Electrochemistry Plan 1. Electrode processes. Electrode potential. 2. Different types of electrodes. 3. Cell potential. 4. G alvanic cells. Cell potential.

Electrochemistry

Plan1. 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.

Assistant Kozachok S.S prepared

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.

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 KСl, 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.

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

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.

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

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. It’s usually an inverted U-tube filled with concentrated solution of inert electrolyte. The essential requirements of electrolyte are:

a) The mobility of the anion and cation of the electrolyte should be almost same.

b) The ions of the electrolyte are not involved in electrochemical change.

c) The ions do not react chemically with the species of the cell.

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:

a) Salt bridge completes the electrical circuit.

b) 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 SO4

2- ) is prevented by using salt bridge, which provides a passage for the flow of the charge in the internal circuit.

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:

1. 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.

2. 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.

3. The salt bridge which separates the two half cells is indicated by two vertical lines.

Electrode Potential and E.M.F. of a galvanic cellElectrode Potential

The flow of electric current in an electrochemical cell indicates that a potential difference exists between two electrodes.

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).

Forming the double layer

1- metal2 - solutionCu

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.

The mechanism of the double layer forming For example of the copper electrode is dipped into CuSO4

solution.The chemical potential of copper’s ions in the metal and

in the solution is not equal. The chemical potential of copper’s ions in the metal at the given temperature is stable value, the chemical potential of copper’s ions in the solution depends on the solution’s concentration.

If the at the given concentration of solution the chemical potential of copper’s 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 metal’s surface. Sulfate’s anions are attracted to metal’s surface, they courses the negative charge. These processes cause the double electrical layer and related with it the potential difference.

Cu

CuSO4

(-)++ ++ ++ +

- - -- - - -- -

-

-----

Forming the double layer

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

The electrode potential may be of two types:

1. Oxidation potential:The tendency of an electrolyte to lose electrons

or to get oxidised

E.M.F. or Cell Potential of a CellThe 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.

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.

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.)

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.

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.

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:

2. Measurement of Electrode Potential of Cu2+|Cu Electrode

In this case, the hydrogen has greater tendency to lose electrons. Therefore, oxidation occurs at hydrogen electrode and reduction occurs at copper electrode.

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.

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.

Application of the Electrochemical

(electromotive) series:

1. 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)

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.

Dependence of electrode and cell potentials on concentration: Nernst equation

The electrode potentials depend on the concentration of the electrolyte solutions.

or ] (aq)Mn [ log n

0.059 E 0 E

Nerst equationNerst equation

a

nF

RTЕЕ dOxdOx ln0

Re/Re/

an

ЕЕ dOxdOx lg059.00

Re/Re/

thatF

RT,059.03.2

dDcCbBaA ne

In general, for an electrochemical cell reaction:

The Nerst equation may be written as:

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.

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:

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.

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.

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.

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.

1) Primary cells

2) Secondary cells

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.

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:

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:

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

The metal - metal ion electrode consists of а metal in contact with its ions in solution.

An example is а piece of silver metal immersed in а solution of silver nitrate. The diagram for such an electrode serving as а cathode (it would appear at the right in а cell diagram) is: Ag+

(aq) Ag(s)and the cathode half-reaction is: Ag+(aq) + e-Ag(s)in which the electrons соmе 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 а cell diagram), and its half-

reaction equation is:Ag(s) Ag+(aq) + е-

In general the first reference electrons can be represented as:Mz+/M. The half reduction reaction is:Mz+ + ze MFollowing convention the half reaction that occurs on the electrode

is written as a reduction reaction

Nernst equation for these type electrodes is:

] (aq)Mn [ log n

0.059 E 0 zE

Mn

0 a log n

0.059 E zE

] (aq)[Mn

] [Mn(s) log

n

0.059 E

z0

E

z

EMn

Mn(s)0

a

a log

n

0.059 E

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 (n=1) 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 )

Ha

2. In the metal-insoluble salt-anion electrode, а metal is in contact with one of its insoluble salts and also with а solution containing the anion of the salt.

