Electrolytes , Electolytic Cell and Electrochemical Cell

04/11/2020 Vinod 0 Comments

Electrolytes :-

Electrolytes are substances that produce an electrically conducting solution when dissolved in a polar solvent , such as water . The dissolved electrolytes separates into cations and anions , which disperse uniformly through the solvent .

Examples – NaCl , HCl , H₂SO₄ , HNO₃ , CaCl₂ , NaOH , CH₃COOH , KNO₃ etc .

Strong Electrolytes :-

Strong electrolytes are solutes or solutions I.e. electrolytes that completely dissociates in solution . The solution will contain only ions and no molecules of the electrolytes .

  • Strong electrolytes are good conductors of electricity but only in aqueous solution .

Examples – NaCl , HCl , H₂SO₄ , HNO₃ , CaCl₂

Weak Electrolytes :-

Weak electrolytes are electrolytes that does not completely dissociate in aqueous solution . The solution will contain both ions and molecules of the electrolytes .

  • Weak electrolytes are only partially ionize in water usually 1 % to 10%

Examples – CH₃COOH , NH₄OH , H₃PO₄ , H₂CO₃ etc .

Non-electrolytes :-

Non-electrolytes are substance that does not exist in an ionic form in aqueous solution . Non-electrolytes tend to be poor electrical conductors and don’t readily dissociate into ions when dissolved .

Example – Ethanol , Sugar , Chloroform .

Electrolysis :-

It is a technique that uses a direct electric current ( DC ) to drive an otherwise non -spontaneous chemical reaction .

  • The site where electrolysis occurs is in an electrolytic cell , which is a type of electrochemical cell .

Electrolytic cell :-

ElectrolyticCells
Fig . 1

In an electrolytic cell external source of voltage is used to bring about a chemical reaction . The electrochemical process are of great importance in the laboratory and chemical industry .

  • One of the simplest electrolytic cell consists of two copper strips dipping in an aqueous solution of copper sulphate .
  • If a DC voltage is applied to the two electrodes , then Cu²⁺ ions discharge at the cathode ( negative charged ) and the following reaction takes place –

Cu²⁺ (aq ) + 2e⁻ → Cu ( s )

Copper metal is deposited on the cathode . At the anode , Copper is converted into Cu²⁺ by the reaction –

Cu ( s ) → Cu²⁺ ( s ) + 2e⁻

Thus Copper is dissolved ( oxidised ) at anode and deposited ( reduced ) at cathode .

  • This is the basis for an industrial process in which impure copper is converted into copper of high purity .
  • Sodium and magnesium metals are produced by the electrolysis of their fused chlorides and Aluminium is produced by electrolysis of Aluminium oxide in presence of cryolite .

Example – Electrolysis of liquified NaCl

NaCl → Na⁺ + Cl⁻

2Na⁺ + 2e → 2Na ( on cathode ) , Reduction

2 Cl⁻ → Cl₂ + 2e ( on anode ) , Oxidation


2Na⁺ + 2Cl⁻ → 2 Na + Cl₂

Faraday’s laws of Electrolysis :-

Michael Faraday was the first scientist who describe the quantitative aspects of electrolysis .

Faraday proposed two laws called Faraday’s law of electrolysis

First law :-

The mass of the substance deposited or liberated at any electrode is directly proportional to the quantity of electricity or charge passed .

In the mathematical form this law can be represented as follows

m α Q

m = ZQ

where m is the mass of substance in gram ( g ) , Q is charge measured in coulombs ( C ) and Z is the proportionality constant in g / C and is also known as electrochemical equivalent .

Electrochemical equivalent :-

It is the mass of substance produced at the electrode during electrolysis by one coulomb of charge I . e .

Z = m / Q

Since Q = It

Where I is current at time t

Hence m = ZQ

Or m = Z I t

If Equivalent weight of a substance is E , then

Z = E / 96500

Or. m = W = E I t / 96500

Second law :-

It states that – when the same quantity of electricity is passed through different electrolytes , the masses of different ions liberated at the electrodes are directly proportional to their chemical equivalents ( equivalent weights ) that means –

The electrochemical equivalent ( Z )of an element is directly proportional to its Equivalent weight ( E )

Or Z α E

1 Faraday ( F ) = charge of an electron × Avogadro’s number

Or 1F = 1.6021 x 10⁻¹⁹ C x 6.02 x 10²³ mol⁻¹

= 96487 or 96500

One mole of different ions liberated at the electrodes required amount of current electricity = n F

Where n is the valency bof ions .

