Electrochemistry: Electrode Potential

• Electrochemistry studies electricity production from spontaneous chemical reactions
• Electrochemistry studies the use of electrical energy for non-spontaneous chemical transformations

Electrical Energy and Chemical Energy Interactions

• Electrical energy converts into chemical energy
• Chemical energy converts into electrical energy

Electrode Potential (E)

• Electrode potential is the tendency of an electrode to lose or gain electrons when in contact with its ions in solution
• Potential measures the energy per unit charge from drives an oxidation/reduction reaction

Standard Electrode Potential (E°)

• Standard electrode potential is measured under standard conditions
• Standard temperature is 25°C
• Standard pressure is 1 atm
• Standard concentration is 1 Molar
• It is not possible to determine experimentally the potential of a single electrode (half-cell)
• Standard reduction potentials are determined by measuring the potential difference between two electrodes combined to form a complete cell

Standard Hydrogen Electrode (SHE)

• Standard reduction potentials are measured with reference to SHE

Salt Bridge

• Salt bridge connects oxidation and reduction half-cells in a galvanic cell (voltaic cell)
• Salt Bridge maintains electrical neutrality within the internal circuit, preventing rapid reaction to equilibrium

Electrochemical Energy Systems

• Batteries are energy storage devices
• Fuel cells are energy conversion devices

Batteries/Cells

• A common use of galvanic cells is to generate portable electrical energy
• Galvanic cells are also popularly known as batteries
• A battery consists of two or more galvanic cells connected in series
• A battery is an arrangement of electrochemical cells used as an energy source
• The basis of an electrochemical cell is an oxidation-reduction reaction
• A useful battery should be light and compact for easy transport
• A useful battery should have reasonably long life whether in use or not
• A useful battery has voltage that does not vary much during its use

Types of Commercial Cells

• Primary batteries
• Secondary batteries
• Reserve batteries

Primary Cells

• The cell reaction is irreversible
• Reactions in the forward direction cannot be reversed
• Reactants convert into products and the cell discharges
• Primary cells cannot be recharged
• Examples include: Lechlanche cell (dry cell), mercury cell, Daniell cell, and alkaline dry cell

Lechlanche Cell (Dry Cell)

• Known as a dry battery because it lacks a liquid electrolyte
• The zinc vessel acts as the anode
• The cathode is a graphite rod at the center of the cell
• The electrolyte is a paste of NH4Cl, ZnCl2, and MnO2
• Traces of acetylene black and graphite powder are added
• Starch is added to thicken the mixture
• Cell representation: Zn(s)/ZnCl2(aq), NH4Cl(aq)/MnO2/C(s)
• Cathode and anode reaction occur as follows:
  • Anode: Zn(s) → Zn2+(aq) + 2e-
  • Cathode: 2MnO2(s) + H2O + 2e- → Mn2O3(s) + 2OH-
• Overall reaction: Zn(s) + 2MnO2(s) + H2O → Zn2+(aq) + Mn2O3(s) + 2OH-
• Secondary reactions occur when hydroxyl ions from the cathode react with NH4Cl, producing NH3.
• NH3 reacts with Zn2+ ions, forming [Zn(NH3)2]Cl2(s)
• EMF is 1.2 V to 1.5 V
• Secondary/side reactions include:
  • 2NH4Cl + 2OH- → 2NH3 + 2Cl- + 2H2O
  • Zn2+ + 2NH3 + 2Cl- → [Zn(NH3)2]Cl2

Lechanche Cell Applications

• The Lechanche cell is a primary cell that gives a voltage of 1.5 V
• These cells are used in transistors, tape recorders, toys, and portable electronic gadgets

Alkaline Battery

• These batteries use an alkaline electrolyte
• Modification of dry cell, in a zinc-MnO2 battery
• NH4Cl replaced by KOH or NaOH which serves as the electrolyte
• Replacing NH4Cl with KOH or NaOH prevents corrosion of zinc since zinc does not dissolve in a basic medium
• The battery has a zinc cylinder body, acting as the anode
• The graphite cathode is surrounded by a paste of MnO2
• KOH is mixed with zinc powder surrounding the MnO2 layers

