Mixtures, Water, & Solutions

1. Mixtures, Water, & Solutions Chapter 2,15, & 16

2. Water Properties • Water is a polar molecule. • Hydrogen bonds form between water molecules. • High surface tension • Low vapor pressure • Maximum density at 4oC (as a liquid!) • What would happen to aquatic life if ice were denser than water?

3. Solutions (Homogeneous Mixtures) • Solute: becomes dispersed in the solvent • Solvent: dissolves the solute • Aqueous Solutions:substance dissolved in water. (water = solvent)

4. Examples Kool – Aid Solute – powered substance Solvent – water Chocolate Milk Solute – chocolate powder or syrup Solvent – milk

5. Identify the solvent and the solute • A teaspoon of sugar is dissolved in 200.2g of water. • Sterling silver is made by adding small amounts to copper to pure silver. • Jell-O consists of solid particles that were dissolved and then left suspended in water. • NaCl(aq)

6. “Like Dissolves Like” • Dissolving is a physical process in which particles of a solute are held apart by particles of the solvent. • RULE: Like Dissolves Like • Nonpolar solvents (ex. oil) dissolve nonpolar compounds. • Polar solvents (ex. water) dissolve polar compounds and most ionic compounds (ex. NaCl).

7. “Like Dissolves Like” Examples • Oil (nonpolar) does not dissolve in water (polar). • Alcohols (ex. CH3CH2OH – ethanol) have both a polar and a nonpolar end. • Dissolve polar and nonpolar solutes, but NOT ionic compounds.

8. What happens when compounds dissolve? Dissolving Ionic Compounds • NaCl (s)  Na+ (aq) + Cl- (aq)   • Dissociation in Water causes Ions to formed.

9. What happens when compounds dissolve? Dissolving Covalent Compounds • C12H22O11 (s)  C12H22O11 (aq) • NO dissociation because NO ions • Sucrose dissolves in water because sugar is polar (-OH group), but dissociation does not occur. Sucrose molecules are simply separated from each other. No ions are formed

10. Electrolytes • Electrolyte: a compound that conducts an electric current when it is in an aqueous state. • Mobile ions are required for the conduction of electric current. • Ex. Ionic Compounds (salts), acids, and bases. • Nonelectrolytes: cannot conduct electricity • Ex. glucose, alcohol

11. Solubility • Solubility: The amount of solute that dissolves in a given quantity of a solvent at a specified temperature and pressure. Concentration of Solute in a Solution • Concentrated: Relatively more solute in the solution. • Dilute: Relatively less solute in the solution. • Example: A 0.02M solution is more dilute than a 2M solution.

12. Saturation • Unsaturated Solution: the amount of solute dissolved is less than the maximum that could be dissolved. • Ex. Earth’s oceans = unsaturated salt solution • The ocean can hold a LOT more salt than it actually has in it! • Saturated Solution: the solution holds the maximum amount of solute. • Supersaturated Solution: contains more solute than the usual maximum amount and is unstable (may release solute suddenly). • Created by dissolving solute in the solution at a high temperature then slowly cooling the solution.

13. SolubilityCurves • How many gramsof potassiumchloride would dissolve in 100 g of water at 90oC? • How much woulddissolve in 200 g of water? 55 grams 55 grams x 2 = 110 grams

14. Solubility Curves • A solution has 132 g of NaNO3in 100 g of waterat 75oC.Is the solution saturated, unsaturated, orsupersaturated? • What about 95 gof KNO3 in 100gwater at 50oC? UNSATURATED – falls BELOW the line of saturation SUPERSATURATED – falls ABOVE the line of saturation

15. Factors Affecting Solubility • Temperature • Solids dissolving in liquids solubility increases with increasing temperature • Gases dissolving in liquids solubility decreases with increasing temperature • Carbonated soda • Thermal pollution from industry affecting dissolved oxygen in lakes

16. Factors Affecting Solubility • Pressure • Solids and Liquids little effect • Gases solubility increases under increased pressure. - ex. carbonated beverages • Stirring / Solute Particle Size / Solute Surface Area (crushing)

17. Calculating the Concentration of Solutions • Molarity (M) = moles of solute liters of solution • 3MNaCl = “three molar solution of sodium chloride”

18. Molarity Practice • What is the molarity of a solution in which 67 g of NaCl are dissolved in 1 L of solution? • How many grams of KNO3 should be used to prepare 2 L of a 1 M solution?

