Transcript Class Notes

Reactions in Aqueous Solution
Chapter 4
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Describing Chemical Reactions
 A chemical reaction is the process by which one or more substances
are changed into one or more different substances
reactants
products
(NH4)2Cr2O7(s)  N2(g) + Cr2O3(s) + 4H2O(g)
“The reactant ammonium dichromate yields the products nitrogen, chromium (III) oxide and water”
 A CHEMICAL EQUATION represents, with symbols and
formulas, the identifies and relative amounts of the reactants and
products in a chemical equation
2
Indications of a Chemical Reaction
1.
Evolution of heat and light is strong evidence that a chemical
reaction has taken place! But, the evolution of heat or light by itself is not
necessarily a sign of a chemical change since many physical changes also release
either heat or light.
Production of gas! (aka bubbles when two substances are mixed)
3. Formation of precipitate! A solid that is produced as a result of a chemical
2.
reaction in solution and that separates from the solution is known as a precipitate
4.
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Color Change!
Characteristics of Chemical Equations
The equation must represent all reactants and products.
2. The equation must contain the correct formulas for the reactants
and products
3. The law of conservation of mass MUST be satisfied!!
1.
Law of conservation of mass – atoms are neither created nor
destroyed in ordinary chemical reactions
To equalized numbers of atoms, coefficients are added in front of the
formulas where necessary
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Types of Chemical Reactions
A+B↔C
C ↔A + B
A + BC ↔ AC + B
AB + CD ↔ AD + CB
 Synthesis (Combination)
 Decomposition
 Single Replacement
 Precipitation Reactions (Double
Replacement Reactions)
Acid + Base ↔ salt + water  Neutralization Reactions (Acid/Base)
Change of oxidation state  Redox Reactions
Hydrocarbon + O2 ↔ CO2 + H2O  Combustion
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PRECIPITATION REACTIONS
Double Replacement Reactions
AB + CD ↔ AD + BC
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A solution is a homogenous mixture of 2 or more
substances
The solute is(are) the substance(s) present in the
smaller amount(s)
The solvent is the substance present in the larger
amount
Solution
Solvent
Solute
Soft drink (l)
H2O
Sugar, CO2
Air (g)
N2
O2, Ar, CH4
Soft Solder (s)
Pb
Sn
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aqueous solutions
of KMnO4
An electrolyte is a substance that, when dissolved in
water, results in a solution that can conduct electricity.
A nonelectrolyte is a substance that, when dissolved,
results in a solution that does not conduct electricity.
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nonelectrolyte
weak electrolyte
strong electrolyte
Conduct electricity in solution?
Cations (+) and Anions (-)
Strong Electrolyte – 100% dissociation
NaCl (s)
H 2O
Na+ (aq) + Cl- (aq)
Weak Electrolyte – not completely dissociated
CH3COOH
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CH3COO- (aq) + H+ (aq)
Ionization of acetic acid
CH3COOH
CH3COO- (aq) + H+ (aq)
A reversible reaction. The reaction can
occur in both directions.
Acetic acid is a weak electrolyte because its
ionization in water is incomplete.
10
Hydration is the process in which an ion is surrounded
by water molecules arranged in a specific manner.
d-
d+
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H2O
Nonelectrolyte does not conduct electricity?
No cations (+) and anions (-) in solution
C6H12O6 (s)
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H 2O
C6H12O6 (aq)
Precipitation Reactions
Precipitate – insoluble solid that separates from solution
precipitate
Pb(NO3)2 (aq) + 2KI (aq)
PbI2 (s) + 2KNO3 (aq)
molecular equation
Pb2+ + 2NO3- + 2K+ + 2I-
PbI2 (s) + 2K+ + 2NO3-
ionic equation
Pb2+ + 2IPbI2
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PbI2 (s)
net ionic equation
K+ and NO3- are spectator ions
Precipitation of Lead Iodide
Pb2+ + 2I14
PbI2 (s)
PbI2
Solubility is the maximum amount of solute that will dissolve
in a given quantity of solvent at a specific temperature.
