Download Chapter 4 Aqueous Reactions and Solution Stoichiometry

Chemistry, The Central Science, 11th edition
Theodore L. Brown; H. Eugene LeMay, Jr.;
and Bruce E. Bursten
Chapter 4
Aqueous Reactions and
Solution Stoichiometry
Solution Chemistry and
the Hydrosphere
Temperature and Salinity
(halinity) in the Oceans
Solution Terms
• Solutions are homogeneous mixtures
of two or more substances.
• Solvent – the substance present in a solution
in the greatest proportion (in number of
moles).
• Solute - the substance dissolved in the
solvent.
Aqueous Solution (aq) – a solution where
water is the solvent.
Ionic Dissociation
• When an ionic
substance dissolves
in water, the solvent
pulls the individual
ions from the crystal
and solvates them.
• This process is called
dissociation or
ionization.
NaCl(s) - Na+(aq) + Cl-(aq)
Ionic Dissociation
• An electrolyte is a
substances that
dissociates into ions
when dissolved in
water.
• Since it is in water
(solvent) these are
referred to as
aqueous solutions
Aqueous Solutions
Water is the dissolving medium
Partial Charges on
the water molecule
104.5o
Aqueous Solutions
Water is the dissolving medium
Partial Charges on the
water molecule
red= area of negative
charge
blue= area of positive
charge
Some Properties of Water
•
Water is “bent” or V-shaped.
•
Water is a molecular compound.
•
Water is a polar (separaton of partial
charge) molecule.
•
Hydration occurs when ionic
compounds dissolve in water.
Hydration of a solute Ion (Mg2+)
Electrolytes
Dissolved sugars would be
examples of nonelectrolytes
• An electrolyte is a
substances that
dissociates into ions
(charged species)
when dissolved in
water.
• A nonelectrolyte may
dissolve in water, but
it does not dissociate
into ions when it does
so.
Electrolytes and
Nonelectrolytes
Soluble ionic
compounds tend
to be electrolytes.
Electrolytes and
Nonelectrolytes
Molecular
compounds tend to
be nonelectrolytes,
except for acids and
bases.
H3COH (methanol) (aq)
Electrolytes
• A strong electrolyte
dissociates completely
when dissolved in
water.
• A weak electrolyte
only dissociates
partially when
dissolved in water.
Electrolytes
Strong - conduct
current efficiently
Examples: Aqueous
solutions of NaCl,
HNO3, HCl
Electrolytes
Weak - conduct
only a small
current
(vinegar,
tap water)
Non Electrolyte
A solution in which
no ionization occurs.
There is no
conduction of
electrical current.
Examples: Aqueous
solutions of sugar,
ethelyne glycol
Strong Electrolytes Are…
• Strong acids
• Strong bases
Strong Electrolytes Are…
• Strong acids
• Strong bases
• Soluble ionic salts
Solubility Guide
4.20 Using the solubility guide line (Table 4.1), predict the
solubility of the following.
a) Ni(OH)2
b) PbBr2
c) Ba(NO3)2
d) AlPO4
e) AgCl
f) CaCO3
g) (NH4)2CO3
Precipitation Reactions
When one mixes ions
that form compounds
that are insoluble (as
could be predicted by
the solubility
guidelines), a
precipitate is formed.
Metathesis (Exchange) Reactions
• Metathesis comes from a Greek word that
means “to transpose.”
AgNO3 (aq) + KCl (aq)
AgCl (s) + KNO3 (aq)
Metathesis (Exchange) Reactions
• It appears the ions in the reactant
compounds exchange, or transpose, ions.
AgNO3 (aq) + KCl (aq)
AgCl (s) + KNO3 (aq)
Solution Chemistry
• It is helpful to pay attention to exactly
what species are present in a reaction
mixture (i.e., solid, liquid, gas, aqueous
solution).
• If we are to understand reactivity, we
must be aware of exactly what is
changing during the course of a
reaction.
Molecular Equation
The molecular equation lists the reactants
and products in their molecular form. The
(aq) indicates that the ions are dissociated in
aqueous solution.
AgNO3 (aq) + KCl (aq)
AgCl (s) + KNO3 (aq)
Ionic Equation
• In the ionic equation all strong electrolytes (strong
acids, strong bases, and soluble ionic salts) are
dissociated into their ions.
• This more accurately reflects the species that are
found in the reaction mixture.
Ag+ (aq) + NO3- (aq) + K+ (aq) + Cl- (aq)
AgCl (s) + K+ (aq) + NO3- (aq)
Net Ionic Equation
• To form the net ionic equation, cross out anything
that does not change from the left side of the
equation to the right.