19. Making Dilutions • When you dilute a solution, you increase the amount of solvent (the number of solute particles stays the same). • M1 xV1 = M2 xV2 • M = molarityV = volume (in mL or L  must be same for both)

20. Practice • A chemist starts with 50.0 mL of a 0.40 MNaCl solution and dilutes it to 1000. mL. What is the concentration of NaCl in the new solution? • If you dilute 175 mL of a 1.6 M solution of LiCl to 1.0 L, determine the new concentration of the solution. • A chemist wants to make 500. mL of 0.050 MHCl by diluting a 6.0 MHCl solution. How much of that solution should be used?

21. Solution Compared to Pure Solvent • Freezing-Point Depression – presence of solute particles disrupts formation of orderly pattern found in solid phase. • Ex. Antifreeze • Vapor-Pressure Lowering (nonvolatile solute) • Boiling-Point Elevation – lowers vapor pressure • For aqueous solutions, BP will be higher than 100oC

22. Acids & Bases

23. Acids • Sour taste • Examples: • HCl (stomach acid), H2SO4 • HC2H3O2 (acetic acid) – (vinegar = acetic acid & water) • Citric acid (lemon juice) • Releases H+ when dissolved in water, producing hydronium ions! H+ + H2O  H3O+ • Electrolyte

24. Acids (cont’d) • React with metals to produce H2 gas. • Ex. HCl + Zn  ZnCl2 + H2 • When diluting acids, always slowly pour the acid into water while stirring. • Acid/Base Indicators: • Turns litmus paper red. • Phenolphthalein does not change color.

25. Bases • Bitter taste, slippery feel • Examples: NaOH, Mg(OH)2 (milk of magnesia), NH3 (ammonia), soap, household cleaners • Releases OH- (hydroxide ions) when dissolved in water. • Electrolyte. • Acid/Base Indicators: • Turns litmus paper blue. • Phenolphthalein turns bright pink.

26. Naming Common Acids Textbook Table 9.5

27. 3 Definitions of Acids & Bases 1) Arrhenius Theory • Acids: ionize to produce H+ ions in aqueous solution • Monoprotic: HNO3, HCl, HC2H3O2 • Diprotic: H2SO4, H2SO3 • Triprotic: H3PO4 • Bases: dissociate to produce OH- ions in aqueous solution • NaOH, KOH, Ca(OH)2, Al(OH)3

28. Definitions (cont’d) 2) Brønsted-Lowry Theory • Hydrogen ion (H+) = a proton • Acids: proton donors – ex. HCl • Bases: proton acceptors – ex. NH3

29. Conjugate Acids and Bases • Show the direction of H+ transfer. • Label: Acid, Base, Conjugate Base, Conjugate Acid • NH3 (aq) + H2O (l) NH4+ (aq) + OH- (aq) • HCl (g) + H2O (l) H3O+ (aq) + Cl- (aq)

30. More Examples + WS • Show the direction of H+ transfer. • Label: Acid, Base, Conjugate Base, Conjugate Acid • H2SO4 + OH- HSO41- + H2O • HSO41- + H2O SO42- + H3O+

31. Definitions (cont’d) 3) Lewis Theory • Acids: electron-pair acceptor • Bases: electron-pair donor • HCl (g) + H2O (l) H3O+ (aq) + Cl- (aq)

32. Strong Acids and Bases • Strong Acids: completely ionize in water • Ex. HCl, HBr, HI, HNO3, H2SO4, HClO3 • Strong Bases: completely dissociate into ions in water • Ex. NaOH, LiOH, KOH, Ca(OH)2, Sr(OH)2, Ba(OH)2

33. Weak Acids and Bases • Weak Acids: only some molecules ionize in water • Ex: acetic acid (less than 0.5% of molecules ionize) • Weak Bases: do not completely dissociate into ions in water • Ex: ammonia (only 0.5% of molecules dissociate)

34. Concentrated vs. Strong • “Concentrated” – refers to the amount dissolved in solution. • “Strong” – refers to the fraction of molecules that ionize.   • For example, if you put a lot of ammonia into a little water, you will create a highly concentrated solution. However, since only 0.5% of ammonia molecules ionize in water, this basic solution will not be very strong.

35. pH Scale

36. pH Scale • pH Scale: logarithmic scale in which [H3O+] is expressed as a number from 0 to 14. • [OH-] = [H3O+] when pH = 7.

