Transcript + H 2 O (l) + 2e

Batteries and
Corrosion
1
Commercial Voltaic Cells
Voltaic Cells are convenient energy sources
Batteries is a self-contained group of voltaic cells arranged in series.
Advantage: Portable
Disadvantage: Very Expensive (.80€ / Kwatt-h)
Need cells in series to provide power
The Processes occurring during the
discharge and recharge of a lead-acid
battery. When the lead-acid battery
is discharging (top) it behaves like a
voltaic cell: the anode is negative
(electrode-1) and the cathode is
positive (electrode-2). When it is
recharging (bottom), it behaves like
an electrolytic cell; the anode is
positive (electrode-2) and the
cathode is negative (electrode-1).
2
Dry Cell or LeClanche Cell
Dry Cells
Invented in the 1860’s the common dry cell or LeClanche cell, has become a familiar household
item. An active zinc anode in the form of a can house a mixture of MnO 2 and an acidic electrolytic
paste, consisting of NH4Cl, ZnCl2, H2O and starch powdered graphite improves conductivity. The
inactive cathode is a graphite rod.
Anode (oxidation)
Zn(s) g Zn2+(aq) = 2eCathode (reduction). The cathodic half-reaction is complex and even today, is still
being studied. MnO2(s) is reduced to Mn2O3(s) through a series of steps that may
involve the presence of Mn2+ and an acid-base reaction between NH4+ and OH- :
2MnO2 (s) + 2NH4+(aq) + 2e- g Mn2O3(s) + 2NH3(aq) + H2O (l)
The ammonia, some of which may be gaseous, forms a complex ion with Zn2+,
which crystallize in contact Cl- ion:
Zn2+(aq) + 2NH3 (aq) + 2Cl-(aq) g Zn(NH3)2Cl2(s)
Overall Cell reaction:
2MnO2 (s) + 2NH4Cl(aq) + Zn(s) g Zn(NH3)2Cl2(s) + H2O (l) + Mn2O3(s)
Ecell = 1.5 V
Uses: common household items, such as portable radios, toys, flashlights,
Advantage; Inexpensive, safe, available in many sizes
Disadvantages: At high current drain, NH3(g) builds up causing drop in voltage,
short shelf life because zinc anode reacts with the acidic NH4+ ions.
3
Alkaline Battery
Alkaline Battery
The alkaline battery is an improved dry cell. The half-reactions are similar, but the
electrolyte is a basic KOH paste, which eliminates the buildup of gases and maintains the
Zn electrode.
Anode (oxidation)
Zn(s) + 2OH- (aq) g ZnO(s) + H2O (l) + 2eCathode (reduction).
2MnO2 (s) + 2H2O (l) + 2e- g Mn(OH)2(s) + 2OH-(aq)
Overall Cell reaction:
2MnO2 (s) + H2O (l) + Zn(s) g ZnO(s) + Mn(OH)2(s)
= 1.5 V
Ecell
Uses: Same as for dry cell.
Advantages: No voltage drop and longer shell life than
dry cell because of alkaline electrolyte; sale ,amu sizes.
Disadvantages; More expensive than common dry cell.
4
Mercury Button Battery
Mercury and Silver batteries are similar.
Like the alkaline dry cell, both of these batteries use zinc in a basic
medium as the anode. The solid reactants are each compressed with
KOH, and moist paper acts as a salt bridge.
Half reactions: E°Cell = 1.6 V
Anode:
Zn(s) + 2OH-(aq) g ZnO(s) + H2O(l) + 2eCathode (Hg): HgO (s) + 2H2O(l) + 2e- g Hg(s) + 2OH-(aq)
Cathode (Ag): Ag2O (s) + H2O(l) + 2e- g 2Ag(s) + 2OH-(aq)
Advantage:
Small, large potential,
silver is nontoxic.
Disadvantage:
Mercury is toxic, silver is
expensive.
5
Lead Storage Battery
Lead-Acid Battery. A typical 12-V lead-acid battery has six cells
connected in series, each of which delivers about 2 V. Each cell
contains two lead grids packed with the electrode material: the
anode is spongy Pb, and the cathode is powered PbO2. The grids
are immersed in an electrolyte solution of 4.5 M H2SO4. Fiberglass
sheets between the grids prevents shorting by accidental physical
contact. When the cell discharges, it generates electrical energy as a
voltaic cell.
