Lacture OF Electrochemistry

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Electrochemistry

• Maruthupandi M INDIAN_TN_MDU

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Electrochemistry

• Plan
• 1. Electrode processes. Electrode potential.
• 2. Different types of electrodes.
• 3. Cell potential.
• 4. Galvanic cells. Cell potential or EMF.
• 5. The kinetics of electrochemistry processes.

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• The changes in which electrical energy is
  produced as a result of chemical change. The
  devices used to produce electrical energy from
  chemical reactions are called electrical cells,
  galvanic or voltic cells.
• In these cells, oxidation and reduction reaction
  reactions occur in separate containers called
  half cells and the red-ox reaction is
  spontaneous.
• The arrangement consists of two beakers, one of
  with contains 1,0 M solution of zinc sulphate and
  the other 1,0 M solution of copper sulphate. A
  zinc rod is dipped into ZnSO4 solution while a
  copper rod is dipped into CuSO4 solution. These
  metallic rods are known as electrodes.

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• The metallic rods in the beaker are connected to
  the ammeter by means of an insulated wire through
  a key. Ammeter is used to know the passage of
  current which moves in opposite direction to the
  flow of electrons. The solution in the two
  beakers are connected by an inverted U-tube
  containing saturated solution of some electrolyte
  such as K?l, KNO3, NH4OH which does not undergo a
  chemical change during the process.
• The two openings of the U-tube are plugged with
  some porous material such as glass wood or
  cotton. The U-tube which connects the two glass
  beakers is called a salt-bridge.

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When the circuit is completed by inserting the
key in the circuit, it is observed that electric
current flows through external circuit as
indicated by the ammeter. The following
observations are made Therefore, the current
flows from copper to zinc N.B. The flow of
electric current is taken opposite to the flows
of electrons
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There observation can be explained as During the
reaction, zinc is oxidized to Zn2 ions which go
into the solution. Therefore, the zinc rod
gradually loses its weight. The electrons
released at the zinc electrode move towards the
other electrode through outer circuit. Here,
these are accepted by Cu2 ions of CuSO4 solution
which are reduced to copper.
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• The zinc electrode where electrons are released
  or oxidation occurs s called anode while the
  copper electrode where electrons are accepted or
  reduction occurs is called cathode.
• The two containers involving oxidation and
  reduction half reactions are called half cells.
  The zinc rod dipping into a ZnSO4 solution is
  oxidation half cell and the copper electrode
  dipping into a CuSO4 solution is reduction half
  cell

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• N.B. The galvanic cells which consists of the
  zinc rod dipping into a ZnSO4 solution and the
  copper electrode dipping into a CuSO4 solution is
  Daniell cell.
• Its formula is
• Salt bridge and its function. Its usually an
  inverted U-tube filled with concentrated solution
  of inert electrolyte. The essential requirements
  of electrolyte are
• The mobility of the anion and cation of the
  electrolyte should be almost same.
• The ions of the electrolyte are not involved in
  electrochemical change.
• The ions do not react chemically with the species
  of the cell.

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• Generally, salts like KCl, KNO3, etc. are used.
  The seturated solutions of these electrolytes are
  prepared in agar agar jelly or gelatin. The jelly
  keeps the electrolyte in semi-solid phase and
  thus prevents mixing.
• The important functions of the salt bridge are
• Salt bridge completes the electrical circuit.
• Salt bridge maintains electrical neutrality of
  two half cell solution.
• The accumulation of charges in the two half cells
  (accumulation of extra positive charge in the
  solution around the anode according to the
  realizing of Zn2 in excess and accumulation of
  extra negative charge in the solution around the
  catode due to excess of SO42- ) is prevented by
  using salt bridge, which provides a passage for
  the flow of the charge in the internal circuit.

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• REPRESENTATION OF AN ELECTROCHEMICAL CELL
• An electrochemical cells or galvanic cell
  consists of two electrodes anode and cathode.
  The electrolyte solution containing these
  electrodes are called half cells.
• The following conventions are used in
  representing an electrochemical cell
• A galvanic cell is represented by writing the
  anode (where oxidation occurs) on the left hand
  side and cathode (where reduction occurs) on the
  right hand side.
• The anode of the cell is represented by writing
  metal or solid phase first and then the
  electrolyte (or the cation of the electrolyte)
  while the cathode is represented by writing the
  electrolyte first and then metal or solid phase.
  
