Titration - Acid-Base and Oxidation/Reduction (Redox)

Acid-Base Titration

In this section, the assumption will be that the compound that is boldfaced will be the one that is the analyte. Suggested methods, titrants, and indicators will be given in the description that follows.
Strong Acids
Namely hydrochloric and sulfuric acids, the base used to titrate should be a primary standard. This rules out solutions, and solid sodium carbonate works well here. Utilizing between 0.05 and 0.1 g of sodium carbonate will titrate a small volume of hydrochloric acid in the vicinity of 0.2 M. The sodium carbonate must be dissolved in distilled water before the titration commences. Methyl orange is the indicator to use here; it will turn from an orange color (basic from the sodium carbonate solution) to a light pink at the endpoint.
Since vinegar is a weak acid (roughly 4% acetic acid), it needs to be titrated with a strong base. Sodium hydroxide and potassium hydroxide are the easiest to use. If the basic solutions need to be standardized, KHP is an excellent acidic salt to use for a primary standard. Phenolphthalein works as a suitable indicator, turning the solution from clear to light pink.

Redox Titration - Potassium Permanganate

Potassium permanganate is well known as an excellent oxidizing agent.  This is due in large part to the fact that manganese has several oxidation states, the largest of which (+7) is present in potassium permanganate.  Manganese is often reduced to manganese(IV) (as in manganese (IV) oxide) or a soluble manganese(II) salt.  Here are some findings on the benefits and drawbacks of using potassium permanganate as an oxidizing agent for a variety of compounds.
• Copper(I) salts
Namely, the consideration was given towards copper(I) chloride.  Although this compound is green when in solution, it was decided to see what happens since it is possible to oxidize copper(I) to copper(II).  This procedure was halted fairly quickly when it was observed that the acidification of the copper(I) solution oxidizes it to copper(II) (as noted by the light blue color of solution) before any potassium permanganate was added.  In truth, the color of both copper(I) and copper(II) would have probably been difficult to overcome.

• Iron(II) salts
Most redox titration labs out there utilize an iron(II) salt, frequently iron(II) sulfate heptahydrate.  This compound is used for two reasons.  First, it has a large molar mass, making it a good (but not the best) primary standard.  Second, its solution is near colorless, so the pink/purple endpoint is easy to detect.  Color proves to be an issue when dealing with many other compounds.  Iron(II) salts are an excellent choice for a titration utilizing potassium permanganate.
• Hydrogen peroxide
Hydrogen peroxide is a powerful reducing agent in the presence of a strong oxidizing agent such as potassium permanganate. For this reason dilute concentrations of both compounds must be used. A common laboratory experiment is to experimentally determine the concentration of commercially sold hydrogen peroxide (about 3%) via titration with potassium permanganate (usually prepared to be approximately 0.02 M). Concentrated solutions of these reactants will react explosively and must be avoided.
• Lead(II) salts
While lead(II) can oxidize to lead(IV), this is not a feasible reaction for a number of reasons.  Firstly, most lead(II) salts are insoluble in water; truly only lead(II) nitrate can be considered.  Secondly, the acidification of the solution will immediately produce a precipitate.  If sulfuric acid is used, lead(II) sulfate will begin to precipitate from the solution.  Hydrochloric acid will not produce much insoluble lead(II) chloride, in particular when it is heated, but the acid will react with the potassium permanganate yielding chlorine gas.  Consequently, lead(II) is not an option.

• Sodium Oxalate
In the oxalate ion, carbon has a +3 charge.  When reacted with potassium permanganate the carbon is oxidized to the +4 ion in the form of carbon dioxide gas.  Titrations involving the oxalates are a little more involved, but the results were very good.  The reaction below describes the overall redox reaction:

5 C2O4-2 + 2 MnO4- + 16 H+ → 2 Mn+2 + 10 CO2 + 8 H2O

In the trial run performed, 0.51 g of sodium oxalate was titrated against a 0.1M potassium permanganate solution.  The potassium permanganate was a premade solution obtained from a supplier.  You must gently heat the sodium oxalate solution (most sources say between 55°C and 60°C) before beginning the titration.  14.80 mL (0.00148 mol) of potassium permanganate was needed to reach the endpoint.  Stoichiometrically, 0.0015(22) mol should be needed to fully titrate the original amount of oxalate.  From this, it was verified that sodium oxalate is excellent for use in a redox titration with potassium permanganate.
• Tin(II) chloride dihydrate
At first, one might think tin(II) chloride dihydrate would work well as a reducing agent for potassium permanganate.  After all, tin(II) can be oxidized to tin(IV).  However, there are several problems that arise.  First, tin(II) chloride dihydrate, while soluble in water, establishes an equilibrium with Sn(OH)Cl:

SnCl2 (aq) + H2O (ℓ) ⇌ Sn(OH)Cl (s) + HCl (aq)

Addition of chloride ions should shift the equilibrium to the left and maintain a colorless solution of tin(II) ions.  Several sources confirmed that this was possible, but in practice it is not so simple.  Firstly, hydrochloric acid cannot be used to acidify the solution and provide chloride ions.  This is due to the acid's willingness to react with potassium permanganate to produce chlorine gas.  To avoid this, sodium and potassium chloride were used.  After the addition of copious amounts of each salt, the solution remained cloudy white.  After performing several trials the correct stoichiometric ratio of tin(II) needed to react with permanganate was never verified.  In the first trial, 0.66 g of tin(II) chloride dihydrate was massed, an equivalent to 0.0029(2) mol.  9.32 mL (0.000932 mol) of potassium permanganate was needed to reach the endpoint.  Stoichiometrically, 0.0011(6) mol should have been needed to fully titrate the tin.  In the second trial, 0.52 g of tin(II) chloride dihydrate was massed, an equivalent to 0.0023(0) mol.  7.96 mL (0.000796 mol) of potassium permanganate was needed to reach the endpoint.  Stoichiometrically, 0.00092(15) mol should have been needed to fully titrate the tin.  From this data, it was concluded that tin(II) chloride dihydrate was not a good reducing agent.  It is assumed that this is because of a certain quantity of SnOCl that refused to reverse back to tin(II) chloride.

Redox Titration - Iodometry

Iodometry is a special type of redox reaction that takes advantages of the color change of elemental iodine to a colorless iodide ion (or vice versa). Like potassium permanganate reactions, it is self-indicating in nature.


Titrations involve solution stoichiometry. To see a more detailed explanation of what solution stoichiometry is, and how to solve these problems, visit the stoichiometry page.