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Chapter 4
The Major Classes of Chemical Reactions
4-1
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The Major Classes of Chemical Reactions
4.1 The Role of Water as a Solvent
4.2 Writing Equations for Aqueous Ionic Reactions
4.3 Precipitation Reactions
4.4 Acid-Base Reactions
4.5 Oxidation-Reduction (Redox) Reactions
4.6 Elements in Redox Reactions
4-2
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Figure 4.1
4-3
Electron distribution in molecules of H2 and H2O.
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Figure 4.2
4-4
The dissolution of an ionic compound.
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Figure 4.3
4-5
The electrical conductivity of ionic solutions.
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Sample Problem 4.1 Determining Moles of Ions in Aqueous Ionic Solutions
PROBLEM:
How many moles of each ion are in the following solutions?
(a) 5.0 mol of ammonium sulfate dissolved in water
(b) 78.5 g of cesium bromide dissolved in water
(c) 7.42x1022 formula units of copper(II) nitrate dissolved in water
(d) 35 mL of 0.84 M zinc chloride
PLAN: We have to relate the information given and the number of moles of
ions present when the substance dissolves in water.
H2O
SOLUTION: (a) (NH4)2SO4(s)
5.0 mol (NH4)2SO4
4-6
2NH4+(aq) + SO42-(aq)
2 mol NH4+
1 mol
(NH4)2SO4
= 10. mol NH4+
5.0 mol SO42-
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Sample Problem 4.1
Determining Moles of Ions in Aqueous Ionic
Solutions
continued
H2O
Cs+(aq) + Br-(aq)
(b) CsBr(s)
mol CsBr
78.5 g CsBr
= 0.369 mol CsBr
212.8 g CsBr
H2O
(c) Cu(NO3)2(s)
= 0.369 mol Cs+
= 0.369 mol Br-
Cu2+(aq) + 2NO3-(aq)
= 0.123 mol Cu2+
mol Cu(NO3)2
7.42x1022 formula
= 0.123 mol Cu(NO3)2
units Cu(NO3)2 6.022x1023 formula units
= 0.246 mol NO
(d) ZnCl2(aq)
35 mL ZnCl2
H2O
Zn2+(aq) + 2Cl-(aq)
1L
0.84 mol ZnCl2
103mL
L
= 2.9x110-2 mol Zn2+
4-7
= 2.9x110-2 mol ZnCl2
= 5.8x10-2 mol Cl-
3
-
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Sample Problem 4.2
Determining the Molarity of H+ Ions in Aqueous
Solutions of Acids
PROBLEM: Nitric acid is a major chemical in the fertilizer and explosives
industries. In aqueous solution, each molecule dissociates and the
H becomes a solvated H+ ion. What is the molarity of H+(aq) in 1.4M
nitric acid?
PLAN:
Use the formula to find the molarity of H+.
+
SOLUTION: One mole of H (aq) is released per mole of nitric acid (HNO3)
H2O
HNO3(l)
H+(aq) + NO3-(aq)
1.4M HNO3(aq) should have 1.4M H+(aq).
4-8
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Writing Equations for Aqueous Ionic Reactions
The molecular equation
shows all of the reactants and products as intact, undissociated compounds.
The total ionic equation
shows all of the soluble ionic substances dissociated into ions.
The net ionic equation
eliminates the spectator ions and shows the actual
chemical change taking place.
4-9
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Figure 4.4
4-10
A precipitation reaction and its equation.
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Figure 4.6
The reaction of Pb(NO3)2 and NaI.
NaI(aq) + Pb(NO3)2 (aq)
PbI2(s) + NaNO3(aq)
2NaI(aq) + Pb(NO3)2 (aq)
PbI2(s) + 2NaNO3(aq)
2Na+(aq) + 2I-(aq) + Pb2+(aq) + 2NO3-(aq)
PbI2(s) + 2Na+(aq) + 2NO3-(aq)
2NaI(aq) + Pb(NO3)2(aq)
PbI2(s) + 2NaNO3(aq)
double displacement reaction (metathesis)
4-11
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Predicting Whether a Precipitate Will Form
1. Note the ions present in the reactants.
2. Consider the possible cation-anion combinations.
3. Decide whether any of the ion combinations is insoluble.
See Table 4.1 (next slide) for solubility rules.
4-12
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Table 4.1 Solubility Rules For Ionic Compounds in Water
Soluble Ionic Compounds
1. All common compounds of Group 1A(1) ions (Li+, Na+, K+, etc.) and
ammonium ion (NH4+) are soluble.
2. All common nitrates (NO3-), acetates (CH3COO- or C2H3O2-) and most
perchlorates (ClO4-) are soluble.
3. All common chlorides (Cl-), bromides (Br-) and iodides (I-) are soluble,
except those of Ag+, Pb2+, Cu+, and Hg22+.
Insoluble Ionic Compounds
1. All common metal hydroxides are insoluble, except those of Group 1A(1)
and the larger members of Group 2A(2)(beginning with Ca2+).
2. All common carbonates (CO32-) and phosphates (PO43-) are insoluble,
except those of Group 1A(1) and NH4+.
