Entry Test Preparation 2015, Chemistry BOOK 1 Chapter # 10 Electrochemistry , Theory and Key Points

Entry Test Preparation 2015

Chemistry

Chapter # 10

electrochemistry

Electrochemistry:

Electrochemistry is concerned with the conversion of electrical energy into chemical energy in electrolytic cells as well as the conversion of chemical energy into electrical energy in galvanic or voltaic cells.

Electrochemical Cells:

An electrochemical cell is a system consisting of electrodes that are submerged into an electrolyte and in which a chemical reaction uses or generates electric current.

Electrolytic Cell:

An electrolytic cell is an electrochemical cell in which electric current is used to drive a non-spontaneous reaction. Examples of electrolytic cells are Down’s cell and Nelson’s cell, etc.

Electrolytic Conduction:

Electrolytes in the form of solution or in the fused state have the ability to conduct electricity. The current is carried by ions having positive and negative charges, which are produced in the solution or in fused state due to ionization of the electrolyte.

Using inert electrodes (Platinum or graphite)

Electrolyte	Cathode	Anode
PbBr2(l)	Pb (s)	Br2 (g)
NaCl(l)	Na (s)	Cl2 (g)
NaCl (aq)	H2 (g)	Cl2 (g)
CuCl2 (aq)	Cu (s)	O2 (g)
CuSO4 (aq)	Cu (s)	O2 (g)
KNO3 (aq)	H2 (g)	O2 (g)
NaOH (aq)	H2 (g)	O2 (g)
H2SO4 (aq)	H2 (g)	O2 (g)

When electrodes take part in the reaction

Electrolyte	Copper Cathode	Copper Anode
CuSO4 (aq)	Cu deposits	Cu (s) dissolves to form Cu2+ ions
Electrolyte	Silver Cathode	Silver Anode
AgNO3 (aq) and HNo3 (aq)	Ag deposits	Ag (s) dissolves to form Ag+ ions

Electrolysis

Electrolysis is the process in which a chemical reaction takes place at the expense of electrical energy. Electrolysis is used for the extraction of elements and for the commercial preparation of several compounds. It is also used for electroplating.

Examples of Electrolysis:

Extraction of sodium by the electrolysis of fused sodium chloride (Down’s Cell)

NaCl (s)  Na⁺ + Cl^-

At cathode      Na⁺ (l) + e^-  Na (s)

At anode         Cl^- (l)  Cl (g) + e^-

Cl (g) + Cl (g)  Cl2 (g)

Caustic Soda is obtained on industrial scale by the electrolysis of concentrated aqueous solution of sodium chloride using titanium anode and mercury or steel cathode (Nelson’s Cell)

At anode         2Cl(g)  Cl₂(g) + 2e

At cathode

2H₂O(l) + 2e  H₂ (g) + 2OH (aq)

2Na+ (aq) + 2Cl (aq) + 2H₂O (l)

Cl₂ (g) + H₂ (g) + 2Na⁺ (aq) + 2OH (aq)

Magnesium and calcium metals are extracted by the electrolysis of their fused chlorides.

Aluminium is extracted by electrolyzing fused bauxite (Al₂O_3.2H₂O) in the presence of fused cryolite (Na₃AlF₆).

Anodized aluminium is prepared by making it an anode in an electrolytic cell containing sulphuric acid or chromic acid, which coats a thin layer of oxide on it. The aluminium oxide layer resists attack by corrosive agents. The freshly anodized aluminium is hydrated and can absorb dyes.

Electrolytic cell can also be used for the purification of copper where impure copper is made the anode and a thin sheet of pure copper is made the cathode. Copper sulphate solution is used as an electrolyte.

Copper, silver, nickel or chromium plating is done by various types of electrolytic cells.

VOLTAIC OR GALVANIC CELL

A voltaic or galvanic cell is an electrochemical cell in which a spontaneous reaction generates electric current. A voltaic or a galvanic cell consists of two half-cells that electrically connected. Each half cell is a portion of the total cell in which a half reaction takes place. The left half cell consists of a stripe of zinc metal dipped in 1M solution of zinc sulphate giving the following equilibrium

Zn(s)  Zn²⁺ (aq) + 2e^-

The right half – cell is a copper metal stripe that dips in to 1M copper sulphate solution and the equilibrium here is presented as follows:

Cu(s)  Cu²⁺ (aq) + 2e^-

The following half – cell reaction occur at two electrodes:

At anode         Zn(s)  Z²⁺ (aq) + 2e-

At cathode      Cu²⁺ (aq) + 2e^-   Cu(s)

The overall voltaic cell reaction is the sum of these two half cell reactions.

Zn(s) + Cu²⁺ (aq)  Zn²⁺ (aq) + Cu(s)

It can be noted that reduction occurs at the copper electrode and oxidation occurs at the zinc electrode.

On the other hand, if the external circuit is replaced by a source of electricity that opposes the voltaic cell, the electrode reactions can be reversed. Now the external source pushes the electrons in the opposite direction and supplies energy or work to the cell so that the reverse non-spontaneous reaction occurs. Such a cell is called a reversible cell.

For the zinc copper cell, the half cell reactions are reversed to give

Zn²⁺ (aq) + 2e^-    Zn(s)

Cu(s)  Cu²⁺ (aq) + 2e^-

and the overall reaction becomes

Zn²⁺ (aq) + Cu(s)  Zn (s) + Cu²⁺ (aq)

Now, oxidation occurs at the copper electrode and reduction takes place at the zinc electrode.

