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Titration (also called titrimetry[1] and volumetric analysis) is a common laboratory method of quantitative chemical analysis to determine the concentration of an identified analyte (a substance to be analyzed). A reagent, called the titrant or titrator,[2] is prepared as a standard solution of known concentration and volume; the titrant reacts with a solution of analyte (which may also be called the titrand[3]) to determine the analyte's concentration. The volume of titrant that reacted with the analyte is called the titration volume.

History and etymology[edit]

The word "titration" descends from the French word tiltre (1543), meaning the "proportion of gold or silver in coins or in works of gold or silver"; i.e., a measure of fineness or purity. Tiltre became titre,[4] which thus came to mean the "fineness of alloyed gold",[5] and then the "concentration of a substance in a given sample".[6] In 1828, the French chemist Gay-Lussac first used titre as a verb (titrer), meaning "to determine the concentration of a substance in a given sample".[7]

Volumetric analysis originated in late 18th-century France. François-Antoine-Henri Descroizilles (fr) developed the first burette (which was similar to a graduated cylinder) in 1791.[8][9][10] Joseph Louis Gay-Lussac developed an improved version of the burette that included a side arm, and coined the terms "pipette" and "burette" in an 1824 paper on the standardization of indigo solutions;[11] the first true burette was invented in 1845 by the French chemist Étienne Ossian Henry (1798–1873).[12][13] A major breakthrough in the methodology and popularization of volumetric analysis was due to Karl Friedrich Mohr, who redesigned the burette into a simple and convenient form, and who wrote the first textbook on the topic, Lehrbuch der chemisch-analytischen Titrirmethode (Textbook of analytical chemistry titration methods), published in 1855.[14][15]


Analysis of soil samples by titration

A typical titration begins with a beaker or Erlenmeyer flask containing a very precise volume of the analyte and a small amount of indicator (such as phenolphthalein) placed underneath a calibrated burette or chemistry pipetting syringe containing the titrant. Small volumes of the titrant are then added to the analyte and indicator until the indicator changes color in reaction to the titrant saturation threshold, reflecting arrival at the endpoint of the titration. Depending on the endpoint desired, single drops or less than a single drop of the titrant can make the difference between a permanent and temporary change in the indicator; when the endpoint of the reaction is reached, the volume of reactant consumed is measured and used to calculate the concentration of analyte by


Preparation techniques[edit]

Typical titrations require titrant and analyte to be in a liquid (solution) form. Though solids are usually dissolved into an aqueous solution, other solvents such as glacial acetic acid or ethanol are used for special purposes (as in petrochemistry).[17] Concentrated analytes are often diluted to improve accuracy.

Many non-acid–base titrations require a constant pH throughout the reaction. Therefore, a buffer solution may be added to the titration chamber to maintain the pH.[18]



Titration curves[edit]

A typical titration curve of a diprotic acid titrated with a strong base. Shown here is oxalic acid titrated with sodium hydroxide. Both equivalence points are visible.


In an acidbase titration, the titration curve reflects the strength of the corresponding acid and base. For a strong acid and a strong base, the curve will be relatively smooth and very steep near the equivalence point; because of this, a small change in titrant volume near the equivalence point results in a large pH change and many indicators would be appropriate (for instance litmus, phenolphthalein or bromothymol blue).


Titrations between a weak acid and a weak base have titration curves which are highly irregular; because of this, no definite indicator may be appropriate and a pH meter is often used to monitor the reaction.[22]


Types of titrations[edit]

There are many types of titrations with different procedures and goals; the most common types of qualitative titration are acid–base titrations and redox titrations.

Acid–base titration[edit]

Methyl orange
Indicator Color on acidic side Range of color change
Color on basic side
Methyl violet Yellow 0.0–1.6 Violet
Bromophenol blue Yellow 3.0–4.6 Blue
Methyl orange Red 3.1–4.4 Yellow
Methyl red Red 4.4–6.3 Yellow
Litmus Red 5.0–8.0 Blue
Bromothymol blue Yellow 6.0–7.6 Blue
Phenolphthalein Colorless 8.3–10.0 Pink
Alizarin yellow Yellow 10.1–12.0 Red

Acid–base titrations depend on the neutralization between an acid and a base when mixed in solution. In addition to the sample, an appropriate pH indicator is added to the titration chamber, reflecting the pH range of the equivalence point; the acid–base indicator indicates the endpoint of the titration by changing color. The endpoint and the equivalence point are not exactly the same because the equivalence point is determined by the stoichiometry of the reaction while the endpoint is just the color change from the indicator. Thus, a careful selection of the indicator will reduce the indicator error. For example, if the equivalence point is at a pH of 8.4, then the Phenolphthalein indicator would be used instead of Alizarin Yellow because phenolphthalein would reduce the indicator error. Common indicators, their colors, and the pH range in which they change color are given in the table above;[23] when more precise results are required, or when the reagents are a weak acid and a weak base, a pH meter or a conductance meter are used.

