ELECTRO CHEMISTRY

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Office: C610 and E-mail: harikrishnaiitm@gmail.com  

Dr Harikrishna Erothu, PhD Associate Professor

Center for Advanced Energy Studies (CAES) K L University, A. P.

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ELECTRO CHEMISTRY: Single electrode potential and its

measurement, Electrochemical cells, EMF series, Nernst

equation, Cell emf measurement, Reversible and

irreversible cells, Concentration cells, Reference

electrodes-Determination of pH using glass electrode.

Storage devices: Chemistry, construction and engineering

aspects of primary (mercury battery) and secondary (lead-

Acid cell, Ni-Metal hydride cell, Lithium cells) and fuel cells:

Hydrogen-Oxygen fuel cell and advantages of fuel cell.

Fuels – Types of fuels, Calorific value, Determination of

Calorific value

Redox reactions are those involving the oxidation and reduction of species.

OIL – Oxidation Is Loss of electrons.

RIG – Reduction Is Gain of electrons.

Oxidation and reduction must occur together. They cannot exist alone.

Redox Reactions

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Redox Reactions

LEO  the  lion  says  GER!  

GER!  

•  OIL RIG

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•  Oxidation Half-Reaction: Zn(s) → Zn2+(aq) + 2 e–. •  The Zn loses two electrons to form Zn2+.

Redox Reactions

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Redox Reactions

•  Reduction Half-Reaction: Cu2+(aq) + 2 e– → Cu(s) •  The Cu2+ gains two electrons to form copper.

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•  Overall: Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)

Redox Reactions

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Electrochemical Cells •  Electrodes: are usually metal strips/wires connected by an

electrically conducting wire.

•  Salt Bridge: A U-shaped tube that contains a gel permeated with a solution of an inert electrolyte.

•  Anode: The electrode where oxidation takes place.

•  Cathode: The electrode where reduction takes place.

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§  OXIDATION - loss of electron(s) by a species; increase in oxidation number; increase in oxygen.

§  REDUCTION - gain of electron(s); decrease in oxidation number; decrease in oxygen; increase in hydrogen.

§  OXIDIZING AGENT - electron acceptor; species is reduced.

§  REDUCING AGENT - electron donor; species is oxidized.

Terminology for Redox Reactions

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•  Convention for expressing the cell:

Anode Half-Cell || Cathode Half-Cell

Electrode | Anode Soln || Cathode Soln | Electrode

Zn(s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu(s)

Pt(s) | H2 (1 atm) | H+ (1 M) || Fe3+(aq), Fe2+(aq) | Pt(s)

•  Electrons flow from anode to cathode. Anode is placed on left by convention.

Electrochemical Cells

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•  The standard potential of any galvanic cell is the sum of the standard half-cell potentials for the oxidation and reduction half-cells.

E°cell = E°oxidation + E°reduction

•  Standard half-cell potentials are always quoted as a reduction process. The sign must be changed for the oxidation process.

Electrochemical Cells

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§  The standard half-cell potentials are determined from the difference between two electrodes.

§  The reference point is called the standard hydrogen electrode (S.H.E.) and consists of a platinum electrode in contact with H2 gas (1 atm) and aqueous H+ ions (1 M).

§  The standard hydrogen electrode is assigned an arbitrary value of exactly 0.00 V.

Electrochemical Cells

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Electrochemical Cells

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Electrochemical Series

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Electrochemical series Useful for:

ü  Predicting the oxidising or reducing ability ü  The more positive value of E, the greater the tendency to be

reduced i.e. oxidising ability ü  The more negative value of E, the better reducing ability ü Under standard conditions, any substance in the table will

spontaneously oxidise any substance lower that in the Table

Predicting the Cell EMF: E˚ = E˚cat+ E˚an

•  Predicting the Feasibility of Reaction •  Predicting whether a metal will displace another metal from

its salt solution or not

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Electro Chemical Cell: It’s a device, used to convert chemical energy into electrical energy and vice versa. These electrochemical cells are classified into two types. 1.  Galvanic or Voltaic cells: Electrochemical cells, which

convert chemical energy into electrical energy. Eg: Daniel cell, Dry cell, etc. 2. Electrolytic cells: Electrochemical cells, which convert electrical energy into chemical energy. Eg: Lead acid battery, Nickel cadmium battery etc.,

Electro Chemistry  

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Single electrode Potential: Defined as the potential developed at the interface between the metal and the solution, when a metal is dipped in a solution containing its own ions. It is represented as E.

