Electrochemistry

1
Chapter six
Electrochemistry
2
6-1 Oxidation-reduction Concepts

• Oxidation Numbers
• Oxidation-reduction reaction
• Oxidizing agent and reducing agent
• Redox couple

3
? Oxidation Numbers
- The charge on an atom or a monatomic ion - The
charge that an atom in a substance would have if
the shared pair of electrons belonged to the more
electronegative atom in the bond
4
Rules for Assigning an Oxidation Number (Ox)
General rules
1. For an atom in its elemental form Ox 0
(O2,Cl2, H2, ) (not combine with different
element) 2. For a monatomic ion Ox ion
charge Ca2, 2 Br-, -1 3?? ? Ox.
charge of molecule or ion.
Sum of oxidation states 0 in neutral compounds
Sum of oxidation states charge of the ion
H2SO4, Cr2O72-
5
4. Rules for specific atoms or periodic table
groups
Except with F, K OF2 KO2
(1) For oxygen Ox -2 in most compounds
Ox -1 in peroxides (H2O2)


(2) For hydrogen Ox 1 in combination with
nonmetals
Ox -1 in combination with metals
with electropositive element (i.e., Na, K) H
-1
(3) For fluorine Ox -1 in all compounds
(4) For Group 1A Ox 1 in all compounds
(5) For Group 2A Ox 2 in all compounds
(6) For Group 7A Ox -1 in most compounds
6
Calculation of Oxidation Numbers
7
Example 1
3
-2
Al2S3
Al is a monatomic ion with a 3 charge, so its
oxidation state is 3 (Rule 2).
When combined with metals in binary compounds, S
is a monatomic ion with a 2- charge,
so its oxidation state is -2 (Rule 2).
2 Al at 3 each 6 3 S at -2 each -6
sum 0
Reminder Nonmetals (like sulfur) as well as
metals not in group I or II can have many
oxidation states, so they must be carefully
analyzed.
8
Example 2
1
-2
4
Na2CO3
Na is a monatomic ion with a 1 charge, so its
oxidation state is 1 (Rule 4).
Oxygen has a -2 oxidation state (Rule 1).
2 Na at 1 each 2 3 O at -2 each -6
sum -4
The overall sum must be 0 for a compound (Rule
2), so carbon must have a 4 oxidation state.
9
? Oxidation-reduction (redox) reaction
Redox reaction A reaction in which one or more
electrons are transferred from one atom to
another.
Redox reaction A reaction in which oxidation
numbers of some elements are changed in a
reaction process.
10
2FeCl3 SnCl2 2FeCl2 SnCl4
Redox reaction
2Fe3 Sn2 2Fe2 Sn4
2Mg O2 2MgO
Electron transfer Cu(s) 2 Ag(aq) ? Cu2(aq)
2 Ag(s)
Loss of Electrons OXIDATION (LEO)
Gain of Electrons REDUCTION (GER)
11
Half- reaction
Oxidation reaction
- loss of electrons
- oxidation number increases
Redox reaction
A g ?A g e-
Reduction reaction
Half- reaction
- gain in electrons
- decrease in oxidation number
Fe2 2e- ?Fe
Redox process always occurs together. In redox
process, one cant occur without the other.
12
(No Transcript)
13
Oxidation
Reduction
14
? Oxidizing agent and reducing agent
Oxidizing agent
- electron acceptor species is reduced.
Reducing agent
- electron donor, species is oxidized.
15

• An oxidizing agent is a species that oxidizes
  another species it is itself reduced.
• A reducing agent is a species that reduces
  another species it is itself oxidized.

