ELECTRO CHEMISTRY UNIT -I

Contents: Conductance-Electrolyte in solution Conductance-Specific, Equivalent and Molar conductance Ionic mobilities Kolrausch’s Law Application of conductance

EMF Galvanic Cells Types of Electrodes Nernst equation Concentration Cells Galvanic series Potentiometric titrations Numerical problems. Batteries Fuel cells

Conductance The reciprocal of the resistance is called conductance. It is denoted by C. C = 1 / R Conductors allows electric current to pass through them. Examples are metals, aqueous solution of acids, bases and salts etc. Unit of conductance is ohm-1 or mho or Siemen (S)

Classification of Conductors These may be divided into three main categories; they are: I) Gaseous II) Metallic or electronic III) Electrolytic. Gases conduct electricity with difficulty and only under the influence of high potentials or if exposed to the action of certain radiations.

Metallic or electronic conductors : Conductors which transfer electric current by transfer of electrons, without transfer of any matter, are known as metallic or electronic conductors. Metals such as copper, silver, aluminum, etc., non-metals like carbon (graphite - an allotropic form of carbon) and various alloys belong to this class.

Classification of Conductors…. Electrolytic conductors : (a) Conductors like aqueous solutions of acids, bases and salts in which the flow of electric current is accompanied by chemical decomposition are known as electrolytic conductors. b)The substances whose aqueous solutions do not conduct electric current are called non-electrolytes. Solutions of cane sugar, glycerine, alcohol, etc., are examples of non-electrolytes.

Electrolytes: Substances whose solution in water conducts electric current. Conduction takes place by the movement of ions. Examples are salts, acids and bases. Substances whose aqueous solution does not conduct electricity are called non electrolytes. Examples are solutions of cane sugar, glucose, urea etc.

Types of Electrolytes Strong electrolyte are highly ionized in the solution. Examples are HCl, H2SO4, NaOH, KOH etc Weak electrolytes are only feebly ionized in the solution. Examples are H2CO3, CH3COOH, NH4OH etc

Difference between electronic & electrolytic conductors (3) Conduction increases with increase in temperature (3) Conduction decreases with increase in temperature (2) Flow of electricity is due to the movement of ions (2) Conduction is due to the flow of electron (1)Flow of electricity takes place by the decomposition of the substance. (1) Flow of electricity take place without the decomposition of substance. Electrolytic conductors Electronic conductors

Specific Resistance: Resistance refers to the opposition to the flow of current. For a conductor of uniform cross section (a) and length (l); Resistance R, Where ‘ ρ’ is called resistivity or specific resistance.

Specific conductance or conductivity Conductance of unit volume of cell is specific conductance.    If l = 1 cm and a = 1 cm2, then R = ρ                             Κ= 1/ρ, Κ = kappa - the specific conductance           ρ = a/l. R or 1/ρ = 1/a.1/R K = 1/a×C (1/z = cell constant) Specific conductance = cell constant x Conductance The unit of specific conductance is ohm-1 cm-1.

Specific conductance or conductivity Representation of specific conductance Specific conductance depend on the nuber of ions present in unit volume (1 ml ) solution

Equivalent Conductance(λ / /\): It is the conductance of one gram equivalent of the electrolyte dissolved in V cc of the solution. /\ = KV In case, if the concentration of the solution is c g equivalent per litre, then the volume containing 1 g equivalent of the electrolyte will be 1000/C. So equivalent conductance /\ = K 1000 / c /\ = k × 1000/N          Where N = normality                 The unit of equivalent conductance is ohm-1 cm-2 equi-1.

Molar conductance The molar conductance is defined as the conductance of all the ions produced by ionization of 1 g mole of an electrolyte when present in V mL of solution. It is denoted by. Molar conductance     Λ m = k ×V                                  Where V is the volume in mL containing 1 g mole of the electrolyte. If c is the concentration of the solution in g mole per litre, then Λ m = k × 1000/c It units are ohm-1 cm2 mol-1.                Equivalent conductance =  (Molar conductance)/n                Where            n = (Molecular mass) / (Equivalent mass)

Effect of Dilution on Conductivity Specific conductivity decreases on dilution. Equivalent and molar conductance both increase with dilution and reaches a maximum value. The conductance of all electrolytes increases with temperature.