The origin of electrode Potential: When a metal is dipped in a solution containing its own ions, the metal may undergo oxidation by losing electrons or the metal ions in solution may undergo reduction and get deposited on the metal surface. Consider a metal ‘M’ is dipped in a solution containing its ions Mn+. The tendency of metal to pass into solution (oxidation) can be represented as

Simultaneously the metal ions from the solution tend to deposit on the metal as metal atoms (reduction)

M Mn+ + ne-

Mn+ + ne- M  

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When a metal undergoes oxidation, it loses positive ions into solution leaving behind a layer of negative charges on its surface. This layer attracts +ve charges and forms an electric double layer (EDL) because of the formation of EDL, electrode potential arises.

When metal ions undergo reduction depositing metal atoms on the metallic surface, the metal surface becomes positively charged. The accumulated positive charge on the metal surface attracts a layer of –ve charges and forms an electrical double layer or Helmholtz EDL which causes the origin of electrode potential.

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The electrode potentials of any metal electrodes can be determined by standard hydrogen electrode (SHE). SHE is coupled with the electrode whose electrode potential is to be determined and the electrode potential of the electrode is determined by fixing the electrode potential and SHE as zero (at all temperatures)

Measurement of electrode potential  

Example: Consider the determination of Single electrode potential of Zinc electrode using Standard Hydrogen electrode. To determine the Single electrode potential of Zinc electrode it is coupled with Standard Hydrogen electrode as follows:

The electrode potential of Zinc electrode can be calculated as Ecell = Ecathode – Eanode Ecell = ESHE – EZn 0.76 = 0 - EZn Ezn = -0.76 V

ü  The electrode potentials can also be determined by using secondary reference electrode such as calomel electrode and Ag/Agcl electrode.

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Electron Motive Force (EMF): Defined as the potential difference between the two electrodes of a galvanic cell which causes the flow of current from an electrode with higher reduction potential to the electrode with lower reduction potential. It is denoted as E cell.

E cell = E right – E left E cell = E cathode – E anode

Standard electrode potential: Defined as potential developed at the interface between the metal and the solution, when a metal is dipped in a solution containing its own ions of unit ion concentration at 298 K. (If the electrodes involve gases then it is one atmospheric pressure) It is denoted as E0.

Problems to be done…  

Nernst’s equation for electrode potential:  

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Galvanic cells consisting of same metal electrodes dipped in same metal ionic solution in both the half cells but are different in the concentration of the metal ions. Eg: Consider the following concentration cell constructed by dipping two copper electrodes in CuSO4 solutions of M2 molar and M1 molar where M2 > M1 The two half-cells are internally connected by a salt bridge and externally connected by a metallic wire through voltmeter or ammeter.

Concentrations cells

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The electrode, which is dipped in less ionic concentration solution (M1) act as anode and undergoes oxidation. The electrode, which is dipped in more ionic concentration (M2) act as cathode and undergoes reduction.

At anode : Cu (S) Cu2+ (M1) + 2e-

At cathode : Cu2+ (M2) + 2e- Cu (S)

NCR Cu2+ (M2) Cu2+ (M1)

E of cell = E cathode – E anode.

!

E cell = [Eo + 0.0591 log (M2) ] – [Eo + 0.0591 log (M1)] n n Where (M2) > (M1)

Problems to be done…  

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ΔG  =  -­‐nFEoCell  

ΔEo    >  0  then  the  reac<on  is  spontaneous  (ΔG    <  0  )        

ΔEo    <  0  then  the  reac<on  is  non-­‐spontaneous  (ΔG    >  0  )      

Ecell  =  Eocell  +  0.0591/n  log(1/K)    Ecell=  0  at  equilibrium    Then  LogK  =  nEocell  /0.0591  