An example is the so-called silver - silver chloride electrode, written as а cathode as:

Cl-(aq) AgCl(s) Ag(s)

for which the cathode half-reaction is:

AgCl (s) + е- Ag(s) + Cl- (aq)EAg,AgCl | Cl- = E0

Ag+/AgCl - 0.059 lg aCl-

E0Ag+/AgCl = 0,2224

A

0 a log n

0.059 E zE

Silver - silver chloride electrode

Ag, is covered by the layer of nonsoluble AgCl

КCl solutionKCl, AgCl | Ag

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, HgThe 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.

3. An inert oxidation-reduction electrode

It consists of а strip, wire, or rod of an inert materiel (Pl, Au, Ir…) in contact with а solution, which contains ions of а 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:

1.Simple: Fe2+, Fe3+ | Pt Fe3+ + e → Fe2+

Ox

Red0

a

a log

n

0.059 E E

3

2

Fe

Fe0

a

a log

n

0.059 E E

SScheme of ox-red electrode cheme of ox-red electrode ((the third reference electrode)the third reference electrode)

Pt

FeCl3 + FeCl2

Fe3+ , Fe2+ | Pt

Fe3+ + e = Fe2+

2

3

2323 lg059.00

//Fe

FeFeFeFeFe a

aEE

2. Complex ox-red electrode there is changing the charge and the composition of the ions

Mn2+, MnO4-, H+ | Pt MnO4

- + 8H++5e→Mn2++4H2O

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 quinonhydrone’s crystalls in this solution.

Quinonhydrone is a crystalline product which consists of quinone (benzoquinone) С6Н4О2 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.

Quinonhydrone electrode ((the third reference electrodethe third reference electrode))

С6Н4О2, С6Н4(ОН)2, H+ | Pt

С6Н4О2 + 2Н+ + 2е = С6Н4(ОН)2

НOHC

OHHCcell аа

aEE

2

)(0

246

246lg2

059.0

Hcell aEE lg059.00

Including that the activity of Including that the activity of quinone and hydroquinone is equal quinone and hydroquinone is equal in the seturated solution, we havein the seturated solution, we have

Quinonhydrone

The scheme of quinonhydrone cell The scheme of quinonhydrone cell with one electrolytewith one electrolyte

Pt, Н2 | quinhydr, H+ | KCl | KCl,Hg2Cl2| Hg

Ecell = E quinhydr - Ecalomel

The glass electrode for pH measurements. The cell consists of a glass indicator electrode and а saturated calomel reference electrode, both immersed in the solution whose pH is to be determined. The indicator electrode consists of а thin, рН-sensitive glass membrane sealed onto one end of а heavy-walled glass or plastic tube. А small volume of dilute hydrochloric acid saturated with silver chloride is contained in the tube (in some electrodes this solution is а buffer containing chloride ion). A silver wire in this solution forms а silver/silver chloride reference electrode, which is connected to one of the terminals of а 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.

The make the cell with glass and calomel electrodes and measuring its e.m.f can be determined pH of solution.

glass electrode’s 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.

Naglass aconstE lg059.0

The mechanism of the diffusion potential

HCl 1 н HCl 0.1 н

H+

Cl-

+++++++++

---------

2

1lg059.0

a

a

nuu

uuE

AnKt

AnKtcell

where a1 > a2

The scheme of concentrated cellThe scheme of concentrated cell

(+)Ag | AgNO3 || AgNO3 | Ag(-)

C1 >

C2

_ +

AgNO3, C2 AgNO3, C1

NO3-

Ecell = E2 Ag+/Ag – E1 Ag+/Ag =

= 0.059 lg (a2 / a1)

The electric circuit with transfer and without it

The electric circuit without transfer is when the electrodes are dipped in one solution.

Examples: Pt|H2|HCl|AgCl, Ag – hydrogen-silver-silver 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: Zn|ZnSO4||CuSO4|Cu – Daniell cell

Pt|H2|HCl||KCl|Hg2Cl2,Hg – hydrogen-calomel cell