An Introductory Course of Quantitative Chemical Analysis eBook

This eBook from the Gutenberg Project consists of approximately 188 pages of information about An Introductory Course of Quantitative Chemical Analysis.
of carbonic acid from the carbonate is a disadvantage with many indicators; barium hydroxide solutions may be prepared which are entirely free from carbon dioxide, and such solutions immediately show by precipitation any contamination from absorption, but the hydroxide is not freely soluble in water; ammonia does not absorb carbon dioxide as readily as the caustic alkalies, but its solutions cannot be boiled nor can they be used with all indicators.  The choice of a solution must depend upon the nature of the work in hand.

A !normal acid solution! should contain in one liter that quantity of the reagent which represents 1 gram of hydrogen replaceable by a base.  For example, the normal solution of hydrochloric acid (HCl) should contain 36.46 grams of gaseous hydrogen chloride, since that amount furnishes the requisite 1 gram of replaceable hydrogen.  On the other hand, the normal solution of sulphuric acid (H_{2}so_{4}) should contain only 49.03 grams, i.e., one half of its molecular weight in grams.

A !normal alkali solution! should contain sufficient alkali in a liter to replace 1 gram of hydrogen in an acid.  This quantity is represented by the molecular weight in grams (40.01) of sodium hydroxide (NaOH), while a sodium carbonate solution (Na_{2}Co_{3}) should contain but one half the molecular weight in grams (i.e., 53.0 grams) in a liter of normal solution.

Half-normal or tenth-normal solutions are employed in most analyses (except in the case of the less soluble barium hydroxide).  Solutions of the latter strength yield more accurate results when small percentages of acid or alkali are to be determined.

INDICATORS

It has already been pointed out that the purpose of an indicator is to mark (usually by a change of color) the point at which just enough of the titrating solution has been added to complete the chemical change which it is intended to bring about.  In the neutralization processes which are employed in the measurement of alkalies (!alkalimetry!) or acids (!acidimetry!) the end-point of the reaction should, in principle, be that of complete neutrality.  Expressed in terms of ionic reactions, it should be the point at which the H^{+} ions from an acid[Note 1] unite with a corresponding number of Oh^{-} ions from a base to form water molecules, as in the equation

H^{+}, Cl^{-} + Na^{+}, Oh^{-} —­> Na^{+}, Cl^{-} + (H_{2}O).

It is not usually possible to realize this condition of exact neutrality, but it is possible to approach it with sufficient exactness for analytical purposes, since substances are known which, in solution, undergo a sharp change of color as soon as even a minute excess of H^{+} or Oh^{-} ions are present.  Some, as will be seen, react sharply in the presence of H^{+} ions, and others with Oh^{-} ions.  These substances employed as indicators are usually organic compounds of complex structure and are closely allied to the dyestuffs in character.

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An Introductory Course of Quantitative Chemical Analysis from Project Gutenberg. Public domain.
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