Download Aqueous Solutions

Chapter 4
Type of Chemical Reactions and Solution
Stoichiometric
Water, Nature of aqueous solutions, types of
electrolytes, dilution.
Types of chemical reactions: precipitation,
acid-base and oxidation reactions.
Stoichiometry of reactions and balancing
the chemical equations.
1
Aqueous Solutions
Water is the dissolving
medium, or solvent.
2
Figure 4.1: (Left) The water molecule is polar.
(Right) A space-filling model of the water
molecule.
3
Figure 4.2: Polar water molecules interact with
the positive and negative ions of a salt assisting
in the dissolving process.
4
Some Properties of Water
 Water
is “bent” or V-shaped.
 The O-H bonds are covalent.
 Water is a polar molecule.
 Hydration occurs when salts dissolve in
water.
5
Figure 4.3: (a) The ethanol molecule contains a polar O—H bond similar to
those in the water molecule. (b) The polar water molecule interacts strongly
with the polar O—H bond in ethanol. This is a case of "like dissolving like."
6
A Solute
 dissolves
 changes
in water (or other “solvent”)
phase (if different from the solvent)
 is
present in lesser amount (if the same
phase as the solvent)
7
A Solvent
 retains
its phase (if different from the
solute)
 is
present in greater amount (if the same
phase as the solute)
8
General Rule for dissolution
Like dissolve like
Polar dissolve polar (water dissolve ethanol)
Non-polar dissolve nonpolar (benzene
dissolve fat)
9
Figure 4.5: When solid NaCl dissolves, the Na+
and Cl- ions are randomly dispersed in the water.
10
Electrolytes
Strong - conduct current efficiently
NaCl, HNO3
Weak - conduct only a small current
vinegar, tap water
Non - no current flows
pure water, sugar solution
11
Figure 4.4:
Electrical
conductivity
of aqueous
solutions.
12
Acids
Strong acids -
dissociate completely to produce
H+ in solution
hydrochloric and sulfuric acid
HCl , H2SO4
Weak acids -
dissociate to a slight extent to give
H+ in solution
acetic and formic acid
CH3COOH, CH2O
13
Bases
Strong bases - react completely with water to
give OH ions.
sodium hydroxide
Weak bases - react only slightly with water
to give OH ions.
ammonia
14
Figure 4.6:
HCl(aq) is
completely
ionized.
15
Figure 4.7:
An aqueous
solution of
sodium
hydroxide.
16
Figure 4.8: Acetic
acid
(HC2H3O2)
exists in water
mostly
as
undissociated
molecules. Only a
small percentage of
the molecules are
ionized.
17
Molarity
Molarity (M) = moles of solute per volume of
solution in liters:
moles of solute
M  molarity 
liters of solution
6 moles of HCl
3 M HCl 
2 liters of solution
18
Common Terms of Solution
Concentration
Stock - routinely used solutions prepared in
concentrated form.
Concentrated - relatively large ratio of solute
to solvent. (5.0 M NaCl)
Dilute - relatively small ratio of solute to
solvent. (0.01 M NaCl): (MV)initial=(MV)Final
19
Figure 4.10: Steps involved in the preparation of
a standard aqueous solution.
20
Figure 4.12:
Dilution Procedure (a) A measuring pipet is used
to transfer 28.7mL of 17.4 M acetic acid solution to a volumetric flask. (b) Water is
added
to
the
flask
to
the
calibration
mark.
(c) The resulting solution is 1.00 M acetic acid.
21
Practice Example
How many moles are in 18.2 g of CO2?
22
Practice Example
Consider the reaction
N2 + 3H2 = 2NH3
How many moles of H2 are needed to
completely react 56 g of N2?
23
Practice Example
How many grams are in 0.0150 mole of
caffeine C8H10N4O2
24
Practice Example
A solution containing Ni2+ is prepared by
dissolving 1.485 g of pure nickel in nitric
acid and diluting to 1.00 L. A 10.00 mL
aliquot is then diluted to 500.0 mL. What is
the molarity of the final solution?
(Atomic
weight:
Ni
=
58.70).
25
Practice Example
Calculate the number of molecules of
vitamin A, C20H30O in 1.5 mg of this
compound.
26
Practice Example
What is the mass percent of hydrogen in
acetic acid HC2H3O2
27
Types of Solution Reactions
 Precipitation
reactions
AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq)
 Acid-base
reactions
NaOH(aq) + HCl(aq)  NaCl(aq) + H2O(l)
 Oxidation-reduction
reactions
Fe2O3(s) + Al(s)  Fe(l) + Al2O3(s)
28
Simple Rules for Solubility
Most nitrate (NO3) salts are soluble.
