1. Electrochemistry: Electrochemistry is the study of production of electricity from the energy released during a spontaneous chemical reaction and the use of electrical energy to bring about non-spontaneous chemical transformations.
2. Ohm’s Law: It states that the potential difference (V) across the conductor is directly proportional to the current (I) flowing through it. Mathematically,
V I or V = IR
where R is a constant called resistance of the conductor. Ohm’s law is obeyed by both the metallic, as well as electrolytic conductors.
3. Resistance (R): It is the property of a substance by which it obstructs the flow of electric current through it. The electrical resistance (R) of any object is directly proportional to its length (l) and inversely proportional to its area of cross-section (A).
Thus,
where r (rho) is a constant of proportionality called specific resistance or resistivity.
If l =1 cm and A =1 cm2 then R = r.
Thus, resistivity may be defined as the resistance offered by a conductor of 1 cm length with area of cross-section equal to 1 cm2, i.e., it is the resistance of 1 cm3 of the conductor.
Units:
Its SI unit is ohm metre (W m).
4. Conductance (G): It is the reciprocal of resistance and may be defined as the ease with which the electric current flows through a conductor.
G =
Its SI unit is Siemen (S).
1 S = ohm–1 (mho)
5. Conductivity (k): It is the reciprocal of resistivity (r).
If l = 1 cm and A = 1 cm2, then k = G.
Hence, conductivity of an electrolytic solution may be defined as the conductance of a solution of 1 cm length with area of cross-section equal to 1 cm2.
Alternatively, it may be defined as the conductance of 1 cm3 of the solution of an electrolyte.
Units: k
The SI unit of conductivity is S m–1.
6. Factors Affecting Metallic Conductance
Electrical conductance through metal is called metallic or electronic conductance and is due to the movement of electrons. It depends on:
(a) The nature and structure of metal.
(b) The number of valence electrons per atom.
(c) Temperature (it decreases with increase in temperature).
7. Factors Affecting Electrolytic Conductance
Electrolyte: An electrolyte is a substance that dissociates in solution to produce ions and hence conduct electricity in dissolved or molten state.
Examples: HCl, NaOH, KCl (Strong electrolytes).
CH3COOH, NH4OH (Weak electrolytes).
The conductance of electricity by ions present in the solution is called electrolytic or ionic conductance. The following factors govern the flow of electricity through a solution of electrolyte.
(a) Nature of electrolyte or interionic attractions: Lesser the solute–solute interactions, greater will be the freedom of movement of ions and higher will be the conductance.
(b) Solvation of Ions: Larger the magnitude of solute–solvent interactions, greater is the extent of solvation and lower will be the electrical conductance.
(c) The nature of solvent and its viscosity: Larger the solvent–solvent interactions, larger will be the viscosity and more will be the resistance offered by the solvent to flow of ions and hence lesser will be the electrical conductance.
(d) Temperature: As the temperature of electrolytic solution rises solute–solute, solute–solvent and solvent–solvent interactions decrease, which results in the increase of electrolytic conductance.
8. Difference between Metallic and Electrolytic Conductance
S.No. | Metallic Conductance | Electrolytic Conductance |
(i) | Movement of electrons is responsible for conduction. | Movement of ions is responsible for conduction. |
(ii) | Does not involve transfer of matter. | Matter moves in the form of ions. |
(iii) | Decreases with increase in temperature as kernels start vibrating which produce hindrance in the flow of electrons. | Increases with increase in temperature due to decrease in interionic attraction or increase in dissociation. |
9. Measurement of Conductance: As we know, k
The value of k could be known, if we measure l, A and R. The value of the resistance of the solution between two parallel electrodes is determined by using ‘Wheatstone’ bridge method (Fig. 3.1).
It consists of two fixed resistance R3 and R4, a variable resistance R1 and the conductivity cell having the unknown resistance R2. The bridge is balanced when no current passes through the detector. Under these conditions,
Knowing the values of R1, R3 and R4 the resistance of the solution, R2 is determined. The reciprocal of R2 gives the conductance of the solution.
10. Molar Conductivity (Lm): It may be defined as the conducting power of all the ions produced by dissolving one gram mole of an electrolyte placed between two large electrodes at one centimetre apart.
Mathematically,
where, V is the volume of solution in cm3 containing 1 gram mole of electrolyte and c is the molar concentration.
