Presented by : Arvind Singh Heer MSc-I (Sem-I) Inorganic Chemistry MITHIBAI COLLEGE

REDOX REACTION

CONTENT Introduction. Oxidizing agent and Reducing agent. Identifying OX, RD, SI Species. Mechanisms and their characteristics. Balancing Redox Reactions. Redox reaction in aqueous media.

INTRODUCTION What is redox reaction? Redox reaction include all chemical reactions in

which atoms have their oxidation state changed; in general, redox reactions involve the transfer of electrons between species.

The term "redox" comes from two concepts involved with electron transfer: reduction and oxidation. It can be explained in simple terms:

• Oxidation is the loss of electrons or an increase in oxidation state by a molecule, atom, or ion.

• Reduction is the gain of electrons or a decrease in oxidation state by a molecule, atom, or ion.

OXIDIZING AND REDUCING AGENT

An oxidizing agent (also called an oxidant or oxidizer) is a chemical compound that readily transfers oxygen atoms or a substance that gains electrons in a redox chemical reaction. 

An oxidizing agent oxidizes other substances and gains electrons; therefore, its oxidation state will decrease.

An reducing agent (also called a reductant or reducer) is the element or a compound in a redox reaction that reduces another species. In doing so, it becomes oxidized, and is therefore a electron donor in redox reaction.

A reducing agent reduces other substances and lose electrons; therefore, its oxidation state will increase.

In simple terms:  * The oxidizing agent is reduced. 

* The reducing agent is oxidized.  * All atoms in a molecule can be assigned an oxidation number. This number changes when an oxidant acts on a substrate.  * Redox reactions occur when electrons are exchanged.

A mnemonic for differentiating the reactions is "OIL RIG": Oxidation Is Loss, Reduction Is Gain (of electrons) or "LEO the lion says GER" (Lose Electrons: Oxidation, Gain Electrons: Reduction).  

EXAMPLECa0 + 2 H+1Cl-1 Ca+2Cl-1

2 + H20

Since Ca0 is being oxidized and H+1 is being reduced, the electrons must be going from the Ca0 to the H+1.

Since Ca0 would not lose electrons (be oxidized) if H+1 weren’t there to gain them, H+1 is the cause, or agent, of Ca0’s oxidation. H+1 is the oxidizing agent.

Since H+1 would not gain electrons (be reduced) if Ca0 weren’t there to lose them, Ca0 is the cause, or agent, of H+1’s reduction. Ca0 is the reducing agent.

IDENTIFYING OX, RD, SI SPECIES

Ca0 + 2 H+1Cl-1 Ca+2Cl-12 + H20

Oxidation = loss of electrons. The species becomes more positive in charge. For example, Ca0 Ca+2, so Ca0 is the species that is oxidized.

Reduction = gain of electrons. The species becomes more negative in charge. For example, H+1 H0, so the H+1 is the species that is reduced.

Spectator Ion = no change in charge. The species does not gain or lose any electrons. For example, Cl-1 Cl-1, so the Cl-1 is the spectator ion.

MECHANISMS AND THEIR CHARACTERISTICS

The Outer Sphere Electron Transfer Mechanism = electron transfer with no change of coordination spherea)Example: Co(NH3)6

3+ + Cr(bipy)32+ Co(NH3)6

2+ + Cr(bipy)3

3+

oxidant reductant reduced oxidized

b)The rates of reaction depend on the ability of the electron to “tunnel” through the ligands from one metal to the other

Tunneling = moving through an energy barrier (the ligands) that is normally too high to allow the electron to pass through. This is a quantum mechanical process having to do with the wave nature of e-.

Ligands with p or p orbitals good for bonding more easily allow tunnelling (CN-, F-) than those that don’t (NH3).

c)The ligands don’t change in Outer Sphere electron transfer, but the M—L bond distances do

High Oxidation # = short bond distance Low Oxidation # = longer bond distance

d)The stronger the ligand field, the less favored reduction is (or more favored oxidation is), because more energy is gained by losing high energy eg* electrons (NH3 > H2O)

Co(NH3)63+ + e- Co(NH3)6

2+ Eo = +0.108 VCo(H2O)6

3+ + e- Co(H2O)62+ Eo = +1.108 V

The Inner Sphere Electron Transfer Mechanism = tunneling of an electron through a bridging ligand.

