John C. Kotz
Paul M. Treichel
John Townsend
http://academic.cengage.com/kotz
Chapter 3
Chemical Reactions
John C. Kotz • State University of New York, College at Oneonta
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3
CHEMICAL REACTIONS
Chapter 3
Reactants: Zn + I2
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Product: ZnI2
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Chemical Equations
Depict the kind of reactants and
products and their relative amounts in
a reaction.
4 Al(s) + 3 O2(g) f 2 Al2O3(s)
The numbers in the front are called
stoichiometric coefficients
The letters (s), (g), and (s) are the
physical states of compounds.
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Reaction of Phosphorus with Cl2
Notice the stoichiometric coefficients and the
physical states of the reactants and products.
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Reaction of Iron with Cl2
Notice the stoichiometric coefficients and the
physical states of the reactants and products.
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Chemical Equations
4 Al(s) + 3 O2(g)
f 2 Al2O3(s)
This equation means
4 Al atoms + 3 O2 molecules
f
2 “molecules” of Al2O3
4 moles of Al + 3 moles of O2
f
2 moles of Al2O3
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Chemical Equations
• Because the same
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atoms are present in a
reaction at the
beginning and at the
end, the amount of
matter in a system
does not change.
• The Law of the
Conservation of
2HgO(s) f 2 Hg(liq) + O2(g)
Matter
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Chemical Equations
Because of the principle of
the conservation of
matter,
an equation must be
balanced.
It must have the same
number of atoms of the
same kind on both sides.
Lavoisier, 1788
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PLAY MOVIE
Balancing
Equations
___ Al(s) + ___ Br2(s) f ___ Al2Br6(s)
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Balancing
Equations
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____C3H8(g) + _____ O2(g) f
_____CO2(g) + _____ H2O(g)
____B4H10(g) + _____ O2(g) f
___
B2O3(g) + _____ H2O(g)
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Chemical Equilibrium
• Chemical reactions are reversible.
• Ammonia can be produced from the elements in
the Haber process
N2(g) + 3 H2(g) f 2 NH3(g)
• But NH3 can also be decomposed to the elements
2 NH3(g) f N2(g) + 3 H2(g)
• In a process to make NH3, the reaction can come
eventually to equilbrium.
N2(g) + 3 H2(g) e 2 NH3(g)
• Double arrows indicate equilibrium
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Reaction Reversibility
Stalactites and
stalagmites in caves
depend on a reversible
chemical reaction
Ca2+(aq) + 2 HCO3–(aq)
e
CaCO3(s) + CO2(g) + H2O(s)
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Reaction Reversibility
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Chemical Equilibrium
Once equilibrium is achieved, reaction
continues, but there is no net change in
amounts of products or reactants.
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Reactions
in Aqueous Solution
Many reactions involve ionic compounds, especially
reactions in water — aqueous solutions.
KMnO4 in water
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K+(aq) + MnO4-(aq)
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An Ionic Compound, CuCl2, in Water
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Aqueous Solutions
How do we know ions are
present in aqueous solutions?
The solutions conduct
electricity!
They are called
ELECTROLYTES
HCl, CuCl2, and NaCl are
strong electrolytes.
They dissociate completely (or
nearly so) into ions.
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Aqueous Solutions
HCl, CuCl2, and NaCl are strong
electrolytes. They dissociate completely
(or nearly so) into ions.
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Aqueous Solutions
Acetic acid ionizes only to a small extent, so it is a weak
electrolyte.
CH3CO2H(aq) e CH CO
3
2 (aq)
+ H+(aq)
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Aqueous Solutions
Acetic acid ionizes only to a small
extent, so it is a weak
electrolyte.
CH3CO2H(aq)
e
CH3CO2-(aq) + H+(aq)
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Aqueous Solutions
Some compounds
dissolve in water but
do not conduct
electricity. They are
called nonelectrolytes.
Examples include:
sugar
ethanol
ethylene glycol
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Water Solubility of Ionic Compounds
If one ion from the “Soluble
Compd.” list is present in a
compound, the compound is
water soluble.
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Water Solubility
of Ionic Compounds
Common minerals are often formed with anions that
lead to insolubility:
sulfide
fluoride
carbonate
oxide
Iron pyrite, a sulfide
Azurite, a copper
carbonate
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Orpiment,
arsenic sulfide
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Chemical Reactions in Water
We will look at
EXCHANGE
REACTIONS
AX + BY
AY + BX
The anions
exchange places
between cations.
