Electrolysis is defined as a process of decomposing ionic compounds into their elements by passing a direct electric current through the compound in a fluid form. The cations are reduced at cathode and anions are oxidized at the anode.

Describe the components of electrolysis

The main components that are required for conducting electrolysis are an electrolyte, electrodes, and some form of external power source is also needed. Additionally, a partition such as an ion-exchange membrane or a salt bridge is also used but this is optional. These are used mainly to keep the products from diffusing near the opposite electrode.

An acidified or salt-containing water can be decomposed by passing electric current to their original elements hydrogen and oxygen. Molten sodium chloride can be decomposed to sodium and chlorine atoms.

Electrolysis is usually done in a vessel named ‘electrolytic cell’ containing two electrodes (cathode and anode) connected to a direct current source and an electrolyte which is an ionic compound undergoing decomposition, in either molten form or in a dissolves state in a suitable solvent. Generally, electrodes that are made from metal, graphite and semiconductor materials are used. However, the choice of a suitable electrode is done based on chemical reactivity between the electrode and electrolyte as well as the manufacturing cost.

Describe the Electrolytic Process

In the process of electrolysis, there is an interchange of ions and atoms due to the addition or removal of electrons from the external circuit. Basically, on passing current, cations move to the cathode, take electrons from the cathode (given by the supply source-battery), and is discharged into the neutral atom. The neutral atom, if solid, is deposited on the cathode and if gas, move upwards. This is a reduction process and the cation is, reduced at the cathode.

At the same time anions, give up their extra electrons to the anode and is oxidized to neutral atoms at the anode. Electrons released by the anions travel across the electrical circuit and reach the cathode completing the circuit. Electrolysis involves a simultaneous oxidation reaction at anode and a reduction reaction at the cathode.

For example, when electric current, is, passed through molten sodium chloride, the sodium ion is attracted by the cathode, from which, it takes an electrode and becomes a sodium atom.

Chloride ion reaches the anode, gives its electron, and become chlorine atom to form chlorine molecule.

Na⁺(in electrolyte) + e^–(from cathode) → Na …. At Cathode

Cl^–(from electrolyte) → e^– + Cl → Cl₂…. At Anode

Electrolysis process, while useful to get elemental forms from compounds directly, it can also be used indirectly in the metallurgy of alkali and alkaline earth metals, purification of metals, deposition of metals, preparation of compounds etc.

Explain Cell Potential or Voltage

The minimum potential needed for the electrolysis process depends on their ability of the individual ions to absorb or release electrons. It is also sometimes described as decomposition potential or decomposition voltage which is the minimum voltage (difference in electrode potential) between anode and cathode of an electrolytic cell that enables electrolysis to occur.

The voltage at which electrolysis is thermodynamically preferred is the difference of the electrode potentials as calculated using the Nernst equation. Applying additional voltage, referred to as overpotential, can increase the rate of reaction and is often needed above the thermodynamic value. It is especially necessary for electrolysis reactions involving gases, such as oxygen, hydrogen or chlorine.

This ability is, measured as an electrode potential of the ions present in the electrolytic cell. The cell potential is the sum of the potential required for the reduction and oxidation reaction. The potential involved in various redox reactions is available in literature as standard reduction potential.

Reaction with positive redox cell potentials only will be feasible as per thermodynamic Gibbs free energy (or standard potential). Generally, the electrolysis is thermodynamically controlled.

In electrolysis, a potential equal to or slightly more than that, is, applied externally. The ions, which are stable and not reacting, are made to undergo reaction in the presence of externally applied potential. External potential hence makes an unfavourable reaction to take place. In electrolysis, chemical bonds connecting atoms are either made or broken and so, electrolysis involves the conversion of electrical energy into chemical energy.

Draw a labeled diagram of apparatus used for the electrolysis of dilute sulphuric acid

Explain the Product of Electrolysis

Electrolysis of only two ions (cation and anion) present in a single electrolyte is direct. Electrolysis will produce products present in the compound. When more than one cation and anions are present, each ion will compete for reduction and oxidations. Reactions with more positive redox potentials will be, reduced or oxidized, in preference, to others.

So, in spite of multiple redox couples present, only one can be reduced or oxidized. Sometimes the ions that are reduced or oxidized may depend on their relative amount. In other words, the redox reaction and electrolysis may become kinetically controlled. In such cases, the product of analysis may differ on the relative concentration of the various ions present in the electrolyte.

For example, electrolysis of aqueous sodium chloride may give different products-

Hydrogen and chlorine,
Hydrogen and oxygen and
Hydrogen, oxygen and chlorine.

State factors affecting Electrolysis

The factors that may affect the electrolysis are;

i) The nature of the electrode
ii) Nature and state of the electrolyte

iii) Nature and electrode potential of ions present in the electrolyte and

iv) Overvoltage at the electrodes.