• The only things left in the equation are those things
that change phase (i.e., chemically react (precipitate)
during the course of the reaction.
• Those things that didn’t change (and were deleted
from the net ionic equation) are called spectator ions.
Ag+(aq) + NO3-(aq) + K+(aq) + Cl-(aq)
AgCl (s) + K+(aq) + NO3-(aq)
Net Ionic Equation
• To form the net ionic equation, cross out the
spectator ions (that not change from the left side of
the equation to the right).
• The only things left in the equation are those things
that change (i.e., chemically react) during the course
of the reaction.
Ag+(aq)
+
Cl (aq)
AgCl (s)
Writing Net Ionic Equations
1. Write a balanced molecular equation.
2. Dissociate all strong electrolytes.
3. Cross out anything that remains
unchanged from the left side to the
right side of the equation.
4. Write the net ionic equation with the
species that remain.
Writing Net Ionic Equations
Molecular
Equation
Ionic Equation
Net Ionic
Equation
AgNO3 (aq) + KCl (aq)
AgCl (s) + KNO3 (aq)
Ag+ (aq) + NO3- (aq) + K+ (aq) + Cl- (aq)
AgCl (s) + K+ (aq) + NO3- (aq)
Writing Net Ionic Equations
Molecular
Equation
Ionic Equation
Net Ionic
Equation
AgNO3 (aq) + KCl (aq)
AgCl (s) + KNO3 (aq)
Ag+ (aq) + NO3- (aq) + K+ (aq) + Cl- (aq)
AgCl (s) + K+ (aq) + NO3- (aq)
Ag+(aq) + Cl-(aq)
AgCl (s)
1.
2.
3.
4.
Write a balanced molecular equation.
Dissociate all strong electrolytes.
Cross out anything that remains unchanged from the left side
to the right side of the equation.
Write the net ionic equation with the species that remain.
Write the net ionic equation for the reaction of MgSO4 with BaCl2.
Acids
• Arrhenius defined acids
as substances that
increase the
concentration of H+
when dissolved in water.
• Brønsted and Lowry
defined acids as proton
donors.
Acid-Base Reactions
• Bronsted-Lowry acids are proton (H+)
donors.
• Bronsted-Lowry bases are proton
acceptors.
• Free hydrogen ions don’t exist in water
because they strongly associate with a
water molecule to create a hydronium
ion (H3O+) (a hydrated proton).
Acid-Base Reactions
• A neutralization reaction takes place
when an acid reacts with a base and
produces a solution of a salt and water.
• A salt is made up of the cation
characteristic of the base and the anion
characteristic of the acid.
• Example: HCl + NaOH ---> NaCl + H2O
Acids
There are only seven
strong acids:
•
•
•
•
•
•
•
Hydrochloric (HCl)
Hydrobromic (HBr)
Hydroiodic (HI)
Nitric (HNO3)
Sulfuric (H2SO4)
Chloric (HClO3)
Perchloric (HClO4)
Bases
• Arrhenius defined bases
as substances that
increase the
concentration of OH−
when dissolved in water.
• Brønsted and Lowry
defined them as proton
acceptors.
Bases
The strong bases
are the soluble
metal salts of
hydroxide ion:
•
•
•
•
Alkali metals
Calcium
Strontium
Barium
Strong Acids and Bases
• A strong acid or strong base is completely
ionized in aqueous solution.
• HCl, HBr, HI, HNO3, HClO3, HClO4 and
H2SO4 are all strong acids. All other acids are
assumed to be weak acids.
• A weak acid or weak base only partially
ionizes (dissociates) in aqueous solution.
• Amphiprotic substances can behave as either
a proton acceptor or a proton donor. Water is
an example.
Types of Acid/Base Equations
• Molecular Equations have reactants and
products written as undissociated (not
ionized) molecules.
HCl(aq) + NaOH(aq) ---> NaCl(aq) + H2O(l)
• Overall Ionic Equations show all the
species, both ionic and molecular present
in aqueous solution for the reaction.
H+ + Cl- + Na+ + OH- --> Na+ + Cl- + H2O
Continued
• Strong acids and strong bases are
written as the corresponding ions in an
overall ionic equation.
• Net Ion Equations describe the actual
chemical reaction occurring.
H+ + OH- ----> H2O
• The Na+ and Cl- ions are spectator
ions in this reaction, because they are
unchanged by the reaction (as is the
solvent).
Neutralization Reactions
Generally, when solutions of an acid and a base are
combined, the products are a salt and water.