Most alkali (group 1A) salts and NH4+ are soluble.
Most Cl, Br, and I salts are soluble (NOT Ag+, Pb2+, Hg22+)
Most sulfate salts are soluble (NOT BaSO4, PbSO4, HgSO4,
CaSO4)
5. Most OH salts are only slightly soluble (NaOH, KOH are
soluble, Ba(OH)2, Ca(OH)2 are marginally soluble)
6. Most S2, CO32, CrO42, PO43 salts are only slightly soluble.
1.
2.
3.
4.
29
Figure
4.13:
When
yellow
aqueous
potassium
chromate
is
added
to
a
colorless barium
nitrate solution,
yellow barium
chromate
precipitates.
30
Describing Reactions in Solution
Precipitation
1.
Molecular equation (reactants and
products as compounds)
AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq)
2.
Complete ionic equation (all strong
electrolytes shown as ions)
Ag+(aq) + NO3- (aq) + Na+ (aq) + Cl(aq)
AgCl(s) + Na+ (aq) + NO3- (aq)
31
Describing Reactions in Solution
(continued)
3.
Net ionic equation (show only
components that actually react)
Ag+(aq) + Cl(aq)  AgCl(s)
Na+ and NO3 are spectator ions.
32
Performing Calculations for
Acid-Base Reactions
1.
2.
3.
4.
5.
6.
List initial species and predict reaction.
Write balanced net ionic reaction.
Calculate moles of reactants.
Determine limiting reactant.
Calculate moles of required reactant/product.
Convert to grams or volume, as required.
Remember: n H+ = n OH-
(MV) H+ = (MV) OH33
Neutralization Reaction
acid + base
salt + water
HCl (aq) + NaOH (aq)
NaCl (aq) + H2O
H+ + Cl- + Na+ + OH--
Na+ + Cl- + H2O
H+ + OH-
H2O
34
4.3
Key Titration Terms
Titrant - solution of known concentration
used in titration
Analyte - substance being analyzed
Equivalence point - enough titrant added to
react exactly with the analyte
Endpoint - the indicator changes color so you
can tell the equivalence point has been
reached.
movie
35
Oxidation-Reduction Reactions
(electron transfer reactions)
2Mg (s) + O2 (g)
2Mg
O2 + 4e-
2MgO (s)
2Mg2+ + 4e- Oxidation half-reaction (lose e-)
2O2-
Reduction half-reaction (gain e-)
2Mg + O2 + 4e2Mg + O2
2Mg2+ + 2O2- + 4e2MgO
36
37
Redox Reactions
• Many practical or everyday examples of
redox reactions:
–
–
–
–
–
–
–
Corrosion of iron (rust formation)
Forest fire
combustion
Charcoal grill
Natural gas burning
Batteries
Production of Al metal from Al2O3 (alumina)
Metabolic processes
38
Rules for Assigning Oxidation States
1. Oxidation state of an atom in an element = 0
2. Oxidation state of monatomic element = charge
3. Oxygen = 2 in covalent compounds (except in
peroxides where it = 1)
4. H = +1 in covalent compounds
5. Fluorine = 1 in compounds
6. Sum of oxidation states = 0 in compounds
Sum of oxidation states = charge of the ion
39
40
Zn (s) + CuSO4 (aq)
Zn
ZnSO4 (aq) + Cu (s)
Zn2+ + 2e- Zn is oxidized
Cu2+ + 2e-
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?
Cu (s) + 2AgNO3 (aq)
Cu
Ag+ + 1e-
Cu(NO3)2 (aq) + 2Ag (s)
Cu2+ + 2eAg Ag+ is reduced
Ag+ is the oxidizing agent
41
IF7
Oxidation numbers of all
the elements in the
following ?
F = -1
7x(-1) + ? = 0
I = +7
NaIO3
Na = +1 O = -2
3x(-2) + 1 + ? = 0
I = +5
K2Cr2O7
O = -2
K = +1
7x(-2) + 2x(+1) + 2x(?) = 0
Cr = +6
42
Balancing by Half-Reaction
Method
1.
Write separate reduction, oxidation
reactions.
2.
For each half-reaction:
 Balance elements (except H, O)
 Balance O using H2O
 Balance H using H+
 Balance charge using electrons
43
Balancing by Half-Reaction
Method (continued)
3.
If necessary, multiply by integer to
equalize electron count.
4.
Add half-reactions.
5.
Check that elements and charges are
balanced.
44
Half-Reaction Method Balancing in Base
1.