3. All common sulfides are insoluble except those of Group 1A(1), Group
2A(2) and NH4+.
4-13
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Sample Problem 4.3
Predicting Whether a Precipitation Reaction
Occurs; Writing Ionic Equations
PROBLEM: Predict whether a reaction occurs when each of the following pairs of
solutions are mixed. If a reaction does occur, write balanced
molecular, total ionic, and net ionic equations, and identify the
spectator ions.
(a) sodium sulfate(aq) + strontium nitrate(aq)
(b) ammonium perchlorate(aq) + sodium bromide(aq)
PLAN:
write ions
SOLUTION:
(a) Na2SO4(aq) + Sr(NO3)2 (aq)
2NaNO3(aq) + SrSO4(s)
2Na+(aq) +SO42-(aq)+ Sr2+(aq)+2NO3-(aq)
2Na+(aq) +2NO3-(aq)+ SrSO4(s)
combine anions & cations
check for insolubility
Table 4.1
eliminate spectator ions
for net ionic equation
4-14
SO42-(aq)+ Sr2+(aq)
(b) NH4ClO4(aq) + NaBr (aq)
SrSO4(s)
NH4Br (aq) + NaClO4(aq)
All reactants and products are soluble so no reaction
occurs.
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Table 4.2 Selected Acids and Bases
Acids
Bases
Strong
Strong
hydrochloric acid, HCl
sodium hydroxide, NaOH
hydrobromic acid, HBr
potassium hydroxide, KOH
hydroiodic acid, HI
calcium hydroxide, Ca(OH)2
nitric acid, HNO3
strontium hydroxide, Sr(OH)2
sulfuric acid, H2SO4
barium hydroxide, Ba(OH)2
perchloric acid, HClO4
Weak
hydrofluoric acid, HF
phosphoric acid, H3PO4
acetic acid, CH3COOH (or
HC2H3O2)
4-15
Weak
ammonia, NH3
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Sample Problem 4.4
PROBLEM:
Writing Ionic Equations for Acid-Base Reactions
Write balanced molecular, total ionic, and net ionic equations for
each of the following acid-base reactions and identify the spectator
ions.
(a) strontium hydroxide(aq) + perchloric acid(aq)
(b) barium hydroxide(aq) + sulfuric acid(aq)
PLAN:
SOLUTION:
reactants are strong acids and (a) Sr(OH)2(aq)+2HClO4(aq) 2H2O(l)+Sr(ClO4)2(aq)
bases and therefore completely
Sr2+(aq) + 2OH-(aq)+ 2H+(aq)+ 2ClO4-(aq)
ionized in water
2H2O(l)+Sr2+(aq)+2ClO4-(aq)
products are
2OH-(aq)+ 2H+(aq)
2H2O(l)
water
spectator ions
4-16
(b) Ba(OH)2(aq) + H2SO4(aq)
2H2O(l) + BaSO4(aq)
Ba2+(aq) + 2OH-(aq)+ 2H+(aq)+ SO42-(aq)
2H2O(l)+Ba2+(aq)+SO42-(aq)
2OH-(aq)+ 2H+(aq)
2H2O(l)
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Figure 4.7
Start of titration
Excess of acid
4-17
An acid-base titration.
Point of
neutralization
Slight excess of
base
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Sample Problem 4.5
PROBLEM:
Finding the Concentration of Acid from an
Acid-Base Titration
You perform an acid-base titration to standardize an HCl solution by
placing 50.00 mL of HCl in a flask with a few drops of indicator
solution. You put 0.1524 M NaOH into the buret, and the initial
reading is 0.55 mL. At the end point, the buret reading is 33.87 mL.
What is the concentration of the HCl solution?
PLAN:
SOLUTION:
NaOH(aq) + HCl(aq)
volume(L) of base
multiply by M of base
mol of base
(33.87-0.55) mL x
NaCl(aq) + H2O(l)
1L
= 0.03332 L
103 mL
molar ratio
mol of acid
divide by L of acid
0.03332 L X 0.1524 M = 5.078x10-3 mol
NaOH
Molar ratio is 1:1
5.078x10-3 mol HCl
M of acid
0.05000 L
4-18
= 0.1016 M HCl
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Figure 4.8
An aqueous strong acid-strong base reaction on the atomic scale.
4-19
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Figure 4.9
4-20
The redox process in compound formation.
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Table 4.3 Rules for Assigning an Oxidation Number (O.N.)
General rules
1. For an atom in its elemental form (Na, O2, Cl2, etc.): O.N. = 0
2. For a monoatomic ion: O.N. = ion charge
3. The sum of O.N. values for the atoms in a compound equals zero. The
sum of O.N. values for the atoms in a polyatomic ion equals the ion’s charge.