Half reactions: E°Cell = 2.0 V
Anode: Pb(s) + SO42- g PbSO4 (s) +2 eCathode (Hg): PbO2 (s) + SO42- + 4H+ + 2e- g
PbSO4 (s) + 2 H2O
Net: PbO2
(s) +
Pb(s) + 2H2SO4 g PbSO4 (s) + 2 H2O
Note hat both half-reaction produce Pb2+ ion, one through
oxidation of Pb, the other through reduction of PbO2. At both
electrodes, the Pb2+ react with SO42- to form PbSO4(s)
6
E° = 0.356
E° = 1.685V
E°Cell = 2.0 V
Nickel-Cadmium Battery
Battery for the Technological Age
Rechargeable, lightweight “ni-cad” are used for variety of cordless appliances.
Main advantage is that the oxidizing and reducing agent can be regenerated
easily when recharged. These produce constant potential.
Half reactions: E°Cell = 1.4 V
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Anode:
Cd(s) + 2OH-(aq) g Cd(OH)2 (s) + 2e-
Cathode:
2Ni(OH) (s) + 2H2O(l) + 2e- g Ni(OH)2 (s) + 2 OH-(aq)
Lithium Ion Batteries
Chemistry
Graphite (-), cobalt or manganese (+)
Nonaqueous electrolyte
Features
+ 40% more capacity than NiCd
+ Flat discharge (like NiCd)
+ Self-discharge 50% less than NiCd
Expensive
8
Fuel Cells
9
Fuel Cells; Batteries
Fuel Cell also an electrochemical device for converting
chemical energy into electricity.
In contrast to storage battery, fuel cell does not need to involve a
reversible reaction since the reactant are supplied to the cell as needed
from an external source. This technology has been used in the Gemini,
Apollo and Space Shuttle program.
Half reactions: E°Cell = 0.9 V
Anode:
2H2 (g) + 4OH-(aq) g 4H2O(l) + 4e-
Cathode:
O2 (g) + 2H2O(l) + 4e- g
Advantage:
Clean, portable and product is water.
Efficient (75%) contrast to 20-25% car,
35-40% from coal electrical plant
Disadvantage:
Cannot store electrical energy, needs
continuous flow of reactant, Electrodes
are short lived and expensive.
10
4OH-(aq)
Corrosion
Not all spontaneous redox reaction are beneficial.
Natural redox process that oxidizes metal to their oxides and
sulfides runs billions of dollars annually. Rust for example is
not the direct product from reaction between iron and oxygen
but arises through a complex electrochemical process.
Rust: Fe2O3 • X H2O
Anode: Fe(s) g Fe+2 + 2eE° = 0.44 V
Cathode:
O2 (g) + 4H+ + 4e- g 2H2O (l)
E° = 1.23 V
Net: Fe+2 will further oxidized to Fe2O3 • X H2O
11
Conditions for Corrosion
Conditions for Iron Oxidation:
Iron will oxidize in acidic medium
SO2 g H2SO4 g H+ + HSO4+
Anions improve conductivity for oxidation.
Cl- from seawater or NaCl (snow melting) enhances rusting
Conditions for Prevention:
Iron will not rust in dry air; moisture must be present
Iron will not rust in air-free water; oxygen must be present
Iron rusts most rapidly in ionic solution and low pH (high H+)
The loss of iron and deposit of rust occur at different placm on object
Iron rust faster in contact with a less active metal (Cu)
Iron rust slower in contact with a more active metal (Zn)
12
Iron Corrosion; Chemistry
Most common and economically destructive
form of corrosion is the rusting of iron. Rust
is not a direct product of the reaction
between iron and oxygen but arises through
complex electrochemical process. The
features of a voltaic cell can help explain this
process.
Iron will not rust in dry air; moisture must be present.
Iron will not rust in air-free water; oxygen must be present
Iron rusts most rapidly in ionic solutions and at low pH (High H+)
The loss of iron and the
depositing of rust often occur at
different places on the same
object.
Iron rust faster in contact with a
less active metal (such as Cu)
and more slowly in contact with
a more active metal (such as Zn).
13
Corrosion Prevention
14
Electroplating
Electroplating is the deposition of a metallic
coating onto the surface of an object by
putting a negative charge onto the object and
immersing it into a solution which contains a
salt of the metal to be deposited.
15
Explanation
Silver electroplating.
The anode is a silver
bar and the cathode is
an iron spoon.
16
The item to be coated is placed into a
container containing a solution of one or
more metal salts. The item is connected
to an electrical circuit, forming the
cathode (negative) of the circuit while an
electrode typically of the same metal to
be plated forms the anode (positive).
When an electrical current is passed
through the circuit, metal ions in the
solution take up excess electrons at the
item. The result is a layer of metal on the
item. However, considerable skill and
craft-technique is required to ensure an
evenly-coated finished product.