Electrochemistry : Electrolysis

Electrochemistry : Electrolysis

Electrochemistry is the branch of physical chemistry which deals with the relationship between electrical energy and chemical changes taking place in redox reactions i.e., how chemical energy produced in a redox reaction can be converted into electrical energy or how electrical energy can be used to bring about a redox reaction which is otherwise non-spontaneous. Chemical changes involving production or consumption of electricity are called electrochemical changes.

Electrolytes and Electrolysis.

(1) Definition : “The substances whose aqueous solution undergo decomposition into ions when electric current is passed through them are known as electrolytes and the whole process is known as electrolysis or electrolytic decomposition.”

Solutions of acids, bases, salts in water and fused salts etc. are the examples of electrolytes. Electrolytes may be weak or strong. Solutions of cane sugar, glycerine, alcohol etc., are examples of non-electrolytes.

(2) Electrolytic cell or Voltameter : The device in which the process of electrolysis or electrolytic decomposition is carried out is known as electrolytic cell or voltameter. Following are the important characteristics of voltameter,

(i) Voltameter consist of a vessel, two electrodes and electrolytic solution.

(ii) Voltameter convert electrical energy into chemical energy i.e., electrical energy is supplied to the electrolytic solution to bring about the redox reaction (i.e., electrolysis) which is non- spontaneous and takes place only when electrical energy is supplied.

(iii) In voltameter, both the electrodes are suspended in only one of the electrolytic solution or melt of the electrolyte in the same vessel.

(iv) The electrodes taken in voltameter may be of the same or different materials.

(v) The electrode on which oxidation takes place is called anode (or +ve pole) and the electrode on which reduction takes place is called cathode (or –ve pole)

(vi) During electrolysis in voltameter cations are discharged on cathode and anions on anode.

(vii) In voltameter, outside the electrolyte electrons flow from anode to cathode and current flow from cathode to anode.

Anode $\xrightleftharpoons[\text{Flow of current}]{\text{Flow of electrons}}$ Cathode

(viii) For voltameter, E_cell = – ve and ΔG = + ve.

(3) Mechanism of electrolysis

(i) The net chemical change that takes place in the cell is called the cell reaction, which is non-spontaneous in case of electrolytic cell.

(ii) The mechanism of electrolysis can be easily explained on the basis of ionisation theory according to which, the electrolytes are present in the form of ions in solution and the function of electricity is only to direct these ions to their respective electrodes.

(iii) During electrolysis cations move towards the cathode (–vely charged) while anions move towards the anode (+vely charged).

(iv) The anions on reaching the anode give up their electrons and converted into the neutral atoms.

At anode : A^– → A + e^– (Oxidation)

On the other hand cations on reaching the cathode take up electrons supplied by battery and converted to the neutral atoms.

At cathode : B⁺ + e^– → B (Reduction)

This overall change is known as primary change and products formed is known as primary products.

(v) The primary products may be collected as such or they undergo further change to form molecules or compounds. These are called secondary products and the change is known as secondary change.

(vi) The products of electrolysis depend upon,

(a) Nature of electrolyte,

(b) Concentration of electrolyte,

(d) Nature of electrodes used,

Electrodes which do not take part in chemical change involved in a cell are known as Inert electrode. eg. Graphite, Platinum.
Electrodes which take part in chemical change involved in a cell knows as active electrode. eg. Hg, Au, Cu, Fe, Zn and Ni electrode etc.

(e) Over voltage : For metal ions to be deposited on the cathode during electrolysis, the voltage required is almost the same as the standard electrode potential. However for liberation of gases, some extra voltage is required than the theoretical value of the standard electrode potential (E^o) This extra voltage required is called over voltage or bubble voltage.

(vii) The deposition of different ions at the electrodes takes place only for the time of electricity is passed and stops as soon as electricity is switched off.

Note :

The electrolyte as a whole remains neutral during the process of electrolysis as equal number of charges are neutralised at the electrodes.

If the cathode is pulled out from the electrolytic solution of the cell then there will be no passage of current and ions will show simply diffusion and state moving randomly.

If electrode is active at cathode, metal goes on depositing on cathode and at anode metal is dissolved.

