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CLASS 11-PCB . CHEMISTRY . CHEMISTRY PART II . REDOX REACTION

Chapter 7 : Redox Reaction

Ch 7

CHEMISTRY

CLASS 11-PCB

Concept of Oxidation Reduction Redox Reactions Oxidation Number

Oxidation

Oxidation is a chemical process where a substance gains oxygen or an electronegative element, or loses hydrogen or an electropositive element. This process results in an increase in the oxidation state of the substance.

Oxidation is defined as the addition of oxygen or an electronegative element to a substance, or the removal of hydrogen or an electropositive element from a substance.

In oxidation, oxygen or an electronegative element is added to a substance, or hydrogen or an electropositive element is removed from it. This causes the substance to lose electrons, increasing its oxidation number.

Examples:

CH94(g) + 2O92(g) 121212; CO92(g) + 2H92O(l) (Addition of oxygen)
Mg(s) + F92(g) 121212; MgF92(s) (Addition of electronegative element)
2H92S(g) + O92(g) 121212; 2S(s) + 2H92O(l) (Removal of hydrogen)
2K94Fe(CN)94 + H92O92 121212; 2K93Fe(CN)93 + 2KOH(aq) (Removal of electropositive element)

Reduction

Reduction is a chemical process where a substance loses oxygen or an electronegative element, or gains hydrogen or an electropositive element. This process results in a decrease in the oxidation state of the substance.

Reduction is defined as the removal of oxygen or an electronegative element from a substance, or the addition of hydrogen or an electropositive element to a substance.

In reduction, oxygen or an electronegative element is removed from a substance, or hydrogen or an electropositive element is added to it. This causes the substance to gain electrons, decreasing its oxidation number.

Examples:

2HgO(s) 121212; 2Hg(l) + O92(g) (Removal of oxygen)
2FeCl93(aq) + H92(g) 121212; 2FeCl92(aq) + 2HCl(aq) (Removal of electronegative element)
CH92=CH92(g) + H92(g) 121212; CH930CH93(g) (Addition of hydrogen)
2HgCl92(aq) + SnCl92(aq) 121212; Hg92Cl92(s) + SnCl94(aq) (Addition of electropositive element)

Redox Reaction

A redox reaction is a chemical reaction in which oxidation and reduction occur simultaneously. One substance loses electrons (oxidation) while another gains electrons (reduction), resulting in the transfer of electrons between the substances.

A redox reaction is a chemical reaction that involves both oxidation and reduction processes occurring simultaneously.

In a redox reaction, the substance that loses electrons undergoes oxidation, and the substance that gains electrons undergoes reduction. These processes happen at the same time, maintaining the balance of electrons in the reaction.

Redox Reaction in Terms of Electron Transfer

In a redox reaction, electrons are transferred from one reactant to another. The substance that loses electrons is oxidised, and the substance that gains electrons is reduced. This simultaneous process of electron loss and gain is the essence of redox reactions, enabling the transformation of chemical species through electron exchange.

In a redox reaction, oxidation and reduction occur simultaneously through the transfer of electrons from one substance to another. The substance that loses electrons undergoes oxidation, while the substance that gains electrons undergoes reduction. This electron transfer is the fundamental process in redox reactions.

During a redox reaction, the electron transfer occurs as follows:

The oxidised substance loses electrons, increasing its oxidation state.
The reduced substance gains electrons, decreasing its oxidation state.
This transfer maintains the balance of charge and mass in the reaction.
Electron transfer is the driving force behind energy changes and chemical transformations in redox processes.

In a redox reaction, one substance loses electrons (oxidation) and another gains electrons (reduction). For example, in the reaction between sodium and chlorine:

Sodium (Na) atoms lose electrons to form Na⁺ ions (oxidation).
Chlorine (Cl₂) molecules gain electrons to form Cl^- ions (reduction).
The transfer of electrons occurs simultaneously, resulting in the formation of sodium chloride (NaCl).

Oxidation and Reduction Half Reactions

Oxidation and reduction half reactions represent the two complementary parts of a redox reaction. The oxidation half reaction shows the loss of electrons, while the reduction half reaction shows the gain of electrons. These half reactions can be written separately to analyze and balance redox equations effectively.

