Chemistry II
Aqueous Reactions and
Solution Chemistry
Chapter 4
In this chapter we will consider
chemical processes that occur in
aqueous solutions: precipitation
reactions, acid base reactions
and oxidation – reduction
reactions.
 We
will then consider
concentrations and how the
concepts of stoichiometry can
be applied to concentrations.
Water has many properties that
allow it to help support life.
One of these properties is
that it can dissolve a wide
variety of materials. For this
reason water is often referred
to as what?
Water has many properties that
allow it to help support life.
One of these properties is
that it can dissolve a wide
variety of materials. For this
reason water is often referred
to as what?

The universal solvent
What are aqueous solutions?
What are aqueous solutions?
 Solutions
in which water is the
dissolving medium.
Section 1: General Properties of
Aqueous Solutions
Define solution:
Define solution:
A
homogeneous mixture of
two or more substances.
What is the difference between
solvent and solute?
What is the difference between
solvent and solute?
 The
substance that is present in
greater quantity is usually called the
solvent.
 The other substances in the solution
are solutes. Solutes are dissolved in
the solvent.
Electrolytic Properties

Pure water is not a good conductor of
electricity. The presence of ions causes
aqueous solutions to become good
conductors. Ions carry the charge from
one electrode to the next.
What is an electrolyte?
What is an electrolyte?
A
substance whose aqueous
solutions contains ions.
What is a non-electrolyte?
What is a non-electrolyte?
A
substance that does not
form ions in solution.
Ionic Compounds in water
What does it mean when we say
“ an ionic solid dissociates”
into its component ions as it
dissolves?”
What does it mean when we say
“ an ionic solid dissociates”
into its component ions as it
dissolves?”
 Each
ion separates from the
solid structure and disperses
throughout the solution.
What is a polar molecule. Explain
the significance of this fact
with respect to the
dissociation of ionic solids.
What is a polar molecule. Explain
the significance of this fact
with respect to the
dissociation of ionic solids.
What is a polar molecule. Explain the
significance of this fact with respect to
the dissociation of ionic solids.
A
polar molecule has one end
that has a partial positive
charge and a partial negative
charge.
 Although
water is an electrically
neutral molecule, one end of the
molecule is electron rich and
carries a partial negative charge.
The hydrogen side of the
molecule has a partial positive
charge.
 When
ionic compounds dissolve
the anions are surrounded by the
water molecules so that the
hydrogen side of the molecule
surrounds the anion. The cations
are surrounded by the oxygen
side of the water molecule. This
configuration stabilizes the ions in
solution.
How can we predict the charges
of the ions present in
solution?
How can we predict the charges
of the ions present in
solution?
 Remember the formulas and
charges of the common ions.
 i.e., Na2SO4 will separate into two
Na+ ions and one SO42- ions.
 In solution for every one sodium
sulfate three ions are formed.
Molecular Compounds in Water

When a molecular compound dissolves
in water, the solution usually consists
of intact molecules dispersed
throughout the solution. Usually
molecular compounds are nonelectrolytes. An exception to this rule is
acids
In what way do acids appear to
not follow the rule?
In what way do acids appear to
not follow the rule?
 Acids
are molecular
compounds that will
disassociate or ionize into ions
in aqueous solutions.
+
HCl → H + Cl
The two categories of
electrolytes are strong and weak
electrolytes
What are strong electrolytes?
What are strong electrolytes?
 Those
solutes that exist in
solution completely or nearly
completely as ions.
 Most ionic compounds and some
acids and bases are strong
electrolytes.
What are weak electrolytes?
 Those solutes that exist in
solution in the form of
molecules but only partially
disassociated into ions.
Can you determine if a solute is
a strong or weak electrolyte
by how well it dissolves?
Can you determine if a solute is
a strong or weak electrolyte
by how well it dissolves?
 No, for example acetic acid
(vinegar) is very soluble in water,
but only partially dissociates into
ions.
How can we indicate that an
electrolyte is a weak
electrolyte?
How can we indicate that an
electrolyte is a weak
electrolyte?
We can use double arrows to show that the
reaction is significant in both directions.
HC2H3O2(aq) ↔ H+(aq) + C2H3O2-(aq)

* The state of equilibrium between molecules and
ions varies from one weak electrolyte to another.
How do chemists indicate the
ionization of strong
electrolytes?
How do chemists indicate the
ionization of strong
electrolytes?
With the use of a single arrow.
HCl (aq) → H+(aq) + Cl-(aq)
The single arrow indicates that the ions have
no tendency to recombine to molecules.