(4) Preferential discharge theory

(i) According to this theory “If more than one type of ion is attracted towards a particular electrode, then the ion is discharged one which requires least energy or ions with lower discharge potential or which occur low in the electrochemical series”.

(ii) The potential at which the ion is discharged or deposited on the appropriate electrode is termed the discharge or deposition potential, (D.P.). The values of discharge potential are different for different ions.

(iii) The decreasing order of discharge potential or the increasing order of deposition of some of the ions is given below,

For cations : Li⁺, K⁺, Na, Ca²⁺, Mg²⁺, Al³⁺, Zn²⁺, Fe²⁺, Ni²⁺, H⁺, Cu²⁺, Hg²⁺, Ag⁺, Au³⁺

For anions : SO₄^2–, NO₃^–, OH^–, Cl^–, Br^–, I^–

(5) Electrolysis and electrode processes : The chemical reactions which take place at the surface of electrode are called electrode reactions or electrode processes. Various types of electrode reactions are described below,

(i) Electrolysis of molten sodium chloride : Molten sodium chloride contains Na⁺ and Cl^– ions.

NaCl_(l) ⇌ Na_(l)⁺ + Cl^–_(l)

On passing electricity, Na⁺ ions move towards cathode while Cl^– ions move towards anode. On reaching cathode and anode following reactions occur.

At cathode : Na⁺_(l) + e^– → Na_(l)↓ (Reduction, primary change)

At anode : Cl^–_(l) – e^– → Cl (Oxidation, primary change)

$\[\overset{\bullet }{\mathop{Cl}}\,+\overset{\bullet }{\mathop{Cl}}\,\to C{{l}_{2}}(g)\downarrow$ (Secondary change)

Overall reaction : 2Na⁺_(l) + 2Cl^–_(l)2Na_(l) + Cl_2(g)

The sodium obtained at the cathode is in the molten state.

(ii) Electrolysis of molten lead bromide : Molten lead bromide contains Pb²⁺_(l) and Br^– ions.

PbBr_2(l) ⇌ Pb²⁺_(l) + 2Br_(l)^–

On passing electricity, Pb²⁺ ions move towards cathode while Br^– ions move towards anode. On reaching cathode and anode following reactions occur,

At cathode : $\[Pb_{(l)}^{2+}+2{{e}^{-}}\to P{{b}_{(l)}}\downarrow$ (Reduction, primary change)

At anode : $\[Br_{(l)}^{-}-{{e}^{-}}\to \overset{\bullet }{\mathop{Br}}\,$ (Oxidation, primary change)

$\[\overset{\bullet }{\mathop{Br}}\,+\overset{\bullet }{\mathop{Br}}\,\to B{{r}_{2}}(g)\downarrow$ (Secondary change)

Overall reaction : $\[Pb_{(l)}^{2+}+2Br_{(l)}^{-}\xrightarrow{\text{Electrolysis}}P{{b}_{(l)}}+B{{r}_{2}}(g)$

The lead obtained at the cathode is in the molten state.

(iii) Electrolysis of water : Water is only weakly ionized so it is bad conductor of electricity but the presence of an acid (H₂SO₄) increases the degree of ionization of water. Thus, in solution, we have three ions, i.e., H⁺, OH⁻ and SO₄²⁻ produced as follows,

$\[{{H}_{2}}O\rightleftharpoons H_{(aq)}^{+}+O{{H}^{-}}$

$\[{{H}_{2}}SO(aq)\to 2{{H}^{+}}(aq)+SO_{4}^{2-}(aq)$

When electric current is passed through acidified water, ions move towards cathode, OH⁻ and SO₄²⁻ ions move towards anode. Now since the discharge potential of OH^– ions is much lower than that of SO₄^2- ions, therefore, OH^– ions are discharged at anode while SO₄^2- ions remain in solution. Similarly, H⁺ ions are discharged at the cathode. The reactions, occurring at the two electrodes may be written as follows,

At cathode : H⁺(aq) + e⁻ → H (Reduction, primary change)

$\[\overset{\bullet }{\mathop{H}}\,+\overset{\stackrel{\scriptscriptstyle\to}{\leftarrow}\stackrel{\scriptscriptstyle\to}{\leftarrow}\bullet }{\mathop{H}}\,\to {{H}_{2}}(g)$ (Secondary change)