Units:
= ohm–1 cm2 mol–1 or S cm2 mol–1
11. Variation of Conductivity and Molar Conductivity with Concentration:
Conductivity decreases with the decrease in concentration, this is because the number of ions per unit volume that carry the current in the solution decreases on dilution.
Molar conductivity (Lm = k × V) increases with the decrease in concentration or increase in dilution. This is because the total volume V of solution containing one mole of electrolyte increases with increase in dilution. It has been found that the decrease in k on dilution of a solution is more than compensated by increase in its volume.
Graphical representation of the variation of
Limiting Molar Conductivity (Lom ): The limiting value of molar conductivity when the concentration approaches zero is known as limiting molar conductivity or molar conductivity at infinite dilution.
It is possible to determine the molar conductivity at infinite dilution (Lom ) in case of strong electrolyte by extrapolation of curve of Lm vs
The mathematical relationship between Lm and
where
12. Kohlrausch’s Law: It states that the limiting molar conductivity of an electrolyte can be represented as the sum of the individual contributions of the anion and cation of the electrolyte.
In general, if an electrolyte on dissociation gives n+ cations and n– anions then its limiting molar conductivity is given by
Here,
Applications of Kohlrausch’s Law:
(a) Calculation of molar conductivities of weak electrolyte at infinite dilution: For example, molar conductivity of acetic acid (weak acid) at infinite dilution can be obtained from the knowledge of molar conductivities at infinite dilution of strong electrolytes like HCl, CH3COONa and NaCl as illustrated below:
i.e.,
(b) Determination of degree of dissociation of weak electrolytes:
Degree of dissociation
(c) Determination of dissociation constant (K) of weak electrolytes:
Also,
∴ K =
(d) Determination of solubility of sparingly soluble salts:
13. Electrochemical Cells: An electrochemical cell is a device in which chemical energy of the redox reaction is converted into electrical energy. The redox reaction is carried out in an indirect manner and the decrease in free energy during chemical reaction appears as electrical energy.
The simplest electrochemical cell is Daniel cell or Galvanic cell in which a zinc rod is placed in a solution of Zn2+ ions (say, ZnSO4) in the left container and a bar of copper metal is immersed in a solution of Cu2+ ions (say, CuSO4) in the right container. The two metals which act as electrodes are connected by a metallic wire through a voltmeter. The two solutions are joined by an inverted U-tube containing semi-solid paste of either KCl, KNO3 or NH4Cl in gelatin or agar-agar jelly. This arrangement of U-tube is called salt bridge (Fig. 3.3).
The overall cell reaction,
Zn(s) + Cu2+ (aq) → Zn2+ (aq) + Cu(s)
can be split into two half cells. The deflection in the voltmeter indicates the flow of current through the external circuit. The conventional current flows through the outer circuit from copper metal to zinc metal, which implies flow of electrons from zinc to copper bar.
(a) At zinc electrode, the metal undergoes oxidation and releases two electrons.
Zn(s) → Zn2+ (aq) + 2e– (oxidation)
Because oxidation is taking place, the electrode behaves as anode. These electrons travel through wire and reach the copper metal.
(b) Cu2+ (aq) + 2e– → Cu(s) (reduction)
The above reaction occurs at the copper electrode. Electronation takes place which is a reduction process and that is why it acts as cathode.
As a result of the two half cell reactions, zinc metal dissolves in anode solution to form Zn2+ ions, while the Cu2+ ions are discharged at the cathode by accepting two electrons and are deposited at cathode. The electrical neutrality is maintained in two half cells using a salt bridge. The anions of the inert electrolyte in the salt bridge migrate to the anodic chamber and cations to the cathodic chamber.
As a result, as the reaction progresses, copper bar gains weight whereas zinc rod loses weight. As a consequence, the cell continues to function till either zinc metal or copper ions in solution are consumed fully.
Since electrons are released at anode, it acquires negative polarity and cathode becomes positive because it needs electrons for the reduction of +ve ions. This observation is against the usual electrolytic cell where anode is +ve and cathode is –ve.
Salt Bridge and Its Functions
A commonly used form of salt bridge consists of a glass U-tube containing semi-solid paste of either KCl, KNO3 or NH4Cl in gelatin or agar-agar jelly.
The electrolytes that are often used in salt bridge are called inert electrolytes which are supposed:
(a) not to interact chemically with either of the solutions present in anodic or cathodic chamber.
(b) not to interfere with overall cell reaction.
(c) only those electrolytes can be used in a salt bridge in which mobility of ions is almost the same.
