Redox Titrations – Introduction and Basic Concepts
Redox titrations are based on chemical reactions in which oxidation and reduction take place simultaneously. In these titrations, a solution of known concentration (called the titrant) is added to another solution of unknown concentration until the reaction reaches completion, which is marked by a distinct end point.
- These reactions involve the transfer of electrons between reacting species.
- Oxidizing agents lose oxygen (or gain electrons), while reducing agents gain oxygen (or lose electrons).
Oxidation and Reduction
Oxidation can be defined in two ways:
- Old definition: Combination of a substance with oxygen.
- Modern definition: Loss of electrons. (Example: Na → Na+ + e–)
Reduction can be defined as:
- Old definition: Removal of oxygen from a compound.
- Modern definition: Gain of electrons. (Example: Cl + e– → Cl–)
In any redox reaction, oxidation and reduction always occur together — when one substance loses electrons, another gains them.
Oxidation Number or Oxidation State (OS)
The oxidation number of an element in a compound represents the number of electrons gained or lost by the atom.
- A positive oxidation state means loss of electrons.
- A negative oxidation state means gain of electrons.
Example: In potassium permanganate (KMnO4), the total charge is zero, so the sum of oxidation numbers of all atoms equals zero.
For KMnO4: K (+1) + Mn (x) + 4 × O (–2) = 0 → Mn = +7
Balancing Simple Redox Reactions
Redox reactions are balanced by separating them into two half-reactions — one for oxidation and the other for reduction.
Example: Copper and silver ions
- Ag+ + e– → Ag (Reduction)
- Cu → Cu2+ + 2e– (Oxidation)
Balance the number of electrons transferred, then combine both half-reactions to get the overall balanced redox equation.
Equivalent Weight of Oxidizing and Reducing Agents
The equivalent weight of a substance depends on the chemical reaction it undergoes, not just on its molecular weight.
Example 1: FeSO4 oxidized to Fe2(SO4)3 by KMnO4 in acidic medium.
Example 2: In acidic medium, MnO4– (permanganate ion) reduces to Mn2+. In neutral medium, it reduces to MnO2, and in basic medium, to MnO42–.
Theory of Redox Titrations
Redox titrations use a special type of electrochemical cell that has two electrodes:
- Indicator Electrode: Usually a platinum (Pt) electrode that transfers electrons between the solution and the external circuit.
- Reference Electrode: Has a fixed potential, such as the Standard Hydrogen Electrode (SHE) or Standard Calomel Electrode (SCE).
The potential difference between these electrodes changes as the titration proceeds, helping to identify the end point.
Redox Indicators
Redox indicators are substances that change color due to a reversible oxidation-reduction process near the equivalence point.
- The oxidized and reduced forms of the indicator have different colors.
- The color change occurs sharply near the end point.
The Nernst equation describes the relationship between the potential and the ratio of oxidized to reduced forms of the indicator.
Types of Indicators in Redox Titrations
1. Self Indicators
In some titrations, the titrant itself acts as an indicator.
Example: Potassium permanganate (KMnO4) acts as its own indicator. The solution remains colorless until the end point, when a faint pink color appears due to excess KMnO4.
2. External Indicators
External indicators react with the solution outside the titration flask to show color change.
Example: Potassium ferricyanide gives a blue color with ferrous ions during the titration of Fe2+ with dichromate. The color disappears when all Fe2+ ions are oxidized to Fe3+.
3. Internal or Redox Indicators
These indicators participate in the reaction and show different colors in their oxidized and reduced forms. Most are organic dyes that become colorless when reduced.
4. Potentiometric Method
In some redox titrations, instead of visual color change, the electrode potential is measured using instruments. This method is especially useful for colored or dilute solutions.
Oxidation–Reduction Curves
During a redox titration, the potential (E) of the solution changes gradually, then rapidly near the equivalence point. The shape of the redox curve helps determine the end point accurately using instruments or indicators.
Common Redox Titrations
1. Iodimetry and Iodometry
- Iodimetry: Uses standard iodine solution as the titrant to determine reducing agents like Na2S2O3 (thiosulphate) or SnCl2.
- Iodometry: Involves titration of iodine liberated during a reaction, usually with thiosulphate as titrant.
Indicator: Starch solution, which forms a blue complex with iodine that disappears at the end point.
2. Bromatometry
Potassium bromate (KBrO3) acts as a strong oxidizing agent in acidic solution. It is used to determine substances like arsenic, antimony, and iodides.
Indicator: Methyl orange (changes from red to colorless).
3. Dichrometry
Potassium dichromate (K2Cr2O7) is another strong oxidizing agent used in acidic medium. It is more stable than KMnO4 and not affected by chloride ions.
Indicator: Diphenylamine (colorless to deep blue at the end point).
Use: Determination of ferrous salts (Fe2+).
4. Titration with Potassium Iodate (KIO3)
Potassium iodate acts as a reliable oxidizing agent and can be used as a primary standard. It is often used in iodometric methods to standardize thiosulphate solutions.
Common Oxidizing and Reducing Agents
Oxidizing Agents:
- Potassium permanganate (KMnO4)
- Potassium dichromate (K2Cr2O7)
- Potassium iodate (KIO3)
- Potassium bromate (KBrO3)
- Ceric sulphate (Ce(SO4)2)
- Iodine (I2)
Reducing Agents:
- Ferrous sulphate (FeSO4)
- Sodium thiosulphate (Na2S2O3)
- Stannous chloride (SnCl2)
- Oxalic acid (H2C2O4)
Among these, compounds like potassium iodate, potassium bromate, potassium dichromate, and arsenious oxide can serve as primary standards — their standard solutions can be prepared directly and used for accurate titrations.
Detailed Notes:
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