Download Chapter 20 Oxidation-Reduction Reactions

Redox!
a.k.a. Oxidation-Reduction
a.k.a Electrochemistry
Review of Oxidation numbers
The charge the atom would have in a molecule (or an
ionic compound) if electrons were completely transferred.
1. Free elements (uncombined state) have an oxidation
number of zero.
Na, Be, K, Pb, H2, O2, P4 = 0
2. In monatomic ions, the oxidation number is equal to
the charge on the ion.
Li+, Li = +1; Fe3+, Fe = +3; O2-, O = -2
3. The oxidation number of oxygen is usually –2. In H2O2
it is –1.
4.4
Rules for Assigning Oxidation Numbers
1) The oxidation number of any
uncombined element is zero.
2) The oxidation number of a
monatomic ion equals its charge.
0
0
1
1
2 Na  Cl 2  2 Na Cl
Rules for Assigning Oxidation Numbers
3) The oxidation number of oxygen in
compounds is -2, except in peroxides,
such as H2O2 where it is -1.
4) The oxidation number of hydrogen in
compounds is +1, except in metal
hydrides, like NaH, where it is -1.
1
2
H2O
Rules for Assigning Oxidation Numbers
5) The sum of the oxidation numbers of the
atoms in the compound must equal 0.
1
2
H2O
2(+1) + (-2) = 0
H
O
2
2 1
Ca(O H ) 2
(+2) + 2(-2) + 2(+1) = 0
Ca
O
H
Rules for Assigning Oxidation Numbers
6) The sum of the oxidation numbers in
the formula of a polyatomic ion is equal
to its ionic charge.
? 2
N O3

? 2
S O4
2
X + 3(-2) = -1
N
O
X + 4(-2) = -2
S
O
 X = +5
 X = +6
Electron Transfer Reactions
• Electron transfer reactions are oxidationreduction or redox reactions.
• Results in the generation of an electric
current (electricity) or be caused by
imposing an electric current.
• Therefore, this field of chemistry is often
called ELECTROCHEMISTRY.
Electrochemical processes are oxidation-reduction
reactions in which:
•
the energy released by a spontaneous reaction is
converted to electricity or
•
electrical energy is used to cause a nonspontaneous
reaction to occur
0
0
2Mg (s) + O2 (g)
2Mg
2+ 2-
2MgO (s)
2Mg2+ + 4e- Oxidation half-reaction (lose e-)
O2 + 4e-
2O2- Reduction half-reaction (gain e-)
19.1
Terminology for Redox Reactions
• OXIDATION—loss of electron(s) by a species;
increase in oxidation number; increase in oxygen.
• REDUCTION—gain of electron(s); decrease in
oxidation number; decrease in oxygen; increase in
hydrogen.
• OXIDIZING AGENT—electron acceptor; species is
reduced. (an agent facilitates something; ex. Travel
agents don’t travel, they facilitate travel)
• REDUCING AGENT—electron donor; species is
oxidized.
You can’t have one… without the other!
• Reduction (gaining electrons) can’t happen
without an oxidation to provide the electrons.
• You can’t have 2 oxidations or 2 reductions in
the same equation. Reduction has to occur at
the cost of oxidation
LEO the lion says GER!
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GER!
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Another way to remember
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x s o
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Oxidation and Reduction (Redox)
Redox currently says that electrons
are transferred between reactants
•
Mg
+
S→
Mg2+
+
S2-
•The magnesium atom changes to a magnesium ion by losing 2
electrons, and is thus oxidized
•The sulfur atom is changed to a sulfide ion by gaining 2
electrons, and is thus reduced.
Oxidation and Reduction (Redox)
0
1
0
1
2 Na  Cl 2  2 Na Cl
Each sodium atom loses one electron:
1
0
Na  Na  e