CH3COOH (aq) + NaOH (aq)
CH3COONa (aq) + H2O (l)
Neutralization Reactions
When a strong acid reacts with a strong base, the net
ionic equation is…
HCl (aq) + NaOH (aq)
NaCl (aq) + H2O (l)
Neutralization Reactions
When a strong acid reacts with a strong base, the net
ionic equation is…
HCl (aq) + NaOH (aq)
NaCl (aq) + H2O (l)
H+ (aq) + Cl- (aq) + Na+ (aq) + OH-(aq)
Na+ (aq) + Cl- (aq) + H2O (l)
Neutralization Reactions
When a strong acid reacts with a strong base, the net
ionic equation is…
HCl (aq) + NaOH (aq)
NaCl (aq) + H2O (l)
H+ (aq) + Cl- (aq) + Na+ (aq) + OH-(aq)
Na+ (aq) + Cl- (aq) + H2O (l)
H+ (aq) + OH- (aq)
H2O (l)
Problem
Write a balanced molecular equation and a net ionic
equation for the following reactions:
a.
Solid magnesium hydroxide reacts with a
solution of sulfuric acid.
Gas-Forming Reactions
• Some metathesis reactions do not give the
product expected.
• In this reaction, the expected product (H2CO3)
decomposes to give a gaseous product
(CO2).
CaCO3 (s) + HCl (aq)
CaCl2 (aq) + CO2 (g) + H2O (l)
Gas-Forming Reactions
When a carbonate or bicarbonate reacts with
an acid, the products are a salt, carbon
dioxide, and water.
CaCO3 (s) + HCl (aq)
NaHCO3 (aq) + HBr (aq)
CaCl2 (aq) + CO2 (g) + H2O (l)
NaBr (aq) + CO2 (g) + H2O (l)
Gas-Forming Reactions
Similarly, when a sulfite reacts with an acid,
the products are a salt, sulfur dioxide, and
water.
SrSO3 (s) + 2 HI (aq)
SrI2 (aq) + SO2 (g) + H2O (l)
Gas-Forming Reactions
• This reaction gives the predicted product, but
you had better carry it out in the hood, or you
will be very unpopular!
• But just as in the previous examples, a gas is
formed as a product of this reaction.
Na2S (aq) + H2SO4 (aq)
Na2SO4 (aq) + H2S (g)
Write the products and net ionic equation for the
following
HNO3(aq) + Ca(OH)2(aq) →
FeO(s) + HClO4(aq) →
Write the products and net ionic equation for the reaction
of solid Strontium carbonate with aqueous perchloric
acid.
Oxidation-Reduction Reactions
• An oxidation occurs
when an atom or ion
loses electrons.
• A reduction occurs
when an atom or ion
gains electrons.
• One transfer cannot
occur without the
other also occurring.
Oxidation-Reduction Half Reactions
• When copper wire is
immersed in a solution
of silver nitrate it is
oxidized.
• Cu ---> Cu2+ + 2 e• Ag+ + e- ---> Ag
• Silver ion is reduced,
and the copper metal
is oxidized.
Oxidation-Reduction Reactions
Fe2O3(s) + Al(s)
Fe(l) + Al2O3(s)
M
Oxidized
X
Reduced
Loses e-
Gains e-
Oxidation State
Increases
Reduction State
Increases
Reducing Agent
Oxidizing Agent
e-
M+
X-
Oxidation Numbers
To determine if an oxidation-reduction
reaction has occurred, we assign an
oxidation number to each element in a
neutral compound or charged entity.
Rules for Assigning
Oxidation States
1. The oxidation number of elements in a neutral
molecule sum to zero or sum to charge of the ion in
an ion.
2. Oxidation state of an atom in an element = 0
3. Oxidation state of monatomic ion = charge
4. Fluorine = 1 in all compounds
5. Hydrogen = +1, Oxygen = 2 in most compounds
(except in peroxides where oxygen = 1)
6. Unless combined with O, or F the halogens are -1.
Determining Oxidation Numbers
SO2
CrO42NH3
ClO3SF6
Cl2
Oxygen is -2 and Sulfur is +4
Displacement Reactions
• In displacement reactions,
ions in solution are
displaced (or replaced) by
oxidation of another
element.
Zn(s) + 2 HBr(aq)
Fe(s) + Ni(NO3)2(aq)
ZnBr2(aq) + H2(g)
Fe(NO3)2(aq) + Ni(s)
Displacement Reactions
In this reaction,
silver ions oxidize
copper metal.
Cu (s) + 2 Ag+ (aq)
Cu2+ (aq) + 2 Ag (s)
Displacement Reactions
The reverse reaction,
however, does not
occur (under normal conditions).