Example, KCl, K2SO4 , etc.
A salt bridge carries out two important functions:
(a) It allows only flow of ions through it. Thus, the circuit is completed.
(b) It also maintains the electrical neutrality.
14. Cell Diagram or Representation of an Electrochemical Cell: The following conventions or notations are applied for writing the cell diagram in accordance with IUPAC recommendations: The Daniel cell is represented as follows:
Zn(s) | Zn2+ (c1) || Cu2+ (c2) | Cu (s)
(a) Anode half cell is written on the left hand side while cathode half cell on right hand side.
(b) A single vertical line separates the metal from aqueous solution of its own ions.
Zn(s) | Zn2+ (aq); | Cu2+ (aq) | Cu(s) |
Anodic chamber | Cathodic chamber |
(c) A double vertical line represents salt bridge which allows the passage of ions through it but prevents the mixing of two solutions.
(d) The molar concentration (c) is placed in brackets after the formula of the corresponding ion.
(e) The value of EMF of the cell is written on the extreme right of the cell. For example,
Zn(s) | Zn2 + (1M) || Cu2 + (1M) | Cu (s) EMF = + 1.1 V
(f) If an inert electrode like platinum is involved in the construction of the cell, it may be written along with the working electrode in bracket, say for example, when a zinc anode is connected to a hydrogen electrode.
Zn(s) | Zn2+ (c1) || H+ (c2) | H2 (Pt)
15. Reversibility of Daniel Cell:
(a) When external voltage is less than 1.10 V, electrons flow from Zn to Cu but current flows from Cu to Zn, i.e., in opposite direction. Zinc dissolves at anode and copper deposits at cathode [see Fig. 3.4(a)]
(b) When external voltage applied is less than 1.10 V and is increased slowly, it is observed that the reaction continues to take place till the external voltage attains the value 1.10 V. When this is so, reaction stops altogether and no current flows [see Fig. 3.4(b)].
(c) If the value of external voltage exceeds the voltage of Daniel cell (1.10 V), the reaction takes place in opposite direction, i.e., the cell functions like an electrolytic cell [see Fig. 3.4(c)].
16. Electrode Potential: It may be defined as the tendency of a metal, when it is placed in contact with its own ions to either lose or gain electrons and in turn become positively or negatively charged.
The electrode potential will be named as oxidation or reduction potential depending upon whether oxidation or reduction has taken place.
OR
Characteristics:
(a) Both oxidation and reduction potentials are equal in magnitude but opposite in sign.
(b) The reduction potential shows an increase with increasing concentration and decrease with decreasing concentration of ions in a solution.
(c) It is not a thermodynamic property, so values of E are not additive.
17. Standard Hydrogen Electrode (SHE): It is a reference electrode and its reduction potential is arbitrarily assigned as zero volt at all temperature.
The standard hydrogen electrode consists of a platinum electrode coated with platinum black. The electrode is dipped in an acidic solution having 1M concentration of H+ ions. Pure hydrogen gas at 1 bar pressure is continuously bubbled through the solution at a temperature of 298 K (Fig. 3.5).
The hydrogen electrode can act both ways—as an anode or as a cathode.
Acting as anode — oxidation takes place,
H2 (g) → 2H+ (aq) + 2e–
Acting as cathode — reduction takes place,
2H+ (aq) + 2e– → H2(g)
Representation of SHE
Pt(s) | H2 (g) | H+ (aq) (c = 1 M)
18. Standard Electrode Potential (Eo): It may be defined as the electrode potential of an electrode determined relative to standard hydrogen electrode under standard conditions. The standard conditions taken are:
(a) 1 M concentration of each ion in the solution.
(b) A temperature of 298 K.
(c) 1 bar pressure for each gas.
19. Cell Potential or EMF of a Cell: The difference between the electrode potentials of two half cells is called cell potential. It is known as electromotive force (EMF) of the cell if no current is drawn from the cell.
Ecell = Ereduction – Eoxidation = Ecathode – Eanode
Since anode is put on left and cathode on right, therefore it follows
Ecell = ER – EL
For a Daniel cell,
= 0.34V – (–0.76V)
1.10V
20. Nernst Equation: It relates electrode potential with the concentration of ions.
For an electrode reaction, Mn+ (aq) + ne– → M(s)
Nernst equation can be written as
where,
R = 8.314 JK–1 mol–1
T = Temperature in kelvin
n = No. of electrons gained
F = Faraday constant (96500 C mol–1)
Substituting the value of R and F, we get
OR
Thus, the reduction potential increases with the increase in the concentration of ions.
