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 Chocolate Milk 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” RULE: 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 0.02 M solution is more dilute than a 2 M 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. Solubility Curves
How many grams of potassium chloride would dissolve in 100 g of water at 90oC?
How much would dissolve in 200 g of water?
55 grams
55 grams x 2 = 110 grams

14. Solubility Curves
A solution has 132 g of NaNO3 in 100 g of water at 75oC. Is the solution saturated, unsaturated, or supersaturated?
What about 95 g of KNO3 in 100g water 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
3M NaCl = “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 x V1 = M2 x V2
M = molarity V = volume (in mL or L  must be same for both)

20. Practice
A chemist starts with 50.0 mL of a 0.40 M NaCl 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 M HCl by diluting a 6.0 M HCl 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 Anion Ending Example Acid Name
-ide
Chloride, Cl-
Hydro-(stem)-ic acid
Hydrochloric acid
-ite
Sulfite, SO32-
(stem)-ous acid
Sulfurous acid
-ate
Phospate, PO43-
(stem)-ic acid
Phosphoric acid

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 Acidic Neutral Basic 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14
[H3O+]
[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 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 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 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 → HCO 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 M KOH?
8. Calculate the pH if the [H3O +] is 1.2 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 0.100M NaOH 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?

51. Equilibrium

52. Factors Affecting Chemical Reaction Rates
Temperature
Higher temperature = faster reaction (ex. baking a cake)
Lower temperature = slower reaction (ex. batteries in refrigerator)
Concentration of Reactants
Increased concentration = higher reaction rate (higher frequency of collisions between reactants).

53. Factors Affecting Chemical Reaction Rates
Smaller Particle Size/Greater Surface Area
Increases the amount of reactant exposed for reacting.
Catalysts/Inhibitors
Added to a reaction to speed it up (or slow it down).
Is not permanently changed or used up during the reaction.
A catalyst speeds up reaction by lowering activation energy.
An inhibitor interferes with the catalyst.
Enzymes = biological catalysts

54. LeChâtelier’s Principle
Regaining Equilibrium

55. 3 Stresses that Upset Equilibrium
Concentration: Changing the amount, or concentration, of reactants or products in a system at equilibrium disturbs the equilibrium.
Ex. Removing products to increase yield
Chicken Egg Production: Chicken egg
If the farmer removes the product (egg), the chicken produces more.
Ex. H2CO3 (aq) CO2 (aq) + H2O (l)
Add H2CO3 SHIFT _____________
Add CO2 SHIFT ___________
Remove CO2 SHIFT ___________

56. 3 Stresses that Upset Equilibrium
2. Temperature: Increasing the temperature causes the equilibrium to shift in the direction that absorbs heat. Think of heat as a product or reactant and treat like concentration changes. Ex. 2SO2 (g) + O2 (g) 2SO3 (g) + heat Increase temperature SHIFT ____________ Decrease temperature SHIFT __________

57. 3 Stresses that Upset Equilibrium
3. Pressure: A change in the pressure on a system affects only an equilibrium that has an unequal number of moles of gaseous reactants and products. Ex. Formation of Ammonia N2 (g) + 3H2 (g) 2NH3 (g) Increase pressure, shifts to side with fewer mol of gas SHIFT _________ Decrease pressure, shifts to side with more mol of gas SHIFT __________

58. Equilibrium
The rate of the forward reaction is equal to the rate of the reverse reaction.
The concentration of products and reactants stays the same, but the reactions are still running.
Shown with the double arrow.

59. Measuring equilibrium
At equilibrium the concentrations of products and reactants are constant.
We can write a constant that will tell us where the equilibrium position is.
Keq = [Products]coefficients [ [Reactants]coefficients
Keq = equilibrium constant
Square brackets [ ] means concentration in molarity (moles/liter)

60. Writing Equilibrium Expressions
General equation
nA + mB xC + yD
Keq = [C]x [D]y [A]n [B]m

61. Practice
Write the equilibrium expressions for the following reactions.
3H2(g) + N2(g) 2NH3(g)
2H2O(g) H2(g) + O2(g)

62. Calculating Equilibrium
Keq is the equilibrium constant, it is only effected by temperature.
Calculate the equilibrium constant for the following reaction H2(g) + N2(g) 2NH3(g)
If at 25ºC there 0.15 mol of N2 , 0.25 mol of NH3, and 0.10 mol of H2 in a 2.0 L container.

63. What it tells us If Keq > 1 Products are favored
More products than reactants at equilibrium
If Keq < 1 Reactants are favored

64. Practice
1. At 1500K, an equilibrium mixture of N2, O2, and NO gases consists of 6.4 mol/L of N2, 1.7 mol/L of O2, and mol/L of NO. What is the equilibrium constant for the system? What is favored at equilibrium?
N2 (g) + O2 (g) 2NO (g)
2. At equilibrium at 2500K, [HCl] = mol/L and [H2] = [Cl2] = mol/L for the reaction. Find the value of Keq. What is favored at equilibrium?
H2 (g) + Cl2 (g) HCl (g)

65. Warm Up What types of compounds would ionize in water?
What would happen if we increase temperature in a solution?
What mass of Na2SO4 is needed to make 2.5L of 2.0M solution?
If the pH is 9, what is the concentration of hydroxide ion?
What is the pH of a solution with a concentration of 0.01 M hydrochloric solution?
Why does oil and water not mix?
Write the equilibrium expression.

66. On your test… All work/scratch work on green sheet
For #13 please fill in:
Form A: A
Form B: A
When you are finished..staple green sheet to back of gradecam and bring up front
Work on packet for final exam