Ch 4: Types of Chemical
Reactions and Solutions
Stoichiometry
Do Now:
1.
True or false:
a.
b.
2.
Most ionic substances need an aqueous medium in
order to react.
Reactions occur at the level of atoms and molecules,
therefore count of each reactant atoms/molecules/ions
needs to be known in order to determine how much
product can be produced.
If the above statements are true
a.
b.
How can you determine the amount of particles
involved in a reaction if you have 28 mL of NaCl(aq)
If you know that 50 g of CaS is dissolved in 250 mL of
water, how would you determine the volume to use if
you needed 0.21 mol of CaS for a reaction?
4.1 Water, the Common
Solvent
A. Structure of water
1. Oxygen's electronegativity is high
(3.5) and hydrogen's is low (2.1)
2. Water is a bent molecule (shape)
3. Water is a polar molecule (charge differential)
B. Hydration of Ionic Solutes
1. Positive ions attracted to the oxygen end of
water
2. Negative ions attracted to the hydrogen end of
water
3. Particle dissociates (or separate) into ions.
a.
b.
c.
Level of dissociation depends on Ksp (solubility
product constant) value.
HIGH (+) Ksp value (soluble) Salt, CuCl2, Ba(NO3)2
LOW Ksp value (insoluble) BaSO4, Ca(OH)2
C. Hydration of Polar Solute Molecules
(Covalent)
1. Negative end of polar solute molecules are
attracted to water's hydrogen
2. Positive end of polar solute molecules are
attracted to water's oxygen
3. Particle stays together. Ex: sugar, alcohol
D. "Like Dissolves Like"
1. Polar and ionic compounds dissolve in
polar solvents like water
2. Nonpolar compounds like fats dissolve in
nonpolar solvents like acetone
NOTICE: DAILY QUIZ

From Today Until the Time that Your
Class Learns How to Behave and Be
Considerate Towards Each Other,
You Will Take a Daily Quiz
◦ Starting at the Bell
◦ Timed at 4-6 minutes
◦ Any talking or commenting during quiz
will earn you an automatic ZERO
◦ No time extensions
◦ DO NOT START UNTIL TOLD TO DO SO.
4.2 The Nature of Aqueous Solutions:
Strong and Weak Electrolytes
A. Definition of Electrolytes
1. A substance that when dissolved in
water produces a solution that can
conduct an electric current
B. Strong electrolytes conduct current
very efficiently
1. Completely ionized when dissolved in
water
a. Soluble Ionic compounds
b. Strong acids (HNO3(aq), H2SO4(aq), HCl(aq))
c. Strong bases (KOH, NaOH)
C. Weak electrolytes conduct only a
small current
1. Slightly ionized in solution
a. Weak acids (organic acids - acetic, citric,
butyric, malic)
b. Weak bases (ammonia)
D. Nonelectrolytes conduct no
current
1. No ions present in solution
a. alcohols, sugars
4.3 The Composition of
Solutions
A. Molarity
1. Moles of solute per liter of solution
M = molarity = moles of solute/liters
of solution
2. Steps
a.
b.
c.
Convert Mass/volume/particles into moles
Convert volume (may be based on mass and density) of
solution into Liters.
Divide by Liter of solution (NOT volume of water used)
Example:
1. Calculate the molarity of a solution prepared by
dissolving 11.5 g of solid NaOH in enough water to
make a 1.50L of solution
2. Calculate the molarity of a solution prepared by
dissolving 0.959 L of gaseous HCl in enough water
to make 26.8 mL of solution.
B. Concentration of Ions in Solution
1. Ionic compounds dissociate in solution
2. multiplying the molarity by the number
of ions present indicates concentration
of ions**
a. **Individuals IONS react, not complete
formula units.
Example
1. Give the concentration of each type
of ion in the following examples
a. 0.50 M Co(NO3)2
b. 1 M Fe(ClO4)3
C. Moles from Concentration
1. Liters of solution x molarity = moles
of solute
Example:
1. Calculate the number of moles of
Cl- ions in 1.75 L of 1.0 X 10-3 M
ZnCl2
2. Typical blood serum is about
0.14 M NaCl. What volume of
blood contains 1.0 mg NaCl
D. Solutions of Known Concentration
1.
Standard solution - a solution whose concentration is accurately
known
Preparation of Standard solutions
2.
How much x How strong x What does it weigh?
L x mol/L x g/mol = grams required to prepare the standard
3.
a.
b.
c.
d.
e.
f.
g.
h.
Volumetric flask-used to prepare single, specific volume of
known molarity.
Determine mass of solute
Mix with small amount of distilled water in beaker (or directly into
volumetric flask **difficult)
Transfer to volumetric flask (funnel)
Rinse beaker multiple times with small amounts of distilled water
and transfer to volumetric flask (ensures all particles of solute are
transferred. (skip if solute added to flask directly)
Rinse funnel into volumetric flask
Add sufficient water to rise ALMOST to level of calibration mark
(meniscus).
Allow any water/solution clinging to neck to settle down.
Add last several mL of distilled water dropwise until calibration
mark (meniscus).
Example:
1. To analyze the alcohol content of a certain wine, a
chemist needs 1.00L of aqueous 0.200 M K2Cr2O7
solution. How much solid must be weighed out to make
this solution?
E. Dilution
1. Dilution of a volume of solution with
water does not change the number of
moles present
2. Solving dilution problems
M1V1 = M2V2
Example:
1. What volume of 16M sulfuric acid
must be used to prepare 1.5 L of a
0.10 M H2SO4 solution?
Homework
Page 180-181
Questions (9, 12, 15, 17,)
(19, 21, 23 a, b, 25 )
DO NOW
Take out HW, then do the
practice problems below.
Write the dissociation equation for
the dissolution of Ammonium
phosphate in water.
 If 543.0 g of the above compound
is dissolved to form 345 mL of
solution

