1. Pre-AP 4/25 Turn in the solution poster you were given yesterday
Please pick up the blank paper from the side table.
Today we will start discussing the unit on solutions.
You will need something to write with and take notes on.
We will perform a lab on Thursday covering the topics discussed today.

2. Quick Warm-Up
I am going to display a table that classifies matter. You have seen this table before (at the beginning of the year).
I need you to copy this table on the blank piece of paper that you picked up and then with a marker or highlighter, you need to circle the portion of the table that relates to this unit.
I also need you to list TWO EXAMPLES OF EACH TYPE OF MATTER ON THE BACK OF THIS PAPER.

3. Pre-AP 4/27 Pick up the lab from the side table.
Take out something to write with and something to take notes on.
We will do a brief introduction to solutions today. And then you will perform a lab.
Hw: finish the lab questions.

4. Unit: Solutions

5. Types of Mixtures

6. Recall Heterogeneous mixture-non-uniform composition
Homogeneous mixture-uniform composition
AKA a solution

7. Classification of Matter
yes no
Can it be physically separated?
MIXTURE
PURE SUBSTANCE
Is the composition uniform? no
yes
Can it be chemically decomposed? no
yes
Homogeneous Mixture
(solution)
Heterogeneous Mixture
Compound
Element

8. Solution
A solution is a mixture that is so well mixed that is the same throughout, even down to the molecules. And those molecules can be separated physically
Ex: carbonated water, coke, rubbing alcohol, 14-carat gold, air

9. Parts of a Solution
Solution-Homogeneous mixture of two or more substances in a single phase
Solvent-dissolving medium
What something is being dissolved into
Solute-substance dissolved
If filtered, would pass through filter paper

10. Concentration of Solutions (we will learn to calculate later)
The more solute you have, the more concentrated the solution.
Example: sweet tea, unsweet tea, semi-sweet tea

11. Mixtures that are NOT solutions: Suspensions
A temporary mixture in which the solvent particles are so large that they settle out unless the mixture is constantly stirred or agitated.
Particles can be separated using filter paper
Example-muddy water

12. Mixtures that are NOT solutions: Colloids
A mixture that has larger particles. The particles come in clusters, not single molecules (like in solutions) and they don’t settle (like in suspensions)
Examples: mayonnaise, egg whites, milk

13. Colloids vs solutions
Tyndall effect-scattering of light by colloidal particles is a transparent medium
Can be used to distinguish between colloids and solutions

14. Types of Mixtures Solutions Colloids Suspensions Homogenous
Heterogeneous
Particle size: molecular (smallest)
Particle size: Slightly larger in clusters
Particle size: larger particles (often visible)
Do not separate on standing
Particles settle out
Cannot be separated by filtration
Can be separated by filtration
Do not scatter light
Scatter light (Tyndall effect)
May scatter light, but are not transparent
Table 13-3, page 398

15. Solutions
Recall-A homogenous mixture of 2 or more substances in a single phase
Doesn’t settle
Cannot be filtered
Does not exhibit Tyndall effect
Solutes can be classified according to whether they yield molecules or ions in solution

16. Electrolytes
When an ionic compound dissolves in water, the cations and anions separate from each other
Ions are free to move and make it possible for an electric current to pass through the solution
Electrolyte-a substance that dissolves in water to give a solution that conduct electricity
Examples-any soluble ionic compound

17. Nonelectrolytes
A solution containing only neutral solute molecules does not conduct electric current because it does not contain any mobile charged particles.
Nonelectrolyte-a substance that dissolves in water to give a solution that does not conduct an electric current.
Example-sugar

18. Electrolyte vs. Nonelectrolyte- +
sugar - +
acetic acid - +
salt
Non-
Electrolyte
Weak
Electrolyte
Strong
Electrolyte
solute exists as
molecules only
solute exists as
ions and
molecules
solute exists as
ions only
View animation online.

19. Electrolyte or Nonelectrolyte?
SO2
Calcium carbonate dissolved in water
Glucose
KCl
Tap water
Distilled water
Gatorade

20. Pre-AP: 5/1 Take out a piece of paper and something to write with.
WE will review what you learned last week and then go over the solution process.
I will not see you again until Thursday. I plan to give you no HW tonight, however if we don’t finish, you will have HW.

21. REVIEW…
1. Define homogeneous mixture and describe the components of the mixture. List three examples.
2. What is the difference between a colloid, a suspension, and a solution? Describe each.
3. How can you tell the difference in a colloid and a solution?
4. Define electrolyte, nonelectrolyte, and a weak electrolyte. Give an example of each and draw the disassociation of each in water.
5. Draw three water molecules demonstrating Hydrogen Bonds.
6. Water is polar. Why?
7. What is surface tension and why does it occur?