For a general electrochemical reaction of the type:
Nernst equation can be given as
Substituting the values of R and F we get
21. Equilibrium Constant from Nernst Equation: For a Daniel cell, at equilibrium
or
But at equilibrium,
∴
In general,
22. EMF and Gibbs Free Energy: The work done by a reversible galvanic cell is equal to decrease in its free energy.
Mathematically, DrG = –nFEcell
If concentration of all the reacting species is unity, then,
From
23. Concentration Cells: If two electrodes of the same metal are dipped separately into two solutions of the same electrolyte having different concentrations and the solutions are connected through salt bridge, such cells are known as concentration cells. In these cells, oxidation takes place on the electrode with lower concentration (c1) while reduction takes place on the electrode with higher concentration (c2). For example,
H2 | H+(c1) || H+(c2) | H2; Cu | Cu2+(c1) || Cu2+(c2) | Cu
Zn | Zn2+ (c1) || Zn2+ (c2) | Zn
The EMF of concentration cell at 298 K is given by
24. Electrochemical Series: The arrangement of various standard half-cells in the order of their decreasing standard reduction potential values is known as electrochemical series.
Table 3.1: Standard Electrode Potentials at 298K
Note: Ions are present as aqueous species and H2O as liquid; gases and solids are shown by g and s.
l A negative value of Eo means that the redox couple is stronger reducing agent than H+/H2. e.g., Mg (– 2.36).
l A positive value of Eo means that the redox couple is weaker reducing agent than H+/H2. e.g., Br2 (1.09).
25. Electrolysis: The process of decomposition of an electrolyte when electric current is passed through its aqueous solution or fused state is called electrolysis.
The process of electrolysis of a substance is governed by Faraday’s laws of electrolysis.
(a) Faraday’s first law of electrolysis
“The amount of any substance deposited or liberated at the electrode is directly proportional to the quantity of electricity passing through the electrolyte.
If w grams of the substance deposited on passing Q coulombs of electricity, then
w Q or w I × t [ Q = I × t]
or w = Z × I × t
where, Z is a constant of proportionality known as electrochemical equivalent of the substance deposited.
l Electrochemical equivalent (Z): If I = 1 ampere and t = 1 second, then
w = Z
Thus, the electrochemical equivalence may be defined as the amount of the substance deposited by passing one ampere of current for one second or by passing one coulomb of charge through the electrolyte.
l 1 Faraday = Quantity of electricity carried by 1 mole of electrons.
(6.023 × 1023 mol–1 × 1.6 × 10–19 C = 96472 C mol–1 ≅ 96500 C mol–1)
l If n mol of electrons are involved in an electrode reaction, then
n × 96500 C of charge will deposit = M g of the element
1 C of charge will deposit
But 1 C of charge deposit mass of element = Z g
where E is the equivalent mass of the element and is equal to
or E = 96500 × Z
l Equivalent mass: The mass of an element deposited by passing 96500 C of charge.
(b) Faraday’s second law of electrolysis
“When same quantity of electricity is passed through different electrolytes, the amount of different substances deposited at the electrodes is directly proportional to their equivalent masses.
Mathematically,
26. Commercial Cells (Batteries): Batteries are the electrochemical cells used commercially to generate electricity. “Any battery consists of two or more than two galvanic cells connected in series where the chemical energy of the redox reactions is converted into electrical energy. There are mainly two types of batteries:
(a) Primary cells (Batteries): These cells are not chargeable because the electrode reaction occurs only once and after the use over a period of time the cells become dead and cannot be reused.
The most familiar example of this type of cell is the dry cell (known as Leclanche cell after its discoverer) which is used commonly in watches, radios, calculators, etc. It consists of a zinc container that also acts as anode and the cathode is a carbon (graphite) rod surrounded by powdered manganese dioxide and carbon. The space between the electrodes is filled by a moist paste of NH4Cl and ZnCl2 (Fig. 3.6).
The electrode reactions are:
Anode: Zn(s) → Zn2+ + 2e–
Cathode:
The cell has a potential of nearly 1.5 V.