The oxidation half reaction involves the loss of electrons by a substance, indicating oxidation. The reduction half reaction involves the gain of electrons by another substance, indicating reduction. These half reactions occur simultaneously and together represent the overall redox process.

The oxidation half reaction involves a species losing electrons, for example:

Zn(s) 121212; Zn²⁺ + 2e^-

The reduction half reaction involves a species gaining electrons, for example:

Cu²⁺ + 2e^- 121212; Cu(s)

Together, these half reactions combine to form the overall redox reaction.

The half reaction involving loss of electrons is called the oxidation half reaction, and the half reaction involving gain of electrons is called the reduction half reaction.

Oxidising Agent and Reducing Agent

In redox reactions, the oxidising agent and reducing agent play crucial roles:

Oxidising Agent: Accepts electrons from another species and is itself reduced.
Reducing Agent: Donates electrons to another species and is itself oxidised.

These agents facilitate the electron transfer that drives the redox process.

In redox reactions, the oxidising agent is the substance that causes oxidation by accepting electrons and itself gets reduced. The reducing agent causes reduction by donating electrons and itself gets oxidised. Both agents are essential for the electron transfer process.

The oxidising agent causes oxidation by accepting electrons and undergoing reduction. Conversely, the reducing agent causes reduction by donating electrons and undergoing oxidation. Both agents are essential for the redox reaction to proceed, as they enable the transfer of electrons between reactants.

Oxidising Agent: The species that oxidises another substance and itself gets reduced; it accepts electrons.

An oxidising agent is a substance that causes another substance to lose electrons (oxidation) by accepting those electrons itself. During this process, the oxidising agent is reduced. It plays a crucial role in redox reactions by facilitating the oxidation of other substances.

An oxidising agent is a chemical species that gains electrons during a redox reaction, thereby oxidising another substance. It undergoes reduction itself. For example, in the reaction between hydrogen and fluorine, fluorine acts as the oxidising agent by accepting electrons.

For example, in the reaction between hydrogen and fluorine, fluorine acts as the oxidising agent. It gains electrons from hydrogen, causing hydrogen to be oxidised to H⁺ ions, while fluorine is reduced to fluoride ions (F^-).

Reducing Agent: The species that reduces another substance and itself gets oxidised; it donates electrons.

A reducing agent is a substance that causes another substance to gain electrons (reduction) by donating electrons itself. During this process, the reducing agent is oxidised. It is essential in redox reactions as it facilitates the reduction of other substances.

A reducing agent is a chemical species that loses electrons during a redox reaction, thereby reducing another substance. It undergoes oxidation itself. For example, in the reaction between hydrogen and fluorine, hydrogen acts as the reducing agent by donating electrons.

For example, in the reaction between hydrogen and fluorine, hydrogen acts as the reducing agent. It donates electrons to fluorine, causing fluorine to be reduced to fluoride ions, while hydrogen is oxidised to H⁺ ions.

Oxidation Number

Oxidation number, also known as oxidation state, is a concept used to keep track of electron transfer in chemical reactions. It is a hypothetical charge assigned to an atom in a compound or ion based on a set of rules that consider the electronegativity of atoms and the nature of chemical bonds. This number helps identify which atoms are oxidised and which are reduced during a reaction.

Oxidation number (O.N.) is a numerical value assigned to an atom in a compound or ion that represents its degree of oxidation or reduction. It is determined based on a set of rules considering the electronegativity of atoms and the nature of chemical bonds. Oxidation numbers help in identifying which atoms are oxidised and which are reduced in a reaction.

Oxidation number is assigned according to the following rules:

Atoms in their elemental form have an oxidation number of zero.
Oxygen usually has an oxidation number of -2, except in peroxides (-1), superoxides (-1/2), and when bonded to fluorine (+2 or +1).
Hydrogen usually has an oxidation number of +1, except when bonded to metals in binary compounds where it is -1.
Fluorine always has an oxidation number of -1.
Other halogens usually have an oxidation number of -1 when in halide ions, but positive oxidation numbers when combined with oxygen.
The sum of oxidation numbers in a neutral compound is zero.
The sum of oxidation numbers in a polyatomic ion equals the charge on the ion.

These rules allow chemists to systematically determine the oxidation states of atoms in compounds and ions, which is essential for understanding and balancing redox reactions.

Oxidation number (O.N.) is the charge an atom would have if electrons in bonds were assigned to the more electronegative atom according to a set of rules.