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Examples of Insoluble Compounds
CdS
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PbS
Ni(OH)2
Al(OH)3
Problem 4.20
Characterize the following compounds as
(a) soluble or (b) insoluble in water:
1. CaCO3
2. ZnSO4
3. Hg(NO3)2
4. HgSO4
5. NH4ClO4
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Writing Net Ionic Equations
1. Write the balanced molecular equation.
2. Write the ionic equation showing the strong electrolytes
completely dissociated into cations and anions.
3. Cancel the spectator ions on both sides of the ionic equation
4. Check that charges and number of atoms are balanced in the
net ionic equation
Write the net ionic equation for the reaction of silver nitrate
with sodium chloride.
AgNO3 (aq) + NaCl (aq)
Ag+ (aq) + NO3- (aq) + Na+ (aq) + Cl- (aq)
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Ag+ (aq) + Cl- (aq)
AgCl (s) + NaNO3 (aq)
AgCl (s) + Na+ (aq) + NO3- (aq)
AgCl (s)
Predict what happens when a potassium hydroxide solution
is mixed with a solution of sodium chloride. Write a net ionic
equation for the reaction.
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What do we expect to see if we add copper (II) sulfate to
sodium hydroxide? Write the molecular equation, ionic
equation, and net ionic equation.
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What do we expect to see if we add copper (II) sulfate to
sodium hydroxide? Write the molecular equation, ionic
equation, and net ionic equation.
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EXTRA PRACTICE
Predict what happens when a potassium phosphate solution
is mixed with a solution of calcium nitrate. Write a net ionic
equation for the reaction.
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Predict what happens when a silver nitrate solution is mixed
with a solution of potassium hydroxide. Write a net ionic
equation for the reaction.
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Types of Chemical Reactions
A+B↔C
C ↔A + B
A + BC ↔ AC + B
AB + CD ↔ AD + BC
 Synthesis (Combination)
 Decomposition
 Single Replacement
 Precipitation Reactions (Double
Replacement Reactions)
Acid + Base ↔ salt + water  Neutralization Reactions (Acid/Base)
Change of oxidation state  Redox Reactions
Hydrocarbon + O2 ↔ CO2 + H2O  Combustion
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NEUTRALIZATION REACTIONS
Acid/Base Reactions
Acid + Base ↔ Salt + Water
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Properties of Acids
Have a sour taste. Vinegar owes its taste to acetic acid. Citrus
fruits contain citric acid.
Cause color changes in plant dyes.
React with certain metals to produce hydrogen gas.
2HCl (aq) + Mg (s)
MgCl2 (aq) + H2 (g)
React with carbonates and bicarbonates
to produce carbon dioxide gas
2HCl (aq) + CaCO3 (s)
CaCl2 (aq) + CO2 (g) + H2O (l)
Aqueous acid solutions conduct electricity.
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Properties of Bases
Have a bitter taste.
Feel slippery. Many soaps contain bases.
Cause color changes in plant dyes.
Aqueous base solutions conduct electricity.
Examples:
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Arrhenius acid is a substance that produces H+ (H3O+) in water
Arrhenius base is a substance that produces OH- in water
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Hydronium ion, hydrated proton, H3O+
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A Brønsted acid is a proton donor
A Brønsted base is a proton acceptor
base
acid
acid
base
A Brønsted acid must contain at least one ionizable
proton!
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Monoprotic acids
HCl
H+ + Cl-
HNO3
H+ + NO3H+ + CH3COO-
CH3COOH
Strong electrolyte, strong acid
Strong electrolyte, strong acid
Weak electrolyte, weak acid
Diprotic acids
H2SO4
H+ + HSO4-
Strong electrolyte, strong acid
HSO4-
H+ + SO42-
Weak electrolyte, weak acid
Triprotic acids
H3PO4
H2PO4HPO4231
H+ + H2PO4H+ + HPO42H+ + PO43-
Weak electrolyte, weak acid
Weak electrolyte, weak acid
Weak electrolyte, weak acid
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Identify each of the following species as a Brønsted acid, base,
or both. (a) HI, (b) CH3COO-, (c) H2PO4-
HI (aq)
H+ (aq) + I- (aq)
CH3COO- (aq) + H+ (aq)
H2PO4- (aq)
CH3COOH (aq)
H+ (aq) + HPO42- (aq)
H2PO4- (aq) + H+ (aq)
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Brønsted acid
H3PO4 (aq)
Brønsted base
Brønsted acid
Brønsted base
Problem 4.32
Identify each of the following as either a
(a) Brønsted acid, (b) Brønsted base, or (c) both.
