1. Chapter 12 Solutions

2. Types of Mixtures: THINK SOLUTE PARTICLE SIZE!!!
Solutions: solute particles too small to see (<1nm)
Colloids: solute particles intermediate in size – they remain suspended throughout the solvent. (1-1000nm)
Examples: smoke, fog, foams
Suspensions: solute particles are so large they settle out upon standing. (>1nm)
The Tyndall Effect: Is it a true solution or a colloid???
The particles in a colloid will scatter light
A true solution will transmit light

3. The Tyndall Effect: Is it a true solution or a colloid???

4. Chapter 12, Table 12.10, Comparison of Solutions, Colloids, and Suspensions

6. Solutions
are homogeneous mixtures of two or more substances.
contain a solute, usually in smaller quantity, and uniformly dispersed in another substance called the solvent, usually present in greater quantity.

7. Solutes in a true solution
Solutes are not visible and
cannot be separated by filtration.
can be separated by evaporation.
can give color to a solution.

8. Types of Solutions – some examples
Solid in liquid: sugar in water
Solid in solid: metal alloys; 14-karat gold is a mixture of pure gold with silver and copper to make it more durable
Gas in liquid:
CO2 in water = carbonation,
NH3 in water = cleaning solution
Liquid in liquid: alcohol in water

9. Types of Solutes and Solvents
solutes and solvents can be solids, liquids, or gases.
Alloy: a solid solution of two or more metals

10. Solute–Solvent Attractions THE SOLUBILITY RULE: “like dissolves like”
The expression “like dissolves like” describes the polarities of the solute and solvent particles needed to form a solution.
NONPOLAR DISSOLVES NONPLAR
POLAR DISSOLVE POLAR AND IONIC
Attractive forces only occur when they are similar in polarity.

11. Water as a Solvent Water, the most common solvent, is a polar molecule
partial charges on H and O.
forms hydrogen bonds between molecules.

12. Solutions: Ionic solutes
Na+ & Cl− ions on the surface of a NaCl crystal are attracted to partial charges in polar water molecules.
oxygen atoms attract Na+ ions.
hydrogen atoms attract Cl− ions
Ions are hydrated by surrounding H2O molecules.

13. Solutions: Polar Solutes
Polar molecules such as CH3—OH are soluble because of the polar –OH group that forms hydrogen bonds with the polar solvent water.
Some polar solvents include water, H2O, acetic acid (vinegar), CH3COOH, and alcohols.

14. Solutions: Nonpolar Solutes
Solutes that are nonpolar, such as iodine (I2), oil, or grease, do not dissolve in the polar solvent water.
There are no attractions between the nonpolar solute particles and the polar solvent.
Oil and water don’t mix!
Nonpolar molecules will only form a solution with other nonpolar solvents.
Nonpolar solvents include hydrocarbons like grease, oil, gasoline.

15. Learning Check
Will the following solutes dissolve in water? A. Na2SO4 B. gasoline (nonpolar) C. I2 D. HCl

16. Solution Will the following solutes dissolve in water?
Na2SO4 Yes, the solute is ionic.
gasoline (nonpolar) No, the solute is nonpolar.
I2 No, the solute is nonpolar.
HCl Yes, the solute is polar.
Most polar and ionic solutes dissolve in water because water is a polar solvent.

17. The solution process & solids
To increase the rate of dissolution of the SOLID solute:
crush solute to increase Surface Area
stir to disperse
heat to increase movement

18. The solution process & gases
To increase the rate of dissolution of a solute gas, increase the pressure & decrease the temperature.
Just think soda!!!
Cold = low temperature
Lid on tightly = high pressure
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 a liquid.

19. Chapter 12, Unnumbered Figure, Page 370

20. Solutes: Strong Electrolytes
In water, strong electrolytes
completely dissociate.
produce ions.
conduct an electric current.
Examples: DISSOCIATION EQNs
100% ions
NaCl(s) Na+(aq) + Cl− (aq)
CaBr2(s) Ca2+(aq) + 2Br−(aq)
H2O
H2O
A strong electrolyte in an aqueous solution will completely dissociate into ions.