◦ determine the molarity of the
solution.
◦ Determine the concentration of each
type of ion dissolved in solution.
Topic: Types of Chemical
Reactions
AIM: To differentiate, predict
products and solve stoichiometry
problems for three different
categories of reactions.
Consider a 1.00L sample each of two solutions of 1.00 M Potassium
chromate and 2.00 M Barium nitrate.
1. Write the dissociation equation for each.
2. Draw a particulate diagram for each solution (separate)
3. Write a balanced chemical equation for the reaction
between Potassium chromate and Barium nitrate.
4. Which, if any, is the limiting reactant? Justify
5. Draw a particulate diagram for the final reaction mixture.
HW: Textbook Page 182
# 29, 31, 33, 37, 40, 42, 43
4.4: Type of Chemical
Reactions (Rxn)
Millions of chemical reactions
 Can be categorized into 3 primary
groups

◦ Precipitation reactions
◦ Acid-Base reactions
◦ Oxidation-Reduction (Redox)
reactions
Figure 4.14: Reactant Solutions:
(a) Ba(NO3)2(aq) and (b) K2CrO4(aq)
4.5: Precipitation Reactions
A. Precipitation Reactions
1. A reaction in which a solid (precipitate) is formed.
B. Process
1.
a.
Dissociation
Ionic compounds dissolve in water and the ions separate
and move independently
Ex 1: AgNO3(aq) + NaCl(aq)  products
2. Determination of Products
a.
b.
c.
Recombination of ions
Elimination of reactants as products
Identifying the precipitate
a.
b.
Dependent on ability to separate products successfully
If there are NO insoluble products, reaction has NOT
taken place.
If both products are insoluble, reaction has NOT taken
place.
3. Determine whether reaction takes place (solubility
rules).
c.
3a. Solubility Rules
Practice Set 1: Justify whether
each of these reactions will occur.
1.
NaCl(aq) + KNO3(aq) 
2.
Na2SO4
3.
KNO3 + BaCl2 
4.
KOH + Fe(NO3)3 
5.
Ag2SO4 +
6.
KCl + AgNO3 
+
Pb(NO3)2 
CaCl2
4.6 Describing Reactions in
Solution
A.
The Molecular Equation
1. Gives the overall reaction stoichiometry, not
necessarily the actual forms of reactants and
products in solution.
Example: Na2(CO3) + Ca(NO3)2  ????
i.
B.
C.
The Complete Ionic Equation
1. Represents as ions all reactants and products that
are strong electrolytes
The Net Ionic Equation ***
1. Includes only those components that take part in
the chemical change
2. Does not indicate spectator ions.
i.
D.
Don’t EVER give this as your answer.
***ONLY valid form of representing reaction on AP exam.
ALL species MUST be written AS their
prescribed behavior in water—ions or formula
units/molecules.
Practice: Write the COMPLETE
and the NET ionic reaction
1.
2.
Aqueous potassium chloride is
added to aqueous silver nitrate to