Problems to be done…on Equilibrium constant K and change in free energy, ΔG  

Important Problem (8 M)  

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Different types of single electrodes i) Metal-Metal ion electrode: These electrodes consist of a metal dipped in a solution of its own ions. Example: Zn/Zn++, Cu/Cu++, Ag/Ag+ ii) Metal-Metal salt electrode These electrodes consists of a metal in contact with it’s salt Example: Calomel electrode (Hg/Hg2Cl2/Cl-), Silver–Silver Chloride electrode (Ag/AgCl/Cl-), Lead–Lead sulphate electrode (Pb/PbSO4/SO4

2-) iii) Gas electrode Example: H2 electrode (H2/Pt/H+), Chlorine electrode (Pt/Cl2/Cl-) iv) Amalgam electrode Example: Lead amalgam electrode (Pb-Hg/Pb+) v) Oxidation– Reduction electrode Example: Pt/Fe2+ Fe3+, Pt/Ce3+ Ce4+, Pt/Sn2+ Sn4+

vi) Ion selective electrode Example: Glass electrode

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Reference Electrodes Standard electrodes with reference to these, the electrode potentials of any other electrode can be determined. The Reference Electrodes can be classified into two types i) Primary reference electrodes Ex: Standard hydrogen electrode ii) Secondary reference electrodes Ex: Calomel and Ag/AgCl electrodes

Construction, working and limitations of standard hydrogen electrode (SHE):

Limitations of SHE: i) The construction of SHE is difficult. ii) It is very difficult to maintain the concentration of H+ as 1M and Pressure H2 gas at 1atm iii) Platinum electrode is poisoned by the impurities of the gas iv) It cannot be used in the presence of oxidizing agents

The electrode reaction is:

It consists of platinised platinum foil fused to the glass tube. Mercury is placed at the bottom of the tube and a copper wire is used for electrical connections. The platinum foil is immersed in a solution containing unit molar hydrogen ions. Pure hydrogen gas is bubbled about the electrode through the H2 gas inlet at 1atm pressure. The electrode is represented as Pt/H2(g)/H+

If the concentration of H+ is 1M, H2 gas bubbled at 1atm pressure and at temperature 298 K, then the electrode is called standard hydrogen electrode. And the electrode potential is arbitrarily fixed as zero.

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Construction and working of Calomel electrode

Calomel electrode consisting of a glass container at the bottom of which mercury is placed above which a layer of mercury and mercurous chloride (called Calomel) is placed with 3/4th of bottle is filled with saturated KCl solution. Electrode potential of the cell depends on the concentration of KCl used.

The calomel electrode can be represented as

The calomel electrode acts as both anode and cathode depending upon the other electrode used. The platinum wire is used for electrical connections. Salt bridge is used to couple with other half cell.

Advantages of Calomel Electrode: ü  It is simple to construct as a reference electrode. ü  The electrode potential is reproducible and stable.

When it acts as anode, the electrode reaction:

When it acts as cathode, the electrode reaction:

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Construction and working of Silver-Silver chloride electrode Ag/AgCl electrode is a metal-metal salt electrode. It consists of narrow glass tube at the bottom of which agar is placed above which saturated solution of KCl is placed. The silver wire is used for electrical connections and it is coated electrolytically with AgCl

The cell can be represented as below: Ag(s) / AgCl (s) / Saturated KCl Electrode acts as both anode and cathode depending on the other electrode used. When it acts as anode, the electrode reaction is

When it acts as cathode, the electrode reaction is

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Construction of Glass Electrodes: Measurement of pH by glass electrodes

Ion Selective Electrode: which responds to specific ions only and develop potential against that ion while ignoring the other ions present in the solution. Eg: Glass electrode Glass electrode is a pH sensitive electrode widely used for pH determinations. It consists of a long glass tube at the bottom of which a thin and delicate glass bulb, made up of special type of glass (12 % Ba2O, 6% of Cao, 72% of SiO2) with low melting point and high electrical conductance. The glass bulb is filled with 0.1 M HCl and Ag-AgCl is used as a internal reference electrode. A Pt wire is used for electrical contact.