Cu2+ (aq)
+ 2 Ag (s)
x
Cu (s) + 2 Ag+ (aq)
Activity Series
For example, scrap iron is
used to reduce Cu2+ in
Berkeley Pit (mine) waters.
That is any element on the
list can be oxidized by the
ions of the elements below it
on the list.
Try Sample Exercise 4.10
(p. 142).
Activity Series
Using the Activity Series, write a balanced chemical equation
for the following reactions: a) zinc metal is added to a solution
of Magnesium sulfate, b) aluminum metal is added to a solution
of cobalt(II) sulfate.
Balancing redox reactions by method of half-reactions.
1.
2.
3.
4.
5.
6.
7.
8.
9.
Separate the reactions into oxidation and reduction processes
Work with one (ox or red) first
Balance number of non-oxygen, non-hydrogen atoms first.
Then balance oxygen with water
Then balance hydrogen with H+
Then balance charge with electrons.
Then balance other half-reaction using steps 3 through 6.
Balance numbers of electrons between both half reactions and then
add together.
If you need to change to a basic solution then add the same number
of hydroxide ions (OH-) to each side as H+ ions in balanced equation.
These rules for balancing redox reaction are on page 847 in
the text but you can also find them on the course website.
See “A Closer Look” and “Strategies in Chemistry”, p. 143.
Balancing by Half-Reaction
Method in Acid Solution
1. Write separate reduction, oxidation halfreactions.
KMnO4 + Cu+  Cu2+ + Mn2+
Ox. rxn:
•
Red. Rxn
•
Cu+  Cu2+
Cu+  Cu2+ + eKMnO4  Mn2+ + K+
KMnO4 + 5e-  Mn2+ + K+
Balance number of non-oxygen, non-hydrogen
atoms first.
3. For each half-reaction, balance oxygen with water
KMnO4 + 5e-  Mn2+ + K+ + 4 H2O
4. Then balance hydrogen with H+
8H+ + KMnO4 + 5e-  Mn2+ + K+ + 4 H2O
5. Balance numbers of electrons between both half reactions and
then add together.
5(Cu+  Cu2+ + e-)
+ (KMnO4 + 8 H+ + 5 e-  Mn2+ + K+ + 4 H2O)
5 Cu+ + KMnO4 + 8 H+ + 5 e-  5 Cu2+ + 5 e- + Mn2+ + K+ + 4H2O
Practice
Balance the following equation in acid solution:
Br-(aq) + MnO4-(aq)
Br2( ) + Mn2+(aq)
Br-(aq) + MnO4-(aq)
→
Br2(l) + Mn2+(aq)
Balance the following oxidation-reduction reaction in
acidic and then for basic solution.
Ag(s) + CN- + O2(g)  Ag(CN)2-(aq) + H2O(l)
Ag(s)  Ag(CN)2-(aq)
Ag(s) + 2 CN-  Ag(CN)2-(aq)
Ag(s) + 2 CN-  Ag(CN)2-(aq) + eO2(g)  H2O(l)
O2(g)  2 H2O(l)
O2(g) + 4 H+  2 H2O(l)
O2(g) + 4 H+ + 4 e-  2 H2O(l)
4(Ag(s) + 2 CN-  Ag(CN)2-(aq) + e-)
O2(g) + 4 H+ + 4 e-  2 H2O(l)
4 Ag(s) + 8 CN- + O2(g) + 4 H+  4 Ag(CN)2-(aq) + 2 H2O(l)
To convert to basic solution (add hydroxide to each side):
4 Ag(s) + 8 CN- + O2(g) + 4 H+ + 4 OH-  4 Ag(CN)2-(aq) + 2 H2O(l)+ 4 OH4 Ag(s) + 8 CN- + O2(g) + 4 H2O(l) 4 Ag(CN)2-(aq) + 2 H2O(l) + 4 OH4 Ag(s) + 8 CN- + O2(g) + 2 H2O(l) 4 Ag(CN)2-(aq) + 4 OH-
How much of a solute dissolves?
Concentration – ratio of the quantity of
solute to either the mass or volume of the
solution or solvent in which the solute is
dissolved.
Consequently, there are many different
concentration terms depending on:
1. What is used to identify the quantity of
solute (e.g. moles, mass, volume, etc.)
2. Whether the denominator of the ratio is
the solvent or solution.
3. Whether mass or volume is used as the
unit in the denominator.
Molarity
• Two solutions can contain the same
compounds but be quite different because the
proportions of those compounds are different.
• Molarity is one way to measure the
concentration of a solution.