21. Dissociation of ionic solutes
The separation (dissolution) of ions that occurs when an ionic compound dissolves.
EXAMPLES: Dissociation Equations
CaCl2 (s) → Ca+2 (aq) + 2Cl-1 (aq)
Al2(SO4)3 (s) → 2Al+3(aq) + 3 SO4-2 (aq)
AlPO4 (s) → Al+3(aq) + PO4 -3 (aq)

22. You Try: Dissociation of Ionic Compounds
Na2SO4 (s) → 2Na+1(aq) +1SO4-2 (aq)
AlCl3 (s) → 1Al+3(aq) + 3 Cl-1 (aq)

23. How many moles of ions formed?
EXAMPLES: Dissociation Equations
CaCl2 (s) → 1Ca+2 (aq) + 2Cl-1 (aq)
Al2(SO4)3 (s) → 2Al+3(aq) + 3SO4-2 (aq)
AlPO4 (s) → 1Al+3(aq) + 1PO4 -3 (aq)
Na2SO4 (s) → 2Na+1(aq) +1SO4-2 (aq)
AlCl3 (s) → 1Al+3(aq) + 3 Cl-1 (aq)

24. Ionization of molecular acids
Some molecular compounds can form ions (where there were none) in solutions if the solvent is polar.
These are also strong electrolytes.
EX:
HCl (g) → H+1 (aq) + Cl-1 (aq) OR
HCl (g) + H2O (l) → H3O+1 (aq) + Cl-1 (aq)
H3O+1 = the hydronium ion = hydrated H+1

25. Solutes, Weak Electrolytes
In water, weak electrolytes
dissolve mostly as molecules.
produce only a few ions.
conduct a weak current.
Examples:
HF(g) + H2O H3O+(aq) + F−(aq)
NH3(g) NH4+(aq) + OH−(aq)
A weak electrolyte forms mostly molecules and a few ions in an aqueous solution.

26. Solutes, Nonelectrolytes
In water, nonelectrolytes
dissolve as molecules.
do not produce ions.
do not conduct a current.
Example:
C12H22O11(s) C12H22O11(aq)
H2O
A nonelectrolyte dissolves in water only as molecules,
which do not ionize.

27. Classifying Solutes

28. Check for understanding
1. Write the equation for the formation of a solution for each of the following.
2. Indicate whether solutions of each of the following contain only ions, only molecules, or
mostly molecules and a few ions.
a. Na2SO4(s), a strong electrolyte
b. sucrose, C12H22O11(s), a nonelectrolyte
c. acetic acid, CH3COOH(l), a weak electrolyte
d. Hydrochloric acid, H2SO4 (l), a strong electrolyte

29. Solutions of Electrolytes & Nonelectrolytes
a. An aqueous solution of Na2SO4(s) contains only the ions Na+ and SO42−.
b. A nonelectrolyte such as sucrose, C12H22O11(s), produces only molecules when it dissolves in water.
c. A weak electrolyte such as CH3COOH (l) produces mostly molecules and a few ions when it dissolves in water.
H2O
Na2SO4(s) Na+ (aq) + SO42–(aq)
H2O
C12H22O11(s) C12H22O11(aq)
H2O
CH3COOH(l) H+(aq) + CH3COO–(aq)

30. Solubility
For every combination of solute and solvent, there is a limit to the amount of solute that can dissolve.
Factors:
Nature of solute and solvent: polarity!
Nature of solvent: polarity!
The temperature
The pressure if a gas
Defined: The amount of substance required to form a saturated solution with a specific amount of solvent at a specified temperature.

31. Solubility & Solution Equilibrium
Eventually, solution equilibrium is reached:
The solute particles dissolve and return to the crystal at equal rates.
Defined:
The physical state in which the opposing processes of dissolution and crystallization of a solute occur at equal rates.

32. Solubility & Solution Equilibrium
Solutions can be:
Unsaturated: contains less than max. amt. of solute possible at that temp.
Saturated: contains max. amt. of solute possible at that temp.
Supersaturated: contains more than max. amt. of solute possible at that temp.

33. Chapter 12, Unnumbered Figure 4, Page 367

34. Supersaturated Solutions
If solution is saturated, additional solute will settle to the bottom
Heat and stir and more solute is forced to enter the solution
Solute remains dissolved even after solution cools down but will re-crystalize easily – it is unstable

35. Solubility Curve SOLUBILITY:
Given in grams of solute per 100.g of water at 20.C.
Determined experimentally and reported in reference tables.
Temperature must be specified.
For gases, pressure must also be specified

36. Practice Using the Solubility Curve
What is the solubility of NH3 at 50°C?
What is the solubility of KCl at 50°C?
What is the solubility of KNO3 at 50°C?
Which substances have lower solubility at higher temps? Why?

