Name:____________________________________________
Objectives:
1.
Students will understand the meaning of molarity, how to prepare a solution of a known molarity, and how to calculate the molarity of a solution.
2.
Students will become familiar with the electrical properties of aqueous solutions and the effects of strong electrolytes and nonelectrolytes.
3.
Based on molarities, students will be able to state the concentrations of all ions present in aqueous solutions.
4.
Students will be able to express the following interactions in equation form: ionization/dissociation of strong acids and bases, and neutralization.
5.
Students will be able to consult solubility tables to predict the formation of precipitates and to represent those reactions in net ionic form.
6.
Students will be able to identify the substances oxidized and reduced and to identify the oxidizing and reducing agents in a redox reaction.
7.
Students will become familiar with practical applications of redox reactions.
• Solution – _________________________________ of two or more substances.
• Solute – the substance being __________________________
• Solvent – the substance ____________________________________________
• Concentration of a solution – the quantity of a ____________________in a given quantity of
______________________
– A concentrated solution contains a relatively ______________amount of
______________ vs. the ____________________
– A dilute solution contains a relatively ____________concentration of
________________________________________________
– “Concentrated” and “dilute” aren’t very _____________________
 Molarity (M), or molar concentration – the amount of solute, in
_____________per _______________of __________________________.
Formula: Molarity = moles of solute
liters of solution

• A solution that is 0.35 M sucrose contains _______________________________ in each ________________________________.
• Keep in mind that molarity signifies moles of solute per liter of ______________ , not liters of _________________.
Formula: Molarity = moles of solute
liters of solution
Example #1: What is the molarity of a solution prepared by dissolving 5.0 mol of NaCl in enough water to make 2 L of solution?
Example #2: What is the molarity of a solution prepared by dissolving 50.75 g of AgNO
3
in enough water to make 1500 mL of solution?
Example #3: How many grams of sodium carbonate are needed to prepare 0.250 L of a 0.300 M solution?
2

Arrhenius’s theory:
• Certain substances _______________________ into __________________ and
__________________when dissolved in water.
– These ions allow
_____________________________ to flow.
• A strong electrolyte _______________________________________.
– A strong electrolyte is present in solution almost exclusively as ________.
– Strong electrolyte solutions are __________conductors.
 Example: __________________________________________________
• A weak electrolyte _________________________________________.
– Weak electrolyte solutions are ______________ conductors.
– Different weak electrolytes dissociate to different extents.
 Example: ______________________________________________
• A nonelectrolyte ____________________________________________.
– Nonelectrolytes are present in solution as almost all _________________.
– Nonelectrolyte solutions __________________ conduct electricity.
 Example: ___________________________________________
3

a.
Dissolution (Dissociation) Equations for Strong Electrolytes
Write cations and anions as individual particles with the appropriate amount of each.
NaCl (aq) 
MgSO
4
(aq) 
Na
2
SO
4
(aq)  b. Calculating Ion Concentrations in Solution
Multiply the molarity (M) by the coefficient of each ion in the dissolution equation to find the ion concentration!
Na
2
SO
4
(aq) 
Example 4.2 from pg. 77 _
Potassium dichromate, K
2
Cr
2
O
7
, is used in the tanning of leather. A flask containing 125 mL of solution is labeled 0.145 M K
2
Cr
2
O
7
. a) What is the molarity of each ion in solution? b) A sample containing 0.200 moles of K + is added to the solution. Assuming no volume change, what is the molarity of the new solution?
4

5
1. Precipitate –
2. Simple solubility rules – see handout! (Or look at pg. 13, 14 in reference book) s = ____________________ -- _________________________________ ss= i = __________________________________________ --
_______________________________________________________
Example #1: Are the following compounds soluble or insoluble?
NaNO
3
Ba(OH)
2
NaSO
4
CaSO
4
Example #2: Using the solubility rules and your knowledge of double replacement reactions, predict the products of the following reactions and then indicate if the compounds are soluble or insoluble.
A) Cu(NO
3
)
2
( ) + (NH
4
)
2
SO
4
( ) 
B) FeCl
3
( ) + AgNO
3
( ) 
 Net Ionic Equations –
 Spectator Ions –
***see Net Ionic Equation handout for further explanation!!