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Pb(NO3) 2(aq) + 2 KI(aq)
f PbI2(s) + 2 KNO3 (aq)
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Precipitation Reactions
The “driving force” is the formation
of an insoluble compound — a
precipitate.
Pb(NO3)2(aq) + 2 KI(aq) f
2 KNO3(aq) + PbI2(s)
BaCl2(aq) + Na2SO4(aq) f
BaSO4(s) + 2 NaCl(aq)
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Net Ionic Equations
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Pb(NO3)2(aq) + K2CrO4(aq)
f PbCrO4(s) + 2 KNO3(aq)
This is the “complete equation”
Because Pb(NO3)2 and K2CrO4 are strong electrolytes we should
write
Pb2+(aq) + 2 NO3-(aq) + 2 K+(aq) + CrO42-(aq)
f PbCrO4(s) + 2 K+(aq) + 2 NO3-(aq)
This is the “ionic equation”
Question: do we need to include the K+ and NO3- ions?
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Net Ionic Equations
Ionic equation:
Pb2+(aq) + 2 NO3-(aq) + 2 K+(aq) + CrO42-(aq)
f PbCrO4(s) + 2 K+(aq) + 2 NO3-(aq)
The NO3- and K+ ions are SPECTATOR IONS — they
do not participate. Could have used Na+ instead of K+.
We leave the spectator ions out —
Pb2+(aq) + CrO42-(aq) f PbCrO4(s)
to give the NET IONIC EQUATION
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ACIDS
An acid f H3O+ in water
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The
Nature of
Acids
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ACIDS
An acid f H3O+ in water
Some strong acids are
HCl
H2SO4
HClO4
HNO3
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hydrochloric
sulfuric
perchloric
nitric
HNO3
Weak Acids
WEAK ACIDS = weak
electrolytes
CH3CO2H
acetic acid
H2CO3
carbonic acid
H3PO4
phosphoric acid
HF
hydrofluoric acid
Acetic acid
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ACIDS
Nonmetal oxides can be acids
CO2(aq) + H2O(s)
f H2CO3(aq)
SO3(aq) + H2O(s)
f H2SO4(aq)
and can come from burning coal and
oil.
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BASES
Table 3.2
Base f OH- in water
NaOH(aq) f Na+(aq) + OH-(aq)
NaOH is
a strong
base
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Ammonia, NH3
An Important Base
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BASES
Metal oxides are bases
CaO(s) + H2O(s)
f Ca(OH)2(aq)
CaO in water. Indicator
shows solution is basic.
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Know the strong
acids & bases!
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Acid-Base Reactions
• The “driving force” is the formation of water.
NaOH(aq) + HCl(aq) f
NaCl(aq) + H2O(liq)
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• Net ionic equation
OH-(aq) + H3O+(aq)
f 2 H2O(s)
• This applies to ALL reactions
of STRONG acids and bases.
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See Active Figure 3.14
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Acid-Base Reactions
• A-B reactions are sometimes called
NEUTRALIZATIONS because the solution is
neither acidic nor basic at the end.
• The other product of the A-B reaction is a SALT, MX.
HX + MOH f MX + H2O
Mn+ comes from base & Xn- comes from acid
This is one way to make compounds!
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Gas-Forming Reactions
This is primarily the chemistry of metal carbonates.
CO2 and water f H2CO3
H2CO3(aq) + Ca2+ f
2 H+(aq) + CaCO3(s) (limestone)
Adding acid reverses this reaction.