Another type of primary cell is the mercury cell, consisting of zinc–mercury amalgam as anode and a paste of HgO and carbon as the cathode. The electrolyte is a paste of KOH and ZnO. The electrode reactions are:
Anode: Zn(Hg) + 2OH– → ZnO(s) + H2O + 2e–
Cathode: HgO(s) + H2O + 2e– → Hg (l) + 2OH–
The cell potential is approximately 1.35 V and remains constant as the ionic concentration of the solution is not changed during its life.
(b) Secondary cells (Batteries): A secondary battery is rechargeable and can be used again and again. It is recharged by passing current through it from an external source. Most familiar example of secondary cell is the lead storage battery commonly used in automobiles and invertors. It consists of a lead anode and a grid of lead packed with lead dioxide (PbO2) as cathode.
A 38% solution of H2SO4 is used as an electrolyte (Fig. 3.7).
The cell reactions when the battery is in use, are:
At anode: Pb (s) + SO42– (aq) → PbSO4(s) + 2e–
At cathode: PbO2 (s) + SO42– (aq) + 4H+ (aq) + 2e– → PbSO4(s) + 2H2O (l)
The overall reaction is:
Pb(s) + PbO2(s) + 2H2SO4 (aq) → 2PbSO4 (s) + 2H2O(l)
On recharging the cell, operated like an electrolytic cell; the reaction is reversed and PbSO4(s) on anode and cathode is converted into Pb and PbO2, respectively.
At anode: PbSO4(s) + 2H2O(l) → PbO2(s) + SO42– (aq) + 4H+ + 2e–
At cathode: PbSO4 (s) + 2e– → Pb(s) + SO42– (aq)
Another important secondary cell is the nickel–cadmium cell which has longer life than the lead storage cell but is costly. Here, the overall reaction during discharge of the battery is
Cd(s) + 2Ni(OH)3 (s) → CdO(s) + 2Ni(OH)2 (s) + H2O(l)
27. Fuel Cells: Fuel cells are those cells which produce electrical energy directly from the combustion of fuels such as hydrogen, carbon monoxide or methane. The most successful fuel cell, H2–O2 cell utilises the reaction between hydrogen and oxygen to produce water. Hydrogen and oxygen are bubbled through a porous carbon electrode in the cell into concentrated aqueous sodium hydroxide. Catalysts are incorporated into the electrode (Fig. 3.8). The electrode reactions are
Anodic reaction: 2H2 (g ) + 4OH– (aq) → 4H2O(l) + 4e–
Cathodic reaction: O2 (g) + 2H2O(l) + 4e– → 4OH– (aq)
Overall reaction 2H2 (g) + O2 (g) → 2H2O(l)
Advantages of Fuel Cells:
(a) It is a pollution-free device since no harmful products are formed.
(b) Its efficiency is about 75% which is considerably higher than conventional cells.
(c) These cells are light in weight as compared to electrical generators to produce corresponding quantity of power.
(d) It is a continuous source of energy if the supply of gases is maintained.
28. Corrosion: The process of slow eating up of metals by gases and water vapours present in atmosphere due to the formation of certain compounds like oxides, sulphides, carbonates, etc. is called corrosion. Corrosion of iron is known as rusting. Chemically, rust is hydrated ferric oxide, Fe2O3 . xH2O. Corrosion may be considered as an electrochemical phenomenon. According to electrochemical theory of rusting, the impure iron surface behaves like a small electrochemical cell in the presence of moisture containing oxygen or carbon dioxide. Such a cell is called corrosion cell or corrosion couple. In these miniature corrosion cells, pure iron acts as anode, impure surface area acts as cathode and moisture having dissolved carbon dioxide or oxygen acts as electrolyte.
At anode, oxidation of iron takes place. Thus, Fe enters into the solution as Fe2+ ions leaving behind electrons which are pushed into cathodic area.
Fe → Fe2+ + 2e–;
At cathode, the electrons are picked up by the H+ ions which are produced from H2CO3 (formed due to dissolution of CO2 in moisture) or from H2O.
H2CO3
H+ ions, thus formed, reduces the dissolved oxygen as the net reaction at the cathodic area is
The overall reaction of the corrosion cell can be obtained by adding equations (i) and (ii)
The ferrous ions so formed move through water and come at the surface where these are further oxidised by atmospheric oxygen to ferric ions and form rust which is hydrated ferric oxide (Fig. 3.9).
29. Prevention of Corrosion: The metal surface can be protected against corrosion by the following methods:
(a) Barrier protection: A thin film is introduced between iron and atmospheric oxygen, carbon dioxide and moisture. The following methods are adopted for depositing thin film on metal surface:
(i) By covering the surface with paint or a thin film of grease.
