Electrochemistry-Dr. Surendran Parambadath

ELECTROCHEMISTRY

Dr. SURENDRAN PARAMBADATH (M.Sc, M.Phil, M.Tech)

Formerly: Post Doctoral Research Associate,Nano-Information Materials Research Laboratory,

Pusan National University, Busan-South Korea

Currently: Assistant ProfessorGovt. Polytechnic College, Perinthalmanna

Electrochemistry is the study of

Inter-convention of electrical energy

and chemical energy

Electrical Energy

Chemical Energy

Electrolytic Cell

This device can convert electrical energy in to chemical energy

Electrochemical Cell

This device can convert chemical energy in to electrical energy

Conductors and Insulators

This classification is based on their ability to allow electric current to pass through them.

Conductors are those substances which allow electric current to pass through them. Examples: Metals, Alloys, Graphite, Ionic compounds in fused or dissolved state.

Insulators are those substances which do not allow electric current to pass through them. Examples: Glass, Wood, Paper, Organic Compounds etc.

Types of Conductors

1. Metallic Conductors Eg: Metals and Alloys

2. Electrolytic Conductors Eg: NaCl, CuSO4, etc

3. Semi Conductors Eg: Ge doped with Ar or P.

4. Super Conductors Eg: Mercury at 4K

Metallic Conductors Electrolytic Conductors

1. Due to movement of electrons 1. Due to movement of ions

2. No Chemical Change takes place 2. Electrolysis takes place

3. No transfer of matters 3. Transfer of matter in the forms of ions

4. Conductance of metals decreases with increase of temperature

4. Conductance of metals increases with increase of temperature

Electrolytes are substances which conduct electric current through them either in the molten state or in the dissolved state. Eg: NaCl, H2SO4, KOH, HNO3 etc

Non electrolytes are substances which do not conduct electricity in the fused state or in dissolved state. Eg: Sugar. Urea, alcohol etc.

Electrolytes

Strong electrolytes: Electrolytes that dissociate almost completely into ions even at moderate concentration are called strong electrolytes.They have high conductivity.

Eg: HCl, HNO3, H2SO4 etc. NaOH, KOH etc, NaCl, CuSO4 etc.

Weak electrolytes: Electrolytes which dissociate into ions partially at moderate concentrations. They have low conductivity.

Eg: Acetic acid, Oxalic acid, NH4OH etc.

ElectrolysisElectrolysis is the process of decomposition of an electrolyte by passage of electric current.

+- battery

Na (l)

electrode half-cell

electrode half-cell

Na+

Cl-

Cl-

Na+

Na+

Na+ + e- Na 2Cl- Cl2 + 2e-

Cl2 (g) escapes

NaCl (l)

(-)

Cl-

(+)

All rights reserved. http://academic.pgcc.edu/~ssinex/E_cells.ppt. 11

+-battery

Na (l)

electrode half-cell

electrode half-cell

Molten NaCl

Na+

Cl-

Cl- Na+

Na+

Na+ + e- Na 2Cl- Cl2 + 2e-

Cl2 (g) escapes

Observe the reactions at the electrodes

NaCl (l)

(-)

Cl-

(+)

Faraday’s Law of Electrolysis1st Law………….The mass of substance discharged at an electrode during electrolysis is directly proportional to the quantity of electricity passed through the electrolyte.

m Q ………………………m = Zit

m = mass in grams of substance dischargedQ= quantity of electricity in coulombst= time in secondsZ= Electrochemical equivalent

ECE may be defined as the mass of the substance discharged by passing one coulomb of electricity.

2nd Law…….

The law states that when the same quantity of electricity is passed through different electrolytes connected in series, the amount of substance discharged at the electrodes are directly proportional to their chemical equivalent.

E = Equivalent weight m1/m2 = E1/E2

battery- +

+ + +- - -

e-

e-e- e-

Electroplating or Electrodeposition

The process of depositing a superior metal on am inferior metal by passing electric current is called electroplating.

The base metal object, which is to be plated is made the cathode in the electrolytic cell.

The rod of pure metal to be deposited on the object is made the anode.

The electrolyte is a solution of a soluble salt of the superior metal.

