Chemists employ many types of redox titrations, the common element being a combination of oxidation (electron loss) and reduction (electron gain) reactions. The chemistry can be quite challenging, but the concepts are rather simple.
Reduction refers to the chemical process of gaining electrons. For example, silver ions can undergo reduction to silver metal (metals always have a charge of zero):
Ag⁺(aq) + e⁻ ' Ag(s).
Note that the silver(I) ion gains an electron, but its charge is decreased (from +1 to 0), thus the term "reduction."
Oxidation is the opposite process from reduction. It refers to the chemical process of gaining electrons. A typical example is the conversion of iron metal to iron(III) ions:
Fe(s) ' Fe³⁺ + 3e⁻.
"Redox" is the shortened form of "reduction-oxidation reaction." This is the generic term used in describing the processes previously discussed. Oxidation and reduction always occur as a pair: Nothing can be oxidised unless something else is reduced, and vice versa.
Redox titrations are frequently used to determine the amount of a given substance in a compound or mixture. An example is the determination of the amount of hydrogen peroxide (H₂O₂) in an over-the-counter peroxide solution. In this case, a strong oxidiser is needed, such as potassium permanganate (KMnO₄):
5 H₂O₂(aq) + 6 HCl(aq) + 2 KMnO₄(aq) ' 5 O₂(g) + 2 MnCl₂(aq) + 8 H₂O(l).
If the concentration of the permanganate solution is known, the chemist can calculate the amount of H₂O₂ in the original sample.
In addition to the analysis of hydrogen peroxide, other common redox titrations performed in chemistry teaching labs include analysis of household bleach by titration with sodium thiosulfate, analysis of iron in iron ore by titration with potassium permanganate and analysis of vitamin C in orange juice by titration with iodine.
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