Whether a cell operates as a voltaic or an electrolytic cell, the electrode at which reduction occurs is called the cathode while the electrode at which oxidation takes place is called the anode.

ELECTRODE POTENTIAL

Electrode potential is developed when a metal is dipped into a solution of its own ions. The potential set up when an electrode is in contact with one molar solution of its own at 298 K is known as Standard Electrode Potential or Standard Reducing Potential of the element. It is represented by E.

Standard electrode potential of hydrogen has arbitrarily been chosen as zero while the standard electrode potentials of other elements can be found by comparing them with standard hydrogen electrode potential.

The electrochemical Series

When elements are arranged in the order of their standard electrode potentials on the hydrogen scale, the resulting list is known as Electrochemical Series.

Application of Electrochemical Series:

·       Prediction of the Feasibility of a Chemical Reaction

·       Calculation of the Voltage or Electromotive Force (emf) of Cells

·       Comparison of Relative Tendency of Metals and non Metals to get oxidized or reduced

·       Relative Chemical Reactivity of Metals

·       Reaction of Metals with Dilute Acids

·       Displacement of one metal by another from its solution

Primary and secondary Cells

Voltaic or Galvanic cells which cannot be recharged are called primary cells and the one which can be recharged are called secondary cells.

Lead Accumulator:

It is commonly used as a car (automobiles) battery. It is secondary or a storage cell. Passing a direct current through it must charge it. The charged cell can then produce electric current when required. The cathode of a fully charge lead accumulator Is lead oxide, PbO₂ and its anode is metallic lead. The electrolyte is 30% sulphuric acid solution (density 1.25 g cm^-3).

At the anode, the lead atoms release two electrons each to be oxidized to Pb²⁺ ions, which combine with SO₄^2- ions present in the electrolyte and get deposited on the cathode as PbSO4.

Pb(s)Pb²⁺(aq)+2^e-

Pb²⁺(aq)+SO₄^2-(aq) PbSO₄(s)

At the cathode, the electrons from the anode are accepted and PbO2 and hydrogen ions from the electrolyte then under go a redox reaction to produce lead ions and water as follows:

PbO₂(s) + 4H⁺ (aq) + 2e^-  Pb²⁺ (aq) + 2H₂O (l)

The Pb²⁺ ions then combine with SO₄^2- ions and they both deposit at the cathode as PbSO₄. When both electrodes are completely covered with PbSO₄ deposits, the cell wll cease to discharge any more current until it is recharged.

During the process of recharging, the anode and the cathode of the external electrical source are connected to the anode and the cathode of the cell respectively. The overall reaction is

2PbSO₄(s)+2H₂O(l)  Pb(s)+PbO₂(s)+4H⁺(aq)+2SO₄^2-(aq)

Alkaline Battery:

It is a dry alkaline cell, which uses zinc and manganese dioxide as reactants. Zinc rod serves as the anode and manganese dioxide functions as the cathode. The electrolyte, however contains potassium hydroxide and is therefore basic alkaline).

Silver Oxide Battery:

These tiny and rather expensive batteries have become popular as power sources in electronic watches auto exposure cameras and electronic calculators. The cathode is of silver oxide, Ag₂O, and the anode is of zinc metal.

Nickel Cadmium Cell:

A strong cell that has acquired wide spread use in recent years is the NICAD or Nickel Cadmium battery. The anode is composed of cadmium, which undergoes oxidation in an alkaline electrolyte.

The cathode is composed of NiO₂ which undergoes reduction.

Fuel cells:

These are other means by which chemical energy may be converted into electrical energy. Gaseous fuels undergo reactions such as H₂ and O₂.These are used in space vehicles. The electrodes are hollow tubes made of porous compressed carbon coated with platinum which acts as a catalyst. Electrolyte is KOH. H2 is oxidized to water and O2 is reduced to hydroxide ions. Fuel cell is operated at high temperature so that the water produced as a product of the cell reaction evaporates and may be condensed and used as drinking water for an astronaut.

oxidation Number or State:

it is the apparent charge on an atom of an element in a molecule or an ion. It may be positive or negative or zero.

Rules for assigning oxidation number:

The oxidation number of all elements in the free state is zero.

The oxidation number of an ion consisting of a single element is the same as the charge on the ion.

For example, the oxidation number of

K⁺, Ca²⁺, Al³⁺, S^2- are +1, +2, +3, -2, respectively.

The oxidation number of hydrogen in all its compounds (except metal hydrides) is + 1, in metal hydrides it is – 1.

The oxidation number of oxygen in all its compound (except in peroxides, Of₂ and super oxides) is – 2. it is – 1 in peroxides and +2 in OF₂. in case of super oxide it is – ½.

·       In neutral molecules, the algebraic sum of the oxidation number of all the elements is zero.

·       In ions, the algebraic sum of oxidation number equals the charge on the ion.

Oxidation:

·       Loss of electrons causes an increase in oxidation number.

·       Loss of hydrogen or gain of oxygen is called oxidation.

·       Generally oxidation takes place at the anode.

Reduction:

·       Gain of electrons causes a decrease in oxidation number

·       Gain of hydrogen or loss of oxygen is called reduction.

·       Generally reduction takes place at the cathode.

Cell potential:

The cell potential is the sum of the reduction potential of the specie that is being reduced and the oxidation potential of that specie which is being oxidized. In other word, the cell potential is the difference of the reduction potentials of both the species which are taken as electrodes.

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