Thus

1 mole of Ag⁺ = 1 x F = F

1 mole of Cu²⁺ = 2 x F = 2F

1 mole of Al³⁺ = 3 x F = 3F

1 mole of Mg²⁺ = 2 x F = 2F

For the electrode reaction

Ag⁺ + e → Ag

Cu²⁺ + 2e → Cu

Al³⁺ + 3e → Al

Mg²⁺ + 2e → Mg

Electrochemical cell :-

It is a device capable of either generating electrical energy from chemical reactions .

Examples – Galvanic cell and voltaic Cell and Daniel cell .

Galvanic cell or voltaic cell :-

A Galvanic cell is an Electrochemical cell that converts the chemical energy of a spontaneous Redox reaction into electrical energy . In this device the Gibbs energy of the spontaneous Redox reaction is converted into electrical work which may be used for running a motor or other electrical gadgets like heater , fan , geyser etc .

voltaic 20cell
Fig . 2 . Galvanic Cell .
  • The Galvanic cell construct the pattern of Daniel cell by taking combinations of different half cells . Each half cell consists of a metallic electrode dipped into an electrolyte .
  • The two half cells are connected by a metallic wire through a voltmeter and a switch externally .
  • The electrolytes of the two half cells are connected internally through a Salt Bridge . A Salt Bridge is a U – shaped containing KOH or other electrolyte solidified by agar -agar gel . It is used to make electrical connection of both the electrolytes of the cell . It is used to maintain in electrical neutrality of charges in both the beakers .
  • Sometimes both the electrodes dip in the same electrolyte solution and in such cases we don’t require a Salt Bridge .
  • At each electrode – electrolyte interfare there is a tendency of metal ions from the solution to deposit on the metal electrode trying to make it positively charged . At the same time metal atoms of the electrode have a tendency to go into the solution as ions and leave behind the electrons at the electrode trying to make it negatively charged .

Daniel Cell :-

Daniel Cell having electrodes of zinc and copper dipping in the solutions of their respective salt . This cell converts the chemical energy liberated during the Redox reaction

Zn ( s ) + Cu²⁺ ( aq ) + Cu ( s )

This reaction is a combination of two half reactions , whose addition gives the overall cell reaction –

1 . Cu²⁺ + 2e⁻ → Cu ( s ) , Reduction half reaction

2 . Zn ( s ) → Zn²⁺ + 2e⁻ , Oxidation half reaction .

These reactions occur in two different portions of the Daniel cell . The reduction half reaction occurs on the Copper electrode while the oxidation half reaction occurs on the Zinc electrode . These two portions of the Cell are also called half Cells or redox couples .

Electrode potential :-

A potential difference develops between the electrode and the electrolytes which is called electrode potential . When the concentrations of all the species involved in half – cell is unity then the electrode potential is known as standard electrode potential .

  • The potential difference between two electrodes of a Galvanic Cell is called the Cell potential and is measured in Volt .

It is also called the electromotive force ( emf ) of the Cell when no current is drawn through the Cell .

  • The emf of the Cell is positive and is given by the potential of the half – Cell on the right hand side minus the potential of the half – Cell on the left hand side .

I . e . Ecell = Eright – Eleft

Example – Cell reaction

Cu ( s ) + 2Ag⁺ ( aq ) → Cu²⁺ + 2Ag ( s )

Half Cell reactions

2Ag+ ( aq ) + 2e⁻ → 2Ag ( s ) , Reduction on cathode .

Cu ( s ) → Cu²⁺ ( aq ) + 2e⁻ , Oxidation on anode .

  • Silver electrode acts as a cathode and Copper electrode acts as an anode . The Cell can be represented as –

Cu ( s ) | Cu²⁺ ( aq ) | | Ag⁺ ( aq ) | Age ( s )

then the emf of the Cell

Ecell = Eright – Eleft = EAg⁺ / Ag – ECu²⁺ / Cu

Measurement of Electrode potential :-

Standard hydrogen electrode 33
Fig . 3 . Standard Hydrogen Electrode .

According to convention a half – Cell called a standard hydrogen electrode represented by Pt ( s )| H₂ ( g ) | H⁺ ( aq ) is assigned a zero potential at all temperature corresponding to the reaction

H⁺ ( aq ) + e⁻ → 1 / 2 H₂ ( g )

  • The standard hydrogen electrode consists of a Platinum electrode coated with Platinum black .The electrode is dipped in an acidic solution and pure hydrogen gas is bubbled through it . The concentration of both the reduced and oxidised forms of hydrogen is maintained at Unity . The pressure of hydrogen gas is 1 bar and concentration of hydrogen ion in the solution is 1 molar .
  • If the concentrations of the oxidised and reduced forms of the species in the right of Cell are unity then the Cell potential is equal to standard electrode potential , E₁ of the given half Cell

E = E₁ – E₂

As E₂ for standard hydrogen electrode is zero .