At anode : OH⁻(aq) + e⁻ → OH (Oxidation, primary change)

4OH → 2H₂O(l) + O₂ (Secondary change)

Overall reaction : 4H⁺(aq)+4OH⁻(aq)→2H₂(aq) + 2H₂O(l) + O₂(g)

(iv) Electrolysis of aqueous copper sulphate solution using inert electrodes : Copper sulphate and water ionize as under,

CuSO₄(aq) → Cu²⁺(aq) + SO₄²⁻(aq)(almost completely ionized)

H₂O(l) ⇌ H⁺(aq) + OH⁻(aq) (only slightly ionized)

On passing electricity, Cu²⁺(aq) and H⁺(aq) move towards the cathode while SO₄²⁻ ions and OH⁻ ions move towards the anode. Now since the discharge potential of Cu²⁺ ions is lower than that of H⁺ ions, therefore, Cu²⁺ ions are discharged at cathode. Similarly, OH⁻ ions are discharged at the anode. The reactions, occurring at the two electrodes may be written as follows,

At cathode : Cu²⁺ + 2e⁻ → Cu (Reduction, primary change)

At anode : 4OH⁻ − 4e⁻ → 4OH (Oxidation, primary change)

4OH → 2H₂O(l) + O₂(g) (Secondary change)

(v) Electrolysis of an aqueous solution of copper sulphate using copper electrode : Copper sulphate and water ionize as under:

CuSO₄(aq) → Cu²⁺(aq) + SO₄²⁻(aq)(almost completely ionized)

H₂O(l) ⇌ H⁺(aq) + OH⁻ (weakly ionized)

On passing electricity, Cu²⁺ and H⁺ move towards cathode while OH⁻ and SO₄²⁻ ions move towards anode. Now since the discharge potential of Cu²⁺ ions is lower than that of H⁺ ions, therefore, Cu²⁺ ions are discharged in preference to ions at cathode. Unlike electrolysis of CuSO₄ using platinum electrodes, no ions are liberated here, instead, anode itself undergoes oxidation (i.e., loses electrons) to form Cu²⁺ ions which go into the solution. This is due to the reason that Cu is more easily oxidised than both OH⁻ and SO₄²⁻ ions. The reactions, occuring at the two electrodes may be written as follows,

At cathode : Cu²⁺(aq) + 2e⁻ → Cu(s) (Reduction, primary change).

At anode : Cu(s) → Cu²⁺(aq) + 2e⁻ (Oxidation, primary change).

Products of electrolysis of some electrolytes

Electrolyte	Electrode	Product at cathode	Product at anode
Aqueous NaOH	Pt or Graphite	H₂	O₂
Fused NaOH	Pt or Graphite	Na	O₂
Aqueous NaCl	Pt or Graphite	H₂	Cl₂
Fused NaCl	Pt or Graphite	Na	Cl₂
Aqueous CuCl₂	Pt or Graphite	Cu	Cl₂
Aqueous CuCl₂	Cu electrode	Cu	Cu oxidised to Cu²⁺ ions
Aqueous CuSO₄	Pt or Graphite	Cu	O₂
Aqueous CuSO₄	Cu electrode	Cu	Cu oxidized to Cu²⁺ ions
Dilute H₂SO₄	Pt electrode	H₂	O₂
Conc. H₂SO₄	Pt electrode	H₂	Peroxodisulphuric acid (H₂S₂O₈)
Aqueous AgNO₃	Pt electrode	Ag	O₂
Aqueous AgNO₃	Ag electrode	Ag	Ag oxidised to Ag⁺ ions
Acidified water	Pt electrode	H₂	O₂
Fused PbBr₂	Pt electrode	Pb	Br₂

Note :

Alkali metals, alkaline earth metals and other metals having E^o lower than hydrogen cannot be obtained during electrolysis of their aqueous salt solutions because of their strong electropositive nature.

The electrolysis of aqueous solution of electrolytes is some what more complex because of the ability of water to be oxidised as well as reduced.

(6) Application of electrolysis: Electrolysis has wide applications in industries. Some of the important applications are, as follows,

(i) Production of hydrogen by electrolysis of water.