List any two industrial application of electrolysis

There are many industrial applications of electrolysis. The most common applications are as follows:

Extraction of metals
Purification of metals
Electroplating of metals

Define the term standard electrode potential

Standard electrode potential (E°) is a measure of the tendency of an electrode to undergo reduction (gain electrons) or oxidation (lose electrons) compared to a standard hydrogen electrode (SHE) under standard conditions. It is a fundamental property used to determine the relative strength of an oxidation-reduction (redox) reaction.
The standard electrode potential is defined as the potential difference between the electrode and the standard hydrogen electrode when both are in their standard states (1 M concentration, 1 atm pressure, and 25°C temperature). The standard hydrogen electrode is assigned an arbitrary standard potential of 0 volts.
The standard electrode potential is typically represented in units of volts (V) or millivolts (mV). The sign of the standard electrode potential indicates the direction of electron flow during a redox reaction. If the standard electrode potential is positive, it indicates a tendency for the electrode to undergo reduction (gain electrons), while a negative value indicates a tendency for oxidation (loss of electrons).
The standard electrode potential values provide a basis for comparing the reactivity and relative strengths of different redox couples. The more positive the standard electrode potential, the stronger the reducing agent (a species that readily donates electrons) associated with that electrode. Conversely, the more negative the standard electrode potential, the stronger the oxidizing agent (a species that readily accepts electrons) associated with that electrode.

State three uses of the standard electrode potential

Standard electrode potentials are used in various applications, including electrochemical cells, batteries, corrosion studies, and the determination of the feasibility and directionality of redox reactions. They provide valuable information about the thermodynamics and kinetics of redox processes and aid in the understanding and prediction of electrochemical behavior.

State the law of independent migration of ions

The law of independent migration of ions, also known as the independent migration principle, states that in an electrolyte solution, each ion moves independently towards its respective electrode during electrolysis. This principle was proposed by Friedrich Kohlrausch in the 19th century.

According to the law of independent migration of ions:

Each ion in an electrolyte solution carries its own charge and mass.
The movement of an ion is not influenced by the presence or movement of other ions in the solution.
The speed of migration of an ion is directly proportional to its charge and inversely proportional to its mass.
The total current flowing through the electrolyte is the sum of the currents carried by individual ions.

In other words, the law states that during electrolysis, the different ions present in the electrolyte solution move towards their respective electrodes independently, without interacting with or affecting the movement of other ions. The rate of migration of an ion depends on its charge and mass, with ions of higher charge and lower mass migrating faster.

This law has significant implications in electrolysis and electrochemical processes. It allows for the prediction and understanding of the movement of ions in solutions, which is crucial for the design and operation of various electrochemical cells, such as batteries, fuel cells, and electrolytic cells. The law of independent migration of ions forms the basis for the quantitative analysis of electrolyte solutions and plays a key role in fields such as electrochemistry, analytical chemistry, and materials science.

Define electroplating

Electroplating is a process that involves the deposition of a metal coating onto a surface using electrolysis. It is a widely used technique to improve the appearance, durability, and functionality of objects by adding a layer of metal.

In electroplating, a metal object to be plated, known as the substrate or the cathode, is immersed in a solution, known as the electrolyte or plating bath. The electrolyte contains ions of the metal to be plated, which are usually in the form of a salt. The metal ions in the electrolyte are attracted to the substrate by an electric field.

The electroplating process involves several key steps:

Preparation: The substrate is cleaned thoroughly to remove any dirt, grease, or oxide layers from its surface. This ensures good adhesion of the plated metal.
Electrolyte Solution: The electrolyte solution is prepared by dissolving a suitable salt of the metal to be plated in a solvent. The choice of electrolyte depends on the metal being plated and the desired properties of the plated surface.
Electrolysis: The substrate is connected to the negative terminal (cathode) of a power supply, while a metal electrode of the plating material is connected to the positive terminal (anode). When the power supply is turned on, an electric current passes through the electrolyte, causing metal ions from the electrolyte to be attracted to the substrate. The metal ions gain electrons at the substrate surface, leading to their reduction and the deposition of a metal layer on the substrate.

Control of Parameters: During electroplating, various parameters such as current density, temperature, plating time, and agitation are carefully controlled to achieve the desired thickness, smoothness, and quality of the plated metal layer.

