ELECTRICAL CELL

 

Week: THREE                    Date:     20-24/05/2019                                                   Time:

Period:                                 Duration: 1 HR 20 MIN.                                                 Average age of learners: 16YEARS

Subject:                               CHEMISTRY                                                                        Class: SS TWO

Topic:                                    ELECTROCHEMICAL CELLS

Sub topic:  Reference materials:

(1) ESSENTIAL CHEMISTRY, TONALD PUBLISHERS, I. O ODESINA

(2) NEW SCHOOL CHEMISTRY, AFRICAN FIRST PUBLISHERS, OSEI YAW ABABIO

(3) INTERNET

Instructional materials:

Entry behavior: The students have been taught chemical reaction

Behavioural objective: At the end of the lesson the students should be able to:

i.                     Define electrochemical cell

ii.                   Differentiate between galvanic cell and electrolytic cells

iii.                  Define electrode potential

iv.                 Solve problems involving e.m.f of a cell

v.                   Define primary and secondary cell with relevant examples.

CONTENT

ELECTROCHEMICAL CELLS

Electrochemical cell is the set – up in which chemical energy is converted to electrical energy. It consist of two half cells: an oxidation half – cell reaction and a reduction half-cell reaction. The overall redox reaction result in a flow of electrons i.e an electric current.

An example of a electrochemical cell is a zinc electrode dipping into a solution of ZnSO_4, connected to a copper electrode dipping into a solution of CuSO₄. The two solutions are separated by a porous partition. The porous partition allows electrical contact but prevents excessive mixing of the electrolyte by inter diffusion.

The atoms at the zinc electrodes undergo oxidation and loose two electrons each to form zinc ions (Zn²⁺) which go into the solution. The zinc electrode becomes negatively charged and functions as the negative electrode or anode. The copper (ii) ions become reduced by gaining two electrons each to form metallic copper which is deposited on the copper on the copper electrode. The copper electrode thus become positively charged and functions as the positive electrode or cathode.

At the zinc electrode (anode)

Zn(s)                   Zn2+    +    2e-  (oxidation)

The anode slowly becomes reduced in  size as the metallic zinc is converted to zinc ions.

At the copper electrodes (cathode)

Cu2+  +  2e-             Cu (reduction)

In the electrochemical cells oxidation always occurs at the anode and reduction at the cathode. Electrons flow from the anode to cathode, the negative electrode is the anode, while the positive electrode is the cathode.

Difference in Electrolytic Cell and Galvanic Cell:

Electrolytic Cell	Galvanic cell
Electrical energy is converted into chemical energy. These types of cell involve spontaneous chemical reactions like batteries.	Chemical energy is converted into electrical energy. These types of cell involve non-spontaneous chemical reaction and even need external source for electron flow.
Anode positive electrode. Cathode negative electrode	Anode negative electrode. Cathode positive electrode.
Ions are discharged on both the electrodes.	Ions are discharged only on the cathode
If the electrodes are inert, concentration of the electrolyte decreases when the electric current is circulated	Concentration of the anodic half-cell increases while that of cathodic half-cell decreases when the two electrodes are joined by a wire
Both the electrodes can be fitted in the same compartment	The electrodes are fitted in different compartment

 

STANDARD ELECTRODE POTENTIALS

The standard electrode potential of metal ions is measured by connecting it to the 2H⁺_aq          H_g system by a salt bridge and voltameter to show the reading .

Thus, the standard electrode potential of a metal ions is the potential difference set-up between the metal and one molar solution of its ions at 298K.

THE E.M.F OF A CELL

When two half-cells are joined together through a salt bridge, the e.m.f of a cell formed is the algebraic difference between the two potential difference.

                                E^o_Total    =  E^o_reduction  -   E^o_oxidation

                                E^o_Total    =  E^o_right  -   E^o_left

                                E^o_Total    =  E^o_cathode  -   E^o_anode

 A positive E.M.F implies that the reaction is thermodynamically feasible.

We can predict the livelihood of a reaction if two systems in the electrochemical series are linked by cells.

 

Remember, the system which is lower in the series will lose electron, and the one higher in the series will gain electron e.g

A system containing Co2+aq/ Cos and Ni2+aq/ Nis. Eo of Ni = - 0.26 and Co = - 0.28

Co      Co²⁺  +  2e^-               (oxidation)

Ni²⁺   +   2e^-       Ni               (reduction)

Co  + Ni²⁺           Co²⁺  +   Ni

E^o_Total    =  E^o_reduction  -   E^o_oxidation

                = - 0.26  -  (-0.28)

                = 0.02V

EXAMPLE

Consider the reaction,

2Ag⁺ + Cd →   2Ag + Cd²⁺

The standard electrode potentials for Ag⁺ --> Ag and Cd²⁺ --> Cd couples are 0.80 volt and -0.40 volt, respectively.

(i) What is the standard potential E^o for this reaction?

(ii) For the electrochemical cell in which this reaction takes place which electrode is negative electrode?

Solution:             

(i) The half reactions are:

2Ag⁺  + 2e^- →   2Ag.

 Reduction

Cathode)

E^o_Ag⁺/Ag =0.80  volt          (Reduction potential)

Cd → Cd²⁺  + 2e^-,

Oxidation

(Anode)

E^o_Cd⁺/Cd = -0.40 volt               (Reduction potential)

or     E^o_Cd⁺/Cd² = +0.40 volt

E^o = E^o_Cd⁺/Cd² + E^o_Ag⁺/Ag = 0.40+0.80 = 1.20  volt

(ii) The negative electrode is always the electrode whose reduction potential has smaller value or the electrode where oxidation occurs. Thus, Cd electrode is the negative electrode.

