3.2.5.5 Variable oxidation states

Explain how are Vanadium species in oxidation states IV, III and II formed?

Vanadium species in oxidation states IV, III and II are formed by the reduction of vanadate(V) ions by zinc in acidic solution.

Vanadium(V)

Yellow solution, VO2+​

Vanadium(IV)

Blue solution, VO2+

Vanadium(III)

Green solution, V3+

Vanadium(II)

Violet solution, V2+

Reduction of VO2+​ (Vanadium(V)) to VO2+ (Vanadium(IV))

Yellow to blue

Reduction of VO2+ (Vanadium(IV)) to V3+ (Vanadium(III))

Blue to green

Reduction of V3+ (Vanadium(III)) to V2+ (Vanadium(II))

Green to violet

What is The redox potential, E⦵

It indicates how easily an ion is reduced

Ions with higher, more positive E⦵ values are

more easily reduced to lower oxidation states.

E⦵ values are measured under standard conditions, so the actual redox potential can vary. So what is E⦵ affected by

1.) Ligands - E⦵ assumes metal ions are surrounded by water ligands in aqueous solution.

-The type of ligand bonded to the metal ion can alter the redox potential by stabilizing certain oxidation states.

2.) pH - Acidic conditions provide excess H+ ions needed for reduction of some metal ions. -For example: VO2+(aq) + 2H+(aq) + e- ➔ VO2+(aq) + H2O(l)

-Alkaline conditions favour oxidation reactions that consume OH- ions instead. For example:

-Cr(OH)3(s) + 5OH-(aq) ➔ CrO42-(aq) + 4H2O(l) + 3e-

In general, more acidic solutions

give higher, more positive redox potentials.

The reduction of [Ag(NH3 )2 ]+ (Tollens' reagent) to metallic silver

1.) Aldehydes are reducing agents and can donate electrons to [Ag(NH3)2]+, which is reduced to metallic silver (Ag).

2.) The half-equation for the reduction is:

[Ag(NH3)2]+ + e − → Ag(s) + 2NH3

-The aldehyde itself is oxidized to a carboxylate ion in the process, usually resulting in a colourless solution.

What oxidising agent is commonly used in redox titrations

Potassium permanganate due to its strong oxidizing ability and distinct colour change

Titration procedure

1.) Use a pipette to transfer a measured volume of reducing agent into a conical flask.

2.) Add dilute sulfuric acid to the conical flask to provide excess acidic conditions.

3.) Fill a burette with MnO4-(aq) ions and slowly add it to the conical flask while swirling continuously. The purple MnO4- ions react to produce a colourless solution of Mn2+ ions.

4.) Record the volume of MnO4-(aq) needed to reach the end point. The end point is indicated by the sudden appearance of a purple solution, indicating an excess of MnO4- ions.

5.) Repeat the titrations until concordant titres are obtained (within 0.10 cm3 of each other).

The reduction half-equation

MnO4- + 8H+ + 5e- ➔ Mn2+ + 4H2O

Determining the Mass of Iron in Iron Tablets

1.) Dissolve a crushed iron tablet in dilute sulfuric acid in a conical flask, ensuring all Fe2+ ions are fully dissolved.

2.) Using a burette, add a standardized KMnO4​ solution to the flask.

3.) Titrate until a faint pink colour persists, indicating the endpoint.

4.) Record the volume of KMnO4KMnO4​ used.

(Image has the half equation)

Determining the Percentage of Iron in Steel

1.) Dissolve a measured steel sample in sulfuric acid to produce Fe2+ ions.

2.) Transfer the solution to a conical flask.

3.) Titrate with standardized KMnO4​ solution until a persistent pink colour appears.

4.) Record the volume of KMnO4​ used.

Calculation

Percentage of iron = Mass of Fe/Mass of steel sample

x100

Calculating the Mr​ of Hydrated Ammonium Iron(II) Sulfate and Ethanedioic Acid (C2​ O4^2−​)

1.) Dissolve a known mass of hydrated ammonium iron(II) sulfate or ethanedioic acid in sulfuric acid to ensure full dissolution.

2.) Titrate with standardized KMnO4​ solution, adding until a faint pink endpoint is reached.

3.) Record the volume of KMnO4KMnO4​ used.

-Calculate the moles of the substance based on the titration reaction:

For ethanedioic acid:

C2O4^2− + MnO4^− + 8H+ → 2CO2 + Mn2++4H2OC2​O42−​+MnO4−​+8H+→2CO2 ​+ Mn2+ + 4H2​O Calculate the Mr​ by dividing the mass of the compound by the moles calculated from the titration data.

Finding the Concentration of Hydrogen Peroxide (H2​O2​) in Hair Bleach

1.) Add a measured volume of the H2O2H2​O2​ solution to a conical flask and add sulfuric acid.

2.) Titrate with standardized KMnO4​ solution until a persistent pink colour indicates the endpoint.

3.) Record the volume of KMnO4KMnO4​ used.

-Calculate the concentration of H2​O2​ based on the reaction:

5H2O2 + 2MnO4−+ 6H+ → 5O2 + 2Mn^2+ + 8H2O

-Use the volume and concentration of KMnO4​ to find the moles of H2​O2​ and calculate its concentration in the bleach.