Titration and pH Measurement


Titration is the quantitative addition of a solution of known concentration to a solution of unknown concentration until the reaction between them is complete to determine the concentration of the second solution. An acid–base titration is the quantitative determination of the concentration of an acid or a base. Titration of an acid with a base requires that the pH, or relative concentrations of the two reactants, be monitored. pH can be assessed by litmus paper or by indicators, for example, phenolphthalein, but these methods lack precision. Typically, pH measurement in the laboratory is done by measuring the cell potential of that sample in reference to a standard hydrogen electrode. A plot of the pH of an acidic (or basic) solution as a function of the amount of added base (or acid) is a titration curve. From this, the endpoint or equivalent points can be determined.

Key Concepts:

  • Acid–base titrations determine the concentration of an acidic (or basic) solution (the analyte) by the delivery of a measured volume of base (or acid) of known concentration (the titrant).

  • The progress of an acid–base titration is often monitored by a titration curve.

Keywords: acid; base; electrode; neutralisation; pH; titration; equivalence point; logarithmic scale

Figure 1.

Comparisons of the pH measurement capability of litmus paper, the indicator phenolphthalein and a pH meter at different stages of a strong acid–strong base titration. (a) Just before the equivalence point, the litmus paper indicates that the solution in the beaker is acidic by its bright red colour. This conclusion is confirmed by the fact that the phenolphthalein added to the solution in the beaker is colourless. The pH meter indicates that the precise solution pH is 3.40, which is equal to a H3O+ concentration of 4.0×10–4 mol L−1. (b) At the equivalence point, the litmus paper fails to display any appreciable colour change, suggesting a neutral solution; however, quantitation of the H3O+ concentration is problematic. As the phenolphthalein has not yet reached its colour transition range, and is therefore still colourless, it indicates that the solution is either acidic or neutral. Only the pH meter is capable of indicating that the solution is neutral (pH 7.0). (c) Just beyond the equivalence point, the litmus paper indicates that the solution in the beaker is basic by its bright blue colour. The phenolphthalein verifies the pH increase with a pink colour change. In theory, this is the titration endpoint. In practice, however, it is very difficult for the experimenter not to overshoot this point and add too much base. Thus, depending on the experimenter's dexterity, the pink colour may or may not denote the true end point. Finally, as it has done at each of the three titration points, the pH meter again measures and indicates the precise solution pH of 10.60, which is equal to a H3O+ concentration of 2.5×10–11 mol L−1.

Figure 2.

Curves for the titration of 100 mL of 0.1 mol L−1 sodium hydroxide with a strong acid (100 mL of 0.1 mL−1 hydrochloric acid, blue curve) and with a weak acid (100 mL of 0.1 mol L−1 acetic acid, red curve). Note that the weak acid's pH changes least rapidly near the half‐neutralisation point. Here, the concentration of weak acid approximates the concentration of its conjugate base, and the ability of the solution to resist pH changes, known as buffering, is optimal.



Glauber (1658) Tractatus de natura salium. 2 vols.

Macca C and Solda L (2002a) pH‐static techniques in volumetric analysis I. A theoretical reappraisal of pH‐static acid‐base titrations. Electroanalysis 14(1): 57–61.

Macca C and Solda L (2002b) pH‐static techniques in volumetric analysis II. Experimental study of pH‐static acid‐base Titrations. Electroanalysis 14(1): 63–70.

Naumann R, Allexander‐Weber C, Eberhardt R, Giera J and Spitzer P (2002) Traceability of pH measurements by glass electrode cells: performance characteristic of pH electrodes by multi‐point calibration. Analytical and Bioanalytical Chemistry 374(5): 778–786.

Spitzer P and Werner B (2002) Improved reliability of pH measurements. Analytical and Bioanalytical Chemistry 374(5): 787–795.

Vonau W (2010) Electrochemical pH sensors for special applications. TM‐Technisches Messen 77(3): 162–172.

Further Reading

Albert A and Sargent EP (1962) Ionization Constants of Acids and Bases. New York: Wiley.

Bell RP (1973) The Proton in Chemistry, 2nd edn. Ithaca, NY: Cornell University Press.

Galster H (1991) pH Measurement: Fundamentals, Methods, Applications, Instrumentation. Weinheim: VCH.

Kotyk A and Slavik J (1989) Intracellular pH and Its Measurement. Boca Raton, FL: CRC Press.

Lunelli B and Scagnolari F (2009) pH Basics. Journal of Chemical Education 86(2): 246–250.

Miao YQ, Chen JR and Fang KM (2005) New technology for the detection of pH. Journal of Biochemical and Biophysical Methods 63(1): 1–9. DOI: 10.1016/j.jbbm.2005.02.001.

Perrin DD and Dempsey B (1974) Buffers for pH and Metal Ion Control. London: Science Paperbacks, Chapman and Hall.

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Jennings, Patricia A, Mullen, Christine A, and Roy, Melinda(Dec 2010) Titration and pH Measurement. In: eLS. John Wiley & Sons Ltd, Chichester. http://www.els.net [doi: 10.1002/9780470015902.a0002700.pub2]