Examples for Superior Metals

Cr, Ni, Ag & Au

Examples for Inferior Metals

Fe & Cu

1. To protect the inferior metal object from corrosion.2. To increase the resistance to chemical attack3. To improve its physical appearance so as to make it

more attractive.4. To modify hardness5. To repair damaged part of the machinery.6. To strengthen light weight non metallic like wood,

glass, leather, cloth etc.7. To obtain conducting surfaces, eg copper plating on

wooden or plastic radar antenna masts.

Electroplating is done to achieve …….

Electroplating an object with Nickel

Step: 1 Wash the object with an organic solvent to remove any grease or oil on it. Then wash with dilute sulphuric acid to remove oxide film from the surface.Finally wash with chromic acid or detergent to clean the surface thoroughly.

Step: 2 The metal surface should be rough so that the deposit sticks firmly and permanently.

Step 3: The electrolyte is nickel sulphate solution containing nickel chloride or nickel ammonium sulphate solution.

Step 4: pH of the electrolytic solution is maintained between 4 to 5.

Step 5: The cleaned object to be plated is made the cathode of the electrolytic cell, and pure nickel plate or block, the anode.

Reaction:

At Anode,Ni Ni2+ + 2e-

At Cathode, Ni2+ + 2e- Ni

The Electrolytes used for electroplating should be

1.Highly soluble

2.Stable towards oxidation, reduction or hydrolysis

3.A good conductor

AnodizingAnodizing is a process of coating a base metal like Aluminium, or Magnesium with a thin uniform and protective oxide film.

Anode: Base metalElectrolyte: Chromic acid, dil. H2SO4 or Phosphoric acidCathode: Graphite rod or lead sheetThe anode coating being thicker than the natural oxide film, it has greater resistance to corrosion and mechanical injury.

By addition of suitable dyes and pigment to the electrolyte, brightly colored, lustrous surface coating are obtained.

+++

–––

“Cells” are containers of liquid with electrodes:

• In “electrolytic cells”, electricity is used to force chemicals to undergo a redox reaction

• In “galavanic cells”, electricity is produced spontaneously from a redox reaction

Source or use of electricity

Molten or aqueous

chemicals

Cell Electrode

Overview

Types of cells

An apparatus that allows a redox reaction to occur by transferring electrons through an external connector.

Product favored reaction

> voltaic or galvanic cell --> electric current

Reactant favored reaction

> electrolytic cell ---> electric current used to cause chemical change.

Batteries are voltaic cells

The device in which chemical energy is converted into electrical energy is called galvanic cell.

Working w.r.t: Oxidation-Reduction Reaction.

GALVANIC CELLVOLTAIC CELL

ELECTROCHEMICAL CELL

1. Two half cells, namely zinc half cell and copper half cell. In the former is a zinc rod dipped in a ZnSO4 solution and the

latter is a copper rod dipped in a CuSO4 solution. The two metallic rods are called electrodes.

2. The two half cells are connected externally by a metallic wire to a galvanometer through a key and internally by a salt bridge.

3. The salt bridge is an inverted U-tube containing an arouse solution of an inert salt like KCl, KNO3, NH4NO3 to which some agar-agar or gelatin has been added to convert it into a semi solid, ie gel. The ends of the U-tube are plugged with glass wool.

Daniel Cell

1.Permits the passage of electric current internally,2.Maintains the electrical neutrality of the solution,3.Prevents intermixing of the solutions,4.It does not take part in cell reaction.

FUNCTIONS OF SALT BRIDGE

Zn(s) / Zn2+(aq) // Cu2+

(aq)/Cu(s)

Working of the Daniel Cell

.)(2

.)(22

2

IICueCu

IeZnZn

Anode

Cathode

Zinc plate is eaten away and Copper deposits on the copper plate. Electrons produced at the zinc anode flow through the outer circuit to the copper cathode.

Electric Current is assumed to flow from copper to zinc, ie, from positive terminal to negative terminal.

