UNIT-I

ELECTROCHEMISTRY

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Dr.A.Geetha

Associate Professor & HEAD

Department of Chemistry

Kongu Engineering College

UNIT-I

ELECTROCHEMISTRY

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Introduction - cells - types - representation of galvanic cell – electrode potential – Nernst equation (derivation of cell EMF) –�calculation of cell EMF from single electrode potential – reference electrodes: construction, working and applications of standard�hydrogen electrode, standard calomel electrode, glass electrode – EMF series and its applications - potentiometric titrations�(redox) – conductometric titrations – mixture of weak and strong acid vs strong base

Electrochemistry

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Electrical Energy and Chemical Energy

ELECTROCHEMICAL CELL

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ELECTROLYTIC CELL GALVANIC CELL

EE to CE CE to EE

(DANIEL CELL, VOLTAIC CELL)

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• CELL & BATTERY

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Dry cell

CE to EE

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Lead Acid storage battery

Both process occur

Charging-EE to CE

Discharging-CE to EE

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Cell and Half Cell

KEY TERMS IN ELECTROCHEMISTRY

Conductor:

Material which conduct electric current.

Ex: Metals & Graphite

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Conductance:

The ability of the material to conduct eletric current .

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Current:

The flow of electrons through a wire or any conductor.

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Non conductor(Insulators):

Material which do not conduct electric current.

Ex: Plastics and Wood

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Semi conductors:

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Common Components

• Electrodes:

is a material/rod/strip which conducts electrons

conduct electricity between cell and surroundings Anode

is an electrode where oxidation takes place.

Cathode

is an electrode where reduction takes place.

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• Electrolyte:

is a water soluble substance forming ions in solution

and conduct an electric current

mixture of ions involved in reaction or carrying charge

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REDOX

• Oxidation-reduction: “Redox”

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• Electrochemistry:
• study of the interchange between chemical change and electrical work (It deals the amount of energy produced or consumed during chemical change)

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• Electrochemical cells:
• systems utilizing a redox reaction (Chemical change) to produce or use electrical energy

Redox Review

• Redox reactions: electron transfer processes

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• Oxidation: loss of one or more e^-
• Reduction: gain of one or more e^-

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LEOGER

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• Oxidation numbers: imaginary charges

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Oxidation-reduction

• Oxidation is loss of e^-
• Oxidation Number increases (more positive)

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• Reduction is gain of e^-
• Oxidation Number decreases (more negative)

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• Oxidation involves loss OIL
• Reduction involves gain RIG

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Redox

• Oxidation is loss of e^-

causes reduction

“reducing agent” Ex:FAS

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• Reduction is gain of e^-

causes oxidation

“oxidizing agent” Ex:KMnO₄

• Single electrode:

Half cell

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• Electrode Potential:

It is the tendency of an electrode to lose or gain electrons

Oxidation potential

Reduction potential

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• Single Electrode Potential: E

It is the tendency of an electrode to lose or gain electrons

when it is dipped in their own salt solution.

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• Standard electrode potential:E⁰

It is the tendency of an electrode to lose or gain electrons

when it is dipped in their own salt solution of 1M concentration at 25⁰C

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Types of Electrochemical cells

• Voltaic (galvanic) cells:

a spontaneous reaction generates

electrical energy (Chemical change

produces Electrical energy)

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• Electrolytic cells:

absorb free energy from an

electrical source to drive a

nonspontaneous reaction

(Electrical energy utilized to bring chemical change)

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Electrolysis

• Forcing a current through a cell to produce a chemical change for which the cell potential is negative.

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ELECTROCHEMICAL CELL or

Voltaic (galvanic or daniel) cells:

Introduction

An electrochemical cell is a device in which a redox reaction is utilized to get electrical energy. An electrochemical cell is also commonly referred to as voltaic or galvanic cell. The electrode where reduction occurs is called cathode. The electrode where oxidation occurs is called anode.