1. PO432. ClO23. NH4+
4. HCO3-
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Neutralization Reaction
acid + base
HCl (aq) + NaOH (aq)
H+ + Cl- + Na+ + OH-
H+ + OH-
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salt + water
NaCl (aq) + H2O
Na+ + Cl- + H2O
H2O
Neutralization Reaction Involving a Weak
Electrolyte
weak acid + base
HCN (aq) + NaOH (aq)
HCN + Na+ + OH-
HCN + OH-
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salt + water
NaCN (aq) + H2O
Na+ + CN- + H2O
CN- + H2O
Neutralization Reaction Producing a Gas
acid + base
2HCl (aq) + Na2CO3 (aq)
2H+ + 2Cl- + 2Na+ + CO32-
2H+ + CO32-
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salt + water + CO2
2NaCl (aq) + H2O +CO2
2Na+ + 2Cl- + H2O + CO2
H2O + CO2
Types of Chemical Reactions
A+B↔C
C ↔A + B
A + BC ↔ AC + B
AB + CD ↔ AD + BC
 Synthesis (Combination)
 Decomposition
 Single Replacement
 Precipitation Reactions (Double
Replacement Reactions)
Acid + Base ↔ salt + water  Neutralization Reactions (Acid/Base)
Change of oxidation state  Redox Reactions
Hydrocarbon + O2 ↔ CO2 + H2O  Combustion
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OXIDATION REDUCTION REACTIONS
Redox Reactions
Synthesis Reactions:
A+B↔C
Decomposition Reactions:
C ↔A + B
Single Replacement Reactions: A + BC ↔ AC + B
Combustion Reactions:
hydrocarbon + O2 ↔ CO2 + H2O
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Oxidation-Reduction Reactions
(electron transfer reactions)
2Mg
O2 + 4e40
2Mg2+ + 4e- Oxidation half-reaction (lose e-)
2O2Reduction half-reaction (gain e-)
2Mg + O2 + 4e2Mg2+ + 2O2- + 4e2Mg + O2
2MgO
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Zn (s) + CuSO4 (aq)
Zn
Zn2+ + 2e- Zn is oxidized
Cu2+ + 2e-
ZnSO4 (aq) + Cu (s)
Zn is the reducing agent
Cu Cu2+ is reduced Cu2+ is the oxidizing agent
Copper wire reacts with silver nitrate to form silver metal.
What is the oxidizing agent in the reaction?
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Oxidation number
The charge the atom would have in a molecule (or an
ionic compound) if electrons were completely transferred.
1. Free elements (uncombined state) have an oxidation
number of zero.
Na, Be, K, Pb, H2, O2, P4 = 0
2. In monatomic ions, the oxidation number is equal to
the charge on the ion.
Li+, Li = +1; Fe3+, Fe = +3; O2-, O = -2
3. The oxidation number of oxygen is usually –2. In H2O2
and O22- it is –1.
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4.4
4. The oxidation number of hydrogen is +1 except when
it is bonded to metals in binary compounds. In these
cases, its oxidation number is –1.
5. Group IA metals are +1, IIA metals are +2 and fluorine is
always –1.
6. The sum of the oxidation numbers of all the atoms in a
molecule or ion is equal to the charge on the
molecule or ion.
7. Oxidation numbers do not have to be integers.
Oxidation number of oxygen in the superoxide ion,
O2-, is –½.
-
HCO3
What are the oxidation numbers
of all the elements in HCO3- ?
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O = –2
H = +1
3x(–2) + 1 + ? = –1
C = +4
The Oxidation Numbers of Elements in their Compounds
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What are the oxidation numbers of
all the elements in each of these
compounds?
