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AP Chapter 4 Notes
Solutions & Chemical Reactions
Most reactions in the natural world take place in solution.
a common solvent and it:
•
many substances
•
has a
shape
•
is formed by
bonds
•
and is a
molecule with
Dissolving occurs when
the ionic solid
if the solvent is water).
is
charges
forces between the polar water and
the solid apart. (Also called
.
Solubility is the amount which actually
or
in the case of ionic solids. Solubility varies greatly and depends on:
•
between the ions
•
attractive forces between the
and
Non-ionic substances can dissolve in water if they are
.
Which leads to the phrase “
”
.
.
Nature of Aqueous Solutions: some basic information
Solution is a
mixture whose
vary.
Solute is the substance
in the solvent.
Solvent is the
.
Electrical conductivity is the useful ability of a solution
to
.
Strong Electrolytes are those electrolytes, when dissolved in
water,
.
Weak Electrolytes are electrolytes which
.
Non-Electrolytes
conduct electric current.
Strong Electrolytes are considered to have
dissociation or very
nearly so. We show it by
. We can determine which salts
are strong electrolytes by checking the list of
.
Also strong Arrhenius acids (those that produce
in solution) will
completely dissociate and so will strong Arrhenius bases (those that
produce
in solution).
Weak Electrolytes are:
1.
or
soluble salts which do not completely
ionize in water....
2.
Weak acids (see solubility rules)
3.
Weak bases (see solubility rules)
Factors affecting solubility:
a)
b)
c)
d)
e)
Mrs. Goff expects me
to memorize these!
Reactions often occur when 2 solutions are mixed. We need to know:
•
•
nature of the reactants (solid, liquid, gas, aqueous soln)
the amount of chemicals present or
.
Molarity
is a unit of solution concentration. We use the letter
M=
EX 4.1 (pg 139)
Calculate the molarity of a solution prepared by dissolving 11.5 g of solid
NaOH in enough water to make 1.5 L of solution.
EX 4.2 (pg 140)
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.
1 M of a strong electrolyte like NaCl means we actually have:
EX 4.3 (pg 140)
Give the concentration of each type of ion in the following solutions:
a. 0.50 M Co(NO3)2
b. 1 M Fe(ClO4)3
EX 4.4 (pg 141)
Calculate the number of moles of Cl- ions in 1.75 L of 1.0 x 10-3 M ZnCl2.
EX 4.5 (pg 142)
Typical blood serum is about 0.14 M NaCl. What volume of blood
contains 1.0 mg NaCl
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To make a standard solution:
a) calculate the amount of solute in
grams needed for a desired volume
b) put a small amount of distilled
water in to dissolve
c) fill to the marking on the
volumetric flask with distilled water
A standard solution is a solution in which the concentration
is
. It is used to analyze or calculate the
concentration of an
solution.
EX 4.6 (pg 142)
To analyze the alcohol content of a certain wine, a chemist needs
1.00 L of an aqueous 0.200 M K2Cr2O7 solution. How much
solid K2Cr2O7 must be weighed out to make this solution?
Dilutions occur when you add
We use the little formula
to a concentrated molarity.
to prepare a dilution.
for example: How would you prepare 500 ml of a 0.25 M HCl solution
from a concentrated (12 M) HCl solution?
EX 4.7 (pg 145)
What volume of 16 M sulfuric acid must be used to prepare 1.5 L of a
0.10 M H2SO4 solution?
Driving forces for Chemical Reactions




.
.
.
 Precipitate
A
or precipitate forms as 2 solutions are mixed.
K2CrO4(aq)+Ba(NO3)2(aq)
.
EX 4.8 (pg 151)
Using the solubility rules, predict what will happen when the following
pairs of solutions are mixed.
a.
KNO3 (aq) and SrCl2 (aq)
b.
Na2SO4 (aq) and Pb(NO3)2 (aq)
c.
