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The measurement of pH can become difficult at extremely acidic or alkaline conditions, such as below pH 2.5 (ca. 0.003 mol/dm 3 acid) or above pH 10.5 (above ca. 0.0003 mol/dm 3 alkaline). This is due to the breakdown of the Nernst equation in such conditions when using a glass electrode. There are several factors that contribute to this problem.
The Henderson–Hasselbalch equation can be used to estimate the pH of a buffer solution by approximating the actual concentration ratio as the ratio of the analytical concentrations of the acid and of a salt, MA. The equation can also be applied to bases by specifying the protonated form of the base as the acid. For example, with an amine,
Therefore, the buffer regions will be centered at about pH 1.3 and pH 4.3. The buffer regions carry the information necessary to get the pK a values as the concentrations of acid and conjugate base change along a buffer region. Between the two buffer regions there is an end-point, or equivalence point, at about pH 3.
A pH indicator is a halochromic chemical compound added in small amounts to a solution so the pH (acidity or basicity) of the solution can be determined visually or spectroscopically by changes in absorption and/or emission properties. [1] Hence, a pH indicator is a chemical detector for hydronium ions (H 3 O +) or hydrogen ions (H +) in the ...
Acid strength is the tendency of an acid, symbolised by the chemical formula , to dissociate into a proton, , and an anion, . The dissociation or ionization of a strong acid in solution is effectively complete, except in its most concentrated solutions. Examples of strong acids are hydrochloric acid , perchloric acid , nitric acid and sulfuric ...
It is the quantity of base (usually potassium hydroxide (KOH)), expressed as milligrams of KOH required to neutralize the acidic constituents in 1 gram of a sample. [1][2][3][4] The acid value measures the acidity of water-insoluble substances like oils, fats, waxes and resins, which do not have a pH value. The acid number is a measure of the ...
The pKa of KHP is 5.4, so its pH buffering range would be 4.4 to 6.4; however, due to the presence of the second acidic group that bears the potassium ion, the first pKa also contributes to the buffering range well below pH 4.0, which is why KHP is a good choice for use as a reference standard for pH 4.00. [8] [9]
An example of an alkalimetric titration involving a strong acid is as follows: H 2 SO 4 + 2 OH − → SO 4 2-+ 2 H 2 O. In this case, the strong acid (H 2 SO 4) is neutralized by the base until all of the acid has reacted. This allows the viewer to calculate the concentration of the acid from the volume of the standard base that is used.