Illustrative Example The resistance of 0.01N NaCl solution at 250C is 200 ohm. Cell constant of conductivity cell is unity. Calculate the equivalent conductance and molar conductance of the solution. Solution: Conductance of the cell=1/resistance =1/200 =0.005 S. Specific conductance=conductance x cell constant =0.005 x 1 =0.005 S cm-1

Equivalent Conductance = Specific conductance x (1000/N) Solution Cont. Equivalent Conductance = Specific conductance x (1000/N) = 0.005 x 1000/0.01 = 500 ohm-1cm2eq-1 Molar Conductivity = Equivalent conductivity x n-factor = 500 x 1 = 500 ohm-1mol-1cm2

Kohlrausch’s Law “Limiting molar conductivity of an electrolyte can be represented as the sum of the individual contributions of the anion and cation of the electrolyte.” Where are known as ionic conductance of anion and cation at infinite dilution respectively.

Application of Kohlrausch’s law 1. It is used for determination of degree of dissociation of a weak electrolyte. Where, represents equivalent conductivity at infinite dilution. represents equivalent conductivity at dilution v. 2. For obtaining the equivalent conductivities of weak electrolytes at infinite dilution. 3. Determination of the solubility of a sparingly soluble salt

Thus conductance is given by product of charge and velocity. Ionic Mobilities: Equivalent conductivity of an electrolyte at any dilution is directly proportional to the charged carried by the ions and their velocities. Thus conductance is given by product of charge and velocity. At infinite dilution charge is constant ( because of complete ionisation) therefore at infinite dilution equivalent conductivity depends only on ionic velocities.

where u = ionic velocity , u – and u + are the velocities of ions Ionic mobilities cont.. Λ∞ / λ ∞ α u λ ∞ = K u Λ+ α u + or Λ- α u – λ + = K u+ or Λ = K u – where u = ionic velocity , u – and u + are the velocities of ions λ ∞ = equivalent conductivity at infinite dilution K is proportionality constant

Electromotive Force The maximum potential difference between the electrodes of a voltaic cell is referred to as the electromotive force (emf) of the cell, or Ecell. It can be measured by an electronic digital voltmeter which draws negligible current. 23 2

Electrons are driven (“pushed”) through conducting wire in the direction of anode  cathode by cell force Origin of cell force is maximum electric potential difference between electrodes or electromotive force (Ecell) or cell potential Potential difference - difference in electrical potential (electrical pressure) between two electrodes; standard unit of cell potential difference is the Volt 24

"9" Calomel Electrode Hg0l , Hg2Cl2 (sat’d) , KCl (X M = 0.09M)

Saturated Calomel Electrode H2 (1 bar) potentiometer, eo = + 0.9415 V KCl / agar salt bridge inert Pt (s) electrode 2 Cl- 2 H+ 2 e- KCl (sat aq) KCl (s) Hg2Cl2 (s) Hg (l) Pt wire

(½ Hg2Cl2 (s) + e – ----------> Hg (l) + Cl – aq) In standard cell notation this cell would be described as Pt / H2 (g) / soln with a H+ // KCl (sat) / Hg2Cl2 (s) / Hg (l) / Pt The anode half-cell reaction is ½ H2 (g) ----------> H + (aq) + e – The cathode half-cell reaction is (½ Hg2Cl2 (s) + e – ----------> Hg (l) + Cl – aq) The net cell reaction is ½ H2 (g) + ½ Hg2Cl2 (s) ------> Hg (l) + Cl – (aq) + H + (aq)

E left = E0CE + 0.0591 / n log [Hg2Cl2] / [Hg0]2[Cl-]2 Anodic Side Hg2Cl2 ↔ Hg22+ + 2Cl- Hg22+ + 2e- ↔ 2Hg0 Hg2Cl2+ 2e- ↔ 2Hg0 + 2Cl- E left = E0CE + 0.0591 / n log [Hg2Cl2] / [Hg0]2[Cl-]2 =0.268 + (0.0591 / 2) log [1/1*(0.09)2] = 0.329 volt.