37. ANSWERS 30g 42g 88g The gases! HCl, NH3, and SO2
WHY? The solubility of gases ↓ as temp ↑

38. Solubility Rules

39. Practice Using the Solubility Rules
Determine if the following are soluble…
NaNO3
CaCO3
AgCl
Fe(OH)3
BaCl2
PbSO4

40. ANSWERS NaNO3 = YES = (aq) CaCO3 = NO = (s) AgCl = NO Fe(OH)3 = NO
Determine if the following are soluble…
NaNO3 = YES = (aq)
CaCO3 = NO = (s)
AgCl = NO
Fe(OH)3 = NO
BaCl2 = YES
PbSO4 = NO

41. Review
As temperature increases, the solubility of gases in liquids __________________
The substance dissolved is called the ____________
A mixture whose particles separate is a __________
A mixture with particles too small to see is a_______
A mixture that scatters light is a________________ and exhibits the _______________ effect.
A substance that d0es not dissolve in a polar solvent is probably __________________
To carry an electric current a solution must contain __________________

42. Review
Liquid solutes and solvents that are not soluble in each other are _immiscible__ and examples include….
A(n) __________________ is a solution whose solute and solvent are both solid metals.
The solubility of an ionic compound such as CuCl2 would __________________ with increasing temp.
A solution that contains the max. amount of dissolved solute is __________________
A solute molecule that is surrounded by solvent molecules is __________________
Name three alloys:

43. ANSWERS Decrease Solute Suspension solution Colloid, Tyndall nonpolar
immiscible
alloy
Increase
Saturated
Hydrated/solvated

44. Enthalpy of Solution KI (s) + energy → K+ (aq) + I- (aq)
The amount of energy absorbed as heat when a specific amount of solute dissolves in a solvent.
EX: The dissolving process
KI (s) + energy → K+ (aq) + I- (aq)
THIS PROCESS IS ENDOTHERMIC
NaOH (s) → Na+ (aq) + OH- (aq) + energy
THIS PROCESS IS EXOTHERMIC

45. Precipitation Reactions (from Ch 8) & Net Ionic Equations
Precipitation Reactions Review:
Double Replacement: C + C → C + C
Reactants are in solution: (aq)
The ions trade partners
There is no change in the charge on the ions
Often form a precipitate: use solubility rules
Net Ionic Equations:
Include only the compounds and ions that change.

46. Precipitation Reactions & Net Ionic Equations
EX: Reactants are aqueous solutions of zinc nitrate and ammonium sulfide:
Net Ionic Equation: includes only those compounds and ions that change

47. Precipitation Reactions & Net Ionic Equations
EX: Reactants are aqueous solutions of potassium sulfate and barium nitrate.
Net ionic equation:

48. Concentration of Solutions
Molarity: the number of moles of solute in one liter of solution.
𝒂𝒎𝒐𝒖𝒏𝒕 𝒐𝒇 𝒔𝒐𝒍𝒖𝒕𝒆 (𝒎𝒐𝒍𝒆𝒔) 𝒗𝒐𝒍𝒖𝒎𝒆 𝒐𝒇 𝒔𝒐𝒍𝒖𝒕𝒊𝒐𝒏 (𝒍𝒊𝒕𝒆𝒓𝒔)
Molarity (M)=

49. Concentration of Solutions
For example, a 0.25 M NaOH solution
(this is read as 0.25 molar)
contains 0.25 moles of the solute NaOH
in every liter of solution = 𝟎.𝟐𝟓 𝒎𝒐𝒍𝒆𝒔 𝑵𝒂𝑶𝑯 𝟏 𝒍𝒊𝒕𝒆𝒓 𝒐𝒇 𝒔𝒐𝒍𝒖𝒕𝒊𝒐𝒏 = 𝟎.𝟐𝟓 𝒎𝒐𝒍 𝟏 𝑳
Anytime you see the abbreviation M
immediately think of it as 𝒎𝒐𝒍𝒆𝒔 𝑳𝒊𝒕𝒆𝒓

50. Preparing Solutions
Note: a 1 M solution is NOT made by adding 1 mole of solute to 1 Liter of solvent because the final volume may be > 1 Liter!
Instead, dissolve 1 mole of solute in less than 1 Liter of solvent first, then add enough solvent to bring the total volume up to 1 Liter.