Purpose of the Salt Bridge

Finally, we must understand the purpose and use of salt bridge. During reaction we have seen that Zinc ions are produced by losing electrons and this Zinc ions comes out in the solution, due to this the net positive charge of the zinc rod beaker increases.

On the same time, the overall negative charge on copper side beaker increases because Cu atom gets deposited on copper rod.

The salt bridge helps to prevent the net accumulation of positive and negative charges on both the sides. Doing so the negative ions from the salt bridge enter Zinc beaker side to reduce the net positive charge. The positive ions from the salt bridge enters copper side beaker to decrease net negative charge there.

If this was not done then due to accumulation of net positive and negative charges on both the sides the redox reaction will come to an end.

Thus we can tell that although salt bridge do not participate in the reaction directly but it help to maintain continuity of reaction.
 Significance of Salt Bridge

The following are the func­tions of the salt bridge:                               

It helps to complete the connection of both half cells.

It prevents diffusion of the solutions in both the half cells.

It helps to build electrical neutrality.

It avoids liquid-liquid junction potential. (The potential difference arising when two liquids are in contact to each other.)

Note:

Two parallel vertical lines in a cell reaction indicate the salt bridge.

Zn|Zn²⁺||Cu²⁺|Cu

Salt bridge can be replaced by a porous partition which allows the migration of ions without allowing the solution to intermix.

PRIMARY AND SECONDARY CELLS

PRIMARY CELLS: These are cells that produce electric currents by using up the chemicals of which they are composed and have to be replaced after some time i.e cannot be recharge and do not supply a steady current. Examples are radio battery, dry cell, alkaline battery, button battery, Le clanche cell, Daniel cell, torch battery, bicycle battery

SECONDARY CELLS: They are cells that are rechargeable by passing a direct current through them. It supply a steady current and these cells last long. Examples are Lead-acid (storage) cell, lead-acid accumulator, lithium-ion battery, battery in rechargeable lamp, car battery, nickel-metal hydride cell, phone battery, nickel-cadmium cell, (hydrogen-oxygen) fuel cell, Aluminium-air cell

 Daniel cell

The cells which produce electrical energy from chemical reactions are known as galvanic or voltaic cells. A galvanic cell in which one electrode is Zn plate in   (or   ion) solution and the other is Cu plate in   ion) solution is called the Daniel cell.

A schematic diagram of the cell is shown below:

Daniell cell

The two half reactions,

I.e.

 At anode

 at cathode

are made to take place simultaneously with the electron transfer occurring through an extemal conducting wire. The two solutions (i.e.,  ) are linked by an inverted U-shaped tube as salt bridge. The tube is filled with a solution of an electrolyte like KCl,   to which gelatin or agar-agar has been added to obtain a semi solid paste. Sometimes, a porous partition is used in place of salt bridge, which separates the two solutions.

At anode, metallic zinc passes into solution as   ions and liberates electrons. while at cathode, ions gain these electrons and are discharged as metallic copper. Hence net reaction may be represented as:

It is the chemical energy of this reaction which is converted into electrical energy.

Introduction and working method of Leclanché Cell

Leclanché Cell was invented by Frances Scientist Georges Leclanché in 1866.^[1][pdf] There is a glass container which is filled with ammonium chloride. in the solution, a rod of mercury-coated zinc is submerged which acts as a negative pole. a jar containing foramen made of ceramic is dipped in solution, in which the carbon rod is kept inside. which works in positive duality. in this jar, a mixture of manganese dioxide and granulated carbon is filling.^[2] manganese dioxide acts as a depolarizer and transmits granular carbon electrons to ammonium ions.

When the carbon and zinc rod is added to the circuit then in contact with electrolytes, zinc atoms are ionized by giving two electrons.

                                                              Zn → Zn²⁺ + 2e^-

Ammonium chloride (NH₄Cl) in the electrolyte is thus ionized:

                                                       2NH₄Cl → 2NH₄⁺ + 2Cl^-

Zinc ions make zinc chloride combined with chlorine ions.

                                                        Zn²⁺ + 2Cl^- → ZnCl₃

Ammonium ions (NH₄⁺) flow towards the carbon rod and become Indifferent by taking an electron from the carbon rod, ammonium is converted into hydrogen gas.

                                                       2NH₄⁺ + 2e^- → 2NH₃ + H₂

Manganese dioxide converts hydrogen gas into water, which does not cause the action of polarization.

                                                     2MnO₂ + H₂ → Mn₂O₃ + H₂O

Full Reaction

            Zn + 2NH₄Cl + 2MnO₂ → ZnCl₂ + 2NH₃ + Mn₂O₃ + H₂O + Energy ^[3]

Thus, the excess of electrons on the zinc rod and the lack of electrons on the carbon rod decreases, whereby the potentiality of carbon rod becomes more than zinc rod potential. So the electrons start flowing towards the carbon rod from zinc rod. The Electromotive force of the cell is 1.5 volt.

Note: Manganese dioxide is in the solid state. It can not quickly convert hydrogen into water. Some hydrogen gas starts to accumulate in the carbon rod, resulting in partial polarity and the electric current becomes dim. but if the cell is rested, the hydrogen frozen on the carbon rod becomes oxidized in water. due to this defect, this cell is used in the telephone, telegram, electric bell.^[4]

PRESENTATION

Step I: The teacher explains electrochemical cell

Step II: The teacher differentiates between galvanic cell and electrolytic cells

Step III: The teacher leads the students to find the e.m.f of a cell

Step IV: the teacher explain primary and secondary cells

Step V: The students chorus examples of primary and secondary cells

EVALUATION

The teacher assesses the lessons by asking the following questions:-

1.       Define electrochemical cell

2.       State 2 differences between galvanic  cell and electrolytic cell

3.       Define electrode potential

4.       Differentiate between primary and secondary cell

ASSIGNMENT

   Read about electrolysis