KMn04 with iron(II) ammonium sulphate hexahydrate-redox titration.
1-what is the purpose of adding H2SO4 to iron sample before the titration?
2-why is there no indicator used to determine the end point for this titration?
3-suggest another possible oxidising agent and reducing agent can be used to replace potassium permanganate & iron (II) ammonium sulphate haxahydrate in this exp.
H2SO4 is added to make sure the solution is acid. KMnO4 behaves differently in acid vs basic solutions.
No separate indicator is needed since KMnO4 serves as its own indicator.
#3 is confusing to me. I don't know what the question is asking.
1. The purpose of adding H2SO4 to the iron sample before the titration is to create an acidic environment. Iron (II) ions (Fe2+) are more stable in acidic solutions and do not easily oxidize to iron (III) ions (Fe3+) in the presence of oxygen in the air. By adding H2SO4, the iron (II) ions are stabilized and protected from oxidation before the actual titration.
2. In this titration, no indicator is used to determine the endpoint because the reaction between potassium permanganate (KMnO4) and iron (II) ammonium sulphate hexahydrate is self-indicating. KMnO4 is a strong oxidizing agent, and it reacts completely with iron (II) ions, turning the solution from colorless to a pink or purple color. The disappearance of this color change indicates that all the iron (II) ions have been oxidized. Therefore, the color change of the reaction itself serves as an indicator for the endpoint of the titration.
3. Another possible oxidizing agent that can be used is cerium (IV) sulfate (Ce(SO4)2). It can oxidize iron (II) ions to iron (III) ions in a similar manner to potassium permanganate. A possible reducing agent that can be used is sodium thiosulfate (Na2S2O3). It can reduce cerium (IV) ions back to cerium (III) ions, completing a redox reaction. However, it's important to note that the specific conditions and requirements of the experiment should be considered before substituting the oxidizing and reducing agents, as different agents may have different reaction rates and specificities.
1. The purpose of adding H2SO4 to the iron sample before the titration is to acidify the solution. Acidification helps to ensure that the reduction of iron(II) to iron(III) proceeds smoothly. In the presence of acidic conditions, iron(II) reacts with permanganate ions (MnO4-) to form iron(III) ions (Fe3+) and the reduced product, Mn2+.
2. In this particular titration, no indicator is used to determine the end point because potassium permanganate (KMnO4) itself acts as a self-indicating agent. KMnO4 is a purple-colored solution, and during the titration, it reacts with iron(II) ions (Fe2+) to form a colorless solution of iron(III) ions (Fe3+) and Mn2+ ions. The solution changes from purple to colorless when all the iron(II) ions are oxidized. Thus, the disappearance of the purple color indicates the end point of the titration.
3. Another possible oxidizing agent that could be used in place of potassium permanganate is hydrogen peroxide (H2O2). Hydrogen peroxide is commonly used as an oxidizing agent in various reactions. It could potentially oxidize iron(II) ions to iron(III) ions. However, the titration conditions, including the concentration of hydrogen peroxide, pH, and stoichiometry of the reaction, would need to be carefully established.
As for a possible reducing agent to replace iron(II) ammonium sulfate hexahydrate, one example could be sodium sulfite (Na2SO3). Sodium sulfite is often used in redox reactions as a reducing agent. It could potentially reduce an oxidizing agent, such as potassium permanganate. Again, the specific reaction conditions would need to be optimized for accurate titration.