Common metals used for electroplating include gold, silver, nickel, copper, chromium, and zinc. Each metal provides different properties and advantages, such as corrosion resistance, conductivity, aesthetic appeal, and improved durability

List uses of electroplating

Electroplating has various applications across industries due to its ability to enhance the appearance, durability, and functionality of objects. Some common uses of electroplating include:

Decorative Applications: Electroplating is extensively used for decorative purposes. Objects such as jewelry, watches, silverware, and decorative items are often electroplated with precious metals like gold, silver, or platinum to enhance their appearance and provide a luxurious finish.
Corrosion Protection: Electroplating is employed to provide a protective coating on metals, preventing them from corrosion. Metals like zinc, nickel, and chromium are commonly electroplated onto surfaces such as steel, iron, or aluminum, creating a barrier that shields the underlying metal from exposure to corrosive environments.
Electrical Conductivity: Electroplating is used to improve the electrical conductivity of surfaces. For example, copper is electroplated onto circuit boards to enhance their conductivity and facilitate the flow of electrical currents.
Wear Resistance: Electroplating can improve the wear resistance of objects by adding a layer of hard and durable metal. For instance, objects like tools, engine components, and machine parts can be electroplated with materials like chromium or nickel, which provide resistance against abrasion and wear.
Reflectivity and Optics: Electroplating is employed in the production of mirrors, reflective surfaces, and optical components. A thin layer of metal, such as aluminum or silver, is electroplated onto glass or other substrates to create a highly reflective surface or to manipulate light in optical devices.
Automotive Applications: Electroplating is widely used in the automotive industry for various purposes. It is employed to provide corrosion resistance and aesthetic finishes on parts such as bumpers, grills, emblems, and trim. Additionally, electroplating is used in the production of engine components, brake parts, and connectors to enhance their durability and performance.
Medical and Pharmaceutical Uses: Electroplating is utilized in medical and pharmaceutical applications. For instance, surgical instruments are electroplated to improve their corrosion resistance, cleanliness, and biocompatibility. Electroplating is also used in the production of pharmaceutical packaging materials to prevent interactions between the drug and the container.

List two applications of electroplating

Electroplating finds applications in a wide range of industries, including automotive, aerospace, electronics, jewelry, and decorative arts. It is used to enhance the appearance of objects, provide protection against corrosion, reduce friction and wear, improve conductivity, and achieve desired surface properties.

Explain faraday’s first law of electrolysis

Faradays first electrolysis Law states that the quantity of reaction taking place in terms of mass of ions formed or discharged from an electrolyte is proportional to the amount of electric current passed. Since electric current (ampere) is the number of coulombs (Q) flowing in one second,

Mass of the ions formed or reacted (m)∝ electric current α Q, or

m ∝ Q

m = CQ

Where, C is a proportionality constant, called the chemical equivalent of the element.

For a flow of, 1 Coulomb of charges for one second, m = C

The proportionality constant is equal to the mass of the substance involved in the reaction. Z is the electrochemical equivalent mass of one coulomb charge.

One coulomb of charge corresponds to a mass of one equivalent.

i) Electric current and Charge (Q)

Electric current is measured in ampere and it is the charges flowing per unit time (seconds).

I = Q

Quantity of Charges flowing (Q) = It = ampere × seconds

m = CIt

ii) Number of electrons–Charge Q of electrons – Faraday – Equivalent mass of substances

But the charges are associated with electrons. Every electron carries a charge of 1.6 × 10^-19 coulombs.

Charges carried by one equivalent/ mole number of electrons (Q)

= 6.02 × 10²³ × (1.6 × 10^-19) coulombs

= 96485 coulombs

≈ 96500C = 1 Faraday = 1F

One equivalent (or mole or Avogadro’s) of electrons ∝ 96485 coulombs ≈ 96500C = 1 Faraday = 1F

iii) Mass of substance undergoing electrolysis

Faraday Law says, m = C ×Q or m = C I t.

When one coulomb corresponds to one electrochemical equivalent mass (C) of the substance, one equivalent (or mole) of electrons flowing per second, will correspond to 96485 Equivalents mass.

This, 96485 electrochemical equivalents

= C × 96485

=Equivalent weight of substance in gram.

So,

Electrochemical equivalent of a substance = C

= Equivalent weight of substance in gram = E

96485 96485

Equivalent weight of a substance is, related to its molecular weight.

Equivalent mass of a substance = Molecular weight of the substance

Valency or charge

For every, one mole of electrons or charges (1Faraday) or 96485 ampere sec, passing through an electrolyte, one equivalent mass of the electrolyte is reacted, discharged/deposited etc.

One Avogadro’s number of electrons = Charge of 1 Faraday

= 96485 coulombs = 96485 ampere sec

= 1equivalent mass (reacted / deposited / neutralized)

Mass of the substance (gm) = Equivalent weight (gm) × Coulombs

96485

Equivalent weight (gm) × ampere, sec

96485 ampere,sec

Or, m = EQ = EIt

96485 96485

For example, on passing electric current through a copper sulphate electrolyte solution, copper ions gets discharged and deposit on the anode.

Cu²⁺ + 2e– →Cu

Each copper ion needs two electrons for its reduction reaction to copper atom. More the electrons passing through the copper sulphate, solution more will be the copper ions being, deposited on the anode. Hence there is a direct relationship between the mass of material being, reduced and the numbers of electrons flowing into the electrolyte.