51. Molarity Problems – Finding Molarity
Given: g NaCl, 3.50L of solution, find Molarity
Given: 5.85g of KI, 0.125L of solution, find Molarity
Antifreeze is a solution of ethylene glycol in water. If 4.50L of antifreeze contains 27.5 grams of ethylene glycol, C2H6O2, what is the concentration of the solution?

52. Molarity Problems – Finding Molarity
What is the molarity of a solution that contains moles of H2SO4 in 2.50 L of solution?
What is the molarity of a solution prepared by dissolving 25.0 g of HCl (g) in in enough water to make mL of solution?

53. Molarity Problems – Molarity Given
If a reaction requires 146.3g of NaCl, what volume of 3.00M NaCl should be used?
How many grams of NaCl would you need to prepare 500. mL of a 2.00 M solution?

54. Molarity Problems – Molarity Given
A reaction requires 23.4g of K2CrO4. I have a 6.0M stock solution. How much of it should I use for the reaction?
How many mL of 0.54 M AgNO3 would contain 0.34g of the solute?

55. Dilutions: Changing Volume
In a dilution, the amount of solute does not change; only the volume of the solution changes. C1V1 = C2V2 where C is the concentration of the solution (it can be molarity or percent concentration), and V is the volume of the solution.
Grams or moles
of solute in
concentrated solution
Grams or moles of solute in diluted solution =

56. Dilutions: Changing Volume
When water is added to a concentrated solution, there is no change in the number of particles. The solute particles spread out as the volume of the diluted solution increases.

57. Molarity of a Diluted Solution
What is the molarity of a solution when 75.0 mL of a 4.00 M KCl solution is diluted to a volume of 500. mL?
STEP 1 Prepare a table of the concentrations and volumes of the solutions.

58. Molarity of a Diluted Solution
What is the molarity of a solution when 75.0 mL of a 4.00 M KCl solution is diluted to a volume of 500. mL?
STEP 2 Rearrange the dilution expression to solve for the unknown quantity.

59. Molarity of a Diluted Solution
What is the molarity of a solution when 75.0 mL of a 4.00 M KCl solution is diluted to a volume of 500. mL?
STEP 3 Substitute the known quantities into the dilution expression and calculate.

60. Strong Electrolytes STRONG BASES INCLUDE STRONG ACIDS INCLUDE HCl HBrHI
HNO3
HClO3
HClO4
H2SO4
NaOH
KOH
LiOH
Ba(OH)2
Ca(OH)2

61. Weak Electrolytes WEAK ACIDS INCLUDE WEAK BASES INCLUDE HF CH3COOH
H3PO4
H2CO3
NH3

62. REVIEW QUESTIONS Distinguish b/w ionization and dissociation:
Distinguish b/w electrolyte and nonelctrolyte:
Distinguish b/w strong and weak electrolytes:
Write the formula for aluminum chloride:
Write equation that shows the dissociation of aluminum chloride (a salt) in water:
How many moles of ions are released in solution when aluminum chloride dissolves?
Is aluminum chloride soluble?

63. AlCl3 (s) → Al+3 (aq) + 3Cl-1 (aq)
Answers
Distinguish b/w ionization and dissociation: use ionization for covalent compounds, dissociation for ionic compounds
Distinguish b/w electrolyte and nonelctrolyte: electrolytes yield ions in solution, non-electrolytes do not.
Distinguish b/w strong and weak electrolytes: strong electrolytes ionize completely, weak electrolytes do not
Write the formula for aluminum chloride: AlCl3
Write eqn that shows the dissociation of aluminum chloride (a salt) in water:
AlCl3 (s) → Al+3 (aq) + 3Cl-1 (aq)
How many moles of ions are released in solution when aluminum chloride dissolves? 4
Is aluminum chloride soluble? yes