6
Writing Net Ionic Equations:
1.
Write a balanced reaction based on the reactants given.
2.
Determine which compounds are soluble and which are insoluble in water and indicate this using (s) or (aq).
3.
For the soluble compounds, replace the formulas with the ions each compound would dissociate into. a.
The insoluble compounds will stay together!
4.
Cancel out spectator ions (ions that appear the same on both sides of the reaction).
5.
Rewrite the equation with the ions left over to yield your net ionic equation.
Example: Predict the products of the following reactions to decide if a precipitate will form. Indicate which compounds would be soluble and which would be insoluble. Finally, write the full ionic and net ionic equation. a. Ba(NO
3
)
2
and Na
2
CO
3
Balanced Reaction:
Full Ionic Equation:
Net Ionic Equation: b. MgCl
2
and AgNO
3
Balanced Reaction:
Full Ionic Equation:
Net Ionic Equation:
END OF MATERIAL FOR PART 1 TEST

A. Properties of Acids and Bases
Acids
 Taste:_____________
 React with indicators: turn blue litmus _______; turn phenolphthalein ______
 React with certain metals to form ______________
 Corrosive to metals and ______________________________
 Neutralize bases to form_________________________________________
Bases
 Taste:_____________
 Feel: ___________________
 React with indicators: turn red litmus _______; turn phenolphthalein _______
 Corrosive to _____________________________
 Neutralize acids to form_____________________________________
B. Arrhenius Definitions of Acids and Bases
• Acids produce ___________in water solution
• Bases produce ___________in water solution
Examples: HCl (aq) 
NaOH (aq) 
 Arrhenius definitions are limited!!!
 Definition: The Hydronium Ion:
In aqueous solution, H + does not exist! Note: In all problems, [H + ] = [H
3
O + ]
H + + H
2
O  ______________(hydronium ion)
7

8
C. Strong Acids/Bases
1. Strong Acids: THERE ARE ONLY SIX!
Strong acids are ______________________in solution!
(Use single arrow; no equilibrium established)
HNO
3

2. Strong Bases: Group I metals + OH
Some examples:
–
Some Group II metals + OH –
Strong bases are ______________________ in solution!
(Use single arrow; no equilibrium established)
Sr(OH)
2


D. Weak Acids/Bases
1.
Weak acids are __________________________________ in water. a.
Use double arrow; equilibrium established
 Weak organic acids contain the –COOH group
 One of the best known is acetic acid = CH
3
COOH
2.
Weak bases are __________________________________ in water a.
Use double arrow; equilibrium established
• One of the best known weak bases is ammonia (NH
3
)
• NH
3
(aq) + H
2
O(liq) 
E. Bronsted-Lowry Theory
The most general theory for determining acids and bases in aqueous solution is the BRØNSTED -
LOWRY theory of acids and bases
DEFINITIONS:
• ACIDS ________________ H + IONS (proton donor)
• BASES ________________ H + IONS (proton acceptor)
The Brønsted-Lowry definition means NH
3 is a ______________ in water – and water is itself an
________________________
NH
3
+ H
2
O ↔ NH
4
+ + OH -
F. Conjugate Pairs
 NH
3
/ NH
4
+ is a________________________________ o Related by the gain or loss of ___________.
9

Every acid has a conjugate__________, formed when H + is _______________from the acid.
Every base has a conjugate___________, formed when H + is _______________ to the base.
10
Example #1: Write the Bronsted-Lowry equations for the weak acid HNO
2
and the weak base NH
3
, identifying the conjugate acid-base pairs in each equilibrium.
Example #2: Write the Bronsted-Lowry equations for the weak acid HCO
3
and the weak base NH
3
, identifying the conjugate acid-base pairs in each equilibrium.
G. Amphiprotic (amphoteric) species
 Some substances can function as both an __________OR a __________ depending on what they are reacted with. o Can ______________ OR _______________H + o They are called amphiprotic or amphoteric
E.g. H
2
O + H
2
O 

 Acids fall into 2 categories based on their formulas
1.
Binary acids – H + with one other element (HCl, HF, H
2
S, etc.)
2.
H + with a polyatomic ion
Naming Binary Acids
“hydro____________ic acid”
Ex.) HCl = hydrochloric acid
***In the blank, put the root of the non-metal element’s name
[chlorine  “chlor”]
HI =
Naming Acids w/ Polyatomic Acids
 Change the ending of the polyatomic ion as shown below: o “-ate” ending  “-ic” ending o “-ite” ending  “-ous” ending
 Add the word “acid” after the name of the polyatomic ion with the changed ending
Examples:
HNO
3
–
HNO
2
–
Mixed Naming Acids Practice: Name the acids based on the appropriate set of rules:
H
2
CO
3
HCl
HBr HF
H
3
PO
HClO
2
4
H
2
CH
SO
3
3
COOH
11

12
A common way to express acidity & basicity is with pH (the “____________” of hydrogen). pH =
***pH calculations look at the pH, pOH, concentration of H + ion in solution [H + ], and concentration of
OH in solution [OH ]
A. The pH Scale
Basic solution
Neutral
Acidic solution
B. Ion Concentration Scale
Neutral solution:
[H+] [OH-] =
Acid solution:
[H+] [OH-] [H+] >
Basic solution:
[H+] [OH-] [OH-] >
B
B