form a silver chloride precipitate
and aqueous potassium nitrate
Aqueous potassium hydroxide is
mixed with aqueous iron (III)
nitrate to form a precipitate of
iron (III) hydroxide and aqueous
potassium nitrate.
3. Lithium iodide solution is mixed
with ammonium sulfate solution.
4. Calcium carbonate and Barium
hydroxide are added to water.
4.7: Stoichiometry of
Precipitation Reactions
Practice
1.
Calculate the mass of PbSO4(s)
formed when 1.25 L of 0.0500 M
Pb(NO3)2 and 2.00 L of 0.0250 M
Na2SO4 are mixed.
a) Based on question 2, determine the
moles of excess reactant left over.
b) Determine the moles of each ion in
solution AFTER the reaction is
complete.
c) Determine the concentration of each
ion at the end of the reaction.

Strategy
◦ Is this a limiting reactant problem? How
do you know?
◦ What conversions are required?
 Do we know how to make the conversion?
◦ Will ALL Limiting Reactant ions be used
up? Why or why not?
◦ What portion of Excess Reactant ions will
be left over?
 Both ions in equal concentrations? Why or why
not?
◦ What does concentration (Molarity)
depend on?
Practice
2. Determine the mass of the
precipitate formed with 50. mL of
1.75 M Na2(CO3) is mixed with 75 mL
of 2.5 M Ca(NO3)2
a. Based on question 2, determine
the moles of excess reactant left
over.
b. Determine the moles of each ion
in solution AFTER the reaction is
complete.
c. Determine the concentration of
each ion at the end of the reaction.
4.8 Acid-Base Reactions
AIM:
1. To understand the different concepts of an
ACID vs. BASE
2. To understand the net ionic reaction for acidbase reactions
3. To learn how to titrate to achieve
neutralization reaction.
DO NOW:
1. How do you recognize an acid from formula?
A base from formula?
2. Identify the products of the following
reaction. Determine whether reaction will
take place.
Sulfuric acid reacts with Sodium hydroxide
A. Defining Acids vs. Bases
2 Definitions
1. Arrhenius—
a. Acids: substance that produces H+ ions
in solution
i.
Strong acids: virtually every molecule ionizes
-Example: Hydrochloric, Nitric, Sulfuric,
Perchloric Acid
b. Base: substance that produces OH- ions
in solution
i.
2.
Strong bases: Sodium hydroxide, Lithium
hydroxide
Bronsted-Lowry:
a. Acid: substance that is a proton donor
b. Base: substance that is a proton acceptor
(discuss in detail in ch. 14)
List of Strong acids and Bases

Strong acids—completely
ionized in water to give one
or more H+ (protons) per
acid molecule
◦
◦
◦
◦
◦
◦
HI
HBr
HClO4
HCl
HClO3
H2SO4
 H+(aq) + HSO4-(aq) (HSO4- is a
weak acid that contributes
additional protons)
◦ HNO3
◦ Everything else is a WEAK acid
(doesn’t dissociate fully)