Rules for specific atoms or periodic table groups
1. For Group 1A(1):
O.N. = +1 in all compounds
2. For Group 2A(2):
O.N. = +2 in all compounds
3. For hydrogen:
O.N. = +1 in combination with nonmetals
4. For fluorine:
O.N. = -1 in combination with metals and boron
5. For oxygen:
O.N. = -1 in peroxides
O.N. = -2 in all other compounds(except with F)
O.N. = -1 in combination with metals, nonmetals
(except O), and other halogens lower in the group
6. For Group 7A(17):
4-21
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Sample Problem 4.6
PROBLEM:
PLAN:
Determining the Oxidation Number of an Element
Determine the oxidation number (O.N.) of each element in these
compounds:
(a) zinc chloride (b) sulfur trioxide (c) nitric acid
The O.N.s of the ions in a polyatomic ion add up to the charge of the
ion and the O.N.s of the ions in the compound add up to zero.
SOLUTION:
(a) ZnCl2. The O.N. for zinc is +2 and that for chloride is -1.
(b) SO3. Each oxygen is an oxide with an O.N. of -2. Therefore the
O.N. of sulfur must be +6.
(c) HNO3. H has an O.N. of +1 and each oxygen is -2. Therefore
the N must have an O.N. of +5.
4-22
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Figure 4.10
Highest and lowest
oxidation numbers of
reactive main-group
elements.
4-23
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Figure 4.11
A summary of terminology for oxidation-reduction
(redox) reactions.
e-
X
4-24
transfer
Y
or shift of
electrons
X loses electron(s)
Y gains electron(s)
X is oxidized
Y is reduced
X is the reducing agent
Y is the oxidizing agent
X increases its
oxidation number
Y decreases its
oxidation number
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Sample Problem 4.7
Recognizing Oxidizing and Reducing Agents
PROBLEM: Identify the oxidizing agent and reducing agent in each of the following:
(a) 2Al(s) + 3H2SO4(aq)
Al2(SO4)3(aq) + 3H2(g)
(b) PbO(s) + CO(g)
(c) 2H2(g) + O2(g)
Pb(s) + CO2(g)
2H2O(g)
PLAN: Assign an O.N. for each atom and see which atom gained and which atom
lost electrons in going from reactants to products.
An increase in O.N. means the species was oxidized (and is the reducing
agent) and a decrease in O.N. means the species was reduced (is the
oxidizing agent).
SOLUTION:
0
+1 +6 -2
(a) 2Al(s) + 3H2SO4(aq)
+3 +6 -2
0
Al2(SO4)3(aq) + 3H2(g)
The O.N. of Al increases; it is oxidized; it is the reducing agent.
The O.N. of H decreases; it is reduced; H2SO4 is the oxidizing agent.
4-25
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Sample Problem 4.7
continued
+2 -2
Recognizing Oxidizing and Reducing Agents
+2 -2
(b) PbO(s) + CO(g)
0
+4 -2
Pb(s) + CO2(g)
The O.N. of C increases; it is oxidized; CO is the reducing agent.
The O.N. of Pb decreases; it is reduced; PbO is the oxidizing agent.
0
0
(c) 2H2(g) + O2(g)
+1 -2
2H2O(g)
The O.N. of H increases; it is oxidized; it is the reducing agent.
The O.N. of O decreases; it is reduced; it is the oxidizing agent.
4-26
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Figure 4.12
4-27
An active metal displacing hydrogen from water.
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Figure 4.13
4-28
Displacing one metal with another.
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Figure 4.14
The activity series
strength as reducing agents
of the metals.
4-29
Li
K
Ba
Ca
Na
Mg
Al
Mn
An
Cr
Fe
Cd
Co
Ni
Sn
Pb
H2
Cu
Hg
Ag
Au
can displace H
from water
can displace H
from steam
can displace H
from acid
cannot displace H
from any source
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Sample Problem 4.8
PROBLEM:
Identifying the Type of Redox Reaction
Classify each of the following redox reactions as a combination,
decomposition, or displacement reaction, write a balanced
molecular equation for each, as well as total and net ionic
equations for part (c), and identify the oxidizing and reducing
agents:
(a) magnesium(s) + nitrogen(g)
(b) hydrogen peroxide(l)
magnesium nitride (aq)
water(l) + oxygen gas
(c) aluminum(s) + lead(II) nitrate(aq)
aluminum nitrate(aq) + lead(s)
PLAN: Combination reactions produce fewer products than reactants.
Decomposition reactions produce more products than reactants.
Displacement reactions have the same number of products and reactants.
4-30
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Sample Problem 4.8
continued
Identifying the Type of Redox Reaction
0
(a) Combination
3Mg(s) +
0
+2 -3
N2(g)
Mg3N2 (aq)
Mg is the reducing agent; N2 is the oxidizing agent.
+1 -1
(b) Decomposition
+1 -2
0
H2O2(l)
H2O(l) + 1/2 O2(g)
2 H2O2(l)
2 H2O(l) +
or
O2(g)
H2O2 is the oxidizing and reducing agent.
(c) Displacement
0
+2 +5 -2
Al(s) + Pb(NO3)2(aq)
2Al(s) + 3Pb(NO3)2(aq)
+3 +5 -2
Al(NO3)3(aq) + Pb(s)
2Al(NO3)3(aq) + 3Pb(s)
Pb(NO3)2 is the oxidizing and Al is the reducing agent.
4-31
0