• The salt bridge which separates the two half
  cells is indicated by two vertical lines.

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• Electrode Potential and E.M.F. of a galvanic cell
• Electrode Potential
• The flow of electric current in an
  electrochemical cell indicates that a potential
  difference exists between two electrodes.

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• If the metal has relatively high tendency to get
  oxidised, its atom will lose electrons readily
  and form Cu2 ions, which go into the solution.
  The electrons lost on the electrode would be
  accumulated on the metal electrode and the
  electrode acquires a slight negative charge with
  respect to the solution. Some of the Cu2 ions
  from the solution will take up electrons and
  become Cu atoms. After some time, an equilibrium
  will be established as
• When such an equilibrium is attained, it results
  in separation of charges (negative on the
  electrode with respect to the solution).

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Forming the double layer
1- metal 2 - solution
Cu
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• Similarly, if the metal ions have relatively
  greater tendency to get reduced, they will take
  electrons from the electrode. As a result, a net
  positive charge will be developed on the
  electrode with respect to the solution. This will
  also result into separation of charges (positive
  on the electrode with respect to the solution).
• Due to separation of charges between the
  electrode and the solution, an electrical
  potential is set up between metal electrode and
  its solution.
• The electrical potential difference set up
  between the metal and its solution is known as
  electrode potential.

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• The mechanism of the double layer forming
• For example of the copper electrode is dipped
  into CuSO4 solution.
• The chemical potential of coppers ions in the
  metal and in the solution is not equal. The
  chemical potential of coppers ions in the metal
  at the given temperature is stable value, the
  chemical potential of coppers ions in the
  solution depends on the solutions concentration.
• If the at the given concentration of solution the
  chemical potential of coppers ions in the
  solution is greater than the chemical potential
  of these ions in the metal . Then at the dipping
  of the metal in the solution some quantity of
  Cu2 ions are hydrated and transferred on the
  metal according to that positive charge forming
  on the metals surface. Sulfates anions are
  attracted to metals surface, they courses the
  negative charge. These processes cause the double
  electrical layer and related with it the
  potential difference.

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Forming the double layer
Cu
(-)

• - - -
• -
• -
• -

- - - - -
CuSO4
-
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The electrode potential may be of two types 1.
Oxidation potential The tendency of an
electrolyte to lose electrons or to get oxidised

• 2. Reduction potential. The tendency of an
  electrode to gain electrons or to get reduced.

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• E.M.F. or Cell Potential of a Cell
• The difference between the electrode potentials
  of the two electrodes constituting an
  electrochemical cell is known as electromotive
  force (e.m.f.) or cell potential of a cell. This
  acts as a driving force for the cell reaction.
  The potential difference is expressed in volts.

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• Therefore, the cell potential or e.m.f. arises
  from the difference in the tendencies of the two
  ions to get reduced.
• It is equal to the reduction potential for the
  substance that actually undergoes reduction minus
  the reduction potential of the substance that
  undergoes oxidation.
• Thus, e.m.f. of a cell may be defined as the
  potential difference between two electrodes of
  the cell when either no or negligible current is
  allowed to flow in the circuit.

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• Standard electrode potential
• Since a half cell in an electrochemical cell can
  work only in combination with the other half cell
  and does not work independently, it is not
  possible to determine the absolute electrode
  potential of an electrode. We can, therefore,
  find only the relative electrode potential.
• This difficulty can be solved by selecting one of
  the electrodes as a reference electrode and
  arbitrarily fixing the potential of this
  electrode as zero. For this purpose, reversible
  hydrogen electrode has been universally accepted
  as a reference electrode. It is called standard
  hydrogen electrode (S.H.E) or normal hydrogen
  electrode (N.H.E.)

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• Standard hydrogen electrode. It consists of
  platinum wire sealed in a glass tube and has a
  platinum foil attached to it. The foil is coated
  with finely divided platinum and acts as platinum
  electrode. It is dipped into an acid solution
  containing H ions in 1 M concentration (1M
  HCl). Pure hydrogen gas at 1 atmospheric pressure
  is constantly bubbled into solution at constant
  temperature of 298K. The surface of the foil acts
  as a site for the reaction.