Substitution links the reactants e- transfer Separation of products This reaction could be followed by ion

exchange and UV-Vis

BALANCING REDOX REACTIONS

STEP 1. Split Reaction into 2 Half-Reactions STEP 2. Balance Elements Other than H & O STEP 3. Balance O by Inserting H2O into eqns. as

necessary STEP 4. Balance H with H+ or H2O (see 4a, 4b) STEP 5. Balance Charge by Inserting Electrons as

needed STEP 6. Multiply Each 1/2 Reaction by Factor needed

to make no. of Electrons in each 1/2 Reaction Equal STEP 7. Add Eqns. & Cancel Out Duplicate terms,

where possible

STEP 4a. In ACID: Balance H by Inserting H+, as needed

STEP 4b. In BASE: Balance H by (i) inserting 1 H2O for each missing H & (ii) inserting same no. of OH- on OTHER SIDE OF REACTION as H2Os added in (i)

EXAMPLE Complete and Balance Following Reaction: CuS (s) + NO3

- (aq) Cu2+(aq) + SO42- (aq) + NO (g)

STEP1. Split into 2 Half-Reactions a.1 CuS Cu2+ + SO4

2- b.1 NO3

- NO STEP 2. Balance Elements Other than H & O (It is already balanced). STEP 3. Balance O by inserting H2O into equations as necessary. a.3 CuS + 4H2O Cu2+ + SO4

2- b.3 NO3

- NO + 2H2O STEP 4. ACIDIC, so Balance H by inserting H+ as needed. a4. CuS + 4H2O Cu2+ + SO4

2- + 8H+

b4. NO3- + 4H+ NO + 2H2O

STEP 5. Balance Charge by inserting Electrons, where necessary

a5. CuS + 4H2O Cu2+ + SO42- + 8H+ + 8e-

b5. NO3- + 4H+ + 3e- NO + 2H2O

STEP 6. Multiply each Eqn. by factor to make No. of Electrons in Each 1/2 Reaction the Same

a6. Multiply by 3x 3CuS + 12H2O 3Cu2+ + 3SO4

2- + 24H+ + 24e- b6.Multiply by 8x 8NO3

- + 32H+ + 24e- 8NO + 16H++ 24e-

STEP 7. Add Eqns. and Cancel Out Duplicated Terms (a7 + b7) 8H+

3CuS + 12H2O + 8NO3- + 32H+ + 24 e- 3Cu2+ + 3SO4

2- + 24H+ + 8NO +16H2O +24e-

4H2O So, the final, balanced reaction is:

3CuS(s) + 8 NO3-(aq) + 8H+ (aq) 3Cu2+(aq) + 3 SO4

2-(aq) + 8NO(g) + 4H2 O(l)

Redox Reactions in Aqueous Media

There are three main methods of predicting redox reactions in aqueous solutions, summarising the thermodynamic stabilities of oxidation states of elements in aqueous solutions.

1. LATIMER DIAGRAM2. VOLT EQUIVALENT [Also known as Frost (Free energy oxidation

state) diagramThese two are usually restricted to extremes of pH=0 or pH=14 solutions.

3. Pourbaix diagramIt expresses the variation in stabilities of oxidation states as a function of pH between pH values 0 to 14.

Latimer Diagrams It is a list of various oxidation states of an element arranged in

descending order from left to right.

The numerical value of standard potential is written over a horizontal line connecting species of element in different oxidation states.

Most oxidised form is on the left.

Species to the right are in successively lower oxidation states.

Available for all the elements exhibiting more than one oxidation state.

Latimer diagram for Manganese [acidic medium]

The Latimer diagram for Mn illustrates its standard reduction potentials (in 1M HCl) in oxidation states from +7 to 0.

It compresses into shorthand notation all the standard potentials for redox reactions of element Mn.

Values for multi-electron reactions can be also calculated by first adding ∆Gº (nFEº) values and then dividing by the total no of electrons.

Calculation of values for multi-electron reactions by first adding ΔG°(=-nFE°) values and then dividing by the total number of electrons

for the 5-electron reduction of MnO4- to Mn2+, we

write

for the three-electron reduction of MnO4-(aq) to

MnO2(s),

Thermodynamically stable & unstable oxidation states

An unstable species on a Latimer diagram will have a lower standard potential to the left than to the right.

2 MnO4-3 → MnO2 + MnO4

2- ; MnO4-3 is unstable.