Oxidation number (O.N.) is a useful concept to keep track of electron transfer in redox reactions. It helps in identifying which atoms are oxidised and which are reduced by assigning a numerical value to each atom based on a set of rules. These rules consider the electronegativity of atoms and the nature of bonds, allowing chemists to systematically balance redox equations and understand reaction mechanisms.

Rules for Oxidation Number:

Atoms in elemental form have O.N. = 0.
Oxygen usually has O.N. = -2, except in peroxides (-1), superoxides (-1/2), and when bonded to fluorine (+2 or +1).
Hydrogen usually has O.N. = +1, except when bonded to metals in binary compounds (-1).
Fluorine always has O.N. = -1.
Other halogens usually have O.N. = -1 when in halide ions, but positive O.N. when combined with oxygen.
The sum of oxidation numbers in a neutral compound is zero.
The sum of oxidation numbers in a polyatomic ion equals the charge on the ion.

Examples of Oxidation Numbers in Ions:

Na⁺: +1
Mg²⁺: +2
Fe³⁺: +3
Cl^-: -1
O^2-: -2

According to Oxidation Number Concept:

Oxidation is an increase in oxidation number.
Reduction is a decrease in oxidation number.
Oxidising agent increases the oxidation number of an element.
Reducing agent decreases the oxidation number of an element.
Redox reactions involve changes in oxidation numbers of species.

Key Terms

Oxidising Agent: A substance that gains electrons and gets reduced in a redox reaction, causing another substance to be oxidised.

Reducing Agent: A substance that loses electrons and gets oxidised in a redox reaction, causing another substance to be reduced.

Redox Reaction: A chemical reaction involving simultaneous oxidation and reduction processes.

Oxidation Number: The hypothetical charge assigned to an atom in a compound based on electron distribution rules.

Half Reaction: An equation showing either the oxidation or reduction process separately, including electron transfer.

Electrochemical Cell: A device that converts chemical energy into electrical energy through redox reactions occurring in separate compartments.

Salt Bridge: A pathway that allows ion flow between two half cells in an electrochemical cell, maintaining electrical neutrality.

Oxidation Number: The charge that an atom appears to have in a compound or ion.

The oxidation number is a formalism used to describe the degree of oxidation or reduction of an atom in a chemical compound. It is assigned according to specific rules, such as:

Elements in their natural state have an oxidation number of zero.
Oxygen usually has an oxidation number of -2, except in peroxides and when bonded to fluorine.
Hydrogen usually has an oxidation number of +1, except when bonded to metals where it is -1.
The sum of oxidation numbers in a neutral compound is zero, and in ions equals the ion charge.

This concept helps in understanding electron transfer and balancing redox reactions.

Key Facts

Electron releasing tendency among Zn, Ag, and Cu is Zn > Cu > Ag.
Oxidation state is often used interchangeably with oxidation number.
Stock notation expresses oxidation number using Roman numerals in parentheses after the metal symbol, e.g., Sn(II)Cl₂ for stannous chloride.

Mnemonics

Concept: Redox reaction

Mnemonic: OiL RiG

Interpretation: Oxidation Is Losing electrons, Reduction Is Gaining electrons.

Concept: Redox reaction

Concept: Redox reaction

Mnemonic: OiL RiG

Mnemonic: OiL RiG

Interpretation: Oxidation Is Losing (electrons), Reduction Is Gaining (electrons).

Interpretation: Oxidation Is Losing (electrons), Reduction Is Gaining (electrons).

This mnemonic helps remember the fundamental concept of redox reactions: Oxidation involves the loss of electrons, while reduction involves the gain of electrons. This is essential for understanding electron transfer in chemical reactions.

Solved Examples

Example 1: Identify oxidation and reduction in the reaction:

Zn(s) + CuSO₄(aq) 121212; ZnSO₄(aq) + Cu(s)

Solution:

Oxidation: Zn 121212; Zn²⁺ + 2e^- (Zn loses electrons)

Reduction: Cu²⁺ + 2e^- 121212; Cu (Cu gains electrons)

Zn is oxidised and Cu²⁺ is reduced.

Example 2: Calculate oxidation number of S in H₂SO₄.