Each chlorine atom gains one electron:
0

1
Cl  e  Cl
LEO says GER :
Lose Electrons = Oxidation
1
0
Na  Na  e

Sodium is oxidized
Gain Electrons = Reduction
0

1
Cl  e  Cl
Chlorine is reduced
LEO says GER :
- Losing electrons is oxidation, and the
substance that loses the electrons is
called the reducing agent.
- Gaining electrons is reduction, and
the substance that gains the electrons is
called the oxidizing agent.
Mg is the
reducing
agent
Mg is oxidized – loses e-
Mg(s) + S(s) → MgS(s)
S is the oxidizing agent
S is reduced – gains e-
Not All Reactions are Redox Reactions
- Reactions in which there has been no
change in oxidation number are not
redox reactions.
Examples:
1 5 2
1
1
1
1
1 5 2
Ag N O 3 (aq)  Na Cl (aq )  Ag Cl ( s )  Na N O 3 (aq )
1 2 1
1
6 2
1
6 2
1
2
2 Na O H (aq )  H 2 S O 4 (aq )   Na 2 S O 4 (aq)  H 2 O(l )
Reducing Agents and Oxidizing Agents
• An increase in oxidation number = oxidation
• A decrease in oxidation number = reduction
1
0
Na  Na  e

Sodium is oxidized – it is the reducing agent
0

1
Cl  e  Cl
Chlorine is reduced – it is the oxidizing agent
Trends in Oxidation and Reduction
Active metals:
Lose electrons easily
Are easily oxidized
Are strong reducing agents
Active nonmetals:
Gain electrons easily
Are easily reduced
Are strong oxidizing agents
Identifying Redox Equations
In general, all chemical reactions can
be assigned to one of two classes:
1) oxidation-reduction, in which
electrons are transferred:
• Single-replacement, combination,
decomposition, and combustion
2) this second class has no electron
transfer, and includes all others:
• Double-replacement and acidbase reactions
Identifying Redox Equations
In an electrical storm, oxygen and
nitrogen react to form nitrogen monoxide:
N2(g) + O2(g) → 2NO(g)
YES!
•Is this a redox reaction?
•If the oxidation number of an element
in a reacting species changes, then
that element has undergone either
oxidation or reduction; therefore, the
reaction as a whole must be a redox.
Balancing Redox Equations
It is essential to write a correctly
balanced equation that represents
what happens in a chemical reaction
• Fortunately, two systematic methods
are available, and are based on the
fact that the total electrons gained in
reduction equals the total lost in
oxidation. The two methods:
1) Use oxidation-number changes
2) Use half-reactions
Using half-reactions
A half-reaction is an equation showing
just the oxidation or just the reduction that
takes place
they are then balanced separately, and
finally combined
Step 1: write unbalanced equation in ionic
form
Step 2: write separate half-reaction
equations for oxidation and reduction
Step 3: balance the atoms in the halfreactions
Using half-reactions
continued
•Step 4: add enough electrons to one side
of each half-reaction to balance the charges
•Step 5: multiply each half-reaction by a
number to make the electrons equal in both
•Step 6: add the balanced half-reactions to
show an overall equation
•Step 7: add the spectator ions and balance
the equation
CHEMICAL CHANGE --->
ELECTRIC CURRENT
•To obtain a useful
current, we separate the
oxidizing and reducing
agents so that electron
transfer occurs thru an
external wire.
This is accomplished in a GALVANIC or
VOLTAIC cell. http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/galvan5.swf
A group of such cells is called a battery.
Galvanic Cells
anode
oxidation
cathode
reduction
-
+
spontaneous
redox reaction
19.2
Galvanic Cells
The difference in electrical
potential between the anode
and cathode is called:
• cell voltage
• electromotive force (emf)
• cell potential
Cell Diagram
Zn (s) + Cu2+ (aq)
Cu (s) + Zn2+ (aq)
[Cu2+] = 1 M & [Zn2+] = 1 M
Zn (s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu (s)
anode
cathode
19.2
Standard Electrode Potentials
Zn (s) | Zn2+ (1 M) || H+ (1 M) | H2 (1 atm) | Pt (s)
Anode (oxidation):
Zn (s)
Cathode (reduction): 2e- + 2H+ (1 M)
Zn (s) + 2H+ (1 M)
Zn2+ (1 M) + 2eH2 (1 atm)
Zn2+ + H2 (1 atm)
19.3
Standard Electrode Potentials
Standard reduction potential (E0) is the voltage associated
with a reduction reaction at an electrode when all solutes
are 1 M and all gases are at 1 atm.
Reduction Reaction
2e- + 2H+ (1 M)
H2 (1 atm)
E0 = 0 V
Standard hydrogen electrode (SHE)
19.3
•
E0 is for the reaction as
written
•
The more positive E0 the
greater the tendency for the
substance to be reduced
•
The half-cell reactions are
reversible
•
The sign of E0 changes
when the reaction is
reversed
•
Changing the stoichiometric
coefficients of a half-cell
reaction does not change
the value of E0
19.3