In a the next few sections we will learn
how to predict if a compound is a
strong electrolyte, weak electrolyte or
non-electrolyte. For now, in general,
soluble ionic compounds are always
strong electrolytes.
How can we identify compounds
are being ionic?
How can we identify compounds
are being ionic?
Ionic compounds are composed of metals
and nonmetals
 NaCl
 FeSO4
 Al(NO3)3
 NH4Br

The diagram on the right represents an
aqueous solution of one of the following
compounds: MgCl2, KCl, or K2SO4. Which
solution does the drawing best
represent?
If you were to draw diagrams (such as
that shown on exercise 4.1)
representing aqueous solutions of each
of the following ionic compounds, how
many anions would you show if the
diagram contained six cations?
 (a) NiSO4, (b) Ca(NO3)2, (c) Na3PO4, (d)
Al2(SO4)3

Homework

Page 145 1-9
Section 2_ Precipitation
Reactions
What are precipitation reactions?
What are precipitation reactions?
Reactions
that result in
the formation of an
insoluble product.
What is a precipitate?
What is a precipitate?
 An
insoluble solid formed by a
reaction in solution.
When do precipitation reaction
occur?
 Precipitation
reactions
occur when certain pairs of
oppositely charged ions
attract each other so
strongly that they form an
insoluble ionic solid.
For example
Pb(NO3)2(aq) + 2KI(aq)→ PbI2(s) + 2KNO3(aq)
Solubility Guidelines for Ionic
compounds
What is solubility ?
What is solubility ?
 The
amount of substance that
can be dissolved in a given
quantity of solvent.
 Any substance with a solubility
less than 0.001 mol/L is
referred to as insoluble.

Experimental observations have led to
guidelines for predicting solubility
(page 118 Table 4.1)
Sample exercise 4.2
Classify the following ionic compounds
as soluble or insoluble in water:
 (a) sodium carbonate (Na2CO3),
 (b) lead sulfate (PbSO4).

Classify the following compounds as
soluble or insoluble in water:
 (a) cobalt(II) hydroxide,
 (b) barium nitrate,
 (c) ammonium phosphate.

How can we predict whether a
precipitate forms when we mix
aqueous solutions of two
strong electrolytes?
 1.)
note the ions present in the
reactants.
How can we predict whether a
precipitate forms when we mix
aqueous solutions of two
strong electrolytes?
 1.)
note the ions present in the
reactants.
 2.) Consider possible
combinations of cations and
anions.
How can we predict whether a
precipitate forms when we mix
aqueous solutions of two
strong electrolytes?
 3.)
Use solubility guidelines to
determine if any combinations
are insoluble
How can we predict whether a
precipitate forms when we mix
aqueous solutions of two
strong electrolytes?
 For
example, will a precipitate
form when Mg(NO3)2 and NaOH
are mixed?
 Precipitation
reactions are a
type of double replacement
reactions. They are also known
as exchange or metathesis
reactions.
AX + BY → AY + BX
AgNO3 + KCl → AgCl + KNO3
Exchange (Metathesis)
Reactions
 In
exchange reactions the
chemical formulas of the
products are based on the
charges of the ions.
Sample exercise 4.3
(a) Predict the identity of the precipitate
that forms when solutions of BaCl2 and
K2SO4 are mixed.
 (b) Write the balanced chemical
equation for the reaction.

 (a)
What compound
precipitates when solutions
of Fe2(SO4)3 and LiOH are
mixed?
 (b) Write a balanced
equation for the reaction.
 Will
a precipitate form when
solutions
of Ba(NO3)2 and KOH are
mixed?
Ionic Equations
In writing chemical equations for
reactions in aqueous solutions, it is
often helpful to know if the dissolved
substances are present mainly as
molecules or as ions. For example:
Molecular EquationPb(NO3)2(aq) + 2KI(aq) → PbI2(s) + 2KNO3(aq)