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• The electrode potential of an electrode can be
  determined by connecting this half cell with a
  standard hydrogen electrode. The electrode
  potential of the standard hydrogen electrode is
  taken as zero.
• The electrode potential of a metal electrode as
  determined with respect to a standard or normal
  hydrogen electrode is called standard electrode
  potential (E0). Standard electrode potentials are
  always associated with the reduction occurring at
  the electrodes.

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• In this case, the electrons flow from zinc
  electrode to hydrogen electrode and therefore,
  the zinc electrode acts as anode and S.H.E. acts
  as a cathode. The cell may be represented as

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• 2. Measurement of Electrode Potential of Cu2Cu
  Electrode
• In this case, the hydrogen has greater tendency
  to lose electrons. Therefore, oxidation occurs at
  hydrogen electrode and reduction occurs at copper
  electrode.

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• It may be noted that it is not always convenient
  to use standard hydrogen electrode as reference
  electrode because of experimental difficulties in
  its preparation and use.

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• N.B. The standard electrode potentials given in
  the following table are measured in their
  standard states when the concentration of the
  electrolyte solutions are fixed as 1M and
  temperature is 298 K.

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• Application of the Electrochemical
  (electromotive) series
• Relative strengths of oxidising and reducing
  agents.
• The substances which have lower reduction
  potentials are stronger reducing agents while
  have higher reduction potentials are stronger
  oxidations agent.
• 2. Calculation of the E.M.F. of the cell.
• E0cell E0(cathode) - E0(anode)

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• 3. Predicting feasibility of the reaction.
• In general, a red-ox reaction is feasible only if
  the species to release electrons must have lower
  reduction potential as compared to the species
  which is to accept electrolytes.
• 4. To predict whether a metal can liberate
  hydrogen from acid or not.
• In general, only those metals can liberate
  hydrogen from the acid which have negative values
  of reduction potentials , - E0 values.

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• Dependence of electrode and cell potentials on
  concentration Nernst equation
• The electrode potentials depend on the
  concentration of the electrolyte solutions.

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or
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Nerst equation

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• In general, for an electrochemical cell reaction
• The Nerst equation may be written as

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• The value of a, b, c, d and n are obtained from
  the balanced cell reactions.
• N.B. It must be remembered that while writing the
  Nerst equation for the overall cell reaction, the
  log term is the same as the expression for the
  equilibrium constant for the reaction. However,
  some books use the expression in the reverse form
  as the expression for the equilibrium constant
  but, sign after E0 is changed.

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• Equilibrium constant from Nernst equation
• The e.m.f. of the cell may be used to calculate
  the equilibrium constant for the cell reaction.
  At equilibrium, the electrode potentials of the
  two electrodes become equal so that e.m.f. of the
  cell is zero. Consider the following redox
  reaction

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Significance of Kc. The value of Kc gives the
extent of the cell reaction. If the value of Kc
is large, the reaction proceeds to large extent.
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• Electrochemical cell and free energy
• In electrochemical cells, the chemical energy is
  converted into electrical energy. The cell
  potential is related to free energy change. In an
  electrochemical cell, the system does work by
  transferring electrical energy through an
  electric circuit.

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• Where is the standard free energy for
  the reaction.
• Significance. The above equation helps us to
  predict the feasibility of the cell reaction. For
  a cell reaction to be spontaneous, must be
  negative. This means that E must be positive for
  a spontaneous cell reaction.

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• SOME COMMERCIAL CELLS
• One of the main uses of galvanic cells is the
  generation of portable electrical energy. These
  cells are also popularly known as batteries. The
  term battery is generally used for two or more
  galvanic cells connected in series. Thus, a
  battery is an arrangement of electrochemical
  cells used as an energy source. The basis of an
  electrochemical cell is an oxidation-reduction
  reaction.
• Types of commercial cells.
• Primary cells
• Secondary cells

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• Primary cells. In these cells, the electrode
  reactions cannot be reversed by an external
  electric energy source. In these cells, reactions
  occur only once and after use thaey become dead.
  Therefore, they are not chargeable. Examples are
  dry cell, mercury cell.
• Secondary cells (storage cells or accumulators).
  In the secondary cells, the reaction can be
  reversed by an external electric source.
  Therefore, these cells can be recharged by
  passing electric current and used again and
  again. Examples are lead storage battery and
  nickel-cadmium storage cell.
• The most popular example is of lead storage cell
  which is used in automobiles.

