1. Units 15 & 16 Solutions & Acids and Bases

2. Solutions All solutions are composed of two parts: The solute and the solvent. The substance that gets dissolved is the solute The substance that does the dissolving is the solvent (Usually present in the larger amount) ***A solution may exist as a solid, liquid or gas depending on the state of the solvent.

3. Types of Solutions Gas in liquid Example – soda water Solute: carbon dioxide (gas) Solvent: water (liquid) Solid in solid Example – Steel Solute: carbon Solvent: iron

4. Types of Solutions Gas in gas Example - Air Solute: Oxygen Solvent: Nitrogen

5. Types of Solutions Liquid in liquid Example – Vinegar Solute: Acetic acid Solvent: Water Solid in liquid Example – Ocean Water Solute: Sodium Chloride (solid) Solvent: Water (liquid)

6. Aqueous Solution Any mixture where water is the solvent. Something is dissolved in water

7. Soluble - a substance that dissolves in another substance. Insoluble - a substance that does not dissolve in another substance. Solubility

8. Immiscible - two liquids that are insoluble in each other. Miscible - two liquids that are soluble in each other.

9. Solvation – a process that occurs when an ionic solute dissolves in a solvent Solvation (Hydration) Solvation Video

10. Solvents of a specific polarity or type will dissolve solute of similar polarities or types! Aqueous solutions of ionic compounds: The charged ends of the water molecules attract the positive and negative ions making up an ionic solid, forcing them to separate. Aqueous solutions of molecular compounds: Molecular compounds that have polar sections easily form aqueous solutions with water. “Like Dissolves Like”

11. Solubility – the maximum amount of solute that will dissolve in a given quantity of solvent at a specific temperature and pressure to produce a saturated solution Units for solubility: grams of solute per 100 g solvent Example: At 20˚C, NaNO 3 has a solubility of 74 g/100 g H 2 O Solubility

12. Saturated Solution - contains the maximum amount of dissolved solute Unsaturated Solution - contains less than the maximum amount of dissolved solute Supersaturated Solution – contains more solute than can theoretically be dissolved at a given temperature

13. Supersaturated Solutions How can you dissolve more solute than possible??

14. Supersaturation A supersaturated solution is created when a warm, saturated solution is allowed to cool without the precipitation of the excess solute Testing for saturation: add crystal of a solid and watch for crystallization.

15. Supersaturated Sodium Acetate One application of a supersaturated solution is the sodium acetate “heat pack.” Sodium acetate has an ENDOthermic heat of solution.

16. Supersaturated Sodium Acetate Sodium acetate has an ENDOthermic heat of solution. NaCH 3 CO 2 (s) + heat ----> Na + (aq) + CH 3 CO 2 - (aq) Therefore, formation of solid sodium acetate from its ions is EXOTHERMIC. Na + (aq) + CH 3 CO 2 - (aq) ---> NaCH 3 CO 2 (s) + heat

17. Solubility Curves Solubility of a solid generally increases with increasing temperature The higher the temperature, the greater amount of solute that will be dissolved in the solvent Solubility can be represented in a chart called a solubility curve

18. Solubility Curves

19. Factors that affect solubility Temperature- Generally, as temperature increases, more solid solute will dissolve in the same amount of liquid solvent. The opposite is true for gaseous solutes.

20. Concentration –the measure of the amount of solute dissolved in a given quantity of solvent

21. Solution Concentration Expressing concentration: Concentration DescriptionRatio Percent by mass Percent by volume Molarity

22. Percent By Mass Percent by mass -ratio of the solute’s mass to the solution’s mass expressed as a percent. Example: An aquarium contains 3.6 g NaCl per 100.0 g of water. What is the percent by mass of NaCl in the solution?

23. Percent By Volume Percent by volume-ratio of the volume of the solute to the volume of the solution expressed as a percent. Example: What is the percent by volume of ethanol in a solution that contains 35 mL of ethanol dissolved in 115 mL of water?