For very strong bases, such as organolithium reagent, metal amides, and hydrides, water is generally not a suitable solvent and indicators whose pKa are in the range of aqueous pH changes are of little use. Instead, the titrant and indicator used are much weaker acids, and anhydrous solvents such as THF are used.[24][25]

Phenolphthalein, a commonly used indicator in Acid and Base Titration

The approximate pH during titration can be approximated by three kinds of calculations. Before beginning of titration, the concentration of is calculated in aqueous solution of weak acid before adding any base. When the number of moles of bases added equals the number of moles of initial acid or so called equivalence point, one of hydrolysis and the pH is calculated in the same way that the conjugate bases of the acid titrated was calculated. Between starting and end points, is obtained from the Henderson-Hasselbalch equation and titration mixture is considered as buffer. In Henderson-Hasselbalch equation the [acid] and [base] are said to be the molarities that would have been present even with dissociation or hydrolysis. In a buffer, can be exactly calculated but the dissociation of HA, the hydrolysis of and self-ionization of water must be taken into account.[26] Four independent equations must be used:[27]

In the equations, and are the moles of acid (HA) and salt (XA where X is the cation), respectively, used in the buffer, and the volume of solution is V. The law of mass action is applied to the ionization of water and the dissociation of acid to derived the first and second equations; the mass balance is used in the third equation, where the sum of and must equal to the number of moles of dissolved acid and base, respectively. Charge balance is used in the fourth equation, where the left hand side represents the total charge of the cations and the right hand side represents the total charge of the anions: is the molarity of the cation (e.g. sodium, if sodium salt of the acid or sodium hydroxide is used in making the buffer).[28]

Redox titration[edit]

Redox titrations are based on a reduction-oxidation reaction between an oxidizing agent and a reducing agent. A potentiometer or a redox indicator is usually used to determine the endpoint of the titration, as when one of the constituents is the oxidizing agent potassium dichromate; the color change of the solution from orange to green is not definite, therefore an indicator such as sodium diphenylamine is used.[29] Analysis of wines for sulfur dioxide requires iodine as an oxidizing agent. In this case, starch is used as an indicator; a blue starch-iodine complex is formed in the presence of excess iodine, signalling the endpoint.[30]

Some redox titrations do not require an indicator, due to the intense color of the constituents. For instance, in permanganometry a slight persisting pink color signals the endpoint of the titration because of the color of the excess oxidizing agent potassium permanganate.[31] In iodometry, at sufficiently large concentrations, the disappearance of the deep red-brown triiodide ion can itself be used as an endpoint, though at lower concentrations sensitivity is improved by adding starch indicator, which forms an intensely blue complex with triiodide.

Color of iodometric titration mixture before (left) and after (right) the end point

Gas phase titration[edit]

Gas phase titrations are titrations done in the gas phase, specifically as methods for determining reactive species by reaction with an excess of some other gas, acting as the titrant. In one common gas phase titration, gaseous ozone is titrated with nitrogen oxide according to the reaction

O3 + NO → O2 + NO2.[32][33]



Complexometric titration[edit]

Complexometric titrations rely on the formation of a complex between the analyte and the titrant. In general, they require specialized complexometric indicators that form weak complexes with the analyte; the most common example is the use of starch indicator to increase the sensitivity of iodometric titration, the dark blue complex of starch with iodine and iodide being more visible than iodine alone. Other complexometric indicators are Eriochrome Black T for the titration of calcium and magnesium ions, and the chelating agent EDTA used to titrate metal ions in solution.[35]

Zeta potential titration[edit]

Zeta potential titrations are titrations in which the completion is monitored by the zeta potential, rather than by an indicator, in order to characterize heterogeneous systems, such as colloids.[36] One of the uses is to determine the iso-electric point when surface charge becomes zero, achieved by changing the pH or adding surfactant. Another use is to determine the optimum dose for flocculation or stabilization.[37]


An assay is a form of biological titration used to determine the concentration of a virus or bacterium. Serial dilutions are performed on a sample in a fixed ratio (such as 1:1, 1:2, 1:4, 1:8, etc.) until the last dilution does not give a positive test for the presence of the virus. The positive or negative value may be determined by visually inspecting the infected cells under a microscope or by an immunoenzymetric method such as enzyme-linked immunosorbent assay (ELISA); this value is known as the titer.[38]