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• Each battery consists of a number of voltaic
  cells connected in series. Three to six such
  cells are generally combined to get 6 to 12 volt
  battery. In each cell, the anode is a grind of
  lead packed with divided spongy lead and the
  cathode is a grind of lead packed with PbO2.
• The electrolyte is aqueous solution of sulfuric
  acid (38 by mass) having a density 1,30 g ml-1
  sulfuric acid. When the lead plates are kept for
  sometimes, a deposit of lead sulphate is formed
  on them.
• At the anode, lead is oxidised to Pb2 ions and
  insoluble PbSO4 is formed. At the cathode PbO2 is
  reduced to to Pb2 ions and PbSO4 is formed.
• The following reactions take place in the lead
  storage cell

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• During the working of the cell, PbSO4 is formed
  at each electrode and sulphuric acid is used up.
  As a result, the concentration of H2SO4 decreases
  and the density of the solution also decreases.
  When the density of H2SO4 falls below 1.2 g
  ml-1, the battery needs recharge.
• Recharge the Battery
• The cell can be charge by passing electric
  current of a suitable voltage in the opposite
  direction. The electrode reaction gets reversed.
  As a result, the flow of electrons gets reversed
  and lead is deposited on anode and PbO2 on the
  cathode. The density of sulphuric acid also
  increases. The reaction can be written as

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• The most important types of electrodes are
• 1. The first reference electrode Metal-metal ion
  electrodes and gas-ion electrodes
• 2. The second reference electrode
  Metal-insoluble salt-anion electrodes
• 3. The third reference electrode inert
  "oxidation-reduction" electrodes
• 4. Membrane electrodes

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• The metal - metal ion electrode consists of ?
  metal in contact with its ions in solution.
• An example is ? piece of silver metal immersed in
  ? solution of silver nitrate. The diagram for
  such an electrode serving as ? cathode (it would
  appear at the right in ? cell diagram) is
  Ag(aq) ? Ag(s)
• and the cathode half-reaction is Ag(aq)
  e-?Ag(s)
• in which the electrons ??m? from the external
  circuit. When this electrode serves as an anode,
  it is diagramed as Ag(s) ? Ag(aq)
• (as it would appear at the left in ? cell
  diagram), and its half-reaction equation is
• Ag(s) ? Ag(aq) ?-
• In general the first reference electrons can be
  represented as
• Mz/M. The half reduction reaction is
• Mz ze ? M
• Following convention the half reaction that
  occurs on the electrode is written as a reduction
  reaction

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• Nernst equation for these type electrodes is

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• The gas-ion electrode (Standard hydrogen
  electrode)
• Hydrogen electrode that works at the following
  conditions
• 1, pH 101,3 kPa, T 250C 298K is
  called standard.
• Electrochemical potential of this electrode
  depends on the hydrogen ions concentration.
  However the standard potential of this electrode
  equals o and the valency of hydrogen equals 1
  (n1) Nernst equation is
• Ecell 0,059 log H - 0,059 pH
• Measuring of pH to use potentiometric method of
  the determination of hydrogen ions concentration.
  
• This method is based on the measuring of e.m.f of
  the cell which consists of the reduction
  electrode (calomel) and the electrode which has
  dependence on the hydrogen ions concentration
  (gas-ion electrode, glass electrode,
  quinonhydrone electrode )

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• 2. In the metal-insoluble salt-anion electrode, ?
  metal is in contact with one of its insoluble
  salts and also with ? solution containing the
  anion of the salt.
• An example is the so-called silver - silver
  chloride electrode, written as ? cathode as
• Cl-(aq) ? AgCl(s) ? Ag(s)
• for which the cathode half-reaction is
• AgCl (s) ?- ? Ag(s) Cl- (aq)
• EAg,AgCl Cl- E0Ag/AgCl - 0.059 lg aCl-
• E0Ag/AgCl 0,2224

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Silver - silver chloride electrode
Ag, is covered by the layer of nonsoluble AgCl
?Cl solution
KCl, AgCl Ag
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• Calomel electrode consists of the mercury and
  calomel past that is dipped in potassium chloride
  solution. It is often used as a reference
  electrode to determine the standard electrode
  potential ( more often than hydrogen electrode).
  Its scheme is
• Cl-Hg2Cl2, Hg
• The half reaction is
• Hg2Cl2 2e ? 2Hg 2Cl-
• Ecell E0 - 0.059 lg aCl-
• As a rule to use the calomel electrodes that
  contain
• 0,1 M, 1 M and saturated solution of potassium
  chloride. Their standard potential at 298K equal
• 0,337 0,2801 0,2412 V.