64. Put 160 grams of sodium acetate in a flask and add 30 mL of water
Put 160 grams of sodium acetate in a flask and add 30 mL of water.Put the flask on a hot plate, heat it gently and stir until the crystals of sodium acetate dissolve. Use a small amount of water to rinse down the inside of the flask.Remove the flask from the heat and let it cool slowly without disturbing it.Add one or two crystals (that’s right, it only takes a single crystal) to the liquid in the flask. Don’t take your eyes off of the liquid as beautiful crystals begin to form inside the flask.Feel the flask… it’s warm!  - See more at:
If you attempt to dissolve sugar in water, you reach a point where you cannot dissolve any more sugar. This is called a saturated solution. However, if you heat this solution, more sugar will dissolve. When the solution is cooled, the sugar will remain in solution. This is called a supersaturated solution, which is very unstable and will crystallize easily. If you had really wanted to fool your mom when you added more sugar to the Kool-Aid, you should have heated the liquid to dissolve the extra sugar and then allowed it to cool. Your supersaturated Kool-Aid would have been "super sweet!" The process of crystallization gives off heat. It’s said to be exothermic. That’s why the solution is used in hand warmers (the old-style liquid-type of hand warmers).
How do Hand Warmers Work?  Commercially available hand warmers use a supersaturated solution of sodium acetate.These products consist of a concentrated aqueous salt solution together with a flexible metallic activator strip (usually stainless steel) in a sealed, flexible container. Sodium acetate and calcium nitrate are examples of suitable salts. These salts are much more soluble in hot water than in cold water. The flexible metal strip is bent back and forth a few times, whereupon a white cloud of crystals begins to precipitate. Within seconds, the entire pack is filled up with solid crystalline needles of sodium acetate without any solution left, and the temperature rises to 130°F for about 30 minutes. Because heat is released upon this precipitation, it is called an exothermic reaction (the opposite is called an endothermic reaction). Supercooled liquids can be cooled below their normal freezing point without turning solid. Then, at the flick of button, the supercooled liquid is triggered to solidify (crystallize) and at the same time release large amounts of heat. Salt solutions that have been processed in such a way that their temperature can be lowered well below their solidification (or melting) temperature and still remain in liquid are defined as supercooled or metastable liquids. The triggering device initiates the rapid solidification of the solution. In the case of salt solutions that release or absorb large amounts of energy during phase changes (common table salt sodium chloride does not do this), the solidification process is a rapid crystallization that releases a large amount of heat at the salt solution's normal melting temperature. The activator is a thin metal piece with ridges and a specially roughened surface. The flexing causes metal-to-metal contact that releases one or more very tiny particles of metal from the roughened surface. This acts as a nesting site for one crystal deposited from the solution and (voila!) all of the crystals fall out instantly. These heat packs are reusable because, by reheating the pack in boiling water for a few minutes, the salt re-dissolves and the pack again contains a clear solution. Best of all, the activator strip can be reused dozens of times!
- See more at:

65. Colligative Properties of Solutions
Properties that depend on the concentration and presence of solutes but not on their identity.
Colligative properties are proportional to the number of moles of particles the solute makes in solution.
The vapor pressure of a solvent containing a nonvolatile solute is LOWER than the vapor pressure of a pure solvent. (evaporates more slowly)
The greater the concentration of a solution, the greater its osmotic pressure.
A nonvolatile solute will elevate the Boiling Point
A nonvolatile solute will depress the Freezing Point

66. Why is the boiling point of a solution higher than the BPt of pure water?
Attractive forces exist b/w the solvent and solute particles.
It takes more E for solvent particles to overcome these forces and change state.
For example, the BPt of water ↑ 0.51 ˚C for every mole of solute particles.

67. Why is the freezing point of a solution lower than the FPt of pure water?
Attractive forces exist b/w the solvent and solute particles.
more E must be taken from the solution in order to change state and form a solid.
the FPt of water ↓ 1.86 ˚C for every mole of solute particles.

68. Why would sugar (a nonelectrolyte) have less of an effect than an ionic salt like MgCl2?
A solute that dissociates into several particles has a greater colligative effect than a solute that does not dissociate.

69. Why would sugar have less of an effect than an ionic salt like MgCl2?
When sugar dissolves, the molecule does not dissociate.
When 1 mole of MgCl2 dissolves, 3 moles of ions are produced:
MgCl2 (s) → Mg+2 (aq) + 2Cl-1 (aq)

70. dissociation equations…
How many moles of ions are produced when K3N dissolves?
K3N (s) → 3K+1 (aq) + N-3 (aq)
NaCl?
NaCl (s) → Na+1 (aq) + Cl-1 (aq)