C. pH Relationships
 pH =
 pOH =


Water Dissociation Constant (Auto-ionization Constant of Water)
1. For any sample of water molecules:
H
2
O (liquid) + H
2
O (liquid) 
2. K w
=
13

Example #1: A sample of tap water has a [H + ] = 2.8 x 10 -6 M. What is the [OH ]?
Example #2: Calculate the pH, pOH and [OH ] of an acid solution whose [H + ] is 1.8 x 10 -4 M.
Example #3: If the pH of Coke is 3.12, what is the concentration of the hydrogen ion in solution?
Also, is the solution acidic or basic?
D. pH of strong acids and strong bases
 Strong acids dissociate completely in aqueous solution:
E.g. HCl 
 For a strong acid, [H + ] can be calculated from the ___________________ of the acid
Example #4:
Calculate the [H + ], pH, pOH, and [OH ] of a 0.15 M solution of the strong acid, HNO
3
.
14

 Strong bases dissociate completely in aqueous solution:
E.g. Sr(OH)
2

 For a strong base, [OH ] can be calculated from the ________________ of the base
Example #5: Calculate the pH, pOH, [H + ], and [OH ] of a solution made by dissolving 0.0500 moles of
Ba(OH)
2
- a strong base – in 5.00 L of water.
 Neutralization reactions –
• The products of this neutralization are a _________ (_______________________) and
_______________
• It is a __________________________________reaction.
Example:
Titration Definitions
 A ______________________is a carefully controlled neutralization reaction. o Titrations are used to determine the concentration of an unknown acid or base.
 A ______________is the piece of glassware used in a titration.
 The ________________is the substance of ____________concentration used to determine the ___________________concentration of the other substance.
15

16
 ___________________ - substance that changes color at a certain ________ -- is added to tell us when the neutralization is complete
 Example: Phenolphthalein undergoes a color change between pH 8 and 10
 The _______________________________is the point in the titration where the neutralization is complete: o _________________ = ________________
 The _______________________is the point where the ____________________ changes color. o If the indicator is chosen correctly, these two points are identical!
Strong Acid and Base Titration: Net Ionic Equation
Acid + Base  Salt + Water
HCl + NaOH  H
2
O + NaCl
Full Ionic Equation:
Net Ionic Equation:
Strong Acid and Base Titration: Short-cut Equation
***At the equivalence point (end point):

17
Example#1:
HCl is titrated with NaOH.
1. Write the complete balanced neutralization equation and the net ionic equation.
2. How many milliliters of 0.0195 M HCl are required to titrate 25.00 mL of 0.0365 M NaOH?
Example #2:
Ba(OH)
2
is titrated with HCl.
1. Write the complete balanced neutralization equation and the net ionic equation.
2. How many milliliters of 0.0108 M Ba(OH)
2
are needed to titrate 25.00 mL of 0.0213 M HCl?
END OF MATERIAL FOR PART 2 TEST
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1.
Definitions:
 Oxidation –
 Reduction –
 Oxidation number –
 Oxidation-reduction (redox) reaction –
 Oxidizing agent –
 Reducing agent –
Example: Mg + Cu 2+  Cu + Mg 2+
2. Rules for determining oxidation numbers a. (Neutral species) b. Monatomic ions (groups 1, 2, 17) c. (Oxygen) d. (Hydrogen)
18

e. Sum of oxidation numbers for neutral and charged species:
Example: CaCl
2
Example #1: Assign oxidation numbers to the elements in the following species:
Ca C
2
O
4
Cr
2
O
7
2-
N
2
O N
2
O
4
ClO 1ClO
4
1-
3. Identifying Redox Reactions – look for changes in oxidation numbers!!
Al (s) + Fe
2
O
3
(s)  2 Fe (liq) + Al
2
O
3
Who gets oxidized?
Who gets reduced?
19

20
Example #3: Identify the element reduced, the element oxidized, the reducing agent, and the oxidizing agent:
A) Fe 2+
(aq)
+ Cr
2
O
7
2-
(aq)
+ H +
(aq)
 Fe 3+
(aq)
+ Cr 3+
(aq)
+ H
2
O
(l)
Element reduced:
Element oxidized:
Reducing agent:
Oxidizing agent:
B) 3 Cl
2 (g)
+ 2 Cr(OH)
3 (aq)
+ 10 OH -  2CrO
4
-
(aq)
+ 6Cl -
(aq)
+ 8H
2
O
(l)
Element reduced:
Element oxidized:
Reducing agent:
Oxidizing agent:
END OF MATERIAL FOR PART 3 TEST
(END OF UNIT 4)
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