24. Molarity (M) = moles of solute liters of solution Moles of solute dissolved in 1 liter of solution Example: 0.23 M solution = 0.23 moles of solute dissolved in 1 L of solution

25. M is read as “molar” when next to a number 4 M HCl = 4 molar hydrochloric acid Keep in mind that the liters of solution takes into account the volume of the solute AND the volume of the solvent

26. Example: What is the molarity of a solution that contains 0.65 mol of CuCl 2 in 500 mL of water?

27. Example: What is the molarity of a solution that contains 5.10 g of glucose (C 6 H 12 O 6 ) in 100.5 mL of solution?

28. Preparing Molar Solutions How is a solution of known molarity made? Convert moles of solute to grams and measure the amount out. Add solvent so that the total volume of the solution is 1 L. For any volume other than 1 L we must adjust the amount of solute needed by multiplying it by the fraction of a liter of solution we need.

29. Example Example: How many grams of CaCl 2 would be dissolved in 1.0 L of water to make a 0.10 M solution of CaCl 2 ?

30. Example Example: How many grams of NaOH are in 250 mL of a 3.0 M NaOH solution?

31. Dilutions are used to decrease the concentration (or molarity) of a solution M 1 V 1 =M 2 V 2

32. Steps to Performing a Dilution 1. Calculate how many mL of the original (stock) solution to start with 2. Measure out the volume of stock solution (using a graduated cylinder or a pipet) and place in appropriately sized volumetric flask 3. Add water to the mark on flask

34. Example: What volume, in milliliters, of 2.00 M CaCl 2 is needed to make 0.50 L of 0.300 M CaCl 2 solution?

35. Colligative Properties Colligative means “depending on the collection.” Depends only on the number of dissolved particles, not on the identity of dissolved particles. Includes vapor pressure lowering, boiling point elevation, freezing point depression, and osmotic pressure

36. Electrolytes and Colligative Properties Electrolyte: Soluble ionic compounds. When they dissolve in solution, they dissociate into their component ions and conduct electricity. Ex: NaCl (s)  Na + (aq) + Cl - (aq) Covalent molecules in aqueous solution: Covalent particles do not dissociate when in solution, so the # of molecules = the # of particles.

37. Boiling Point Elevation Boiling occurs when vapor pressure equals atmospheric pressure. Boiling point of a solution is higher than the boiling point of the pure solvent. Dissolving substances increases the boiling point of a solvent. Examples of Colligative Properties

38. Freezing Point Depression Freezing point of a solution is lower than the freezing point of the pure solvent. Dissolving substances lowers the freezing point of a solvent. Ex: Icy pavement - throw down CaCl 2 or NaCl, and the water will then freeze at a lower temperature

39. Ex: Antifreeze: a solution of ethylene glycol in water 1. Prevents car’s radiator from freezing in the winter. 2. Prevents car’s radiator from boiling over in the summer The more ethylene glycol in the water, the lower the freezing point, and the higher the boiling point.

40. Acids & Bases

41. Properties of Acids Physical: Taste sour Chemical: React with metals to produce H 2 gas Neutralized when reacted with a base Litmus Indicator: Turns blue litmus paper red Ions in Solution: H +, H 3 O + (hydronium ion)

42. Properties of Bases Physical: Taste Bitter Slippery Chemical: Neutralized when reacted with an acid Do not react with metals Why are bases used as drain cleaners? Litmus Indicator: Turns red litmus paper blue Ions in Solution: OH -1

43. Arrhenius Acids & Bases Arrhenius Model: ACIDS:Acids contain the H + ion Ex.) HCl, HBr, HNO 3 BASES: Bases contain the OH -1 ion Ex.) NaOH, KOH, Ca(OH) 2

44. Bronsted-Lowry Acids & Bases Bronsted-Lowry Model: For every acid, there must be a base Acid = proton donor Base = proton acceptor HCl (aq) + NH 3 (aq)  NH 4 + (aq) + Cl -1 (aq)

45. Conjugate Pairs NH 3 / NH 4 + is a conjugate pair — related by the gain or loss of H + NH 3 / NH 4 + is a conjugate pair — related by the gain or loss of H + Every acid has a conjugate base, formed when H+ is removed from the acid. Every acid has a conjugate base, formed when H+ is removed from the acid. Every base has a conjugate acid, formed when H+ is added to the base. Every base has a conjugate acid, formed when H+ is added to the base.