Cu(s) / Cu2+(aq) // Ag+

(aq)/Ag(s)

Mg(s) / Mg2+(aq) // Ni2+

(aq)/Ni(s)

Fe(s) / Fe2+(aq) // Au3+

(aq)/Au(s)

Al(s) / Al3+(aq) // Sn2+

(aq)/Sn(s)

OTHER GALVANIC CELLS

Cu Zn

+- batterye-

e-

(-) (+)

Na+

Cl-

Cl-

Cl-

Na+

Na+

AnodeCathode

AnodeCathode

(+) (-)

Sign of the Electrodes in Electrolytic Cell and Galvanic Cell

galvanic electrolytic

needpowersourcetwo

electrodes

produces electrical current

anode (-)cathode (+)

anode (+)cathode (-)

salt bridge vessel

conductive medium

Comparison of Electrochemical Cells

No salt bridge

Electrolytic Cell Galvanic Cell

1 Electrical Energy is converted into chemical energy

Chemical Energy is converted into electrical energy

2 Electrical energy brings about a redox reaction

Electrical energy is generated by a redox reaction

3 Anode is positive while cathode is negative

Anode is negative while cathode is positive

4 Redox reaction takes place in the same container

Oxidation and reduction reactions are carried out separately

5 No salt bridge is required Salt bridge is generally required

6 Ions are discharged at both the electrodes

Ions are discharged at the cathode while anode is consumed.

Difference between Electrolytic Cell and Galvanic Cell

Some Commercial Cells

One of the main uses of electrochemical cells is the generation of portable electrical energy.

Two or more cells are connected in series to form a battery which acts as a source of electrical energy.

A commercial Cell must fulfill the following requirements.

It should be compact and light and easy to transport.Its voltage must not vary much during use.It should have reasonably long life both when being used or not being used.

There are two category of energy producing galvanic

cells

1.Primary Cells&

2.Secondary Cells

1.Primary Cells (Disposable)

Zinc carbon (flashlights, toys)Heavy duty zinc chloride (radios, recorders)Alkaline (all of the above)Lithium (photoflash)Silver, mercury oxide (hearing aid, watches)

Battery (Ancient) History

1800 Voltaic pile: silver zinc

1836 Daniell cell: copper zinc

1859 Planté: rechargeable lead-acid cell

1868 Leclanché: carbon zinc wet cell

1888 Gassner: carbon zinc dry cell

1898 Commercial flashlight, D cell1899 Junger: nickel cadmium cell1946 Neumann: sealed NiCd1960s Alkaline, rechargeable NiCd1970s Lithium, sealed lead acid1990 Nickel metal hydride (NiMH)1991 Lithium ion1992 Rechargeable alkaline1999 Lithium ion polymer

1.Primary Cells (Disposable)

In primary cells the redox reaction occurs only once and the cell becomes dead since the chemical reactions in these are not reversible.

Daniel cell, mercury cell, Dry cell etc….

Daniel Cell in the commercial form consists of a zinc electrode dipping in zinc sulphate solution contained in a porous pot. The pot is placed ia a cylindrical copper vessel containing copper sulphate solution.

Copper Vessel (Cathode)ZnSO4 Solution

Porous Pot

CuSO4 Solution

CuSO4 Crystals

+-

Zinc Rod (Anode)

...

...

...

...

...

...

...

...

...

...

..

...

...

...

...

...

...

...

...

...

...

..

...

...

...

...

...

...

...

...

...

...

..

...

...

...

...

...

...

...

...

...

...

..

….

Daniel Cell (Commercial Form)

The porous pot allows the passage of only ions from one solution to another and serves the purpose of salt bridge in the conventional galvanic cell. When connections are made as shown electrons flow from zinc to copper and current is assumed to flow from copper to zinc.

Zn(s) / Zn2+(aq) // Cu2+

(aq)/Cu(s)

Zn(s) + CuSO4 (aq) ZnSO4 (aq) + Cu(s)

The e.m.f of the cell is 1.1 V

Dry Cell Battery

Anode (-)

Zn ---> Zn2+ + 2e-

Cathode (+)

2 NH4+ + 2e- 2 NH3

+ H2

Give the net E0 of the complete recation

Give the net E0 of the complete recation

NS

Alkaline Battery

Nearly same reactions asin common dry cell, but under basic conditions.