Eº = +4.27 − 0.274 = +3.997V ; (spontaneous disproportionation) Which Mn species are unstable with respect to

disproportionation? MnO4

-3 ; 5+ → 6+, 4+ Mn3+ ;  3+→4+, 2+ So stable species are MnO4

-, MnO2, MnO42-, Mn2+, Mn0

Thermodynamically unstable ions can be quite stable kinetically.

Disproportionation In most redox reactions atoms of one element are

oxidized and atoms of a different element are reduced.  In some redox reactions a single substance can be both

oxidized and reduced. These are known as disproportionation reactions.

Example : Decomposition reaction of H2O2

2 H2O2(aq) → 2 H2O(l) + O2(g) The decomposition reaction of hydrogen peroxide

produces oxygen and water. Oxygen is present in all parts of the chemical equation

and as a result it is both oxidized and reduced.

Frost Diagrams In a Frost diagram, we plot ΔG°⁄F (=

nE°) vs. oxidation number.  The zero oxidation state is assigned

a nE° value of zero. Contains the same information as in

a Latimer diagram, but graphically shows stability and oxidizing power.

Stable and unstable oxidation states can be easily identified in the plot.

The standard potential for any electrochemical reaction is given by the slope of the line connecting the two species on a Frost diagram. 

What You Can Learn From a Frost Diagram

Thermodynamic stability is found at the bottom of the diagram. Thus, the lower a species is positioned on the diagram, the more thermodynamically stable it is (from a oxidation-reduction perspective)Mn (II) is the most stable species.

Any species located on the upper left side of the diagram will be a strong oxidizing agent. MnO4

- is a strong oxidizer. Any species located on the upper left side

of the diagram will be a strong oxidizing agent. MnO4

- is a strong oxidizer. The information obtained from a Frost

diagram is for species under standard conditions (pH=0 for acidic solution and pH=14 for basic solution). 

Pourbaix Diagrams Plots of E versus pH for various couples in oxidation of an

element.

The Pourbaix diagram is a type of predominance diagram -- it shows the predominate form in an element will exist under a given set of environmental conditions.

These diagrams give a visual representation of the oxidizing and reducing abilities of the major stable compounds of an element and are used frequently in geochemical, environmental and corrosion applications.

 Like Frost diagrams, Pourbaix diagrams display thermodynamically preferred species.

How to Read a Pourbaix Diagram

Vertical lines separate species that are in acid-base equilibrium.

Non vertical lines separate species related by redox equilibria. Horizontal lines separate species in

redox equilibria not involving hydrogen or hydroxide ions.

Diagonal boundaries separate species in redox equilibria in which hydroxide or hydrogen ions are involved.

Dashed lines enclose the practical region of stability of the water solvent to oxidation or reduction.

What You Can Learn From a Pourbaix Diagram

Any point on the diagram will give the thermodynamically most stable form of that element at a given potential and pH condition.

Strong oxidizing agents and oxidizing conditions are found only at the top of Pourbaix diagrams.Strong oxidizing agents have lower boundaries that are also high on the diagram.Permanganate is an oxidizing agent over all pH ranges. It is very strongly oxidizing at low pH.

A species that ranges from the top to the bottom of the diagram at a given pH will have no oxidizing or reducing properties at that pH.

Reducing agents and reducing conditions are found at the bottom of a diagram and not elsewhere.Strong reducing agents have low upper boundaries on the diagram. Manganese metal is a reducing agent over all pH ranges and is strongest in basic conditions.

When the predominance area for a given oxidation state disappears completely above or below a given pH and the element is in an intermediate oxidation state, the element will undergo disproportionation MnO4

2- tends to disproportionate.

REFRENCES1. Rayner- Canham, G. Descriptive Inorganic Chemistry; Freeman:New

York,1996; Chapter 9

2. Douglas, B; McDaniel, D.; Alexander, J. Concepts and Models of Inorganic Chemistry, 3rd ed.; Wiley & Sons: New York, 1994; Chapter 8.

3. J. Kotz, P. Treichel, J. Townsend, D. Treiche, Chemistry & Chemical Reactivity, 9th ed. ; Cengage Learning.

4. J.E. Huheey, E.A. Keiter, R.L. Keiter, O.K. Medhi, Inorganic Chemistry: Principles of Structure and Reactivity, 4th ed. ; Pearson Education.

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