Titrimetry, Titration Classifications, Standard solutions, Equivalence Point

Written by: J.A.D.I.Thathsarani, BSc.(Hons) special in Chemistry, University of Kelaniya (UOK)

Introduction of Titrimetric analysis

What is Titration?

The process of addition of standard solution (titrant) from the burette into the solution of unknown concentration (analyte) taken in a conical flask is known as titration. Following requirements should be completed for a titrimetric analysis.

  1. The reaction must be stoichiometric.
  2. The reaction must be relatively fast. (most ionic reactions satisfy this condition)
  3. It must go to completion with no side reactions.
  4. It should alter a physical or chemical property at the completion of the reaction
  5. An indicator should be available to sharply define the endpoint of the reaction.
Titration setup

Titrimetric methods are generally accomplished of high precision and wherever applicable retain noticeable advantages over gravimetric methods. It necessities simpler apparatus, are generally, quickly performed.

The following apparatus and solutions are required for titration.

  1. Calibrated measuring containers including burettes, pipettes and measuring flasks
  2. Substances of known purity for the preparation of the standard solutions
  3. A visual indicator or an instrumental method for detecting the completion of the reaction

To carry out the titration, these steps must be followed.

  1. Add the set volume of unknown solution (analyte) into a clear titration flask using the pipette
  2. Add a few drops of indicator into that flask
  3. Fill the burette with the known solution (titrant)
  4. Record the starting volume of the burette
  5. Slowly add the titrant solution from the burette to the analyte in the flask
  6. Need to slow the flask or use a stir plate
  7. Stop adding titrant solution when reaching to endpoint (appropriate colour change).
  8. Record the Final volume (subtract the Final volume from the initial reading)

In general, each titration is to be performed three times, first to get a rough reading, the other two for accurate determination of the equivalence point. For the calculation, the average (arithmetic mean) of second and third readings is used.

Titration calculation

Nevertheless, of the type of reaction used in the titration experiment, the subsequent calculation is established on the known stoichiometry with which the titrated substance (unknown) reacts with the titrant. While it can be accomplished in several ways, usage of the following simple formula is recommended,

C1V1 = C2V2


  • C1 - concentration of analyte (the unknown)
  • V1 - the volume of the unknown sample taken for titration
  • C2 - concentration of titrant (known)
  • V2 - the volume of the titrant solution consumed for titration until the equivalence point is reached (obtained from experimental data)
titration calculation

Note: The values of the C and V on both sides of the equation must be in the same units. In the case where the stoichiometry of the titration reaction differs from a simple 1:1 ratio, it must be taken into account in the calculation as well.


An unknown sample of sulfuric acid (H2SO4) was titrated with the known KOH solution. It was found that 12 mL of the KOH (C = 0.1 molL-1) was necessary for just complete neutralization of 10 mL of the unknown sample of H2SO4. What is the concentration of the sulfuric acid in the sample?

Equation : H2SO4 + 2KOH → K2SO4 + 2H2O

Calculation : nH2SO4 KOH

  • C1 . V1 = C2 .V2
  • C1 = C2 . V2/ V1
  • C1 = 0.1 * 12/10 = 0.12

The concentration of H2SO4 would be 0.12 mol L-1. However, it goes from the reaction above that 2 moles of KOH are required for neutralization of 1 mol of H2SO4. Therefore, the result, 0.12, must be divided by 2, giving the concentration of sulfuric acid in the unknown sample 0.06

Classification of titrations

Classification of titrations

Standard solutions in titrations

In a titration, certain chemicals are used frequently in defined concentrations as reference solutions. Such substances are referred to as primary standards or secondary standards.

Primary standard

A compound that can be weighed and diluted to get an extract concentration solution is known as the primary standard.

Characteristics of primary standard

  1. They must be of known purity, preferably 100%
  2. Reaction with the reagent to be analyzed should be stoichiometric, complete, fast and selective
  3. They should be easy to handle (weighing and dissolving)
  4. High molecular weight (to minimize weighing errors)
  5. Readily available, inexpensive and easy to dry
  6. Stable in conditions used (in the air & solution)
  7. Should not absorb water (hygroscopic) or carbon dioxide
  8. The titration error should be negligible or easy to determine accurately by experiment.
  9. In practice, an ideal primary standard is difficult to obtain and a compromise between the above ideal requirements is usually necessary.

Example of Primary standards

Bases - Following Bases are used for standardizing acids

  • Sodium carbonate (Na2CO3) - commonly used but the low molecular weight
  • 4- aminopyridine - high purity and stability
  • Sodium tetraborate (Na2B4O7)

Acids - Following Acids are used for standardizing bases

  • Benzoic acid C6H5COOH
  • Potassium hydrageniodate KH(IO3)2
  • Potassium hydrogen phthalate KH(C8H4O4) - most commonly used. High MW (204.2 gmol-1). High purity, thermally stable and reacts fast with NaOH and KOH.
  • 2-Furonic acid - stronger acid than KHP

Other common primary standards

  • For complex formation reactions: Ag, AgNO3, NaCl, various metals (eg: spectroscopically pure zinc, magnesium, copper and manganese)
  • For precipitate reactions: Ag, AgNO3, NaCl, KCl, KBrO3
  • For oxidation-reduction reactions: K2Cr2O7, KBrO3, KIO3, KH(IO3)2, Na2C2O4, As2O3 and pure iron

NaOH, KOH, HCl, HNO3, H2SO4, H3PO4, KMnO4, Na2S23 are not primary standards

Hydrated salts, as a rule, do not make good standards because of the difficulty of efficient drying. However, those salts which do not effloresce, such as sodium tetraborate Na2B4O7.10H2O and copper sulphate CuSO4.5H2O are found by experiment to be satisfactory secondary standards.

Secondary standards

A solution with an approximate concentration is prepared and the exact concentration is established using a primary standard. (standardized)

Equivalence point - Theoretical endpoint

The point where enough titrant (stoichiometric amount) is added to completely react with the titrant (analyte).

The product of the equivalence point volume (Veq) and the titrant's concentration (CT), gives the moles of titrant react with the analyte.

Moles titrant = Veq * CT

Knowing the stoichiometry of titration reaction , we can calculate the moles of the analyte.

Unfortunately, in most titrations, we ordinarily have no obvious indication that an equivalence point has been reached. Instead, we stop adding titrant when we reach an endpoint of our choosing. Often this endpoint is indicated by a change in the colour of the substance added to the solution containing the analyte. Such substances are known as indicators.

The difference between the endpoint volume and equivalence point volume is known as titration error.

The indicator and the experimental conditions should be so selected that the difference between the visible endpoint and the equivalence point is as small as possible.

Minimizing Titration Error

Blank Titration

Titrate the solvent without analyte

Blank titration

Back Titration

A titration in which a reagent is added to a solution containing the analyte, and the excess reagent remaining after its reaction with the analyte is determined by titration.

Add an excess of the standard to the analyte (known amount, but it's too much). Use a secondary standard for titrating excess of the first standard.

Back Titration

Select an easily observable property

  • Observing color change of the indicator
  • Precipitation or Dissolution
  • Detection of color change (Absorption of Light)- Spectrometer method
  • Detection of change in voltage or current (Electrochemistry)

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