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Construction

• Electrochemical Cells are made up of two half-cells, each consisting of an electrode which is dipped in an electrolyte. The same electrolyte can be used for both half cells. These half cells are connected by a salt bridge which provides the platform for ionic contact between them. A salt bridge minimizes or eliminates the liquid junction potential.
• The practical application of an electrochemical or galvanic cell is the Daniel cell.
• It consists of a Zn electrode dipping in ZnSO₄ solution and a Cu electrode dipping in CuSO₄ solution.

Cell reaction

• Anode : Zn → Zn²⁺ + 2e^- (Oxidation)
• Cathode : Cu²⁺ + 2e^- → Cu (reduction)
• Overall : Zn + Cu²⁺ → Zn²⁺ + Cu (Redox)
• Representation of Daniel cell : Zn / Zn²⁺ || Cu²⁺ / Cu
• Cell EMF : 1.1 V

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• EMF of a cell:

Electro Motive Force

The difference of potential which causes the flow of electrons from one electrode of higher potential to the other electrode of lower potential.

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Representation of a galvanic cell

Electrochemical Series or EMF Series

• The various electrodes are arranged in the order of their increasing values of standard reduction potential on the hydrogen scale is referred to as emf or electrochemical series.

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Applications of Electrochemical Series

• To Find Reactivity of Metals
• As we move down in the electrochemical series reactivity of metal decreases
• Alkali metals and alkaline earth metals at the top are highly reactive. They can react with cold water and evolve hydrogen. They dissolve in acids forming salts.
• Metals like Fe, Pb, Sn, Ni and Co which lie a little down in the series, do not react with cold water but react with steam and evolve hydrogen.
• Metals like Cu, Ag and Au which lie below the hydrogen are less reactive and do not evolve hydrogen from water.

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• For Studying displacement reaction

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• Elements having higher reduction potential will gain electrons and that having lower reduction potential will lose electrons. Hence element higher in electrochemical series can displace an element placed lower in electrochemical series from its salt solution.

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Example

Can zinc displaces copper from its salt solution?

Zn displaces Cu from CuSO₄, because, zinc is placed higher in

electrochemical series while Cu is placed lower in electrochemical

series. Hence zinc can easily displace copper from CuSO₄.

Zn+CuSO₄ --------> ZnSO₄ + Cu

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• For choosing elements as Oxidizing Agents

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• The elements which have more electron-accepting tendency are oxidizing agents. The strength of an oxidizing agent increases as the value of reduction potential becomes more and more positive. Elements at the bottom of the electrochemical series have higher (+ve) reduction potential. So they are good oxidizing agents. Thus, oxidizing power increases from top to bottom in the series.

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Example- F₂ is a stronger oxidant than Cl₂, Br₂ and I_2.

Cl₂ is a stronger oxidant than Br₂ and I_2.

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• For choosing elements as Reducing Agents

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• The elements which have more electron losing tendency are reducing agents. The power of reducing agent increases as the value of reduction potential becomes more and more negative. Elements at the top of the electrochemical series have higher (-ve) reduction potential. So they are good reducing agents. Thus, reducing power decreases from top to bottom in the series.

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Example-

The element like Zn, K, Na, Fe, etc. are good reducing agent.

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• Displacement of hydrogen from dilute acids by metals
• The metal which can provide electrons to H⁺ ions present in dilute acids for reduction evolve hydrogen from dilute acids. The metal having negative values of reduction potential possesses the property of losing an electron or electrons.
• Thus, the metals occupying top positions in the electrochemical series readily liberate hydrogen from dilute acids and on descending in the series, tendency to liberate hydrogen gas from dilute acids decreases.
• The metals which are below hydrogen in the electrochemical series like Cu, Hg, Au and Pt do not evolve hydrogen from dilute acids.

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Example

Zinc reacts with dil.H₂SO₄ to give H₂ but Ag does not. Why?

Zn+H₂SO₄ --------> ZnSO₄ + H₂ ; E⁰ _Zn = -0.76 volts

Ag+H₂SO₄ --------> No reaction; E⁰ _Ag = +0.80 volts

The metal with a positive reduction potential will not displace

hydrogen from an acid solution.