Alkaline Battery Electrode Reactions

• At the Anode: Zn (s) + 2OH- (aq) → ZnO(s) + H2O (l) + 2e-
• At the Cathode: 2MnO2 (s) + H2O (l) + 2e- → Mn2O3(s) + 2OH- (aq)
• Cell reaction: Zn(s) + 2MnO2(s) → ZnO(s) + Mn2O3(s)
• Alkaline batteries have a longer life than dry cells because the zinc does not corrode, giving a voltage of 1.5 V
• They are mainly used in cameras, calculators, and watches

Mercury Battery

• Tiny cells with specialized medical and advanced electronics applications
• The anode is zinc amalgam with a solution of Zn in mercury
• The cathode uses mercuric oxide paste with some graphite
• Zinc oxide and potassium hydroxide form the electrolyte
• A cellulose paper separates the anode and cathode and acts as absorbent for the electrolyte
• The entire cell is covered in stainless steel case

Mercury Battery Electrode Reactions

• Anode: Zn amalgam(s) + 2OH- (aq) → ZnO(s) + H2O(l) + 2e-
• Cathode: HgO(s) + H2O (l) + 2e- → Hg (l) + 2OH- (aq)
• Net Cell Reaction: Zn amalgam (s) + HgO(s) → ZnO(s) + Hg(l)
• They are small and expensive
• Used in pacemakers, hearing aids, digital watches, etc

Mercury Cell Design and Properties

• In a dry cell, ZnCl2 combines with NH3 to form complex [Zn(NH3)2Cl2]; otherwise, pressure from NH3 can crack the cell
• Mercury cells provide a constant voltage because the KOH electrolyte is not consumed in the reaction

Lithium-MnO2 Batteries

• Primary batteries, with no recharge capability
• Lithium metal or compounds are used as the anode
• Specially treated MnO2 crystals are used for the cathode

Lithium-MnO2 Batteries - Electrolyte

• Lithium highly reactive with water and non-aqueous solvents, these solvents cannot be used as electrolyte
• Lithium salt solution in propylene carbonate and dimethoxy salt is used as the electrolyte
• Solvents such as thionyl chloride containing lithium compounds like LiCl, LiBr, LiAlCl4, LiSO3CF3, LiClO4 can be used
• Cell representation: Li / Non-aqueous solvent containing lithium salts / MnO2

Lithium-MnO2 electrode reactions

• Anode: Li → Li⁺ + e¯
• Cathode: Mn(IV)O2 + Li+ + e → Li+Mn(III)02
• The net cell reaction Mn(IV)O2 + Li → Li+Mn(ΙΙΙ)Ο2

Lithium-MnO2 Applications

• Lithium batteries are found in clocks, calculators, toys, digital cameras, watches, medical equipment like artificial pacemakers, thermometers, and remote car locks

Reserve Battery

• Made of magnesium and cuprous chloride
• Magnesium acts as anode; cuprous chloride as cathode
• Activated by adding sea or tap water
• Works on the principle of depositing copper metal onto the magnesium anode
• Mg2+ + 2e− ⇌ Mg(s) E° (V) = –2.356
• Cu2+ + 2e− ⇌ Cu(s) E° (V) = +0.34

Reserve Battery Electrode reactions

• Anode: Mg → Mg2 + 2e-
• Cathode: 2Cu + 2e- 2Cu
• Overall cell reaction: Mg + 2CuCl MgCl₂ + 2Cu
• Other salts include silver chloride, lead chloride, cuprous iodide, or copper thiocyanate can be used instead of cuprous chloride
• Cell potential ranges from 1.5 to 1.6 volts

Secondary Cells

• Reactions can be reversed via an external electrical energy source; therefore, cells can be recharged
• Common secondary cells include: Ni-Cd batteries and lead acid batteries used in cars