37. pH Scale AcidicNeutralBasic 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 100 10-1 10-2 10-3 10-4 10-5 10-6 10-7 10-8 10-9 10-10 10-11 10-12 10-13 10-14 [H3O+] 10-14 10-13 10-12 10-11 10-10 10-9 10-8 10-7 10-6 10-5 10-4 10-3 10-2 10-1 100 [OH-]

38. pH equations • pH= - log [H3O+] • pOH = - log [OH-] • pH + pOH = 14 • [H+] x [OH-] = 1 x 10-14

39. pH – Examples • What is the pH if [HCl] = 1 x 10-4 M? • What is the [H+] if the pH = 9? • What is the pH if [NaOH] = 1 x 10-2 M • What is the concentration of [OH-] if the pOH is 3? • What is the concentration of [H+] if the pOH is 10?

40. pH calculations – easy practice • What is the pH of the solution? • [H3O+] = 1 x 10-4 M • [H+] = 1 x 10-10 M • [HCl] = 1 x 10-2 M • What is the concentration of H3O+ if the pH is 5? • What is the concentration of H+ if the pH is 11?

41. (cont’d) • What is the pOH of each solution? • [OH-] = 1 x 10-4 M • [NaOH] = 1 x 10-10 M • What is the pH of a solution if the pOH is 4? • What is the pH of each solution? • [OH-] = 1 x 10-8 M • [KOH] = 1 x 10-3 M • What is the [H3O+] in a solution if [OH-] = 1 x 10-3 M? • What is the [OH-] in a solution if [H3O+] = 1x 10-5 M?

42. pH equations • pH= - log [H3O+] • pOH = - log [OH-] • pH + pOH = 14 • [H+] x [OH-] = 1 x 10-14

43. Practice – Using the equations Find the pH of the following solutions. Is the solution acidic or basic? • 0.01 M HCl • 0.050 M Ca(OH)2 • 2.6 x 10-12 M Mg(OH)2 • 1 x 10-7 M HC2H3O2 • Find the concentration of hydrogen ions if the pH is 3. • Find the concentration of hydroxide ions if the pH is 5.6. • Find the [H3O+] in a solution if [OH-] = 3 x 10-6 M

44. Warm Up • In the reaction CO2-+ H2O → HCO3-1+ OH -1, label the acid, base, conjugate acid and conjugate base. 2. If the hydrogen ion concentration of a solution is 10-10M, is the solution acidic, basic, or neutral?. 3. In a neutral solution, the [H+] would have to be ____. 4. What is the best description for a solution with a hydroxide-ion concentration of 1 x 10-4M? 5. Which type of solution is one with a pH of 13? 6. Which of these solutions is the most basic? a.[H ] = 1 x 10-2M b.[H ] = 1 x 10-11M c.[OH ] = 1 x 10-4M d.[OH ] = 1 x 10-13M 7. What volume of 0.200M HCl will neutralize 10.0 mL of 0.400 M KOH? 8. Calculate the pH if the [H3O +] is1.2 x 10-3M. 9. Calculate the [OH -] if the pOH is 13.5. 10. Calculate the [H3O+] if the pOH is 8.

45. Neutralization Reaction • ACID + BASE  SALT + WATER • Salt: ionic compound formed from the negative part of the acid and the positive part of the base. • Example:   2HCl + Mg(OH)2MgCl2 + 2H2O • What type of reaction is this? Synthesis, Decomposition, Single Replacement, Double Replacement, or Combustion?

46. Complete and Balance the Neutralization Reactions • HCl + NaOH • HC2H3O2 + Ca(OH)2 • HBr + Al(OH)3

47. Acid-Base Titration • Uses a neutralization reaction to determine the concentration of an acid or base. • Standard Solution: the reactant that has a known molarity • Endpoint: the point at which the unknown has been neutralized.

48. Titration Examples • Example #1) 8.0 mL of 0.100MNaOH is used to neutralize 20.0 mL of HCl. What is the molarity of HCl?

49. Titration Examples (cont’d) • Example #2) A 0.1M Mg(OH)2 solution was used to titrate an HBr solution of unknown concentration. At the endpoint, 21.0 mL of Mg(OH)2 solution had neutralized 10.0 mL of HBr. What is the molarity of the HBr solution?

50. Titration Practice Example #3) What is the molarity of an Al(OH)3 solution if 30.0 mL of the solution is neutralized by 26.4 mL of a 0.25 M HBr solution?