Strong Bases—
completely ionized in
water to give one or
more OH- ions per base
molecule
NaOH
KOH
LiOH
RbOH
CsOH
Group 2 hydroxides—not
fully soluble, but forms
strong bases.
◦ Ca(OH)2
◦ Ba(OH)2
◦ Sr(OH)2
◦
◦
◦
◦
◦
◦
BONUS QUESTION: Notice that Mg(OH)2 is missing from the list of
strong bases. Yet, Mg(OH)2 is the active ingredient of many antacids?
Is this a good choice for an antacid? Why or why not? How is it able to
neutralize stomach acid (HCl)?
B. Acid-Base Reactions
Net ionic equation for STRONG acid-STRONG base reactions
1.
2.
H+(aq) + OH-(aq)  H2O(l)
The hydroxide ion can be assumed to completely react
with even a weak acid in solution
Ex: OH- (aq) + HC2H3O2(aq)
 H2O(l) + C2H3O2-(aq)
a) Notice Acetic Acid shown as molecule, not ions. Why?
i.
b)
Weak Acids and Bases must be written as molecular formulas
because they DO NOT dissociate fully. **write species as their
prescribed behavior in water.
Also true for weak base with hydrogen ion.
i.
Practice: write neutralization reaction between ammonium
hydroxide and hydrochloric acid.
Often called NEUTRALIZATION reaction
3.
a)
Why?
b)
But may result in a NON-neutral final reaction mixture.
Why?
Practice
Determine whether the following
reactions are Acid-Base
neutralization? If yes, predict
product.
1. HCl
+ NaOH
2. HNO3
+ CaSO4
3. Ca(OH)2
+ C6H12O6
4. H3PO4
+ KOH
C. Stoichiometry Calculations
for Acid-Base Reactions
1. List the species present in the combined solution
before any reaction occurs; decide what reaction will
occur
2. Write the balanced net ionic equation for this
reaction
3. Calculate the moles or reactants
a. For reactions in solution, use volumes of the
original solutions and their molarities
4. Determine the limiting reactant where appropriate
5. Calculate the moles of the required reactant or
product
6. Convert to grams or volume of solution as required
Practice Problem 1
What volume of a 0.100 M HCl
solution is needed to neutralize
25.0 mL of 0.350 M NaOH
What volume of a 0.100 M HCl
solution is needed to neutralize
25.0 mL of 0.350 M NaOH
 Strategy?

◦
◦
◦
◦
Determine
Determine
Determine
Determine
Molarity
moles of OH- present.
moles of H+ needed.
H+ to Acid ratio
volume based on given
Practice Problem 2
In a certain experiment, 28.0 mL of a 0.250 M
HNO3 and 53.0 mL of 0.320 M KOH are
mixed.
a. Calculate the volume of water formed in
the resulting reaction.
b. What is the concentration of H+ or OH- ions
in excess after the reaction goes to
completion.
c. Is the resulting solution acidic, basic or
neutral. Explain.
d. What is the molarity of the solution in
terms of the
i.
ii.
resulting salt?
Excess H+ or OH- ions?
In a certain experiment,
28.0 mL of a 0.250 M
HNO3 and 53.0 mL of
0.320 M KOH are
mixed.
a.
Calculate the volume
of water formed in
the resulting reaction.
b. What is the moles of
H+ or OH- ions in
excess after the
reaction goes to
completion.
c.
Is the resulting
solution acidic, basic
or neutral. Explain.
d. What is the molarity
of the solution in
terms of the
i.
ii.
resulting salt?
Excess H+ or OH- ions?
Strategy
 Is it a Limiting Reactant
problem?
 Determine moles of one
reactant needed with respect
to the other.
 Determine moles water
formed
◦ Determine mass, then volume
(DH2O = 1.0 g/mL)
◦ Not a gas, cannot use 22.4
L/mol

Determine excess ion
amount used and remaining.
◦ Determine acid/base/neutral
depending on excess ion
remaining

