Oxidation-Reduction Reactions

1
Oxidation-Reduction Reactions
Chapter 4 and 18
2

• ______________________ 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

_______ half-reaction (____ e-)
_______ half-reaction (____ e-)
3
Review Oxidation a species is oxidized when it
loses one or more electrons, and it is called a
reducing agent Reduction a species is reduced
when it gains one or more electrons, and it is
called an oxidizing agent Oxidation and reduction
always occur together, never in isolation. If
something gains electrons, something else had to
lose them.
4
Oxidation Number
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 ______.

Na, Be, K, Pb, H2, O2, P4 0

1. In monatomic ions, the oxidation number is equal
   to the __________________.

Li, Li 1 Fe3, Fe 3 O2-, O -2

1. The oxidation number of oxygen is usually
   _________. In H2O2 and O22- it is __________.

5
Oxidation Number

1. The oxidation number of hydrogen is ___ except
   when it is bonded to metals in binary compounds.
   In these cases, its oxidation number is ___.
   (LiAlH4)

1. Group IA metals are ___, IIA metals are ___ and
   fluorine is always ___.

6. The sum of the oxidation numbers of all the
atoms in a molecule or ion is equal to
________________________.
6
Oxidation Number
H C O
7
Balancing Redox Equations
Balance the oxidation of Fe2 to Fe3 by Cr2O72-
in acid solution

1. Write the unbalanced equation for the reaction in
   ionic form.

1. Separate the equation into two half-reactions.

Oxidation
Reduction

1. Balance the atoms other than O and H in each
   half-reaction.

8
Balancing Redox Equations

1. For reactions in acid, add H2O to balance O atoms
   and H to balance H atoms.

1. Add electrons to one side of each half-reaction
   to balance the charges on the half-reaction.

1. If necessary, equalize the number of electrons in
   the two half-reactions by multiplying the
   half-reactions by appropriate coefficients.

9
Balancing Redox Equations

1. Add the two half-reactions together and balance
   the final equation by inspection. The number of
   electrons on both sides must cancel.

Oxidation
Reduction

1. Verify that the number of atoms and the charges
   are balanced.

14x1 2 6x2 24 6x3 2x3

1. For reactions in basic solutions, add OH- to both
   sides of the equation for every H that appears
   in the final equation.

10
Electrochemical Cells
_______ __________
_______ __________
spontaneous redox reaction
11
Electrochemical Cells

• The difference in electrical potential between
  the anode and cathode is called
• ____________________
• ____________________
• ____________________

Cell Diagram
Cu2 1 M Zn2 1 M
Zn (s) Zn2 (1 M) Cu2 (1 M) Cu (s)
anode
cathode
12
Standard Electrode Potentials
Zn (s) Zn2 (1 M) H (1 M) H2 (1 atm) Pt
(s)
Anode (oxidation)
Cathode (reduction)
13
Standard Electrode Potentials
____________ ____________ ____________ (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.
Any time you see º, think standard state
conditions
Reduction Reaction
E0 0 V
Standard hydrogen electrode (SHE)
14
Standard Electrode Potentials
Zn (s) Zn2 (1 M) H (1 M) H2 (1 atm) Pt
(s)
15
Standard Electrode Potentials
Pt (s) H2 (1 atm) H (1 M) Cu2 (1 M) Cu
(s)
Anode (oxidation)
Cathode (reduction)
16

• E0 is for the reaction as written
• The more positive E0 the greater the tendency for
  the substance to be reduced
• The more negative E0 the greater the tendency for
  the substance to be oxidized
• Under standard-state conditions, any species on
  the left of a given half-reaction will react
  spontaneously with a species that appears on the
  right of any half-reaction located below it in
  the table (the diagonal rule)

17

• 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

18
Can Sn reduce Zn2 under standard-state
conditions?
How do we find the answer? Look up the Eº
values in Table. Zn2(aq) 2e- gt Zn(s) (Is
this oxidation or reduction?) Which reactions in
the table will reduce Zn2(aq)?
19
What is the standard emf of an electrochemical
cell made of a Cd electrode in a 1.0 M Cd(NO3)2
solution and a Cr electrode in a 1.0 M Cr(NO3)3
solution?
Cd is the stronger oxidizer Cd will oxidize Cr
x 2
Anode (oxidation)
Cathode (reduction)
x 3
20
Spontaneity of Redox Reactions
DG -nFEcell
n number of moles of electrons in reaction
96,500 C/mol
DG0 -RT ln K
21
Spontaneity of Redox Reactions
If you know one, you can calculate the other If
you know K, you can calculate ?Eº and ?Gº If you
know ?Eº, you can calculate ?Gº
22
Spontaneity of Redox Reactions
Relationships among ?G º, K, and Eºcell
23
Oxidation
n ___
Reduction
K ________________
24
2Mg 2Mg2 6e-
Oxidation
n ?
Reduction
25
2Mg 2Mg2 6e-
Oxidation
Reduction
n __
___ V
___ X (96,500 J/V mol) X ___ V
DG0 _______ kJ/mol
26
The Effect of Concentration on Cell Emf
DG DG0 RT ln Q
DG -nFE
-nFE -nFE0 RT ln Q
_____________ equation
At 298K
27
The Nernst equation enables us to calculate E as
a function of reactants and products in a
redox reaction.
28
Oxidation
n ___
Reduction
2e- Fe2 2Fe
E ____________
E ___ 0
________________
29
Batteries
Dry cell
Leclanché cell
Anode
Cathode
30
Batteries
Mercury Battery
Anode
Cathode
31
Batteries
Lead storage battery
Anode
Cathode
32
Batteries
Solid State Lithium Battery
33
Batteries
A ______ ______ is an electrochemical cell that
requires a continuous supply of reactants to keep
functioning
Anode
Cathode
34
Corrosion
35
Cathodic Protection of an Iron Storage Tank
36
_______________ is the process in which
electrical energy is used to cause a
nonspontaneous chemical reaction to occur.
37
Electrolysis of Water
38
Electrolysis and Mass Changes
charge (C) current (A) x time (s)
1 mole e- 96,500 C
So what is the charge on a single electron?