46. Conjugate Pairs Identify the conjugate acid-base pairs: HCl in water NaOH in water NH 3 in water

47. Types of Acids Monoprotic and Polyprotic Acids Acids can contain 1 or more hydrogens that are acidic **Not ALL hydrogens are acidic (Ex. Vinegar) Identify the following as monoprotic or polyprotic: HNO 3, H 2 SO 4, HClO, HClO 4, H 3 PO 4, HC 2 H 3 O 2

48. Strength of Acids/Bases Strengths of Acids Strong Acid  Give off LOTS of H + 100% Dissociation Strong Acids: HCl, HI, HBr, HNO 3, H 2 SO 4, HClO 4 That’s it!  Everything else is “weak” Weak Acid  Give off smaller amounts of H + Equilibrium occurs (breaks apart and then recombines) Not all H + ions separate (not 100% dissociation)

49. Strength of Acids/Bases Strengths of Bases Strong Base  Give off LOTS of OH -1 100% Dissociation Generally, Group I, II Hydroxides (except H, Be, Mg) Ex.) Ca(OH) 2, NaOH Everything else is “weak” Weak Base  Give off smaller amounts of OH -1 Equilibrium occurs (breaks apart and then recombines) Not 100% dissociation

50. Strength of Acids/Bases Strong or weak vs. concentrated and dilute Strong/weak tells you how much it dissociates Concentrated/dilute indicates the concentration (amount of solute in the solvent)

51. pH pH & pOH pH tells us the acidity or basicity of a solution Based on measuring the [H + ] (a.k.a. [H 3 O + ]) pH Scale Ranges 0 to 14 Acid ~ 0 to 7 Bases ~ 7 to 14

52. Definition: Hydronium Ion In aqueous solution, H + does NOT exist! Note: In problems, [H + ] = [H 3 O + ] H + + H 2 O  H 3 O + (hydronium ion)

53. pH pH = - log [H 3 O + ] pOH = - log [OH - ] Make sure you have the negative sign! Find the “log” function on your calculator! pH + pOH = 14 [H + ][OH - ] = 1.0 x 10 -14

55. pH Calculations What is the pH and pOH for a solution with a H + concentration of 3.0 x 10 -6 M H + ?

56. pH Calculations What is the H + and OH - concentration of blood with a pH of 7.40?

57. Neutralization Reactions Neutralization reaction: Reaction in which acid and base react to neutralize one another Acid + Base  Water + Salt ***Salt = Any ionic compound formed as a by-product of an acid-base reaction

58. Neutralization Acid-base Titration: Definition: Lab technique which allows you to get moles of acid and base EXACTLY equal to another Complete neutralization Allows you to calculate the concentration of an unknown acid or base

59. Definitions The titrant is the substance of known concentration used to determine the unknown concentration of the other substance. An indicator- substance that changes color at a certain pH—is added to tell us when the neutralization is complete. Example: Phenolphthalein undergoes a color change between pH 8 and 10 clear in acid Light pink in neutral Dark pink in base

60. Neutralization Procedure: Add known volume of acid or base to Erlenmeyer flask Add a known concentration of the other to a buret Add an indicator to the flask Slowly dispense titrant (what you’re adding with a buret) into the flask Stop when 1 drop of titrant causes the indicator to switch from one color to another

61. Neutralization Equivalence point: pH at which amount of acid = base Indicator: Compound that changes color due to a change in pH Common Indicators and pH Range Litmus: 5.5 to 8.0 (red= acid, blue = base) Phenolphthalein: 8.2 to 10.6 (colorless to magenta) End point: Point at which the volume of titrant added makes the amount of acid and base are equal and the indicator changes color