Ravi Verma Ppt

8/3/2019 Ravi Verma Ppt

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NAME: RAVI R. VARMA

ROLL NO: 925

CLASS : BE-4SUBJECT: E.M.C.

SEMINAR ON

POTENTIAL OF AN ELECTRODE AND EVALUATION OF CELLPOTENTIAL

GUIDED AND PROGRESSIVELY INITIATED BY:

PROF. V. V. MATHANE SIR


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BACKGROUND: Different reactivity of metals.

Example of magnesium and copper to understand thereactivity.

Electrode potential and its aspects.

Need of standard value to calculate electrode potential Standard hydrogen electrode.

Standard electrode potentials.

Electrode potentials and its measurement.

Calculation of cell potentials.

Example.


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When metals react, they give away electrons and form positiveions. This particular topic sets about comparing the ease with

which a metal does this to form hydrated ions in solution.

For example we compare the tendency to lose electrons inmagnesium and copper.

Consider the two reactions.

These reactions shows the tendency of the two metals to loseelectrons.


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There will be some tendency for the magnesium atoms toshed electrons and go into solution as magnesium ions.

The electrons will be left behind on the magnesium.


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Some of them will be attracted enough that theywill reclaim their electrons and stick back on tothe piece of metal

A dynamic equilibrium will be established whenthe rate at which ions are leaving the surface isexactly equal to the rate at which they are joining

it again. At that point there will be a constant negative

charge on the magnesium, and a constant numberof magnesium ions present in the solution

around it.


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Simplifying the diagram to get rid of the "bites"out of the magnesium, you would be left with a

situation like this:

How would this be different if you used a piece ofcopper instead of a piece of magnesium?


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In the magnesium case, there is a lot of differencebetween the negativeness of the metal and thepositiveness of the solution around it. In the

copper case, the difference is much less. Thispotential difference could be recorded as a

voltage - the bigger the difference between thepositiveness and the negativeness, the bigger the

voltage.

Unfortunately, that voltage is impossible tomeasure!

So consecutively there was need for a standardthing to compare these values.


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Going back to the magnesium and copper equilibria:

All we need to know is that the magnesium

equilibrium lies further to the left than the copperone. We need to know that magnesium shedselectrons and forms ions more readily than copperdoes.

That means that we don't need to be able to measurethe absolute voltage between the metal and thesolution. It is enough to compare the voltage with astandardised system called a reference electrode.


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The standard hydrogen electrode


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What is happening?

As the hydrogen gas flows over the porous platinum, anequilibrium is set up between hydrogen molecules and hydrogenions in solution. The reaction is catalysed by the platinum.

This is the equilibrium that we are going to compare all theothers with.

Standard conditions

The position of any equilibrium can be changed by changingconditions. That means that the conditions must be standardisedso that you can make fair comparisons.

The hydrogen pressure is 1 bar (100 kPa). (You may find 1atmosphere quoted in older sources.) The temperature is 298 K(25C).

The concentration of the hydrogen ions in solution is alsoimportant. Changing concentrations is one of the ways ofchanging the position of an equilibrium. Throughout this topic,all ion concentrations are taken as being 1 mol dm-3.


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Magnesium has a much greater tendency to form its ions than hydrogen does. The

position of the magnesium equilibrium will be well to the left of that of the hydrogen

equilibrium.

That means that there will be a much greater build-up of electrons on the piece of

magnesium than on the platinum

Connection of SHE with an electrode of magnesium.


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In Case of copper, there is less difference between the electrical

charges on the two electrodes, so the voltage measured will be less. Thistime it is only 0.34 volts.

This emf of a cell measured under standard conditions called StandardElectrode Potential and is given the symbol Ecell.


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The values that we have just quoted for the twocells are actually the standard electrode

potentials of the Mg2+ / Mg and Cu2+ / Cu systems.

The emf measured when a metal / metal ionelectrode is coupled to a hydrogen electrodeunder standard conditions is known as thestandard electrode potential of that metal / metal

ion combination.


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The voltage existing between an electrode and thesolution or electrolyte in which it is immersed.

The equilibrium potential difference between twoconducting phases in contact, most often an

electronic conductor such as a metalor semiconductor on the one hand, and an ionicconductor such as an electrolyte solution (asolution containing ions) on the other.

In short, it is the interfacial potential differencebetween the electrode and the electrolyte.
http://www.answers.com/topic/semiconductorhttp://www.answers.com/topic/electrolytehttp://www.answers.com/topic/electrolytehttp://www.answers.com/topic/semiconductor


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The interfacial potential difference is usually the

consequence of the transfer of some chargecarriers from one conducting phase to the other.

For a metal in contact with its metal ionsofvalencez, the potential difference Ecan be

expressed in terms of a standard potential E andthe concentration c of these ions in solutionthrough the Nernst equation, where R is the gasconstant, Tis the absolute temperature, and Fis

the Faraday.
http://www.answers.com/topic/valencyhttp://www.answers.com/topic/nernst-equationhttp://www.answers.com/topic/nernst-equationhttp://www.answers.com/topic/valency


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Calculation of the half cell potential using standardoxidation or reduction potential and the NernstEquation.

Write one as oxidation and other as reduction. Theoxidation potentials above hydrogen in the series have

a positive sign while those below these have negativesign.

The total cell reaction and cell e.m.f are obtained bythe algebraic sum of the half cell reactions and the

potentials.A positive cell potential indicates that cell reactionwritten is spontaneous.


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For Zn Zn+2 + 2e Ecell = 0.762.

Therefore for zinc electrode, we calculate the halfcell potential from nernst equation as

Zn Ecell= -0.7915For Cu+2 +2e Cu Ecell = 0.345

Similarly from nernst equation

Cu Ecell =+0.286

Therefore the total cell potential= Zn Ecell + Cu Ecell.


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0.7915+0.286 = +1.07756 volts.

For the reaction

Zn + Cu+2 Cu + Zn+2Cell e.m.f = +1.07756 volts.

This reaction is spontaneous.

In this way we can find out the feasibility of the reactionsby using the electrode potentials.


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References:

Class discussion

Introduction to electrometallurgy andcorrosion.

By sharan and naryain.

Internet search


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Thank youand

Good luck to all.

JAI HIND !!!!

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Ravi Verma Ppt