Anode (-): Zn + 2 OH- ZnO + H2O + 2e-

Cathode (+): 2 MnO2 + H2O + 2e- Mn2O3 + 2 OH-

NS

Mercury Battery

Anode:Zn is reducing agent under basic conditions

Cathode:HgO + H2O + 2e- ---> Hg + 2 OH-

NS

2. Secondary cells (Rechargeable)Nickel cadmiumNickel metal hydrideAlkalineLithium ionLithium ion polymerLead acid

In a secondary cell, the chemical reactions taking place are reversible and can be reversed by passing electricity.

Since these cells can be recharged, they can be used again and again.

A battery consists of two or more voltaic cells connected in series.

Lead Storage Battery

Anode (-) Pb + HSO4- ---> PbSO4 + H+ + 2e- Eo = +0.36 V

Cathode (+) PbO2 + HSO4- + 3 H+ + 2e- ---> PbSO4 + 2 H2O E0 = +1.68 V

Lead Storage Battery

The lead acid accumulator used in automobiles consists of 3 to 6 cells to get a voltage of 6 to 12.

The cell has anode made of spongy lead presses in to grids and cathode made of lead dioxide, PbO2 presses into grid made of lead.

A number of lead plates are connected in parallel and a number of lead dioxide plates are also connected in parallel.

The plates are arranged alternately, separated by thin perforated plastic or fibre glass sheets.

The whole arrangement is suspended in the electrolyte which is dilutee sulphuric acid of density 1.31 gml-1, taken in a plastic or hard rubber vessel.

Working of Lead Storage Cell

When discharging a lead storage cell, At AnodeLead loses electrons which flow through the wire to the cathode.Pb Pb2+ + 2e-

The lead ions combine with the sulphate ions from sulphuric acid forming a precipitate of lead sulphate. Pb2+ + SO4

2- PbSO4

At CathodeThe electrons flowing from anode react with PbO2 of cathode and PbO2 is reduced to Pb2+ in presence of H+ ions from H2SO4.PbO2 + 4H+ + 2e- Pb2+ + 2 H2O

The lead ions formed at the cathode react with sulphate ions forming a precipitate of lead sulphate. Pb2+ + SO4

2- PbSO4

Overall Reaction

Pb + PbO2 + 4H+ + 2SO42- 2PbSO4 + 2H2O + Energy

Total EMF = 2 V

When you charge a battery, you are forcing the electrons backwards (from the + to the -). To do this, you will need a higher voltage backwards than forwards. This is why the ammeter in your car often goes slightly higher while your battery is charging, and then returns to normal.

In your car, the battery charger is called an alternator. If you have a dead battery, it could be the battery needs to be replaced OR the alternator is not charging the battery properly.

Charging a Battery

Charging of Lead Storage Cell

During discharging, both the electrodes get covered with PbSO4 and the dilute sulphuric acid is consumed and its density falls from 1.31 to 1.2 g/ml.

When recharging an external e.m.f greater than 2 volts is passed from a generator to recharge the cell. The positive pole of the generator is connected to positive pole of the storage cell.

At Anode (+ve terminal)

PbSO4 + 2H2O PbO2 + 4H+ + SO42- + 2e-

At Cathode (-ve terminal)

PbSO4 + 2e- Pb + SO42-

Overall reaction

2PbSO4 + 2H2O + Energy Pb + PbO2 + 4H+ + 2SO42-

Ni-Cad BatteryAnode (-) : Cd + 2 OH- ---> Cd(OH)2 + 2e-

Cathode (+) : NiO2 + 2H2O + 2e- ---> Ni(OH)2 + 2OH-

Cell representation: Cd CdO KOH NiO2 Ni

ELECTRODE POTENTIAL

When a metal rod is dipped in its own salt solution, two changes

occurs.1.The metal atoms convert to metal ions.

M Mn+ + ne- (Oxidation)

++++

++++

--------

1.The metal ions in solution gains electrons from metal leaving a positive charge on the metal.

Mn+ + ne- M (Reduction)

+++++++

-----

-----

Whatever may be the process, an electrical double layer generates in between the metal and the solution. This electrical double layer generates a potential difference.

The potential difference set up between the metal and its ions in the solution is called electrode potential.

It is a measure of the tendency of an electrode to lose or gain electrons when it is in contact with its own ions in solution.

i) If oxidation takes place at the electrode, the potential is called oxidation potential. ii) If reduction takes place at the electrode, the potential is called reduction potential.