Balance as in acid.
2.
Add OH that equals H+ ions (both
sides!)
3.
Form water by combining H+, OH.
4.
Check elements and charges for balance.
45
Balancing Redox Equations
•Example: Balance the following redox reaction:
•Cr2O72- + Fe2+
Cr3+ + Fe3+ (acidic soln)
1) Break into half reactions:
Cr2O72Cr3+
Fe2+
Fe3+
46
Balancing Redox Equations
2) Balance each half reaction:
Cr2O72Cr3+
Cr2O72Cr2O72-
2 Cr3+
2 Cr3+ + 7 H2O
Cr2O72- + 14 H+
6 e- + Cr2O72- + 14 H+
2 Cr3+ + 7 H2O
2 Cr3+ + 7 H2O
47
Balancing Redox Equations
2) Balance each half reaction (cont)
Fe2+
Fe2+
Fe3+
Fe3+ + 1 e-
48
Balancing Redox Reactions
3) Multiply by integer so e- lost = e- gained
6 e- + Cr2O72- + 14 H+
Fe2+
2 Cr3+ + 7 H2O
Fe3+ + 1 e-
x6
49
Balancing Redox Reactions
3) Multiply by integer so e- lost = e- gained
6 e- + Cr2O72- + 14 H+
6 Fe2+
2 Cr3+ + 7 H2O
6 Fe3+ + 6 e-
4) Add both half reactions
Cr2O72- + 6 Fe2+ + 14 H+
2 Cr3+ + 6 Fe3+ + 7 H2O
50
Balancing Redox Reactions
5) Check the equation
Cr2O72- + 6 Fe2+ + 14 H+
2 Cr
7O
6 Fe
14 H
+24
2 Cr3+ + 6 Fe3+ + 7 H2O
2 Cr
7O
6 Fe
14 H
+ 24
51
Balancing Redox Reactions
• Procedure for Basic Solutions:
– Divide the equation into 2 incomplete half
reactions
• one for oxidation
• one for reduction
52
Balancing Redox Reactions
– Balance each half-reaction:
different
•
•
•
•
•
•
•
balance elements except H and O
balance O atoms by adding H2O
balance H atoms by adding H+
add 1 OH- to both sides for every H+ added
combine H+ and OH- on same side to make H2O
cancel the same # of H2O from each side
balance charge by adding e- to side with greater overall
+ charge
53
Balancing Redox Equations
– Multiply each half reaction by an integer so that
• # e- lost = # e- gained
– Add the half reactions together.
• Simply where possible by canceling species
appearing on both sides of equation
– Check the equation
• # of atoms
• total charge on each side
54
Balancing Redox Reactions
Example: Balance the following redox
reaction.
NH3 + ClOCl2 + N2H4 (basic soln)
1) Break into half reactions:
NH3
ClO-
N2H4
Cl2
55
Balancing Redox Reactions
2) Balance each half reaction:
NH3
N2H4
2 NH3
N2H4
2 NH3
N2H4 + 2 H+
+ 2 OH+ 2 OH2 NH3 + 2 OH
2 NH3 + 2 OH
-
2
H2O
N2H4 + 2 H2O
N 2 H 4 + 2 H 2O + 2 e 56
Balancing Redox Reactions
2) Balance each half reaction:
ClOCl2
2 ClOCl2
2 ClOCl2 + 2 H2O
2 ClO- + 4 H+
Cl2 + 2 H2O
+ 4 OH+ 4 OH2 ClO- + 4 H2O
Cl2 + 2 H2O + 4 OH2 ClO- + 2 H2O
Cl2 + 4 OH2 e- + 2 ClO- + 2 H2O
Cl2 + 4 OH57
Balancing Redox Reactions
3) Multiply by integer so # e- lost = # e- gained
2 NH3 + 2 OHN2H4 + 2 H2O + 2 e2 e- + 2 ClO- + 2 H2O
Cl2 + 4 OH4) Add both half reactions
2 NH3 + 2 OH- + 2ClO- + 2 H2O
N2H4 + 2 H2O + Cl2 + 4 OH-
58
Balancing Redox Reactions
5) Cancel out common species
2 NH3 + 2 OH- + 2 ClO- + 2 H2O
6) Check final equation:
2 NH3 + 2 ClO2N
6H
2 Cl
2O
-2
N2H4 + 2 H2O + Cl2 + 4 OH-
2
N2H4 + Cl2 + 2 OH2N
6H
2 Cl
2O
-2
59
60
61
Practice Example
In the following the oxidizing agent is:
5H2O2 + 2MnO4- + 6H+  2Mn2+ + 8H2O + 5O2
a. MnO4b. H2O2
c. H+
d. Mn2+
e. O2
62
Practice Example
Determine the coefficient of Sn in acidic
solution
Sn + HNO3  SnO2 + NO2 + H2O
1
63
Practice Example
The sum of the coefficients when they are
whole numbers in basic solution:
Bi(OH)3 + SnO22-  Bi + SnO32-
13
64
http://www.chemistrycoach.com/balancing_redox_in_acid.htm#Bal
ancingRedoxEquationsinAcidicorBasicMedium
65
http://www.chemistrycoach.com/tutorials-5.htm#Oxidation-Reduction
66
• http://www.sstdt.org.sa/arc-e.htm
67
QUESTION
An unknown substance dissolves readily in
water but not in benzene (a nonpolar solvent).