Measuring the endpoint of a titration[edit]

Different methods to determine the endpoint include:[39]

  • Indicator: A substance that changes color in response to a chemical change. An acid–base indicator (e.g., phenolphthalein) changes color depending on the pH. Redox indicators are also used. A drop of indicator solution is added to the titration at the beginning; the endpoint has been reached when the color changes.
An elementary pH meter that can be used to monitor titration reactions
  • Conductivity: A measurement of ions in a solution. Ion concentration can change significantly in a titration, which changes the conductivity. (For instance, during an acid–base titration, the H+ and OH ions react to form neutral H2O.) As total conductance depends on all ions present in the solution and not all ions contribute equally (due to mobility and ionic strength), predicting the change in conductivity is more difficult than measuring it.
  • Precipitation: If a reaction produces a solid, a precipitate will form during the titration. A classic example is the reaction between Ag+ and Cl to form the insoluble salt AgCl. Cloudy precipitates usually make it difficult to determine the endpoint precisely. To compensate, precipitation titrations often have to be done as "back" titrations (see below).
  • Amperometry: Measures the current produced by the titration reaction as a result of the oxidation or reduction of the analyte. The endpoint is detected as a change in the current; this method is most useful when the excess titrant can be reduced, as in the titration of halides with Ag+.

Endpoint and equivalence point[edit]


There is a slight difference between the endpoint and the equivalence point of the titration; this error is referred to as an indicator error, and it is indeterminate.[41]

Back titration[edit]


Graphical methods[edit]

The titration process creates solutions with compositions ranging from pure acid to pure base. Identifying the pH associated with any stage in the titration process is relatively simple for monoprotic acids and bases; the presence of more than one acid or base group complicates these computations. Graphical methods,[43] such as the equiligraph,[44] have long been used to account for the interaction of coupled equilibria; these graphical solution methods are simple to implement, however they are infrequently used.

Particular uses[edit]

A titration is demonstrated to secondary school students.

Acid–base titrations[edit]

Redox titrations[edit]

  • Winkler test for dissolved oxygen: Used to determine oxygen concentration in water. Oxygen in water samples is reduced using manganese(II) sulfate, which reacts with potassium iodide to produce iodine; the iodine is released in proportion to the oxygen in the sample, thus the oxygen concentration is determined with a redox titration of iodine with thiosulfate using a starch indicator.[47]


  • Karl Fischer titration: A potentiometric method to analyze trace amounts of water in a substance. A sample is dissolved in methanol, and titrated with Karl Fischer reagent; the reagent contains iodine, which reacts proportionally with water. Thus, the water content can be determined by monitoring the potential of excess iodine.[50]

See also[edit]


