CURVE Solubility  how much solute dissolves in a given amt. of solvent at a given temp. below unsaturated:solution could hold more solute; below line on saturated:solution has “just right” amt. of solute; on line supersaturated:solution has “too much” solute dissolved in it; above the line

55. ToTo Sol. ToTo Solids dissolved in liquids Gases dissolved in liquids As T o, solubility

56. Sometimes you'll need to determine how much additional solute needs to be added to a unsaturated solution in order to make it saturated. For example,30 grams of potassium nitrate has been added to 100 cm 3 of water at a temperature of 50ºC.

57. How many additional grams of solute must be added in order to make it saturated? From the graph you can see that the solubility for potassium nitrate at 50ºC is 84 grams

58. If there are already 30 grams of solute in the solution, all you need to get to 84 grams is 54 more grams ( 84g-30g )

59. Solubility Table LeMay Jr, Beall, Robblee, Brower, Chemistry Connections to Our Changing World, 1996, page 517 shows the dependence of solubility on temperature 0 10 20 30 40 50 60 70 80 90 100 Solubility vs. Temperature for Solids Solubility (grams of solute/100 g H 2 O) KI KCl 20 10 30 40 50 60 70 80 90 110 120 130 140 100 NaNO 3 KNO 3 HClNH 4 Cl NH 3 NaCl KClO 3 SO 2 gases solids

60. Classify as unsaturated, saturated, or supersaturated. per 100 g H 2 O 80 g NaNO 3 @ 30 o C 45 g KCl @ 60 o C 50 g NH 3 @ 10 o C 70 g NH 4 Cl @ 70 o C =unsaturated =saturated =unsaturated =supersaturate d 0 10 20 30 40 50 60 70 80 90 100 Solubility vs. Temperature for Solids Solubility (grams of solute/100 g H 2 O) KI KCl 20 10 30 40 50 60 70 80 90 110 120 130 140 100 NaNO 3 KNO 3 HClNH 4 Cl NH 3 NaCl KClO 3 SO 2 gases solids

61. So sat. pt. @ 40 o C for 500 g H 2 O = 5 x 66 g = 330 g 120 g < 330 g unsaturated saturation point @ 40 o C for 100 g H 2 O = 66 g KNO 3 Per 500 g H 2 O, 120 g KNO 3 @ 40 o C Solubility vs. Temperature for Solids Solubility (grams of solute/100 g H 2 O) KI KCl 20 10 30 40 50 60 70 80 90 110 120 130 140 100 NaNO 3 KNO 3 HClNH 4 Cl NH 3 NaCl KClO 3 SO 2 gases solids 0 10 20 30 40 50 60 70 80 90 100

62. (A) Per 100 g H 2 O, 100 g Unsaturated; all solute NaNO 3 @ 50 o C. dissolves; clear solution. (B) Cool solution (A) very Supersaturated; extra slowly to 10 o C. solute remains in solution; still clear. Describe each situation below. (C) Quench solution (A) in Saturated; extra solute an ice bath to 10 o C. (20 g) can’t remain in solution, becomes visible.

63. 1.a. 80 gb.42 g 2.b. 42 g KNO 3 = 25g KNO 3 = 60 g 100 g H 2 O x g H 2 O 3.a. 7.7 mg at 30 o C b.9.2 mg 20 o C 4.9.2 mg O 2 = x mg O 2 1,000 g H 2 O 1,000 g H 2 O= =0.92 mg O 2