Molecules of what type are present in the
substance?
1) Neither polar nor nonpolar
2) Polar
3) Either polar or nonpolar
4) Nonpolar
5) none of these
68
ANSWER
2)
Polar
Section 4.1 Water, the Common Solvent (p. 134)
The solubility rule for molecular compounds is
“like dissolves like”, that is, polar molecules
dissolve in polar solvents and nonpolar
molecules dissolve in nonpolar solvents.
69
QUESTION
How many grams of NaCl are contained in
350. mL of a 0.250 M solution of sodium
chloride?
1) 41.7 g
2) 5.11 g
3) 14.6 g
4) 87.5 g
5) None of these
70
ANSWER
2)
5.11 g
Section 4.3 The Composition of Solutions
(p. 140)
Volume (L) times concentration (mol/L) gives
moles. Moles are then converted to grams.
71
QUESTION
What volume of 18.0 M sulfuric acid must be
used to prepare 15.5 L of 0.195 M H2SO4?
1) 168 mL
2) 0.336 L
3) 92.3 mL
4) 226 mL
5) None of these
72
ANSWER
1)
168 mL
Section 4.3 The Composition of Solutions
(p. 140)
Use the dilution formula, M1  V1 = M2  V2, M is
in mol/L and V is in L.
73
QUESTION
The net ionic equation for the reaction of
aluminum sulfate and sodium hydroxide
contains which of the following species?
3+
1) 3Al (aq)
–
2) OH (aq)
–
3) 3OH (aq)
3+
4) 2Al (aq)
5) 2Al(OH)3(s)
74
ANSWER
3)
–
3OH (aq)
Section 4.6 Describing Reactions in Solution
(p. 154)
The net ionic equation is found by canceling the
spectator ions from the total ionic equation.
75
QUESTION
Which of the following is a strong acid?
1) HF
2) KOH
3) HClO4
4) HClO
5) HBrO
76
ANSWER
3)
HClO4
Section 4.2 The Nature of Aqueous Solutions:
Strong and Weak Electrolytes (p. 136)
Memorization of the list of strong acids will allow
one to determine the difference between strong
acids and weak acids.
77
QUESTION
All of the following are weak acids except:
1) HCNO.
2) HBr.
3) HF.
4) HNO2.
5) HCN.
78
ANSWER
2)
HBr.
Section 4.2 The Nature of Aqueous Solutions:
Strong and Weak Electrolytes (p. 136)
Knowing the list of strong acids will allow one to
determine which acids are strong and which are
weak.
79
QUESTION
Which of the following is not a strong base?
1) Ca(OH)2
2) KOH
3) NH3
4) LiOH
5) Sr(OH)2
80
ANSWER
3)
NH3
Section 4.2 The Nature of Aqueous Solutions:
Strong and Weak Electrolytes (p. 136)
Knowing the list of strong bases will allow one to
determine which bases are strong and which are
weak.
81
QUESTION
The interaction between solute particles and
water molecules, which tends to cause a salt to
fall apart in water, is called:
1) hydration.
2) polarization.
3) dispersion.
4) coagulation.
5) conductivity.
82
ANSWER
1)
hydration.
Section 4.1 Water, the Common Solvent (p. 134)
Hydration is the process of water molecules
surrounding and stabilizing ions so that they can
be pulled into solution.
83
QUESTION
The concentration of a salt water solution that
sits in an open beaker decreases over time.
1) True
2) False
84
ANSWER
2) False
Section 4.3 The Composition of Solutions
(p. 140)
The amount of water decreases over time so the
concentration (mol NaCl/volume of water)
increases over time.