Molarity (M) =
moles of solute
volume of solution in liters
Concentration of Solutions
• Molarity (M) = moles of solute per
volume of solution in liters:
M
molarity
3 M HCl
moles of solute
liters of solution
6 moles of HCl
2 liters of solution
Problem
What is the molarity of nitrate ions
in a 0.0525 M solution of aluminum
nitrate?
Other Common Units of Concentration
• Molality (m): moles solute/kg
solvent
• ppm: parts per million; mg solute/kg
solution
• ppb: mg solute/kg solution
• ppt: ng solute/kg solution
• % by weight:
(grams of solute/total g solution) x
100%
• Mole fraction:
mole solute/total moles in solution
Average Concentrations of the 11 Major Constituents of Seawater.
Ions
g/kg
mmol/kg
mmol/L
Na+
10.781
468.96
480.57
K+
0.399
10.21
10.46
Mg2+
1.284
52.83
51.14
Ca2+
0.4119
10.28
10.53
Sr2+
0.00794
0.0906
0.0928
Cl-
19.353
545.88
559.40
SO42-
2.712
28.23
28.93
HCO3-
0.126
2.06
2.11
Br-
0.0673
0.844
0.865
B(OH)3
0.0257
0.416
0.426
F-
0.00130
0.068
0.070
Total
35.169
1119.87
1147.59
Mixing a Solution
• To create a solution of a
known molarity, one
weighs out a known mass
(and, therefore, number of
moles) of the solute.
• The solute is added to a
volumetric flask, and
solvent is added to the line
on the neck of the flask.
Problem
What is the molarity of an aqueous solution
prepared by adding 36.5 g of barium chloride
(208.233g/mol) to enough water to make 750.0
mL of solution?
Determining the Number of
Moles of Solute
#
of
moles
of
solute
Molarity =
# of liter of solution
• # moles Solute = (Molarity)(# Liters of Solution)
• n = M x volume (in Liters)
Problem
How many grams of aluminum nitrate (212.996
g/mol) are required to make 500.0 mL of a
0.0525 M aqueous solution?
Dilution
• One can also dilute a more concentrated
solution by
 Using a pipet to deliver a volume of the solution to
a new volumetric flask, and
 Adding solvent to the line on the neck of the new
flask.
Dilution
The molarity of the new solution can be determined
from the equation
Mc
Vc = Md
Vd,
where Mc and Md are the molarity of the concentrated and dilute
solutions, respectively, and Vc and Vd are the volumes of the
two solutions.
Practice
Hydrochloric acid is obtained in 12.0 M stock
solution. What volume of stock solution is
required to make 500.0 mL of a 0.145 M dilute
solution?
Using Molarities in
Stoichiometric Calculations
Problem
What mass of barium sulfate
(233.390g/mo) is produced when 100.0
mL of a 0.100 M solution of barium
chloride is mixed with 100.0 mL of a 0.100
M solution of iron (III) sulfate?
What mass of barium sulfate (233.390g/mo) is produced when 100.0 mL of
a 0.100 M solution of barium chloride is mixed with 100.0 mL of a 0.100 M
solution of iron (III) sulfate?
Titration
Titration is an
analytical
technique in
which one can
calculate the
concentration
of a solute in
a solution.
Titration Terms
• A titration is a volumetric analytical method
used to determine the concentration of an
unknown solution by reacting it with a
standard solution.
• A standard solution is a solution of known
concentration.
• The equivalence point in a titration is reached
when enough standard solution has been
added to completely react with the unknown
solution.
• The end point in a titration is reached when
an indicator changes color or a specific pH
(acid/base) or potential (redox) is reached.
Stoichiometry Calculations
H2SO4 + 2 NaOH -->
Na2SO4 + 2 H2O
What is the
concentration of
sulfuric acid if
15.00ml of it reacts
with 18.45 mL of a
0.0973 M NaOH
solution?
Titration Example
H2SO4 + 2 NaOH ---> Na2SO4 + 2 H2O
NaOH
H2SO4
End Point
H2SO4(aq) + 2 NaOH(aq) --> Na2SO4(aq) + 2 H2O(l)
What is the concentration of sulfuric acid if 15.00 ml of it reacts
with 18.45 mL of a standardized 0.0973 M NaOH solution?
Prob. 114 The arsenic in a 1.22g sample of a pesticide was
converted to AsO43- by chemical treatment. It was then
titrated with Ag+ to form Ag3AsO4 as a precipitate.
a) What is the oxidation state of arsenic in AsO43- ?
b) If it took 25.0 mL of 0.102 M Ag+ to reach the end point for
this titration, what is the mass percentage of arsenic in the
pesticide.
25.0 mL of 0.102 M Ag+ to reach the end point