Lead Acid Battery

• Consists of a rectangular ebonite/polymeric case containing 5 M sulphuric acid (37%)
• Lead grid electrodes separated by microporous polyethylene
• The anode is coated with spongy lead, and the cathode has a paste of spongy lead and lead dioxide in 1:1 ratio
• Six anode/cathode pairs are placed in series, terminals welded, and the case is sealed
• Cell Representation: Pb–PbSO4 / H2SO4 (5 M) / Pb – PbSO4–PbO2

Lead Acid Battery Electrode Reactions

• At anode Pb(s) + SO₄²⁻(aq) ⇄ PbSO₄(s) + 2e⁻
• At cathode PbO₂(s) + SO₄²⁻(aq) + 4H⁺(aq) + 2e⁻ ⇄ PbSO₄(s) + 2H₂O (l)
• Net cell reaction: Pb(s) + 2SO₄²⁻(aq) + PbO₂(s) + 4H⁺ (aq) ⇄ 2PbSO₄(s) + 2H₂O (l)
• Each cell has a voltage of 2 V, and total voltage for six cells in series is nearly 12 volts
• Water is produced, diluting the sulphuric acid during discharge; discharge extent is checked by measuring H2SO4 density using a hydrometer, then recharged
• Lead is gradually converted into insoluble lead sulphate
• When both electrodes become lead sulphate, the battery is fully discharged and cannot be recharged
• Lead storage cells function as both a voltaic and electrolytic cell depending on whether electricity is being used out or recharged into the battery

Nickel Cadmium Battery

• Secondary alkaline storage battery: anode grid of spongy cadmium with cadmium hydroxide
• The cathode is nickeloxyhydroxide NiO(OH) paste with a small amount of graphite for electrical conductivity
• Electrolyte: 5 M KOH solution
• Cell representation: Cd(s)/Cd(OH)2/KOH(5M)/NiO(OH)(s)/Ni(OH)2(s)

Nickel Cadmium reaction details

• Chemistry
• Reactions occur during discharging
• At anode (Cd0 Cd+2): Cd + 2OH- → Cd(OH)2 + 2e-
• At cathode (Ni+3 Ni+2): 2NiO(OH) + 2H2O + 2e- → 2Ni(OH)2 + 2OH-
• Overall reaction during discharging 2NiO(OH) + Cd + 2H2O → 2Ni(OH)2 + Cd(OH)2
• Reactions occur during recharging: 2Ni(OH)2 + Cd(OH)2 → 2NiO(OH) + Cd + 2H2O

Nickel Cadmium Properties

• Commonly abbreviated NiCd or nicad
• Nickel oxide hydroxide (NiOOH) as Cathode
• Alkaline potassium hydroxide (KOH) as electrolyte
• 1.4 V Cell potential (EMF)
• The NiCd battery has 1.4 Volt
• The energy density is about double that of lead acid batteries
• Overcharge of NiCad batteries also causes damage
• It undergoes hundreds of discharge-recharge cycles
• Toxicity of Cadmium discourages its use
• Cd cell voltage is 1.4 V, compared to 1.5 V for primary alkaline cells
• The NiCd needs self re-sealing safety vents to prevent damage
• The use of Cadmium in consumer products is now deprecated on environmental grounds
• It is used in calculators, electronic flash units, cordless shavers, transistors, etc.

Nickel Metal Hydride properties

• Rechargeable batteries like nickel cadmium cells
• However, hydrogen-absorbing alloys like zirconium hydride ZrH2, vanadium hydride VH2, titanium hydride TiH2 adsorbed on an alloy of LaNi5 and TiAl replace the cadmium anode
• A nickel foil with coated with nickel hydroxide and nickel oxyhydroxide acts as the cathode
• Polypropylene separator
• 5 M potassium hydroxide electrolyte solution
• Electrodes and electrolyte are in a sealed container
• Cell Representation: ΜΗ /КОН (5 M)/Ni(OH)2, NiO(OH)

Nickel Metal Hydride reactions

• At anode: MH(s) + OH- (aq) ⇄ M(s) + H₂O (l) + e⁻
• NiO(OH)(s) + + H₂O (1) ⇄ Ni(OH)₂ (aq) + OH⁻ (aq)
• Net cell reaction: MH(s) + NiO(OH) (s) ⇄ Ni(OH)₂ (aq) + M(s)
• These batteries are inexpensive and used in cameras, medical instruments, electric toothbrushes, razors, mobile phones, pagers, etc