Glass electrode  

The glass electrode can be represented as

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Determination of pH using glass electrode To determine pH of unknown solution, the glass electrode is combined with secondary reference electrode such as calomel electrode and the glass-calomel electrode assembly is dipped in the solution whose pH is to be determined. The two electrodes are connected to potentiometer or pH meter.

The combined electrodes can be represented as Hg(l) / Hg2Cl2(S) / Saturated KCl //solution of unknown pH /glass/0.1M HCl/Ag/AgCl(s)

The EMF of the above cell is given by

E cell = E cathode – E anode

E cell = E calomel – E glass

= ESCE -EG

= 0.2422 – (Eo Glass - 0.0591 pH )

E cell = 0.2422- EoG + 0.0591 pH

pH = E cell + EoG - 0.2422/0.0591

(EG=E°G-0.0591pH)  

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Difference between Electrochemical Cell and Electrolytic Cells

Electrochemical Cell Electrolytic Cell It converts chemical energy into electrical energy.

It converts electrical energy into chemical energy.

It is based upon the redox reactions which are spontaneous.

The redox reactions are non-spontaneous and take place only when energy is supplied.

The chemical changes occurring in the two beakers are different.

Only one chemical compound undergoes decomposition.

Anode (-ve) - Oxidation takes place.

Anode (+ve) - Oxidation takes place.

Cathode (+ve) - Reduction takes place.

Cathode (-ve) - Reduction takes place

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A battery is an Electrochemical cell or often several electrochemical cells which are connected in series to produce a constant voltage to run portable goods.

Introduction: Batteries

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Mainly batteries are of 3 types: 1. Primary Battery is the one in which cell reactions are irreversible. Eg. Lechlaunche cell, Zn-Carbon Battery, Hg battery.

2. Secondary Battery is the one in which cell reactions can be reversed by passing external EMF in opposite direction. i.e. It can be used for many cycles of charging and discharging. Eg. Lead – Acid battery, Li-ion battery, Ni-Cd, Ni-metal-hydride (NiMH)

3. Flow battery is the one in which all the constituents of the battery flow throughout the battery. Eg. Hydrogen-Oxygen Fuel cell.

§  Known as Ruben- Mallory cell (Inventors)

§  Mercury batteries use either pure mercury (II) oxide (HgO - also called

mercuric oxide) or a mixture of HgO with some graphite as cathode; the

graphite also helps prevent collection of mercury into large droplets.

§  The anode is made of amalgamated zinc (Zn) powder and

§  Separated from the cathode with a layer of paper or other porous material

soaked with electrolyte (paste of ZnO+KOH).

Discharge:

§  Zinc is oxidized (loses electrons) to become zinc oxide (ZnO)

§  Mercuric oxide gets reduced (gains electrons) to form elemental mercury

Mercury Battery (Primary Battery)

The half-reaction at the

cathode is:

HgO + H2O + 2e− → Hg +

2OH−

Hg cell potential: + 1.35 V

Two half-reactions occur at the anode.

T h e f i r s t c o n s i s t s o f a n

electrochemical reaction step:

Zn + 4OH− → Zn(OH)4−2 + 2e−

followed by the chemical reaction

step:

Zn(OH)4−2 → ZnO + 2OH− + H2O

yielding an overall anode half-reaction

of: Zn + 2OH− → ZnO + H2O + 2e−

The overall reaction for the battery is

Zn + HgO → ZnO + Hg

2OH−

2e-  

Electrolyte Sodium hydroxide:

Ø  It provides constant voltage at low discharge currents

Eg: Hearing aids, calculators, and electronic watches

Potassium hydroxide:

Ø  It provides constant voltage at higher currents

Ø  Better performance at lower temperatures

Ø  Suitable for applications where current surge

Eg: Photographic cameras with flash, and watches with a

backlight

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Storage cell: It can act both as voltaic cell and electrolytic cell. When it functions as voltaic cell, it supplies electric current and the process is known as discharging. When it functions as electrolytic cell, it receives electric current and this process is known as charging. It can be used for a large no. of cycles of charging and discharging. The best example for storage cell is lead acid battery or lead acid accumulator.