Ion Selective Electrode: Ex: Glass Electrode

The glass electrode is hydrogen ion responsive electrode. A very thin membrane of the special glass is sealed on to the end of a glass tube. Which is smooth and square. The membrane potential can be developed by exchange of ions between glass membrane and the solution. Cell representation Ag(s)AgCl(s) Cl- //glass membraneelectrolyte E cell = E GE – E SCE

Quinhydrone Electrode It is a secondary electrode Quinhydrone is a combination of quinone and hydroquinone Q + 2H + + 2e-  H2Q Pt (s) quinhydrone (test solution) // KCl (sat) / Hg2Cl2 (s) / Hg (l)

E QE = E0QE + (0.0591/n) log [Q]*[H+]2 / [H2Q] = 0.699 + 0.0591/2 log [H+]2 = 0.699 + 0.0591 log [H+] =0.699 - 0.0591 pH

-nFEcell = -nFE°cell + RTln(Q) Nernst Equation Recall, in general: DG = DG° + RTln(Q) However: DG = -nFEcell -nFEcell = -nFE°cell + RTln(Q) Ecell = E°cell - (RT/nF)ln(Q) Ecell = E°cell - (0.0591/n)log(Q) The Nernst Equation

Concentration Cells Consider the cell presented on the left. The 1/2 cell reactions are the same, it is just the concentrations that differ.

Concentration Cells (cont.) Ag+ + e- Ag E°1/2 = 0.80 V • What if both sides had 1 M concentrations of Ag+? • E°1/2 would be the same; therefore, E°cell = 0.

Concentration Cells (cont.) Anode: Ag Ag+ + e- E1/2 = ? V Cathode: Ag+ + e- Ag E1/2 = 0.80 V Ecell = E°cell - (0.0591/n)log(Q) 0 V 1 Ecell = - (0.0591)log(0.1) = 0.0591 V

Galvanic Cell Types of electrochemical systems electrolytic - chemical reaction which occurs when electrical current is passed through solution voltaic/galvanic - spontaneous reactions able to generate a supply of electricity (e.g., batteries) / which converts chemical energy to electrical energy. 37

Contd.. An electrochemical cell is a system consisting of electrodes that dip into an electrolyte in which a chemical reaction either uses or generates an electric current. A voltaic, or galvanic, cell is an electrochemical cell in which a spontaneous reaction generates an electric current. An electrolytic cell is an electrochemical cell in which an electric current drives an otherwise non spontaneous reaction. 38

Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s) A complete redox reaction takes place in a galvanic cell Overall reaction separated into half-reactions which take place at the anode and cathode Given the following reaction: Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s) the two half-reactions are: oxidation half-reaction: Zn(s)  Zn2+(aq) + 2e- (anode) reduction half-reaction: Cu2+(aq) + 2e-  Cu(s) (cathode) Electrode reactions: Anode: site of oxidation; electrons originate there; neg. pole of cell (anions migrate toward) Cathode: site of reduction; electrons consumed there; pos. pole of cell (cations migrate toward) 39

Figure : Atomic view of voltaic cell Figure : Atomic view of voltaic cell. In a voltaic cell, two half-cells are connected in such a way that electrons flow from one metal electrode to the other through an external circuit. The electrons flow through the external circuit to the copper electrode where copper ions gain the electrons to become copper metal. 40

Figure Two electrodes are connected by an external circuit Daniel Cell Design 41

The Electrochemical Cell

Zinc is (much) more easily oxidized than Copper Maintain equilibrium electron densities Add copper ions in solution to Half Cell II Salt bridge only carries negative ions This is the limiting factor for current flow Pick a low-resistance bridge

The Electrochemical Series Most wants to reduce (gain electrons) Gold Mercury Silver Copper Lead Nickel Cadmium But, there’s a reason it’s a sodium drop Iron Zinc Aluminum Magnesium Sodium Potassium Lithium Most wants to oxidize (lose electrons)

Reactions at electrodes Left: Galvanic cell. Electrons are deposited on the anode (so it is neg) and collected from the cathode (so it is positive) Right: Electro-lytic cell. Elec-trons are forced out of the anode (positive) and into the cathode (negative)

Zn(s) Zn2+(aq) Cu2+(aq) Cu(s) Cell Notation Cell notation is used to describe structure of galvanic cell For the Zn/Cu cell, the galvanic cell notation is: Zn(s) Zn2+(aq) Cu2+(aq) Cu(s) = phase boundary = salt bridge anode reaction: to the left of the salt bridge cathode reaction: to the right of the salt bridge both half-cell reactions in order of spontaneous reaction Zinc solid reacts to form zinc(II) ion at the anode Copper(II) ion reacts to form copper metal at the cathode 46

Potentiometric titrations Principle It measures the change in potential , can be used for all kinds of titration : 1- acid base 2-redox 3-complexometry.