KOH (aq) and Fe(NO3)3 (aq)
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We show reactions as either
molecular:
ionic:
net ionic:
EX 4.9 (pg 153)
For each of the following reactions, write the molecular equation, the
complete ionic equation, and the net ionic equation.
a) Aqueous potassium chloride is added to aqueous silver nitrate to form
a silver chloride precipitate plus aqueous potassium nitrate.
b) Aqueous potassium hydroxide is mixed with aqueous iron(III) nitrate
to form a precipitate of iron(III) hydroxide and aqueous potassium
nitrate.
EX 4.10 (pg 154)
Calculate the mass of solid NaCl that must be
added to 1.50 L of a 0.100 M AgNO3 solution
to precipitate all the Ag+ ions in the form of
AgCl.
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Remember:
The stoichiometry of precipitate reactions follows the
same steps as we learned already:
1. Write balanced equation 2. Convert to moles
3. Use mole-to mole ratio
4. Answer question
EX 4.11 (pg 156)
When aqueous solutions of Na2SO4 and Pb(NO3)2 are mixed, PbSO4
precipitates. Calculate the mass of PbSO4 formed when 1.25 L of 0.0500
M Pb(NO3)2 and 2.00 L of 0.0250 M Na2SO4 are mixed.
 Acid Base Reactions
We first must know the definition of acids and bases. Actually, there are 3
different definitions:
Arrhenius Acid: a substance that produces
in water
Arrhenius Base: a substance that produces
in water
Bronstead-Lowry Acid: a substance that
Bronstead-Lowry Base: a substance that
Neutralization reactions between acids and bases always form
.
.
.
&
EX 4.12 (pg 158)
What volume of a 0.100 M HCl solution is needed to neutralize 25.0 mL
of 0.350 M NaOH?
EX 4.13 (pg 160)
In a certain experiment, 28.0 mL of 0.250 M HNO3 and 53.0 mL of 0.320
M KOH are mixed. Calculate the amount of water formed in the resulting
reaction. What is the concentration of H+ or OH- ions in excess after the
reaction goes to completion?
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Titrations
Titrations are laboratory procedures that allow us to quantitatively
neutralize acids & bases.
Volumetric analysis: technique for determining
. by titration.
Titration: uses a
to deliver a measured quantity of a solution of
concentration (titrant) into the substance to be analyzed (analyte)
Equivalence Point: the point at which enough
exactly react with the
.
Indicator: a weak organic acid which changes
equivalence point.
End Point: the point when the indicator
has been added to
at or near
.
EX 4.14 (pg 162)
A student carries out an experiment to standardize a sodium hydroxide
solution. To do this, the student weighs out 1.3009 g sample of potassium
hydrogen phthalate (KHC8H4O4 or KHP–molar mass 204.22 g/mol). The
student dissolves the KHP in distilled water, adds phenolphthalein as an
indicator, and titrate the resulting solution with the sodium hydroxide
solution to the phenolphthalein endpoint. The difference between the final
and initial buret readings indicates that 41.20 mL of the sodium hydroxide
solution is required to react exactly with the 1.3009 g KHP. Calculate the
concentration of the sodium hydroxide solution.
EX 4.15 (pg 163)
An environmental chemist analyzed the effluent (waste) from an industrial
process known to produce the compounds carbon tetrachloride (CCl4) and
benzoic acid (HC7H5O2), a weak acid that has one acidic hydrogen atom
per molecule. A sample of the effluent weighing 0.3518 g was shaken
with water, and the resulting aqueous solution required 10.59 mL of
0.1546 M NaOH for neutralization. Calculate the mass percent of
HC7H5O2 in the original sample.
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Oxidation Reduction Reactions
Are reactions in which one or more
are
.
These reactions are also called
reactions. Most
reactions for
are redox such as:
We first have to look at oxidation states and how to assign them.
Oxidation State Rules:
1. with the same element, like
have an oxidation state of zero
since there is no difference in
.
2. with different elements like
have the shared electrons
assigned to the more electronegative element
3. Oxidation number for an atom in its
is zero
4. Oxidation number for a monatomic ion is the same as
.