22. The Solution Process

23. Factors Affecting Dissolution Rate
3 major factors affect how quickly a SOLID solute can be dissolved
1. Surface area of solute
The more surface area that is available for contact with the solvent, the more quickly a solute will go into solution
Example-sugar cube vs. granulated sugar
2. Agitation of solution
Increases contact of solvent with solute, speeding up dissolution
3. Temperature of solvent
Increases the kinetic energy of the solvent particles, so collisions between the solvent & solute are more frequent & more forceful. Helps to separate the solute particles & disperse them more quickly.

24. Solubility
Why is there a limit to how much solute you can dissolve in a solvent?
Solution equilibrium (Ksp) -physical state in which the opposing processes of dissolution & crystallization of a solute occur at equal rates.
Even though nothing appears to be happening in a solution at equilibrium, solute particles are re-entering the solid phase at the same rate others are leaving it.

25. Saturated vs. Unsaturated
Saturated solution-a solution that contains the maximum amount of dissolved solute
If more solute is added to the solution, it will settle to the bottom of the container & does not dissolve.
Unsaturated solution-a solution that contains less solute than a saturated solution under the existing conditions.

26. Supersaturated
When a saturated solution is cooled, excess solute usually falls out of solution.
Sometimes excess solute doesn’t separate, and the solution contains more solute than is should be able to at that temp.
Supersaturated solution-a solution that contains more dissolved solute than a saturated solution under the same conditions.

27. Saturation of Solutions
UNSATURATED SOLUTION
more solute dissolves
SATURATED SOLUTION
no more solute dissolves
SUPERSATURATED SOLUTION
becomes unstable, crystals form
concentration

28. Solubility OF SOLIDS
Maximum grams of solute that will dissolve in 100 g or 100 mL of solvent at a given temperature
Varies with temp
Based on a saturated soln
Note: rate as which dissolution occurs is not related to solubility.

29. 5/4: Pre-AP Turn in the practice solubility curve you received Monday
Pick up the solubility curve and the molarity practice worksheet from the side table.
You will need something to write with, a periodic table, and a calculator.
We will review how to read solubility curves, learn to calculate molarity and molarity by dilution, and then quiz over everything you learned so far.
HW: molarity practice worksheet
FYI: You will take your test over solutions on Monday. You will have a review and practice stoichiometry problems for HW over the weekend.

30. Solubility Curve Shows the dependence of solubility on temperature
THIS IS A “RECIPE” FOR A SATURATED SOLUTION.
SHOWS HOW MUCH SOLUTE YOU ADD TO 100 mL OF WATER AT CERTAIN TEMPERATURES TO OBTAIN A SATURATED SOLUTION.
FYI: 1 grams = 1 mL

31. Solubility OF GASES Solids are more soluble at...
high temperatures.
Gases are more soluble at...
1. low temperatures &
2. high pressures (Henry’s Law).
EX: nitrogen narcosis, the “bends,” soda

32. Solute-Solvent Interaction
“Like Dissolves Like”
NONPOLAR
POLAR

33. Liquid Solutes and Solvents
Immiscible-liquid solutes and solvents that are not soluble in each other
Example-oil and water (nonpolar and polar)
Miscible-liquids that dissolve freely in one another in any proportion
Example-ethanol and water (both polar)

34. Effect of Pressure on Gas Solubility
Henry’s law-the solubility of a gas in a liquid is directly proportional to the partial pressure of that gas on the surface of the liquid.
Example-carbonated beverages
At a bottling plant, CO2 is forced into a solution of flavored water by increasing the CO2 pressure to 5-10 atm and the placing the pressurized solution in a can or bottle. When the pressure is reduced to atmospheric pressure by opening the container, some of the gas rapidly comes out of solution as bubbles.

35. Heats of Solution
Energy can be either absorbed or released during the formation of a solution.
Dissolve KI in water  beaker feels cold (endothermic)
Dissolve LiCl in water  beaker feels warm (exothermic)
Heat of solution-net amount of heat energy absorbed or release when a specific amount of solute dissolves in a solvent.
Exothermic-H (Enthalpy) is negative; energy released; beaker would feel warm
Endothermic-H (Enthalpy) is positive; energy absorbed; beaker would feel cold

36. Concentration of Solutions

37. Concentration
A measure of the amount of solute in a given amount of solvent or solution.
2 ways to express concentration:
Molarity
Molality (we will never use this)

38. Solution Calculations
In order to perform stoichiometric calculations with solutions, we must know the nature of the solution and the amounts of chemical present.
M = moles of solute
liters of solution

39. Molarity Number of moles of solute in one liter of solution.
Represented by “M”
Units are mol/L or M or [ ]
You may be asked to convert from moles to grams or grams to moles.