Ionic Equation
Pb+2(aq)+2NO3(aq)-+2K+(aq)+ 2I-(aq)→

PbI2(s) + 2K+(aq) + 2NO3-(aq)
What is a complete ionic
equation?
 An
equation such as the one
above written with all strong
electrolytes written as ions.
What are spectator ions?
 Ions
that appear in identical
forms on both the reactant and
product side of the equation.
What are spectator ions?
Pb+2(aq)+2NO3(aq)-+2K+(aq)+ 2I-(aq)→
PbI2(s) + 2K+(aq) + 2NO3-(aq)
What is a net ionic equation?
What is a net ionic equation?
When the spectator ions
are cancelled out we are
left with the net ionic
equation.
Pb+2(aq) + 2I-(aq)→ PbI2(s)
Note:
if every ion in a
complete ionic equation is
a spectator ion, then no
reaction occurs.
 The
net ionic equation includes
only the ions and molecules
directly involved in the reaction.
How can net ionic equations be
used?
 They
can be used to show
similarities between large
numbers of reactions
 A net ionic equation shows
that more than one set of
reactions can lead to the same
net reaction.
What are the steps for writing
net ionic equations?
What are the steps for writing
net ionic equations?
 1.)
Write a balanced molecular
equation for the reaction.
 2.) Rewrite the equation to show
the ions that form only strong
electrolytes are written in ionic
form.
 3.) Identify and cancel spectator
ions.
Sample Exercise 4.4
Write the net ionic equation for
the precipitation reaction that
occurs when solutions of
calcium chloride and sodium
carbonate are mixed.
 Write
the net ionic equation for
the precipitation reaction that
occurs when aqueous
solutions of silver nitrate and
potassium phosphate are
mixed.
4.3 Acid Base Reactions
 Acids
and bases happen to
be common electrolytes.
Acids
What are acids?
What are acids?
 Acids
are substances that ionize in
aqueous solutions to form hydrogen
ions, increasing the concentration of
hydrogen ions in solution.
 Because hydrogen ions are just a
proton, acids are known as proton
donors.
 The
number of hydrogen ions
produced depends on the
number of hydrogen atoms
present. Acids like HCl and
HNO3 yield one hydrogen ion
per molecule. Some acids such
as H2SO4 yield 2 hydrogen
ions.
What is the difference between a
monoprotic an diprotic acid?
What is the difference between a
monoprotic an diprotic acid?
 Acids
•
•
that yield one hydrogen ion are
monoprotic such as HCl
HCl → H+(aq) + Cl-(aq)
Acids that yield two hydrogen ions are
diprotic i.e., H2SO4
H2SO4 → 2H+(aq) + SO42-(aq)
Bases
What are bases?
What are bases?
Bases are substances that accept H+ ions
thereby reducing the number of H+ ions in
solution.
 Bases produce OH- ions when dissolved in
water.
 When bases are dissolved in water they
release OH- and create more OH- ions by
bonding to all of the available H+ ions.

What are bases?
Some bases do not contain OH NH3 is an example
NH3 + H2O → NH4+ + OH
Strong And Weak Acids and
Bases
What are strong and bases?
What are strong and bases?
 Acids
and bases that are
strong electrolytes are called
strong acids or bases.
What are weak acids and bases?
What are weak acids and bases?
 Acids
and bases that are
only partly ionized in
solution.
Table 4.2 on page 122 lists the
strong acids and bases. These
should be committed to memory.
What are the strong acids and
bases.
Table 4.2 on page 122 lists the
strong acids and bases. These
should be committed to memory.
What are the strong acids and
bases.
ACIDS
 HCl
 HBr
 HI

HClO3
HClO4
HNO3
H2SO4
Table 4.2 on page 122 lists the
strong acids and bases. These
should be committed to memory.
What are the strong acids and
bases.
Bases
 The alkali group hydroxides
 Alkaline earth metals Ca, Sr and Ba
hydroxides.

Identifying Strong and Weak
Electrolytes
Is a soluble ionic compound a strong
electrolyte, weak electrolyte or a
nonelectrolyte?
Is a soluble ionic compound a strong
electrolyte, weak electrolyte or a
nonelectrolyte?
 All
soluble ionic compounds are
strong electrolytes.
How can you tell if a soluble
molecular compound is a strong
electrolyte, weak electrolyte or
nonelectrolyte?
 All strong acids and bases are
strong electrolytes
 All weak acids and bases are
weak electrolytes
 All other soluble molecular
compounds are nonelectrolytes.
The following diagrams represent aqueous
solutions of three acids (HX, HY, and HZ) with
water molecules omitted for clarity. Rank them
from strongest to weakest.