Solution:

Let oxidation number of S = x

H = +1, O = -2

Sum of oxidation numbers = 0

2(+1) + x + 4(-2) = 0

2 + x - 8 = 0

x = +6

Oxidation number of S is +6.

Practice Set

Level 1: Define oxidation and reduction with examples.
Level 2: Explain the role of oxidising and reducing agents in a redox reaction.
Level 3: Calculate the oxidation number of chlorine in ClO₄^- ion.

Answer Key

Level 1: Oxidation is addition of oxygen or removal of hydrogen; reduction is removal of oxygen or addition of hydrogen. Examples: Oxidation - CH94 + O92 121212; CO92 + H92O; Reduction - 2HgO 121212; 2Hg + O92.
Level 2: Oxidising agent accepts electrons and gets reduced; reducing agent donates electrons and gets oxidised.
Level 3: In ClO₄^- ion, sum of oxidation numbers = -1; O = -2; let Cl = x; x + 4(-2) = -1; x - 8 = -1; x = +7.

Types of Redox Reactions Balancing Electrode Processes

Types of Redox Reactions

Redox reactions are chemical reactions that involve the transfer of electrons between two species. These reactions include various types based on how substances combine, break down, or replace each other during the process of oxidation and reduction.

Redox reactions are classified into:

Combination Reactions
Decomposition Reactions
Displacement Reactions
Disproportionation Reactions

Combination Reactions

Combination reactions are those in which two or more reactants combine to form a single product. These reactions involve the simultaneous oxidation and reduction of substances to form a new compound.

These involve two or more substances combining to form a single compound.

For example, when carbon reacts with oxygen, they combine to form carbon dioxide:

C(s) + O₂(g) → CO₂(g)

This is a typical combination redox reaction where carbon is oxidised and oxygen is reduced.

Example:

C(s) + O92(g) 121212; CO92(g)

Decomposition Reactions

Decomposition reactions involve the breakdown of a compound into two or more simpler substances, at least one of which is in the elemental form. These reactions also involve oxidation and reduction processes.

For example, when sodium hydride is heated, it decomposes into sodium metal and hydrogen gas:

2NaH(s) → 2Na(s) + H₂(g)

In this reaction, sodium is reduced and hydrogen is oxidised.

These involve the breakdown of a compound into two or more components, at least one in elemental form.

Decomposition reactions are chemical processes where a single compound breaks down into two or more simpler substances. At least one of the products formed is in the elemental form, meaning it exists as a pure element rather than a compound. These reactions often require an input of energy such as heat, light, or electricity to proceed, as breaking chemical bonds requires energy.

Decomposition reactions involve the breakdown of a compound into two or more components, where at least one of the products is in the elemental state. These reactions often require energy input such as heat, light, or electricity to proceed.

In decomposition reactions, a compound splits into multiple products, with at least one product being an element in its natural state. These reactions are typically endothermic, meaning they absorb energy from the surroundings to break the bonds within the compound. Common examples include thermal decomposition, electrolysis, and photolysis. The energy input facilitates the breaking of chemical bonds, resulting in simpler substances or elements.

Example:

2NaH(s) 121212; 2Na(s) + H92(g)

Displacement Reactions

Displacement reactions are a type of chemical reaction where an atom or ion in a compound is replaced by another atom or ion. This replacement occurs because the displacing element is more reactive than the element it replaces. Displacement reactions are common among metals and halogens, where a more reactive metal or halogen displaces a less reactive one from its compound.

Displacement reactions are chemical reactions where an atom or ion in a compound is replaced by another atom or ion. This type of reaction is common in metals and halogens, where a more reactive element displaces a less reactive element from its compound.

In displacement reactions, a more reactive element displaces a less reactive element from its compound. This process involves the transfer of electrons, where the displacing element undergoes oxidation and the displaced element undergoes reduction. These reactions are important in metallurgy and chemical synthesis, as they allow extraction and purification of elements.

These involve replacement of an atom or ion by another atom or ion from a compound.

Displacement reactions involve the replacement of an atom or ion in a compound by another atom or ion. The displacing species must be more reactive to successfully replace the other. This reaction can be represented by the general equation where element X displaces element Y from compound YZ, forming compound XZ and releasing element Y.

In displacement reactions, an element in its free state replaces another element from a compound. The general form of a displacement reaction is:
X + YZ 121212; XZ + Y
where X is the displacing element and YZ is the compound.