Zinc Air Battery specifications

• Anode made of zinc plates
• Perforated carbon plates with water repellents act as a cathode.
• Sodium hydroxide (5 M) or potassium hydroxide electrolyte
• Also has a vent for air or O2 to enter the cell.
• Cell reactions:
  • Anode: 2 Zn (s) + 4 OH⁻ (aq) → 2ZnO(s) + 2H₂O (l) + 4e⁻
  • Cathode: O₂ (g) + 2H₂O (1) + 4e⁻ → 4OH⁻(aq)
  • Overall reaction: 2Zn(s) + O2(g) → 2ZnO (s)

Zinc Air Battery uses

• 1.65 volts of electricity are produced
• Has properties of both fuel cell and batteries
• Electrically non-rechargeable; can be recharged mechanically
• Zinc-oxide converts back into zinc metal, which recharges the battery
• The zinc metal can be reused
• Cheap batteries that are replacing mercury batteries
• Varies from watches, hearing aids and cameras to electric propulsion for vehicles

Batteries Applications

• Batteries are used in motorized equipment, power tools, transistors, electronic calculators, commercial and industrial portable products
• Batteries are used in medical instrumentation, emergency lighting, toys, cordless and wireless telephones.

Lithium Ion Battery specifications

• The highest energy density of all batteries
• Lithium metal cannot be used with the traditional aqueous electrolytes
• Lithium is vigorously corrosive between Li and water
• Has flammable product

Lithium Ion Properties

• High electrochemical potential, lightest weight of all metals
• Lithium metal is explosive
• Lithium-{cobalt, manganese, nickel} dioxide is used safely
• Metallic lithium has risk of explosion

Lithium Ion Differences

• Lithium batteries are not rechargeable, but Li-ion batteries are rechargeable
• Lithium batteries use lithium in its pure metallic form
• Li-ion batteries use lithium compounds, more stable than elemental lithium

Lithium Ion materials

• Lithium Cobalt-oxide (LiCoO2)
• Lithium Manganese-oxide (LiMn2O4)
• Lithium Nickel-oxide (LiNiO2)

Lithium Ion Construction

• Anode: Layered carbon or graphite
• Non-metallic, stores, and exchanges lithium ions
• Cathode: Layered LiCoO2, lithium cobalt oxide
• Electrolyte: Non-aqueous electrolyte, mixture of organic carbonate and complexes of Li ions such as LiPF

Lithium Ion EMF/Voltage

• 3.7 V EMF/Voltage
• Reactions include:
• During charging: LiCoO2 Li1-xCoO2 + xLi+ + xe-
• During discharging: 6C + xLi+ + xe- LixC6
• Net cell reaction is
• During charging: LiCoO2 + 6C Li1-xCoO2 + LixC6

Lithium Ion Properties and Applications

• High voltage and lightweight batteries
• Used in cell phones, note PC, portable LCD TV, Semiconductor driven audio etc
• It is smaller in size
• Three times the voltage of Ni-Cd batteries

Fuel Cells Explained

• Fuel cell converts the chemical energy of fuels to electricity
• Fuel + Oxygen, create an oxidation product and electricity
• In a fuel cell, chemical substances are fed into the cells whenever energy is desired

Comparison of Fuel Cells to Batteries

• Energy Source for Fuel Cells: continuous supply of hydrogen or methanol
• Store internally Batteries energy in the electrodes
• Fuel Cells convert chemical energy into electricity and into a continuous reaction
• Batteries use electrochemical reaction
• Lifetime of FC: indefinitely with a fuel supply
• Battery Life Time is based on Limited charge cycles before degradation
• Fuel cells are eco-friendly, using green hydrogen
• Batteries use lithium-ion, which may cause concern

Proton exchange membrane fuel cell(PEMFC) & Direct methanol fuel cells (DMFCs):