Determine molarity of salt
formed and H+ or OH- ion.
Don’t forget to ADD the
volume of water formed
from reaction to total volume
from aqueous solutions.
◦ Is the volume of water formed
significant?
D. Acid-Base Titrations
1. Vocabulary
a. Titrant - Solution of known concentration
b. Analyte - Solution of unknown
concentration
c. Equivalence point - Point at which the
amount of titrant added to analyte results
in perfect neutralization
d. Indicator - a substance added at the
beginning of the titration that changes
color at the equivalence point
e. Endpoint - the point at which the
indicator changes color
2. Requirements for a successful
titration
a. the exact reaction between titrant and
analyte must be known
b. the reaction must proceed rapidly
c. the equivalence point must be marked
accurately (select the appropriate
indicator)
d. the volume of titrant required to reach
the equivalence point must be known
accurately
e. for acid-base titrations, the titrant
should be a strong acid or a strong base
Video 1
Video 2
4.9-4.10: Redox Reactions
AIM:
1. To differentiate between reduction and
oxidation of elements
2. To learn how to balance Redox reactions.
DO NOW:
1. Identify the oxidation state of oxygen,
Aluminum, Carbon, Hydrogen and
Nitrogen.
---in elemental form?
---in compound form?
4.9: Oxidation-Reduction
Reactions
1.
Reactions in which one or more electrons
are transferred are known as oxidationreduction or REDOX reactions
a. Example:
2Na(S) + Cl2(g)  2NaCl(s)
b. Most reactions used for ENERGY
PRODUCTION (photosynthesis, combustion,
cellular respiration) are redox reactions
c. Apparent when products or reactants are
ionic OR both elements and compounds as
reactants / products
d. Covalent compounds can also be
represented by assumed or actual electron
transfers.
2. Oxidation States
AKA oxidation numbers
Helps keep track of electrons in redox
reactions ** especially with covalent
compounds
c. History/Rules
a.
b.
i.
in covalent compounds, oxidation numbers
for elements are based on arbitrary
assignment given to certain elements.
ii. In partial sharing of electrons, element with
stronger affinity is assumed to have all the
electrons.
Ionic compounds: oxidation number is
equal to charge of ion.
e. In electrically neutral compounds, sum of
oxidation numbers must equal ZERO. In
ions, the sum must equal the charge on
the ion.
d.
Table 4.2: Rules for Assigning
Oxidation States ** MEMORIZE
3. Examples of oxidation
numbers
Determine oxidation number for
each atom in the following
1. CO2
2. SF6
3. Fe3O4
4. NO35. PO4-3
6. NH4+1
4. Characteristics of Redox
Reactions
a.
b.
c.
Characterized by TRANSFER of
electrons
Ionic compounds—literal transfer
Example:
2Na(s) + Cl2(g)  NaCl(s)
Covalent compounds—presumed
transfer
Example:
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)
5. Oxidation VS. Reduction
a. Oxidation: an increase in the
oxidation state (a loss of electrons)
Example: 2Na  2Na+1
b. Reduction: A decrease in oxidation
state (a gain of electrons)
Example Cl2  2Cl-1
Practice:
Determine oxidation states of each
element in the following reactions.
Determine which elements are
oxidized and which are reduced
1. 2Al(s) + 3I2(s)  2AlI3(s)
2. 2PbS(s) + 3O2(g)  2PbO(s) + 2SO2(g)
3. PbO(s) + CO(g)  Pb(s) + CO2(g)
4.10: Balancing REDOX
Reactions

Useful to separate into 2 half
reactions
◦ For oxidation and reduction
◦ Ex: Ce4+ + Sn2+  Ce3+ + Sn4+


Balance each half reaction separately
Equalize number of electrons
transferred.
◦ **differs for acidic vs. basic media

Add up half reactions to obtain full
balanced reaction
2
Acidic Media
+
2
+
Add OH- to both sides of equation
(equal to H+)
Form H2O on the side containing
H+ and OH- ions
Eliminate number of H2O
appearing on both sides
Check that elements and
charges are balanced
BASIC MEDIA