Lead-­‐acid  ba6ery    

Lead-acid battery : The oldest rechargeable battery in existence, invented by the French physician Gaston Planté in 1859, lead-acid was the first rechargeable battery for commercial use. Applications: Cars, wheelchairs, scooters, golf carts and UPS systems.

Construction: Large number of anodes and cathodes are arranged alternatively in a series separated by insulators. The entire set up is immersed in dilute sulphuric acid solution.

Anode: lead plate Cathode: lead dioxide plate Electrolyte: 25% H2SO4 solution

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Discharging: Reactions taking place during discharging During discharging, it acts as voltaic cell and supplies electrical energy. Anode: Pb → Pb2+ + 2 e- (oxidation)

Pb2+ + SO42- → PbSO4 ↓

Cathode: PbO2 + 4 H+ + 2 e- → Pb2+ + 2H2O (Reduction)

Pb2+ + SO42- → PbSO4 ↓

Net reaction: Pb + PbO2 + 4 H+ + 2 SO4

2- → 2 PbSO4 ↓ + 2 H2O + Energy (=2 V)

Charging: 2 PbSO4 ↓ + 2 H2O + Energy (> 2 V) → Pb + PbO2 + 4 H+ + 2 SO4

2-

The following points can be noticed from the above reaction: ü  The concentration of sulphuric acid decreases in course of reaction. ü  Both the electrodes are covered with lead sulphate. As lead sulphate

is insoluble in sulphuric acid, it acts as a protective layer and prevents the further corrosion of lead.

ü  The cell reactions can easily be reversed by passing emf just above the voltage of the cell i.e. > 2 V.

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§  Alkaline and rechargeable batteries and commercialised in 1990 (a lot lighter than lead-acid batteries).

§  Anode: Metal hydrides like MH and MH2 Cathode: NiO(OH) and Electrolyte: Aq. KOH

§  Separator contains a thin layer of polypropylene

Nickel-Metal Hydride Battery (NiMH)

Nickel-Metal Hydride Battery  

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Ø  The battery offers an EMF of 1.3 V. Ø  Long shelf life and cycle life Ø  High capacity and rapid recharge capability Ø  Poor charge detention capacity Ø  Applications in laptop computers, cellular phones and automobiles

Redox Reaction for a Ni-MH Battery

Nickel-Metal Hydride Battery (NiMH)

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The newest rechargeable battery is based on the migration of Li+ ions and LIBs use an intercalated lithium compound as the electrode material

ü  During discharge, Li+ ions carry the current from the negative to the positive electrode, through the non-aqueous electrolyte

ü  During charging, an external electrical power source applies a higher voltage forcing the current to pass in the reverse direction. The lithium ions then migrate from the positive to the negative electrode, where they become embedded in the porous electrode material in a process known as intercalation

Lithium ion Battery (LIB)

+ - e- e-

e- e-

Li+ conducting electrolyte LiCoO2 LixC6 Graphite

Charge

Discharge Li+

Li+

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Lithium ion Battery (LIB)

In LIB, Li+ ions are transported to and from the cathode or anode with the transition metal, cobalt (Co) in LixCoO2 being oxidized from Co3+ to Co4+ during charging, and reduced from Co4+ to Co3+ during discharge.

The positive electrode half-reaction (with charging being forwards) is

The negative electrode half-reaction is

LiCoO2 Li1-xCoO2 + xLi+ + xe-

xLi+ + xe- + 6C LixC6

Applications: ü  Portable devices: Mobile phones and smartphones, laptops and

tablets, digital cameras and camcorders, electronic cigarettes, handheld game consoles and torches (flashlights)

ü  Power tools: Cordless drills and a variety of garden equipment including whipper-snippers and hedge trimmers

ü  Electric vehicles: Because of their light weight LIBs are used for a wide range of electric vehicles such as aircraft, electric cars, Pedelecs, advanced electric wheelchairs, radio-controlled models, etc.

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Fuel cells typically have about 40% conversion to electricity; the remainder is lost as heat, can be used to drive turbine generators.