MCO3 + acid f CO2 + salt
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Gas-Forming
Reactions
PLAY MOVIE
CaCO3(s) + H2SO4(aq) f
2 CaSO4(s) + H2CO3(aq)
Carbonic acid is unstable and forms CO2 & H2O
H2CO3(aq) f CO2 + water
(Antacid tablet has citric acid + NaHCO3)
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Oxidation-Reduction Reactions
Section 3.9
Thermite reaction
Fe2O3(s) + 2 Al(s)
f
2 Fe(s) + Al2O3(s)
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EXCHANGE: Precipitation Reactions
EXCHANGE
Gas-Forming
Reactions
REACTIONS
REDOX
REACTIONS
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EXCHANGE
Acid-Base
Reactions
REDOX REACTIONS
REDOX = reduction & oxidation
O2(g) + 2 H2(g) f 2 H2O(s)
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REDOX
REACTIONS
REDOX = reduction & oxidation
Corrosion of aluminum
2 Al(s) + 3 Cu2+(aq) f 2 Al3+(aq) + 3 Cu(s)
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REDOX REACTIONS
Cu(s) + 2 Ag+(aq) f Cu2+(aq) + 2 Ag(s)
In all reactions if
something has been
oxidized then
something has also
been reduced
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REDOX REACTIONS
Cu(s) + 2 Ag+(aq) f Cu2+(aq) + 2 Ag(s)
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Why Study Redox Reactions
Batteries
Corrosion
Fuels
Manufacturing metals
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REDOX REACTIONS
Redox reactions are characterized by
ELECTRON TRANSFER between an electron
donor and electron acceptor.
Transfer leads to—
1.
increase in oxidation number of
some element = OXIDATION
2. decrease in oxidation number of some
element = REDUCTION
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OXIDATION NUMBERS
The electric charge an element APPEARS to
have when electrons are counted by some
arbitrary rules:
1.
2.
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Each atom in free element has ox. no. = 0.
Zn
O2
I2
S8
In simple ions, ox. no. = charge on ion.
-1 for Cl+2 for Mg2+
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OXIDATION NUMBERS
3.
O has ox. no. = -2
(except in peroxides: in H2O2, O = -1)
4.
Ox. no. of H = +1
(except when H is associated with a metal
as in NaH where it is -1)
5.
Algebraic sum of oxidation numbers
= 0 for a compound
= overall charge for an ion
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OXIDATION NUMBERS
NH3
N =
ClO-
Cl =
H3PO4
P =
MnO4-
Mn =
Cr2O72-
Cr =
C3H8
C =
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Oxidation
number of F
in HF?
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Recognizing a Redox Reaction
Corrosion of aluminum
2 Al(s) + 3 Cu2+(aq) f 2 Al3+(aq) + 3 Cu(s)
Al(s) f Al3+(aq) + 3 e• Ox. no. of Al increases as e- are donated by the metal.
• Therefore, Al is OXIDIZED
• Al is the REDUCING AGENT in this balanced halfreaction.
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Recognizing a Redox Reaction
Corrosion of aluminum
2 Al(s) + 3 Cu2+(aq) f 2 Al3+(aq) + 3 Cu(s)
Cu2+(aq) + 2 e- f Cu(s)
• Ox. no. of Cu decreases as e- are accepted by the ion.
• Therefore, Cu is REDUCED
• Cu is the OXIDIZING AGENT in this balanced halfreaction.
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Recognizing a Redox Reaction
Notice that the 2 half-reactions add up to
give the overall reaction
—if we use 2 mol of Al and 3 mol of Cu2+.
2 Al(s) f 2 Al3+(aq) + 6 e3 Cu2+(aq) + 6 e- f 3 Cu(s)
----------------------------------------------------------2 Al(s) + 3 Cu2+(aq) f 2 Al3+(aq) + 3 Cu(s)
Final eqn. is balanced for mass and charge.
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Examples of Redox Reactions
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Metal + halogen
2 Al + 3 Br2 f Al2Br6
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Examples of Redox Reactions
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Nonmetal (P) + Oxygen
f P4O10
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Metal (Mg) + Oxygen
f MgO
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Recognizing a Redox
Reaction
See Table 3.4
Reaction Type
Oxidation
In terms of oxygen
gain
loss
In terms of halogen
gain
loss
In terms of electrons
loss
gain
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Reduction
Common Oxidizing and
Reducing Agents
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See Table 3.4
Metals
(Cu) are
reducing
agents
HNO3 is an
oxidizing
agent
Cu + HNO3 f
Cu2+ + NO2
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Metals
(Na, K,
Mg, Fe)
are
reducing
agents
2 K + 2 H 2O f 2
KOH + H2
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Examples of Redox Reactions
Metal + acid
Mg + HCl
Mg = reducing agent
H+ = oxidizing agent
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Metal + acid
Cu + HNO3
Cu = reducing agent
HNO3 = oxidizing agent