Classify each of the following dissolved
substances as a strong electrolyte,
weak electrolyte, or nonelectrolyte:
CaCl2, HNO3, C2H5OH (ethanol), HCOOH
(formic acid), KOH.
Neutralization Reactions and
Salts
What is a neutralization
reaction?
What is a neutralization
reaction?
 When
a solution of an acid and
a base are mixed and the pH
of the mixture is neither acidic
or basic.
What are the products of a
neutralization reaction?
What are the products of a
neutralization reaction?
 The
products of a neutralization
reaction are always a salt and
water.
 HCl + NaOH→ NaCl + H2O
What is the definition of a salt?
What is the definition of a salt?
 Any
ionic compound whose
cation comes from a base and
whose anion comes from an
acid.
What is the net ionic equation for
all neutralization reactions?
What is the net ionic equation for
all neutralization reactions?
H+(aq) + OH-(aq)→ H2O(l)
What type of reaction is a
neutralization reaction?
A
double replacement (also
known as a metathesis reaction
or exchange reaction)
Mg(OH)2 + 2HCl → MgCl2 + 2H2O
Sample exercise 4.7
(a) Write a balanced molecular equation
for the reaction between aqueous
solutions of acetic acid (CH3COOH)
and barium hydroxide, Ba(OH)2.
 (b) Write the net ionic equation for this
reaction.

(a) Write a balanced molecular equation
for the reaction of carbonic acid
(H2CO3) and potassium hydroxide
(KOH).
 (b) Write the net ionic equation for this
reaction.

Acid Base Reactions with gas
formation
There are two bases besides OHthat react with H+. Two of these
include the sulfide and
carbonate ions.
There are two bases besides OHthat react with H+. Two of these
include the sulfide and
carbonate ions.
Both of these react with acids to
form gases.
2HCl + Na2S → H2S + 2NaCl
HCl + NaHCO3 → NaCl + H2CO3
Both NaHCO3 ( Sodium
carbonate) and Na2HCO3
(Sodium Bicarbonate) are used
as acid neutralizers and
antacids.
 Read
chemistry at work
antacids.
 Homework
Page 146
22- 34 evens.
Section 4
Oxidation Reduction Reactions
What is corrosion?
 The
conversion of a metal
into a metal compound by a
reaction between the metal
and some substance in its
environment.
What is corrosion?
 When
a metal corrodes it lose
electrons and forms cations
 Ca + 2HCl → CaCl2 + H2
 0
2+
 I______________l
 Calcium is oxidized because it lost
electrons
What is corrosion?
+ 2HCl → CaCl2 + H2

+1
0

l_______________l
 Hydrogen is reduced
because it gained electrons.
 Ca
What are oxidation – reduction
(redox) reactions?
 Reactions
in which electrons
are transferred between
reactants.
What is oxidation?
When
an atom becomes
positively charged.
When it has lost electrons
Loss of electrons by a
substance is called
oxidation.
The term oxidation is used
because the first reactions of
this sort to be studied were
reactions with oxygen.
What is reduction?
 Gain
of electrons from a
substance.
 The oxidation of one substance
is always accompanied by the
reduction of another
substance.
Oxidation Numbers
What are oxidation numbers?
 The oxidation number of an
atom in a substance is the
actual charge of the atom if
it is a monoatomic ion
or it is the hypothetical
charge assigned using a set
of rules
Rules for assigning oxidation
numbers
 1.)
For an atom in the
elemental form, the
oxidation number is always
zero.
 H2, Ca, O2
Rules for assigning oxidation
numbers
 2.)
For any monatomic ion,
the oxidation number equals
the charge on the ion
 K+ = 1 +
 S2- = 2-
Rules for assigning oxidation
numbers
 3.)
non-metals usually have
negative oxidation numbers.
oxygen is usually 2- w/ the
exception of the peroxide ion
(O2) which has the oxidation
number 1
a.)
Rules for assigning oxidation
numbers
b.)
hydrogen has an oxidation
number of 1+ when bonded to a
nonmetal [(HCl) H 1+ ; Cl 1-]
and has a oxidation of 1- when
bonded to a metal
[ CaH2 – Ca 2+ , H 1-]
Rules for assigning oxidation
numbers
 The sum of the oxidation
numbers of all atoms in a
nuetral compound is zero.
 The sum of the oxidation
numbers in a polyatomic ion
equals the charge of the ion.
Sample exercise 4.8 page 128
Determine the oxidation number of
sulfur in each of the following:
 (a) H2S,
 (b) S8,
 (c) SCl2,
 (d) Na2SO3,
 (e) SO42–.