The general form of a displacement reaction is:
X + YZ 121212; XZ + Y
Here, X is the displacing element, and YZ is the compound containing element Y. The reaction proceeds when X is more reactive than Y, allowing X to replace Y in the compound, resulting in the formation of XZ and the release of Y in its elemental form.

General form:

X + YZ 121212; XZ + Y

Metal Displacement Reactions

Metal displacement reactions are a type of redox reaction where a more reactive metal displaces a less reactive metal from its compound. This occurs because metals have different tendencies to lose electrons, which is reflected in their position in the reactivity series. The more reactive metal loses electrons (oxidation) and replaces the less reactive metal, which gains electrons (reduction).

Metal displacement reactions occur when a more reactive metal displaces a less reactive metal from its compound. This is based on the reactivity series of metals, where metals higher in the series can replace those lower from their compounds.

In these reactions, the metal that is higher in the reactivity series can replace a metal that is lower from its compound. This is because the more reactive metal has a greater tendency to lose electrons and form positive ions, thereby displacing the less reactive metal from its compound.

A metal displaces another metal from its compound if it is more reactive.

A metal displaces another metal from its compound if it is more reactive. This means that the metal with higher reactivity will undergo oxidation by losing electrons and enter into solution as ions, while the less reactive metal will be reduced and precipitate out as a solid.

For example, zinc metal can displace copper from copper sulfate solution because zinc is more reactive than copper:
Zn(s) + CuSO₄(aq) 121212; ZnSO₄(aq) + Cu(s)

For example, zinc metal can displace copper from copper sulfate solution because zinc is more reactive than copper:
Zn(s) + CuSO₄(aq) 111; ZnSO₄(aq) + Cu(s)
In this reaction, zinc loses electrons to form Zn²⁺ ions and goes into solution, while copper ions gain electrons and deposit as copper metal.

Example:

Zn(s) + CuSO₄(aq) 121212; ZnSO₄(aq) + Cu(s)

Non-metal Displacement Reactions

Non-metal displacement reactions involve the replacement of one non-metal by another more reactive non-metal, or the displacement of hydrogen from water or acids by metals. These reactions demonstrate the reactivity differences among non-metals and metals with hydrogen.

Non-metal displacement reactions involve the displacement of hydrogen from water or acids by metals, or displacement among halogens. For example, alkali metals displace hydrogen from water, producing hydroxides and hydrogen gas.
Examples:
2Na(s) + 2H₂O(l) 121212; 2NaOH(aq) + H₂(g)
Mg(s) + 2H₂O(g) 121212; Mg(OH)₂(s) + H₂(g)

For example, alkali metals such as sodium displace hydrogen from water, producing metal hydroxides and hydrogen gas:
2Na(s) + 2H₂O(l) 111; 2NaOH(aq) + H₂(g)
Magnesium reacts with steam (water vapour) to form magnesium hydroxide and hydrogen gas:
Mg(s) + 2H₂O(g) 111; Mg(OH)₂(s) + H₂(g)
These reactions show how metals can displace hydrogen from water or acids depending on their reactivity.

These include displacement of hydrogen from water or acids by metals.

Metals such as calcium, magnesium, and zinc can displace hydrogen from water or acids due to their higher reactivity. This displacement results in the formation of metal hydroxides or salts and the release of hydrogen gas. For example, calcium reacts with water to form calcium hydroxide and hydrogen gas:

Ca(s) + 2H₂O(l) → Ca(OH)₂(aq) + H₂(g)

Similarly, zinc reacts with hydrochloric acid to produce zinc chloride and hydrogen gas:

Zn(s) + 2HCl(aq) → ZnCl₂(aq) + H₂(g)

This property is utilized in laboratory preparation of hydrogen gas and demonstrates the reducing nature of these metals.

Examples:

2Na(s) + 2H₂O(l) 111; 2NaOH(aq) + H₂(g)

Mg(s) + 2H₂O(g) 111; Mg(OH)₂(s) + H₂(g)

Examples:

2Na(s) + 2H92O(l) 121212; 2NaOH(aq) + H92(g)

Mg(s) + 2H92O(g) 121212; Mg(OH)92(s) + H92(g)

Disproportionation Reactions

Disproportionation reactions are special redox reactions where a single element undergoes both oxidation and reduction simultaneously, producing two different products containing the element in different oxidation states.

