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A.P. Chemistry
Chapter 4: Part 2
Types of Chemical Reactions & Solution
Stoichiometry
4.8- 4.10
4.8 Acid-Base Aqueous Reactions (p. 157)
Definition of Acid- substances that ionize in water to produce hydrogen ions
Definition of Base- substances that ionize in water to produce hydroxide ions
Acids
Bases
Sour taste
bitter taste
Cause color changes in plant dyes
feel slippery
React w/ certain metals (Zn, Mg, Fe),
cause color changes in plant dyes
Producing hydrogen gas
aqueous solns conduct electricity
React w/ carbonates & bicarbonates,
Producing carbon dioxide gas
Aqueous solns conduct electricity
Strong Acids
HCl, HBr, HI, HNO3,
acids
H2SO4, HClO4
Strong Bases
NaOH, KOH
Weak Acids
HF, HNO2, H3PO4, HC2H3O2, organic
Weak Bases
NH3
Neutralization Reactions- a reaction between an
acid and a base, generally producing a salt and
water
Acid + base  salt + water
HCl(aq) + NaOH(aq)  NaCl (aq) + H2O(l)
H+(aq) + OH-(aq)  H2O(l)
net ionic
Calculations using Titration Data: M1V1 =
M2V2
M = mol/L Volume in liters (be sure to convert
mL  L) (p. 145-147)
Problem: What volume of 16 M sulfuric acid
must be used to prepare 1.5 L of a 0.10 M
H2SO4 solution?
Titration- a solution of accurately known concentration, called a
standard solution, is added gradually to another solution of
unknown concentration, until the chemical reaction between the
two solutions is complete. The equivalence point or stoichiometric
point is the point at which the acid has completely reacted with or
been neutralized by the base. Indicators are substances that have
distinctly different colors in acid or base media and are used to
determine the equivalence point. The point at which the indicator
changes color is the endpoint.
Example: What volume of a 0.100 M HCl solution is needed to
neutralize 25.0 mL of 0.350 M NaOH?
Example: in a titration experiment, a student finds that 35.18 mL of a
KOH solution are needed to neutralize 0.5468 of potassium acid
phthalate (KHP), KHC8H4O4, a monoprotic acid. What is the molarity
of the KOH?
4.9 Oxidation- Reduction Reactions
Oxidation-Reduction Aqueous Chemistryelectron-transfer reactions (p. 164)
Definition of Oxidation- loss of electrons
Definition of Reduction- gain of electrons
Determining Oxidation Number (p. 167)
free elements have an oxidation state of 0
monatomic ions: oxidation state is equal to the charge
oxygen: is –2 in most compounds; -1 in hydrogen peroxides
or peroxide ion
hydrogen: is +1 , except when it is bonded with a metal in a
binary compound; hydride ion is –1
fluorine: -1 in all compounds
other halogens: -1 when monatomic ions, positive
oxidative states in oxyanions
sum of oxidation numbers in a neutral molecule = 0
sum of oxidation numbers in a polyatomic ion must equal
the charge on the ion
(rare) oxidation numbers need not be integers, for ex.,
Oxygen in O2-1 is –1/2
Example: Assign oxidation states to all atoms in
the following:
a. CO2
b. SF6
c. NO3-1
Writing Half-Reactions- a Half-reaction explicitly shows
the electrons being lost or gained by the different
chemical species in a chemical reaction; the sum of the
half-reactions gives the overall reaction. Electrons are
the product in the oxidation half-reaction; they are the
reactant in a reduction half-reaction. (p. 171)
2Na(s) + Cl2(g)  2NaCl(s)
oxidation: 2 Na  2Na+1 + 2ereduction: Cl2 +
2e-  2Cl-1
Example: Write the half-reactions for the following
reaction:
2Al(s) + 3I2(s)  2AlI3(s)
Different Types of REDOX Reactions (p. 164)
Combination (synthesis), decomposition (analysis), combustion, and single
replacement reactions are all types of Redox reactions. There is also a
reaction type called disproportionation in which an element in one
oxidative state is simultaneously oxidized and reduced.(p. 950)
Double replacement reactions, acid-base neutralization reactions, and
precipitation reactions do not involve changes in oxidative state, and
therefore, will not be redox reactions.
Predicting Displacement Reactions
Example: Which of these reactions will not occur?
Al + CuCl2 
Cu + AlCl3 
Cl2 + NaBr 
Br2 + NaCl 
4.10 Balancing Oxidation- Reduction Equations
Balancing Redox Reactions – In Acid Solution
• Write the ½ reactions
• Balance all elements except for H & O
• Balance O, using H2O
• Balance H, using H+
• Balance charges, using e• Balance # of electrons in the 2 reactions
• Add the two ½ reactions, cancel like terms
• Check that all species are balanced
In Base Solution (p. 176-177)
• Do all of the steps above (for Acid Solution)
• Add OH- to both sides of the equation equal to
the # of H+ (this forms water)
• Cancel H2O that appears on both sides of
equation
• Check that all species are balanced
(NOTE: the example on p. 137-8 of your textbook is
the redox reaction in the Redox Titration Lab and
was a question on the 2007 AP Chemistry exam.)
Redox Titration- when a solution containing an
oxidizing agent is titrated with a solution
containing a reducing agent. The equivalence
point is reached when the reducing agent is
completely oxidized by the oxidizing agent. Two
common oxidizing agents are KMnO4 and
K2Cr2O7; the colors of the reduced species is
distinctly different from the polyatomic ions.
Example: A 16.42 mL volume of 0.1327 M KMnO4
solution is needed to oxidize 25.00 mL of a FeSO4
solution in an acidic medium. Determine the
molarity of the FeSO4 solution.