(ii) By electroplating iron with some non-corrosive metals such as nickel, chromium, copper, etc.
(b) Sacrificial protection: In this method, iron surface is covered with a more electropositive metal than iron which gets oxidised in preference to iron. In such a situation the more electropositive metal loses electrons instead of iron and thus this metal is sacrificed at the cost of iron, hence the name sacrificial protection. Iron is generally coated with zinc and this process is called galvanization.
(c) Electrical protection: This is also a case of sacrificial protection. This method is used for the protection of underground water pipes or iron tanks. In this method, the exposed surface of iron is protected by connecting it to a block of some active metal such as magnesium, aluminium or zinc (Fig. 3.10). This more electropositive metal acts as anode and lose electrons in preference to iron. The iron surface acts as cathode. This method, therefore, is also called cathodic protection. The electrons released at the anode are accepted by H+ ions of water at the surface of iron. More electropositive metal is consumed gradually in the process and needs periodical replacement.
30. Products of Electrolysis: Under the influence of electric current through molten electrolytes or their aqueous solutions, ions move towards oppositely charged electrodes. Many times the electrode products differ. For example, the electrolysis of molten sodium chloride yields sodium metal at the cathode and chlorine gas is liberated at the anode.
At anode, Cl– → Cl + e– Oxidation
Cl + Cl → Cl2 (g)
At cathode, Na+ + e– → Na (s) Reduction
However, when a concentrated aqueous solution of sodium chloride is electrolysed, H2 gas at cathode and Cl2 gas at anode are obtained. This is because water is preferably reduced at cathode.
At cathode, 2H2O(l) + 2e– → H2(g) + 2OH– (aq)
This happens because the standard reduction potential of water is greater than the standard reduction potential of Na+ ion.
Na+ (aq) + e– → Na(s) Eocell = – 2.71 V
2H2O(l) + 2e– → H2 (g) + 2OH – (aq) Eocell = – 0.83 V
At anode, however, Cl2 gas is liberated because of over potential of oxygen. In fact, the remaining solution after electrolysis yields solid NaOH on evaporating. Thus,
(a) Electrolysis of dilute H2SO4
During electrolysis of dilute H2SO4, the products are H2(g) at cathode and O2(g) at anode:
If H2SO4 is concentrated then the following reaction occurs at anode
(b) Electrolysis of aqueous copper sulphate using inert electrodes (Pt)
In this, copper is deposited at cathode and oxygen is liberated at anode.
At anode
Water, having low Eo, would be preferably oxidised at anode instead of
Cu2+ ions have greater reduction potential, copper metal is deposited at the cathode.
At cathode,
(c) If CuSO4 is electrolysed between two copper electrodes (active electrodes), the Cu2+ ions discharge at the cathode (negatively charged) and the following reaction occurs
Cu2+ (aq) + 2e– → Cu (s)
Thus, copper metal is deposited at cathode. At the anode, copper is converted into Cu2+ ions with the following change:
Cu (s) → Cu2+ (aq) + 2e–
Thus, copper is dissolved (oxidised) at anode and deposited (reduced) at cathode. This forms the basis of an industrial process in which impure copper is converted into copper of high purity. The impure copper is made as anode that dissolves on passing current and pure copper is deposited at cathode.
(d) Electrolysis of aqueous sodium bromide: Like aqueous NaCl, the electrode products are Br2 (l) at anode and H2 (g) at cathode.
Conclusions: 1. Cathodic reaction will be one which has higher
2. Anodic reaction will be one which has higher
Important Formulae
1.
where, R = Resistance
A = Area of cross-section of the electrodes
r = Resistivity
2. k =
where, k = Conductivity or specific conductance
3.
where,
M = Molarity of the solution.
4.
where,
5.
where, a = Degree of dissociation
6. For a weak binary electrolyte AB
where, K = Dissociation constant
7. Nernst equations for a general electrode reaction:
Mn+ + ne– → M
8. Nernst equation for a general electrochemical reaction:
9. log
where, Kc = Equilibrium constant
10.
where, DrG° = Standard Gibbs energy of the reaction
11. Q = I × t
where Q = Quantity of charge in coulombs
I = Current in amperes
t = Time in seconds
12. m = Z × I × t
where m = Mass of the substance deposited at the electrodes
Z = Electrochemical equivalent
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