For example, hydrogen peroxide decomposes into water and oxygen:

2H₂O₂ → 2H₂O + O₂

In this reaction, oxygen in H₂O₂ has an oxidation state of -1, while in water it is -2 (reduction) and in oxygen gas it is 0 (oxidation).

Another example is the reaction of phosphorus in alkaline medium:

P₄ + 3OH^- + 3H₂O → PH₃ + 3H₂PO₂^-

Here, phosphorus is both oxidised and reduced, illustrating disproportionation.

Examples:

2H₂O₂ 111; 2H₂O + O₂

P₄ + 3OH^- + 3H₂O 111; PH₃ + 3H₂PO₂^-

In these reactions, an element in one oxidation state is simultaneously oxidised and reduced.

In disproportionation reactions, a single element undergoes both oxidation and reduction simultaneously, resulting in the formation of two different products containing the element in different oxidation states.

Examples:

2H92O92 121212; 2H92O + O92

P94 + 3OH^- + 3H92O 121212; PH93 + 3H92PO92^-

Balancing Redox Reactions

Balancing redox reactions ensures that the number of atoms and the charges are equal on both sides of the chemical equation. Two common methods are used:

Oxidation Number Method

Write correct chemical formulas for reactants and products.
Assign oxidation numbers to all elements and identify those that change.
Calculate the increase and decrease in oxidation numbers and balance them by multiplying appropriately.
In acidic medium, add H⁺ ions; in basic medium, add OH^- ions to balance charges.
Add H₂O molecules to balance hydrogen and oxygen atoms.
Verify that atoms and charges are balanced on both sides.

Half Reaction Method

Write the ionic equation and separate it into oxidation and reduction half reactions.
Balance atoms other than oxygen and hydrogen in each half reaction.
Balance oxygen atoms by adding H₂O and hydrogen atoms by adding H⁺ (acidic medium).
Balance charges by adding electrons.
Multiply half reactions to equalize the number of electrons transferred.
Add the half reactions and cancel electrons to get the balanced equation.
For basic medium, add OH^- ions to both sides to neutralize H⁺ ions and form water.

Oxidation Number Method:

Write correct chemical formulas for reactants and products.
Assign oxidation numbers to all elements and identify those that change.
Calculate the increase and decrease in oxidation numbers and balance them by multiplying appropriately.
In acidic medium, add H⁺ ions; in basic medium, add OH^- ions to balance charges.
Add H₂O molecules to balance hydrogen and oxygen atoms.
Verify that atoms and charges are balanced on both sides.

Half Reaction Method:

Write the ionic equation and separate it into oxidation and reduction half reactions.
Balance atoms other than oxygen and hydrogen in each half reaction.
Balance oxygen atoms by adding H₂O and hydrogen atoms by adding H⁺ (acidic medium).
Balance charges by adding electrons.
Multiply half reactions to equalize the number of electrons transferred.
Add the half reactions and cancel electrons to get the balanced equation.
For basic medium, add OH^- ions to both sides to neutralize H⁺ ions and form water.

Oxidation Number Method

Write correct formulas for reactants and products.
Assign oxidation numbers and identify changes.
Balance changes in oxidation numbers by multiplying.
Add H⁺ or OH^- ions to balance charges in acidic or basic medium.
Add H₂O to balance hydrogen and oxygen atoms.
Check that atoms and charges are balanced.

Half Reaction Method

Write ionic equation and separate into oxidation and reduction half reactions.
Balance atoms other than O and H.
Balance O by adding H₂O and H by adding H⁺ (acidic medium).
Balance charges by adding electrons.
Multiply half reactions to equalize electrons and add them.
For basic medium, add OH^- to both sides to neutralize H⁺.

Electrode Processes

Electrode processes refer to the oxidation and reduction reactions occurring at electrodes during electrochemical reactions. These processes are fundamental to the operation of electrochemical cells and involve electron transfer at the interface between an electrode and an electrolyte.

Direct Redox Reactions

Direct redox reactions occur when oxidation and reduction half reactions take place in the same container. An example is the electrolytic cell, where an external power source drives the redox reaction, such as zinc being oxidised and copper ions being reduced in a single solution.