Loss of 2 e-1 oxidation
reducing agent
oxidizing agent
Gain of 2 e-1 reduction
16
Reducing agent
Losses one or more electrons
Causes reduction
Undergoes oxidation
Becomes more positive or less negative
Oxidizing agent
Gains one or more electrons
Causes oxidation
Undergoes reduction
Becomes more negative or less positive
17
Example
CH4 (g) 2 O2 (g) CO2 (g) 2H2O
(l) C H O -4 1 0 4 1 -2 Which species
is oxidized ? (lost electrons/Ox state became
more positive) Which species is reduced ?
(gained electrons/Ox state became more negative)
18
? Redox couple (Pair of Electrons Oxidation and
Reduction)
The overall reaction is split into two
half-reactions, one involving oxidation and one
reduction.
8H MnO4- 5Fe2 Mn2 5Fe3
4H2O Reduction 8H MnO4- 5e- Mn2
4H2O Oxidation 5Fe2 5Fe3 5e-
19
2Fe3 Sn2 2Fe2 Sn4
Reduction Fe3 e- ?Fe2
Oxidation Sn2 - 2e- ?Sn4
Redox couples
The oxidized and reduced states of each substance
taking part in a half-reaction form a redox
couple.
Notation of redox couple
Oxidized species / reduced species
Sn4 / Sn2 Fe3 / Fe2
20
Redox couple half-reactions
Oxidized state n e -
Reduced state
Ox n e -
Red
21
6-2 Voltaic Cells (Primary Cells)
Electron transfer Zn(s) CuSO4 (aq) ? ZnSO4
(aq) Cu(s)
Loss of Electrons OXIDATION (LEO)
Gain of Electrons REDUCTION (GER)
22
What could happen when we put a piece of zinc
metal into the solution of copper sulfate?
23

• Copper is deposited on the zinc
• The copper plates out onto the
• surface of the zinc metal
• The blue copper (?)ions are gradually replaced by
  colorless zinc ions
• Chemical energy (reaction enthalpy)
• is released as heat

24
CHEMICAL CHANGE ? ELECTRIC CURRENT
With time, Cu plates out onto Zn metal strip, and
Zn strip disappears.
Zn is oxidized and is the reducing agent Zn(s)
? Zn2(aq) 2e- Cu2 is reduced and is the
oxidizing agent Cu2(aq) 2e- ? Cu(s)
25
CHEMICAL CHANGE ? ELECTRIC CURRENT (2)
Oxidation Zn(s) ? Zn2(a q) 2e- Reduction
Cu2(a q) 2e- ? Cu(s) -----------------------
----------------------------- Zn(s) Cu2(aq) ?
Zn2(aq) Cu(s)
26
CHEMICAL CHANGE ? ELECTRIC CURRENT (2)
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 VOLTAIC cell.
• (also called GALVANIC cell)
• A group of such cells is called a battery.

27
Voltaic Cell
A device in which chemical energy is changed to
electrical energy.
6-2.1 Cu- Zn Primary Cell
28


()
(-)
29
ANODE (-)
CATHODE ()
Negative electrode generates electron Oxidation
Occur
Positive electrode accepts electron Reduction
Occur
Zn half-cell
Cu half-cell
Zn2 / Zn
Cu2 / Cu
Electrons travel thru external wire. Salt
bridge allows anions and cations to move
between electrode compartments. This maintains
electrical neutrality.
30
Voltaic Cells
A voltaic cell consists of two half-cells.

• Each half-cell is a portion of the
  electrochemical cell in which a half-reaction
  takes place.

• A simple half-cell can be made from a metal strip
  dipped into a solution of its metal ion.
• For example, the zinc-zinc ion half cell consists
  of a zinc strip dipped into a solution of a zinc
  salt.

31

• Another simple half-cell consists of a copper
  strip dipped into a solution of a copper salt.

• In a voltaic cell, two half-cells are connected
  in such a way that electrons flow from one metal
  electrode to the other through an external
  circuit.

32
As long as there is an external circuit,
electrons can flow through it from one electrode
to the other.

• Because zinc has a greater tendency to lose
  electrons than copper, zinc atoms in the zinc
  electrode lose electrons to form zinc ions.

• The electrons flow through the external circuit
  to the copper electrode where copper ions gain
  the electrons to become copper metal.

33
The two half-cells must also be connected
internally to allow ions to flow between them.

• Without this internal connection, too much
  positive charge builds up in the zinc half-cell
  (and too much negative charge in the copper
  half-cell) causing the reaction to stop.
• Figure A and B show the two half-cells of a
  voltaic cell connected by salt bridge.

