Transcript Electrochemistry

Electricity from
Chemical
Reactions
Electrochemistry
• The production of electrical energy from
chemical reactions
• Redox reactions involve the transfer of
electrons
• Redox means that reduction and oxidation
are occurring simultaneously
Reduction
• Occurs when there is a decrease in
oxidation number Zn2+  Zn
• Gains electrons
• Loses Oxygen
• Converting a complex substance into a
simpler form i.e. smelting iron to produce
the pure metal iron
Oxidation
• Occurs where there is an increase in
oxidation number Zn  Zn2+
• Loses electrons
• Gains oxygen
• The reaction used to describe the reaction of
any substance with oxygen
Determining Oxidation Numbers
• The atoms in elements have an Oxidation
Number of zero eg Fe, C, Cl2
• For a neutral molecule, the sum of the
oxidation numbers are zero eg CO2
• For a monatomic ion, the oxidation number
is the same as it’s charge Cl – , Na+
Determining Oxidation Numbers
• Oxygen usually takes – 2 in compounds. In
peroxides (H2O2 & BaO2) it is – 1
• Hydrogen takes + 1 in compounds, except
in hydrides (NaH, CaH2) where it takes – 1
Determining Oxidation Numbers
• For a polyatomic ion, the sum of the
oxidation numbers of its component atoms
is the same as its charges
• For polyatomic molecules or ions, the, most
electronegative element has a negative
oxidation number and the least
electronegative element has a positive
oxidation number
Redox Half Reactions
• Consider the reaction when a strip of zinc is
dropped in a solution of Copper Sulphate
• Zn(s) + Cu 2+(aq)  Zn2+(aq) + Cu(s)
• Electrons are transferred from zinc atoms to
copper ions
• Reaction occurs spontaneously, that is with
no external force or energy being applied
Redox Half Reactions
• Redox reactions consist of two half
reactions
• Oxidation Zn(s)  Zn2+(aq) + 2e–1
• Reduction Cu 2+(aq) + 2e–1  Cu(s)
• It is possible to use redox reactions to
produce electricity
Galvanic Cells
• Also called Electrochemical Cells
• Achieved by separating the half equations
into half cells
• Transferred electrons are forced to pass
through an external circuit
• Such an apparatus is called a Galvanic Cell
Galvanic Cells
zinc
–
Flow of electrons
+
copper
Salt
bridge
2+
Zn
Negative Electrode
(ANODE)
Cu
2+
Positive Electrode
(CATHODE)
Standard Electrode Potentials
• The electrical potential of a galvanic cell is
the ability of the cell to produce an electric
current.
• Electrical potential is measured in volts
• Cannot measure the electrode potential of
an isolated half cell
• Can measure the difference in in potential
between two connected half cells
Standard Electrode Potentials
• Electrical potential of a cell results from
competition between 2 half cells for
electrons
• Half cell with the greatest tendency to
attract electrons will undergo REDUCTION
• Other half cell will lose electrons and
undergo OXIDATION
Standard Electrode Potentials
• The Reduction Potential of a half cell is a
measure of the tendency of the oxidant to
accept electrons and so undergo reduction
• The difference between the reduction
potentials of the two half cells is called the
Cell Potential Difference
Standard Electrode Potentials
• The Standard Cell Potential Difference
(E0 cell) is the measured cell potential
difference when the concentration of each
species = 1M, pressure = 1 atm and Temp =
25 C
• E0 cell = E0 oxidant – E0 reductant
Standard Electrode Potentials
• A Standard Hydrogen Half cell is used as a
comparative measure the reduction
potentials of other cells
• The SHE is given a value of 0.00 V
• All other half cells are given a reduction
potential value in comparison to this SHE
by being connected to it
Standard Hydrogen Electrode
Glass
sleeve
Platinum wire
Salt Bridge to
Other half-cell
1.00M
Acid solution
Platinum electrode
H2 gas
(1 Atm)
Standard Hydrogen Electrode
• SHE is used to measure reduction potential
of other cells
• If a species accepts electrons more readily
than hydrogen, its electrode potential is
positive
• If a species accepts electrons less readily
than hydrogen, its electrode potential is
negative