Oxidation and reduction occur in the same container, e.g., electrolytic cell where Zn is oxidised and CuSO₄ is reduced.

In an electrolytic cell, oxidation and reduction reactions take place simultaneously within the same container. For example, when a zinc strip is immersed in a copper sulfate (CuSO₄) solution, zinc atoms lose electrons and go into solution as Zn²⁺ ions (oxidation), while Cu²⁺ ions in the solution gain electrons and deposit as copper metal on the electrode (reduction). This direct redox process involves electron transfer within a single vessel, driven by an external electric current.

Indirect Redox Reactions

Indirect redox reactions occur when the oxidation and reduction half reactions take place in separate containers. These containers are connected by a salt bridge that allows ion flow to maintain electrical neutrality. This setup is typical of electrochemical cells, where chemical energy is converted into electrical energy through spontaneous redox reactions occurring in different compartments.

Oxidation and reduction occur in separate containers connected by a salt bridge, e.g., Daniel cell with Zn and Cu electrodes.

In an indirect redox reaction, oxidation and reduction occur in separate containers connected by a salt bridge. For example, in the Daniel cell, a zinc electrode is placed in a zinc sulfate solution, and a copper electrode is placed in a copper sulfate solution. The two solutions are connected by a salt bridge, allowing ions to flow and maintain charge balance. Zinc undergoes oxidation, releasing electrons, while copper ions undergo reduction by gaining electrons. This separation of half reactions enables the generation of electrical energy.

Redox Couple

A redox couple consists of the oxidised and reduced forms of a chemical species involved in a redox reaction. It represents the pair of species that can interconvert by gaining or losing electrons. For example, the Zn²⁺/Zn couple includes zinc ions (oxidised form) and zinc metal (reduced form), which participate in oxidation and reduction half reactions respectively.

The oxidised and reduced forms of a substance involved in a half reaction.

The oxidised and reduced forms of a substance involved in a half reaction together constitute a redox couple. This pair participates in electron transfer processes where the oxidised form accepts electrons to become reduced, and the reduced form donates electrons to become oxidised. Understanding redox couples is essential for analyzing electrochemical reactions and electrode potentials.

Electrode

An electrode is a conductor, usually a metal or graphite, that provides a surface for oxidation or reduction reactions to occur in an electrochemical cell. It allows the transfer of electrons between the external circuit and the chemical species in the electrolyte. Electrodes are essential components in galvanic and electrolytic cells, facilitating the flow of electric current through redox reactions.

Electrical conductor making contact with non-metallic part of circuit.

An electrode is an electrical conductor, typically made of metal or graphite, that provides a surface for oxidation or reduction reactions to occur in an electrochemical cell. It allows the flow of electrons between the external circuit and the electrolyte, facilitating the redox process. Electrodes are essential components in galvanic and electrolytic cells, enabling the conversion between chemical and electrical energy.

Electrode Potential

Electrode potential is the voltage developed at the interface between an electrode and its electrolyte due to the tendency of the electrode to lose or gain electrons. It reflects the relative ability of a species to be oxidised or reduced. Electrode potentials are measured under standard conditions and are used to predict the direction of redox reactions and the feasibility of electrochemical cells.

Potential associated with each electrode.

The potential associated with each electrode, known as electrode potential, arises from the difference in the tendency of the electrode material to lose or gain electrons compared to a standard reference electrode. It is measured in volts (V) and depends on factors such as concentration, temperature, and pressure. Standard electrode potentials are measured against the standard hydrogen electrode under standard conditions.

Salt Bridge

A salt bridge is a device used in electrochemical cells to connect the oxidation and reduction half-cells. It typically consists of an inverted U-shaped tube filled with an electrolyte solution, such as potassium chloride (KCl) or ammonium nitrate (NH₄NO₃). The salt bridge allows the flow of ions between the two half-cells, maintaining electrical neutrality and completing the circuit without mixing the different solutions directly.

Inverted U-tube containing electrolyte connecting two half cells, maintaining electrical neutrality.

The salt bridge is an inverted U-tube containing an electrolyte solution that connects the two half-cells in an electrochemical cell. It permits the movement of ions to balance the charge as electrons flow through the external circuit, preventing the buildup of charge that would otherwise stop the redox reaction. This maintains electrical neutrality in both half-cells, allowing the cell to function continuously.