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For Calculation of Standard emf of the cell

Standard reduction potential values are given in emf series. From the values E⁰_cell is calculated using formula

E⁰_cell or standard emf of a cell = E⁰_red(cathode) - E⁰_red(anode)

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For predicting spontaneity of the cell reaction

E=positive, G=negative

E⁰_cell > 0 cell reaction is spontaneous (feasible)

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E=negative, G=positive

E⁰_cell < 0 cell reaction is non-spontaneous (not feasible)

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E= 0, G= 0

E⁰_cell = 0 cell reaction is in equilibrium

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For determination of equilibrium constant for a reaction

We know that

-∆⁰_G = RTlnK

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= 2.303RT logK

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log K =

log K = (-∆⁰_G = nFE⁰)

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Thus, from the value of E⁰ for a cell reaction, its equilibrium constant can be calculated.

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Standard Electrodes �or �Reference Electrodes

• A standard electrode is one which always has a standard potential at all temperatures.

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• It is an electrode of standard potential with which potentials of other electrodes are compared (is called a reference electrode).

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• It can acts both as anode or cathode depending upon the nature of other electrode.

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Standard Electrodes or Reference Electrodes

Classification:

1. Primary reference electrodes Ex : Standard Hydrogen electrode
2. Secondary reference electrodes Ex: Calomel and Ag/AgCl electrodes

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Reference

Electrode

Working or Indicator

Electrode

The part of the cell that is kept constant

The part of the cell that contains the unknown solution

Construction and Working of Standard Calomel Electrode (SCE)

• A common reference electrode.
• It consists of a wide glass tube.
• Mercury is placed at the bottom of the glass tube.
• A paste of mercury and mercurous chloride(Calomel) is placed above the mercury. The remaining portion above the paste is filled with a saturated KCl solution of known concentration (0.1N, 1.0N and saturated) .
• A platinum wire is immersed into the mercury to obtain electrical contact.

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Pt wire

Electrode representation: Anode

Pt, Hg_(l), Hg₂Cl₂(s) / KCl(satd)

Electrode representation: Cathode

KCl(satd) / Hg₂Cl₂(s), Hg _(l), Pt

Electrode potential

Measurement of pH using Calomel electrode

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Measurement of pH using Calomel electrode

Hydrogen electrode containing a solution of unknown pH combine with the calomel electrode to set up a complete cell.

We use a saturated calomel electrode as the reference and the complete cell can be represented as :

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E = E_cal- E_H

= 0.2422 + 0.0591 pH (E_H = -0.0591 pH)

pH =

Pt, H₂ (1atm)/ H⁺ (C=?)/ /KCl(satd) / Hg₂Cl₂(s), Hg _(l), Pt

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Potentiometric Titration

• According to Nernst equation, the potential of an electrode depends upon the concentration of the ions in the solution.
• In a titration, there is a change in concentration of ions, which can be noted by measuring the potential of a suitable electrode.

Thus potentiometric titrations are titrations which involve the measurement of electrode potential with the addition of the titrant.

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• It needs two electrodes.

Reference electrode and Indicator electrode

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Potentiometric titrations depend on measurement of emf between reference electrode and an indicator electrode.

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Indicator or working electrode:

The electrode in which the potential changes with change in concentration is known as Indicator electrode.

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Theory

• Potential of an Electrode dipping in solution of eletrolyte depends upon the concentration of active ions.

E= E⁰ + (RT/nF) log C

• Small Change in active ion concentration in the solution changes the electrode potential correspondingly
• Concentration of Active ion decreases electrode potential of indicator electrode decreases
• The potential of Indicator electrode is measured potentiometrically by connecting with a reference electrode (Saturated Calomel Electrode)

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Determination of End point

• The emf of a cell changes by the addition of a small amount of titrant. So concentration of reversible ion in contact with indicator electrode changes.
• Record the change in emf with every small addition of titrant.
• The changes of potential will be slow at first, but at equivalence, the point change will be sharp.
• The values are plotted against corresponding volume changes.
• Change in emf with addition of titrant (⧍E/⧍V) is plotted against volume (V)
• The maximum of the curve gives the end point.