Electromotive Force (EMF)

When two half cells are connected, due to the difference in potential an electric current flows from the electrode of higher potential to the electrode of lower potential.

The difference in potentials of two half cells of a cell is known as electromotive force or emf of the cell or cell potential.

EMF = Ecathode - Eanode

Electrochemical Series

It is an arrangement of elements in the increasing order of their standard reduction potential.

Lithium----------------- -3.05 VPotassiumCalciumSodiumMagnesiumAluminumZincNickelTinHydrogen--------------- 0.00CopperSilverPlatinumGold---------------------- +1.15 V

Metal SRP, Eo

Decreasing tendency to loose electrons

Increasing order of std reduction potential

Characteristics of ECS

1.Metals lying above hydrogen are easily rusted.2. Iron and metals above it decomposes steam,

liberating hydrogen gas.3.Oxides of iron and metals below it are decomposed

easily. 4.Oxides of mercury and metals below it are

decomposed on heating.

Applications of

Electrochemical Series

1. It gives an idea regarding the tendency of elements to lose or gain electrons.

Elements with lower reduction potential have a tendency to lose electrons, that is greater tendency to get oxidized.

So they are good reducing agents.

Elements with higher reduction potential have a tendency to receive electrons, that is greater tendency to get reduced.

So they are good oxidizing agents.

2. Displacement Reaction

An element above in the series can displace an element below it.

In otherwords, an element with lower reduction potential can displace an element with higher reduction potential.

Eg: Zinc has lower reduction potential than Copper. Hence zinc displaces copper from CuSO4 solution.

3. When a cell is constructed, anode should be a metal higher in the series and cathode a metal lower in the series. Eg: When a cell is constructed using zinc and copper, Zn which is higher in the series will be the anode and copper will be the cathode.

Zn Cu+ -

Anode Cathode

4. A metal above hydrogen in the series can displace H2 gas from dilute acid. But a metal below hydrogen cannot liberate H2 gas from acid. 2Na + H2SO4 Na2SO4 + H2

2K + H2SO4 K2SO4 + H2

Ca + H2SO4 CaSO4 + H2

Mg + H2SO4 MgSO4 + H2

FUEL CELL

Fuel cells are galvanic cells in which chemical energy from combustion of fuel such as H2, CO, CH4 (gases) alcohols (liquids) can be converted into electrical energy.

About 75% of the chemical energy can be converted into electrical energy.

Hydrogen-Oxygen

Fuel Cell

H2/O2 as a Fuel

Cars can use electricity generated by H2/O2 fuel cells.H2 carried in tanks or generated from hydrocarbons (fuel)

The cell consists of two electrodes made of porous graphite impregnated with a catalyst Pt, Ag or CuO.

They are placed in aqueous concentrated (35%) solution of NaOH or KOH.

H2 gas and O2 gas are continuously bubbled through the porous electrodes at the anode and cathode respectively at a pressure of 50 atm.

The reaction at the electrodes are,

At anode 2H2 + 4OH- 4H2O + 4e-

At cathode O2 + 2H2O + 4e- 4OH-

Overall reaction 2H2 + O2 2H2O + energyThe cell will produce an emf of about 1 volt. It is used in

military equipments, manned space crafts and submarines.

Methanol-Oxygen

Fuel Cell

Fuel: 1 to 2 molar methanol in water.

To keep the concentration of methanol constant a mixture of effluent and fresh methanol is recycled.

Anode: Oxidation of methanol

CH3-OH + H2O CO2 + 6H+ + 6e-

The protons H+ move from anode to the cathode via the electrolyte.

Cathode: Reduction of oxygen

3/2 O2 + 6H+ + 6e- 3H2O

The overall reactionCH3-OH + 3/2 O2 CO2 + 2H2OThis type of cells are used as energy source in (i) Space vehicles (ii) Submarines (iii) military vehicles (iv) automobiles.

Advantages of Fuel cells

1.It converts energy of the fuel directly to electrical energy.

2.Do not cause pollution problems.3.Fuel cells are light and compact. 4.Efficiency is very high (60-75%)5.Energy supply is continuous and

without any drop.

THANK YOU