85
QUESTION
The following reactions:
–
Pb + 2I  PbI2
4+
–
3+
2Ce + 2I  I2 + 2Ce
+
–
HOAc + NH3  NH4 + OAc
2+
are examples of
86
QUESTION
(continued)
1)
2)
3)
4)
5)
acid-base reactions.
unbalanced reactions.
precipitation, acid-base, and redox
reactions, respectively.
redox, acid-base, and precipitation
reactions, respectively.
precipitation, redox, and acid-base
reactions, respectively.
87
ANSWER
5)
precipitation, redox, and acid-base
reactions, respectively.
Section 4.5 Precipitation Reactions (p. 151)
4+
PbI2 is insoluble, Ce changes to Ce
HOAc is an acid while NH3 is a base.
3+
and
88
QUESTION
Which of the following salts is insoluble in
water?
1) Na2S
2) K3PO4
3) Pb(NO3)2
4) CaCl2
5) All of these are soluble in water.
89
ANSWER
5)
All of these are soluble in water.
Section 4.5 Precipitation Reactions (p. 153)
According to the solubility rules for ionic
compounds, compounds containing Group IA
ions or nitrate ions will always be soluble.
Compounds containing halides are generally
soluble, aside from silver, lead and mercury(I)
halides.
90
QUESTION
2+
When NH3(aq) is added to Cu (aq), a
precipitate initially forms. Its formula is:
2+
1) Cu(NH3)4
2) Cu(NO3)2
3) Cu(OH)2
2+
4) Cu(NH3)4
5) CuO
91
ANSWER
3)
Cu(OH)2
Section 4.4 Types of Chemical Reactions
(p. 148)
Ammonia produces hydroxide ion in water:
+
–
NH3 + H2O  NH4 + OH
92
QUESTION
In the balanced molecular equation for the
neutralization of sodium hydroxide with sulfuric
acid, the products are:
1) NaSO4 + H2O
2) NaSO3 + 2H2O
3) 2NaSO4 + H2O
4) Na2S + 2H2O
5) Na2SO4 + 2H2O
93
ANSWER
5)
Na2SO4 + 2H2O
Section 4.8 Acid-Base Reactions (p. 158)
The salt is made from the anion of the acid and
the cation of the base.
94
QUESTION
What mass of NaOH is required to react exactly
with 25.0 mL of 1.2 M H2SO4?
1) 1.2 g
2) 1.8 g
3) 2.4 g
4) 3.5 g
5) None of these
95
ANSWER
3)
2.4 g
Section 4.8 Acid-Base Reactions (p. 158)
Remember that the reaction is 2NaOH + H2SO4
 Na2SO4 + 2H2O, so there are two moles of
NaOH used per one mole of H2SO4.
96
QUESTION
In which of the following does nitrogen have an
oxidation state of +4?
1) HNO3
2) NO2
3) N2O
4) NH4Cl
5) NaNO2
97
ANSWER
2)
NO2
Section 4.9 Oxidation-Reduction Reactions
(p. 164)
Oxygen almost always has an oxidation state of
–2 when part of a compound. The exception is
when it is part of a peroxide. For example,
hydrogen peroxide H2O2. Then it has an
oxidation state of –1.
98
QUESTION
In the reaction 2Cs(s) + Cl2(g)  2CsCl(s), Cl2 is
1) the reducing agent.
2) the oxidizing agent.
3) oxidized.
4) the electron donor.
5) two of these
99
ANSWER
2)
the oxidizing agent.
Section 4.9 Oxidation-Reduction Reactions
(p. 164)
Metals lose electrons, so they are oxidized,
making the other reactant an oxidizing agent.
100
QUESTION
Which of the following statements is(are) true?
Oxidation and reduction:
1) cannot occur independently of each other.
2) accompany all chemical changes.
3) describe the loss and gain of electron(s),
respectively.
4) result in a change in the oxidation states of
the species involved.
5) 1, 3, and 4 are true
101
ANSWER
5)
1, 3, and 4 are true.
Section 4.9 Oxidation-Reduction Reactions
(p. 167)
(2) is false because certain reactions, such as
double displacement reactions, are not redox
reactions.
102
QUESTION
How many of the following are oxidationreduction reactions?
NaOH + HCl  NaCl + H2O
Cu + 2AgNO3  2Ag + Cu(NO3)2
Mg(OH)2  MgO + H2O
N2 + 3H2  2NH3
103
QUESTION
(continued)
1)
2)
3)
4)
5)
0
1
2
3
4
104
ANSWER
3)
2
Section 4.9 Oxidation-Reduction Reactions
(p. 164)
If an element is found on the reactant’s side, this
is almost always a redox reaction, since an
element usually becomes part of a compound
during a chemical reaction.
105