1. Do Now: Please begin completing the “Error Analysis” section of your Unknown Hydrate laboratory report utilizing the correct empirical formula for the unknown hydrate which is:
AlK(SO4)2 x 12H2O
Objective: Define and discuss the nature of aqueous solutions:
Strong and Weak electrolytes. Discuss Types of chemical
reactions.
HW: Do problems pg. 171 #10, 22, 24 (b and d).
Pg. 172 #36, 44 and 46

2. Types of Chemical Reactions and Solution Stoichiometry

3. Classification of Matter
Solutions are homogeneous mixtures

4. Solute Solvent A solute is the dissolved substance in a solution.
Salt in salt water
Sugar in soda drinks
Carbon dioxide in soda drinks
Solvent
A solvent is the dissolving medium in a solution.
Water in salt water
Water in soda

5. Saturation of Solutions
A solution that contains the maximum amount of solute that may be dissolved under existing conditions is saturated.
A solution that contains less solute than a saturated solution under existing conditions is unsaturated.
A solution that contains more dissolved solute than a saturated solution under the same conditions is supersaturated.
One useful property for characterizing a solution is its
electrical conductivity

6. Definition of Electrolytes and Nonelectrolytes
An electrolyte is:
A substance whose aqueous solution conducts
an electric current.
A nonelectrolyte is:
A substance whose aqueous solution does not
conduct an electric current.
Try to classify the following substances as
electrolytes or nonelectrolytes…

7. Electrolytes?
Pure water
Tap water
Sugar solution
Sodium chloride solution
Hydrochloric acid solution
Lactic acid solution
Ethyl alcohol solution
Pure, solid sodium chloride

8. Answers… ELECTROLYTES: NONELECTROLYTES: Tap water (weak) NaCl solution
HCl solution
Lactate solution (weak)
Pure water
Sugar solution
Ethanol solution
Pure, solid NaCl
But why do some compounds conduct electricity in
solution while others do not…?

9. Ionic CompoundsDissociate
NaCl(s) 
Na+(aq) + Cl-(aq)
AgNO3(s) 
Ag+(aq) + NO3-(aq)
MgCl2(s) 
Mg2+(aq) + 2 Cl-(aq)
Na2SO4(s) 
2 Na+(aq) + SO42-(aq)
AlCl3(s) 
Al3+(aq) + 3 Cl-(aq)

10. Ions tend to stay in solution where they can conduct a current rather than re-forming a solid.
The reason for this is the polar nature of
the water molecule…
Positive ions associate with the negative
end of the water dipole (oxygen).
Negative ions associate with the positive
end of the water dipole (hydrogen).

11. Some covalent compounds IONIZE in solution
Covalent acids form ions in solution, with the
help of the water molecules.
For instance, hydrogen chloride molecules, which are polar, give up their hydrogens to water, forming chloride ions (Cl-) and
hydronium ions (H3O+).

12. Strong acids such as HCl are completely ionized in solution.
Other examples of strong acids include:
Sulfuric acid, H2SO4
Nitric acid, HNO3
Hydriodic acid, HI
Perchloric acid, HClO4

13. Weak acids such as lactic acid usually ionize less than 5% of the time.
Many of these weaker acids
are “organic” acids
that contain a “carboxyl”
group.
The carboxyl group does not easily give up its
hydrogen.

14. Because of the carboxyl group, organic acids are sometimes called “carboxylic acids”.
Other organic acids and their sources include:
Citric acid – citrus fruit
Malic acid – apples
Butyric acid – rancid butter
Amino acids – protein
Nucleic acids – DNA and RNA
Ascorbic acid – Vitamin C
This is an enormous group of compounds; these
are only a few examples.

15. However, most covalent compounds do not ionize at all in solution.
Sugar (sucrose – C12H22O11),
and ethanol (ethyl alcohol – C2H5OH) do not
ionize -
That is why they are nonelectrolytes!

16. Molarity
The concentration of a solution measured in moles of solute per liter of solution.
mol = M L

17. Preparation of Molar Solutions
Problem: How many grams of sodium chloride are needed to prepare 1.50 liters of M NaCl solution?
Step #1: Ask “How Much?” (What volume to prepare?)
Step #2: Ask “How Strong?” (What molarity?)
Step #3: Ask “What is its mass?” (Molar mass is?)
1.500 L
0.500 mol
58.44 g
= 43.8 g
1 L
1 mol

18. Preparation of Molar Solutions
Problem: Calculate the molarity of a solution prepared by
Dissolving 1.56 g of gaseous HCl in enough water to make
26.8 mL of solution.
Try this at your workstations!