NaIO3
IF7
K2Cr2O7
NaIO3
IF7
F = -1
7x(-1) + ? = 0
I = +7
Na = +1 O = -2
3x(-2) + 1 + ? = 0
I = +5
K2Cr2O7
O = -2
K = +1
7x(-2) + 2x(+1) + 2x(?) = 0
Cr = +6
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Problem 4.50
Give the oxidation number for the underlined atoms or in each
of the following species:
a) Mg3N2
b) CsO2
c) CaC2
d) CO32e) C2O42f) ZnO22g) NaBH4
h) WO42-
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Types of Oxidation-Reduction Reactions
Synthesis/Combination Reaction
A+B
0
C
+3 -1
0
2Al + 3Br2
2AlBr3
Decomposition Reaction
C
+1 +5 -2
2KClO3
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A+B
+1 -1
0
2KCl + 3O2
Types of Oxidation-Reduction Reactions
Combustion Reaction
A + O2
B
0
0
S + O2
0
0
2Mg + O2
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+4 -2
SO2
+2 -2
2MgO
Types of Oxidation-Reduction Reactions
Displacement Reaction
A + BC
0
+1
+2
Sr + 2H2O
+4
0
TiCl4 + 2Mg
0
-1
Cl2 + 2KBr
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AC + B
0
Sr(OH)2 + H2 Hydrogen Displacement
0
+2
Ti + 2MgCl2
-1
Metal Displacement
0
2KCl + Br2
Halogen Displacement
The Activity Series for Metals
Hydrogen Displacement Reaction
M + BC
MC + B
M is metal
BC is acid or H2O
B is H2
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Ca + 2H2O
Ca(OH)2 + H2
Pb + 2H2O
Pb(OH)2 + H2
Problem 4.52
Which of the following metals can react with water to produce
H2 (g)?
a) Au
b) Li
c) Hg
d) Ca
e) Pt
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The Activity Series for Halogens
F2 > Cl2 > Br2 > I2
Halogen Displacement Reaction
0
-1
Cl2 + 2KBr
I2 + 2KBr
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-1
0
2KCl + Br2
2KI + Br2
Types of Oxidation-Reduction Reactions
Disproportionation Reaction
The same element is simultaneously oxidized
and reduced.
Example:
reduced
+1
0
Cl2 + 2OHoxidized
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-1
ClO- + Cl- + H2O
Classify each of the following reactions.
Ca2+ + CO32NH3 + H+
Zn + 2HCl
Ca + F2
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CaCO3
NH4+
ZnCl2 + H2
CaF2
Problem 4.54
Predict the outcome of the reactions represented by the
following equations by using the activity series, and balance
the equations.
Cu (s) + HCl (aq) 
I2 (g) + NaBr (aq) 
Mg (s) + CuSO4 (aq) 
Cl2 (g) + KBr (aq) 
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Problem 4.56
Classify the following redox reactions by type:
P4 + 10Cl2  4PCl5
2NO  N2 + O2
Cl2 + 2KI  I2 + 2KCl
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Solution Stoichiometry
Molarity
Dilutions
Gravimetric Analysis
Titrations
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Solution Stoichiometry
The concentration of a solution is the amount of solute
present in a given quantity of solvent or solution.
M = molarity =
moles of solute
liters of solution
What mass of KI is required to make 5.00 x 102 mL
of a 2.80 M KI solution?
M KI
volume of KI solution
5.00x102
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mL x
1L
1000 mL
moles KI
x
2.80 mol KI
1 L soln
x
M KI
166 g KI
1 mol KI
grams KI
= 232 g KI
Problem 4.60
Calculate the mass in grams of sodium nitrate required to
prepare 2.50 x 102 mL of a 0.707 M solution.
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Preparing a Solution of Known Concentration
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Dilution is the procedure for preparing a less concentrated
solution from a more concentrated solution.
Dilution
Add Solvent
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Moles of solute
before dilution (i)
=
Moles of solute
after dilution (f)
MiVi
=
MfVf
How would you prepare 60.0 mL of 0.200 M HNO3
from a stock solution of 4.00 M HNO3?