Electrochemical Series
• The reaction that is higher on the
electrochemical series will occur as it
appears and will reverse the direction of the
reaction that occurs lower on the table
Potential Difference
• Is measured by a volt meter
• Can be estimated by using electrochemical
series
• Connect Mg2+/Mg and Cl2/Cl– half cells
get a potential difference of 3.7V
• Looking at the electrochemical series
Potential Difference
•
•
•
•
Cl2 + 2e–  Cl– has an E0 of 1.36V
–
0
2+
Mg + 2e  Mg has an E of – 2.38V
The potential difference can be calculated
1.36 – (– 2.38) = 3.74V
Galvanic Cells
• Primary Cells
– Produce energy until one component is used up,
then discarded
• Secondary Cells
– Store energy and may be recharged
Primary Cells
• Dry Cells
• Alkaline Cells
• Button Cells
Dry Cells
• The ordinary zinc – carbon cell
• Anode oxidation (–)
– Zn (s)  Zn 2+ (aq) + 2e –
• Cathode oxidation (+)
– 2MnO2 (s) + NH4+ (aq) + 2e–  Mn2O2 (s) +
2NH3 (aq) + H2O (l)
Dry Cells
• The new cell produces about 1.5V
• Once reaction reaches equilibrium its “flat”
Dry Cell
Metal Cap (+)
Mixture of Carbon &
Manganese Dioxide
Ammonium Chloride &
Zinc Chloride Electrolyte
Anode Zinc Case (–)
Cathode
Carbon Rod
Alkaline Cells
• The ordinary zinc – carbon cell
• Anode oxidation (–)
– Zn (s)  Zn 2+ (aq) + 2e –
• Immediately reacts with OH – ions in the
electrolyte to form zinc hydroxide
– Zn (s) + 2OH –(aq)  Zn(OH)2 (s) + 2e –
Alkaline Cells
• Cathode reduction (+)
– 2MnO2 (s) + H2O(l) + 2e–  MnO2 (s) + OH –(aq)
+ H2O (l)
• Five times the life of the dry cell
Alkaline Cell
Metal Cap (+)
Cathode outer
steel case
Potassium Hydroxide
Electrolyte
Powdered
Zinc
Anode
Steel or Brass
Mixture of Carbon &
Manganese Dioxide
Metal Base (–)
Button Cells
• Used in very small applications like
watches, cameras etc.
• Two main types
• Mercury zinc and silver zinc
• Anode Oxidation (–)
– Zn (s) + 2OH –(aq)  Zn(OH)2 (s) + 2e –
Button Cells
• Cathode Reduction (+)
• depends on the type of battery
– HgO(s) + H2O (l) + 2e –  Hg (l) + 2OH –(aq)
– Ag2O(s)+H2O (l) + 2e –  2Ag (s) + 2OH (aq)
• Produce an almost constant 1.35V
Button Cell
Metal Cap (–)
Zinc Powder
Electrolyte
Mercury Oxide
Cathode outer
container of
nickel or steel (+)
Secondary Cells
• Lead – Acid (Car Battery)
• Nickel cadmium Cells
• Fuel Cells
Lead Acid Battery
• Car Batteries p 211-2
• Also called storage batteries or
accumulators
• Each cell produces 2 volts so typical 12 volt
car battery contains 6 cells
• Both electrodes are lead plates separated by
some porous material like cardboard
Lead Acid Battery
• Positive electrode is coated with PbO2 Lead
(IV) Oxide
• The electrolyte is a solution of 4M sulfuric
acid
Lead Acid Battery
• Anode Oxidation (–)
• Pb(s) + SO4 2-  PbSO4 (s) + 2e –
• Cathode Reduction (+)
• PbO2(s) + SO4 2- + 4H+ + 2e –  PbSO4 (s) + 2H2O (l)
• Overall Reaction
• Pb(s) + PbO2(s) + 2H2SO4  2PbSO4 (s) +2H2O (l)
Nickel Cadmium Cells
•
•
•
•
Often called Nicads
Electrodes are Nickel and Cadmium
Electrolyte is Potassium Hydroxide
Reactions involve the hydroxides of the two
metals
Nickel Cadmium Cells
• Anode (Oxidation) (– )
• Cd (s) + 2OH– (aq)  Cd(OH)2 (s) + 2 e–
• Cathode (Reduction) (+)
• NiO-OH (s) + H2O (l) + e–  Ni(OH)2 (s) + OH– (aq)
• Overall Reaction
– Cd (s) +NiO-OH(s) + H2O(l)  Cd(OH)2 (s)+
Ni(OH)2 (s)
Fuel Cells
• Limitation of dry cells looked at so far is
that they contain reactants in small amounts
and when they reach equilibrium.
• Primary Cells are then discarded, secondary
cells are then recharged
• A cell that can be continually fed reactants
would overcome this and allow for a
continual supply of electricity
Fuel Cells
• Fuel cells transform chemical energy
directly into electrical energy
• 60% efficiency
• Space Program uses hydrogen and oxygen
with an electrolyte of Potassium Hydroxide
Fuel Cells
• Anode Oxidation (–)
– H2(g) + 2OH –(aq)  2H2O (l) + 2e–
• Cathode Reduction (+)
– O2(g) + 2H2O(l) + 4e–  4OH–(aq)
• Overall Equation
– H2(g) + O2(g)  2H2O (l)
Hydrogen Oxygen Fuel Cell
–
+
Electrolyte
Oxygen Gas
Inlet
Hydrogen
Gas Inlet
Porous
Anode
Porous
Cathode
Water outlet