(ii) Manufacture of heavy water (D₂O).

(iii) Electrometallurgy : The metals like Na, K, Mg, Al, etc., are obtained by electrolysis of fused electrolytes.

Fused electrolyte	Metal isolated
NaCl + CaCl₂ + KF CaCl₂ + CaF₂	Na Ca
Al₂O₃ + cryolite	Al
MgCl₂ + NaCl + CaCl₂	Mg
NaOH	Na
KCl + CaCl₂	K

(iv) Manufacture of non-metals : Non-metals like hydrogen, fluorine, chlorine are obtained by electrolysis.

(v) Electro-refining of metals : It involves the deposition of pure metal at cathode from a solution containing the metal ions. Ag, Cu etc. are refined by this method.

(vi) Electrosynthesis : This method is used to producing substances through non-spontaneous reactions carried out by electrolysis. Compounds like NaOH, KOH, Na₂CO₃, KClO₃ white lead, KMO₄ etc. are synthesised by this method.

(vii) Electroplating : The process of coating an inferior metal with a superior metal by electrolysis is known as electroplating. The aim of electroplating is, to prevent the inferior metal from corrosion and to make it more attractive in appearance. The object to be plated is made the cathode of an electrolytic cell that contains a solution of ions of the metal to be deposited.

For electroplating	Anode	Cathode	Electrolyte
With copper	Cu	Object	CuSO₄ + dilute H₂SO₄ K[Ag(CN)₂]
With silver	Ag	Object
With nickel	Ni	Object	Nickel ammonium sulphate K[Au(CN)₂]
With gold	Au	Object	ZnSO₄
With zinc	Zn	Iron objects
With tin	Sn	Iron objects	SnSO₄

Thickness of coated layer : Let the dimensions of metal sheet to be coated be (a cm × b cm)

Thickness of coated layer = c cm

Volume of coated layer = (a × b × c) cm³

Mass of the deposited substance = Volume × density = (a × b ×) × dg = (a × b × c) × dg

∴ $\[(a\times b\times c)\times d=\frac{I\times t\times E}{96500}$

Using above relation we may calculate the thickness of coated layer.

Note : Sometimes radius of deposited metal atom is given instead of density.

For example, radius of silver atom = 10⁻⁸ cm; Atomic mass of Ag = 108

Mass of single silver atom $\[=\frac{108}{6.023\times {{10}^{23}}}g$

Volume of single atom $\[=\frac{4}{3}\times \pi \,\,\,{{R}^{3}}\,\,\,=\,\,\frac{4}{3}\times 3.14\times {{({{10}^{-8}})}^{3}}\,\,\,c{{m}^{3}}$

Density of Ag $\[=\frac{\text{Mass}\,\,\text{of}\,\,\text{single}\,\,\text{atom}}{\text{Volume}\,\,\text{of}\,\,\text{single}\,\,\text{atom}}=\frac{108/6.023\times {{10}^{23}}}{\frac{4}{3}\times 3.14\times {{({{10}^{-8}})}^{3}}}$ = 42.82 g / cm³

Faraday’s laws of electrolysis.

The laws which govern the deposition of substances (In the form of ions) on electrodes during the process of electrolysis is called Faraday’s laws of electrolysis. These laws given by Michael Faraday in 1833.

(1) Faraday’s first law : It states that,

“The mass of any substance deposited or liberated at any electrode is directly proportional to the quantity of electricity passed.”

i.e., W ∝ Q

Where, W= Mass of ions liberated in gm,

Q = Quantity of electricity passed in Coulombs = Current in Amperes (I) × Time in second (t)

WI × t or W = Z × I × t

In case current efficiency (ή) is given, then

$\[W=Z\times I\times t\times \frac{\eta }{100}$

Where, Z = constant, known as electrochemical equivalent (ECE) of the ion deposited.