What is the oxidation state of the
element in each of the following:
(a) P2O5
(b) NaH
(c) Cr2O7 2–
(d) SnBr4
(e) BaO2
Oxidation of metals by acids and
salts
 Some
common types of
redox reactions are
combustion reactions and
reactions between metals
and acids or salts.
Oxidation of metals by acids and
salts
 The
common form of an acid
reacting with a metal is
A + BX → AX + B
Zn + 2HCl → Zn Br2 + H2
What do we call these types of
reactions and why are they
classified as redox reactions?
 These
reactions are called
displacement or single
replacement reactions. The ion
in solution is displaced or
replaced through the oxidation
of an element.
Use the reaction between
magnesium and hydrochloric
acid to show that oxidation and
reduction have occurred.
Mg(s) + 2HCl → MgCl2(aq) + H2
0
1+ 1- 2+ 10
l___oxidized__l
l_____reduced_____l
Write the net ionic equation for
the reaction of magnesium and
hydrochloric acid.
Mg(s)+ 2H+(aq)+ 2Cl- → Mg2+(aq) + 2Cl- + H2(g)
Cl- is a spectator ion.
Mg(s)+ 2H+(aq) → Mg2+(aq) + H2(g)
Metals can also be oxidized by
aqueous solutions of various
salts. Show the oxidation –
reduction that occurs when iron
reacts with nickel II nitrate
Fe(s) + Ni(NO3)3(aq) → Fe(NO3)2(aq) + Ni(s)
0
+2 -1
+2 -1
0
l______oxidized_____l
l___________reduced______l
NO3 is the spectator ion.
Net ionic equation
Fe(s) + Ni 2+(aq) → Fe 2+ (aq) + Ni(s)
0
+2
+2
0
l______oxidized_____l
l_____reduced______l
Sample 4.9
Write the balanced molecular
and net ionic equations for the
reaction of aluminum with
hydrobromic acid.
Write the balanced molecular
and net ionic equations for the
reaction between magnesium
and cobalt(II) sulfate
* What is oxidized and what is reduced in
the reaction?
The Activity Series
 Different
metals vary in the
ease with which they are
oxidized.
What is the activity series?
 The
activity series is a list of metals
arranged in order of decreasing ease
of oxidation. The metals at the top of
the table are the most easily oxidized.
 They react the most easily to form
compounds.
What are active metals?
 The
metals at the top of the activity
series are the most easily oxidized
metals.
Which are the noble metals?
 The
metals are the bottom of
the activity series. These
metals are very stable and can
be used to make coins and
jewelry.
How can the activity series be
used to predict the outcome of
reactions?
Any metal on the list can be oxidized by the
ions of the an element below it.
Cu + HCl → No reaction
Copper is not oxidized by hydrogen because
hydrogen is not below copper
Cu + AgNO3 → Ag + Cu(NO3)2
Copper is oxidized by silver because silver
is below copper on the activity series.
Sample Exercise 4.10

Will an aqueous solution of iron(II)
chloride oxidize magnesium metal? If
so, write the balanced molecular and
net ionic equations for the reaction.
Which of the following metals
will be oxidized by Pb(NO3)2:
Zn, Cu, Fe?
Homework
Section 5
Concentrations of Solutions
Define concentration The
amount of solute
dissolved in a given quantity of
solvent.
What is Molarity?
Molarity
(M) expresses the
concentration of a solution
as the number of moles of
solute in a liter of solution.
 Molarity
(M) = moles of solute
volume of soln.(L)
Sample 4.11
Calculate the molarity of a
solution made by dissolving
23.4 g of sodium sulfate
(Na2SO4) in enough water to
form 125 mL of solution.
Calculate the molarity of a
solution made by dissolving
5.00 g of glucose (C6H12O6) in
sufficient water to form exactly
100 mL of solution.
Expressing the Concentration of
an Electrolyte
 When
an ionic compound
dissolves, the relative
concentrations of the ions
introduced into the solution
depend on the chemical
formula.
Expressing the Concentration of
an Electrolyte
1 mol of NaCl – 1 mole Na+ ions
1 mole Cl- ions
1 mol of Na2SO4- 2 mole Na+ ions
21 mole SO4 ions
Sample exercise 4.12
What are the molar
concentrations of each of
the ions present in a
0.025 M aqueous solution of
calcium nitrate?
What is the molar
concentration of K+ ions in a
0.015 M solution of
potassium carbonate?
Interconverting Molarity, Moles
and Volume
Because molarity contains 3
quantities; molarity, moles and
volume. Dimensional analysis
can be used to find any of
these values if we know the
other two.
Calculate the number of moles of
HNO3 in 2.00L of a 0.200M HNO3
# mol HNO3= ( 2 L HNO3) ( .200molHNO3)
1 L HNO3
= .4 mol HNO3
Sample exercise 4.13
How many grams of Na2SO4
are required to make 0.350 L
of 0.500 M Na2SO4?
How many grams of Na2SO4
are there in 15 mL of 0.50 M
Na2SO4?
Dilutions
What are stock solutions?
The concentrated solutions.
When solvent is added to
dilute a stock solution the
number of moles of solute
before dilution is equal to
the number of moles of
solute after dilution.
To prepare 250mL of a 0.100 M
CuSO4 from a stock of 1M
CuSO4…

1st determine the number of moles of CuSO4
we will need in the dilute solution.