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• Fig (a)– Volume of Titrant Vs Emf
• Fig(b) Volume of titrant Vs (⧍E/⧍V)

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Classification of Potentiometric Titration

• Acid –Base titration (HCl vs NaOH)

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• Redox titration(Reduction- Oxidation)

(FeSO₄ vs K₂Cr₂O₇)

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• Precipitation titration (AgNO₃ vs KCl)

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Redox Titration

• Oxidizing Agent--- K₂Cr₂O₇ or KMnO₄
• Reducing Agent--- Ferrous Salt
• Indicator----- Platinum Indicator

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K₂Cr₂O₇ +6 FeSO₄+7 H₂SO₄------------K₂SO₄+ Cr₂(SO₄)₃+3Fe₂(SO₄)₃ +7H₂O

• The initial concentration of an ion Fe²⁺ can be found by titrating with a strong Oxidizing agent such as Cr⁶⁺ . That reaction is carried out half side other half is reference electrode

^Pt_(S)^/Fe2+/Fe3+ /^{Reference electrode(Calomel)}

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on adding K₂Cr₂O₇ from the burette, emf of the cell will increase first slowly, but at the equivalence point,Fe²⁺ will have been totally consumed and the potential will then be controlled by the concentration ratio of Cr⁶⁺/Cr³⁺.

• emf value is recorded for every one mL of Oxidizing agent.
• Plot the graph between Volume and emf. From the graph find the end point.

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3Fe²⁺ +Cr⁶⁺ 3Fe³⁺ +Cr³⁺

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Advantages of Potentiometric Titrations

• Potentiometric titrations can be carried out in colored solutions, where indicators cannot be used
• There is no need of prior information about the relative strength of titrant before the titration

CONDUCTOMETRIC TITRATION

• Volumetric method based on the measurement of conductance of the solution during the titration

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Process

• Taking a solution to be titrated in a beaker kept in a water bath at a constant temperature.
• Conductivity cell is dipped and connected to a conductivity bridge.
• The titrant is added from the burette(Fig)
• Conductance is measured each addition of solution.
• Recorded value is plotted
• From the graph end point is noted.

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Types of Conductometric Titrations

• Acid –Base titration
• Strong Acid Vs Strong Base
• Weak Acid Vs Strong Base
• Mixture of Weak and Strong Acid Vs Strong Base
• Precipitation titration
• Replacement titration
• Redox titration
• Complexometric titration

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Procedure

• Calibrate the instrument by releasing the calibration knob
• Standard Sodium Hydroxide is taken in the burette
• The given acids is made upto 100ml in the standard measuring flask (SMF)
• 20 ml of made up acids + 20 ml of conductivity water are added in 100 ml beaker
• Conductance is noted for addition of every 1ml of Standard Sodium Hydroxide
• Plot a graph between Volume of Standard Sodium Hydroxide Vs Conductance
• End Points I & II are noted from the graph
• Equivalent Weight of Hydrochloric Acid = 36.5
• Equivalent Weight of Acetic Acid = 60

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Strong Acid Vs Strong Base(HCl Vs NaOH)

• Solution of electrolytes conducts electricity due to the presence of ions. The specific conductance of solution is proportional to the concentration of ions in it. The reaction between HCl and NaOH may be represented as

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• When a solution of hydrochloric acid is titrated with NaOH, the fast moving hydrogen ions are progressively replaced by slow moving sodium ions. As a result conductance of the solution decreases. This decrease in conductance will take place until the end point is reached. Further addition of alkali raises the conductance sharply as there is an excess of hydroxide ions.
• A graph is drawn between volume of NaOH added and the conductance of solution. The exact end point is the point of intersection of the two straight lines.

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HCl + NaOH NaCl + H₂O

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Mixture of Weak and Strong Acid Vs Strong Base�(HCl and CH₃COOH Vs NaOH)�

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HCl + NaOH NaCl +H₂O

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CH₃COOH+NaOH CH₃COO^-Na⁺ +H₂O

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Advantages

• This method can be used with very diluted Solution
• This method can be used with Coloured and Turbid solution in which the end point cannot be seen clearly
• This method can be used in which there is no suitable indicator
• Used for acid-base, redox, precipitation titration etc.,

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