E = E₁ – 0 = E₁

The measured emf of the Cell :-

Pt ( s ) | H₂ ( g , 1 bar ) | H⁺ ( aq , 1 M ) || Cu²⁺ ( aq , 1 M ) | Cu is 0·34 and it is also the value for the Standard hydrogen electrode of the half cell corresponding to the reaction

Cu²⁺ ( aq , 1 M ) + 2e⁻ → Cu ( s )

Electrochemical cell and Gibbs energy of the reaction :-

Electrical work done in one second is equal to electrical potential multiplied by total charge passed . If we want to obtain maximum work from a Galvanic cell then charge has to be passed reversibly . The reversible work done by a Galvanic cell is equal to decrease in its Gibbs energy and therefore if the emf of the cell is E and nF is the amount of charge passed and ∆G is the Gibbs energy of the reaction then

∆G = – nF Ecell

The value of ∆G is depends on n . Ecell is an intensive parameter but ∆G is extensive thermodynamic property .

If we write the reaction

Zn ( s ) + Cu²⁺ ( aq ) → Zn²⁺ ( aq ) + Cu ( s )

∆G = – 2F Ecell

When we write the reaction

2 Zn ( s ) + 2 Cu²⁺ ( aq ) → 2 Zn²⁺ ( aq ) + 2Cu ( s )

∆G = – 4F Ecell

If the concentration of all the reacting species is unity , then

Ecell = E’cell

∆G = – nF E’cell

Thus , from the measurement of E’cell we can obtain an important thermodynamic quantity ∆G , standard Gibbs energy of the reaction .

From the latter we can calculate equilibrium constant by the equation

∆G = – R T InK

Battery or Cells :-

Batteries are a collection of one or more cells whose chemical reaction create of flow of electrons in a circuit .

All batteries are made up of three basic components :- an anode ( the negative side ) , a cathode ( the positive side ) and some kind of electrolytes ( substances that chemically with the anode and cathode )

There are mainly two types of batteries :-

1 .Primary batteries and

2 . Secondary batteries .

Primary batteries :-

In the primary batteries , the reaction occurs only once and after use over a period of time battery becomes dead and cannot be reused again .

Examples – Dry cell ( Leclanche cell ) , Galvanic Cell or Voltaic Cell .

CNX Chem 17 05 DryCell
Fig . 4 . Dry Cell .
  • It was first discovered by Leclanche in 1868 it is used commonly in transistors and clocks .
  • The cell consists of a Zinc container that also acts as anode and the cathode is a a carbon ( graphite ) rode surrounded by powdered magnese dioxide and carbon .

The space between the electrodes is filled by a moist paste of Ammonium Chloride and Zinc Chloride .

  • The electrode reactions are complex but they can be written as

Anode : Zn ( s ) → Zn²⁺ + 2e⁻

Cathode : MnO₂ + NH₄⁺ + e⁻ → MnO ( OH ) + NH₃

In the reaction at Cathode magnese is reduced from the +4 Oxidation state to +3 state . Ammonia produced in the reaction forms a complex with Zn²⁺ to give [ Zn ( NH₃ )₄ ]²⁺

The Cell has a potential of nearly 1.5 volt .

Secondary batteries :-

Secondary batteries after use can be recharged by passing current through it in the opposite direction so that it can be used again . A good secondary Cell can undergo a large number of discharging and charging cycles .

Examples – Lead storage battery , Nickel Cadmium Cell .

00002933
Fig . 5 . Lead storage Cell .

It consists of a lead anode and a grid of lead packed with lead oxide ( PbO₂ ) as cathode . A 38% solution of Sulphuric acid is used as electrolytes .

The cell reactions when the battery is in use are given as

Anode : Pb ( s ) + SO₄²⁻ ( aq ) → PbSO₄ ( s ) 2e⁻

Cathode : PbSO₂ ( s ) SO₄²⁻ ( aq ) + 4 H⁺ ( aq ) + 2e⁻ → PbSO₄ ( s ) + 2 H₂O ( l )

i . e . overall Cell reaction consisting of cathode and anode reaction is

Pb ( s ) + PbO₂ ( s ) + 2 H₂SO₄ ( aq ) → 2 PbSO₄ ( s ) + 2 H₂O ( l )

On charging the battery the reaction is reversed and PbSO₄ ( s ) on anode and Cathode is converted into Pb and PbO₂ respectively .

Leave a Reply