Eright= E0+0.0591/n log OX/RED Potentiometry Ecell=ERt-ELf Eleft= E0+0.0591/n log OX/RED Eright= E0+0.0591/n log OX/RED

When it is used It is used when the endpoints are very difficult to determine , either when: 1- very diluted solution. 2-coloured and turbid solution 3-absence of a suitable indicator

Instrument Combined glass electrode ( double function electrode ( Potentiometer / PH meter Magnetic stirrer

AIM: Titration of a Strong acid ( Hydrochloric acid ) against a strong base ( NaOH) REAGENTS REQUIRED: Standard NaOH solution Hydrochloric acid

Procedure: 10.0 ml of the given hydrochloric acid solution is pipetted out into a 100ml beaker and about 40 ml of distilled water is added so that the tips of combined glass electrode is completely immersed in solution.  The potential / pH of the solution is noted before adding the alkali.  Then standardized sodium hydroxide is added from a burette with 1ml increment and the readings (pH) are noted while shaking thoroughly the contents of the beaker during the addition.  Near the equivalence point sodium hydroxide is added drop wise.  To get a neat curve titration is continued with a few more increments.   

Potentiometric titration curve of 0. 1 N NaOH against 0 Potentiometric titration curve of 0.1 N NaOH against 0.1N Hydrochloric acid acid

Calculation The normality of the hydrochloric acid from the formula V1 N1 =  V2 N2.  The concentration of hydrochloric acid (in gm/l) as follow C = Eq. wt x normality of hydrochloric acid

Batteries

Battery A device that stores energy for later release as electricity.

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 cell 1899 Junger: nickel cadmium cell Daniell cell was first reliable source of energy Solved corrosion problems in Voltaic cell

Battery History 1946 Neumann: sealed NiCd 1960s Alkaline, rechargeable NiCd 1970s Lithium, sealed lead acid 1990 Nickel metal hydride (NiMH) 1991 Lithium ion 1992 Rechargeable alkaline 1999 Lithium ion polymer 1957 9v battery introduced 1991 Early lithium battery recall (mobile phone battery burst into flame)

Battery Nomenclature Duracell batteries 9v battery 6v dry cell More precisely Two cells A real battery Another battery A battery contains two or more cells All products here are dry cells (car batteries are wet cells)

Classification Primary Batteries Secondary Batteries

Non rechargeable and are meant for Primary Batteries Non rechargeable and are meant for single use. They are discarded after use

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

Standard Zinc Carbon Batteries Chemistry Zinc (-), manganese dioxide (+) Zinc, ammonium chloride aqueous electrolyte Features Inexpensive, widely available Inefficient at high current drain Poor discharge curve (sloping) Poor performance at low temperatures

Lithium Cells Liquid Cathode Lithium Cell Solid Cathode Lithium Cells ( Li - MnO2) Solid Electrolyte Lithium Cell

Lithium Manganese Dioxide Chemistry Lithium (-), manganese dioxide (+) Alkali metal salt in organic solvent electrolyte Features High energy density Long shelf life (20 years at 70°C) Capable of high rate discharge Expensive

IBM ThinkPad Backup Battery Panasonic CR2032 coin-type lithium-magnesium dioxide primary battery Application: CMOS memory backup Constant discharge, ~0.1 mA Weight: 3.1g 220 mA-h capacity

Heavy Duty Zinc Chloride Batteries Chemistry Zinc (-), manganese dioxide (+) Zinc chloride aqueous electrolyte Features (compared to zinc carbon) Better resistance to leakage Better at high current drain Better performance at low temperature

Battery Nomenclature Duracell batteries 9v battery 6v dry cell More precisely Two cells A real battery Another battery A battery contains two or more cells All products here are dry cells (car batteries are wet cells)

Standard Alkaline Batteries Chemistry Zinc (-), manganese dioxide (+) Potassium hydroxide aqueous electrolyte Features 50-100% more energy than carbon zinc Low self-discharge (10 year shelf life) Good for low current (< 400mA), long-life use Poor discharge curve

Alkaline-Manganese Batteries (2)

Secondary Batteries cycle use. These are rechargeable and are meant for multi cycle use. After every use the electrochemical reaction could be reversed by external application till the capacity fades or lost due to leakage or internal short circuit.