(.250 L) ( .10 mol/ 1L) = .0250 mol CuSO4
To prepare 250mL of a 0.100 M
CuSO4 from a stock of 1M
CuSO4…

Then determine the volume of stock solution
needed
L conc. Soln.= .025 mol CuSO4 ( 1L/ 1 mole CuSO4)
= .025 L of 1 molar CuSO4 = 25 mL
Add 25 mL of 1 molar CuSO4 to a 250 mL volumetric
flask and bring up to volume.
To work the same problem
quickly we can note
Moles of solute in concentrated
soln. = moles of solute in diluted
soln.
Mconc.Vconc= Mdil.Vdil
( 1M) ( Vconc) = (.1M) ( 250 mL)
Vconc = 25 mL
Sample exercise 4.14
 How
many milliliters of 3.0 M
H2SO4 are needed to make
450 mL of 0.10 M H2SO4?
What volume of 2.50 M
lead(II) nitrate solution
contains 0.0500 mol of Pb 2+ ?
How many milliliters of 5.0 M
K2Cr2O7 solution must be
diluted to prepare 250 mL of
0.10 M solution?
If 10.0 mL of a 10.0 M stock
solution of NaOH is diluted to
250 mL, what is the
concentration of the resulting
stock solution?
Solution Stoichiometery and
Chemical Analysis
 Recall
that the coefficients
in a balanced equation give
the relative number of moles
of reactants and products.
Sample exercise 4.15
How
many grams of
Ca(OH)2 are needed to
neutralize 25.0 mL of .10M
HNO3
How many grams of NaOH are
needed to neutralize 20.0 mL
of 0.150 M H2SO4 solution?
How many liters of 0.500 M
HCl(aq) are needed to react
completely with 0.100 mol of
Pb(NO3)2(aq), forming a
precipitate of PbCl2(s)?
Titrations
What is a titration?
A
titration involves combining
a sample of the solution with a
reagent solution of known
concentration called the
standard solution.
What is the equivalence point?
 The
point at which
stoichiometrically equivalent
quantities are brought
together.
How does a chemist know when
the equivalence point is
reached?
 An indicator is used. The indicator
will show pH changes when the
color changes the acid has been
nuetralized. The color change
indicates the end point of the
titration.
Sample exercise 4.16



The quantity of Cl– in a municipal water supply is
determined by titrating the sample with Ag+. The
reaction taking place during the titration is
Ag+(aq) + Cl-(aq) → AgCl(s)
The end point in this type of titration is marked by
a change in color of a special type of indicator. (a)
How many grams of chloride ion are in a sample of
the water if 20.2 mL of 0.100 M Ag+ is needed to
react with all the chloride in the sample? (b) If the
sample has a mass of 10.0 g, what percent Cl–
does it contain?

A sample of an iron ore is dissolved in acid,
and the iron is converted to Fe2+. The sample
is then titrated with 47.20 mL of 0.02240 M
MnO4– solution. The oxidation-reduction
reaction that occurs during titration is as
follows:


(a) How many moles of MnO4– were added to
the solution? (b) How many moles of Fe2+
were in the sample? (c) How many grams of
iron were in the sample? (d) If the sample
had a mass of 0.8890 g, what is the
percentage of iron in the sample?

One commercial method used to peel
potatoes is to soak them in a solution
of NaOH for a short time, remove them
from the NaOH, and spray off the peel.
The concentration of NaOH is normally
in the range of 3 to 6 M. The NaOH is
analyzed periodically. In one such
analysis, 45.7 mL of 0.500 M H2SO4 is
required to neutralize a 20.0-mL sample
of NaOH solution. What is the
concentration of the NaOH solution?

What is the molarity of an NaOH
solution if 48.0 mL is needed to
neutralize 35.0 mL of 0.144 M H2SO4?
(b) How many milliliters of 0.50 M Na2SO4
solution are needed to provide 0.038
mol of this salt?