5. Fluoride in its compounds is always
.
6. Oxygen is generally
in covalent compounds except for:
peroxides like
where it would be
and in OF2 where
oxygen’s Ox # is
.
7. Hydrogen is generally
in non-metallic covalent compounds and
in hydrides like
.
8. In neutral compounds the sum of the oxidation numbers must be
In
polyatomic ions, the sum of the oxidation numbers must be equal to the
on the polyatomic number.
Charges on ions are conventionally written as
Oxidation states are written as
EX 4.16 (pg 167)
Assign oxidation states to all atoms in the following:
a.
CO2
b.
SF6
c.
.
.
NO3-
Occasionally, some really weird oxidation numbers will pop up: Fe3O4
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Now, back to Redox Reactions:
CH4 + 2O2  CO2 + 2H2O
since the oxygen goes from a 0 to a -2, it gains electrons....
and carbon goes from a -4 to a +4, it loses electrons....
A reduction reaction must always be paired with a
 must

Oxidation
oxidation state
 must

.
.
.
Reduction
oxidation state
EX 4.17 (pg 170)
When powdered aluminum metal is mixed with pulverized iodine crystals
and a drop of water is added to help the reaction get started, the resulting
reaction produces a great deal of energy. The mixture bursts into flames,
and a purple smoke of I2 vapor is produced from the excess iodine. The
equation for the reaction is
2Al (s) + 3I2 (s)  2 AlI3 (s)
For this reaction, identify the atoms that are oxidized and reduced, and
specify the oxidizing and reducing agents.
EX 4.18 (pg 170)
Metallurgy, the process of producing a metal from its ore, always involves
oxidation-reduction reactions. In the metallurgy of galena (PbS), the
principle lead-containing ore, the first step is the conversion of lead sulfide
to its oxide (a process called roasting):
2PbS (s) + 3O2 (g)  2PbO (s) + 2SO2 (g)
The oxide is then treated with carbon monoxide to produce the free
metal:
PbO (s) + CO (g)  Pb (s) + CO2 (g)
For each reaction, identify the atoms that are oxidized and reduced, and
specify the oxidizing and reducing agents.
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Now, one way to tackle balancing Redox reactions is by half-reactions
Fe + V2O3  Fe2O3 + VO
1.
Separate into half reactions
2.
Balance all atoms except for Oxygen and Hydrogen
3.
Balance oxygen by adding H2O
4.
Balance hydrogen by adding H+
5.
Balance charges by adding electrons to the more positive side
6.
Multiply one or both reactions so the # electrons lost/gained equal
7.
Add half reactions, cancel identical species
The above is assumed to be in an acidic solution
In a basic solution, the H+ would react with excess OH- to form H2O
2Fe + H2O + V2O3  Fe2O3 + 4H+ + 2VO
EX 4.19 (pg 175)
Potassium dichromate (K2Cr2O7) is a bright orange compound that can be
reduced to a blue-violet solution of Cr3+ ions. Under certain conditions,
K2Cr2O7 reacts with ethyl alcohol (C2H5OH) as follows:
H+ (aq) + Cr2O72- (aq) + C2H5OH (l)  Cr3+ (aq) + CO2 (g) + H2O (l)
Balance this equation using the half-reaction method.
Another method keeps things in tact!
MnO4-+ Fe2+  Fe3++Mn2+
Balance all but oxygen & hydrogen
determine # electrons lost/gained
use multiplying coefficient
find total reactant/product charge
balance charges by adding H+
balance hydrogen by adding H2O
oxygen should be balanced!
EX 4.19 (pg 175) Balance this equation.
H+ (aq) + Cr2O72- (aq) + C2H5OH (l)  Cr3+ (aq) + CO2 (g) + H2O (l)
Assignment Chapter 4:
# 11, 16, 18, 21, 24, 26, 28, 30, 32, 37, 39, 44, 46a, 50b, 52, 53, 58, 61, 63,
66a, 100
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