Litmus Test
The litmus test is a simple method used to determine the acidity or alkalinity of an aqueous solution based on the color change of litmus, a dye derived from lichens. Litmus is sensitive to pH levels, turning red in acidic environments and blue in alkaline conditions, with the intensity of the color reflecting the concentration of hydronium or hydroxide ions present. This color change occurs due to reversible reactions involving litmus compounds that absorb different light wavelengths based on the solution's pH. In neutral solutions, like pure water, litmus remains blue, indicating a balanced concentration of hydronium and hydroxide ions.
While the litmus test serves as a quick indicator of pH, more precise measurements are typically obtained through electronic devices that do not alter the solution's properties. The concept of pH is crucial in chemistry, with a scale ranging from 0 to 14, where a pH of less than 7 indicates acidity, and above 7 indicates alkalinity. Understanding the behavior of litmus and pH can be essential in various scientific fields, including chemistry, biology, and environmental science, where the acidity or alkalinity of solutions plays a significant role in many processes.
Litmus Test
FIELDS OF STUDY: Analytical Chemistry; Inorganic Chemistry; Organic Chemistry
ABSTRACT
Litmus is a complex mixture of compounds isolated from various lichens. The material is sensitive to changes in pH, turning red in acidic solutions and blue in alkaline solutions. The intensity and shade of the color correlate to the pH of the solution.
The Theory behind the Litmus Test
Water-soluble materials, particularly acids and bases, alter the natural amounts of hydrogen ions present in a solution. Pure water undergoes an equilibrium dissociation reaction called autoprotolysis, in which a molecule of water splits into a hydrogen ion (H+) and a hydroxide ion (OH−), according to the following equation:
2H2O ⇌ H+∙H2O + OH−
Since equal amounts of positively charged hydronium ions (H3O+) and negatively charged hydroxide ions are produced, water is neutral, having neither a net negative nor a net positive charge. Thus, it is neither acidic nor alkaline. When a compound is dissolved in water, it may disrupt this equilibrium so that the hydronium ions and the hydroxide ions are no longer in balance. This does not mean that there will be a net negative or positive electrical charge, however, since the material that was dissolved will always have the appropriate number of opposite charges, so the solution will always be electrically neutral. The important feature is that the hydroxide ions and the hydronium ions specifically will not be in balance. For example, dissolving sodium hydroxide (NaOH) in water adds extra hydroxide ions to the solution, but their negative charge is balanced by the positive charges carried by the sodium ions. However, because there are more hydroxide ions than hydronium ions, the autoprotolysis equilibrium cannot be maintained, so a solution of sodium hydroxide in water is alkaline rather than neutral. In the same sense, dissolving an acidic compound, such as sulfuric acid (H2SO4), in water introduces an excess of hydronium ions, creating a solution that is acidic in nature.
The Nature of Litmus
Many aqueous solutions of acid are colorless, exactly like plain, pure water. A method of indicating which solution is acidic, which is alkaline, and which is neutral is required to tell them apart. This is the purpose of the litmus test.
The material known as "litmus" is a combination of several similar compounds isolated from a number of species of lichens. The compounds are sensitive to the acidic or alkaline character of aqueous solutions. They are also water soluble and so can be added to or mixed directly with other solutions in water. Litmus reacts with hydronium and hydroxide ions differently, and as the molecular structures of the litmus compounds change, the materials change color accordingly. In acidic solutions, litmus becomes red, while in alkaline solutions, it becomes blue. The intensity of the color relates directly to the concentration of hydronium and hydroxide ions in the solution. In this way, litmus can be used as a pH indicator.
Litmus is now used only as a quick means of approximating the pH of a solution. Most pH measurements are obtained using electronic instruments that are more precise and do not add contaminants that can alter the pH of the solution.
Litmus and the pH Scale
It is generally inconvenient to use numerical values for hydronium and hydroxide ion concentrations. In neutral water, the concentration of each is just 10−7 molar (M), or moles per liter. Such a numerical value is necessary in precise calculations, but for general purposes, it is much more convenient to describe acidic and alkaline solutions in terms of their respective pH values. The pH of a solution is defined as
pH = −log[H3O+]
where [H3O+] is the molar concentration of hydronium ions in the solution. Accordingly, the pH of pure water is 7. The pOH, which is defined as
pOH = −log[OH−]
is therefore also 7, and the equilibrium constant of water autoprotolysis is 10−7 × 10−7, or 10−14, which has a logarithmic value of 14. Because the constant is an equilibrium value, the pH and the pOH of a solution must always add up to 14, which is an extremely useful rule when carrying out numerical calculations involving pH. As the concentration of hydronium ions increases from the neutral value of 10−7 M, the pH must decrease in value. Thus, acidic solutions have a pH value of less than 7, and alkaline solutions have a pH greater than 7.
The Colors of Litmus
The interaction between litmus and hydronium ions is also an equilibrium process. In solid powder form, litmus appears blue. The color-producing functional groups, or chromophores, of the litmus compounds absorb light such that only blue light is reflected. When litmus becomes protonated—gains a proton—in an acidic solution, as in
litmus + H+ ⇌ litmusH+
the bonds of the chromophores change so that light in the blue region of the spectrum is absorbed, allowing light in the red region to pass through. The more hydronium ions are present, the more the equilibrium between litmus and protonated litmus shifts toward the protonated form, and so the intensity of the red color increases with increasing hydronium ion concentration. But when litmus interacts with hydroxide, it gives up a proton, acting as an acid:
litmus + OH− ⇌ litmus− + H2O
When this happens, the chromophores absorb light from the red region of the spectrum, permitting light from the blue region of the spectrum to pass through. Again, the amount of hydroxide ions present shifts the equilibrium of the litmus-hydroxide reaction, causing the intensity of the blue color to increase with increasing hydroxide ion concentration.
PRINCIPAL TERMS
- acid: a compound that can relinquish one or more hydrogen ions (Brønsted-Lowry acid-base theory) or that possesses vacant atomic orbitals to interact with electron-rich materials (Lewis acid-base theory).
- alkaline: describes a material that tends to increase the concentration of hydroxide ions in an aqueous solution, as well as conditions produced by the presence of bases.
- base: a compound that can relinquish one or more hydroxide ions (Brønsted-Lowry acid-base theory) or that possesses lone pairs of electrons that can interact with electron-poor materials (Lewis acid-base theory).
- neutral: describes a chemical solution that has a pH of 7 and thus is neither acidic nor basic.
- pH indicator: a compound that changes color according to the pH of a solution, or a device that measures the pH of a solution electronically.
- water soluble: describes a compound that can be dissolved by water.
Bibliography
Jones, Mark M., et al. Chemistry and Society. 5th ed. Philadelphia: Saunders Coll., 1987. Print.
MacKay, K. M., R. Ann MacKay, and W. Henderson. Introduction to Modern Inorganic Chemistry. 6th ed. Cheltenham: Nelson, 2002. Print.
Myers, Richard. The Basics of Chemistry. Westport: Greenwood, 2003. Print.
Skoog, Douglas A., Donald M. West, and F. James Holler. Fundamentals of Analytical Chemistry. 9th ed. Boston: Brooks, 2014. Print.
Wink, Donald J., Sharon Fetzer-Gislason, and Sheila McNicholas. The Practice of Chemistry. New York: Freeman, 2004. Print.