Secondary (Rechargeable) Batteries Nickel cadmium Lead acid Nickel metal hydride Alkaline Lithium ion Lithium ion polymer

Nickel Cadmium Batteries Chemistry Cadmium (-), nickel hydroxide (+) Potassium hydroxide aqueous electrolyte Features Rugged, long life, economical Good high discharge rate (for power tools) Relatively low energy density Toxic Used in power tools Use restricted in some countries

Ni Cd Recharging Over 1000 cycles (if properly maintained) Fast, simple charge (even after long storage) C/3 to 4C with temperature monitoring Self discharge 10% in first day, then 10%/mo Trickle charge (C/16) will maintain charge Memory effect Overcome by 60% discharges to 1.1V

NiCd Memory Effect

Lead Acid Batteries Chemistry Features Lead Sulfuric acid electrolyte Least expensive Durable Low energy density Toxic

Lead Acid Recharging Low self-discharge No memory 40% in one year (three months for NiCd) No memory Cannot be stored when discharged Limited number of full discharges Danger of overheating during charging

Lead Acid Batteries Ratings Deep discharge batteries CCA: cold cranking amps (0F for 30 sec) RC: reserve capacity (minutes at 10.5v, 25amp) Deep discharge batteries Used in golf carts, solar power systems 2-3x RC, 0.5-0.75 CCA of car batteries Several hundred cycles

Lithium Ion Batteries Chemistry Features Graphite (-), cobalt or manganese (+) Nonaqueous electrolyte Features 40% more capacity than NiCd Flat discharge (like NiCd) Self-discharge 50% less than NiCd Expensive

Lithium Ion Recharging 300 cycles 50% capacity at 500 cycles

Battery Capacity Type Capacity (mAh) Density (Wh/kg) Alkaline AA 2850 124 Rechargeable 1600 80 Ni Cd AA 750 41 Ni MH AA 1100 51 Lithium ion 1200 100 Lead acid 2000 30 Sources differ as to exact numbers Density should take packaging into account, not just chemistry Consult product specification sheets

Fuel Cells It is an electro chemical cell which converts chemical energy into electrical energy

http://www.fuelcells.org/

Working of a Fuel Cell It operates similarly to a battery, but it does not run down nor does it require recharging As long as fuel is supplied, a Fuel Cell will produce both energy and heat

A Fuel Cell consists of two catalyst coated electrodes surrounding an electrolyte One electrode is an anode and the other is a cathode The process begins when Hydrogen molecules enter the anode The catalyst coating separates hydrogen’s negatively charged electrons from the positively charged protons

Working of a Fuel Cell The electrolyte allows the protons to pass through to the cathode, but not the electrons Instead the electrons are directed through an external circuit which creates electrical current

While the electrons pass through the external circuit, oxygen molecules pass through the cathode There the oxygen and the protons combine with the electrons after they have passed through the external circuit When the oxygen and the protons combine with the electrons it produces water and heat http://www.fuelcells.org/basics/how.html

Chemistry of a Fuel Cell Anode side: 2H2 4H+ + 4e- Cathode side: O2 + 4H+ + 4e- 2H2O Net reaction: 2H2 + O2 2H2O

Individual fuel cells can then be placed in a series to form a fuel cell stack The stack can be used in a system to power a vehicle or to provide stationary power to a building http://www.fuelcells.org/basics/how.html http://www1.eere.energy.gov/hydrogenandfuelcells/pdfs/doe_h2_fuelcell_factsheet.pdf

How can Fuel Cell technology be used? Transportation All major automakers are working to commercialize a fuel cell car Automakers and experts speculate that a fuel cell vehicle will be commercialized by 2010 50 fuel cell buses are currently in use in North and South America, Europe, Asia and Australia Trains, planes, boats, scooters, forklifts and even bicycles are utilizing fuel cell technology as well http://www.fuelcells.org/basics/apps.html http://en.wikipedia.org/wiki/Hydrogen_vehicle

How can Fuel Cell technology be used? Stationary Power Stations Over 2,500 fuel cell systems have been installed all over the world in hospitals, nursing homes, hotels, office buildings, schools and utility power plants Most of these systems are either connected to the electric grid to provide supplemental power and backup assurance or as a grid-independent generator for locations that are inaccessible by power lines http://www.fuelcells.org/basics/apps.html

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