Here, one copper ion needs 2 electrons; so, One mole of copper ions needs 2 moles of electrons.

m = 2 × equivalent weight of copper

When 0.1M MnO₄^2- is oxidized to MnO₄^–, the quantity of electricity required is

a) 96500C
b) 2 × 96500C
c) 9650C
d) 96.50C

MnO₄^2- → MnO₄^–+ e^–

1mole of MnO₄^2- lose 1 mole of electrons or 96500C

⸪ 0.1 mole of MnO₄^2- lose 0.1mole of electrons or 9650C

Answer is (c)

How much electricity in terms of Faraday is required to produce

(i) 20.0 g of Ca from molten CaCl₂?

(ii) 40.0 g of Al from molten Al₂O₃?

i) Ca²⁺+ 2 e^–→ Ca

One mole of Calcium ions gains two moles of electrons or 2 Faraday charge to produce one mole of calcium.

Molecular weight of Calcium is 40.

So, 20 gm of calcium shall need one Faraday of electricity.

ii) Al³⁺⁺+ 3 e^– → Al

One mole of Aluminum ions gains two moles of electrons or 3 Faraday charge to produce one mole of aluminim.

Molecular weight of Aluminum is 27.

So, 40 gm of Aluminum ion shall need 40×3

27 Faraday of electricity

= 4.44F

A solution of Ni(NO3)2 is electrolyzed between platinum electrodes using a current of 5 amperes for 20 minutes. Calculate the weight of nickel reduced on the cathode.

Mass of nickel deposited = E I t

96500

Equivalent eight of nickel = Atomic weight/2 =58.7/2

Mass = 58.7×5×20×60

2×96500 = 1.83gm

The anodic half- cell of lead-acid battery is recharged using electricity of 0.05 Faraday. The amount of

PbSO4 electrolyzed in g during the process is ; (molar mass of PbSO4 = 303g/mole)

a) 22.8
b) 15.2
c) 7.6
d) 11.4

In lead sulphate, lead is in 2⁺state.

One mole (303) of lead sulphate, shall need 2 Faraday electricity.

So, 0.05 Faraday will electrolyze 303×0.05 g of lead sulphate

= 7.6g So the answer is (c)

A cell, Ag | Ag+ || Cu2+ | Cu, initially contains 1 M Ag+ and 1 M Cu2+ ions. What will be the increase in potential of the cell after passing an electric current of 9.65amperes for one hour?

On passing electric current into the cell, the following reaction takes place.

– Cu → Cu²⁺ + 2^e^– ; Ag⁺+ e^– → Ag

i) Change in the concentration of copper and silver ions

Mass deposited= EIt

96500

Mass of copper ions additionally formed

= 63.54×9.65×60×60

2×96500

= 11.43g

= 0.18M

Concentration of copper half-cell = 1-0.18 = 0.82M

Mass of silver ions lost = 108×9.65×60×60

96500

= 38.9g

= 0.36M

Concentration of silver half-cell = 1-0.36 = 0.64M

Calculate the amount of charge transferred when a 5 A current is used for 2 minutes during electrolysis.

2 minutes = 2 × 60 = 120 s

Charge = current × time

Charge = 5 × 120 = 600 C

Sodium and chlorine are produced during the electrolysis of molten sodium chloride:

Na+ + e– → Na

2Cl– → Cl2 + 2e–

9,650 coulombs of charge pass. Calculate the amount of sodium and chlorine produced. Remember that 1 F (faraday) = 96,500 C.

Number of moles of electrons = 9,650 ÷ 96,500

= 0.1 mol

1 mol of electrons are needed to produce 1 mol of sodium – so 0.1 mol of sodium is produced.

2 mol of electrons are needed to produce 1 mol of chlorine – so 0.05 mol (0.1 ÷ 2) of chlorine is produced.

Bromine is produced during the electrolysis of molten lead(II) bromide: 2Br– → Br2 + 2e–

A +current of 13.4 A was used for 0.5 hours. Calculate the mass of bromine produced.

1 F = 96,500 C. Ar of Br2 = 160.

Remember: charge = current × time

Charge = 13.4 × 1,800

= 24,120 C

Remember: One faraday represents one mole of electrons. It is equal to 96,500 coulombs.

Number of moles of electrons = 24,120 ÷ 96,500

= 0.25 mol

Amount of bromine produced = 0.25 ÷ 2

= 0.125 mol

Mass of bromine produced = Ar × mol = 0.125 × 160

= 20 g

During electrolysis of molten sodium chloride, 0.125 mol of chlorine gas was produced. Calculate the volume of chlorine at RTP. Volume = amount of gas × molar volume

Volume of chlorine = 0.125 × 24

= 3 dm3 (or 3,000 cm3)

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