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• 3. An inert oxidation-reduction electrode
• It consists of ? strip, wire, or rod of an inert
  materiel (Pl, Au, Ir) in contact with ?
  solution, which contains ions of ? substance is
  two different oxidation states (oxidation and
  reduction form). The difference between general
  metal electrode and ox-red electrode is that
  ox/red electrode does not take place in ox-red
  reaction which exist in solution but is the
  electrons conductor. For example Pt Sn2, Sn4
  or Pt Fe2, Fe3
• Ox ze ? Red
• There are two types of ox-red electrodes
• Simple Fe2, Fe3 Pt Fe3 e ? Fe2

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Scheme of ox-red electrode (the third reference
electrode)
Pt
Fe3 , Fe2 Pt
Fe3 e Fe2
FeCl3 FeCl2
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• 2. Complex ox-red electrode there is changing the
  charge and the composition of the ions
• Mn2, MnO4-, H Pt MnO4- 8H5e?Mn24H2O
• Example is quinonhydrone electrode.
• It is prepared by the platinum strip or wire
  which is contained in the glass tube. The
  electrode is dipped in the solution with unknown
  pH that is needed to determine and to add some
  quinonhydrones crystalls in this solution.
• Quinonhydrone is a crystalline product which
  consists of quinone (benzoquinone) ?6?4?2 and
  hydroquinone C6H4(OH)2. It is less solubility in
  water and decomposes into quinone and
  hydroquinone in the solution. In the saturated
  solution equal molar mixture of quinone and
  hydroquinone is formed.

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Quinonhydrone electrode (the third reference
electrode)
?6?4?2, ?6?4(??)2, H Pt
?6?4?2 2? 2? ?6?4(??)2
Including that the activity of quinone and
hydroquinone is equal in the seturated solution,
we have
Quinonhydrone
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The scheme of quinonhydrone cell with one
electrolyte
Pt, ?2 quinhydr, H KCl KCl,Hg2Cl2 Hg
Ecell E quinhydr - Ecalomel
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• The glass electrode for pH measurements. The cell
  consists of a glass indicator electrode and ?
  saturated calomel reference electrode, both
  immersed in the solution whose pH is to be
  determined. The indicator electrode consists of ?
  thin, ??-sensitive glass membrane sealed onto one
  end of ? heavy-walled glass or plastic tube. ?
  small volume of dilute hydrochloric acid
  saturated with silver chloride is contained in
  the tube (in some electrodes this solution is ?
  buffer containing chloride ion). A silver wire in
  this solution forms ? silver/silver chloride
  reference electrode, which is connected to one of
  the terminals of ? potential-measuring device.
  The calomel electrode is connected to the other
  terminal.
• Show that the system contains two reference
  electrodes (1) the external calomel electrode
  and (2) the internal silver/silver chloride
  electrode. Although the internal reference
  electrode is part of the glass electrode, it is
  not the pH-sensing element. Instead, it is the
  thin glass membrane at the tip of the electrode
  that responds to pH.

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The make the cell with glass and calomel
electrodes and measuring its e.m.f can be
determined pH of solution. glass electrodes
constant which is depended on the electrode
nature. The constant is fined according to the
graph which is plotted between the Ecell and pH
ordinates.
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The mechanism of the diffusion potential
HCl 1 ?
HCl 0.1 ?

- - - - - - - - -
H
Cl-
where a1 gt a2
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The scheme of concentrated cell
_

()Ag AgNO3 AgNO3 Ag(-) C1 gt C2




Ecell E2 Ag/Ag E1 Ag/Ag 0.059 lg (a2
/ a1)

NO3-


AgNO3, C2
AgNO3, C1
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• The electric circuit with transfer and without it
• The electric circuit without transfer is when the
  electrodes are dipped in one solution.
• Examples PtH2HClAgCl, Ag hydrogen-silver-sil
  ver chloride electrode
• Pb PbSO4 H2SO4(aq) PbO2 Pb the lead
  storage battery
• The electric circuit with transfer is when the
  electrodes are dipped in different solution which
  contact with each other.
• Examples ZnZnSO4CuSO4Cu Daniell cell
• PtH2HClKClHg2Cl2,Hg hydrogen-calomel cell

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Thanks for attention