• A salt bridge is a U shape tube of an electrolyte
  in a gel that is connected to the two half-cells
  of a voltaic cell.

34

• The salt bridge allows the flow of ions but
  prevents the mixing of the different solutions
  that would allow direct reaction of the cell
  reactants.

35
6-2.2 Cell Reaction
The two half-cell reactions, as noted earlier,
are
oxidation half-reaction
reduction half-reaction
36
Note that the sum of the two half-reactions
is the net reaction that occurs in the voltaic
cell it is called the cell reaction.

• Note that electrons are given up at the anode and
  thus flow from it to the cathode where reduction
  occurs.

37
6-2.3 Notation for Voltaic Cells
It is convenient to have a shorthand way of
designating particular voltaic cells.
(-)
()

• The anode (oxidation half-cell) is written on the
  left. The cathode (reduction half-cell) is
  written on the right.

38

• A boundary between different phases (e.g., an
  electrode and a solution) is represented by a
  single vertical line ()

• The anode in a voltaic cell has a negative
• sign because electrons flow from it.

• The cathode in a voltaic cell has a positive sign

39
Notation for Voltaic Cells

• The two electrodes are connected by a salt
  bridge, denoted by two vertical bars.

40
Notation for Voltaic Cells
salt bridge

• The cell terminals are at the extreme ends in the
  cell notation.

41
Notation for Voltaic Cells
salt bridge

• A single vertical bar indicates a phase boundary,
  such as between a solid terminal and the
  electrode solution.

42
2Fe3(c1) Sn2(c2) 2Fe2(c3) Sn4(c4)
Fe3 e
Fe2
Half reaction
Sn2 - 2e
Sn4
Sn4 / Sn2
PtSn4 (c4), Sn2(c2)
PtFe3 (c1), Fe2(c3)
Fe3 / Fe2
Cell notation
PtSn4 (c4), Sn2(c2)
Fe3 (c1), Fe2(c3) Pt
43
When the half-reaction involves a gas, an inert
material such as platinum serves as a terminal
and an electrode surface on which the reaction
occurs.

• hydrogen electrode

44

• The notation for the hydrogen electrode, written
  as a cathode, is

• To write such an electrode as an anode, you
  simply reverse the notation.

45
To fully specify a voltaic cell, it is necessary
to give the concentrations of solutions and the
pressure of gases.
46
Line Notation
solid½Aqueous½½Aqueous½solid Anode on the
left½½Cathode on the right Single line different
phases. Double line salt bridge. If all the
substances on one side are aqueous, a platinum
electrode is indicated. Cu(s)½Cu2(aq)½½Fe2(aq),
Fe3(aq)½Pt(s)
47
6-2.4 Electromotive Force

• The maximum potential difference between the
  electrodes of a voltaic cell is referred to as
  the electromotive force (emf) of the cell,
  denoted E
• E fcathode fanode f -f-
• E is a positive number.

48
The standard emf, E o, is the emf of a cell
operating under standard conditions of
concentration (1 M), pressure (1atm), and
temperature (25 oC).
49
Standard Notation for Electrochemical Cells
ANODE Zn / Zn2 // Cu2 / Cu
CATHODE
REDUCTION
OXIDATION
50
Anode and Cathode

• OXIDATION occurs at the ANODE.
• REDUCTION occurs at the CATHODE.
• Mnemonic O and A are vowels R and C are
  consonants

51
6-2.5 Types of Electrodes
(a) metal/metal ion electrode (b) metal/
insoluble salt electrode (c) gas electrode (d)
redox electrode
52
Types of Electrode (continued)
53
Types of Electrode (continued)
l
l
l
l
l
54

• The types of electrode

Metal-metal ion electrode Zn(s)Zn2( aq
) Cu(s ) Cu2 ( aq ) MMn
Mn ne M Gas electrode
PtH2(p)H(c) 2H 2e H2 Metal-insoluble
salt electrodes PtHg(l)Hg2Cl2(s)Cl- (c)
Hg2Cl2 (s) 2e 2Hg 2Cl- Oxidation-reduction
electrodes Fe3 e Fe2
PtFe3(c1),Fe2(c2)