  1. ^ Whitney, W.D.; Smith, B.E. (1911). "Titrimetry". The Century Dictionary and Cyclopedia; the Century Co. p. 6504.
  3. ^ "Titrand". Science & Technology Dictionary. McGraw-Hill. Retrieved 30 September 2011.
  4. ^ Ortolang: "titre" (in French): "4. a) 1543 tiltre « proportion d'or ou d'argent dans les monnaies, dans les ouvrages d'or et d'argent » " (tiltre: proportion of gold or silver in monies, in works of gold or silver)
  5. ^ Etymology On Line: titrate
  6. ^ WordReference: titre and titer
  7. ^ Gay-Lussac (1828). "Essai des potasses du commerce" [Assays of commercial potash]. Annales de Chimie et de Physique. 2nd series (in French). 39: 337–368. In footnote (1) of p. 340, Gay-Lussac first uses titre as a verb: "Il leur serait plus facile de titrer l'acide sulfurique normal au moyen du carbonate de soude ou de potasse pur; … " ([In determining the concentration of sulfuric acid] it would be easier for them to titrate normal sulfuric acid by means of pure sodium or potassium carbonate; … )
  8. ^ Szabadváry, Ferenc (1993). History of Analytical Chemistry. Taylor & Francis. pp. 208–209. ISBN 2-88124-569-2.
  9. ^ Descroizilles (1795). "Description et usages du Berthollimêtre, …" [Description and uses of the Berthollimeter, …]. Journal des Arts et Manufactures (in French). 1: 256–276.
  10. ^ Wisniak, Jaime (2014). "François Antoine Henri Descroizilles". Revista CENIC Ciencias Químicas. 45 (1): 184–193.
  11. ^ Gay-Lussac (1824). "Instruction sur l'essai du chlorure de chaux" [Instructions on the assaying of chlorinated lime]. Annales de chimie et de physique. 2nd series (in French). 26: 162–175. On pp. 170–171, Gay-Lussac describes various figures that appear in a plate (illustration) that accompanies the article. From p. 170: " F, petite mesure ou pipette de 2 ½ centimètres cubes, … " ( F, small measure or "pipette" of 2 ½ cc., … ) From p. 171: " I, burette destinée à mesurer la teinture d'épreuve: … " ( I, "burette" intended to measure the test dye: … )
  12. ^ Henry, O. (1845). "Nouvelles expériences sur l'essai des potasses du commerce et appareil dit potassimètre pour l'effectuer" [New experiments on the assay of commercial potash and an apparatus called a "potassimeter" to perform it]. Journale de Pharmacie et de Chimie. 3rd series (in French). 7: 214–222. A sketch of Henry's burette appears on p. 218.
  13. ^ See:
  15. ^ Mohr, Karl Friedrich (1855). Lehrbuch der chemisch-analytischen Titrirmethode … , part 1 [Textbook of analytical chemistry titration methods …] (in German). Braunschweig, (Germany): Friederich Vieweg und Sohn. p. 2–20. Page 3 shows Mohr's burette; page 12 shows a burette with a glass stopcock (Glasshahn).
  18. ^ Verma, Dr. N.K.; S.K. Khanna; Dr B. Kapila. Comprehensive Chemistry XI. New Dehli: Laxmi Publications. pp. 642–645. ISBN 81-7008-596-9.
  21. ^ Reger, D.L.; S.R. Goode; D.W. Ball (2009). Chemistry: Principles and Practice (3 ed.). Cengage Learning. pp. 684–693. ISBN 0-534-42012-5.
  22. ^ Bewick, S.; J. Edge; T. Forsythe; R. Parsons (2009). CK12 Chemistry. CK-12 Foundation. pp. 794–797.
  23. ^ "pH measurements with indicators". Retrieved 29 September 2011.
  24. ^ "Titrating Soluble RM, R2NM and ROM Reagents" (PDF). Retrieved 2014-06-04.
  25. ^ "Methods for Standardizing Alkyllithium Reagents (literature through 2006)" (PDF). Retrieved 2014-06-04.
  26. ^ Quantitative Chemical Analysis, 7Ed. by Daniel C. Harris. Freeman and Company 2007.
  27. ^ Skoog, D.A.; West, D.M.; Holler, F.J. (2000). Analytical Chemistry: An Introduction, seventh edition. Emily Barrosse. pp. 265–305. ISBN 0-03-020293-0.
  28. ^ Henry, N.; M.M. Senozon (2001). The Henderson-Hasselbalch Equation: Its History and Limitations. Journal of Chermical Education. pp. 1499–1503.
  30. ^ Amerine, M.A.; M.A. Joslyn (1970). Table wines: the technology of their production. 2 (2 ed.). University of California Press. pp. 751–753. ISBN 0-520-01657-2.
  34. ^ DeMore, W.B.; M. Patapoff (September 1976). "Comparison of Ozone Determinations by Ultraviolet Photometry and Gas-Phase Titration". Environmental Science & Technology. 10 (9): 897–899. Bibcode:1976EnST...10..897D. doi:10.1021/es60120a012.
  36. ^ Somasundaran, P. (2006). "Calculation of Zeta-Potentials from Electrokinetic Data". Encyclopedia of Surface and Colloid Science (2 ed.). CRC Press. 2: 1097. ISBN 0-8493-9607-7.
  37. ^ Dukhin, A.S.; P.J. Goetz (2002). Ultrasound for Characterizing Colloids: Particle sizing, Zeta potential, Rheology. Studies in Interface Science. 15. Elsevier. pp. 256–263. ISBN 0-444-51164-4.
  42. ^ Kenkel, J. (2003). Analytical Chemistry for Technicians. 1 (3 ed.). CRC Press. pp. 108–109.
  43. ^ "The Equligraph: Revisiting an old tool". Retrieved 4 October 2015.
  44. ^ Freiser, H. (1963). Ionic Equilibria in Analytical Chemistry. Kreiger. ISBN 0-88275-955-8.
  46. ^ Remington: the science and practice of pharmacy. 1 (21 ed.). Lippincott Williams & Wilkins. 2005. p. 501. ISBN 0-7817-4673-6.
  48. ^ Biology. 3. London: Taylor & Francis. 1967. p. 52.

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