Fuel cells

q  It consists of two inert porous electrodes made of graphite impregnated with finely divided Pt and 25% KOH solution as electrolyte.

q  Hydrogen gas is bubbled through one inert electrode, acts as anode.

q  Oxygen gas is bubbled through another electrode, acts as cathode.

q  The hydrogen-oxygen fuel cell produces water as a product and hence is an ideal power source for zero-emission vehicles. Hence it is called an eco-friendly battery.

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Redox Reaction in a Hydrogen-Oxygen Fuel Cell At Anode: 2 H2 (g) + 4 OH−(aq) → 4 H2O (l) + 4e− (Oxidation)

At Cathode: O2 (g) + 2 H2O (l) + 4e− → 4 OH− (aq) (Reduction) Net Reaction: 2 H2 (g) + O2 (g) → 2 H2O (l) + 1.0 V

ü  Portable applications include laptops, cellular phones, power tools, military equipment, battery chargers, unattended sensors, and unmanned aerial and underwater vehicles.

ü  Transportation applications including automobiles, buses, utility vehicles, and scooters and bicycles.

ü Utility vehicles powered by fuel cells are forklifts, golf carts, lawn maintenance vehicles, airport movers, wheelchairs, unmanned vehicles, boats, small planes, submarines, small military vehicles.

Fuel cell Applications:  

Fuel: A fuel is defined as naturally occurring or artificially manufactured combustible carbonaceous material, which serves particularly as source of heat and light and also in few cases as a source of raw material. Classification of fuels: Fuels are classified into 2 types. Based on their origin they are classified into a) Primary fuels b) Secondary fuels §  Primary Fuels: These are naturally occurring fuels which serves as

source of energy without any chemical processing. Eg: Wood, Coal, Crude oil, Natural gas, Peat, Lignite, Anthracite. §  Secondary Fuels: - These are derived from primary fuels & serves as

source of energy only after subjecting to chemical processing. Eg: Charcoal, Coke, Producer gas, Petrol, Diesel etc., Based on their physical state, fuel are classified into §  Solid §  Liquid §  Gaseous fuels

Thermal Energy

Calorific value is defined as the amount of heat liberated when a unit mass of fuel burnt completely in the presence of air or oxygen Calorific value is of two types as follows:- •  Higher calorific value (HCV) or Gross calorific value. (GCV) •  Lower calorific value (LCV) or Net calorific value. (NCV)

HCV: It is the amount of heat liberated when a unit mass of fuels burnt completely in the presence of air or oxygen and the products of combustion are cooled to room temperature. Here it includes the heat liberated during combustion and the latent heat of steam. Hence it’s value is always higher than lower calorific value.

LCV: It is amount of heat liberated when a unit mass of fuel is burnt completely in the presence of air or oxygen and the product of combustion are let off completely into air. It does not include the latent heat of steam. Therefore it is always lesser than HCV.

NCV = HCV – Latent heat of steam

Measurement of Calorific Value of a Fuel

Calculation: Mass of the fuel = M Kg Initial temp of the water = t1 0C Final temp of the water = t2 0C Change in temp. = t = (t2 – t1) 0C Specific heat of water = S Water equivalent of calorimeter = W Kg

GCV = W x S x t J/Kg

M

Bomb calorimetric method

Principle: A known amount of the fuel is burnt in excess of oxygen and heat liberated is transferred to a known amount of water. The calorific value of the fuel is then determined by applying the principle of calorimetery

i.e. Heat gained = Heat lost

CALORIFIC VALUE: Calorific value of a fuel is "the total quantity of heat liberated, when a

unit mass (or volume) of the fuel is burnt completely." Units of heat : 1.  'Calorie' is the amount of heat required to raise the temperature of

one gram of water through one degree Centigrade (15-16 °C). 2.  ‘Kilocalorie’ is equal to 1,000 calories. It may be defined as 'the

quantity of heat required to raise the temperature of one kilogram of water through one degree Centigrade. Thus: 1 kcal = 1,000 cal

An ideal fuel should have the following properties: Ø  High calorific value. Ø  Moderate ignition temperature. Ø  Low moisture content. Ø  Low NOn combustible matter. Ø  Moderate velocity of combustion. Ø  Products of combustion not harmful. Ø  Low cost. Ø  Easy to transport.