5.2.3 Redox and electrode potentials

oxidising agent

accepts electrons and gets reduced

reducing agent

donates electrons anf gets oxidised

how to construct half equations

• check for exceptions

• check if the atom undergoing change in oxidation number change is balanced and if not balance

• remember to take into account balancing when calculating oxidation number change

• for oxidation electrons written on products

• reduction electrons on reactants side

how to contrsyt redox equations

• check if both half equations have same numebr of electrons ifnot multiply

• add 2 half equtions

• check charges on both sides

• if not balances add h+/OH- to reactants and water to products

• check if numebr of atoms are bakances

differences with redox titrations

• In many cases, no additional indicator is needed as the reactants themselves undergo colour changes that signal the end point. This is known as a self-indicating titration.

• The titration determines the amount of oxidising agent needed to exactly react with a given quantity of reducing agent (or vice versa).

why are transition metals good redox reagent s

cThey can readily change oxidation states by accepting or donating electrons, making them effective oxidising or reducing agents.

• These changes in oxidation state are often accompanied by distinct colour changes, which serve as built-in visual indicators.

• The colour changes make it easy to identify the end point of the titration without the need for an additional indicator in many cases.

half equations for reduction of permanganate ion

MnO₄^- + 8H⁺ + 5e^- ➔ Mn²⁺ + 4H₂O

half equation for oxidation of fe2+

 Fe²⁺ ➔ Fe³⁺ + e^-

full equation for Redox titrations with potassium permanganate

Titration procedure for determining MnO₄^- concentration

1. Measure a known volume of the reducing agent solution (e.g., Fe²⁺) using a pipette into a conical flask. Add an excess of dilute sulfuric acid to ensure sufficient H⁺ ions for MnO₄^- reduction.

2. Gradually add MnO₄^- solution from a burette to the flask, swirling constantly. Stop when the mixture becomes faintly tinted with the purple MnO₄^- colour, which indicates the end point. Record the volume of MnO₄^- added.

3. Repeat the titration until concordant results are obtained (titre volumes that differ by no more than 0.10 cm³).

4. Calculate the mean volume of MnO₄^- added from the concordant titres.

half equations for iodine-thiosulfate redox titrations

full eqaution for iodine-thiosulfate redox titration

3 main stages of iodine-thiosulfate redox titration

1. A known volume of the oxidising agent is reacted with an excess of acidified potassium iodide (KI) solution. This oxidises some of the iodide ions (I^-) to iodine (I₂).

2. The iodine produced is then titrated with a sodium thiosulfate (Na₂S₂O₃) solution of known concentration. The thiosulfate ions (S₂O₃^2-) reduce the iodine back to iodide ions.

3. The concentration of the original oxidising agent is calculated using the volume and concentration of the sodium thiosulfate solution used in the titration.

how does the colour chnage occur in sodium thiosulfate titration

Here's how the colour change occurs:

1. During the titration, the dark brown colour of the iodine solution gradually fades as the iodine (I₂) is reduced to colourless iodide ions (I^-) by the thiosulfate ions (S₂O₃^2-) from the sodium thiosulfate solution.

2. When the majority of the iodine has been reduced and only a small amount remains, a few drops of starch solution are added to the titration mixture. Starch forms a deep blue-black complex with iodine, making it much easier to detect the presence of even trace amounts of iodine in the solution.

3. As more sodium thiosulfate is added, the blue-black colour persists until all the remaining iodine is reduced to iodide.

4. At the end point of the titration, the final drop of sodium thiosulfate reduces the last trace of iodine, causing the blue-black colour to disappear suddenly and completely, leaving a colourless solution.

his sharp colour change from blue-black to colourless is easily recognisable and indicates that all the iodine has been consumed, marking the end point of the titration.