When a current of 1 Ampere is passed for 1 second (i.e., Q = 1), then, w = z

Thus, electrochemical equivalent (ECE) may be defined as “the mass of the ion deposited by passing a current of one Ampere for one second (i.e., by passing Coulomb of electricity)”. It’s unit is gram per Coulomb. The ECE values of some common elements are,

Element	Hydrogen	Oxygen	Copper	Silver	Iodine	Mercury
Z, g C⁻¹	1.045 × 10⁻⁵	8.29 × 10⁻⁵	3.294 × 10⁻⁴	1.18 × 10⁻³	1.315 × 10⁻³	1.039 × 10⁻³

Note :

Coulomb is the smallest unit of electricity.

96500 Coulombs = 6.23 × 10²³ electrons.

1 Coulomb $\[=\frac{6.023\times {{10}^{23}}}{96500}=6.28\times {{10}^{18}}$ electrons, or 1 electronic charge = 1.6 × 10⁻¹⁹ Coulomb.

(2) Faraday’s second law : It states that,

“When the same quantity of electricity is passed through different electrolytes, the masses of different ions liberated at the electrodes are directly proportional to their chemical equivalents (Equivalent weigths).” i.e.,

$<em>\[\frac{{{W}_{1}}}{{{W}_{2}}}=\frac{{{E}_{1}}}{{{E}_{2}}}\,\,\,or\,\,\,\frac{{{Z}_{1}}It}{{{Z}_{2}}It}=\frac{{{E}_{1}}}{{{E}_{2}}}\,\,\,or\,\,\,\,\frac{{{Z}_{1}}}{{{Z}_{2}}}=\frac{{{E}_{1}}}{{{E}_{2}}}$ (W = Zit)

Thus the electrochemical equivalent (Z) of an element is directly proportional to its equivalent weight (E), i.e.,

E ∝ Z or E = FZ or E = 96500 × Z

Where, Faraday constant = 96500 C mol⁻¹

So, 1 Faraday = 1F = Electrical charge carried out by one mole of electrons.

1F = Charge on an electron × Avogadro’s number.

1F = e⁻ × N = (1.602 × 10⁻¹⁹c) × (6.023 × 10²³ mol⁻¹)

Number of Faraday $\[=\frac{\text{Number}\,\,\text{of}\,\,\text{electrons}\,\,\text{passed}}{\text{6}\text{.023}\times \text{1}{{\text{0}}^{\text{23}}}}$

(3) Faraday’s law for gaseous electrolytic product : For the gases, we use

$\[V=\frac{It\,\,{{V}_{e}}}{96500}$

where, V = Volume of gas evolved at S.T.P. at an electrode

V_e = Equivalent volume = Volume of gas evolved at an electrode at S.T.P. by 1 Faraday charge

Examples :

(a) O₂ M = 32, E = 8; 32 g; 32 g O₂ ≡ 22.4L at S.T.P

[M = Molecular mass, E = Equivalent mass]

8g O₂ ≡ 5.6L at S.T.P; Thus V_e of O₂ = 5.6L.

(b) H₂: M = 2, E = 1; 2g H₂ ≡ 22.4L at S.T.P

1 g H₂ ≡ 11.2L at S.T.p

Thus V_e of H₂ = 11.2L.

(c) Cl₂ ; M =71, E = 35.5; 71g Cl₂ ≡ 22.4L at S.T.P

35.5 g Cl₂ ≡ 11.2L at S.T.P

Thus V_e of Cl₂ = 11.2L.

(4) Quantitative aspects of electrolysis : We know that, one Faraday (1F) of electricity is equal to the charge carried by one mole (6.023 × 10²³) of electrons. So, in any reaction, if one mole of electrons are involved, then that reaction would consume or produce 1F of electricity. Since 1F is equal to 96,500 Coulombs, hence 96,500 Coulombs of electricity would cause a reaction involving one mole of electrons.

If in any reaction, n moles of electrons are involved, then the total electricity (Q) involved in the reaction is given by,

Q = nF = n × 96,500 C

Thus, the amount of electricity involved in any reaction is related to,

(i) The number of moles of electrons involved in the reaction,

(ii) The amount of any substance involved in the reaction.

Therefore, 1 Faraday or 96,500 C or 1 mole of electrons will reduce,

(a) 1 mole of monovalent cation, (b) $\[\frac{1}{2}$ mole of divalent cation,

Electrochemistry : Electrolysis

Products of electrolysis of some electrolytes

Electrolyte

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