40. Molarity example
What is the molarity of a 2.0 L solution that is made from 14.6 g of NaCl?

41. Molarity example
What is the molarity of a HCl solution that contains 10.0 g of HCl in 250 mL of solution?
Answer: 1.1M

42. Molarity example
How many moles of HCl exist in 500. mL of 0.50M solution of HCl?
Answer: 0.25 mol HCl

43. Molarity example
To produce 40.0 g of silver chromate, you will need at least 23.4 g of potassium chromate in solution as a reactant. All you have on hand in the stock room is 5.0 L of 6.0M K2CrO4 solution. What volume of the solution is needed to give you the 23.4 g of K2CrO4 needed for the reaction?
Answer:0.020 L of 6.0M K2CrO4 solution

44. MOLARITY EXAMPLE
To analyze the alcohol content of a certain wine, a chemist needs 1.0 L of an aqueous M K2Cr2O7 solution. How much solid potassium dichromate must be weighed out to make this solution?

45. MOLARITY EXAMPLE: Calculate the molarity of a solution made by dissolving 23.4g of sodium sulfate in enough water to form 125 mL of solution.

46. MOALRITY Ex. How many grams of Na2SO4 are required to make 350 mL of 0
MOALRITY Ex. How many grams of Na2SO4 are required to make 350 mL of 0.50 M Na2SO4?
0.350L mol Na2SO g Na2SO4 = 24.9g
1 L mol Na2SO4

47. Solution Dilution
The concentration calculations that we have done involved preparing a solution from scratch. We started with separate solvent and solute and figured out how much of each you would need to use.
Solutions can be prepared by diluting a more concentrated solution. For example, if you needed a one molar solution you could start with a six molar solution and dilute it.

48. Dilution
When calculating the concentration of a solution by dilution, you use:
FORMULA: M1V1=M2V2
M1=initial molarity (concentrated)
V1=initial volume
M2=final molarity
V2=final volume
*note: number of moles stays the same

49. How to dilute a solution…

50. A standard solution is one where concentration is accurately known
A standard solution is one where concentration is accurately known. Read procedure for using volumetric flasks and types of pipets. We will be using both in several labs this year.

51. Dilution example 1
A chemist starts with 50.0 mL of a 0.40 M NaCl solution and dilutes it to mL. What is the concentration of NaCl in the new solution?

52. Dilution Example 2
How many milliliters of water must be added to 30.0 mL of 9.0 M KCl to make a solution that is 0.50 M KCl?

53. Dilution problem (M1V1 = M2V2) (1.000M)(V1) = (0.250M)(500.0mL)
Dilution Ex. What volume of M KNO3 must be diluted with water to prepare mL of M KNO3?
Dilution problem (M1V1 = M2V2)
(1.000M)(V1) = (0.250M)(500.0mL)
V1 = 125 mL

54. Dilution Example 3
What volume of 16M sulfuric acid must be used to prepare 1.5 L of a 0.10 M sulfuric acid solution?

55. Pre-AP: 5/5 Turn in your solubility curve and the molarity homework
Pick up the review and the solution stoichiometry papers
You will also need a calculator and something to write with
We will practice a few problems over Molarity and Molarity by Dilution and then learn to perform solution stoichiometry (remember this will involve a reaction)
HW: Review
HW: Solution Stoichiometry
HW: Molarity WS

56. Molarity and Molarity by Dilution Review (the following questions are from the molarity practice worksheet due Monday)
# x 10-2 grams of Pb(C2H3O2)4 are dissolved to make 3.5 mL of solution.
#9. I have 340 mL of a 1.5 M NaCl solution. If I boil the water until the volume of the solution is 220 mL, what will the molarity of the solution be?
#10. How much water would I need to add to 500. mL of a 3.5 M KCl solution to make a 1.0 M solution?

57. Solution Stoichiometry
How many grams of silver chromate will precipitate when 150. mL of M silver nitrate are reacted with potassium chromate?
2 AgNO3(aq) K2CrO4(aq)  Ag2CrO4(s) KNO3(aq)

58. Solution Stoichiometry
What volume of M HCl is required to neutralize 20.0 mL of M sodium hydroxide?
HCl(aq) NaOH(aq)  NaCl(aq) H(OH)(l)