• Fuel cells are different in that Chemical Energy is used, creating Electrical Energy
• Proton and Hydroxide ion production
• Anodes are porous Ni plates with Pt black catalyst
• Cathodes are porous Ni plates with Ag Catalyst
• Electrolytes are 25% NaOH or KOH

Proton and Hydroxide reactions in cells

• At anode: CH3OH + H2O → CO2 + 6H+ + 6e- Eo = 0.046 V
• At cathode: 3/2O2 + 6H+ + 6e- → 3H2O Eo = 1.23 V
• Overall CH3OH is Fuel
• The anode is Pt, Pd or Ni, and the cathode is made of Ag or Pt. Methanol and sulphuric acid are used as electrolytes.

Fuel Cell Classifications by Electrolyte

• Proton Exchange Membrane (PEMFC)
• Solid Polymer Electrolyte
• Hydrogen Oxygen
• Alkaline
• Molten Carbonate
• Phosphoric Acid
• Solid Oxide
• Direct Methanol

Hydrogen-oxygen Fuel Cell specifications

• Fuel cells directly convert combustion energy of fuels such as H2, CO, CH4, to electrical energy.
• Hydrogen-oxygen fuel cell consists of porous graphite screens coated with platinum catalyst
• Platinum catalyst is used w/ proton exchange membrane fuel cell (PEMFC) technology
• Polymeric electrolyte is used (polystyrenesulphonic acid or polymeric forms of perfluorosulphonic acid)
• Moisture is supplied to keep the membrane wet by resin
• Cell connections produce a desirable voltage

Hydrogen-oxygen Fuel Cell Reactions

• Proton Exchange Membrane (PEMFC) cell reactions
• At anode : 2H2(g) → 4H+ + 4e− E° = 0 V
• At cathode : O2(g) + 4H+ + 4e− → 2H2O E° = 1.229 V
• Overall reaction : 2H2(g) + O2(g) → 2H2O E° = 1.229 V
• Hydroxide exchange membrane fuel cell
• At anode : 2H2(g) + 4OH−(aq) → 4H2O(l) + 4e− E°=-0.828 V
• At cathode : O2(g) + 2H2O(l) + 4e− → 4OH−(aq) E°=0.401 V
• Overall reaction : 2H2(g) + O2(g) → 2H2O(l)

H2-O2 Fuel Cell

• Solid polymer electrolyte is used
• Used in spacecraft
• EMF of cell at 25 °C is 1.23 V.

Fuel Cell Properties

• Take chemical substances to convert to electrical energy
• Fuel cell differs the need that chemical substances are fed instead of inside the generator
• Fuel cells are used with gaseous or liquid anodic materials

Fuel Cell Mechanics

• In a fuel cell, hydrogen enters where H loses e. The electrodes pick up the electrons and transport to the other anode
• Electrolytes transport appropriate ions
• Cell types are are listed as:

•  Hydrogen-oxygen fuel cells
  

•   Phosphoric acid fuel cells
  

•   Molten carbonate fuel cells
  

•   Solid polymer electrolyte fuel cells
  

•   Solid oxide fuel cells
  

• pure hydrogen, a 'reformer' is employed
• liquid electrolyte used

Direct Methanol Cells

• Electrolytic device converts chemical energy directly
• Electricity and head are produced
• Consists of a porous Nickel plate with Pt black catalyst
• CH3OH runs the Anode, while O2 bubbles a A9 Catalyst
• Electrolyte is NaOH/COH mixture
• CH3 runs Anode function with a cathode that provides O2
• The electrolyte is NaOH mixed at rates of 25% volume for the reaction: CH3OH + H2O → CO2 + 6H+ + 6e-
• O2 used is 3/2 volume for the electrolyte at a rate of 6H+ + 3 H2O

Fuel Cell Technological Applications

• Fuel cell technology has a few applications
• Fuel cell electric vehicles therefore more eco-friendly
• Fuel cell technologies powered the Appolo space mission
• heat and water are byproducts
• They are extremely useful by military for portability
• power other types of electrical equipment
• Used as primary electricity generators