MiVi = MfVf
Mi = 4.00 M Mf = 0.200 M Vf = 0.0600 L
Vi =
MfVf
Mi
Vi = ? L
= 0.200 M x 0.0600 L = 0.00300 L = 3.00 mL
4.00 M
Dilute 3.00 mL of acid with water to a total volume
of 60.0 mL.
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Problem 4.70
Water is added to 25.0 mL of a 0.866 M KNO3 solution until
the volume of the solution is exactly 500 mL. What is the
concentration of the final solution?
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Problem 4.72
You have 505 mL of a 0.125 M HCl solution and you want to
dilute it to exactly 0.100 M. How much water should you add?
(assume that the volumes are additive.)
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Gravimetric Analysis
1. Dissolve unknown substance in water
2. React unknown with known substance to form a precipitate
3. Filter and dry precipitate
4. Weigh precipitate
5. Use chemical formula and mass of precipitate to determine
amount of unknown ion
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A 0.5662 g sample of an ionic compound containing chloride
ions and an unknown metal is dissolved in water, and treated
with excess AgNO3. If 1.0882 g of AgCl precipitate forms, what
is the percent by mass of Cl in the original compound?
35.45g Cl
%Cl 
100%  24.72%
143.4g AgCl
mass of Cl  0.24721.0882g  0.2690g
T hisis theamountof Cl in theoriginalsample,so
0.2690g
%Cl 
100%  47.51%
0.5662g
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Problem 4.78
A sample of 0.6760 g of an unknown compound containing
barium ions (Ba2+) is dissolved in water and treated with an
excess of Na2SO4. If the mass of the BaSO4 precipitate formed
is 0.4105 g, what is the percent by mass of Ba in the original
unknown compound?
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Titrations
In a titration a solution of accurately known concentration is
added gradually added to another solution of unknown
concentration until the chemical reaction between the two
solutions is complete.
Equivalence point – the point at which the reaction is complete
Indicator – substance that changes color at (or near) the
equivalence point
Slowly add base
to unknown acid
UNTIL
the indicator
changes color
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Titrations can be used in the analysis of
Acid-base reactions
H2SO4 + 2NaOH
2H2O + Na2SO4
Redox reactions
5Fe2+ + MnO4- + 8H+
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Mn2+ + 5Fe3+ + 4H2O
What volume of a 1.420 M NaOH solution is required
to titrate 25.00 mL of a 4.50 M H2SO4 solution?
WRITE THE CHEMICAL EQUATION!
H2SO4 + 2NaOH
M
volume acid
25.00 mL x
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acid
2H2O + Na2SO4
rxn
moles red
4.50 mol H2SO4
1000 mL soln
x
coef.
M
moles base
2 mol NaOH
1 mol H2SO4
x
base
volume base
1000 ml soln
1.420 mol NaOH
= 158 mL
Problem 4.86
Calculate the concentration (in molarity) of a NaOH solution if
25.0 mL of the solution are needed to neutralize 17.4 mL of a
0.312 M HCl solution.
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16.42 mL of 0.1327 M KMnO4 solution is needed to
oxidize 25.00 mL of an acidic FeSO4 solution. What is
the molarity of the iron solution?
WRITE THE CHEMICAL EQUATION!
5Fe2+ + MnO4- + 8H+
Mn2+ + 5Fe3+ + 4H2O
M
volume red
red
rxn
moles red
16.42 mL = 0.01642 L
0.01642 L x
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0.1327 mol KMnO4
1L
coef.
V
moles oxid
M oxid
oxid
25.00 mL = 0.02500 L
x
5 mol Fe2+
1 mol KMnO4
x
1
0.02500 L
Fe2+
= 0.4358 M
Problem 4.92
The SO2 present in air is mainly responsible for the acid
rain phenomenon. Its concentration can be determined by
titrating against a standard permanganate solution as
follows:
5SO2 + 2MnO4- + 2H2O  5SO42- + 2Mn2+ + 4H+
Calculate the number of grams of SO2 in a sample of air if
7.37 mL of 0.00800 M KMnO4 solution are required for
the titration.
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