Functions of Salt Bridge

Completes the circuit without mixing solutions.
Maintains electrical neutrality by ion flow.

Representation of Electrochemical Cell

Standard Electrode Potential

Standard electrode potential is the measure of the individual potential of a reversible electrode at standard state, which is 1 molar concentration for each ion participating in the reaction, 1 atmosphere pressure for any gases involved, and a temperature of 298 K (25°C). It is measured relative to the standard hydrogen electrode, which is assigned a potential of zero volts.

Potential of electrode under standard conditions (1 M concentration, 1 atm pressure, 298 K).

The potential of an electrode under standard conditions refers to the voltage developed by a half-cell when all species involved are at unit activity (1 M concentration for solutions, 1 atm pressure for gases) and the temperature is maintained at 298 K (25°C). This potential is called the standard electrode potential and is used to compare the tendency of different electrodes to gain or lose electrons.

Standard Electrode Potential

The standard electrode potential (E) is a quantitative measure of the tendency of a chemical species to be reduced, measured under standard conditions: 1 M concentration for solutions, 1 atm pressure for gases, and temperature at 298 K (25C). It is expressed in volts (V) relative to the standard hydrogen electrode, which is assigned a potential of 0 V. A more positive E value indicates a greater tendency to gain electrons (be reduced), while a more negative value indicates a greater tendency to lose electrons (be oxidized).

The standard electrode potential (E) is the measure of the tendency of a chemical species to be reduced, measured under standard conditions: 1 M concentration for solutions, 1 atm pressure for gases, and temperature at 298 K (25C). It is expressed in volts (V) relative to the standard hydrogen electrode, which is assigned a potential of 0 V.

The standard electrode potential (E) represents the voltage developed by a half-cell under standard conditions when it is connected to the standard hydrogen electrode. It reflects the relative ease with which a species gains electrons (is reduced). The standard hydrogen electrode is assigned a potential of 0 V by convention. Electrode potentials are used to predict the direction of redox reactions and the feasibility of electrochemical cells.

Electrochemical Series

The electrochemical series is a sequence of elements or ions arranged according to their standard electrode potentials (E). It ranks species based on their ability to gain electrons and be reduced. Elements with higher positive electrode potentials are stronger oxidising agents, while those with lower or negative potentials are stronger reducing agents.

Electrochemical Series

The electrochemical series is a list that orders elements and ions by their standard electrode potentials, indicating their relative tendencies to undergo reduction. This series is essential for predicting the outcomes of redox reactions and the feasibility of electrochemical cells.

The electrochemical series is a list of elements or ions arranged in order of their standard electrode potentials (E). It shows the relative tendency of species to gain electrons (be reduced). Elements with higher (more positive) electrode potentials are stronger oxidising agents, while those with lower (more negative) potentials are stronger reducing agents.

The electrochemical series arranges elements or ions in order of their standard electrode potentials (E). It reflects the relative ease with which species gain electrons (are reduced). Species with higher (more positive) electrode potentials are stronger oxidising agents, while those with lower (more negative) potentials are stronger reducing agents.

Arrangement of electrodes in decreasing order of standard reduction potentials.

Electrodes in the electrochemical series are arranged in decreasing order of their standard reduction potentials. This order helps predict which species will be oxidised or reduced in a redox reaction, and determines the direction of electron flow in electrochemical cells.

Arrangement of Electrodes

The arrangement of electrodes in the electrochemical series is based on their standard reduction potentials, listed from highest to lowest. This arrangement aids in understanding the reactivity of metals and the feasibility of displacement and redox reactions.

In the electrochemical series, electrodes are arranged in decreasing order of their standard reduction potentials. This arrangement helps predict the direction of redox reactions, displacement reactions, and the feasibility of electrochemical cells. Metals higher in the series tend to lose electrons and act as reducing agents, while those lower tend to gain electrons and act as oxidising agents.

In the electrochemical series, electrodes are arranged in decreasing order of their standard reduction potentials. Metals higher in the series tend to lose electrons easily and act as reducing agents, while those lower in the series tend to gain electrons and act as oxidising agents. This arrangement is useful for predicting the direction of redox reactions, displacement reactions, and the feasibility of electrochemical cells.

Feasibility of Reaction

Eb0cell = Eb0cathode - Eb0anode

Positive Eb0cell indicates feasible reaction.