19. Serial Dilution
It is not practical to keep solutions of many different concentrations on hand, so chemists prepare more dilute solutions from a more concentrated “stock” solution.
Problem: What volume of stock (11.6 M) hydrochloric acid is needed to prepare 250. mL of 3.0 M HCl solution?
MstockVstock = MdiluteVdilute
(11.6 M)(x Liters) = (3.0 M)(0.250 Liters)
x Liters = (3.0 M)(0.250 Liters)
11.6 M
= L

20. Single Replacement Reactions
A + BX  AX + B
BX + Y  BY + X
Replacement of:
Metals by another metal
Hydrogen in water by a metal
Hydrogen in an acid by a metal
Halogens by more active halogens

21. The Activity Series of the Metals
Lithium
Potassium
Calcium
Sodium
Magnesium
Aluminum
Zinc
Chromium
Iron
Nickel
Lead
Hydrogen
Bismuth
Copper
Mercury
Silver
Platinum
Gold
Metals can replace other metals
provided that they are above the
metal that they are trying to
replace.
Metals above hydrogen can
replace hydrogen in acids.
Metals from sodium upward can
replace hydrogen in water

22. The Activity Series of the Halogens
Fluorine
Chlorine
Bromine
Iodine
Halogens can replace other
halogens in compounds, provided
that they are above the halogen
that they are trying to replace.
2NaCl(s) + F2(g) 
2NaF(s) + Cl2(g)
???
MgCl2(s) + Br2(g) 
No Reaction
???

23. Double Replacement Reactions
The ions of two compounds exchange places in an
aqueous solution to form two new compounds.
AX + BY  AY + BX
One of the compounds formed is usually a
precipitate (an insoluble solid), an insoluble gas that bubbles out of solution, or a molecular compound, usually water.

24. Double replacement forming a precipitate…
Double replacement (ionic) equation
Pb(NO3)2(aq) + 2KI(aq)  PbI2(s) + 2KNO3(aq)
Complete ionic equation shows compounds as aqueous ions
Pb2+(aq) + 2 NO3-(aq) + 2 K+(aq) +2 I-(aq)  PbI2(s) + 2K+(aq) + 2 NO3-(aq)
Net ionic equation eliminates the spectator ions
Pb2+(aq) + 2 I-(aq)  PbI2(s)

25. Solubility Rules – Mostly Soluble
Ion
Solubility
Exceptions
NO3-
Soluble
None
ClO4-
Na+ K+
NH4+
Cl-, I-
Pb2+, Ag+, Hg22+
SO42-
Ca2+, Ba2+, Sr2+, Pb2+, Ag+, Hg2+

26. Solubility Rules – Mostly Insoluble
Ion
Solubility
Exceptions
CO32-
Insoluble
Group IA and NH4+
PO43-
OH-
Group IA and Ca2+, Ba2+, Sr2+
S2-
Groups IA, IIA, and NH4+

27. Oxidation and Reduction (Redox)
Electrons are transferred
Spontaneous redox rxns can transfer energy
Electrons (electricity)
Heat
Non-spontaneous redox rxns can be made to happen with electricity

28. Oxidation and Reduction
An old memory device for oxidation and reduction goes like this…
LEO says GER
Lose Electrons = Oxidation
Gain Electrons = Reduction

29. Oxidation Reduction Reactions (Redox)
Each sodium atom loses one electron:
Each chlorine atom gains one electron:

30. LEO says GER :
Lose Electrons = Oxidation
Sodium is oxidized
Gain Electrons = Reduction
Chlorine is reduced

31. Rules for Assigning Oxidation Numbers Rules 1 & 2
The oxidation number of any uncombined element is zero
2. The oxidation number of a monatomic ion equals its charge

32. Rules for Assigning Oxidation Numbers Rules 3 & 4
3. The oxidation number of oxygen in
compounds is -2
4. The oxidation number of hydrogen in compounds is +1

33. Rules for Assigning Oxidation Number Rule 5
5. The sum of the oxidation numbers in the formula of a compound is 0
2(+1) + (-2) = 0
H O
(+2) + 2(-2) + 2(+1) = 0
Ca O H

34. Rules for Assigning Oxidation Numbers Rule 6
6. The sum of the oxidation numbers in the
formula of a polyatomic ion is equal to
its charge
X + 4(-2) = -2
S O
X + 3(-2) = -1
N O
 X = +5
 X = +6

35. Reducing Agents and Oxidizing Agents
The substance reduced is the oxidizing agent
The substance oxidized is the reducing agent
Sodium is oxidized – it is the reducing agent
Chlorine is reduced – it is the oxidizing agent

36. Trends in Oxidation and Reduction
Active metals:
Lose electrons easily
Are easily oxidized
Are strong reducing agents
Active nonmetals:
Gain electrons easily
Are easily reduced
Are strong oxidizing agents

37. Not All Reactions are Redox Reactions
Reactions in which there has been no change in oxidation number are not redox rxns.
Examples:

38. Balancing Oxidation-Reduction Equations
Half-Reaction Method
separate the full equation into 2 half-reactions: oxidation and reduction
balance each half-reaction separately
add them together to get full balanced equation

39. Steps for acidic solution
Write two separate equations for oxidation and reduction
include any compound containing the element involved
does not need to include everything

40. Steps for acidic solution
For each half-reaction:
balance all elements but H and O
balance O using H2O
balance H using H+
balance charge using electrons (e-)

41. Steps for acidic solution
If needed, multiply one or both half reactions by an integer so that number of electrons is equal in both half reactions

42. Steps for acidic solution
Add the half-reactions together and simplify
Check to be sure all is balanced.

43. Steps for basic solution
Balance using acidic method
Add OH- ions equal to #H+ ions to both sides
Form water on the side of equation that contains both ions
Simplify # water molecules
Check balancing

44. Example

45. Example